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P . BONET-MAURY 79 MECHANISM OF DECOMPOSITION OF WATER BY IONIZING RADIATIONS BY AUGUSTINE 0. ALLEN Chemistry Department, Brookhaven National Laboratory, Upton, Long Island, New York, U S A . Received 14th January, 1952 Irradiated water undergoes a decomposition to molecular H2 and H202, simultaneously with the decomposition to free radicals, H and OH. With gamma rays or hard X-rays, the yield of the niolecular decomposition is about 0.6 Hz molecules formed per 100 eV absorbed, and the yield of free radicals appears to be about 3-5 radical pairs per 100 eV. As the ionization density of the radiation is increased, the molecular yield increases and the free radical yield falls. The decomposition of pure water reverses itself because the free radicals initiate a back reaction between the decomposition products.The rate of the back reaction increases with increasing concentration of dissolved hydrogen but de- creases with increasing concentration of dissolved oxygen or hydrogen peroxide. This80 DECOMPOSITION OF WATER unusual type of kinetics leads to some peculiar phenomena in water radiolysis. The molecular decomposition is ascribed to reactions occurring in the very small regions of high energy density along the charged particle track (hot spots) which correspond to the " ion clusters " formed by fast particles in a gas. Determination of what occurs when high energy radiation acts on water or simple aqueous solutions constitutes a difficult experimental problem. The first requisite for understanding a reaction is establishment of a complete material balance.Only by a knowledge of amounts of all products formed and all sub- stances disappearing can one have reasonable confidence of knowing what is going on. Since hydrogen gas is generally an important reaction product, this demands setting up a refined system for determining small concentrations of gas dissolved in small samples of water, as well as sensitive and precise methods for determination of other dissolved materials. A second necessity is care in purifying the material studied and in establishing reproducibility of results, since the behaviour of most solutions seems to be greatly affected by trace contamin- ations. Another highly desirable condition is that radiation sources should be available which allow reactions to be studied over a wide range of radiation in- tensity as well as radiation quality.During the last five years, researches have been carried out at several laboratories in the United States and Canada which have attempted to meet these difficult requirements. Results of these studies are now beginning to appear in the scientific journals. These researches have greatly enhanced our previous knowledge of radiation decomposition of water, and this paper will attempt to summarize the present status of the subject. DosImTRY.-Radiation reactions are discussed in terms of the yield of molecules reacting per unit energy input. In practice, laboratories working with X-rays or gamma rays generally obtain their yields by comparison with the yield of oxidation of ferrous sulphate in 0-8 N sulphuric acid solution.Unfortunately disagreement exists as to the absolute value of the ferrous sulphate oxidation yield. By comparison with ionization in air, Fricke and co-workers obtained a value of 18.2 molecules oxidized per 100 eV ; 1 Miller later obtained a value of 20.6.2 Calorimetric determination of the heat produced when solutions are irradiated would appear to provide a more direct method. Hochanadel has carried out such a determination with a cobalt gamma ray source at Oak Ridge and finds for the ferrous sulphate solution a yield of only 15.5 molecules oxidized per 100 eV.3 The determinations appear equally sound and it remains for further work to effect a choice between the values. from recent work is that in addition to the free radicals formed from water, a direct decomposition of water occurs to yield gaseous hydrogen independent of what material may be dissolved in the water.I had postulated the molecular reaction in a previous paper,4 largely on the basis of preliminary results obtained by J. A. Ghormley and myself on decomposition of water by X-rays and cathode rays. The molecular reaction has since been verified in numerous systems. Solutions of HCI, CuSO4, KBr and others at various concentrations, placed in the Oak Ridge nuclear reactor, were found to give the same initial yield of hydrogen gas within experimental error.5 Subsequent more accurate work by Dr. Everett Johnson and myself at Brookhaven,6 using hard X-rays from a 2 MeV electrostatic generator with gold target, has shown that a constant yield of hydrogen is obtained from irradiation of a variety of solutions, including KBr, KI, HBr, H202, oxygen-saturated ferrous sulphate, air-saturated water and ceric sulphate solutions.The only difference found was that from solutions containing 0.8 N H2SO4 the yield of hydrogen was consistently 18 % lower than from dilute solutions in pure water. The results were most striking with H202 solutions. Here the H202 rapidly decomposes to oxygen. Throughout the entire reaction, however, the hydrogen still keeps being generated at a constant rate (provided the initial con- centration of peroxide is high enough) which is exactly equal to the rate at which hydrogen is generated in KBr and other unrelated solutions. Changes in the details of preparing the solutions and intentional addition of foreign substances, which had a large effect on the peroxide decomposition yield, did not affect the hydrogen yield.7 THE MOLECULAR YIELD OF WATER DECOMPOSITION.-The main fact which emergesAUGUSTINE 0.ALLEN 81 The hydrogen is accompanied by an equivalent quantity of oxidizing material which appears to consist mostly, if not entirely, of H202. In acid bromide solutions the actual peroxide found in the solution is equivalent to the hydrogen formed. In neutral bromide or iodide solutions the oxidizing agent appears mostly as oxygen rather than peroxide,39 8 but this may be accounted for by the rapid decomposition of peroxide to oxygen under these conditions.7 The production of a constant yield of hydrogen gas was shown by Fricke and Hart for solutions of a number of simple reducing agents, viz.iodide, bromide, nitrite, arsenite, selenite and ferrocyanide.8 In all these cases the yield of hydrogen was the same, and was independent of concentration of dissolved material and of pH. With the above reducing agents other than bromide and iodide the oxidized form of the material appeared in equivalent amount to the hydrogen formed. In these cases the HzOz apparently oxidized the reducing agent quantitatively, this action being perhaps catalyzed by the free radicals present. As the concentration of dissolved materials is reduced to a low level, the hydrogen still initially appears with its characteristic rate but the rate soon slackens and the hydrogen concentration levels off at some definite value.59 6 When no solute is added, the hydrogen concentration produced by X-rays levels of€ at a very low value, generally close to the limits of detection, which is very sensitive to the presence of trace impurities and is practically impossible to reproduce.Evidently water decomposes directly to molecular hydrogen as well as forming free radicals that may react with solutes. In the absence of a sufficiently high concentration of solutes, the hydrogen reacts with radicals and is converted back to water ; but almost any solute capable of oxidation or reduction will destroy the radicals and allow the molecular hydrogen to appear. In a redox system, the radicals may produce equivalent amounts of reduction and oxidation and hence give no net effect, so that the sole overall reaction occurring is the production of molecular hydrogen and an equivalent amount of H202 or reaction products of H202.In HzO2 solutions, the radicals do produce a net reaction, decomposition of the peroxide to oxygen, but as long as peroxide and oxygen are present in sufficient amounts they use up all the radicals and the hydrogen keeps coming out at its characteristic rate. In ceric sulphate solution the overall yield of re- duction of cerium is greater than one would expect from the amount of molecular H202 formed along with the hydrogen gas. The interplay of radicals here is such as to result in some net reduction. Deaerated solutions of ferrous sulphate give an unusually high oxidation yield, the interplay of radicals resulting in some net oxidation.The mechanisms of these reactions are not yet entirely understood. The molecular hydrogen yield was demonstrated by Hart 9 for solutions of formic acid and oxygen. Deaerated formic acid solutions give a high yield of hydrogen, which arises mostly from the acid, as shown by the presence of deuterium in the hydrogen pro- duced by irradiation of solutions of DCOOH in ordinary water. When oxygen is added, the acid is oxidized instead of decomposing and the hydrogen yield falls to the value characteristic of the molecular decomposition of water. The hydrogen from DCOOH + oxygen solutions contains practically no deuterium. The molecular yield of hydrogen from hard X-rays or gamma rays, expressed as a fraction of the yield of ferrous sulphate oxidation, is 0.030 according to Hochanadel,3 0.036 according to Fricke and Hart,8 0.022 according to Hart9 and 0-039 according to Johnson and Allen in pure water, 0.031 in 0-8 N H2S04.6 The results of various laboratories appear to agree within experimental error, except perhaps for Hart’s low value.To convert these ratios into yields of hydrogen molecules formed from water per 1OOV energy input, the above fractions are to be multiplied by either 20.6 or 15.5, depending upon the value accepted for the absolute yield of the ferrous sulphate actinometer. The hydrogen yield thus lies between 0.45 and 0-7 molecules per 100 eV. of hydrogen from decomposition of water by natural alpha rays has been studied by many investigators and is equal to about 1.9 - 0.1 molecules of hydrogen per 100 eV.103 1 1 , 1 2 This number presumably represents the molecular yield of hydrogen from water by this type of radiation since the initial yield does not drop as the hydrogen accumulates, and solutions of HI were found to give nearly the same value as pure water.]] For radiations of ionization density intermediate between natural alpha rays and hard X-rays the avail- able information is only qualitative but strongly suggests that the molecular yield increases continuously with the ionization density.EFFECT OF IONIZATION DENSITY OF RADIATION ON THE MOLECULAR YIELD.-The yield82 DECOMPOSITION OF WATER Results with cyclotron deuterons, protons and alphas have been published by Toulis.13 He found that as the energy of the protons reaching water was increased from 2 to 12 MeV, the peroxide yield per proton instead of increasing linearly with the proton energy, showed a curve which bent over, reached a maximum at 7 MeV and decreased at higher proton energies.Thus the part of the proton beam lying above 7 MeV actually caused more destruction of H202 than formation, while the peroxide formation occurred chiefly in the lower energy parts of the proton tracks. Similar results were obtained with cyclo- tron deuterons and alphas over a comparable range of ionization densities. The resuits of such an experiment must depend on how the HzO2 is mixed into the body of solution, since the local concentrations of peroxide may vary with depth in the water as the mean ionization density of the rays varies. We cannot agree with Toulis in the detailed interpretation that he offers of his results ; but the results themselves are extremely suggestive.Observations have been made of the steady-state level in the decomposition of water with a variety of types of radiation which tend toward the same conclusion. Such observations are not quantitative since the steady-state level is greatly affected by trace impurities and is generally not well reproducible. However, in the Oak Ridge reactor,5 the effect of changing radiation quality on the steady-state level of decomposition of water was demonstrated in a vessel which was kept sealed, thus keeping the chemical con- ditions as constant as possible. When the mean ionization density of the reactor radi- ation was decreased by surrounding the vessel with paraffin, the pressure of gaseous decomposition products of water within the bulb showed a precipitate drop ; when the ionization density was increased by surrounding the vessel with lead, the pressure rose.The change of steady-state decomposition pressure with radiation quality is a good indication that the molecular yield increases with increasing radiation density. In the decomposition of water by the soft beta rays from tritium (average energy 5700 eV) steady- state decomposition levels were obtained which were much higher than those expected from similar experiments with hard X-rays.14 Hart has irradiated solutions of formic acid 9 and of ferrous sulphate,ls in the presence and absence of oxygen, with tritium beta rays as well as gamma rays. An exact com- parison of the molecular yields from the two types of radiation is difficult to make, because of uncertainties in the absolute dosimetry and in the reaction mechanisms, but a higher value for the tritium rays is indicated, and the difference probably amounts to about 20 %.The indications are then that the molecular yield increases with increasing ionization density but more quantitative information on this point is badly needed. decomposition is always the resultant of the rate of decomposition of water (re- ferred to as forward reaction) and the rate of re-formation of water from the decomposition products (back reaction). The forward reaction ordinarily arises entirely as a result of the molecular decomposition of water ; the free radicals formed by water decomposition usually disappear by reaction with the molecular reaction products and are responsible for the back reaction.The kinetics of water decomposition are to be explained by a study of the back reaction, and the most effective way of studying radiation decomposition in water is to irradiate solutions containing H2, 0 2 and H202 at various proportions and concentration levels. Extensive experiments with the mixed fast neutron and gamma radiation in the Oak Ridge reactor5 and later experiments with gamma radiation alone3 have shown that the back reaction is accelerated by increasing the hydrogen concentration, but is decelerated by increasing the concentration of either 0 2 or W202. This leads to some curious phenomena in the decomposition of water.If the water is irradiated in a closed vessel full of pure water, a steady state is soon reached at a low level of product concentrations. If, however, a gas space exists over the water the hydrogen gas produced escapes into the gas phase leaving an excess of H 2 0 2 dissolved in the liquid phase. The excess H 2 0 2 inhibits back reaction and the decomposition proceeds to a much greater extent than in a full vessel. If hydrogen is added to water prior to irradiation the decomposition is completely repressed, but if oxygen or H 2 0 2 is added initially the decomposition proceeds to high levels, with much hydrogen being produced on long irradiations. REACTION KINETICS IN WATER DECOMPOSITION.-The observed yield Of WaterAUGUSTINE 0 . ALLEN 83 A trace of organic impurity in the water will decompose initially with the forma- tion of hydrogen which acts to repress the water decomposition so that formation of oxygen or peroxide is not observed.Traces of certain inorganic impurities, however, may inhibit the back reaction almost entirely so that large amounts of decomposition are seen. These facts explain apparent anomalies in some of the older work on the effect of X-rays on pure water. Risse 16 and Fricke and co- workers 17 irradiated water in filled vessels and observed very little decomposition. Guenther 18 irradiated water in the presence of a very large available gas volume and observed continuing production of hydrogen. These apparently contra- dictory results are exactly what is expected on the present picture.The kinetics of the back reaction in the presence of excess peroxide or oxygen are unfortunately irreproducible.5 Irreproducibility is also characteristic of the kinetics of decomposition of peroxide to oxygen in irradiated peroxide solutions. In systems containing dissolved hydrogen and oxygen or peroxide, with hydrogen not present in excess, conversion of peroxide to oxygen ‘or vice versa always occurs and the kinetics are erratic, However, a solution of H202 and hydrogen with the hydrogen in excess gener- ates little or no oxygen on irradiation, the sole reaction being the disappearance of equimolecular quantities of H2 and H202 with formation of water. Under these conditions reproducible reaction kinetics were found both with nuclear reactor radiation 5 and with gamma radiation.3 The yield of disappearance of hydrogen or peroxide, after correction for the amounts of these materials pro- duced by molecular decomposition of water, is the3 strictly proportional to the hydrogen concentration, inversely proportional to the peroxide concentration and independent of the radiation intensity : where Gr; is the yield of molecular decomposition of water with the radiation used, as evaluated by rate of hydrogen evolution from other solutions such as bromide or HCl.The molecular decomposition equation is 2H20 -- W? 4- H202. Eqn. (1) suggests that a chain reaction is occurring leading to reaction of hydrogen, with peroxide acting to break the chains as well as to carry them on, so that peroxide acts as an inhibitor. The most likely scheme for the chain reaction is (F) H20 r= H -1 OH H + H2Oz -- OH f H20 OH + H2 = H + 1320 The chain-breaking reaction involving peroxide is probably OH + M201= HO2 + H20 which is universally postulated to account for the photochemical or radiation chemical decomposition of peroxide to oxygen.In the presence of excess hydrogen, all the HO;! radicals formed must be reduced back to €3202 and none oxidized to oxygen, since no oxygen is formed in the reaction. This probably occurs because with excess H2 the concentration of hydrogen atoms is kept high by means of reaction (3), so that all HO2 radicals disappear by the reaction (4) Reaction ( 5 ) is not a rate-determining step, since under the conditions postuIated every H02 formed disappears by (5). Furthermore, as long as a reasonable con- centration of HzOz is present all hydrogen atoms not consumed in (5) react by84 DECOMPOSITION OF WATER (2), so that reactions of hydrogen atoms are not rate-determining either.The only rate-determining steps then are reaction (R), which gives the rate at which chains are started, and reactions (3) and (4), the ratio of which gives the proba- bility that the chain will either be carried on by reaction of OH with H2 or ter- minated as a result of reaction of OH with H202. The observed reaction kinetics (eqn. (1)) follow immediately from these considerations, since the rate of any chain reaction equals the rate of initiation times the ratio of the rate of continu- ation to the rate of termination. The constant K is equal to the product G~k3/k4, where GR is the yield of reaction (R).In the papers referred t0,395 the rate expression is derived by the standard steady-state method of homogeneous reaction kinetics, assuming that the bulk concentrations of the different free radicals may be taken as uniform. Objection has been taken to this method of treating radiation chemical reactions, since the concentration of radicals is not in fact uniform but varies as the radicals diffuse away from the regions where they are initially formed. In the above case such considerations make no difference, since the rate-determining step is merely the competition of hydrogen and peroxide molecules for the OH radicals, which depends only on the concentrations of these molecules and not on that of the radicals.The effects of " track overlap " on radiation chemical kinetics, on which considerable speculation has appeared in print, would appear to be important only in cases where the rate of reaction between radicals actually determines the rate of the overall observed reaction. This may be true for radiation-induced polymerization of such solutes as acrylonitrile.19 The course of most radiation- chemical reactions, however, seems to be determined more by competition of dissolved substances for reaction with radicals rather than by reaction between radicals, and in such cases the overlapping of tracks has no special significance. THE FREE RADICAL YIELD IN WATER DECOMPOSITION.-The chemical effect Of any given radiation on water can be characterized by two numbers, GF and GR, one giving the yield of molecular products, the other the yield of free radicals.GF can be readily obtained by measuring the yield of hydrogen produced on irradiation of a variety of solutions. The value of GR is much more difficult to estimate. The difficulty is that whatever action is produced by one radical of the H-OH pair can, in general, be re- versed by the other radical. This is most obvious in reversible redox systems, but even with organic compounds such possibilities exist. Thus with any organic compound we may have the sequence RH + OH = H20 + R; R + H = RH. The yield of conversion of dissolved materials will, in general, give only a lower limit to CR. An attempted way out of the difficulty has been to study a solution containing two solutes, one of which was supposed to react readily with hydrogen atoms and remove them from the system, the other to react with the OH radicals.Then the total chemical change should be a measure of the total free radical yield, providing no subsequent reactions occur to reverse the changes. Possiblity of such reversal still appears, however, to provide an essential difficulty. Attempts to evaluate GR in this way have been made using results obtained on solu- tions of the mixtures formic acid + oxygen, hydrogen + oxygen and ferrous sulphate + oxygen. The formic acid t- oxygen mixture on irradiation yields the reaction 9 HCOOH + 0 2 = H202 -I- C02. The assumed mechanism was H + 0 2 = HO2 OH t HCOOH = HCOO + H20 H02 + HCOO == H202 + C02. According to this mechanism the yield of conversion of formic acid and oxygen should be equal to GR.The difficulty is that if HO2 radicals after being formed react with them- selves or with OH or H, the value of GR will have been underestimated. The radical yield estimated by Hart from this reaction was 2-78 radical pairs per 100 eV for cobalt gamma rays if the dosimetric standard is 15.5, 3.6 if the standard is 20-6, and we may assume that at least this many free radicals are produced.AUGUSTINE 0. ALLEN 85 The formation of H202 in solutions of hydrogen and oxygen was treated by Hochanadel in a similar way.3 The assumed reactions were H + 0 2 = H02 OH 3- H2 = H20 + H 2H02 = H202 + 0 2 and the peroxide yield was taken as a direct measure of the number of free radicals produced. The yield obtained was very close to that found by Hart with formic acid but here again the same difficulty occurs with the reaction of HO2.For the ferrous sulphate + oxygen system, Krenz and Dewhurst 20 and also Hart 15 have used the mechanism H + 0 2 = H02 OH $- Fe2+ = Fe3+ + OH- H+ + HO2 + Fe2+ = Fe3+ + H202 H202 + Fez+ = Fe3+ + OH- + OH. This gives a yield of 4 or 5 radical pairs per 100eV (depending upon the dosimetric standard), which is significantly higher than the free radical yield values found for the formic acid and hydrogen cases. Evidently the mechanisms assumed are not entirely correct in at least one of the above cases, and further work is being carried out on all these reactions. In any case it would appear that the results on any one such system can, in general, give only a lower limit to the free radical yield, and authentic values will be on hand only when the numbers deduced from several different independent reactions are found to show precise agreement.PHYSICAL PICTURE OF THE ACTION OF RADIATIONS ON WATER.-Much has been written about the processes occurring when radiations traverse water, mostly from the viewpoint that water should behave much like a highly compressed perfect gas. Such an assumption is so far from reality that we cannot expect it to furnish a very good picture of the real process, much less to provide quantitative predictions of experiments. However, some sort of physical model seems required by our minds as a framework for our thinking about the observed chemical phenomena, and models of this sort do no harm if not taken too seriously.The rate of energy loss by charged particles passing through matter is nearly independent of the state of chemical combination of the atoms present, but the nature of the excited states formed certainly depends strongly on the state of combination and state of aggregation of the material. We may therefore assume that the distribution of energy loss is the same in liquid water as in a vapour. But the distinction between ionization and excitation, so prominent in the dilute gas, may be largely lost in the liquid state. The upper excitation bands should be broadened in a liquid and merge with the ionized state, while on the other hand a slow electron ejected from a molecule in a liquid by ionization may well recoil, return and recombine with the positive ion thus giving in effect an excited molecule from a process which would be an ordinary ionization in the vapour state.Both ionization and excitation processes, however, result in the dissociation of water molecules to form free radicals, probably mostly H and OH. The best procedure in drawing our picture is therefore to ignore as far as possible the classification of processes in terms of excitation and ionization and to consider rather the distribution of free radicals in terms of distribution of energy loss within the liquid. The distribution of energy loss itself may legitimately be derived from experiments and theories relating to the gas phase. The loss of energy by fast electrons passing through matter occurs not uniformly but in bursts of varying size, leading in gases to the formation of ions in clusters along the track.The size of the clusters can be determined roughly from well- resolved cloud chamber pictures, and counts of the distribution of ions in the clusters have been published by Wilson21 and by Beekman.22 They find that the relative number of clusters of any given size decreases approximately exponen- tially with size of the cluster. The mean size of the clusters is about 2.5 to 386 DhCOMPOSITlON OF WATER ion pairs.23 Thus the “ mean cluster ” corresponds to an energy loss of about 100 eV. Considering a fast electron, say 100 keV, passing through liquid water, we have the picture of some isolated occurrences involving only enough energy loss to dissociate one water molecule into radicals, while other occurrences will involve dissociation of several water molecules located very close together.Assuming that a total of six water molecules are dissociated per 100 eV, a typical “ radical cluster ” will consist of a group of about six H and six OH radicals. Since only a third of the total energy is required to dissociate six water molecules, much energy will appear immediately as heat and the region is perhaps best denoted by the term “ hot spot ”. A typical hot spot arises when an electron of 100 eV energy is ejected from a water molecule. According to Lea’s tables,24 such an electron should have a total range in water of about 30A. This electron will suffer so many collisions that its path will resemble a random difflision path and its course will lie in a region having a diameter equal to about one-third of its total range, or lOA.The six OH radicals formed by it will lie closely along its path, but some question exists as to the location of the six H atoms. Some of these are formed by ejection of an electron from a water molecule and subsequent capture of this electron at a distance. Lea assumed that a free electron in water would diffuse for some time at thermal velocities before being captured and that the mean separation of H and OH was about 150 A.24 There is some evidence that this distance is too large and that the electrons are rapidly captured by liquid water. Bradbury and Tatel25 found that water vapour has a high cross-section for capture of low- energy electrons when the pressure is near the saturation point, but capture ceases when the pressure is low.If water vapour below saturation captures electrons by virtue of incipient molecular association, the liquid itself should capture them immediately. The hydrogen atoms may be no farther apart than the OH radicals. The most that can be said about size of the hot spots is that their radii should lie between 7 A and 150 A, probably closer to the smaller value. Considerable immediate recombination of radicals must occur in such a small region, and when like radicals combine we get the molecular H2 and H202 Nhich is actually observed. Thus the existence of the molecular yield in water decom- position is ascribed to the existence of the hot spots. A calculation of the ratio of the initial rate of recombination to the initial rate of diffusion out of a cluster is readily made if the initial radical distribution is assumed to be Gaussian in distance from the centre.The ratio is equal to 0.054 cc No/Db, where NO is the initial number of radicals, CL and D the reconibina- tion and diffusion coefficients, and b the mean radius. If we take NO = 12 and cc/D = 5 x 10-7 cm, the ratio becomes unity if 0 = 32A, and will be larger if b is smaller. For larger radical clusters (higher No) the ratio will be proportionally greater. As the ionization density of the primary radiation increases, the isolated radical pairs must approach closer and closer to the clusters and eventually will begin to be swallowed up by the clusters. This will result in an increase of the molec- ular yield and decrease of the free radical yield in the irradiation of water.Such a change should not become very prominent until the mean distance between primary events becomes comparable with the mean cluster radius. A quanti- tative theory for the change in molecular yield with ionization density would not be difficult to formulate on the basis of the above considerations, but is perhaps better postponed until detailed experimental data are available. The above formulation has little, if anything, to do with the frequently dis- cussed problem of the diffusion of radicals out of the tracks of fast electrons. The formation of a column of radicals around the track of an electron can hardly be said to occur until the radicals have diffused a distance from the track equal to their mean separation along the track.For a 100 keV electron this distanceAUGUSTINE 0. ALLEN 87 will be of the order of 1000 A, and the process of initial recombination within the radical clusters and loss of the cluster identity by diffusion will be complete by the time this separation has occurred. The assumption of six radical pairs formed per 100 eV is based on the idea that about four pairs appear as free radicals in gamma or X-irradiations, one pair appears as H2 and H202 and one other pair must have disappeared within the radical cluster by recombination between unlike radicals to re-form water. With natural alpha rays the molecular yield is only two H2 molecules per 100 eV and the free radical yield appears to be negligible,26 so that we must conclude that only one-third of the radicals escape from the alpha-ray column as H2 or Hz02. At least one-half of the radicals in the column should, however, combine as like pairs to give molecular products, and if there is segregation within the column as proposed by Lea 24 the fraction will be greater than one-half.We must assume that some H2 and Hz02 molecules formed by initial recombination are destroyed by back reaction to water, occurring as a result of free radical action within the cluster. We then expect that the sum of the molecular yield and free radical yield should increase as the ionization density decreases. The above picture is certainly a grossly over-simplified explanation of the complicated changes that must be occurring within a hot spot. The real processes are probably much less clear-cut than the mere dissociation of molecules to radicals and subsequent recombination of these radicals in pairs. Nevertheless the existence of the hot spots can hardly be doubted, and the connection between these hot spots and the molecular decomposition of water seems highly probable. Work on this paper was performed under the auspices of the United States Atomic Energy Commission. 1 Fricke and Morse, Phil. Mag., 1929, 7, 129. 2 Miller, J . Cheni. Physics, 1950, 18, 79. 3 Hochanadel, J . Physic. Chem., 1952, 56 (in press). 4 Allen, J . Physic. Chem., 1948, 52, 479. 5 Allen, Hochanadel, Ghormley and Davis, 1951, Report AECU-1413 ; J . Physic. 6 Johnson and Allen, paper submitted to the J . Atner. Chern. Soc. 7 Johnson, J. Chem. Physics, 1951, 19, 1204. 8 Fricke and Hart, J . Chem. Physics, 1935, 3, 596. 9 Hart, 1951, Reports AECU-1250 and ANL-4636 ; J . Physic. Chem., 1952,56 (in press). 10 Nurnberger, J . Physic. Chem., 1934, 38, 47. 11 Lanning and Lind, J. Physic. Chem., 1938, 42, 1229. 12 Duane and Scheuer, Le radium, 1913, 10, 33. 1 3 Toulis, 1950, Report UCRL-563. 14 Ghormley and Allen, 1948, Report ORNL-128. 15 Hart, J . Amer. Chem. SOC., 1951, 73, 1891. 16 Risse, Z. physik. Chem. A , 1929, 140, 133. 17 Fricke, Hart and Smith, J . Chem. Physics, 1938, 6, 229. 18 Guenther and Holzapfel, 2. physik. Chem. B, 1939, 44, 374. 19 Dainton and Collinson, Ann. Rev. Physic. Chem., 1951, 2, 99. 20 Krenz and Dewhurst, J . Chem. Physics, 1949, 17, 1337. 21 Wilson, Proc. Roy. Soc. A, 1923, 104, I, 192. 22 Beekman, Physica, 1949, 15, 327. 23 Kara-Michailova and Lea, Proc. Camb. Phil. SOC., 1940, 36, 101. 24 Lea, Actions of Radiations on Living Cells (Cambridge University Press, 1946). 25 Bradbury and Tatel, J. Cliem. Physics, 1934, 2, 835. 26 Dale, Gray and Meredith, Phil. Trans. Roy. SOC. A, 1949, 242, 33. Chrm., 1952, 56 (in press).

ISSN:0366-9033    

DOI:10.1039/DF9521200079

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

RADIATION CHEMISTRY OF PURE ORGANIC COMPOUNDS : BENZENE AND BENZENE-d6* BY SHEFFIELD GORDON -f- AND MILTON BURTON Department of Chemistry, University of Notre Dame, Notre Dame, Indiana, U.S.A. Received 28th January, 1952 In this paper we attempt the use of mass spectrometric and radiation chemical data on the same compounds for interpretation of mechanism of their decomposition. Mass spectrometric data on organic compounds reveal a large number of ions, including many formed by rearrangement from highly excited ions. Products formed by radicals com- plementary to the ions have been found in radiation chemistry. Some effects of impingent particle velocity interpretable in terms of energy transmitted to the individual molecule have also been reported. Energy may be transferred to surrounding molecules or within a molecule so as to favour one type of decomposition and prevent another.In such energy transfer process mechanisms, the same entity may act as both donor and recipient. When ionization and excitation potentials are greatly different the component of a mixture of lower ionization or excitation potential can be a protective agent. In radiolysis of pure liquid C6H6 and of pure liquid C6D6, by 1.5 MV electrons, yields of gaseous pro- ducts are G(H2) = 0.036, G(C2H2) = 0.020 and G(D)2 =- 0.0117, G(C2D2) == 0.0133. Results of radiation chemistry as well as of mass spectrometric measurements indicate two mechanisms of decomposition, by bond rupture and by rearrangement. In a mixture the C6H6 and C6D6 protect each other against radiation. The general concept of pro- tection is discussed in the light of these results. 1.FUNDAMENTALS OF THE RADIATION CHEMISTRY OF PURE ORGANIC COMPOUNDS The phenomena of the radiation-chemical effects of a single high-energy photon or particle include primary and secondary ionization and excitation, charge and excitation transfer, and charge neutralization as well as internal conversion, decomposition processes and free-radical reactions are familiar to the photochemist. The unravelment of some of these phenomena is expedited by studies of organic systems because of the essential simplicity of properly selected organic compounds and because we may, if we wish, select compounds whose molecules (unlike those of water) are not chemically combined with each other.Furthermore, radiation chemists have varied molecular size and structure in a regular way and have thus obtained rather revelatory information about phenomena of great interest. The primary processes of radiation chemistry are ionization and excitation. Ensuant thereon there may be negative-ion formation and there is always neutral- ization. The mass spectrometer used simultaneously as source of energy and analytical instrument has proven a very useful tool for elucidation of some of these processes, particularly such as occur in highly attenuated gases.1 VARIETY OF PRIMARY PROCESSES.-The mass spectrometer indicates that a large number of ions and radical and molecule fragments may be formed in a primary step but with different appearance potentials. However, it must be emphasized that these observations are made on very attenuated gas.Detailed calculations on the basis of the Bethe equation 2 by Hentz 3 show that in an assumed * Contribution from the Radiation Project of the University of Notre Dame, sup- ported in part under U.S. Atomic Energy Commission Contract AT(11-1)-38. -I Present address : Argonne National Laboratory, Lemont, Illinois. 88SHEFFIELD GORDON A N D MILTON BURTON 89 condensed system of hydrogen atoms the energy goes with very high probability to yield the lowest excited and ionized states in ratio dependent on the energy of the incident particle involved. While such a calculation is by no means the last word on such a matter, it does indicate that in a liquid organic system the number of primary products may be considerably more limited than indicated by the mass spectrometer.The mass spectrometer does, however, make certain facts clear. The parent ion is not necessarily the most abundant. This statement has been shown to be particularly true for fluorinated hydrocarbons, in which case Mohler et aZ.4 have suggested the primary removal of an F- ion. Mariner and Bleakney 5 have made similar electron impact studies on formic acid and have also suggested formation of a negative ion as one of the fragments. Mass spectra frequently contain peaks corresponding to ion groupings not present in the parent molecule.6 Such a result means only that in the time before the ion enters the accelerating field, 10-7 sec, it may rearrange and dissociate into other ions.By a careful study of non-integral peaks many metastable ions with half-life of the order of 10-6 sec have been revealed.7 Such ions may de- compose by rupture to yield other ions and radical fragments or may rearrange to give ions and stable molecules. The ions in the latter case need not be present as groupings in the parent ions. By use of a radiotracer technique, Williams and Gevantmans showed that some of the rearrangement fragments found in mass spectroscopy of aliphatic hydrocarbons are indeed produced as free radicals in an early stage of radiolysis of some of these compounds. The diversity of primary products was certainly greater than might be concluded from a necessarily simplified calculation from the Bethe equation. Thus, in a rather fundamental way, it has already been shown in radiation chemistry that, in spite of the enormous energies available for excitation and ionization processes, the excited entities have a comparatively long life and may rearrange before decomposition.mass spectrometer is not a good tool to reveal, for example, the effect of tem- perature on the relative frequency of occurrence of the various primary processes.9 The relative peak heights are changed but such changes are caused by many factors not at all connected with an effect on relative probability of formation of ions. Such information can be obtained only by direct study of the radiation chemistry of compounds involved. Interpretation of the results in such case would be confused by the effect of temperature on possible subsequent free radical reactions.Some effort has been made to find an effect of impingent particle velocity on the relative frequency of the primary processes. Such an effect is indeed existent in mass spectrometry, where the relative ion abundances are very sensitive to the voltage of the impingent electrons. However, the velocities and velocity differences are small and, indeed, trivial in comparison with the velocities of interest in radiation chemistry, namely the velocity of an alpha particle and the approxim- ately 86 times greater velocity of a beta particle of the same energy. The only work reported on this point is indirect and is on liquid systems. Sworski and Burton 10 compared the effects of 1.5 MeV electrons from a Van de Graaff gener- ator and mixed radiation from the Oak Ridge graphite nuclear reactor on a variety of aromatic compounds and obtained the results summarized in table 1 .TABLE 1 .-COMPARISON OF RATIOS OF YIELDS, Hz/CHq, IN RADIOLYSIS OF FACTORS WHICH AFFECT RELATIVE FREQUENCY OF PRZMARY PROCESSES.-The SOME LIQUID HYDROCARBONS P Pie toluene 25.9 16.4 1.58 ethylbenzene 9.7 5.9 1.64 iso-propylbenzene 4.2 2.3 1.83 tert.-butyl benzene 3.5 1.6 2.09 pile irradiation, electron irradiation, e compound90 BENZENE AND BENZENE-& The effect of the slow-particle (neutron) bombardment in the reactor was attenu- ated by gamma radiation and the interpretation of effects was weakened by the choice of liquid compounds for this study. In the liquid state the possibility exists that cage effects will have special influence in slow-particle bombardment, where primary products are formed relatively close together.In some current studies by Dr. S. Schrage a clearer distinction will be sought by studies on methane gas containing radioiodine. In these experiments the relative gamma input in the reactor is markedly decreased by the use of suitable shielding. CHARACTERISTIC CHEMICAL EFFEcTs.-One of the most interesting aspects of radiation chemistry is the relative lack of reactivity of aromatic compounds compared to aliphatic.11 In photochemistry, such a phenomenon is largely unrevealed, primarily because of high excitation potential of aliphatic compounds. However, high-energy radiation is absorbed equally well by both and gives a G value (i.e. yield of molecules per 100 eV input) of about 6 for the decomposition of aliphatic compounds and of about one-hundredth that figure for benzene.I n general, in normal aliphatic compounds, the relative quantity and nature of the products produced is determined by the abundance and nature of the parent groups in the original molecule.ll.12 In a series of studies on fatty acids and other compounds, it was shown by Breger, Sheppard, Burton, V. L., et aZ.13 that for Iarger molecules departures might be noted from the simple rule just cited. For example, radiolysis of palmitic acid led, on this basis, to a disproportionately high yield of pentadecane and carbon dioxide. An explanation suggested for this effect is that in the liquid state deactivation characteristic of a cage effect frequently lowers the excitation energy to a level sufficient only for a decomposition by rearrangement,l4 e.g.However, as has been pointed out,ls such an explanation is inadequate to explain all the curious effects reported. Special structure of the molecules and the inter- relation of the different groupings are also very much involved. ENERGY TRANSFER WITHIN MoLEcuLEs.-Comparative studies of the photo- chemistry and radiation chemistry of a group of alkyl-substituted benzenes (toluene, mesitylene, and ethyl, set-propyl, and tert.-butyl benzenes) 16 have shown that energy absorbed in the benzene ring is transmitted with low probability to the side groups, which decompose preferentially at a bond once removed from the ring, and that energy absorbed initially in the side groups may be transmitted into the ring.The energy held in, or transferred into, the ring is not very effective for decomposition. Studies of the radiation chemistry of pure liquid benzene indicate that the yields of hydrogen and acetylene do not exceed values correspond- ing to G(H2) = 0.036 and G(C2H2) = 0.022 17 and this non-reactivity appears to persist in the substituted compounds. The actual transfer of energy within a molecule, by a process of internal conversion, is well known from photochemistry. This particular series of studies showed not only that the recipient group could be thereby decomposed (a well-known phenomenon) but also that the recipient group might protect an essentially more reactive group from the effects of the energy it absorbed. ENERGY TRANSFER BETWEEN MOLECULES.-In radiation chemistry energy transfer involves both excitation transfer (common in photosensitization processes of photochemistry) and ion transfer.Such transfer processes may be important in mixtures. Provided the states produced by the initial excitation are sufficiently persistent, ionization may be transferred from a molecule of higher ionization potential to one of lower 18 and the latter would thereafter act as if it had been the one initially ionized. The condition for excitation transfer is formally similar. The recipient molecule must have an energy state not too much different in con- tent from that of the donor molecule.19 After energy transfer, the recipient molecule may lose a portion of its energy in collisional processes and even transferS H E F F I E L D GOKDON AND MILTON BURTON 91 energy back to the donor molecule.Ultimately, if the recipient molecule has a sufficiently low excited state its properties and reactivity largely determine the observed chemical effect in the mixture-so far as it is caused by excitation.17 The total chemical effects in a mixture are the resultant of a number of effects including both excitation transfer and ionization transfer. It is possible, with the right relationship of energy levels, that the two effects might actually oppose each other but such opposition does not imply absence of mutual effect. Indeed, it is possible to have mutual protection of the constituents of a mixture, as in the liquid mixture of benzene and cyclohexene,ll which have very nearly the same ionization potentials.20 It has been shown by Manion and Burton 17 that, pro- vided the G(H2) values for the two separate components, A and B, of a mixture are sufficiently different, the combined C(H2) value is given by a simple relation where G(H2, A) and G(H2, B) are the G(H2) values for the pure separate com- ponents, N A and Nn are the respective mole fractions, b is an experimental con- stant and f is the ratio of “ electron numbers ” of molecules A and B.Eqn. (I) has been shown to be consistent with the simple energy transfer mechanism but it cannot be said that the mechanism has been unequivocally established. 2. BENZENE AND BENZENE-& One of the difficulties in the interpretation of radiation-chemical studies of benzene derivatives and benzene mixtures is the uncertainty regarding the mechan- ism of decomposition of benzene and particularly whether free radicals or ultimate molecules are formed in the primary act.There is some information based on bleaching of molybdenum trioxide by benzene gas exposed to light of h < 21 50 A that atomic hydrogen is produced by excitation processes.21 The formation of one part of ethylene and two parts of acetylene in radiation-chemical decomposition of benzene vapour has been interpreted in favour of the free-radical mechanism in that case.17 No other information bearing on detailed mechanism exists. For elucidation of the mechanism of the photochemical and thermal decom- position of acetaldehyde, it has proved very useful to study a mixture of acetal- dehyde and acetaldehyde-d+22 In this work, we have used a similar (but not so extensive) procedure in the study of benzene and benzene-ds.Furthermore, we have determined the mass-spectrometer pattern of benzene-& and have com- pared it with that of ordinary benzene for elucidation of the mechanism of some of the early steps in the ionization portion of the decomposition. EXPERIMENTAL MATERIALS AND HmmrNG.-Benzene, Merck & Co. reagent grade, thiophene-free, was distilled through a 100-theoretical-plats column and the middle fraction boiling at 80.0 C was collected. On subsequent fractional crystallization, the crystals from the fourth freezing were retained ; n:’ = 1.4999. Benzene-d6 was synthesized by Dr. T. J. Sworski according to the method of Ingold, Raisin and Wilson23 (i.e.by four successive equilibrations of benzene with 51 mole % D2S04 in heavy water 24). The sample used had a deuterium content as determined by mass spectrometry of greater than 99 %, corresponding to 5-6 mole % C6DsH impurity. The benzene and benzene-& were degassed, dried over sodium, and used to fill the irradiation cells according to the “ vacuum-filled ” technique described by Manion and Burton.17 The reagents were never distilled through joints or stopcocks lubricated with vacuum grease. Break-offs and glass seals were employed for all gas transportation. The samples were never exposed to the atmosphere after the initial degassing and drying. IRRADIATION CELLS.-cellS used in irradiation were similar to those described by Hentz and Burton 16 except that the windows were ground glass of 5 mils thickness (contrasting with 0.014 to 0.018 in.previously used). The window diameter was 0.9 cm.92 BENZENE A N D B E N Z E N E - d 6 IRRADIATION PROCEDURE.-Technique of irradiation was like that of the previous work of this kind 17 except that cells were cooled with an air stream rather than with drop- ping water. During runs the current was made constant at 1.5 MV and 2.0 & 0-1 PA. For calculation of the amount of energy dissipated in the liquid contents of a cell a cor- rection of 3 % was applied for energy absorbed in the cell window.16 PRODUCT mTERMrNATIoN.-After irradiation, the cell was attached to a vacuum line for extraction and separation of the gaseous contents into two fractions: one non- condensible over liquid nitrogen (- 196" C) and the second non-condensible over ethyl bromide mush (- 120" C). Volumes of gas in these fractions were determined by the semi-micro method of Saunders and Taylor.25 The two separated gas fractions were transferred to glass containers and analyzed on a mass spectrometer for chemical and isotopic composition. They were essentially pure hydrogen and acetylene respectively ; no methane was present in the - 120" fractions.Mass spectra of benzene and benzene-& as well as of products of irradiation of pure benzene were run with a Consolidated 21-102 instrument under the supervision of Mr. Richard Wertzler in the Research Laboratories of the Sinclair Refining Co. at Harvey, Illinois. Product analysis of samples from benzene46 and from the 0.50 electron- fraction samples were run on a 21-103 instrument at Pasadena by Consolidated Engineer- ing Corporation.Certain acetylene results were rejected by us because of evident contamination en route. One sample of product from benzene-& and all other samples not specifically mentioned were analyzed on our own 60" Nier-type instrument (our number 2) by Dr. Sol Davison. DESCRIPTION OF MrxTuRES.-We employ the method of description found applicable to mixtures in earlier work.17 The fraction of electrons in the sample associated with a particular component is called the " electron-fraction '' of that component. In benzene- benzene-& mixtures the electron fraction is practically identical with both volume and mole fraction, for the densities d:: of the pure components are 0.8760 and 0.9456 respectively.RESULTS AND DISCUSSION MASS SPECTRA.-The Bureau of Standards (American Petroleum Institute) data list a total of 41 peaks below 78 in the mass spectrum of benzene. On the basis of the parent peak 78 equal to 100 the total of the other relative peak heights is 127.94 for 70 V ionizing potential. If we exclude all c6 peaks, including c6 itself, the total is still the rather large value of 99.69. It is very improbable that ions of carbon atom content less than 6 would reconstitute benzene in a radiation chemical process. The 100 eV yields for benzene decomposition must be of the order of the G(H2) and G(C2H2) values ; 17 i.e. G N_ 0.04. Since, very roughly, each 100 eV input in high-energy irradiation of benzene is accompanied by formation of four ions (plus electrons), only one in 100 ions (at the most, for this calculation neglects any contribution of excited molecules) is decomposed.A straightforward conclusion consequently is that under conditions of high-energy irradiation of liquid benzene formation of ions in which only a fragment of the ring is present is much less probable than might be suggested by mass-spectral data on the attenuated gas. In mass spectrometry, certain non-integral peaks can be used to distinguish particles of metastable ions 7 with half-life of the order of 10-6 sec. Metastable ions with life I 10-7 sec are unrevealed by such spectra. In the benzene mass spectrum the only non-integral peaks reported are weak ones a t 74.1, 48-1 and 34.7 and doubly charged ion peaks.There is indication of a metastable ion overlap at peak 76. A reasonable interpretation of this result in the light of the re- sistance of benzene to high-energy radiation is that most of the ions produced are the product of decomposition of parent ions (half-life 5 10-7 sec but > 10-9 sec) which would be relatively stabilized in the liquid and even in the vapour state a t ordinary pressures. This stabilization is presumed to occur merely by a process of collisional deactivation before the excess energy content of the ion can cause its decomposition. Such collisions would take place about every 10-9 sec a t 0.1 atm and about every 10-13 sec in the liquid so that all that isSHEFFlELD GORDON AND MILTON BUTON 93 required is a reasonably high probability for energy transfer.In highly resonant benzene the excitation energy is distributed all over the molecule and the change in interatomic distance in any particular region is small; the configuration of the excited state is much like that of the ground state so that, given available energy states for reception of the energy, Franck-Condon restrictions on its transfer will not be great. Such transfers may occur repeatedly, the energy being dissipated throughout a number of molecules or ions until none contains enough energy for decomposition. Benzene thus exhibits a kind of self-protection in radiation chemistry, in which the degree of protection is limited by availability of energy states for re- ception of excess excitation energy of the ions.Comparison of mass spectra of C6H6 and C6D6 gives an indication of the ions stabilized before decomposition. Table 2 is based on the mass spectra of those two compounds.26 The values in the table have been calculated from those data for the two isotopically pure compounds, devoid of C13 content. The sensitivities for appearance of the parent peaks were found to be practically identical. TABLE 2.-RELATIVE ABUNDANCES OF c g IONS IN MASS SPECTRA OF PURE C6Hg AND PURE C6D6 WITH 70V IONIZING VOLTAGE n C6H,+ C6Dtlf C6H/i-!C6D,i ' 6 1 0 0 . 0 100.0 1.0 5 1 4 . 1 5 8.5 1.67 4 6 - 1 9 3.26 1.89 3 1 - 4 7 0.86 1.71 2 4.89 2 . 0 6 2 - 3 7 1 1 - 6 1 0 . 6 6 2.44 0 0.22 0.07 3 . 1 4 a The Bureau of Standards data indicate that for this region normal and metastable peaks overlap.They list peak 76 only. Table 3 compares some raw data on other ion peaks in which we have not yet corrected the values to the base of the isotopically pure (C13-absent) compounds. However, the values correspond so closely to those of the true relative ion abund- ances that the comparisons made on this basis are nevertheless significant. TABLE 3.-uNCORRECTED RELATIVE ABUNDANCES OF SOME SELECTED IONS IN MASS SPECTRA OF PURE C&g AND PURE C6D6 WITH 70V IONIZING VOLTAGE. ALL THE IONS COMPARED HAVE RELATIVE ABUNDANCE GREATER THAN 0.5 172 n C,&' C,D,+ C,,IH,,+/C,,P,l+ 1 0 0 . 0 3.28 0 . 7 1 20.2 2 0 . 0 8 1 7 . 4 2 2.92 5.68 100.0 3 . 4 0.63 1 8 . 5 1 6 . 7 15-65 2-43 5.1 5 1.0 0.96 1.3 1.1 1 . 2 1.1 1 . 2 1.1 The data of tables 2 and 3 and the well-known resistance of benzene to high- energy irradiation evoke the following train of speculation.(i) The zero-point energy of C6D6 carbon-hydrogen vibrations is lower than that of C6H6 by about 1-14 kcal mole-1.27 (ii) This difference in zero-point energies is reflected in an increased resistivity of C6D6 to high-energy radiation (see next section) and in large values of C6Hn+/C6D,+ which almost systematically increase with decrease of n. It is not reflected in the values of C,H,+/C,D,+ shown in table 3, all of which are not greatly different from unity.94 BENZENE A N D BENZENE-& (iii) These differences in the data for the c6 and for the other ions selected suggest a difference in mechanism by which the two groups of ions are formed.The ratio close to unity for the latter group suggests that the mechanism of forma- tion is not very sensitive to zero-point energy. Rearrangement processes 2s would appear to be in this category. It is thus consistent with the facts to suggest that (except for the parent 6-6) those shown in table 3 are formed by rearrange- ment from the parent and those shown in table 2 are formed by successive H or D atom loss in a bond-rupture process from the parent. (iv) The fact that G values are low in spite of the diversity and abundance of ions shown in the mass spectrometer suggests interruption of processes of formation of unstable ions, most of which originate from an ion of common formula. The latter may be formed with very high energy. We may note particlllarly that the relative peak heights with respect to the parents of the doubly-charged parent- molecule ions are approximately 13.6 and 14.6 for CgH62' and CgD62+ respec- tively.Thus, we must accept as real the formation in the primary excitation of a large number of excited states (of unknown abundance) intermediate between the singly charged and the doubly charged ions. Such excited states are either deactivated by collision or decompose, as in the attenuated gas of the mass spectro- meter. Even in condensed systems some small number of them may decompose to yield free atoms, radicals, or even molecules and a new ion. It is because of such processes that benzene yields product when exposed to high-energy radiation. Since more energy would be required for a C-D than for a C-€3 bond rupture and more time for rearrangement of a heavy molecule than of a light one, we may expect that C6D6 would be more resistive to radiation than CGHG.obtained from irradiation of benzene, benzene46 and their mixtures with 1.5 MV eIectrons at 2 PA. The original data from which these calculations were made show that the results were independent of time of irradiation and in good agree- ment. The values G(H2) = 0.035 and G(C2H2) = 0.020 for pure C6H6 are in reasonably good agreement with previously given values 17 G(H2) = 0.036 and G(C2W2) = 0.022. In the values for " pure " C6D6 the third significant figures given for G(hyd) = 0.0117 and G(ac) = 0.0133 are deliberately entered, par- ticularly to emphasize the reality of the difference between the two values.We may presume with reasonable certainty that these values would be respectively practically unchanged for G(D2) and G(C2D2) in isotopically pure CsD6. R E N Z E N E - ~ ~ ON IRRADIATION WITH 1.5 MV ELECTRONS AT 2 pA ; HYD IS TOTAL HYDROGEN ; PRODUCTS RESULTANT FROM IRRADTATTON.-Tabk 4 SUmmariZeS the results TABLE 4.-ANALYSES AND 100 eV YIELDS OF GAS FROM MIXTURES OF BENZENL AND A C IS TOTAL ACETYLENF ." " / , <I __ - G(ac) _ . C6H6, electron Ofexpt. times, G(hyd) - - min € f 2 H D D2 C2H2 C2HD CZDZ fraction 0.0 3~ 60, 40, 40 0,0117 0.0 3.3 96.7 0-0133 0.0 1.7 98-3 0.271 1 40 0.018 30.0 34.9 35.1 0.013 35.4 S.8 55-5 0.50 3 20, 40, 60 0.020 52.1 33.1 14-8 0.014 52.3 21.1 26.1 0-766 1 40 0.026 88*6b 6.4 5.0 0.015 80-2 7-25 12.5 1.00 4 2-60 0.035 - - - 0.020 - - - a Three analyses were rejected because of evidence of Contamination of the samples in transportation.In the sample used for analysis, G values were not determined. Those given do, however, represent the average of three close results. b Because of instrumental difficulty peak heights were uncertain by about 10 % of this analysis ; analyses are thus correspondingly uncertain. While G(H2) > G(C2H2) for benzene, G(D2) < G(C2D2) for benzene-& A reasonable conclusion is that at least two mechanisms are involved in radiolysis of benzene (i.e. both ChH6 and C&) and that these do not contribute in the same way to hydrogen and acetylene production.SHEFFIELD GORDON AND MILTON BURTON 95 The species HD can be formed in significant quantity only by a free radical mechanism ; e.g.Because of zero-point energy differences reaction (2) will occur more frequently than (1) and (3) will be more frequent than (4). Under such circumstances, the ratio of yields HD/D2 should much exceed the mole ratio C6H&6D6, Table 4 shows that this situation is true except for a single case (0.766 electron fraction CgH6) where there is some uncertainty about the analysis. The activation energy differences E1-E2 and &-E3 should be nearly the same and certainly less than the zero-point energy difference between C6H6 and C6D6. No value for the difference in the range 0-1.14 kcal, however, can give the ratio H2 : HD : D2 reported for the 0.5 electron (or mole) fraction mixture if it is assumed that all HI and Dz are formed via reactions (2) and (4).The HI> vield calculated on such a basis turns out to be too high. The data are not adequate for actual calculation of El-E2 (eE4-E3) or of the relative amounts of free radical and ultimate molecule (i.e. rearrangement) decomposition in C6H6 and C6D6 but some relationships may be examined, as in table 5. TABLE 5.-ESTIMATE OF SOURCES OF H2 AND D2 IN 0.5 MOLE FRACTION BENZENE + BENZENE-d6 MIXTURE ON BASIS OF ASSUMPTIONS REGARDING ACTIVATION ENERGY D2, % ~ -~ - E , -E7 HD,% H2, % assu incd , kcal/mole obs. fr:i;l, calc. calc. calc. calc. fry:& ObS. from ( 2 , rearrang. obs. from (4) rearrang. - 33-1 20.9 12.2 52.1 - - 14.77 -- - 0.3 - - I __ 349 17.2 - 7.3 7.4 __ - - 47- I 5-00 -- 5.6 9-7 0.5 - For the calculations of table 5 we assume that the relative contributions of C6H6 and C6D6 to the HD yield are determined by the relative probabilities of H and D formation and that these (for the 50 mole % mixture) are given by the ratio CGH~+/C~D~+ -= 1-67 in table 2.Such assumption is, at best, only ap- proximately valid, 1 -67 2.67 The calculation involves the steps : HD (from (I) calc.) = - HD (obs.), H2 (from (2), calc.) = e(El HD (from (l), calc.), H2 (from rearrangement, calc.) = H2 (obs.) - H2 (from (2), calc.), similar steps for D2, and the assumption that we may justifiably use T N 300" K. Even this Iatter assumption is questionable, for it ignores possible " hot-atom " effects. In spite of the limited usefulness of the results table 5 does show that if a smaII difference in activation energies El - E2 (11 E4 - E3) is assumed, results consistent with our knowledge can be derived.Of the two El - E2 values used, the value 0.3 kcal seems nearer to the truth than 0.5 kcal because it gives lower yields of D2, both by free-atom reaction and by rearrangement, than of H2. The 0.5 kcal figure gives too high a value for yield of D2 by rearrangement. A smaller yield seems more probable principally on the basis of the idea that re- arrangement decompositions of the heavier C6H6 occur more slowly than those of CsH6-an interpretation already shown to be consistent with the data of table 3.96 B E N Z E N E AND BENZENE-& The principal conclusion from table 5 is that the yields of hydrogen in the radiation chemistry of mixtures of benzene and benZene-d6 are consistent with the notion of two simultaneous mechanisms, one of which yields free atoms and the other of which yields ultimate molecules in the primary act.Fig. I , on hydrogen G values, shows that in a mixture of C6H6 and C6D6, the latter protects the former, for tbe H produced (either as H2 or HD) is distinctly below the value predicted merely from the electron fraction of C6H6. Con- versely, the total D yield indicates that this protection may be at the sacrifice of c&6, but the data are not conclusive on this point. Fig. 2 on the other hand demonstrates conclusively that, so far as acetylene yields are concerned, the protection is mutual. One point for total acetylene yield is actually below the G(C2D2) value for pure C6D6. The mechanism of acetylene production is uncertain but the existence of CzHD yield is clear evidence that a free atom mechanism is involved.In explanation FIG. 1.-100 eV yields (G of hydrogen from benzene) bCnZene-& mixtures ; hyd = total hydrogen. / E / ~ I r r o n fraclion, benzevle >. 1.2 1.4 “ 6 1.8 of the ratio G(CZH~)/G(C~H~) N 2 in radiolysis of benzene vapour, Manion and Burton 17 suggested an explanation based on action of “ hot ” hydrogen atoms on the benzene molecule with two acetylene molecules and one ethylene coming from decomposition of each benzene. Ethylene is not a product of liquid radio- lysis so that the mechanism must be quite different in this case. The difficulty here is that the yield of C2HD is rather large for a conventional explanation.An explanation that fits the observations may be summed up in the following state- ments. (i) In the liquid state radiolysis part of the acetylene is formed by a rearrange- ment mechanism and part by the action of free hydrogen atoms. (ii) Interaction of a hydrogen atom with liquid benzene yields only one acetylene molecule. (iii) The acetylene molecule formed by the atomic reaction contains the incident hydrogen atom, This is not the sort of process one would ordinarily imagine. FURTHER COMMENT ON PRoTEc-rIoN.--Tn radiation chemistry the term protection is usually applied to the protective effect of one compound by another, although the protective effect of a benzene ring on the side groups has been demonstrated. In photochemistry, the Franck-Rabinowitch effect 29 is a kind of protection,S H E F F I E L D GORDON AND MILTON BURTON 97 More inclusively defined, protection is the provision of competing means for dis- sipation of excitation energy prior to decomposition.30 Any means which taps off energy from a potentially labile group or increases the probability of such tapping-off is a protective agent. We have seen from consideration of mass spectrometric data that benzene may protect itself by energy-transfer of the type, Similarly, The products, as a result of a series of such energy transfers, are in a potentially less labile state.The best condition for energy transfer is to a level of exactly the same energy content. However, such a transfer is less likely to dissipate energy than the less probable transfer to a slightly different level.In the latter process some small C6H6+* $- C6H6 --f C6H6” f C6H6’. C6H6* -k C6H6 + C6H6” + C6H6’. 4 FIG. 2.-100 eV yields (G) of acetylene from benzene- benzene-& mixtures ; ac = total acetylene. amount of energy is transmitted into translation. Thus, an improved arrange- ment for protection in such a highly resonant substance as benzene is the provision of an additional system with energy levels slightly staggered with respect to the benzene levels. Such a function may be performed by benzene-& for benzene and by benzene for benzene& In the system benzene-benzene-&, ionization potentials and excitation poten- tials are extremely close together. An explanation of protection based on those relationships is satisfactory for substances essentially different in those respects. In this case, mutual protection occurs because the two compounds act to provide for each other more effective and highly competitive processes by which energy can be transferred and dissipated before decomposition has a chance to occur.1 cf. Viallard and Magat, Compt. rend., 1949, 228, 1118, for an excellent brief review 2 Bethe, Ann. Physik, 1930, 5, 325. 3 R. R. Hentz, private communication. 4 Mohler, Bloom, Lengel and Wise, J. Amer. Chem. Soc., 1949, 71, 387. 5 Mariner and Bleakney, Physic. Rev., 1947, 72, 792. 6 Langer, J. Physic. Chem., 1950, 54, 618. and interpretation of some of this work.RADIOLYSIS OF ORGANIC LIQUIDS 7 cf. Hipple, J. Physic. Cliem., 1948, 52, 456 ; Physic. Rev., 1947, 71, 594; Hipple, Fox and Condon, Physic.Rev., 1946, 69, 347 ; Bloom, Mohler, Wise and Wells, J. Res. Nat. Bitr. Stand., 1949, 43, 65. 8 Williams, Jr. and Gevantman, J. Physic. Chem., 1952, 56 (in press). ') cf. Berry, J. Cliem. Physics, 1949, 17, 11 64. 10 Sworski and Burton, J . Amer. Cliem. Suc., 1951, 73, 3890. 11 cf. Burton, J. Physic. Clzem., 1947, 51, 786. 12cf. Honig, Science, 1946, 27, 104; Burton, J . Pliysic. Chem., 1947, 51, 611. 13 cf. particularly Breger, J . Physic. Chem., 1948, 52, 551 ; Burton, V. L., J. Amer. 14 Burton, J. Physic. Chem., 1948, 52, 810. 15 cf. Whitehead, Goodman and Breger, J . Chirn. Phys., 1951,48, 184. 16 Hentz and Burton, J. Ahier-. Clzem. SUC., 1951, 73, 532 ; Sworski, Hentz and Burton, J. Amer. Client. SOC., 1951, 73, 1998 ; Burton, Gordon and Hentz, J . Chim. Phys., 1951, 48, 190. Clzem. SOC., 1949, 71, 4117. 17 Manion and Burton, J . Physic. Clzenz., 1952, 56 (in press). 18cf. Magee and Burton, J. Amer. Chem. SOC., 1951, 73, 523; Magee, J . Physic. Chem., 1952, 56 (in press). 19 Franck-Condon restrictions, of course, apply but such restrictions are not great when the electronically excited state and the non-excited state are of approxi- mately the same configuration. Substances like benzene, which show similar resonance both in the ground and excited states, probably offer such conditions. 20 Price, Proc. Roy. SUC. A, 1940, 174, 207. 21 Krassina, Acta Physicochim., 1939, 10, 193. 22 Zemany and Burton, J. Amer. Chem. Suc., 1951, 73, 499 ; J. Physic. Chem., 1951, 2 3 Ingold, Raisin and Wilson, J . Cliem. Suc., 1936, 916 ; Best and Wilson, J . Chem. 24 The heavy water used (D content 99-4 atom % of total hydrogen) was furnished by 25 Saunders and Taylor, J. Clzem. Physics, 1941, 9, 616. 26 cf. Gordon and Burton, forthcoming publication, for detailed data on C6D6. 27 cf. Randall, Fowler, Fuson, Dangl, Infra-red Determinations of Organic Strirctzrres (D. Van Nostrand & Co., New York, 1949) for characteristic frequencies from which this value is calculated. 28 Rearrangement processes involve rearrangement of atonis and simultaneous change of more than one bond while the atoms of the entity involved are within molecular diameter of each other. Rupture processes involve dissociation of a bond or successive dissociation of bonds, one at a time. An atom or radical escapes in such a process before a second bond breaks. 55, 949. SOC., 1946, 242. the U.S. Atomic Energy Commission. 29 Franck and Rabinowitch, Trairs. Faraday Suc., 1934, 30, 120. 30 Note that protection in this sense is related to decomposition in the same way as is quenching to fluorescence.

ISSN:0366-9033    

DOI:10.1039/DF9521200088

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

RADIOLYSIS OF ORGANIC LIQUIDS THE RADIOLYSIS OF SOME ORGANIC LIQUIDS BY MRS. A. PREVO~T-BERNAS, A. CHAPIRO, C. COUSIN, Laboratoire de Chimie-Physique, Facultk des Sciences de Paris, Paris, France Received 28th Jauuary, 1952 Y. LANDLER AND M. MAGAT A critical review and comparison is made of the possibilities and limitations of two methods which the authors have used up to the present for the determination of the number of free radicals formed in organic liquids by ionizing radiations (polymerization initiation and reaction with diphenyl picrylhydrazyl radical). Attempts are made to correlate the sensitivity of different compounds with their chemical structure and to evaluate the G-values for the radical formation. This leads to an evaluation of the fraction of energy dissipated which is ultimately used for free radical production.TheP R E V O ~ T - B ~ R N A S , CHAPIRO, COUSIN, LANDLER, MAGAT 99 mechanism of the production of free radicals in liquids irradiated in a pile is discussed and a preliminary evaluation of the contributions of different radiations present is given. The possibilities of the determination of the nature of free radicals formed by radi- ations are briefly indicated and some preliminary results are given. While extensive work has been carried out on the radiation chemistry of water and of aqueous solutions there has as yet been comparatively little attention paid to non-aqueous media. The early work, e.g. Kailan 1 and Lind and his school 2 although giving important indications, suffers from the fact that only final products (and sometimes not all of them) were determined.Therefore, on the basis of their observations, it is very djfficult to draw unambiguous conclusions about the mechanism of the primary steps, It is even more dangerous to try to establish a correlation between the radicals produced and the chemical structure of the compounds irradiated. A systematic study of radiation effects on organic substances can alone lead to generalizations and hence to predictions concerning the behaviour of untested compounds and their likely reactions. We think, in particular, that such system- atic invcstigations is best begun by : (i) the determination of the number of free radicals formed per energy unit (ii) the determination of the nature of primary radicals formed in each case.We suppose at this stage that the subsequent reactions of these radicals are known, at least in principle, from general chemical kinetics. absorbed and this according to the type of radiation used; 1. DETERMINATION OF THE NUMBER OF FREE RADICALS FORMED BY IONIZING RADIATION A. METHODS USED AND THEIR LIMITATIONS.-Jt is Clear that in order to deter- mine quantitatively the number of free radicals formed by radiation, the radicals have to be consumed in some kind of reaction which competes favourably with their recombination. At the same time the direct effect on the reactive material must be either kept small by its low concentration or else must be n:easurable independently, with the assumption that no protective effect through excitation or ionization transfer is operative.3 Let us consider the possibility of trapping the radicals R formed by radiation through addition of some stable free radical F, whose disappearance through the reaction F + R + F R (1) could be easily followed.Such a radical could be, e.g., the diphenyl picrylhydrazyl (DPPH) which is highly coloured even in very dilute solutions, thus making colorimetric measurement of concentration easily feasible. Let us assume that reaction (1) requires an activation energy of 5 kcal, while the recombination re- action require no activation energy whatsoever. Assuming that the collision factors are equal, the concentration of F required to use 99 "/o of R by reaction (1) rather than by the recombination reaction (2) is then determined by the equation : ( 3 ) 2R1 -+A (2) e-5000:RT[F] [R] N 100 [R]2 [F] w 105 [Rl.Should the stationary concentration of R be of the order of 10-12 molelcm3 the concentration of F required would only be about 10-7 mole/cm3; this is low enough to make the direct effect negligible. The concentration of R is not100 RADIOLYSJS OF ORGANIC LIQIJIDS the average overall concentration, but the highest local concentration in the im- mediate neighbourhood of the track. In other words, the above calculation applies to free radicals produced by X-, y-, ,&rays and fast neutrons. For slow neutrons acting through the Szilard-Chalmers effect higher concentrations of F may be required and the possibility of its use must be determined experimentally. For x-rays, the local concentration of R may be so high that the required con- centration of F may preclude the assumption that direct effect on F is negligible.In other words, it is possible to determine the number of free radicals formed by X-, y-, ,&rays and fast neutrons by following the disappearance of DPPH, pro- vided that the dosage rate is reasonably low and that the concentration of DPPH is sufficiently high. At lower concentration of DPPH some of the radicals will escape detection, and the apparent value of the number of free radicals formed by radiation will be too low. This is shown in fig. 1 and 2 ; fig. 1 is based on unpublished data of Boag, Chapiro, Ebert and Gray obtained with the 190 KeV X-rays available at the Hammersmith Hospital, the irradiated liquid being CHC13. Fig.2 is based on data of Chapiro, Corval and Cousin 4 and corresponds to the slow neutron irradiation of CH30H in the Zoe pile at Chatillon. FIG. 1.-Moles DPPH x 108 used per cm3 after 1 min irradiation with 190 KV X-ray as a function of initial DPPH concentration (in 10-8 moles/cm3) at various 0 0 0 27.5 r/min ; This very rapid method gives directly the absolute number of free radicals produced but it has several limitations : (i) DPPH reacts thermally and more or less rapidly with double bonds. It can hence be used safely only with saturated and aromatic compounds with the exclusion of all ethylenic groups. (ii) DPPH reacts with water, so that a high degree of dehydration of the tested compounds and careful control measurements are required. It reacts also with organic hydroxyl groups and other labile hydrogens. Another possibility of counting free radicals formed is to let them initiate an easily measurable chain reaction, such as a polymerization.The necessary assumption is that the rate constants of the successive steps (propagation and termination) are well known from independent measurements. The activation energy for the reaction R* t M -+ RM* (propagation reaction) is known in several instances to be of the order of 8 kcal. We find by an argument analogous to that used previously [MI 2 100 e*OOO/RT. [R] > 108 [R]. dose rates. c) 6 Q) 69 r/min; 0 0 0 17.3 r/min. (4)P R E V O $ T - B ~ R N A S , CHAPIRO, COUSIN, LANDLER, MAG AT 101 In a pure monomer its concentration in mole/cm3 is of the order of 10-2.This method can hence be safely used for the same order of radical concentrations (of 10-11 mole/cm3) as the DPPH method.* A control is afforded by the reaction itself: it can be applied with monomer radical concentrations such that the re- action follows the kinetics previously outlined : 5 I M; + M -+ Mi + propagation M; -l- M:, -3 P termination where M is the monomer, and S the solvent. become If the monomer concentration is too low, the predominant termination will RM; + R* -+ M,R (6) 3oo co X / O A 100 200 FIG. 2.-Moles DPPH x 108 used p x cm3 after 12 h irradiation in the Chatillon pile as a function of initial DPPH concentration (in 10-8 moles/cm3). and the order of the reaction with respect to the monomer concentration will be changed. Fig. 3, taken from the work of Chapiro,s shows the validity of the equations derived from the scheme developed above for mixtures of styrene with some aromatic hydrocarbons.The theoretical curves agree well with the experimental results for solutions containing more than 20 % of monomer. With compounds which, by radiolysis, give a larger number of free radicals than the aromatic hydrocarbons, the agreement between theory and experiment is limited to higher monomer concentrations. The method has the advantage of enabling work to be carried out with highly " radiation sensitive " substances giving a large number of radicals/cm3 (provided that the contribution of the monomer itself is relatively small), since the compound tested can be used at low concentrations (5-10 %). This method has, of course, its limitations : (i) the direct effect on the monomer is by no means negligible, however it can easily be taken into account (provided that excitation and ionization transfer do not interfere) ; * According to Gray (J.Chim. Phys., 1951, 48, 172), the concentration of radicals in the immediate vicinity of a 60 KV electron track becomes about 10-9 mole/cm3 at the time of the first collisions between radicals, and drops probably still further to reach very soon the limiting values of 10--10-10--~~.102 RADIOLYSIS OF ORGANIC LIQUIDS (ii) it is limited to monomers and to compounds which are good solvents of monomer and polymer, since complications arise with precipitants and poor solvents, where the kinetic constants and effective concentrations (or activities) of the monomer are different from the normal ones ; (iii) the determination of the absolute value of the free radicals formed depends, as has already been stated, on the values of propagation and termination rate constants. While at the beginning of our work there was reasonable agreement between the values given by different authors 69 7 concerning these constants for styrene, some doubt recently arose8 introducing an uncertainty of a factor as large as 10.B. BASIC ASSuMPTIONS.-h order to correlate the number of free radicals pro- duced by radiolysis with the chemical structure of the compounds irradiated, we have to assume that we are really measuring (i) all the free radicals produced, (ii) only the free radicals. I 25 50 7;5 loo FIG. 3.-Per cent. polystyrene produced after 250 h irradiation with a y-ray source of 400 mg Ra as a function of dilution.Theoretical curves and experimental points. (0 0 0 benzene; 9 0 48 toluene; Q 8 @ m-xylene and 8 8 eJ ethylbenzene.) Let us discuss these assumptions : (i) the assumption that all the free radicals produced are trapped Cprovided the concentration conditions stated above are satisfied) means that the Franck- Rabinovitch cage is not effective in radiation chemistry. This hypothesis has some experimental and theoretical support that will be discussed elsewhere. At any rate the Franck-Rabinovitch cage effect could probably be little effective in polymerizations where the cage is formed by reactive monomer molecules : (ii) the assumption that only free radicals are measured means that neither ions nor excited molecules are responsible for an appreciable part of the chemical effects observed.While nothing precise is known about the possibility of re- actions between DPPH and organic ions, it was shown that addition of benzo- quinone completely inhibits the radiation-induced polymerization, in this case excluding any contribution of an ionic mechanism. The good agreement found between the radiolysis susceptibilities measured by DPPH and by polymerization indicates that DPPH does not react with the ions produced. The possi- bility of a contribution of excited molecules to the reactions observed cannot be excluded as definitely as the contribution of ions. There seems as yet to be no theoretical estimates of the relative importance of free radical production and of electronic excitation by electron impact.The experimental evidence is also scarce and contradictory. The excitation function for the nitrogen molecule for instance, given by Hanle,g seems to indicate that the cross-section decreases much faster with increasing electron velocity than does the ionization cross-section. Hence the probability of excitation by electrons of a rather wide velocityP RE vo # T - B 6 R N A s , c HAP I R 0, c o u s I N , LAND L E R , M A G AT 103 distribution is expected to be smaller than the probabilities of ionization, of dis- socation or of ionization-dissociation. This is in agreement with the observations of Mme Lousteau 10 who found a photon yield of only 10-7 per ion pair for gaseous nitrogen bombarded with x-rays from polonium.Neither could Boag and Gray find any light emission from HzO bombarded by electrons except, of course, the Cerenkov spectrum.11 The situation is quite different for atoms, where the relative probability of excitation rather increases with the electron velocity.12 A high yield of chemically active emission is further suggested by the recent experiments of Richards and Cole13 and Richards and Dee.14 Since no definite decision can be made, we shall consider for the time being, that all the effects observed (DPPH disappear- ance and polymerization) are due to free radicals only, this point being liable to revision, should further experiments disprove our assumption. RESULTS Various types of radiations (ct-, p-, y-, X-rays fast and slow neutrons) have been used in our experiments but it is only with y-rays and with slow neutrons that there are as yet sufficient data to permit a tentative discussion.~-RAYS.-AII our experiments were made with yrays from a Ra source, kindly lent by Mrs. Joliot-Curie. In columns 2 and 3 are given the relative numbers of free radicals produced under identical conditions 15 in the various compounds listed in column 1 as measured by DPPH and polymerization respectively. The values in brackets are to be considered as uncertain because of the perturbation of the reaction by secondary effects ('precipitation or poor solubility of the polymer formed in the polymerization and unusual variations of the reaction with time for DPPH). We sce that in all other cases there is satisfactory agreement between the results obtained by both methods.It appears that the various organic compounds arrange themselves in the following sequence according to increasing susceptibility to radiolysis : aromatic hydrocarbons, aliphatic hydrocarbons, oxygenated compounds, halogenated compounds. Carbon disulphide is the most stable compound so far studied. The particular stability of aromatic hydrocarbons has already been emphasized by Burton and his school,l6 who have also found the sequence: benzene, toluene, ethylbenzene. The NO? group decreases this stability somewhat more than CH3 but much less than chlorine. The introduction of a second C1 atom doubles the effect. This is generally true for an ac- cumulation of chlorine in the compound.Oxygenated compounds are arranged in the sequence : ethers, alcohols, esters, ketones. Some slight stabilization probably occurs in esters containing double bonds. While the results listed in columns 2 and 3 are not dependent on the absolute value of the dose and are hence relatively certain, the results given in columns 4 and 5 are dependent on the dose and are thus only tentative. Absolute dose determination is extremely diffi- cult in our arrangement and the dose was evaluated assuming that the radium source is homogeneously distributed along the axis of a cylinder of the same height as the ampoules whose centres are located on a circle of 4-5 cm radius around the source. This evalu- ation leads to a dosage rate of 1.7 r/min against 3 r/min previously calculated assuming the source concentrated in the centre of the circle.We think that the GR values given now are somewhat closer to reality than the values given previously.lsb The energy utilization yield ( Y %) gives the percentage of the energy absorbed which is used for the ultimate free radical production, assuming that 2 radicals are formed per bond broken and taking for the bond energies the following approximate values : C-C = 80 kcal, C-Cl = 70 kcal, CBr = 65 kcal. Except for acetone, chloroform and carbon tetrachloride, the Y values seem reasonable. A possible explanation for the extremely high values in the last two cases will be proposed at the end of this paper. The agreement between DPPH and polymerization results is less satisfactory if one compares not the relative values but the absolute numbers of free radicals formed in a given compound under identical conditions.A large part of this discrepancy is probably due to the uncertainty in the values of propagation (1,) and termination (k,) rate con- stants of the styrene polymerization. The results are listed in table 1.1 04 compound carbon disulphide benzene styrene toluene m-xylene ethylbenzene acrylonitrile propionitrile nitrobenzene n-heptane iz-octane cyclohexane ether dioxane methanol propanol perdeuteromethanol (CD30D) methyl acrylate methyl methacrylate ethyl acetate vinyl acetate acetone chlorobenzene o-dichlorobenzene ethylbromide 1 : 2-dichloroethane chloroform bromoform carbon tetrachloride RADIOLYSIS OF ORGANIC LIQUIDS TABLE 1 CR relative number of fuee radicals produced per cm3 per unit dose _______~__ (number of free polymerization radicals produced per 100 eV) DPPH 0-7 1.0 1-75 - - - - 2.0 3.5 4.4 5.2 7.4 11.2 13-2 12.5 15.0 13.0 - - 18.4 25.5 12.4 25.4 25.6 33-0 57.0 107.0 (72.0) - - 13.6 15.0 18.7 (50.0) 13.0 27.0 65.0 - - - - (200.0) 0.85 1.8 1.6 3.1 6.3 9.0 2-7 3.9 4.5 9.9 11.4 14.3 24.5 20.0 24.0 30.0 23.0 23.5 27.5 32-0 33.0 50.0 17.5 30.0 28.0 41.0 59.5 57.0 (70.0) Y energy yield in "/, - 3.0 2.8 5.3 11.0 15.0 4.6 6.8 7.8 1742 20.0 24.5 42.0 34-0 41.5 51-5 39.0 40.5 47.5 55.0 57.0 87.0 26.0 45.5 39.0 61.5 88-4 80.0 (105.0) * calculated from molecular weight data.t calculated from polymerization rate data. Table 2 gives the results obtained for styrene from different polymerization data and using the various k,/k,+ values, corrected for 15" C , available in the literature.It also gives the number of free radicals measured with DPPH in benzene which gives nearly the same radical yield as styrene. SLOW NEUTRONS.*-~~OW neutrons are capable of producing free radicals by the follow- ing mechanisms. (i) The Szilard-Chalmers effect : RA -k 11 -.* RA' y -> R + A' (7) A' being a stable or radioactive isotope of A. Two radicals ( A and R) are produced per neutron captured. However, fairly early in our work,17 we found that these primary free radicals could not account for more than a few per cent. of the polymerization occurring in pure styrene or in styrene-halogenated hydrocarbon mixtures. By far the largest part of the radicals are hence produced by the secondary effects.(ii) As can be seen from eqn. (7) y-rays are emitted during or immediately after the capture process (capture y-rays). These y-rays of several MeV are emitted either in a single quantum or in a succession of photons of lower energy (cascade) and are partly absorbed, producing a certain number of free radicals. (iii) If A is a radioactive isotope, it will decay to a stable nucleus, either directly or in several steps and usually with the emission of one or several 8- and y-rays. The energies and number of these " decay fi and y " radiations are known from nuclear data. They are also, at least partly, absorbed in the medium. We thank the Com- missariat de 1'Energie Atomique and particularly M. Ertaud for their helpful co-operation. * These experiments were carried out in the Chatillon Pile, Zoe.PREVO$T-B~RNAS, CHAPIRO, COUSIN, LANDLER, MAGAT 105 TABLE 2 method used no.of free radicals produced per cm3/sec inhibition period of styrene polymerization by benzoquinone * absolute value of polymerization rate assuming at 15" C : 6.5 x 1010 1.1 x 1010 (i) k,/k,: = 0.0043 8 1.7 x 1010 (ii) k,/kt4 = 0.0082 6 0.45 x 1010 (iii) k,/k,& = 0.0105 7 0.29 x 1010 DPPH + benzene (corrected for styrene) 1.3 x 1010 * Assuming one benzoquinone molecule is used per free radical ; this probably leads to a too high value for the number of free radicals, because of the possibility of a co- polymerization reaction, benzoquinone-styrene. j- This value is independent of the ratio of propagation and termination constants.Chain transfer was neglected. The correction that would ensue is smaller than the uncertainty concerning the molecular weight as viscosimetrically determined. (iv) The recoil energy of the capture y's is usually much in excess of the dissociation energy E of the chemical bond ruptured. This excess energy which can be very high (30,000 kcal in the case of H) will be carried off as kinetic energy by R and A (" hot " atoms and radicals) and will be consumed in collisions with surrounding molecules. Free radicals (sometimes " hot ") will be formed as a result of some of these collisions in the process. (8) or by activation of RB, leading for instance to an opening of the double bonds. combination of molecular weight and polymerization rate -f A' (" hot ") + R'B --f A' (less " hot " - > 80 kcal) + R' + B - E, H H A' (" hot ") + RC=CH2 -+ A' (" hot " - > 40 kcal) + RC-CH2 (9) I 1 (v) After A' has spent sufficient of its kinetic energy, it will act as a normal free radical; in our case, it will start a polymerization chain and remain fixed to the polymer.However, if it is a radioactive nucleus, it will decay and eventually become another nucleus, say C , of different valency than A , hence, recreate a free radical. (vi) In addition to slow neutrons, y-rays from the fission process and from the fission products,' as well as a few " fast " neutrons are always present in the pile (" pile y-rays ") and also contribute to free radicals production. In order to separate the contributions of these various effects the following procedure was adopted. (i) The amount of free radicals produced by the Szilard-Chalniers effect proper (two radicals per neutron captured) was calculated from the known cross-sections, irradiation time and the neutron flux in the particular channel.An ampoule containing pure styrene was present in each experimental series for eventual monitoring. (ii) The effect of capture y's was assessed from the known number of neutrons captured and from the efficiency of y-rays of Ra. Since the overall contribution of this effect is rather small, the final result is not appreciably influenced by even relatively large errors which may be due to the assimilation of 5 MeV capture y-rays to 1-2 MeV Ra y-rays and to the neglect of a possible cascade emission. (iii) The contribution of decay P- and y-rays to the overall effect which is important if, e.g., Br and Cl atoms are present, was directly determined for bromine by measuring the rate of polymerization of styrene after addition of known amounts of radioactive (pile-irradiated) ethyl bromide.Fig. 4 shows the time variation of the square of the rate of polymerization (which is proportional to the dosage rate), when pre-irradiated ethyl bromide was added to styrene. The figure shows plainly the period of 33.6 h of 82Br and gives an indication of the 4.5 h period of SOBr. The radical yield is rather low as deduced from this experiment pertaining to ,B-rays of 82Br (e.g. GR = 0.05 in a mixture of 60 % styrene + 40 % ethyl bromide). The con- tributions of /3-rays of 80Br and of 36Cl was calculated on the bases of these GR values, taking into account the amount of energy absorbed.The number of disintegretion * It can be safely assumed that the pile P-rays do not penetrate to the irradiated substance in any significant quantity. D106 RADIOLYSIS OF ORGANIC LIQUIDS and of isomerization y-rays was calculated from the cross-section, neutron flux and nuclear data. The contribution of these radiations (particularly the important effect of isomeriza- tion y-rays of 80Br) to free radical formation in mixtures of styrene and organic bromides and chlorides has been calculated solely from values obtained through application of the corresponding absorption coefficients and of GR values listed in table 1 . There are hence rather large uncertainties in this evaluation.(iv) The contribution of free radicals formed by decay of Br to Kr was calculated in a manner analogous to (i). (v) The determination of the contribution of the pile y-rays and of fast neutrons present in minute amounts became possible only through the development of the DPPH tech- niaue. It is based on the following idea.4 If one irradiates in the pile a solution of For CI their number is very small and was neglected. DPPH in a solvent containing, for ';example, some FIG. 4.-Time dependence of the rate of polymeri- zation of styrene induced by radioactive ethyl- bromide at 37" C. (Correction for thermal poly- merization was made.) hydrogen atoms besides carbon and oxygen, the discoloration of DPPH is produced by all the radicals formed in the processes (i), (ii), (iii) and (v).If, however, hydrogen is replaced completely by deuterium, the effect observed is due entirely to pile y's and fast neutrons. Assum- ing that the y-ray effect is wave length independent, one can hence express the effect of pile y-rays and of fast neutrons as equivalent to the effect of Ra y-rays in some standard arrangement. Knowing the relative radical yields listed in table 1, one can then calculate the number of free radicals produced by pile y-rays and fast neutrons (counted as y-rays) in any particular mixture. In our experiments, we used CH30H and CD30D as solvents," because the latter compound was available in the laboratory. The measurements were fairly reproducible but we had to introduce a correction for the thermal reaction between alcohols and DPPH which is responsible for an appreciable part of the total effect with CD30D and hence decreases the accuracy of determination.A perdeuterated compound which is expected to show no thermal reaction with DPPH is now being prepared and the pile y-effect will be rede- termined. (iv) The difference between the number of free radicals required to explain the ob- served polymerization and the number of free radicals due to the effects (i) to (v) is, for the time being, assumed to arise from collisions of " hot " recoil atoms ; of course, the errors involved in the evaluation of all the other effects accumulate in this value. Only with pure styrene, where no contribution of decay radiation is liable to occur, can this number be assessed with a reasonable margin of error.According to the par- ticular set of propagation and termination constants adopt M it is found that the " hot " atoms form respectively about 2000 or 500 secondary free radicals per neutron captured. If one assumes that only mechanism (8) is operative, the energies necessary to produce these numbers of free radicals would be 85 x 103 kcal and 20 x lo3 kcal respectively while the recoil energy available is of the order of 30 x lo3 kcal. The energy necessary for the production of these numbers of free radicals would be of course smaller if * We have ascertained that the radical yield was the same for CH30H and CD30D (table 1).P R E V O S T - B ~ R N A S , CHAPIRO, COUSIN, LANDLER, MAGAT 107 mechanism (9) contributes to their formation, but, at any rate, the energy utiiization yield remains rather high.Fig. 5 shows for a few mixtures the relative contribution of the different mechanisms to the free radical formation. In pure hydrocarbon systems (styrene and styrenefcyclo- hexene) it is the contribution of the " hot '' recoil atoms which is by far the most pre- dominating. The situation is quite different in styrene+ halogenated hydrocarbon mixtures, where the contribution of decay radiations may become decisive. This effect becomes very important even at low concentrations of halogens. (1) (2) ( 3 ) (4) FIG. 5.--Origin of free radicals in pile irradiated systems. (I) pure styrene ; (2) 60 % styrene + 40 % cyclohexane ; (3) 60 % styrene + 40 % ethyl bromide ; (4) 60 % styrene t 40 % chlorobenzene ; (5) 60 % styrene + 40 % dichlorobenzene.(a) contribution of the Szilard-Chalmers effect. (b) ,, ,, capture y-rays. (c) ,, ,, decay p-rays. (d) ,, ,, hot atoms. (e) ,, ,, disintegration and isomerization y-rays. (f) ,, ,, pile y-rays and fast neutrons. (8) ,, ,, thermal centres. For dichlorobenzene the values of c is the maximum value according to the GR values listed in table 1 for the corresponding compounds. 11. DETERMINATION OF THE NATURE OF FREE RADICALS PRODUCED BY RADIOLYSIS WITH 7-RAYS Our knowledge of the nature of free radicals produced by radiations in liquids is even scantier than our knowledge of their number. Even for water the problem is far from being settled.18 There could a priori be some hope that mass-spectro- metry would at least tell us what free radicals are formed by the mechanism ABww+A+ + B + 2e.Unfortunately, a closer examination of the data shows that in many instances, the reaction goes through an intermediate state, probably an excited- ABf ion of considerable life-time. We think less of relatively rare " metastable " ions than of ions whose formation requires a rather profound reorganization of the original molecule : e.g. C4H9+ ions produced from the 3 : 3-diethylpentane and which give rise to the strongest peak in the mass-spectrum of this compound.19 In the liquid state these intermediate excited ions, of life-time at least 10-10-10-12 sec and probably more, will suffer a large number of collisions which may influence their mode of decomposition. This mechanism of free radical formation is of course not the only one and the overall spectrum of free radicals will in general be different from the one deduced from mass-spectrographic data and will depend on the composition and the state of the surrounding medium.A striking example is given by Hamill, Williams and Voiland20 who found that in the gas phase, at 500 mm Hg pressure, n-pentane on radiolysis yields108 RADJOLYSIS OF ORGANIC LIQUIDS preferentially methyl radicals ; in the liquid phase the main products are vinyl and sec-amyl radicals. The mass spectra would suggest ethyl radicals as main products. The general way to solve the problem of the nature of radicals formed is the same as used for the determination of their total number, i.e. trapping them with reactive entities, such as " stable " free radicals or monomers, and analysis of the final products.This method was successfully used by Dainton21 to show the formation of OH radicals in the radiolysis of water and in an elegant way by Hamill and his co-workers who tagged the radicals with radioactive iodine and bromine. We are now trying to trap the radicals produced with DPPH and to separate and identify the products formed. Fig. 6 reproduces the absorption spectra of mixtures of DPPH derivatives formed by the irradiation of different compounds (e.g. CHC13 and CCl,) in presence of DPPH. With chloroform and methanol at least two substances could be separated by paper chromatography of the irradiated products. We hope to be able to identify them with micro infra-red spectroscopy and eventually to achieve further separations.FIG. 6.-Absorption spectra of DPPH, diphenylpicryl hydrazine and the. products ob- tained by irradiating DPPH in CCI4 and CHC13. Initial DPPH concentration 2x 10-8 moles/cm3. Finally, some indications can also be gained from a detailed study of the kinetics of the disappearance of DPPH during the irradiation. For instance, one observes (with CHC13 and with CC14) a post irradiation effect, the DPPH continu- ing to disappear after irradiation has ceased. This " after-effect " is much more important with CCl4 than with CHC13. From the work of Gunther and his co-workers,23 it is known that CHC13 and CCl4 yield HCl and Cl;! respectively on irradiation. We attribute the after-effect to a relatively slow reaction (as com- pared with the free radical reaction) of HCl and Cl2 with DPPH.Gunther and Cronheim 23 attributed the HCl production in chloroform to a chain reaction of the type : C1* + CHC13 -+ HCl + CCl; CCl; + CHC13 -+ C1* + CC13CHC12 Cl; + CHC13 -+ HCl + CCl., and similar reactions with H atoms.PREVOST- BERNAS, CHAPIRO, COUSIN, L A N D L E R , MAGAT 109 Such a chain would be impossible in the presence of DPPH which is known to act as a chain terminator.24 We have hence to assume that C12 and HCl are also produced in one step, either through CC14 % x ~ % + C12 + CCl2+ -t 2e (1) and CHC13 -xxxx%+ HCl + CC12+ + 2e, (2) respectively or by electron capture,257 26 or else, competing with the " normal reactions ',, CC14 + e + CC1 + Cl2 + C1- cc13 \%A++- --f cc12 + CI2 CC14 \-+V -+ CCl3+ + C1 + 2e CHC13 A\W+ CHC12+ + C1 4- 2e CC14 xxxxxb--f CC13 + C1, (3) (4) reactions (l), ( 3 ) and (4) being relatively more important as compared with the reactions (5), (7), etc., for CC14 than the corresponding reactions with CHC13.Experiments are being carried out to establish this relative importance of the two types of reactions in CCl4 and CHC13. It is important to notice that the energy required for reactions of the type (l), (2), (3), (4) is only slightly larger than that necessary for the reactions (5), (6), (7) (the difference is only of about 13-15 kcal for reactions (4) and (7)), while twice as many radicals are produced. This would explain the abnormally high values of the energy yield Y for these compounds as seen in table 1 which were calculated on the assumption that only (5), (6) and (7) type reactions are occurring. Much further work is of course necessary before these ideas could be con- sidered as more than convenient hypotheses. 1 Kailan, Ber. Wen. Akad., from 1910 onwards. 2 Lind, Chemical Efects of Particles and Electrons (New York, 1928). 3 Burton, J. Chem. Educ., 1951, 404. 4 Chapiro, Corval, Cousin, Compt. rend. (in press). 5 Chapiro, J. Chim. Pltys., 1950, 47, 747. 6 Bamford and Dewar, Faraday SOC. Discussions, 1947, 2, 310. 7 Melville and Valentine, Trans. Faraday Sac., 1950, 46, 210. 8 Matheson, Auer, Bevilacqua and Hart, J. Amer. Chem. Soc., 1951, 73, 1700. 9 Hanle, Physik. Z., 1932, 33, 245. 10 Mme Lousteau, Thesis (Paris, 1951), as yet unpublished. 11 Gray, J. Chim. Plzys., 1951, 48, 172. 12 Bethe, Handbuch der Physik, 2nd ed., 1933, 24/I, 519. 13 Richards and Cole, Nature, 1951. 14 Dee and Richards, Nature, 1951. 15 for details of geometrical arrangement and experimental conditions, see (a) Chapiro, J. Chim. Phys., 1950, 47, 747 ; (b) Chapiro, Compt. rend., 1951, 233, 792. 16 Hentz and Burton, J . Amer. Chem. SOC., 1951, 73, 532. 17 Landler and Magat, Bull. SOC. Chim. Belg., 1948, 57, 381. 18 e.g. Haissinsky, this Discussion. 19 Tables published by Nat. Bur. Stand. (Washington). 20 Hamill, Williams and Voiland, Brookhaven ConJ (Chemistry, 1950). 21 e.g. Dainton, J . Physic. Chem., 1948, 52, 490. 22 Gunther, Von der Horst and Cronheim, Z. Elektrochem., 1928, 34, 616. 23 Cronheim and Gunther, 2. physik. Chem. B, 1930, 9, 201. 24 Bartlett and Kwart, J. Amer. Chem. SOC., 1950, 72, 1051. 25 Warren, Hopwood and Craggs, Proc. Physic. SOC., 1950, 63, 180. 26 Vought, Physic, Rev 1947, 71, 93.

ISSN:0366-9033    

DOI:10.1039/DF9521200098

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

110 GENERAL DISCUSSION G E N E R A L DISCUSSION Dr. N. Miller (Edinburgh Uuiiversity) (partly conimuniccrted) Mr. Wilkinson and I have now completed our studies of the yield for ferrous sulphate oxidation by alpha-particles in aerated 0.8 N sulphuric acid solution. The absolute yield was evaluated by comparison of the chemical change due to the alpha-particles from a given source of polonium with the ionization current due to the same source in an argon-filled ionization chamber. The results are therefore directly dependent on the value assumed for the quantity W (energy released per ion pair formed) for alpha-particles in argon. Values of this quantity in the literature range from 24.9 to 28.3 eV our corresponding G values for ferrous sulphate oxida- tion by alpha-particles being 6.7 to 5.9.We have found a constant yield to be maintained over a 24-fold range of dose rate and independently of the initial ferrous ion concentration from 10-3 M upwards. An interesting feature which emerges from this work a detailed description of which is being prepared for publication is that the same yield is observed if a sheet of mica of about 1.5 mg/ cm2 mass thickness is interposed between the source and the liquid during irradia- tion such that the average linear ion density of the radiation is considerably increased. In this case the ionization measurements are conducted of course through a similar mica sheet. It is therefore apparent that this yield is character- istic of alpha-particle radiation or at any rate of alpha-radiation of energy not greatly different from that observed in natural radioactive transitions.By using solutions containing dissolved polonium we have also studied the effect of dissolved oxygen on the alpha-ray induced oxidation of ferrous ions in 0-8 N sulphuric acid solutions. The yield is reduced by a factor 1.7 by de-aeration this being the case at initial ferrous ion concentrations ranging from 5 x 10-4 to 8 x 10-3 M. Preliminary experiments also indicate that the chemical effects of polonium alpha-rays at a dose rate of 1775 ergs/g min and of 50 kV peak X-rays at a dose rate of 50,600 ergs/g min when administered simultaneously are strictly additive. With regard to the discrepancy between Dr. Hochanadel’s value for the y-ray yield for ferrous sulphate oxidation and those obtained by other workers little more can usefully be said at the moment.A decision can only be reached after more experimental work has been done. The possibility that solutions made up in one laboratory may behave differently from those prepared elsewhere due to traces of dissolved impurities has however been eliminated. This Discussion has provided us with an opportunity for a direct comparison between such solutions made up in different laboratories in that we have just been visited in Edinburgh by Dr. A. 0. Allen and Dr. T. J. Hardwick who brought samples of ferrous sulphate solutions with them. The following table of results obtained by Dr. RELATIVE YIELD VALUES FOR FERROUS SULPHATE OXIDATION IN AFRATED 0.8 N SULPHURIC ACID SOLUTION Brookhaven solution Chalk River solution Edinburgh solution 50 k V peak X-rays C060 ~ ~ - r ~ y ~ _ _ _ ___ ______ NoC1- i 10 3 M C1- 0.97 0-98 1 -02 0.98 0.99 0.98 1 -00 1 -00 H.A. Dewhurst in our laboratory shows that these solutions and those prepared by us behaved very similarly under irradiation with both 50 kV peak X-rays and Co60 y-rays. Each of the figures in the table represents the average of at least 6 separate irradiations. These results providing a useful corollary to those in the experimental and results section of our paper show that the yield for ferrous sulphate oxidation is indeed reproducible and characteristic of the system. The fact that none of these solutions showed any difference in yield when made 10-3 M 111 GENERAL DISCUSSION in chloride ion also shows that no appreciable amounts of organic impurities were present in any case.I am glad that the explanation of the apparent post-irradiation effect reported by Dr. Wild has now been agreed upon. The effect of temperature on the extinc- tion coefficient of ferric ion in 0.8 N sulphuric acid solution was first noticed by Dr. H. A. Dewhurst in our laboratory some months ago and we are in agreement with Dr. Wild as to its magnitude. We also endorse Dr. Wild’s words of warning about the existence of this effect in regard to the use of this mode of analysis in dosimetry. The only genuine after-efTect my colleagues and I have detected in this system in over 5000 irradiations has been that due to the slowness of the Fenton reaction which is only apparent in practice at initial ferrous concentrations at or below the lower limit of the region of concentration independence.The apparent fall-off in yield observed when the initial ferrous ion concentration is below this limit is indeed largely due to the slowness of this reaction. This effect will however doubtless be dealt with in detail in future publications by Dr. Dewhurst of our laboratory and by Dr. Sutton at Leeds who have examined it independently. For the measurement of low doses the DPPH solutions studied by Dr. Chapiro do appear to have certain desirable features. Although the solutions share with the chloroform$water mixtures discussed in our paper the drawback of not being “ air-wall ” it strikes one at once that the “ yield ” in the concentration-independent region is independent of dose rate which is not normally the case in the chloro- form-+water systems.For use for military or civil defence purposes however such criteria as cheapness ready mass production and stability on storage have also to be considered while as a standard for laboratory work it is questionable whether such a method can offer a serious challenge to the aqueous systems previously disc wed. Dr. E. J. Hart (Argo~zrzc National Lab. Chicago) said I wish to point out the merits of the formic acid dosimeter which appears to have been overlooked in the survey of aqueous systems proposed for dosimetry appearing in table 1. Earlier use of this system has been described 12 but more recently 3 it has been successfully used to distinguish between the radicals (H and OH) which escape from the track and the molecular compounds hydrogen and hydrogen peroxide formed by recombination of H and OH radicals within single tracks.In short the use of the formic acid dosimeter enables one to distinguish between the following reactions H20 = H + OH The ferrous sulphate and ceric sulphate systems have not been developed to a point where separate measurements of reactions (1) and (2) can be effected. This refinement in dosimetry is of importance in chemical and biological systems since hydrogen atoms and hydroxyl radicals are much more reactive than hydrogen molecules and hydrogen peroxide. Dr. W. Wild (A.E.R.E.Harwell) said Dr. Miller’s paper focuses attention on the relatively limited range of dose-rate over which present chemical dosimeters can be regarded as satisfactory. In contrast to the situation in radiobiology it seems unlikely that radiation chemistry will require extensive use of very low dose- rate dosimeters. Any requirement in this range will probably be met from current extensive work associated with civil defence. On the other hand there are many chemical problems in the organic field where insensitive analytical methods will demand the application of large doses if adequate answers are to be obtained. Moreover the large dose-rates obtainable from electron beams and big cobalt 1 Fricke and Hart J . Chem. Physics 1934 2 824. 2 Fricke Hart and Smith J.Chem. Physics 1938 6 229. 3 Hart 1951 Report ANL-4636 ; J . Physic. Chem. 1952 56 (in press). GENERAL DISCUSSION 112 sources make possible this desirable extension of the subject. At present there is no adequately tested chemical system applicable to such work. The discovery of such systems would be of great value. It is also perrinent to note the absence of chemical dosimeters usable in organic solvents. Dr. E. J. Hart (Argonne National Lab. Chicago) said The erratic behaviour in the ratio R between the aerated and evacuated ferrous sulphate systems has been the source of considerable concern to us. Organic impurities were suspected as one of the causes of the large variation in results. In accordance with This idea it was found that upon the addition of formic acid the yield of the y-ray oxidation in the presence of air was enormously increased while the yield in the air-free solutions was reduced.As a result of these findings a study was made of the mechanism of the y-ray induced ferrous sulphate + formic acid + oxygen reaction. While the detailed kinetics of this reaction are undoubtedly very complex the existence of an oxidation chain of 30-40 units in length greatly simplified the mathematical treatment of the data. The reaction is initiated by (1 a) (W ( 2 ) ( 3 ) H20 = H + OH H20 = +H2 + iH202 OH + HCOOH = H20 + HCOO H + HCOOH = H2 + HCOO propagated by reactions (2) (4) (3 ( 5 2 ) and (6) HCOO + 0 2 = HO2 + C02 Fe2+ + HO2 = Fe3-1 + HO2- H+ + H02- = H202 Fez+ + H202 = Fe3+ + OH- + OH and terminated by Fe3+ + HCOO = Fe2+ + H+ + C02 Fe2+ + HCOO = Fe3+ + HCOO- Fez+ + OH = Fe3f + OH- Fe3f I- H02 = Fez+ + Hf 4- 0 2 .(4) ( 5 ) ( 5 4 (6) (9) (10) It was further found that ferric sulphate can be efficiently reduced by y-rays in aqueous solutions containing formic acid. The above mechanism provides a simple explanation for deviations of the ratio R from the generally accepted value of about 2.3 when the contaminant is a readily oxidized organic impurity. Prof. Milton Burton (Nutre Dame University Indiana) said The problem of high-intensity actinometry is not in so sad a state as Dr. Wild has suggested. We have merely to make more extensive use of systems in which G values are small.For example we see in the paper by Gordon and myself that benzene46 is even more resistant to radiation than benzene. We need not however confine ourselves to measurement of hydrogen and acetylene production in benzene radiolysis. The value of G (diphenyl production) is very much smaller so that production of (C6H5)2 can be studied when intensity effects are great. Furthermore we can follow a suggestion made by Dr. Maddock (see later). A concentration of 0.2 terphenyl in benzene gives a particularly good scintillation counter. If this physical effect involves specially effective energy transfer from benzene to terphenyl it should be reflected in G values even lower than those heretofore determined. Dr. T. J. Hardwick (N.R.C. Chalk River) said There is an upper limit to the independence of the yield on dose rate in the oxidation of ferrous sulphate in air- saturated 0.8 N sulphuric acid by X- and y-rays.This limit varies slightly with 113 GENERAL DISCUSSION the temperature and the concentration of ferrous ion factors which affect the rate of the ferrous ion + hydrogen peroxide reaction. A decrease in yield at dose rates above 4200 r/min was found on irradiating 10-3 M ferrous sulphate solution with 2000 kVp X-rays at 22" C. However as a pulsed source was used the upper limit of dose-rate independence for continuous radiation may be higher by a factor of 3 or 4. Using a pulsed 50 kVp X-ray source a decrease in yield appeared in the region of dose rate above about 200 r/min. In Ihis case the limit may depend consider- ably on the particular conditions used but where larger dose rates are used experi- ments should be carried out to show the effect of dose rate.Using a continuous source of radiation may increase the critical dose rate by a factor of 3 or 4. The yield on irradiating ceric sulphate is 0.8 N sulphuric acid shows no effect of dose rate up to the limits which were available experimentally. Using a pulsed 2000 kVp X-ray source independence of the yield on dose rate was found up to 45,000 r/min. Using a pulsed 50 kVp X-ray source no effect on the yield was observed up to 140,000 r/min. As pulsed radiation sources were used it may be estimated that for continuous radiation the upper limit of dose-rate dependence must lie beyond dose rates 3 or 4 times higher.Dr. M. Haissinsky (Laboratobe Curie Paris) (partly communicated) In con- nection with Miller and Wilkinson's results I would like to report some experi- ments carried out in our laboratory by Mrs. Anta which show the difficulties encountered in using polonium as ionizing source and also the rather complicated behaviour of ceric salts irradiated by this source. The starting point of this research was the conclusion drawn with Lefort from the results obtained on the irradiation of Ce4f by a-rays of Rn namely that the reduction is in this case mostly due to H202 produced in the primary step. On the other hand Bonet- Maury and Lefort have stated that the x-rays of Po not only do not produce H202 in acid solution but also that there is decomposition of this compound if added previously.It seems interesting therefore to study the action of these rays on ceric ions. Working under somewhat different conditions (the Po salt being homogene- ously dissolved no reagent present in the solutions for H202 detection long irradiations etc.) we have observed that the decomposition of H202 in pure sul- phuric acid is slowed down with time of irradiation and that the quantity remaining in solution tends towards a limit. On irradiating the same solution initially without H202 the latter is formed and the rate of formation also diminishes with time (the decay of Po being negligible during this time). Again a concentration limit is obtained which for the same Po intensity is the same as that obtained previously viz.30 x 1016 molecule/cm3 for about 0.6 mc/cm3. With smaller quantities of Po or in perchloric solution the yields of H202 are even greater and the decomposition curve has here a rather complicated form. As to the irradiated ceric solutions we found that the yield of the reduction diminishes when the dosage rate increases. It depends on the nature of the anion and possibly also on the concentration. Dr. W. Wild (A.E.R.E. Harwell) (comnzunicated) At the Discussion I remarked that Dr. Miller's method of using the ferrous sulphate dosimeter by estimating the ferric ion produced by measurement of light absorption at 302 mp enabled one to take measurements very quickly after termination of an irradiation. Such measure- ments suggested that a considerable post-irradiation increase of ferric ion occurred whose magnitude seemed to be determined by the ferric ion concentration produced by irradiation.Mr. Best has since followed up a suggestion of Dr. Miller's that the observed effect might be a pure temperature effect and this has proved to be correct. It was found that the extinction coefficient of ferric ion in 0.8 N H2SO4 at 302 mp increases by 0.72 % per "C. Published information4 suggests that this 4 Katzin and Gebert Report ANL-4457 (Argonne National Laboratory). GENERAL DISCUSSION 114 peak is a composite one due to sulphate and hydroxy complexes of ferric ion so that the large effect of temperature is not surprising. Temperature records show that the irradiations in question were carried out (in a lead castle at 285 rjminj at temperatures about 10- C lower than the spectro- photometer room.More recent work with the two instruments at the same tem- perature show the absence of any post-irradiation effect provided the ferrous ion concentration is not reduced below 2 x 10-4. If this latter condition is not observed a genuine effect is found which is of small magnitude and is presumably that already noted 5 and ascribed to the slowness of the Fez' $- H202 reaction step in the assumed mechanism. Whilst the above observations do not invalidate the use of this analytical method which is very convenient and accurate if adequate time is allowed for solutions to acquire the temperature for which the spectrophotometer is calibrated they do show that especial care is necessary for the avoidance of unnecessary errors when the absorption cell carrier is not adaptable to thermostatic control as in this instance.Similar remarks are applicable to the use of the ferric thiocyanate method.6 Mr. J. Wright (A.E.R.E. Harwell) said In connection with the use of the benzene-water system for dosimetric purposes we have found a post-irradiation production of phenol which continues for at least 24 h after the end of irradiation and in some cases is as high as 40 % of the phenol present immediately after removal from the source. This effect has been obseried after irradiations in the pile and after irradiation uith a Co y source with solutioix initially aerated and with nitrogen equilibrated solutions using Folin's method of estimating phenol or the 4-amino antipyrelie method described by hlartin.7 The origin of this additional phenol is not known at present but it may be significant that the hydro- gen peroxide produced in the system during irradiation is destroyed during the post-irradiation period.In using the benzene-water system for dosinietric or monitoring purposes it is essential to standardize the time after irradiation at which the phenol concentration is measured and to ensure that changes in irradia- tion conditions do not affect the post-irradiation behaviour. Dr. A. 0. Allen (Brookhaven National Laboratory) said Using 2 MV volt X-rays from the electrostatic generator at Brookhaven we have compared ferrous sulphate oxidalion yields at intensities ranging from about 100 r/min to 10,000.Solutions were irradiated at different distances from the X-ray target and over a wide range of currents. By intercomparison of the results assuming that the ratio of the radiation intensities at the different distances remains constant inde- pendent of the current it could be concluded that the oxidation yield was constant within 5 % over the intensity range studied. In view of our results it appears unlikely that the discrepancy between Hochanadel's value of the yield and other values could be due to the effect of intensity on the yield. A real discrepancy appears to exist between his results 2nd others and further independent deter- minations of the absolute yield in ferrous sulphate oxidation would be most desirable.Prof. Milton Burton ( Uftivcrsity of Notre Dnmc) (co/imunic.nted) The matter of the correct G value for the ferrous sulphate actinomcter is one of great impor- tance for radiation chemists. An explanation must be found for the anomalols situation created by existence of such divergent values. However the funds- mental problem created by the ferrous sulphate actinometer has been obscured by this dispute. That problem is how to account for such high G values whether they be 15 or 20. To that question both experimentalists and theorists must address their efforts. 5 K. H. Krenz and H. A. Dewhurst private communication. Sutton this Discussion. 6 J. Sutton private communication. 7 Martin Anal. Chem. 1949 21 1419. GENERAL DISCUSSION Dr. A. Chapiro (Paris) said Dilute solutions of diphenylpicrylhydrazyl (DPPH) can be used for chemical actinometry of ionizing radiations.This method is very sensitive since differences of a few roentgens can be detected under suitable condi- tions. In one experiment with 190 kV X-rays using bromoform as a solvent a total dose of 35 r administered in 30 sec changed the optical density of the solution by 10 %. In addition the concentration dependence may be interpreted on simple kinetic assumptions and this makes it possible to determine in each case the best experimental conditions to be used. Let us assume that the free radicals formed in the irradiated solvent may either recombine with themselves or combine with DPPH as in the following reaction scheme :8 115 (1) ._ d(DPPH) _ _ _ ~ - V - S .\\\-f 2R* R" + R" -t X R" -+ DPPH + R - DPPH S being the solvent R* the free radicals formed @ the radical yield of the solvent and X a stable molecule.At the steady state k1(R*)2 -'- k2(R*)(DPPH) -L @I = 0 which gives for the reaction rate k9(DPPH)* - 2kl - [ ( X$(DPPH)' 4@1kl ~ )+ - I]. @I kl(W2 k2(R*)(DPPH) k$( D PP H)?' df This equation contains one constant k22,'X'l. For very concentrated solutions and eqn. (7) reduces to I/ -= m. The rate of disappearance of DPPH is equal to the rate of production of free radicals. For very dilute solutions eqn. ( 1 ) can be simplified to I22 V - - (@Z)*(DPPH). kit The constant k?/kl* can be calculated from eqn ( I ) and from the slope of the straight line obtained by plotting V against the concentration of DPPH €or very dilute solutions.Fig. 1 shows the experimental results obtained with chloroform solutions 'irradiated with 190 kV X-rays at 3 different dose rates. The curves are derived from eqn. (I) using k2/kl+ = 66-6 and the experimental points fall close to these theoretical curves. For concentrations of DPPH higher than a given value (DPPH),i,. the re- action rate becomes independent of concentrations the constant rate @I being proportional to the dose rate. (DPPH),i,. is proportional to (@Z)t as can easily be calculated from the kinetic equations. This agrees with the experimental results. It follows that chloroform and bromoform which are the most sensitive solvents for low dose rates cannot be used €or high dose rates since (DPPH),i,.becomes too high for practical colori- metric estimations. However table 1 of our contribution to the present Discus- sions shows that other solvents can be used which give lower radical yields. For 8 A similar interpretation of the dilution effect in water has already been proposed by Dainton and Miller (Proc. Int. Congr. Pure Appl. Chern. (London 1947) 1 77). ( 3 ) 116 GENERAL DISCUSSION instance with methyl acetate dose rates of 5000 rlmin of 37 kV X-rays have been measured. Hydrocarbons could be used for dose rates up to about 200,000 rlmin. The dependence of the chemical yield in chloroform on the type of radiation used is shown in the following table relative chemical yield type of radiation 37 kV X-rays 190 kV X-rays Ra y-rays 1 MeV electrons 4.1 2.41 I 0.4 1 The real energy absorbed per roentgen differs widely for the various radiations listed.An evaluation of the radical yield per unit energy absorbed has shown that the figures for y- and X-rays come close together. There still seems to be a difference between the yield due to fast electrons and that due to y-rays. Finally this actinometry method is very simple. The solutions can be used in the presence of air and the analytical procedure only involves simple colorimetric measurements. FIG. 1. Mr. J. Wright (A.E.R.E. Hurwell) (partly communicated) In quantitative studies with a mixed source of neutrons and gamma radiation one of the important points to establish is whether the chemical effects of the two types of radiation are additive.The evidence on this point presented in the paper is not conclusive but there is additional evidence from studies we have made of the irradiation of ferrous sulphate in 0.8 N H2S04 containing boric acid. In the centre of the pile the rate of oxidation of ferrous ion is a linear function of boron concentration up to 8 x 10-2 M H3B03. The slope of this line leads on the assumption of chemical additivity to a G value for aerated solutions of 3.9 for the additional oxidation produced by the a and Li recoil particles of the 1*B(n aj 7Li reaction. This is much lower than the value of 6.0 given by Dr. Miller9 and this evidence alone suggests that there is less oxidation from the a-particles in the presence of the back- ground pile radiation than from a-particles from a pure a-source.Further evidence was obtained when the ferrous solutions containing H3B03 were irradi- ated in the thermal column of the pile where the oxidation per unit neutron dose in the absence of boron was only one-fifth of that in the centre of the pile. The relation between rate of oxidation and boron concentration appeared to be no longer linear the slope of the curve increasing as boric acid concentration increased. The G values for the additional oxidation due to a and 7Li particles were higher than the 3.9 obtained in the centre of the pile and appeared to increase to a value 9 communicated at this Discussion. 117 GENERAL DISCUSSION which may approach 6 as a greater proportion of the total oxidation is derived from the additional a-particles.Details of this work which is not yet complete will be published separately. It has not been possible so far to control the irradia- tion conditions in the thermal column as carefully as those in the centre of the pile and the spread of results has been somewhat greater. There is no doubt however that the G values for oxidation by the cc and Li particles vary according to the amount of background pile radiation present. Dr. M. Lefort (Institut du Radium Paris) said We can confirm some of the difficulties mentioned by Dr. Wright for the dosimetry in the pile. Mr. J. Pucheault Dr. Haissinsky and myselflo have made a great number of irradiations of boric acid solutions using the French pile ZOE in Chatillon.Like Dr. Wright we used with satisfaction a gold monitor for the determination of the flux of thermal neutrons for each sample because the values given from the control instruments were not suitable for measuring the dose received by the irradiation cell itself. These experiments were made in order to follow the formation of hydrogen per- oxide from boric acid solutions irradiated by the x-rays produced in the solution itself through the nuclear reaction B(n M). It was proposed,ll when this work was begun to use the measurement of the hydrogen peroxide formed as a chemical dosimetry of slow neutrons. Such a method could be adopted only if the forma- tion of H202 was exactly proportional to the flux of thermal neutrons irradiating the solution.From the work we did on boric and borate solutions we find that the following conditions must be fulfilled for such determinations (i) It is neces- sary to correct the results for the action of the known radiation deposit of the pile (y-rays fast neutrons etc.). This correction can be made by measuring the decomposition of different solutions of HzOz placed in the same container as the boron solutions. The reproducibility of results so corrected seems to prove that the assumption made for the correction is reasonable. However it is not certain that the effects of the pile radiation and of the x-rays produced by thermal neutrons are additive. (ii) Solutions containing boron must be either acidic (PH lower than 2) or alkaline (KB02 pH greater than 8).For pH’s between 2 and 8 the results are more complicated and no dosimetry is possible by titration of hydrogen peroxide. This is clearly shown in Dr. Haissinsky’s contribution to the Discus- sion and by the figures given in the publication previously quoted.10 (iii) The G value observed for the formation of H202 even in acid conditions is lower (0-68) than the G value measured when boric acid solutions is irradiated by x-rays from radon (0.9). This is in agreement with the yield for the oxidation of ferrous sulphate in aerated solution mentioned by Dr. Wright as G = 3.9 which is lower than the value of 6.0 given by Dr. Miller using u-rays of polonium. We found too that the reduction of ceric ions occurs in the pile through the B (na) reaction with a yield of G = 2.0 instead of 3.6 with a-rays of radon.12 As we pointed out in (i) the correction for the contribution of the y-rays of the pile may explain these lower values if the two effects are not additive.Prof. Milton Burton (Notre Dame University Indiana) said The track of an ionizing particle contains the residue of primary effects and the secondary effects induced by delta rays. The latter are regions of high concentration of excited molecules and ion pairs and are of the nature of poorly defined “ spurs ” to the primary track. Very little is known experimentally about the composition and internal geography of these spurs. However computation according to Bethe’s method for a system of hydrogen atoms indicates that the spurs are of approxi- mately the same nature for both slow and fast particle irradiation.Thus it follows that the principal parameter which distinguishes effects induced by slow and fast particles is the distance d between spurs the value of d i n liquids being of the order of a few A for slow particle effects and of the order of lOOA for fast ’0 Pucheault Lefort and Haissinsky J. Chim. Phys. 1952 49 286. 11 Bonet-Maury and Deysine Cumpf. rend. 1951 232 1101. 12 Lefort and Haissinsky J . Chim. Phys. 1951 48 368. GENERAL DISCUSSION 13 Dainton and Rowbottom Nature 1952 169 370. 118 particles. The distance d affects competition between chemical processes subse- quent to the early physical effects in a manner characteristic of the particular system affected.When two types of radiation are mixed regions in which d is small are superimposed on regions in which d is large. The effect is not additive. The average value of d is actually decreased by the superimposition of regions of large value of d although not so much as jt would be merely by increasing the intensity of slow particle irradiation. So far as 1 know the theory of such effects has not progressed to a point where we could predict effects in a chemical actino- meter in which a liquid system is employed. On the other hand such difficulties will probably wash out in a gaseous system. Such a gaseous monitor would tend to give purely physical information which would have significance in cases of bio- logical interest when supplemented with data from a liquid actinometer.Dr. S. Schrage in our laboratory is now comparing the effects of electron and neutron irradiation on the system methane + radio-iodine. Dr. A. G. Maddock (Cambridge University) said It appears to me that more careful consideration of the details of degradation of the incident radiation is urgently necessary. Very little direct information on the part played by solvated ions has yet been obtained and when chemical effects with very small G values are under consideration such as the “ molecular ” production of hydrogen de- scribed in Allen’s paper even the less common events in the energy degradation chain may assume importance. For instance some K ionization of the oxygen of the water molecules must take place.There is a definite probability that Auger multiplication leading in some instances to the breaking of both the oxygen-hydrogen bonds will occur. These hydrogen atoms might be conveniently located for molecule formation but the fate of the residual oxygen ion or atom must also be identified. When the energy of the incident electron is very great so that the binding energies of all the electrons of the water are negligible the fraction of ionizing events leading to K ionization must approximate to the ratio of the number of K to the total number of electrons in the molecule i.e. 20 %. When the energy of the incident electron is but a few times the binding energy of the K electrons of the oxygen the fraction of events leading to K ionization is the inverse ratio of the ionization energies of the K and the alternative level-perhaps 1 or 2 %.The probability that K ionization leads to rupture of both hydrogen bonds is much more difficult to estimate. A new theoretical study of the mechanism of thc dissipation of the energy in liquid water of for example incident 1 MeV electrons using stochastic methods would be most valuable. Dr. N. Uri (University of Chicago) (canrrn~mic.utetl) Allen and Miller and Wilkinson deal in their papers with aspects of dosimetry and radiation actino- metry. I would like to make a relatively simple suggestion for the use of the term primary radiation yield and wonder whether this is practicable. If for example a dose of 100 eV leads to 10 radlcal pairs of H and OH (the formation of which might be preceded by ionization) one could compare the energy required to break the chemical bond H .. . OH viz. 5.2 eV with the total dose of 100 eV. As 10 radical pairs are produced the primary radiation yield for the splitting of water would be 0.52. Is such a terminology feasible? On the basis of recent results obtained by Dainton and Rowbottom,l3 it would appear that the primary radiation yield could be close to unity. Dr. J. Weiss (Durhawz University Newcastle) said We have followed with great interest and indeed with admiration the wry interesting work of Dr. Allen and his colleagues. 1 think that it is quite clear that one should expect a greater or smaller extent of primary recombination of the hydrogen atoms in the original tracks the amount of which should depend on the nature of the radiation.system (0.4 M H2SO4) GENERAL DISCUSSION However with hard X-rays and y-rays it seems to me somewhat implausible that the G values for hydrogen evolution should have such a constant value (= 0.65) as Dr. Allen suggests. In some preliminary experiments which have been carried out by Mr. Rigg in this laboratory with X-rays (200 kV) we have found much smaller G values in certain cases Although these are only preliminary results we should like to present them here in the hope that similar experiments will be carried out in other laboratories so that this important question can be clarified in due course. Dr. T. Rigg (Durham University Newcastle) said Dr. Allen in his very interest- ing paper implies that a “ molecular ” decomposition of water into hydrogen and hydrogen peroxide occurs upon irradiation and that the extent of this process will be independent of the solute.The minimum G value for the molecular hydrogen evolved in this process for y-rays or hard X-rays is suggested to be about 0.6 using the value G = 20.6 for the ferrous sulphate actinometry. We have observed similar G values for hydrogen evolution with cerjc salts but for some other solutes we have found in some preliminary experiments sub- stantially smaller yields. Table 1 shows some of the results recently obtained in our laboratory with various solutes. It seems that the solute can apparently influence the rate of formation of molecular hydrogen to an extent which depends upon its oxidizing power so that if some molecular decomposition of water does occur its extent would appear to be possibly less than Dr.Allen suggests. TABLE 1 .-HYDROGEN EVOLUTION FROM AQUEOUS SOLUTIONS IRRADIATED WITH X-RAYS (200 kV) ; mean dose rate 3,200 repimin ; dosimetry based on C = 20.6 for ferrous sulphate in 0.4 M sulphuric acid in the presence of oxygen 1.0 x 10-3 M ceric sulphate (1.0 + 0.1 M H2S04) ; vacuum 1.0 x 10-3 M ceric perchlorate (1 a 0 to 0.1 M HC104) ; vacuum 0.1 M ferric sulphate (pH = 1.7) ; vacuum 0.1 N potassium dichromate ; vacuum pure water ; vacuum 9.0 x 10-3 ferrous sulphate ( 0 2 saturated) 0.1 N potassium permanganate ; vacuum G for H2 It should be noted in particular that in the important case of the ferrous sulphate oxidation the hydroxyl radicals would have no opportunity of reacting with the very small amount of hydrogen present since it has been shown14 that the rate constant for the reaction of hydroxyl radicals with the ferrous ion is about 4 times greater than that of the corresponding reaction with molecular hydrogen.Another important point mentioned by Dr. Allen is that in the irradiation of pure air-free water the rate of hydrogen evolution rapidly falls off due to the accumulation of molecular hydrogen and subsequent back reactions. In this connection it should be emphasized that in radiation experiments on water it is imperative to specify the system completely and to consider (i) the energy and type of radiation (ii) the dose rate and integral dose (iii) the distribution of energy absorption in the liquid (i.e.geometrical considerations) and (iv) the volume ratio 14 Rigg Stein and Weiss Proc. RoJ SOC. A 1952 211 375. 0-15 0-1 0 0-44 0.35 0.30 119 remarks independent of the dose (up to 3.5 x 105 rep) (moles/100 eV) 0.70 0.65 independent of the dose (up to 3.5 x 105 rep) independent of the dose (up to 3-5 x 105 rep) independent of the dose (up to 3.5 x 105 rep) initial yield (maximum) dose = 0-5 x 105 rep independent of the dose (up to 3.5 >i 105 rep) GENERAL DISCUSSION 120 of the liquid and gas phases. Differences in any one of these factors can of course profoundly modify the result obtained. Mr. N. W. Luft (Waltham Abbey) said Could the discrepancies between theoretical and experimental H202 yields in the three representative systems formed in the reaction H + 0 2 -+ HO2? From Robertson's 15 recent value of dealt with by Dr.Allen be partly due to the special properties of the HO2 radical D (H02 + H + 02) = 46 kcal one calculates for gaseous H02 a life of 6 x 10-10 sec at 300" K ; in '' hot spots " it would be smaller. Thus the reverse process would have to be considered too. Prof. I?. S. Dainton (Leeds University) said In his paper Dr. Allen states that an essential difficulty in estimating the free radical yield in water decomposition (GR) is that whatever action is produced in a solute by one radical can in general be reversed by the other and mentions that this difficulty may be overcome by using two suitably chosen solutes present together and each reacting with a different radical.These methods have led to values of GR lying in the range 3-5. It should be pointed out however that it is not necessary to work with 2-solute systems since solutes are known in which the same kind of change may be induced by each radical. The reactions which are initiated are usually chain reactions and examples are the polymerization of vinyl compounds and the radiolysis of hydrogen peroxide. These reactions are of particular interest when they can also be initi- ated photochemically and the quantum yield of the primary act is known since if it can be established that the secondary reactions are identical in both the radio- and photochemical cases comparison of the two overall reaction rates gives the value of GR.This method has been applied by Mr. Rowbottom 16 to the decom- position of concentrated aqueous solutions of hydrogen peroxide taking every precaution to purify the hydrogen peroxide and using the same solution in the same reaction vessel for the radiolysis and photolysis. The value of GR so obtained was 13.4 much higher than the values cited by Allan for other systems either 1-solute or 2-solute. Another exception is the reduction of ceric ion in acid solution which has been studied by Haissinsky and Lefort 17 and yields a value of This discrepancy is perhaps not without significance. On both the H202 decomposition and the CelV reduction the H and OH radicals reinforce one another in their actions (OH radicals are responsible for some of the reduction of CelV which is observed) whereas in the systems from which low GR values have been deduced these radicals could act in opposition.These latter values are therefore to be regarded as minimal. GR. Dr. C. B. Amphlett (A.E.R.E. Harwell) said The problem of solutes com- peting for radicals within the hot-spots and so lowering the molecular hydrogen yield is likely to be very complex. We have to consider competition both for H atoms and for OH radicals and possibly for H202 as well. Although in all cases to date only Hz yields have been measured it cannot be assumed that competi- tion for OH radicals will not affect H2 yields owing to the complicated interplay of radicals in aqueous solutions.The figures quoted by Rigg,ls if substantiated bear out this viewpoint. Thus ceric ion by virtue of its ready reduction by H atoms H202 and possibly OH radicals also would be expected to be a good competitor for radicals produced in hot-spots and yet the molecular hydrogen yield is similar to that quoted by Allen for pure water ; the production of oxygen in the reduction of ceric ion would be expected to favour further competition by virtue of the reaction H + 0 2 -+ HO2. Ferric ion dichromate and permanganate which would not be expected to be more efficient than ceric ion appear to be much more effective competitors. It is noticeable that the three other ions were present 15 Robertson Trans. Faraday SOC. 1952 48 328. 16 Rowbottom and Dainton Nature 1952 169 370.17 Haissinsky and Lefort J. Chim. Phys. 1951 48 368. 18 Rigg this Discussion. 121 GENERAL DISCUSSION at 1000 times the concentration of the ceric ion in these experiments which may explain their greater effect. Results bearing on this problem are of fundamental importance for the understanding of the radiation chemistry of aqueous solutions and require detailed analysis of the kinetics under carefully controlled conditions. Prof. F. S. Dainton and Dr. H. C. Sutton (Leeds University) said In support of his hypothesis that a significant proportion of the products of the primary act consist of hydrogen and hydrogen peroxide Allen has cited several systems both aerated and de-aerated in which molecular hydrogen appears amongst the ultimate products.We have recently been seeking evidence for the production of H202 in equivalent amounts to this hydrogen in ferrous sulphate solutions subjected to X-radiation. The peroxide present initially was estimated from the magnitude of the post-irradiation effects observed in both air-free and aerated solutions when the ferrous sulphate concentration is initially 10-5 N (in 0.8 N H2SO4). At this solute concentration the complete oxidation of the ferrous iron by the H202 formed during the irradiation may take several hours and using the o-phenanthro- line method for [Fez+] estimation we have shown that the rate constant of this reaction is identical with that value published by Barb Baxendale George and Hargrave.19 We have verified that the same curves expressing [Fe3+] as a function of time can be obtained from solutions of Fe2j- Fe37 and H202 made up in com- positions corresponding to that which the irradiated solution (initially containing only Fez-) possessed at the time of cessation of irradiation.In establishing this latter point we made the incidental observation that these curves are unaffected by the presence of excess Fe3+ ion. The experiments in air-free solution are not yet complete but with [Fez+] initial = 1.2 x 10-5 N and a dose rate of 600 r/min of 220 keV X-rays the H202 accounts for about a quarter of the ultimate oxidation yield and corresponds to G H ~ ~ ~ N 0.8. This H202 cannot be due to incomplete removal of air since the residual 0 2 concentration in these experiments was less than 10-9 M.In aerated solutions some of the H202 formed is due to the reaction Fe2+ -I- HO2 + Fe3f + HOa- (1) The maximum contribution which can be made to the H202 yield by this means is obtained by combining eqn. (1) with eqn. (2) (3) and (4) (2) H + OH H20 --\>+ H f 0 2 4 H02 OH + Fez+ -+ OH- + Fe3+ (3) (4) Hence the H202 yield due to (1) cannot exceed 50 % of the initial yield of ferric ion. The observed yield of H202 is in fact 75 % of the initial ferric yield and it is evident that even in aerated solutions a significant proportion of the H202 formed is not due to the dissolved oxygen which is present. Accepting Allen's " hot spot " hypothesis and his figure for GH2 of 0.6 in aerated ferrous sulphate solution it is then possible to calculate a G value for H202 due to " hot spots " (i.e.- 1.5. to origins The result other that than GM dissolved for H202 oxygen). > G for The H2 may value be so significant obtained and is about could be accounted for on the basis of the different spatial distributions of H and OH. Dr. E. Collinson (Leeds University) (comniunicated) The experiments de- tailed in the table below provide further evidence for the formation of hydrogen peroxide during the X-irradiation of both aerated and de-aerated solutions of ferrous ammonium sulphate in agreement with the formation of this substance via the " hot spot " mechanism (cf. the note by Prof. Dainton and Dr. Sutton). The ferrous ion + hydrogen peroxide system is well known to be capable of initiating polymerization of acrylonitrile in aqueous solution and this behaviour was employed as an indication of the presence of hydrogen peroxide in irradiated solutions.19 Barb Baxendale George and Hargrave Trans. Furaday SOC. 195 1 47,462. 122 liquid irradiated 5 x lO-5MFe'f in 0-8 N H2SO4 5 x 10-5MFez" in 0.8 N H2SO4 pure water pure water pure water pure water The irradiations were made with 220 kV X-rays at a dose rate of 180 r/min. Acry- lonitrile monomer was distilled into the irradiated solutions in sufficient quantity to render them 0.7 M in this substance. GENERAL DISCUSSION action after irradiation solution deaerated a n d immediate polymer acrylonitrile acrylonitrile distilled in immediate polymer formation distilled in under vacuum under vacuum solution deaerated a n d polymer visible only acrylonitrile distilled in acrylonitrile distilled i n no polymer formed under vacuum under vacuum solution deaerated ; acry- immediate polymer lonitrile distilled in and ferrous ammonium sul- phate added to give 10-5 M solution state of aeration aerated deaerated aerated deaerated aerated deaerated acrylonitrile distilled in and no polymerization ferrous ammonium sul- phate added to give 10-5 M solution Dr.M. Lefort (Institut du Radium Paris) (communicated) The initial constant yields in the produclion of molecular hydrogen given by Dr. Allen are of great interest.However after a very short time of irradiation hydrogen will escape from the so-called hot spots through the whole bulk of the solution and can then enter into competition with the solutes for reaction with radicals consequently the yield of molecular hydrogen is no longer constant for various solutes. This could be the reason why Dr. Haissinsky and myself found different yields of hydrogen in the oxidation of arsenite salts reduction of ceric ions reduction of chromate ions and oxidation of ferrous sulphate. Therefore the production of H2 and H202 in hot spot regions is only a part of the observed radiochemical action and in many cases a very small part so that we are still confronted with the following problem why most of the chemical effects of the radiations are oxidation processes.Dr. Haissinsky and myself pointed out in 1949,2* what we called " the assymmetry between the action of OH and H radicals ". In 1950 we showed that even the well-established cases where reductions occurred to an important extent could be attributed to OH radicals because the redox potentials of such systems (Ce4+/Ce3+ Crs+/Cr3- etc.) had high values. Prof. Dainton and Dr. Collinson haven given a more precise definition of this concept for comparing the effects of radiation with the redox potential of the involved solutes.21 Comparing the well-known cases of total oxidation total reduction or a steady state between oxidized and reduced forms and their redox potentials the observed chemical effects could be explained by assuming that the irradiated water had itself a redox potential in the region of 0.8 to 1.0 V.Whatever is the real signi- ficance of this equivalent redox potential of irradiated water the value given above well illustrates how much more the irradiated water functions as an oxidizing than a reducing species. Equal amounts of reducing and oxidizing radicals cannot explain such a situation. With Dr. Allen's hot spots in the regions of hizh ion density redusing 20 Haissinsky and Lefort Ccmpt. rend. 1950 230 1 156. 21 Dainton and Collinson AWZ. Rev. Physic. Chem. 1951 2 99. 66 irradiation time (min) 16.5 33 16.5 66 16.5 observation formation after several hours format ion GENERAL DISCUSSION OH + OH ~ H202 H2 f OH = H + H20.H atoms disappear because of the combinations H + H = Hz while the cor- responding reaction maintains the concentration of the oxidizing species. With x-rays we could attribute a large part of the observed chemical eKects to the action of hydrogen peroxide (in the reduction of Ce4-’ and CrO4- and oxidation of As3+ and Fez+). However it is well known that the results are often different with electrons pro- duced from X- and y-rays because regions of high ion density diminish in number and in space as the energy of the electron increases. ‘Therefore the greater part of the chemical effects occurs in the rest of the solution where the distribution is more uniform. Dr. Haissinsky in his contribution showed that even where one might expect an action by H atoms in the solute because the free energy change for it is greater than for the oxidation by OH radicals yet the OH reaction occurs.We have another example in Dr. Amphlett’s paper; in spite of - 1 G being 65 kcal for the reduction of ferric ions by H atoms the reaction occurring is the oxidation. On the other hand when reduction by OH becomes possible (ferric ions complexed with o-phenanthroline EO = 1.1 V) then the reduction is complete and the yield is high (G = 12) although the free energy change of the reaction with OH is about 18 kcal that is to say lower than for redclction of ferric ions by H atoms (65 kcal). So even in the bulk of the solution where we have uniform distribution of radicals the concentration of the oxidizing species seems to be much greater than that oi‘ the reducing ones.Several hypothesis have been proposed to account for these considerations H2+ (Weiss) primary formation of H2 + OH from the ions H2O (Haissinsky and Magat) etc. As yet none of them appears to be sufficiently confirmed by experiments. However we think that the classical explanation of balanced production of OH and H radicals in irradiated water is no longer satisfactory even when corrected for non-uniform distribution as proposed by Dr. Allen. Dr. M. Haissinsky (Paris) said The constancy of the H2 yield claimed by Dr. Allen seems to be somewhat surprising. With Dr. Lefort we have examined several redox systems and it was the peculiar radiochemical behaviour of each of them that struck us no one being similar to the other.Following Allen’s sugges- tion equivalent and constant amounts of H2 and H202 are originated nearly in all aqueous solutions. But the measured yield of H202 is surely variable from one case to another and one must admit that this variation is due at least partially to back reactions. One can then not understand why H2 is not sensitive to these reactions the importance of which depends upon competition conditions. Allen was one of the first lo underline the importance of the reaction Moreover many redox systems are known for which it is necessary to admit a consumption of variable amounts of OH radicals or of an excess of these relative to H atoms ; an equivalent quantity of the latter arising from the reaction allowed by Allen H20 = H + OH combines then to form Ha molecules and the measured yield of the gas could not be the same for various solutions (see for instance the oxidation of FeS04).As these cases are far from being exceptional it would be useful if Dr. Allen would kindly clarify his position. My opinion is on the other hand that the actual experimental data are as yet insufficient and too conflicting to justify gencralizations and over-simplifications a tendency which is now too marked in radiation chemistry. Dr. A. 0. Allen (Brookknscn ATutionul Lob. N. Y.) (cornnzunicated) I agree with Prof. Haissinsky that too many generalizations have been offered in radiation chemistry and that we must avoid oversimplification. We do however find empirically that the same yield of hydrogen gas is produced from a wide im-iety 123 GENERAL DISCUSSION 22 Notre Dame June 11-1 3 1952.124 of dilute solutions provided the same radiation source is used. We have demon- strated this for solutions of KBr KZ HCI CuSO4 Ce(S04)2 H202 0 2 and oxygenated FeS04 while Fricke and Hart had previously demonstrated it also for solutions of nitrite arsenite selenite and ferrocyanide. The constancy of the hydrogen yield refers only to the initial portion of the irradiations and back reactions may reduce the yield after the hydrogen concentration has built up to some extent; some of the low yields observed by Mr. Rigg and Dr. Weiss may possibly be ascribed to this cause. Aside from the hydrogen producing reaction the free radicals produce a variety of reactions in the different solutions as Prof.Haissinsky says and it is precisely this variety which gives rise to uncertainty in the value of the free radical yield GR. Whether the hydrogen is accompanied by an equivalent amount of hydrogen peroxide or by a greater or smaller amount or whether the primary molecular products may include oxygen as well as hydrogen peroxide are questions not yet firmly settled. The interesting work mentioned by Dr. Sutton and Prof. Dainton if confirmed would indicate that hydrogen peroxide is formed in larger amounts than hydrogen. This is the reverse of the effect expected on the theory of Haissinsky and Magat for the formation of molec- ular hydrogen. If the quantities of molecular hydrogen and hydrogen peroxide formed primarily are widely different it is somewhat difficult to explain the fact that hydrogen and oxygen are the main products from irradiation of dilute neutral bromide or iodide solutions with peroxide appearing only to the amount of 17 % of the hydrogen.If excess hydrogen atoms were produced in this system as suggested by Sutton and Dainton one would expect them to reduce the oxygen to hydrogen peroxide and keep the oxygen/hydrogen peroxide ratio in the products below the high observed value. Prof. Milton Burton (Notre Dame University Indiana) said We should avoid use of words which may prejudice our ideas of mechanism. The idea of “ ion clusters ” developed by S. C . Lind now has a very particular meaning in radiation chemistry.We have seen from the paper by Prof. Massey that such ion clusters have physical reality. In a paper which Magee and I will present at the forth- coming symposium on Electron Transfer and Isotopic Reactions,22 we show that ion clusters have a significant but not necessarily exclusive role in radiation chemistry. Thus it is particularly important to avoid use of words which can be confused with ion clusters and in turn confuse interpretation. In Dr. Allen’s paper he has been careful to avoid confusion in speaking of “ formation of ions in clusters ” but may I suggest that the word “ spur ” be used in the future in this sense rather than cluster where spur signifies merely a region of high concentration of ions or excited molecules disposed locally around a primary track and that we reserve the word cluster for the now classical concept of Lind.The term “ hot spot ” has also been employed in the same sense as we use spur but it is evident that hot spot also prejudices our thinking of the nature of both primary and secondary processes. Dr. A. G. Maddock (Cambridge University) said With reference to Prof. M. Burton’s paper I should like to draw attention to a novel form of protection that may become important in aromatic systems. Yesterday’s discussion revealed some difference of opinion as to the magnitude of the photon emission in the ultra-violet arising from the irradiation of aqueous systems with penetrating ionizing radiations. Although such photons can be detected with most if not all liquids for instance by means of a coincidence arrangement with two photo- multiplier tubes I am very doubtful if they can play any important part in aqueous systems.However I find the magnitude of the photon emission from other liquid systems to range over several orders of magnitude. The effect is largest in those systems that have been employed as liquid scintillation detectors for y radiation-there is no doubt that the emission from such systems is much greater than that due to the Cerenkov radiation alone. GENERAL DISCUSSION 125 I have examined the photon emission from some hundred such systems. Most of these consist of very dilute solutions of a fluorescent solute in an aromatic solvent. I find that strong photon emission is excited in systems fulfilling four conditions the solvent must be transparent to the fluorescent radiation of the solute quenching centres such as oxygen must be absent and the solvent and solute must contain extensive systems of conjugated double bonds.The magni- tude of the photon emission is such that there can be no doubt that the energy of excitation of the solvent molecules can be transferred across a chain of solvent molecules to the solute which then radiates a photon. This mechanism of energy dissipation closely similar to the case described by Prof. M. Burton has been examined theoretically by Forster.23 It might be expected that the process would lead to a measurable reduction in radiation decomposition of the solvent or protection. A simple test of this hypothesis would be to compare G(H2) and G(C2H2) for benzene with the same quantities determined for an 0.2 % solution of p-terphenyl in benzene-a solution which exhibits more than 100 times the photon emission of pure benzene.My remarks have so far been confined to two component systems but it should not be supposed that strong scintillations are not found in any pure liquids. Early last year I found that the purest cyclo-octatetraene I could prepare (m.p. - 5 . K ) gave large scintillation pulses nor could these be diminished by chromatographic purification of the compound. I think that careful attention should be paid to the part played by such energy transfers in irradiated systems particularly when aromatic or unsaturated com- pounds are concerned. Dr.E. Collinson (Leeds University) (communicated) In the work reported by Prevost Bernas et aZ. no direct dosimetric measurements were made on the y-irradia- tions. The authors themselves state that a theoretical evaluation of the dose rates was difficult and this raises the query as to whether the rates of polymeriza- tion in all the cases quoted in their fig. 1 were established as dependent on the square root of the dose rate. The importance of this lies in the fact that the results deduced from the polymerization experiments depend on the existence of a mutual termination mechanism When y-ray dose rates are varied by adding or removing sections of a source with otherwise unchanged geometry it is almost certain that the dose rates will not be proportional to the strength of each source employed.Owing to self screening and the greater deviation from a point source the dose rates corresponding to the larger sources will be proportionately lower than those corresponding to the smaller sources. This would mean that results giving an apparent dependence of rate on (dose rate)* would in fact represent a dependence on (dose rate)x where x > $. Mr. B. Coleby (King’s CoZZege Newcastle) said When the effects of X-rays upon steroid compounds were first investigated by Keller and Weiss sterols were converted into the sodium salts of their acid succinyl esters so that the effects of X-irradiation in aqueous systems could be investigated. Later it was ascertained that the same products could be obtained by using 85 % acetic acid or methanol as solvents and these solvents were then used exclusively.In order to gain some idea of the primary reactions in these solvents so that a better idea of the mechanism of the attack of the steroids could be obtained the effects of irradiating diphenylpicrylhydrazyl (DPPH) in these solvents was investigated. As a preliminary step we studied the effects of the presence and absence of moisture and oxygen upon the reaction induced by y-rays. Starting with a solu- tion of DPPH in chloroform at an initial concentration of 10-7 mole/ml the decrease in optical density at a wavelength of 5,300 8 was measured after exposure to an 800 mg source of radium. With dry chloroform in the presence of air the decrease in optical density was linearly proportional to the time of irradiation 23 Forster Naturwiss 1946 33 166; Ann.Physik 1948 2 55. GENERAL DISCUSSION 126 until about 70 % of the initial DPPH had been consumed (irradiation was not continued beyond this point). When a similar solution was irradiated in the presence of water (saturated solution) and in the absence of oxygen (the irradiation vessel was evacuated and then allowed to fill with purified nitrogen) the linear rela- tionship was again obtained and the slope of the line was not altered. It was definitely established that water and oxygen had no appreciable effect upon the rate of the reaction. The great sensitiF ity of the reaction and the ease of measuring the rate of consumption of DPPH colorimetrically are properties which may render this radical suitable for use in dosimetry work.To gain some idea of the primary reactive species produced by irradiation we attempted to isolate the products which they form with DPPH by means of chromatographic separations of the irradiation products on alumina. DPPH in chloroform or methanol solution was irradiated with X-rays (200 kV 15 mA) so that only a small proportion of the DPPH was consumed. The solvent was then of elution chromatography. Elution with petrol + benzene mixtures gave the un- removed under reduced pressure and the residual solids were separated by means changed DPPH followed by ax-diphenyl-P-picryl-hydrazine which was character- ized by its melting point mixed melting point and its ultra-violet absorption spectrum.This is the only irradiation product which we have been able to characterize as yet. The hydrazine and the unchanged DPPH usually account for 95-100 % of the starting material. There is evidence to indicate that some at least of the diphenylpicryl hydrazine is formed by the decomposition of reaction products on the alumina column since e.g. the hydrazine was also isolated during the hromatographic separation of the products obtained by irradiating a solution of DPPH in purificd carbon tetrachloride. When zid-washed alumina was used for this separation oil products were isolated which had characteristic ultra-violet absorption spectra but which decomposed to the hydrazine on standing (even in a desiccator). It is doubtful whether the reaction products are sufficiently stable to pcrmit of the application of the usual chromatographic procedures.Further work is proceed- ing and will be published in due course. Mr. J. Wright (A. E. R. E. Hcn-wll) (cori~nrriiiic'atcrl) In discussing alternative mechanisms of radiation-induced decomposition of benzene based entirely on a study of the gaseous products it is important to ensure that these products are representative of the total decomposition. In work on pile irradiation of solid naphthalene and anthracene in vacuum carried out at Harw ell in collaboration with Miss E. M. Dresel we have found gas (mainly hydrogen) yields which indi- cated much smaller percentage decomposition than that found on examining the irradiated solid. Thus in an irradiation of antliracene in which 0.14 :< of the hydrogen atoms appeared as Hz gas (i.e.1.4 of the anthracene molecules affected if one H atom was derived from each molecule) 8 of the irradiated solid was insoluble in petroleum ether and separation of the soluble fraction on an alumina column gave several bands in addition to that containing unchanged anthracene. In such cases a study of gaseous products alone may be misleading and the mechanism of decomposition cannat usefully be discussed until the major solid and liquid products are identified and a rough material balance established. These observations have a bearing on the much lower G values for gas production in this paper compared with the G values for free radical production given by Dr. Magat in the next paper.Prof. F. S. Dainton (Leeds Uriivcrsiry) said The results described by Dr. Magat and his collaborators raise several interesting questions. In using " radical- catching " agents such as DPPH for determining GR it IS necessary as the authors point out to ensure that all the radicals are caught and therefore to compare radiation yields only under conditions for which it has been previously established that the yields do not depend on either the dose rate or the concentration. How- ever much may also be learnt from the concentration dependent region of the GENERAL DISCUSSION 24 Dainton Research. 1948 1 491. 25 Chapiro J. Clliin. Phys. 1950 47 747. 26 Mayo and U'alling Chem. Rev. 1950 46 277. 27 Walling Briggs Cummings and Mayo J . Amer. Chern. SOC.1950 72 45. 127 rate-concentration graphs. If the mechanism is as summarized by the authors then as the solute concentration is reduced the dose rate exponent will decrease from unity to a limiting value of 0.5 provided the radicals formed in different tracks are freely interacting. The data presented in fig. 1 indicate that the limiting dose rate exponent is certainiy not unity. Three possible explanations of this fact may be offered. Firstly that a substance (unknown) is present which can compete with DPPH for the radicals. Secondly that most of the radicals which recombine have originated in the same track i.e. non-uniform kinetics should be applied. Thirdly that relatively few free radicals are formed and the major part of the energy ab- sorbed is used to excite solvent molecules to a degree which is inadequate to ensure dissociation but which suffices to ensure reaction with DPPH and that such excited molecules may also undergo first order deactivation.This last sug- gestion would also explain the surprisingly high degree of energy utilization observed in some of the systems summarized in table 1. In using the polymerization method a much more extensive investigation is necessary. One must be sure that the kinetics of the polymerization are given by eqn. (HI) p. 751 ref. (2) where GRS is the radical yield for the solvent and GFi is the radical yield for the monomer. In view of the fact that in some cases the values given in column 3 of table 1 of the relative numbers of free radicals differ considerably either from the values cited previously or from the values given by the DPPH method can the authors say which of the two methods they regard as the more reliable and whether eqn.(111) was fully verified for each system for which eqn. (111) should be applicable ? Dr. W. G. Barb (Cowtardds Ltd. Maidenhead) said Prof. Dainton has asked whether some of the polymerization measured by Dr. Magat and his collaborators might perhaps be due to ionic rather than free-radical initiation. This point could be tested experimentally by the use of a mixture of two suitable monomers in place of styrene alone. It is well established that in certain monomer mixtures free-radical and ionic initiators lead to different reaction products ; 26 thus e.g. Walling et a1.27 have shown that mixtures of styrene and methyl methacrylate can give a copolymer nearly pure polystyrene or nearly pure polymethyl methacrylate depending on whether a free radical carbonium or carbanion mechanism is in- volved.A suitable separation and analysis of the product will therefore indicate the contribution of these various processes. Dr. W. Wild (A.E.R.E. Harwell) said There seems little prospect of increasing our understanding of the radiolysis of organic compounds by study of the gaseous and the very complicated condensed products. The programme outlined by M. Magat is a very valuable and encouraging contribution to this difficult field. If one compares this problem with that of carbonyl photochemistry it seems reasonable to assume that progress will require much effort.With these thoughts in mind we began quite independently an examination of the possibilities and validity of the use of diphenylpicryl hydrazyl (DPPH) as a means of capturing counting and characterizing radicals liberated by ioni7ing radiation in organic solvents. Spectrophotometric observations of solutions of DPPH after irradiation lead one to believe that the only process occurring is addition of other radicals to give related hydralines. Subsequent chromatography on alumina and lime columns GENERAL DISCUSSION 28 Bawn and Mellish Tvnns. Faradqy Soc. 1951 47 1216. 128 modifies this conclusion to some degree. After irradiating dilute solutions of DPPH in benzene six bands are found on such columns four coloured and two fluorescent of which one coloured band constitutes 50-60 % of the recoverable products and has a spectrum resembling that of the parent hydrazjne of DPPH.Some of the other bands on elution give spectra showing no resemblance to those of the hydrazyl or hydrazine. As these two substances appear to be stable on such columns and as the total weight of material recoverable from the columns is con- siderably less than that of the DPPH used it is suggested provisionally that some of the DPPH is actually disrupted during irradiation of the solution. Our experi- ence renders it unlikely that any substance of the type R-DPPH would pass through the column and be lost but smaller fragments might pass through. It seems unlikely that such a fragmentation of the DPPH skeleton is the result of direct action of fast electrons at the concentrations employed.Moreover DPPH has been used under very mild conditions to study the rate of radical production from polymerization initiators undergoing thermal decomposition satisfactory proofs of its reliability having been given.28 In our opinion this is no adequate guarantee that the method is reliable in all cases. Two possibilities are conceivable. The severe conditions in the neighbourhood of electron tracks with the attendant possibility of “ hot ” radicals being present may mean that a radical may add on to DPPH and break the weak N-N bond simultaneously in which case three radicals may be consumed to give stable end products. On the other hand another process may occur in some solvents.Benzene for example forms a very stable 1 1 compound with DPPH. Indeed recrystallization from benzene gives a particularly pure sample provided adequate care is taken in the final removal of this solvent. If electronic excitation occurs when ionizing radiations traverse aromatic solvents it may happen that by resonance transfer this energy may be transferred into the solvation shell of the diphenylpicryl hydrazyl and cause its disruption into fragments consuming more than one of the radicals arising from solvent ionization. A direct test of this possibility is to be undertaken shortly. For these reasons it is not yet possible to say that DPPH counts radicals more accurately than by a factor of (say) two or three. This may be a contributory cause to the discrepancies noted by Dr.Magat between the results of the two methods employed. The general qualitative trends of radical production from different solvents are not questioned as they seem consistent with other evidence. It is a relatively easy task to accumulate tens of milligrams of column fractions when using a 0.1 /LA beam of 1.5 MeV electrons. The examination of the principal fraction from benzene by elementary microanalysis and infra-red spectra has revealed a factor which may confuse the identification by both methods of the radicals produced in the solvents. The products are very difficult to purify by recrystallization and our analyses in this one case suggest that some solvents are tenaciously adherent to this type of hydrazine and complete removal is absolutely essential to success.Literature evidence supports this view in the case of diphenyl- picryl hydrazine. The composition of the fraction mentioned suggests the addition of a CsH fragment. From a microanalytical viewpoint the large molecular weight of DPPH is a definite disadvantage of the DPPH method for identifying the usual organic radicals. Experiments are being pursued with simpler molecules and the possible use of some stable molecules should not be ignored. It is of interest that the radical yield from benzene using 1.5-1.8 MeV electrons as deduced from the rate of disappearance of DPPH colour is of the order G = 0.8 for solutions prepared in vacuo and somewhat lower (0.6-0.7) for aerated solutions.These yields are markedly lower than those quoted in table 1 of Magat’s paper. A low value of about 0.5 has been suggested for styrene resulting from molecular weight measurements on polystyrene produced by cobalt y-rays coupled with the assumption that termination proceeds by mutual combination of two growing 129 chains.29 Our yields in carbon tetrachloride are much larger than in benzene but we are not satisfied that the results in halogenated solvents are not confused by purely thermal reactions with DPPH. Despite the most rigorous purification and drying of carbon tetrachloride we have never prepared in vacuo a stable solution of DPPH in this solvent although air-saturated solutions in the same sample have survived many months of dark storage without loss of colour.The role of oxygen clearly requires further study. Dr. A. Chapiro (Paris) (partly communicated) In our polymerization experiments the intensity of the 7-radiation was undoubtedly not quite proportional to the intensity of the source as expressed in curies for reasons indicated by Dr. Collinson. However recent calibrations with small ionization chambers have shown that the errors probably do not exceed 10 %. This degree of inaccuracy may explain the observed small deviations from the theoretical I* plot but are insufficient to throw doubt on the square-root relation between the yield and the intensity. We in- tend howe\er to repeat these experiments with a stronger source that will soon be available and which will permit a greater variation in intensity.We have also observed as Mr. Coleby has done that DPPH solutions in chloro- form give identical results in presence of air and in vacuum. But this has not yet been verified for other solvents particulai ly alcohols. Prof. Dainton’s remark that the limiting dose exponent in the rate against DPPH concentration curves must vary from 0.5 to 1 when the DPPH concentra- tion increases is completely verified by our experiments (see fig. 1 of my discussion remarks). The slope of the initial linear part of the curves is proportional to 1 4 within experimental errors. In the concentration-independent part the rate is proportional to I as can be seen from the following figures rate/Z (rjmin) GENERAL DISCUSSION rate 108 moles/cm3 min 1.34 I 69.0 1.94 1-91 0.525 0.315 27.5 17.3 1.82 The last value is somewhat less reliable since the experimental errors increase for very slow reactions.Dr. M. Magat (Paris) (partly communicated) Prof. Dainton raised several important questions that I shall try to answer in the order they were presented in so far as they have not yet been answered by Dr. Chapiro. It is of course impossible to exclude a priori the possibility of DPPH reacting not only with free radicals but also with excited molecules. This is a limitation of which we are perfectly well aware. However on one hand there is a rather good agreement between the polymerization results for styrene and the DPPH results particularly if one takes into account Dr. Valentine’s remark.On the other hand we have tried to induce a reaction between methanol and DPPH by irradiating the solution with infra-red light of 3 to 4 p in order to excite the OH vibrations. This attempt has been so far unsuccessful. Concerning the polymerization method the equation quoted by Dainton was verified to the following extent. + 30 % methanol mixture and for styrene + benzene cyclohexane and ether (i) The intensity dependence was verified for pure styrene for a 70 % styrene mixtures. I agree that a verification in each individual case would be desirable and we intend to test several more cases in the near future. (ii) The dependence on the amount of solvent added was verified in the con- centration range from 10 % to 80 % for aromatic hydrocarbons and in a some- what smaller range for several other solvents.For polymer precipitants and poor solvents the equation breaks down at lower concentrations of the diluent. The probable reasons for this were amply 29 BNL 141 (T-27) (1951) (Brookhaven National Laboratory Report). 130 GENERAL DISCUSSION discussed by Chapiro 30 and 1 need not repeat them at length. It may be noted that in very dilute solutions of styrene in methanol the intensity exponent also increases to a value between 0.5 and 1. (iii) The difference between the figures of column 3 table 1 and the figures given previously 30 arises from the use of different units. The previous figures were expressed in relative numbers of free radicals per mole while in table 1 we refer to the relative numbers per cm3.(iv) The DPPH method is more direct and is to be preferred whenever it can be used with the reservations made under (i). Unfortunately it is difficult to use with compounds containing double bonds or very motile atoms because of a non-negligible thermal reaction. DPPH is very sensitive to traces of peroxides sometimes difficult to eliminate. The polymerization method is rather doubtful if the diluent is a poor solvent or a precipitant. Hence one or the other method is to be preferred according to the substance investigated. In connection with Dr. Wild’s remarks one could wonder what would be the repercussion on the counting of the free radicals by the suggestion that the DPPH radical is broken up by the reaction with another free radical.Two cases may occur (i) The break is such that a free radical fragment and a stable molecule are formed. The free radical fragment will react with the radiation-produced free radicaIs and the count will remain correct. (ii) The break leads to the formation of a free radical and a bi-radical. Then each DPPH molecule will consume three radiation-produced free radicals ; the real number of radiation-produced radicals will hence be even larger than the one we indicate. We agree with Dr. Wjld that far more work ought to be done along the lines indicated by him before the problem of the DPPH method will be completely solved. I would like to add that since the Discussion we have performed a calibration of our installation with the 275 mc Ra source both using small ionization chambers and by the FeS04 method.According to this calibration all the values listed in columns 4 and 5 (GR and Y ) of our table 1 are to be multiplied by a factor 0.82 which makes the energy yield values more reasonable. The absolute values given in table 2 correspond to an intensity of 1.13 r/min. However there remains a real difference between the GR value of benzene as detcrmined by Dr. Wild with fast electrons (GR 0.6-0.7) and by our group with y-rays (GR = 1.47). Two explanations may be offered either with the intensities used by Dr. Wild the concentration independent region is difficult to reach while still avoiding direct effects or else there exists a difference of the GR values for y-rays and for electrons as was suggested at this Discussion by at least two authors viz.Chapiro and by Alper Ebert Gray Lefort Sutton and Dainton. Prof. Milton Burton (Nofre D a m University Indinno) said Table 1 of the paper by Dr. Magat and his collaborators gives G (free radicals) from benzene N 0.8 contrasted with G(H2) = 0.036 and G(C2H2) = 0-02 in the report by Gordon and myself. Our work indicates that not all of the gaseous product is of free radical origin; we can confidently write for the free radicals involved in such formation G < 0.05. Work by Manion and myself31 indicates that for the re- action benzene molecules + polymer G _N 0.8. Unreported work by Dr. Burr some years ago showed that the polymer aggregates contain on the average 5 or 6 molecules of benrene.Assuming one radical to start and one radical to termina?e that for total free radicals produced G > 0.35. This result is in much better the polymer it appears that for radicals involved in such formation G N 0.3 or agreement with Dr. Magat’s cri 0.8 than I thought during the actual discussion at 30 Chapiro J. Chim. Phys. 1950 47 747 764. 31 Manion and Burton J. Physic. Chem. (May) 1952. 131 1951 47,462 and 591. GENERAL DISCUSSION the meeting. However the discrepancy is real. It may be connected with the limited range of validity cf Dr. Magat's methods indicated by him in his discussion of table 1. In particular there is a special dificulty of radiosensitization when a foreign substance is introduced into benzene to determine the free radical yield from benzene itself.One must be particularly certain that the added substance does not have lower excitation or ionization potential than benzene for if it does energy may transfer to it from the benzene. If further that added substance should be more labile than the benzene (a highly probable situation) it may happen that its presence greatly increases the free radical yield out of all proportion to the concentration of the additive. 1010. Dr. L. Valentine (Birminghanz Ujiiversity) (corninmicafed) In the interesting paper by Prof. Magat and his collaborators there is a iiide scatter in the estimates (table 2) of the absolute numbers of free radicals produced. Account has not been taken however of the fact that the rate constants found by Bamford and Dewar (B and D) and Melville and Valentine (M and V) were calculated assuming that the termination reaction was disproportionation m hilst Mttheson et al.assumed combination. This automatically introduces a factor of 4 2 in the values of kl,,'kl& i.e. a factor of 2 in the estimates of the number of free radicals formed. Evidence is accumulating that the termination step in the polymeriiation of styrene is combination and if the results are recalculated on this basis taking the overall energy of activation to be 6.2 kcal/mole the \slues of kp/A,i become 0.0067 (M and V) and 0-0058 (B and D). The number of free radicals produced per cm3 per sec then becomes 0.9 x 1010 (B and D) and 0.7 x 1010 (M and V) compared with values of 1.7 1010 (Matheson ef a/.) and 1.3 x 1010 (using DPPH).There is much better agreeinent amongst these results than amongst the ones quoted in the paper. The value quoted as being determined from a combination of molec- ular weight and polymerization rate data must also be doubled if disproportiona- tion has been assumed raising it to 2.2 Dr. Philip George (Canikridge University) said The use of ferrous sulphate solutions in dosimetry and the investigation of the reaction mechanism which this has necessitated make it extremely probable that the problem as to whether similar reactions cam occur with the haernoproteins \?/ill attract attention in the very near future since these compounds contain ferrouq or ferric iron in a co-ordination coniplex and play an essential role in metabolic processes.Many of the free radical reactions postulated in ferrous salt dosimetry are identical viith those believed to occur in the reaction of iron salts with hydrogen peroxide,3? and in a recent reL iew 33 the possibility of these reactions occurring with typical haemoproteins has been considered on the basis that the probable influence of this co-ordination on the iron is to lower the ionization potential of its ferrous form in aqueous solution. This leads to the conclusion that reactions of the ferrous form will be some 5 kcal more exothermic and those of the ferric form more endo- thermic by the same amount with the result that all the ferrous ion reactions are equally or more favoured energetically with the ferrous form of haemoproteins.Ionic ferric iron reacting with 0 2 - is exothermic to the extent of about 19 kcal and so a similar reaction occurring with haemoproteins is still likely to be exo- thermic to about 14 kcal and very rapid its reaction with H202 via the anion 02H- is on the other hand endothermic by about 28 kcal and a similar reaction with the ferric form of haemoproteins will be some 5 kcal more endothermic. Furthermore there is no reason to expect such a high frequency factor of the order 1024 as estimated from the ferric iron data for haemoproteins and so on two accounts this reaction is likely to be very slow indeed. Calculation 33 shows 32 Barb Raxendale George and Hargrave Nature 1949 163 6 ; Trans. Faraday Soc. 33 George Ahmices in Catalysis vol. 4 (Academic Press New York 1952) p.367. GENERAL DISCUSSION 36 George Nature 1952 169 612. 132 the bimolecular velocity constant at pH 7.0 and 20" C would probably be of the order 10-19 mole-1 sec- 1. This conclusion is of great significance in deciding a reaction mechanism for the haemoprotein enzymes peroxidase and catalase which nevertheless in their ferric forms react extremely rapidly with hydrogen peroxide. Intermediate com- pounds are formed in these systems and also with the very similar haemoproteins metmyoglobin and methaemoglobin which have long been regarded as enzyme- substrate complexes.34 Recent researches,35? 36 however have shown that certain of these compounds are not of this type but contain the iron in an effective quadrivalent oxidation state for they are reducible to the ferric state in a one equivalent step.The " secondary " compounds of the peroxidases and the metmyoglobin compound show this property. Furthermore hydrogen peroxide or the structurally similar alkyl hydroperoxides are not unique in yielding this type of intermediate compound. With horseradish peroxidase hypochlorous acid hypobromous acid bromate periodate chlorite chlorine dioxide ozone and perdisulphate with silver ions all yield the second intermediate compound. Metmyoglobin gives its intermediate compound with chlorite chloriridate ions and molybdicyanide ions and in the last two cases the reaction is much faster than with hydrogen peroxide.6 The effective redox potential of the metmyoglobin compound is about 1.0 V at pH 6.8 (European convention) and about 0.7 V at pH 11.0 showing that a hydrogen atom is involved in the redox reaction.37 This compound and the secondary com- pound of peroxidase are alike in reacting with reducing agents such as the ferro- cyanide ion more rapidly the higher the hydrogen ion concentration which suggests that a hydrogen atom is similarly involved in the redox reaction of the peroxidase compound .In the radiation chemistry of systems where these haemoproteins are present the formation and reactions of such intermediate compounds should be taken into account along with the favoured ferrous and ferric ion type of radical reactions not only because hydrogen peroxide is formed by irradiation but because these compounds may well be formed by oxidizing entities present in the system before the production of the hydrogen peroxide molecule as such.34 Chance Advances in Enzymology vol. 12 (Interscience New York 1951) ; The Enzymes (Academic Press New York 1951). 35 George and Irvine Nature 1951 168 184 ; Biuchem. J . (in press). 37 George and Trvine unpublished results. 110 G E N E R A L DISCUSSION GENERAL DISCUSSION Dr. N. Miller (Edinburgh Uuiiversity) (partly conimuniccrted) Mr. Wilkinson and I have now completed our studies of the yield for ferrous sulphate oxidation by alpha-particles in aerated 0.8 N sulphuric acid solution. The absolute yield was evaluated by comparison of the chemical change due to the alpha-particles from a given source of polonium with the ionization current due to the same source in an argon-filled ionization chamber.The results are therefore directly dependent on the value assumed for the quantity W (energy released per ion pair formed) for alpha-particles in argon. Values of this quantity in the literature range from 24.9 to 28.3 eV our corresponding G values for ferrous sulphate oxida-tion by alpha-particles being 6.7 to 5.9. We have found a constant yield to be maintained over a 24-fold range of dose rate and independently of the initial ferrous ion concentration from 10-3 M upwards. An interesting feature which emerges from this work a detailed description of which is being prepared for publication is that the same yield is observed if a sheet of mica of about 1.5 mg/ cm2 mass thickness is interposed between the source and the liquid during irradia-tion such that the average linear ion density of the radiation is considerably increased.In this case the ionization measurements are conducted of course, through a similar mica sheet. It is therefore apparent that this yield is character-istic of alpha-particle radiation or at any rate of alpha-radiation of energy not greatly different from that observed in natural radioactive transitions. By using solutions containing dissolved polonium we have also studied the effect of dissolved oxygen on the alpha-ray induced oxidation of ferrous ions in 0-8 N sulphuric acid solutions. The yield is reduced by a factor 1.7 by de-aeration, this being the case at initial ferrous ion concentrations ranging from 5 x 10-4 to 8 x 10-3 M.Preliminary experiments also indicate that the chemical effects of polonium alpha-rays at a dose rate of 1775 ergs/g min and of 50 kV peak X-rays at a dose rate of 50,600 ergs/g min when administered simultaneously are strictly additive. With regard to the discrepancy between Dr. Hochanadel’s value for the y-ray yield for ferrous sulphate oxidation and those obtained by other workers little more can usefully be said at the moment. A decision can only be reached after more experimental work has been done. The possibility that solutions made up in one laboratory may behave differently from those prepared elsewhere due to traces of dissolved impurities has however been eliminated. This Discussion has provided us with an opportunity for a direct comparison between such solutions made up in different laboratories in that we have just been visited in Edinburgh by Dr.A. 0. Allen and Dr. T. J. Hardwick who brought samples of ferrous sulphate solutions with them. The following table of results obtained by Dr. RELATIVE YIELD VALUES FOR FERROUS SULPHATE OXIDATION IN AFRATED 0.8 N SULPHURIC ACID SOLUTION 50 k V peak X-rays: C060 ~ ~ - r ~ y ~ _ _ _ ___ ______ NoC1- i 10 3 M C1-Brookhaven solution 0.97 1 -02 0.99 Chalk River solution 0-98 0.98 0.98 Edinburgh solution 1 -00 1 -00 H. A. Dewhurst in our laboratory shows that these solutions and those prepared by us behaved very similarly under irradiation with both 50 kV peak X-rays and Co60 y-rays. Each of the figures in the table represents the average of at least 6 separate irradiations.These results providing a useful corollary to those in the experimental and results section of our paper show that the yield for ferrous sulphate oxidation is indeed reproducible and characteristic of the system. The fact that none of these solutions showed any difference in yield when made 10-3 GENERAL DISCUSSION 111 in chloride ion also shows that no appreciable amounts of organic impurities were present in any case. I am glad that the explanation of the apparent post-irradiation effect reported by Dr. Wild has now been agreed upon. The effect of temperature on the extinc-tion coefficient of ferric ion in 0.8 N sulphuric acid solution was first noticed by Dr. H. A. Dewhurst in our laboratory some months ago and we are in agreement with Dr.Wild as to its magnitude. We also endorse Dr. Wild’s words of warning about the existence of this effect in regard to the use of this mode of analysis in dosimetry. The only genuine after-efTect my colleagues and I have detected in this system in over 5000 irradiations has been that due to the slowness of the Fenton reaction, which is only apparent in practice at initial ferrous concentrations at or below the lower limit of the region of concentration independence. The apparent fall-off in yield observed when the initial ferrous ion concentration is below this limit is indeed largely due to the slowness of this reaction. This effect will however, doubtless be dealt with in detail in future publications by Dr. Dewhurst of our laboratory and by Dr. Sutton at Leeds who have examined it independently.For the measurement of low doses the DPPH solutions studied by Dr. Chapiro do appear to have certain desirable features. Although the solutions share with the chloroform$water mixtures discussed in our paper the drawback of not being “ air-wall ” it strikes one at once that the “ yield ” in the concentration-independent region is independent of dose rate which is not normally the case in the chloro-form-+water systems. For use for military or civil defence purposes however such criteria as cheapness ready mass production and stability on storage have also to be considered while as a standard for laboratory work it is questionable whether such a method can offer a serious challenge to the aqueous systems previously disc wed.Dr. E. J. Hart (Argo~zrzc National Lab. Chicago) said I wish to point out the merits of the formic acid dosimeter which appears to have been overlooked in the survey of aqueous systems proposed for dosimetry appearing in table 1. Earlier use of this system has been described 12 but more recently 3 it has been successfully used to distinguish between the radicals (H and OH) which escape from the track and the molecular compounds hydrogen and hydrogen peroxide, formed by recombination of H and OH radicals within single tracks. In short, the use of the formic acid dosimeter enables one to distinguish between the following reactions : H20 = H + OH The ferrous sulphate and ceric sulphate systems have not been developed to a point where separate measurements of reactions (1) and (2) can be effected.This refinement in dosimetry is of importance in chemical and biological systems since hydrogen atoms and hydroxyl radicals are much more reactive than hydrogen molecules and hydrogen peroxide. Dr. W. Wild (A.E.R.E. Harwell) said Dr. Miller’s paper focuses attention on the relatively limited range of dose-rate over which present chemical dosimeters can be regarded as satisfactory. In contrast to the situation in radiobiology it seems unlikely that radiation chemistry will require extensive use of very low dose-rate dosimeters. Any requirement in this range will probably be met from current extensive work associated with civil defence. On the other hand there are many chemical problems in the organic field where insensitive analytical methods will demand the application of large doses if adequate answers are to be obtained.Moreover the large dose-rates obtainable from electron beams and big cobalt 1 Fricke and Hart J . Chem. Physics 1934 2 824. 2 Fricke Hart and Smith J. Chem. Physics 1938 6 229. 3 Hart 1951 Report ANL-4636 ; J . Physic. Chem. 1952 56 (in press) 112 GENERAL DISCUSSION sources make possible this desirable extension of the subject. At present there is no adequately tested chemical system applicable to such work. The discovery of such systems would be of great value. It is also perrinent to note the absence of chemical dosimeters usable in organic solvents. Dr. E. J. Hart (Argonne National Lab. Chicago) said The erratic behaviour in the ratio R between the aerated and evacuated ferrous sulphate systems has been the source of considerable concern to us.Organic impurities were suspected as one of the causes of the large variation in results. In accordance with This idea it was found that upon the addition of formic acid the yield of the y-ray oxidation in the presence of air was enormously increased while the yield in the air-free solutions was reduced. As a result of these findings a study was made of the mechanism of the y-ray induced ferrous sulphate + formic acid + oxygen reaction. While the detailed kinetics of this reaction are undoubtedly very complex, the existence of an oxidation chain of 30-40 units in length greatly simplified the mathematical treatment of the data. H20 = H + OH (1 a) H20 = +H2 + iH202 (W OH + HCOOH = H20 + HCOO ( 2 ) H + HCOOH = H2 + HCOO ( 3 ) HCOO + 0 2 = HO2 + C02 (4) Fe2+ + HO2 = Fe3-1 + HO2- ( 5 ) H+ + H02- = H202 ( 5 4 Fez+ + H202 = Fe3+ + OH- + OH (6) The reaction is initiated by propagated by reactions (2) (4) (3 ( 5 2 ) and (6), and terminated by Fe3+ + HCOO = Fe2+ + H+ + C02 Fe2+ + HCOO = Fe3+ + HCOO-Fez+ + OH = Fe3f + OH- (9) Fe3f I- H02 = Fez+ + Hf 4- 0 2 .(10) It was further found that ferric sulphate can be efficiently reduced by y-rays in aqueous solutions containing formic acid. The above mechanism provides a simple explanation for deviations of the ratio R from the generally accepted value of about 2.3 when the contaminant is a readily oxidized organic impurity. Prof. Milton Burton (Nutre Dame University Indiana) said The problem of high-intensity actinometry is not in so sad a state as Dr.Wild has suggested. We have merely to make more extensive use of systems in which G values are small. For example we see in the paper by Gordon and myself that benzene46 is even more resistant to radiation than benzene. We need not however confine ourselves to measurement of hydrogen and acetylene production in benzene radiolysis. The value of G (diphenyl production) is very much smaller so that production of (C6H5)2 can be studied when intensity effects are great. Furthermore we can follow a suggestion made by Dr. Maddock (see later). terphenyl in benzene gives a particularly good scintillation counter. If this physical effect involves specially effective energy transfer from benzene to terphenyl it should be reflected in G values even lower than those heretofore determined.Dr. T. J. Hardwick (N.R.C. Chalk River) said There is an upper limit to the independence of the yield on dose rate in the oxidation of ferrous sulphate in air-saturated 0.8 N sulphuric acid by X- and y-rays. This limit varies slightly with A concentration of 0. GENERAL DISCUSSION 113 the temperature and the concentration of ferrous ion factors which affect the rate of the ferrous ion + hydrogen peroxide reaction. A decrease in yield at dose rates above 4200 r/min was found on irradiating 10-3 M ferrous sulphate solution with 2000 kVp X-rays at 22" C. However as a pulsed source was used the upper limit of dose-rate independence for continuous radiation may be higher by a factor of 3 or 4.Using a pulsed 50 kVp X-ray source a decrease in yield appeared in the region of dose rate above about 200 r/min. In Ihis case the limit may depend consider-ably on the particular conditions used but where larger dose rates are used experi-ments should be carried out to show the effect of dose rate. Using a continuous source of radiation may increase the critical dose rate by a factor of 3 or 4. The yield on irradiating ceric sulphate is 0.8 N sulphuric acid shows no effect of dose rate up to the limits which were available experimentally. Using a pulsed 2000 kVp X-ray source independence of the yield on dose rate was found up to 45,000 r/min. Using a pulsed 50 kVp X-ray source no effect on the yield was observed up to 140,000 r/min. As pulsed radiation sources were used it may be estimated that for continuous radiation the upper limit of dose-rate dependence must lie beyond dose rates 3 or 4 times higher.Dr. M. Haissinsky (Laboratobe Curie Paris) (partly communicated) In con-nection with Miller and Wilkinson's results I would like to report some experi-ments carried out in our laboratory by Mrs. Anta which show the difficulties encountered in using polonium as ionizing source and also the rather complicated behaviour of ceric salts irradiated by this source. The starting point of this research was the conclusion drawn with Lefort from the results obtained on the irradiation of Ce4f by a-rays of Rn namely that the reduction is in this case mostly due to H202 produced in the primary step. On the other hand Bonet-Maury and Lefort have stated that the x-rays of Po not only do not produce H202 in acid solution but also that there is decomposition of this compound if added previously.It seems interesting therefore to study the action of these rays on ceric ions. Working under somewhat different conditions (the Po salt being homogene-ously dissolved no reagent present in the solutions for H202 detection long irradiations etc.) we have observed that the decomposition of H202 in pure sul-phuric acid is slowed down with time of irradiation and that the quantity remaining in solution tends towards a limit. On irradiating the same solution initially without H202 the latter is formed and the rate of formation also diminishes with time (the decay of Po being negligible during this time).Again a concentration limit is obtained which for the same Po intensity is the same as that obtained previously viz. 30 x 1016 molecule/cm3 for about 0.6 mc/cm3. With smaller quantities of Po or in perchloric solution the yields of H202 are even greater and the decomposition curve has here a rather complicated form. As to the irradiated ceric solutions we found that the yield of the reduction diminishes when the dosage rate increases. It depends on the nature of the anion and possibly also on the concentration. Dr. W. Wild (A.E.R.E. Harwell) (comnzunicated) At the Discussion I remarked that Dr. Miller's method of using the ferrous sulphate dosimeter by estimating the ferric ion produced by measurement of light absorption at 302 mp enabled one to take measurements very quickly after termination of an irradiation.Such measure-ments suggested that a considerable post-irradiation increase of ferric ion occurred, whose magnitude seemed to be determined by the ferric ion concentration produced by irradiation. Mr. Best has since followed up a suggestion of Dr. Miller's that the observed effect might be a pure temperature effect and this has proved to be correct. It was found that the extinction coefficient of ferric ion in 0.8 N H2SO4 at 302 mp increases by 0.72 % per "C. Published information4 suggests that this 4 Katzin and Gebert Report ANL-4457 (Argonne National Laboratory) 114 GENERAL DISCUSSION peak is a composite one due to sulphate and hydroxy complexes of ferric ion so that the large effect of temperature is not surprising.Temperature records show that the irradiations in question were carried out (in a lead castle at 285 rjminj at temperatures about 10- C lower than the spectro-photometer room. More recent work with the two instruments at the same tem-perature show the absence of any post-irradiation effect provided the ferrous ion concentration is not reduced below 2 x 10-4. If this latter condition is not observed a genuine effect is found which is of small magnitude and is presumably that already noted 5 and ascribed to the slowness of the Fez' $- H202 reaction step in the assumed mechanism. Whilst the above observations do not invalidate the use of this analytical method which is very convenient and accurate if adequate time is allowed for solutions to acquire the temperature for which the spectrophotometer is calibrated, they do show that especial care is necessary for the avoidance of unnecessary errors when the absorption cell carrier is not adaptable to thermostatic control as in this instance.Similar remarks are applicable to the use of the ferric thiocyanate method.6 Mr. J. Wright (A.E.R.E. Harwell) said In connection with the use of the benzene-water system for dosimetric purposes we have found a post-irradiation production of phenol which continues for at least 24 h after the end of irradiation and in some cases is as high as 40 % of the phenol present immediately after removal from the source. This effect has been obseried after irradiations in the pile and after irradiation uith a Co y source with solutioix initially aerated and with nitrogen equilibrated solutions using Folin's method of estimating phenol or the 4-amino antipyrelie method described by hlartin.7 The origin of this additional phenol is not known at present but it may be significant that the hydro-gen peroxide produced in the system during irradiation is destroyed during the post-irradiation period.In using the benzene-water system for dosinietric or monitoring purposes it is essential to standardize the time after irradiation at which the phenol concentration is measured and to ensure that changes in irradia-tion conditions do not affect the post-irradiation behaviour. Dr. A. 0. Allen (Brookhaven National Laboratory) said Using 2 MV volt X-rays from the electrostatic generator at Brookhaven we have compared ferrous sulphate oxidalion yields at intensities ranging from about 100 r/min to 10,000.Solutions were irradiated at different distances from the X-ray target and over a wide range of currents. By intercomparison of the results assuming that the ratio of the radiation intensities at the different distances remains constant inde-pendent of the current it could be concluded that the oxidation yield was constant within 5 % over the intensity range studied. In view of our results it appears unlikely that the discrepancy between Hochanadel's value of the yield and other values could be due to the effect of intensity on the yield. A real discrepancy appears to exist between his results 2nd others and further independent deter-minations of the absolute yield in ferrous sulphate oxidation would be most desirable.Prof. Milton Burton ( Uftivcrsity of Notre Dnmc) (co/imunic.nted) The matter of the correct G value for the ferrous sulphate actinomcter is one of great impor-tance for radiation chemists. An explanation must be found for the anomalols situation created by existence of such divergent values. However the funds-mental problem created by the ferrous sulphate actinometer has been obscured by this dispute. That problem is how to account for such high G values whether they be 15 or 20. To that question both experimentalists and theorists must address their efforts. 5 K. H. Krenz and H. A. Dewhurst private communication. 6 J. Sutton private communication. 7 Martin Anal. Chem. 1949 21 1419. Sutton this Discussion GENERAL DISCUSSION 115 Dr.A. Chapiro (Paris) said Dilute solutions of diphenylpicrylhydrazyl (DPPH) can be used for chemical actinometry of ionizing radiations. This method is very sensitive since differences of a few roentgens can be detected under suitable condi-tions. In one experiment with 190 kV X-rays using bromoform as a solvent a total dose of 35 r administered in 30 sec changed the optical density of the solution by 10 %. In addition the concentration dependence may be interpreted on simple kinetic assumptions and this makes it possible to determine in each case the best experimental conditions to be used. Let us assume that the free radicals formed in the irradiated solvent may either recombine with themselves or combine with DPPH as in the following reaction scheme :8 S .\\\-f 2R* @I R" + R" -t X kl(W2 R" -+ DPPH + R - DPPH k2(R*)(DPPH) S being the solvent R* the free radicals formed @ the radical yield of the solvent and X a stable molecule.At the steady state k1(R*)2 -'- k2(R*)(DPPH) -L @I = 0, which gives for the reaction rate ._ d(DPPH) _ _ _ ~ - V - k9(DPPH)* - 4@1kl ~ )+ - I]. (1) df 2kl - [ ( X$(DPPH)' This equation contains one constant k22,'X'l. For very concentrated solutions k$( D PP H)?' and eqn. (7) reduces to I/ -= m. The rate of disappearance of DPPH is equal to the rate of production of free radicals. For very dilute solutions eqn. ( 1 ) can be simplified to ( 3 ) I22 kit V - - (@Z)*(DPPH). The constant k?/kl* can be calculated from eqn ( I ) and from the slope of the straight line obtained by plotting V against the concentration of DPPH €or very dilute solutions.Fig. 1 shows the experimental results obtained with chloroform solutions 'irradiated with 190 kV X-rays at 3 different dose rates. The curves are derived from eqn. (I) using k2/kl+ = 66-6 and the experimental points fall close to these theoretical curves. For concentrations of DPPH higher than a given value (DPPH),i,. the re-action rate becomes independent of concentrations the constant rate @I being proportional to the dose rate. (DPPH),i,. is proportional to (@Z)t as can easily be calculated from the kinetic equations. It follows that chloroform and bromoform which are the most sensitive solvents for low dose rates cannot be used €or high dose rates since (DPPH),i,.becomes too high for practical colori-metric estimations. However table 1 of our contribution to the present Discus-sions shows that other solvents can be used which give lower radical yields. For 8 A similar interpretation of the dilution effect in water has already been proposed by This agrees with the experimental results. Dainton and Miller (Proc. Int. Congr. Pure Appl. Chern. (London 1947) 1 77) 116 GENERAL DISCUSSION instance with methyl acetate dose rates of 5000 rlmin of 37 kV X-rays have been measured. Hydrocarbons could be used for dose rates up to about 200,000 rlmin. The dependence of the chemical yield in chloroform on the type of radiation used is shown in the following table : type of radiation 37 kV X-rays 4.1 190 kV X-rays 2.41 1 MeV electrons 0.4 1 relative chemical yield Ra y-rays I The real energy absorbed per roentgen differs widely for the various radiations listed.An evaluation of the radical yield per unit energy absorbed has shown that the figures for y- and X-rays come close together. There still seems to be a difference between the yield due to fast electrons and that due to y-rays. Finally, this actinometry method is very simple. The solutions can be used in the presence of air and the analytical procedure only involves simple colorimetric measurements. FIG. 1. Mr. J. Wright (A.E.R.E. Hurwell) (partly communicated) In quantitative studies with a mixed source of neutrons and gamma radiation one of the important points to establish is whether the chemical effects of the two types of radiation are additive.The evidence on this point presented in the paper is not conclusive, but there is additional evidence from studies we have made of the irradiation of ferrous sulphate in 0.8 N H2S04 containing boric acid. In the centre of the pile, the rate of oxidation of ferrous ion is a linear function of boron concentration up to 8 x 10-2 M H3B03. The slope of this line leads on the assumption of chemical additivity to a G value for aerated solutions of 3.9 for the additional oxidation produced by the a and Li recoil particles of the 1*B(n aj 7Li reaction. This is much lower than the value of 6.0 given by Dr. Miller9 and this evidence alone suggests that there is less oxidation from the a-particles in the presence of the back-ground pile radiation than from a-particles from a pure a-source.Further evidence was obtained when the ferrous solutions containing H3B03 were irradi-ated in the thermal column of the pile where the oxidation per unit neutron dose in the absence of boron was only one-fifth of that in the centre of the pile. The relation between rate of oxidation and boron concentration appeared to be no longer linear the slope of the curve increasing as boric acid concentration increased. The G values for the additional oxidation due to a and 7Li particles were higher than the 3.9 obtained in the centre of the pile and appeared to increase to a value 9 communicated at this Discussion GENERAL DISCUSSION 117 which may approach 6 as a greater proportion of the total oxidation is derived from the additional a-particles.Details of this work which is not yet complete, will be published separately. It has not been possible so far to control the irradia-tion conditions in the thermal column as carefully as those in the centre of the pile and the spread of results has been somewhat greater. There is no doubt, however that the G values for oxidation by the cc and Li particles vary according to the amount of background pile radiation present. Dr. M. Lefort (Institut du Radium Paris) said We can confirm some of the difficulties mentioned by Dr. Wright for the dosimetry in the pile. Mr. J. Pucheault, Dr. Haissinsky and myselflo have made a great number of irradiations of boric acid solutions using the French pile ZOE in Chatillon. Like Dr.Wright we used with satisfaction a gold monitor for the determination of the flux of thermal neutrons for each sample because the values given from the control instruments were not suitable for measuring the dose received by the irradiation cell itself. These experiments were made in order to follow the formation of hydrogen per-oxide from boric acid solutions irradiated by the x-rays produced in the solution itself through the nuclear reaction B(n M). It was proposed,ll when this work was begun to use the measurement of the hydrogen peroxide formed as a chemical dosimetry of slow neutrons. Such a method could be adopted only if the forma-tion of H202 was exactly proportional to the flux of thermal neutrons irradiating the solution. From the work we did on boric and borate solutions we find that the following conditions must be fulfilled for such determinations (i) It is neces-sary to correct the results for the action of the known radiation deposit of the pile (y-rays fast neutrons etc.).This correction can be made by measuring the decomposition of different solutions of HzOz placed in the same container as the boron solutions. The reproducibility of results so corrected seems to prove that the assumption made for the correction is reasonable. However it is not certain that the effects of the pile radiation and of the x-rays produced by thermal neutrons are additive. (ii) Solutions containing boron must be either acidic (PH lower than 2) or alkaline (KB02 pH greater than 8). For pH’s between 2 and 8 the results are more complicated and no dosimetry is possible by titration of hydrogen peroxide.This is clearly shown in Dr. Haissinsky’s contribution to the Discus-sion and by the figures given in the publication previously quoted.10 (iii) The G value observed for the formation of H202 even in acid conditions is lower (0-68) than the G value measured when boric acid solutions is irradiated by x-rays from radon (0.9). This is in agreement with the yield for the oxidation of ferrous sulphate in aerated solution mentioned by Dr. Wright as G = 3.9 which is lower than the value of 6.0 given by Dr. Miller using u-rays of polonium. We found too that the reduction of ceric ions occurs in the pile through the B (na) reaction with a yield of G = 2.0 instead of 3.6 with a-rays of radon.12 As we pointed out in (i) the correction for the contribution of the y-rays of the pile may explain these lower values if the two effects are not additive.Prof. Milton Burton (Notre Dame University Indiana) said The track of an ionizing particle contains the residue of primary effects and the secondary effects induced by delta rays. The latter are regions of high concentration of excited molecules and ion pairs and are of the nature of poorly defined “ spurs ” to the primary track. Very little is known experimentally about the composition and internal geography of these spurs. However computation according to Bethe’s method for a system of hydrogen atoms indicates that the spurs are of approxi-mately the same nature for both slow and fast particle irradiation. Thus it follows that the principal parameter which distinguishes effects induced by slow and fast particles is the distance d between spurs the value of d i n liquids being of the order of a few A for slow particle effects and of the order of lOOA for fast ’0 Pucheault Lefort and Haissinsky J.Chim. Phys. 1952 49 286. 11 Bonet-Maury and Deysine Cumpf. rend. 1951 232 1101. 12 Lefort and Haissinsky J . Chim. Phys. 1951 48 368 118 GENERAL DISCUSSION particles. The distance d affects competition between chemical processes subse-quent to the early physical effects in a manner characteristic of the particular system affected. When two types of radiation are mixed regions in which d is small are superimposed on regions in which d is large. The effect is not additive. The average value of d is actually decreased by the superimposition of regions of large value of d although not so much as jt would be merely by increasing the intensity of slow particle irradiation.So far as 1 know the theory of such effects has not progressed to a point where we could predict effects in a chemical actino-meter in which a liquid system is employed. On the other hand such difficulties will probably wash out in a gaseous system. Such a gaseous monitor would tend to give purely physical information which would have significance in cases of bio-logical interest when supplemented with data from a liquid actinometer. Dr. S. Schrage in our laboratory is now comparing the effects of electron and neutron irradiation on the system methane + radio-iodine.Dr. A. G. Maddock (Cambridge University) said It appears to me that more careful consideration of the details of degradation of the incident radiation is urgently necessary. Very little direct information on the part played by solvated ions has yet been obtained and when chemical effects with very small G values are under consideration such as the “ molecular ” production of hydrogen de-scribed in Allen’s paper even the less common events in the energy degradation chain may assume importance. For instance some K ionization of the oxygen of the water molecules must take place. There is a definite probability that Auger multiplication leading in some instances to the breaking of both the oxygen-hydrogen bonds will occur. These hydrogen atoms might be conveniently located for molecule formation but the fate of the residual oxygen ion or atom must also be identified.When the energy of the incident electron is very great so that the binding energies of all the electrons of the water are negligible the fraction of ionizing events leading to K ionization must approximate to the ratio of the number of K to the total number of electrons in the molecule i.e. 20 %. When the energy of the incident electron is but a few times the binding energy of the K electrons of the oxygen the fraction of events leading to K ionization is the inverse ratio of the ionization energies of the K and the alternative level-perhaps 1 or 2 %. The probability that K ionization leads to rupture of both hydrogen bonds is much more difficult to estimate.A new theoretical study of the mechanism of thc dissipation of the energy in liquid water of for example incident 1 MeV electrons using stochastic methods would be most valuable. Dr. N. Uri (University of Chicago) (canrrn~mic.utetl) Allen and Miller and Wilkinson deal in their papers with aspects of dosimetry and radiation actino-metry. I would like to make a relatively simple suggestion for the use of the term primary radiation yield and wonder whether this is practicable. If for example, a dose of 100 eV leads to 10 radlcal pairs of H and OH (the formation of which might be preceded by ionization) one could compare the energy required to break the chemical bond H . . . OH viz. 5.2 eV with the total dose of 100 eV. As 10 radical pairs are produced the primary radiation yield for the splitting of water would be 0.52.On the basis of recent results obtained by Dainton and Rowbottom,l3 it would appear that the primary radiation yield could be close to unity. Dr. J. Weiss (Durhawz University Newcastle) said We have followed with great interest and indeed with admiration the wry interesting work of Dr. Allen and his colleagues. 1 think that it is quite clear that one should expect a greater or smaller extent of primary recombination of the hydrogen atoms in the original tracks the amount of which should depend on the nature of the radiation. Is such a terminology feasible? 13 Dainton and Rowbottom Nature 1952 169 370 GENERAL DISCUSSION 119 However with hard X-rays and y-rays it seems to me somewhat implausible that the G values for hydrogen evolution should have such a constant value (= 0.65) as Dr.Allen suggests. In some preliminary experiments which have been carried out by Mr. Rigg in this laboratory with X-rays (200 kV) we have found much smaller G values in certain cases Although these are only preliminary results we should like to present them here in the hope that similar experiments will be carried out in other laboratories so that this important question can be clarified in due course. Dr. T. Rigg (Durham University Newcastle) said Dr. Allen in his very interest-ing paper implies that a “ molecular ” decomposition of water into hydrogen and hydrogen peroxide occurs upon irradiation and that the extent of this process will be independent of the solute. The minimum G value for the molecular hydrogen evolved in this process for y-rays or hard X-rays is suggested to be about 0.6, using the value G = 20.6 for the ferrous sulphate actinometry.We have observed similar G values for hydrogen evolution with cerjc salts, but for some other solutes we have found in some preliminary experiments sub-stantially smaller yields. Table 1 shows some of the results recently obtained in our laboratory with various solutes. It seems that the solute can apparently influence the rate of formation of molecular hydrogen to an extent which depends upon its oxidizing power so that if some molecular decomposition of water does occur its extent would appear to be possibly less than Dr. Allen suggests. TABLE 1 .-HYDROGEN EVOLUTION FROM AQUEOUS SOLUTIONS IRRADIATED WITH X-RAYS (200 kV) ; mean dose rate 3,200 repimin ; dosimetry based on C = 20.6 for ferrous sulphate in 0.4 M sulphuric acid in the presence of oxygen remarks G for H2 (moles/100 eV) system 1.0 x 10-3 M ceric sulphate (1.0 + 0.1 M H2S04) ; vacuum 1.0 x 10-3 M ceric perchlorate (1 a 0 to 0.1 M HC104) ; vacuum 0.70 independent of the dose (up to 3.5 x 105 rep) 0.65 independent of the dose (up to 3.5 x 105 rep) 0.1 M ferric sulphate (pH = 1.7) ; vacuum 0-44 independent of the dose (up to 3-5 x 105 rep) 0.1 N potassium dichromate ; vacuum pure water ; vacuum 0.30 initial yield (maximum) 9.0 x 10-3 ferrous sulphate ( 0 2 saturated) 0.1 N potassium permanganate ; vacuum 0.35 independent of the dose (up to 3.5 x 105 rep) 0-15 0-1 0 dose = 0-5 x 105 rep independent of the dose (up to 3.5 >i 105 rep) (0.4 M H2SO4) It should be noted in particular that in the important case of the ferrous sulphate oxidation the hydroxyl radicals would have no opportunity of reacting with the very small amount of hydrogen present since it has been shown14 that the rate constant for the reaction of hydroxyl radicals with the ferrous ion is about 4 times greater than that of the corresponding reaction with molecular hydrogen.Another important point mentioned by Dr. Allen is that in the irradiation of pure air-free water the rate of hydrogen evolution rapidly falls off due to the accumulation of molecular hydrogen and subsequent back reactions. In this connection it should be emphasized that in radiation experiments on water it is imperative to specify the system completely and to consider (i) the energy and type of radiation (ii) the dose rate and integral dose (iii) the distribution of energy absorption in the liquid (i.e.geometrical considerations) and (iv) the volume ratio 14 Rigg Stein and Weiss Proc. RoJ SOC. A 1952 211 375 120 GENERAL DISCUSSION of the liquid and gas phases. Differences in any one of these factors can of course profoundly modify the result obtained. Mr. N. W. Luft (Waltham Abbey) said Could the discrepancies between theoretical and experimental H202 yields in the three representative systems dealt with by Dr. Allen be partly due to the special properties of the HO2 radical, formed in the reaction H + 0 2 -+ HO2? From Robertson's 15 recent value of D (H02 + H + 02) = 46 kcal one calculates for gaseous H02 a life of 6 x 10-10 sec at 300" K ; in '' hot spots " it would be smaller.Thus the reverse process would have to be considered too. Prof. I?. S. Dainton (Leeds University) said In his paper Dr. Allen states that an essential difficulty in estimating the free radical yield in water decomposition (GR) is that whatever action is produced in a solute by one radical can in general be reversed by the other and mentions that this difficulty may be overcome by using two suitably chosen solutes present together and each reacting with a different radical. These methods have led to values of GR lying in the range 3-5. It should be pointed out however that it is not necessary to work with 2-solute systems, since solutes are known in which the same kind of change may be induced by each radical.The reactions which are initiated are usually chain reactions and examples are the polymerization of vinyl compounds and the radiolysis of hydrogen peroxide. These reactions are of particular interest when they can also be initi-ated photochemically and the quantum yield of the primary act is known since, if it can be established that the secondary reactions are identical in both the radio-and photochemical cases comparison of the two overall reaction rates gives the value of GR. This method has been applied by Mr. Rowbottom 16 to the decom-position of concentrated aqueous solutions of hydrogen peroxide taking every precaution to purify the hydrogen peroxide and using the same solution in the same reaction vessel for the radiolysis and photolysis.The value of GR so obtained was 13.4 much higher than the values cited by Allan for other systems either 1-solute or 2-solute. Another exception is the reduction of ceric ion in acid solution which has been studied by Haissinsky and Lefort 17 and yields a value of GR. On both the H202 decomposition and the CelV reduction the H and OH radicals reinforce one another in their actions (OH radicals are responsible for some of the reduction of CelV which is observed) whereas in the systems from which low GR values have been deduced these radicals could act in opposition. These latter values are therefore to be regarded as minimal. Dr. C. B. Amphlett (A.E.R.E. Harwell) said The problem of solutes com-peting for radicals within the hot-spots and so lowering the molecular hydrogen yield is likely to be very complex.We have to consider competition both for H atoms and for OH radicals and possibly for H202 as well. Although in all cases to date only Hz yields have been measured it cannot be assumed that competi-tion for OH radicals will not affect H2 yields owing to the complicated interplay of radicals in aqueous solutions. The figures quoted by Rigg,ls if substantiated, bear out this viewpoint. Thus ceric ion by virtue of its ready reduction by H atoms H202 and possibly OH radicals also would be expected to be a good competitor for radicals produced in hot-spots and yet the molecular hydrogen yield is similar to that quoted by Allen for pure water ; the production of oxygen in the reduction of ceric ion would be expected to favour further competition by virtue of the reaction H + 0 2 -+ HO2.Ferric ion dichromate and permanganate, which would not be expected to be more efficient than ceric ion appear to be much more effective competitors. It is noticeable that the three other ions were present This discrepancy is perhaps not without significance. 15 Robertson Trans. Faraday SOC. 1952 48 328. 16 Rowbottom and Dainton Nature 1952 169 370. 17 Haissinsky and Lefort J. Chim. Phys. 1951 48 368. 18 Rigg this Discussion GENERAL DISCUSSION 121 at 1000 times the concentration of the ceric ion in these experiments which may explain their greater effect. Results bearing on this problem are of fundamental importance for the understanding of the radiation chemistry of aqueous solutions, and require detailed analysis of the kinetics under carefully controlled conditions.Prof. F. S. Dainton and Dr. H. C. Sutton (Leeds University) said In support of his hypothesis that a significant proportion of the products of the primary act consist of hydrogen and hydrogen peroxide Allen has cited several systems both aerated and de-aerated in which molecular hydrogen appears amongst the ultimate products. We have recently been seeking evidence for the production of H202 in equivalent amounts to this hydrogen in ferrous sulphate solutions subjected to X-radiation. The peroxide present initially was estimated from the magnitude of the post-irradiation effects observed in both air-free and aerated solutions when the ferrous sulphate concentration is initially 10-5 N (in 0.8 N H2SO4).At this solute concentration the complete oxidation of the ferrous iron by the H202 formed during the irradiation may take several hours and using the o-phenanthro-line method for [Fez+] estimation we have shown that the rate constant of this reaction is identical with that value published by Barb Baxendale George and Hargrave.19 We have verified that the same curves expressing [Fe3+] as a function of time can be obtained from solutions of Fe2j- Fe37 and H202 made up in com-positions corresponding to that which the irradiated solution (initially containing only Fez-) possessed at the time of cessation of irradiation. In establishing this latter point we made the incidental observation that these curves are unaffected by the presence of excess Fe3+ ion.The experiments in air-free solution are not yet complete but with [Fez+] initial = 1.2 x 10-5 N and a dose rate of 600 r/min of 220 keV X-rays the H202 accounts for about a quarter of the ultimate oxidation yield and corresponds to G H ~ ~ ~ N 0.8. This H202 cannot be due to incomplete removal of air since the residual 0 2 concentration in these experiments was less than 10-9 M. In aerated solutions some of the H202 formed is due to the reaction Fe2+ -I- HO2 + Fe3f + HOa- (1) The maximum contribution which can be made to the H202 yield by this means (2) H f 0 2 4 H02 (3) OH + Fez+ -+ OH- + Fe3+ (4) is obtained by combining eqn. (1) with eqn. (2) (3) and (4) H20 --\>+ H + OH Hence the H202 yield due to (1) cannot exceed 50 % of the initial yield of ferric ion.The observed yield of H202 is in fact 75 % of the initial ferric yield and it is evident that even in aerated solutions a significant proportion of the H202 formed is not due to the dissolved oxygen which is present. Accepting Allen's " hot spot " hypothesis and his figure for GH2 of 0.6 in aerated ferrous sulphate solution it is then possible to calculate a G value for H202 due to " hot spots " (i.e. to origins other than dissolved oxygen). The value so obtained is about - 1.5. The result that GM for H202 > G for H2 may be significant and could be accounted for on the basis of the different spatial distributions of H and OH. Dr. E. Collinson (Leeds University) (comniunicated) The experiments de-tailed in the table below provide further evidence for the formation of hydrogen peroxide during the X-irradiation of both aerated and de-aerated solutions of ferrous ammonium sulphate in agreement with the formation of this substance via the " hot spot " mechanism (cf.the note by Prof. Dainton and Dr. Sutton). The ferrous ion + hydrogen peroxide system is well known to be capable of initiating polymerization of acrylonitrile in aqueous solution and this behaviour was employed as an indication of the presence of hydrogen peroxide in irradiated solutions. 19 Barb Baxendale George and Hargrave Trans. Furaday SOC. 195 1 47,462 122 GENERAL DISCUSSION The irradiations were made with 220 kV X-rays at a dose rate of 180 r/min. Acry-lonitrile monomer was distilled into the irradiated solutions in sufficient quantity to render them 0.7 M in this substance.liquid irradiated 5 x lO-5MFe'f in 0-8 N H2SO4 5 x 10-5MFez" in 0.8 N H2SO4 pure water pure water pure water pure water irradiation time (min) 16.5 33 16.5 66 16.5 66 state of aeration aerated deaerated aerated deaerated aerated deaerated action after irradiation solution deaerated a n d acrylonitrile distilled in under vacuum acrylonitrile distilled in under vacuum solution deaerated a n d acrylonitrile distilled in under vacuum acrylonitrile distilled i n under vacuum solution deaerated ; acry-lonitrile distilled in and ferrous ammonium sul-phate added to give 10-5 M solution acrylonitrile distilled in and ferrous ammonium sul-phate added to give 10-5 M solution observation immediate polymer formation immediate polymer polymer visible only formation after several hours no polymer formed immediate polymer format ion no polymerization Dr.M. Lefort (Institut du Radium Paris) (communicated) The initial constant yields in the produclion of molecular hydrogen given by Dr. Allen are of great interest. However after a very short time of irradiation hydrogen will escape from the so-called hot spots through the whole bulk of the solution and can then enter into competition with the solutes for reaction with radicals consequently the yield of molecular hydrogen is no longer constant for various solutes. This could be the reason why Dr.Haissinsky and myself found different yields of hydrogen in the oxidation of arsenite salts reduction of ceric ions reduction of chromate ions and oxidation of ferrous sulphate. Therefore the production of H2 and H202 in hot spot regions is only a part of the observed radiochemical action and in many cases a very small part so that we are still confronted with the following problem why most of the chemical effects of the radiations are oxidation processes. Dr. Haissinsky and myself pointed out in 1949,2* what we called " the assymmetry between the action of OH and H radicals ". In 1950 we showed that even the well-established cases where reductions occurred to an important extent could be attributed to OH radicals because the redox potentials of such systems (Ce4+/Ce3+ Crs+/Cr3- etc.) had high values.Prof. Dainton and Dr. Collinson haven given a more precise definition of this concept for comparing the effects of radiation with the redox potential of the involved solutes.21 Comparing the well-known cases of total oxidation total reduction or a steady state between oxidized and reduced forms and their redox potentials the observed chemical effects could be explained by assuming that the irradiated water had itself a redox potential in the region of 0.8 to 1.0 V. Whatever is the real signi-ficance of this equivalent redox potential of irradiated water the value given above well illustrates how much more the irradiated water functions as an oxidizing than a reducing species. Equal amounts of reducing and oxidizing radicals cannot explain such a situation.With Dr. Allen's hot spots in the regions of hizh ion density redusing 20 Haissinsky and Lefort Ccmpt. rend. 1950 230 1 156. 21 Dainton and Collinson AWZ. Rev. Physic. Chem. 1951 2 99 GENERAL DISCUSSION 123 H atoms disappear because of the combinations H + H = Hz while the cor-responding reaction maintains the concentration of the oxidizing species. With x-rays we could attribute a large part of the observed chemical eKects to the action of hydrogen peroxide (in the reduction of Ce4-’ and CrO4- and oxidation of As3+ and Fez+). However it is well known that the results are often different with electrons pro-duced from X- and y-rays because regions of high ion density diminish in number and in space as the energy of the electron increases.‘Therefore the greater part of the chemical effects occurs in the rest of the solution where the distribution is more uniform. Dr. Haissinsky in his contribution showed that even where one might expect an action by H atoms in the solute because the free energy change for it is greater than for the oxidation by OH radicals yet the OH reaction occurs. We have another example in Dr. Amphlett’s paper; in spite of - 1 G being 65 kcal for the reduction of ferric ions by H atoms the reaction occurring is the oxidation. On the other hand when reduction by OH becomes possible (ferric ions complexed with o-phenanthroline EO = 1.1 V) then the reduction is complete and the yield is high (G = 12) although the free energy change of the reaction with OH is about 18 kcal that is to say lower than for redclction of ferric ions by H atoms (65 kcal).So even in the bulk of the solution where we have uniform distribution of radicals the concentration of the oxidizing species seems to be much greater than that oi‘ the reducing ones. Several hypothesis have been proposed to account for these considerations H2+ (Weiss) primary formation of H2 + OH from the ions H2O (Haissinsky and Magat) etc. As yet none of them appears to be sufficiently confirmed by experiments. However we think that the classical explanation of balanced production of OH and H radicals in irradiated water is no longer satisfactory even when corrected for non-uniform distribution as proposed by Dr. Allen. Dr. M. Haissinsky (Paris) said The constancy of the H2 yield claimed by Dr.Allen seems to be somewhat surprising. With Dr. Lefort we have examined several redox systems and it was the peculiar radiochemical behaviour of each of them that struck us no one being similar to the other. Following Allen’s sugges-tion equivalent and constant amounts of H2 and H202 are originated nearly in all aqueous solutions. But the measured yield of H202 is surely variable from one case to another and one must admit that this variation is due at least partially, to back reactions. One can then not understand why H2 is not sensitive to these reactions the importance of which depends upon competition conditions. Allen was one of the first lo underline the importance of the reaction OH + OH ~ H202 H2 f OH = H + H20.Moreover many redox systems are known for which it is necessary to admit a consumption of variable amounts of OH radicals or of an excess of these relative to H atoms ; an equivalent quantity of the latter arising from the reaction allowed by Allen H20 = H + OH combines then to form Ha molecules and the measured yield of the gas could not be the same for various solutions (see for instance the oxidation of FeS04). As these cases are far from being exceptional it would be useful if Dr. Allen would kindly clarify his position. My opinion is on the other hand that the actual experimental data are as yet insufficient and too conflicting to justify gencralizations and over-simplifications, a tendency which is now too marked in radiation chemistry. Dr. A. 0. Allen (Brookknscn ATutionul Lob.N. Y.) (cornnzunicated) I agree with Prof. Haissinsky that too many generalizations have been offered in radiation chemistry and that we must avoid oversimplification. We do however find empirically that the same yield of hydrogen gas is produced from a wide im-iet 124 GENERAL DISCUSSION of dilute solutions provided the same radiation source is used. We have demon-strated this for solutions of KBr KZ HCI CuSO4 Ce(S04)2 H202 0 2 and oxygenated FeS04 while Fricke and Hart had previously demonstrated it also for solutions of nitrite arsenite selenite and ferrocyanide. The constancy of the hydrogen yield refers only to the initial portion of the irradiations and back reactions may reduce the yield after the hydrogen concentration has built up to some extent; some of the low yields observed by Mr.Rigg and Dr. Weiss may possibly be ascribed to this cause. Aside from the hydrogen producing reaction, the free radicals produce a variety of reactions in the different solutions as Prof. Haissinsky says and it is precisely this variety which gives rise to uncertainty in the value of the free radical yield GR. Whether the hydrogen is accompanied by an equivalent amount of hydrogen peroxide or by a greater or smaller amount, or whether the primary molecular products may include oxygen as well as hydrogen peroxide are questions not yet firmly settled. The interesting work mentioned by Dr. Sutton and Prof. Dainton if confirmed would indicate that hydrogen peroxide is formed in larger amounts than hydrogen. This is the reverse of the effect expected on the theory of Haissinsky and Magat for the formation of molec-ular hydrogen.If the quantities of molecular hydrogen and hydrogen peroxide formed primarily are widely different it is somewhat difficult to explain the fact that hydrogen and oxygen are the main products from irradiation of dilute neutral bromide or iodide solutions with peroxide appearing only to the amount of 17 % of the hydrogen. If excess hydrogen atoms were produced in this system as suggested by Sutton and Dainton one would expect them to reduce the oxygen to hydrogen peroxide and keep the oxygen/hydrogen peroxide ratio in the products below the high observed value. Prof. Milton Burton (Notre Dame University Indiana) said We should avoid use of words which may prejudice our ideas of mechanism.The idea of “ ion clusters ” developed by S. C . Lind now has a very particular meaning in radiation chemistry. We have seen from the paper by Prof. Massey that such ion clusters have physical reality. In a paper which Magee and I will present at the forth-coming symposium on Electron Transfer and Isotopic Reactions,22 we show that ion clusters have a significant but not necessarily exclusive role in radiation chemistry. Thus it is particularly important to avoid use of words which can be confused with ion clusters and in turn confuse interpretation. In Dr. Allen’s paper he has been careful to avoid confusion in speaking of “ formation of ions in clusters ” but may I suggest that the word “ spur ” be used in the future in this sense rather than cluster where spur signifies merely a region of high concentration of ions or excited molecules disposed locally around a primary track and that we reserve the word cluster for the now classical concept of Lind.The term “ hot spot ” has also been employed in the same sense as we use spur but it is evident that hot spot also prejudices our thinking of the nature of both primary and secondary processes. Dr. A. G. Maddock (Cambridge University) said With reference to Prof. M. Burton’s paper I should like to draw attention to a novel form of protection that may become important in aromatic systems. Yesterday’s discussion revealed some difference of opinion as to the magnitude of the photon emission in the ultra-violet arising from the irradiation of aqueous systems with penetrating ionizing radiations.Although such photons can be detected with most if not all, liquids for instance by means of a coincidence arrangement with two photo-multiplier tubes I am very doubtful if they can play any important part in aqueous systems. However I find the magnitude of the photon emission from other liquid systems to range over several orders of magnitude. The effect is largest in those systems that have been employed as liquid scintillation detectors for y radiation-there is no doubt that the emission from such systems is much greater than that due to the Cerenkov radiation alone. 22 Notre Dame June 11-1 3 1952 GENERAL DISCUSSION 125 I have examined the photon emission from some hundred such systems. Most of these consist of very dilute solutions of a fluorescent solute in an aromatic solvent.I find that strong photon emission is excited in systems fulfilling four conditions the solvent must be transparent to the fluorescent radiation of the solute quenching centres such as oxygen must be absent and the solvent and solute must contain extensive systems of conjugated double bonds. The magni-tude of the photon emission is such that there can be no doubt that the energy of excitation of the solvent molecules can be transferred across a chain of solvent molecules to the solute which then radiates a photon. This mechanism of energy dissipation closely similar to the case described by Prof. M. Burton has been examined theoretically by Forster.23 It might be expected that the process would lead to a measurable reduction in radiation decomposition of the solvent or protection.A simple test of this hypothesis would be to compare G(H2) and G(C2H2) for benzene with the same quantities determined for an 0.2 % solution of p-terphenyl in benzene-a solution which exhibits more than 100 times the photon emission of pure benzene. My remarks have so far been confined to two component systems but it should not be supposed that strong scintillations are not found in any pure liquids. Early last year I found that the purest cyclo-octatetraene I could prepare (m.p. - 5 . K ) gave large scintillation pulses nor could these be diminished by chromatographic purification of the compound. I think that careful attention should be paid to the part played by such energy transfers in irradiated systems particularly when aromatic or unsaturated com-pounds are concerned.Dr. E. Collinson (Leeds University) (communicated) In the work reported by Prevost Bernas et aZ. no direct dosimetric measurements were made on the y-irradia-tions. The authors themselves state that a theoretical evaluation of the dose rates was difficult and this raises the query as to whether the rates of polymeriza-tion in all the cases quoted in their fig. 1 were established as dependent on the square root of the dose rate. The importance of this lies in the fact that the results deduced from the polymerization experiments depend on the existence of a mutual termination mechanism When y-ray dose rates are varied by adding or removing sections of a source with otherwise unchanged geometry it is almost certain that the dose rates will not be proportional to the strength of each source employed.Owing to self screening and the greater deviation from a point source, the dose rates corresponding to the larger sources will be proportionately lower than those corresponding to the smaller sources. This would mean that results giving an apparent dependence of rate on (dose rate)* would in fact represent a dependence on (dose rate)x where x > $. Mr. B. Coleby (King’s CoZZege Newcastle) said When the effects of X-rays upon steroid compounds were first investigated by Keller and Weiss sterols were converted into the sodium salts of their acid succinyl esters so that the effects of X-irradiation in aqueous systems could be investigated.Later it was ascertained that the same products could be obtained by using 85 % acetic acid or methanol as solvents and these solvents were then used exclusively. In order to gain some idea of the primary reactions in these solvents so that a better idea of the mechanism of the attack of the steroids could be obtained the effects of irradiating diphenylpicrylhydrazyl (DPPH) in these solvents was investigated. As a preliminary step we studied the effects of the presence and absence of moisture and oxygen upon the reaction induced by y-rays. Starting with a solu-tion of DPPH in chloroform at an initial concentration of 10-7 mole/ml the decrease in optical density at a wavelength of 5,300 8 was measured after exposure to an 800 mg source of radium.With dry chloroform in the presence of air the decrease in optical density was linearly proportional to the time of irradiation 23 Forster Naturwiss 1946 33 166; Ann. Physik 1948 2 55 126 GENERAL DISCUSSION until about 70 % of the initial DPPH had been consumed (irradiation was not continued beyond this point). When a similar solution was irradiated in the presence of water (saturated solution) and in the absence of oxygen (the irradiation vessel was evacuated and then allowed to fill with purified nitrogen) the linear rela-tionship was again obtained and the slope of the line was not altered. It was definitely established that water and oxygen had no appreciable effect upon the rate of the reaction. The great sensitiF ity of the reaction and the ease of measuring the rate of consumption of DPPH colorimetrically are properties which may render this radical suitable for use in dosimetry work.To gain some idea of the primary reactive species produced by irradiation we attempted to isolate the products which they form with DPPH by means of chromatographic separations of the irradiation products on alumina. DPPH in chloroform or methanol solution was irradiated with X-rays (200 kV 15 mA) so that only a small proportion of the DPPH was consumed. The solvent was then removed under reduced pressure and the residual solids were separated by means of elution chromatography. Elution with petrol + benzene mixtures gave the un-changed DPPH followed by ax-diphenyl-P-picryl-hydrazine which was character-ized by its melting point mixed melting point and its ultra-violet absorption spectrum.This is the only irradiation product which we have been able to characterize as yet. The hydrazine and the unchanged DPPH usually account for 95-100 % of the starting material. There is evidence to indicate that some at least of the diphenylpicryl hydrazine is formed by the decomposition of reaction products on the alumina column, since e.g. the hydrazine was also isolated during the hromatographic separation of the products obtained by irradiating a solution of DPPH in purificd carbon tetrachloride. When zid-washed alumina was used for this separation oil products were isolated which had characteristic ultra-violet absorption spectra, but which decomposed to the hydrazine on standing (even in a desiccator).It is doubtful whether the reaction products are sufficiently stable to pcrmit of the application of the usual chromatographic procedures. Further work is proceed-ing and will be published in due course. Mr. J. Wright (A. E. R. E. Hcn-wll) (cori~nrriiiic'atcrl) In discussing alternative mechanisms of radiation-induced decomposition of benzene based entirely on a study of the gaseous products it is important to ensure that these products are representative of the total decomposition. In work on pile irradiation of solid naphthalene and anthracene in vacuum carried out at Harw ell in collaboration with Miss E. M. Dresel we have found gas (mainly hydrogen) yields which indi-cated much smaller percentage decomposition than that found on examining the irradiated solid.Thus in an irradiation of antliracene in which 0.14 :< of the hydrogen atoms appeared as Hz gas (i.e. 1.4 of the anthracene molecules affected if one H atom was derived from each molecule) 8 of the irradiated solid was insoluble in petroleum ether and separation of the soluble fraction on an alumina column gave several bands in addition to that containing unchanged anthracene. In such cases a study of gaseous products alone may be misleading and the mechanism of decomposition cannat usefully be discussed until the major solid and liquid products are identified and a rough material balance established. These observations have a bearing on the much lower G values for gas production in this paper compared with the G values for free radical production given by Dr.Magat in the next paper. Prof. F. S. Dainton (Leeds Uriivcrsiry) said The results described by Dr. Magat and his collaborators raise several interesting questions. In using " radical-catching " agents such as DPPH for determining GR it IS necessary as the authors point out to ensure that all the radicals are caught and therefore to compare radiation yields only under conditions for which it has been previously established that the yields do not depend on either the dose rate or the concentration. How-ever much may also be learnt from the concentration dependent region of th GENERAL DISCUSSION 127 rate-concentration graphs. If the mechanism is as summarized by the authors then as the solute concentration is reduced the dose rate exponent will decrease from unity to a limiting value of 0.5 provided the radicals formed in different tracks are freely interacting.The data presented in fig. 1 indicate that the limiting dose rate exponent is certainiy not unity. Three possible explanations of this fact may be offered. Firstly that a substance (unknown) is present which can compete with DPPH for the radicals. Secondly that most of the radicals which recombine have originated in the same track i.e. non-uniform kinetics should be applied. Thirdly, that relatively few free radicals are formed and the major part of the energy ab-sorbed is used to excite solvent molecules to a degree which is inadequate to ensure dissociation but which suffices to ensure reaction with DPPH and that such excited molecules may also undergo first order deactivation.This last sug-gestion would also explain the surprisingly high degree of energy utilization observed in some of the systems summarized in table 1. In using the polymerization method a much more extensive investigation is necessary. One must be sure that the kinetics of the polymerization are given by eqn. (HI) p. 751 ref. (2) : where GRS is the radical yield for the solvent and GFi is the radical yield for the monomer. In view of the fact that in some cases the values given in column 3 of table 1 of the relative numbers of free radicals differ considerably either from the values cited previously or from the values given by the DPPH method can the authors say which of the two methods they regard as the more reliable and whether eqn.(111) was fully verified for each system for which eqn. (111) should be applicable ? Dr. W. G. Barb (Cowtardds Ltd. Maidenhead) said Prof. Dainton has asked whether some of the polymerization measured by Dr. Magat and his collaborators might perhaps be due to ionic rather than free-radical initiation. This point could be tested experimentally by the use of a mixture of two suitable monomers in place of styrene alone. It is well established that in certain monomer mixtures free-radical and ionic initiators lead to different reaction products ; 26 thus e.g., Walling et a1.27 have shown that mixtures of styrene and methyl methacrylate can give a copolymer nearly pure polystyrene or nearly pure polymethyl methacrylate, depending on whether a free radical carbonium or carbanion mechanism is in-volved.A suitable separation and analysis of the product will therefore indicate the contribution of these various processes. Dr. W. Wild (A.E.R.E. Harwell) said There seems little prospect of increasing our understanding of the radiolysis of organic compounds by study of the gaseous and the very complicated condensed products. The programme outlined by M. Magat is a very valuable and encouraging contribution to this difficult field. If one compares this problem with that of carbonyl photochemistry it seems reasonable to assume that progress will require much effort. With these thoughts in mind we began quite independently an examination of the possibilities and validity of the use of diphenylpicryl hydrazyl (DPPH) as a means of capturing, counting and characterizing radicals liberated by ioni7ing radiation in organic solvents.Spectrophotometric observations of solutions of DPPH after irradiation lead one to believe that the only process occurring is addition of other radicals to give related hydralines. Subsequent chromatography on alumina and lime columns 24 Dainton Research. 1948 1 491. 25 Chapiro J. Clliin. Phys. 1950 47 747. 26 Mayo and U'alling Chem. Rev. 1950 46 277. 27 Walling Briggs Cummings and Mayo J . Amer. Chern. SOC. 1950 72 45 128 GENERAL DISCUSSION modifies this conclusion to some degree. After irradiating dilute solutions of DPPH in benzene six bands are found on such columns four coloured and two fluorescent of which one coloured band constitutes 50-60 % of the recoverable products and has a spectrum resembling that of the parent hydrazjne of DPPH.Some of the other bands on elution give spectra showing no resemblance to those of the hydrazyl or hydrazine. As these two substances appear to be stable on such columns and as the total weight of material recoverable from the columns is con-siderably less than that of the DPPH used it is suggested provisionally that some of the DPPH is actually disrupted during irradiation of the solution. Our experi-ence renders it unlikely that any substance of the type R-DPPH would pass through the column and be lost but smaller fragments might pass through. It seems unlikely that such a fragmentation of the DPPH skeleton is the result of direct action of fast electrons at the concentrations employed.Moreover DPPH has been used under very mild conditions to study the rate of radical production from polymerization initiators undergoing thermal decomposition satisfactory proofs of its reliability having been given.28 In our opinion this is no adequate guarantee that the method is reliable in all cases. The severe conditions in the neighbourhood of electron tracks with the attendant possibility of “ hot ” radicals being present may mean that a radical may add on to DPPH and break the weak N-N bond simultaneously in which case three radicals may be consumed to give stable end products. On the other hand, another process may occur in some solvents. Benzene for example forms a very stable 1 1 compound with DPPH. Indeed recrystallization from benzene gives a particularly pure sample provided adequate care is taken in the final removal of this solvent.If electronic excitation occurs when ionizing radiations traverse aromatic solvents it may happen that by resonance transfer this energy may be transferred into the solvation shell of the diphenylpicryl hydrazyl and cause its disruption into fragments consuming more than one of the radicals arising from solvent ionization. A direct test of this possibility is to be undertaken shortly. For these reasons it is not yet possible to say that DPPH counts radicals more accurately than by a factor of (say) two or three. This may be a contributory cause to the discrepancies noted by Dr. Magat between the results of the two methods employed. The general qualitative trends of radical production from different solvents are not questioned as they seem consistent with other evidence.It is a relatively easy task to accumulate tens of milligrams of column fractions, when using a 0.1 /LA beam of 1.5 MeV electrons. The examination of the principal fraction from benzene by elementary microanalysis and infra-red spectra has revealed a factor which may confuse the identification by both methods of the radicals produced in the solvents. The products are very difficult to purify by recrystallization and our analyses in this one case suggest that some solvents are tenaciously adherent to this type of hydrazine and complete removal is absolutely essential to success. Literature evidence supports this view in the case of diphenyl-picryl hydrazine.The composition of the fraction mentioned suggests the addition of a CsH fragment. From a microanalytical viewpoint the large molecular weight of DPPH is a definite disadvantage of the DPPH method for identifying the usual organic radicals. Experiments are being pursued with simpler molecules and the possible use of some stable molecules should not be ignored. It is of interest that the radical yield from benzene using 1.5-1.8 MeV electrons, as deduced from the rate of disappearance of DPPH colour is of the order G = 0.8 for solutions prepared in vacuo and somewhat lower (0.6-0.7) for aerated solutions. These yields are markedly lower than those quoted in table 1 of Magat’s paper. A low value of about 0.5 has been suggested for styrene resulting from molecular weight measurements on polystyrene produced by cobalt y-rays coupled with the assumption that termination proceeds by mutual combination of two growing Two possibilities are conceivable.28 Bawn and Mellish Tvnns. Faradqy Soc. 1951 47 1216 GENERAL DISCUSSION 129 chains.29 Our yields in carbon tetrachloride are much larger than in benzene but we are not satisfied that the results in halogenated solvents are not confused by purely thermal reactions with DPPH. Despite the most rigorous purification and drying of carbon tetrachloride we have never prepared in vacuo a stable solution of DPPH in this solvent although air-saturated solutions in the same sample have survived many months of dark storage without loss of colour. The role of oxygen clearly requires further study.Dr. A. Chapiro (Paris) (partly communicated) In our polymerization experiments the intensity of the 7-radiation was undoubtedly not quite proportional to the intensity of the source as expressed in curies for reasons indicated by Dr. Collinson. However recent calibrations with small ionization chambers have shown that the errors probably do not exceed 10 %. This degree of inaccuracy may explain the observed small deviations from the theoretical I* plot but are insufficient to throw doubt on the square-root relation between the yield and the intensity. We in-tend howe\er to repeat these experiments with a stronger source that will soon be available and which will permit a greater variation in intensity. We have also observed as Mr. Coleby has done that DPPH solutions in chloro-form give identical results in presence of air and in vacuum.But this has not yet been verified for other solvents particulai ly alcohols. Prof. Dainton’s remark that the limiting dose exponent in the rate against DPPH concentration curves must vary from 0.5 to 1 when the DPPH concentra-tion increases is completely verified by our experiments (see fig. 1 of my discussion remarks). The slope of the initial linear part of the curves is proportional to 1 4 within experimental errors. In the concentration-independent part the rate is proportional to I as can be seen from the following figures : I rate rate/Z (rjmin) 108 moles/cm3 min 69.0 1.34 1.94 27.5 0.525 1-91 17.3 0.315 1.82 The last value is somewhat less reliable since the experimental errors increase for very slow reactions.Dr. M. Magat (Paris) (partly communicated) Prof. Dainton raised several important questions that I shall try to answer in the order they were presented, in so far as they have not yet been answered by Dr. Chapiro. It is of course impossible to exclude a priori the possibility of DPPH reacting not only with free radicals but also with excited molecules. This is a limitation of which we are perfectly well aware. However on one hand there is a rather good agreement between the polymerization results for styrene and the DPPH results particularly if one takes into account Dr. Valentine’s remark. On the other hand we have tried to induce a reaction between methanol and DPPH by irradiating the solution with infra-red light of 3 to 4 p in order to excite the OH vibrations.This attempt has been so far unsuccessful. Concerning the polymerization method the equation quoted by Dainton was verified to the following extent. (i) The intensity dependence was verified for pure styrene for a 70 % styrene + 30 % methanol mixture and for styrene + benzene cyclohexane and ether mixtures. I agree that a verification in each individual case would be desirable, and we intend to test several more cases in the near future. (ii) The dependence on the amount of solvent added was verified in the con-centration range from 10 % to 80 % for aromatic hydrocarbons and in a some-what smaller range for several other solvents. For polymer precipitants and poor solvents the equation breaks down at lower concentrations of the diluent.The probable reasons for this were amply 29 BNL 141 (T-27) (1951) (Brookhaven National Laboratory Report) 130 GENERAL DISCUSSION discussed by Chapiro 30 and 1 need not repeat them at length. It may be noted that in very dilute solutions of styrene in methanol the intensity exponent also increases to a value between 0.5 and 1. (iii) The difference between the figures of column 3 table 1 and the figures given previously 30 arises from the use of different units. The previous figures were expressed in relative numbers of free radicals per mole while in table 1 we refer to the relative numbers per cm3. (iv) The DPPH method is more direct and is to be preferred whenever it can be used with the reservations made under (i).Unfortunately it is difficult to use with compounds containing double bonds or very motile atoms because of a non-negligible thermal reaction. DPPH is very sensitive to traces of peroxides, sometimes difficult to eliminate. The polymerization method is rather doubtful if the diluent is a poor solvent or a precipitant. Hence one or the other method is to be preferred according to the substance investigated. In connection with Dr. Wild’s remarks one could wonder what would be the repercussion on the counting of the free radicals by the suggestion that the DPPH radical is broken up by the reaction with another free radical. Two cases may occur : (i) The break is such that a free radical fragment and a stable molecule are formed. The free radical fragment will react with the radiation-produced free radicaIs and the count will remain correct.(ii) The break leads to the formation of a free radical and a bi-radical. Then each DPPH molecule will consume three radiation-produced free radicals ; the real number of radiation-produced radicals will hence be even larger than the one we indicate. We agree with Dr. Wjld that far more work ought to be done along the lines indicated by him before the problem of the DPPH method will be completely solved. I would like to add that since the Discussion we have performed a calibration of our installation with the 275 mc Ra source both using small ionization chambers and by the FeS04 method. According to this calibration all the values listed in columns 4 and 5 (GR and Y ) of our table 1 are to be multiplied by a factor 0.82, which makes the energy yield values more reasonable.The absolute values given in table 2 correspond to an intensity of 1.13 r/min. However there remains a real difference between the GR value of benzene as detcrmined by Dr. Wild with fast electrons (GR Two explanations may be offered either with the intensities used by Dr. Wild the concentration independent region is difficult to reach while still avoiding direct effects or else there exists a difference of the GR values for y-rays and for electrons as was suggested at this Discussion by at least two authors viz. Chapiro and by Alper Ebert Gray Lefort Sutton and Dainton. Prof. Milton Burton (Nofre D a m University Indinno) said Table 1 of the paper by Dr. Magat and his collaborators gives G (free radicals) from benzene N 0.8 contrasted with G(H2) = 0.036 and G(C2H2) = 0-02 in the report by Gordon and myself.Our work indicates that not all of the gaseous product is of free radical origin; we can confidently write for the free radicals involved in such formation G < 0.05. Work by Manion and myself31 indicates that for the re-action benzene molecules + polymer G _N 0.8. Unreported work by Dr. Burr some years ago showed that the polymer aggregates contain on the average 5 or 6 molecules of benrene. Assuming one radical to start and one radical to termina?e the polymer it appears that for radicals involved in such formation G N 0.3 or that for total free radicals produced G > 0.35. This result is in much better agreement with Dr.Magat’s cri 0.8 than I thought during the actual discussion at 0.6-0.7) and by our group with y-rays (GR = 1.47). 30 Chapiro J. Chim. Phys. 1950 47 747 764. 31 Manion and Burton J. Physic. Chem. (May) 1952 GENERAL DISCUSSION 131 the meeting. However the discrepancy is real. It may be connected with the limited range of validity cf Dr. Magat's methods indicated by him in his discussion of table 1. In particular there is a special dificulty of radiosensitization when a foreign substance is introduced into benzene to determine the free radical yield from benzene itself. One must be particularly certain that the added substance does not have lower excitation or ionization potential than benzene for if it does energy may transfer to it from the benzene.If further that added substance should be more labile than the benzene (a highly probable situation) it may happen that its presence greatly increases the free radical yield out of all proportion to the concentration of the additive. Dr. L. Valentine (Birminghanz Ujiiversity) (corninmicafed) In the interesting paper by Prof. Magat and his collaborators there is a iiide scatter in the estimates (table 2) of the absolute numbers of free radicals produced. Account has not been taken however of the fact that the rate constants found by Bamford and Dewar (B and D) and Melville and Valentine (M and V) were calculated assuming that the termination reaction was disproportionation m hilst Mttheson et al. assumed combination. This automatically introduces a factor of 4 2 in the values of kl,,'kl& i.e.a factor of 2 in the estimates of the number of free radicals formed. Evidence is accumulating that the termination step in the polymeriiation of styrene is combination and if the results are recalculated on this basis taking the overall energy of activation to be 6.2 kcal/mole the \slues of kp/A,i become 0.0067 (M and V) and 0-0058 (B and D). The number of free radicals produced per cm3 per sec then becomes 0.9 x 1010 (B and D) and 0.7 x 1010 (M and V) compared with values of 1.7 1010 (Matheson ef a/.) and 1.3 x 1010 (using DPPH). There is much better agreeinent amongst these results than amongst the ones quoted in the paper. The value quoted as being determined from a combination of molec-ular weight and polymerization rate data must also be doubled if disproportiona-tion has been assumed raising it to 2.2 Dr.Philip George (Canikridge University) said The use of ferrous sulphate solutions in dosimetry and the investigation of the reaction mechanism which this has necessitated make it extremely probable that the problem as to whether similar reactions cam occur with the haernoproteins \?/ill attract attention in the very near future since these compounds contain ferrouq or ferric iron in a co-ordination coniplex and play an essential role in metabolic processes. Many of the free radical reactions postulated in ferrous salt dosimetry are identical viith those believed to occur in the reaction of iron salts with hydrogen peroxide,3? and in a recent reL iew 33 the possibility of these reactions occurring with typical haemoproteins has been considered on the basis that the probable influence of this co-ordination on the iron is to lower the ionization potential of its ferrous form in aqueous solution.This leads to the conclusion that reactions of the ferrous form will be some 5 kcal more exothermic and those of the ferric form more endo-thermic by the same amount with the result that all the ferrous ion reactions are equally or more favoured energetically with the ferrous form of haemoproteins. Ionic ferric iron reacting with 0 2 - is exothermic to the extent of about 19 kcal, and so a similar reaction occurring with haemoproteins is still likely to be exo-thermic to about 14 kcal and very rapid its reaction with H202 via the anion 02H- is on the other hand endothermic by about 28 kcal and a similar reaction with the ferric form of haemoproteins will be some 5 kcal more endothermic. Furthermore there is no reason to expect such a high frequency factor of the order 1024 as estimated from the ferric iron data for haemoproteins and so on two accounts this reaction is likely to be very slow indeed. Calculation 33 shows 1010. 32 Barb Raxendale George and Hargrave Nature 1949 163 6 ; Trans. Faraday Soc., 33 George Ahmices in Catalysis vol. 4 (Academic Press New York 1952) p. 367. 1951 47,462 and 591 132 GENERAL DISCUSSION the bimolecular velocity constant at pH 7.0 and 20" C would probably be of the order 10-19 mole-1 sec- 1. This conclusion is of great significance in deciding a reaction mechanism for the haemoprotein enzymes peroxidase and catalase which nevertheless in their ferric forms react extremely rapidly with hydrogen peroxide. Intermediate com-pounds are formed in these systems and also with the very similar haemoproteins metmyoglobin and methaemoglobin which have long been regarded as enzyme-substrate complexes.34 Recent researches,35? 36 however have shown that certain of these compounds are not of this type but contain the iron in an effective quadrivalent oxidation state for they are reducible to the ferric state in a one equivalent step. The " secondary " compounds of the peroxidases and the metmyoglobin compound show this property. Furthermore hydrogen peroxide or the structurally similar alkyl hydroperoxides are not unique in yielding this type of intermediate compound. With horseradish peroxidase hypochlorous acid hypobromous acid bromate periodate chlorite chlorine dioxide ozone and perdisulphate with silver ions all yield the second intermediate compound. Metmyoglobin gives its intermediate compound with chlorite chloriridate ions and molybdicyanide ions and in the last two cases the reaction is much faster than with hydrogen peroxide.6 The effective redox potential of the metmyoglobin compound is about 1.0 V at pH 6.8 (European convention) and about 0.7 V at pH 11.0 showing that a hydrogen atom is involved in the redox reaction.37 This compound and the secondary com-pound of peroxidase are alike in reacting with reducing agents such as the ferro-cyanide ion more rapidly the higher the hydrogen ion concentration which suggests that a hydrogen atom is similarly involved in the redox reaction of the peroxidase compound . In the radiation chemistry of systems where these haemoproteins are present the formation and reactions of such intermediate compounds should be taken into account along with the favoured ferrous and ferric ion type of radical reactions, not only because hydrogen peroxide is formed by irradiation but because these compounds may well be formed by oxidizing entities present in the system before the production of the hydrogen peroxide molecule as such. 34 Chance Advances in Enzymology vol. 12 (Interscience New York 1951) ; The 35 George and Irvine Nature 1951 168 184 ; Biuchem. J . (in press). 36 George Nature 1952 169 612. 37 George and Trvine unpublished results. Enzymes (Academic Press New York 1951)

ISSN:0366-9033    

DOI:10.1039/DF9521200110

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

In. INDIRECT ACTION-AQUEOUS SYSTEMS WITH SINGLE SOLUTE MECHANISM OF RADIOCHEMICAL TRANSFORMATIONS IN AQUEOUS DILUTE SOLUTIONS BY M. HAISSINSKY Institut du Radium, Laboratoire Curie, 11 Rue Pierre-Curie, Paris 5e, France Received 23rd January, 1952 Assuming that radiochemical transformations in dilute aqueous solutions are free radical reactions, it is considered that they are essentially due to the oxygenated radiolytic derivatives of water. Oxidation reactions are preponderant, but even reduction reactions are mostly provoked by OH radicals either directly or indirectly through the reaction : H2 + OH = H20 + H. This conclusion and the high yield of molecular hydrogen obtained during some radiochemical reduction and oxidation reactions justify the hy- pothesis that H2 is formed mostly through an elementary process of the type: H20 + e, H2 + 0-. Such a mechanism leads to the formation of two OH radicals and one mole- cule of H2 per ion pair.Several somewhat obscure kinetic peculiarities in radiation chemistry are pointed out. 1. PREPONDERANCE OF OXIDATIVE REACTIONS.-with the fundamental work of Fricke and of Dale and with the hypothesis of Weiss concerning the origin of the components of " activated water ", the radiation chemistry of aqueous solu- tions started on a new path, that of free radicals. This has led to reliable and satisfactory interpretations and reaction schemes at least while the approach was qualitative or semi-quantitative. The theory looked especially fruitful since it enabled many radiochemists to attack various new problems and to formulate generalizations and predictions that were to be confirmed by further experiments.These experiments have not always confirmed the anticipations, especially when quantitative method were employed. It seems now not only that these theories must be modified in some essential points, but that something fundamental for the understanding of radiochemical phenomena in solution is still missing. We shall assume that the principal concepts of these theories are well known, i.e. the origin of the free radicals H and OH with the first, acting as an oxidant, and the latter as a reducing agent, " radiochemical equilibria " or more exactly, the steady states, effects of ionization density, competition, protection and dilution effects and also the kinetic equations, which were proposed by Weiss, Lea, Dainton, etc.* In 1948 we attempted with Lefort to apply these conceptions to the experi- mental data given in the literature, and we were able to make the following statements : 5.637 (i) Radiochemical oxidations appear to preponderate as compared to reduc- tions, particularly with x-rays and simple inorganic compounds where the nature of the reaction products can be identified with certainty.More recent results that we and other workers have obtained confirm this conclusion in the fields of inorganic, organic and bio-chemistry (see 9 2). (ii) No radiochemical equilibria are known with certainty for a-rays. An equilibrium was observed with nitrate solutions,2~ 5 but the speed of reduction * for recent bibliography see Dainton,l Lefort,2 Dainton and Collinson 3 and Burton.4 133134 TRANSFORMATIONS IN AQUEOUS SOLUTIONS was so small that the reaction could be attributed to the &rays or to the p- + y-rays of the active deposit.To date only four are known : (a) H202, 02; (b) 12, I-; (c) AsO43-, AsOz-; (d) NO3-, HN02. It is still possible that with further research other radiochemical equilibria will be discovered under particular conditions,g but this would not change the general observation that the net balance of radiochemical transformations is mostly either a total oxidation or a total reduction. (iii) Reductions usually occur with oxidants possessing a high electrochemical potential. A good example is given by the behaviour of iron complexes. Ferrous salts are quantitatively oxidized by X-9 and cc-rays,lo at least in acid solutions; no reduction of ferric salts takes place 11 (Ei -= 4- 0.77 V).On the other hand the o-phenanthroline ferric complex (Ei == 1.14 V) is reduced by both species of rays 79 8 at pH 6. However, we shall see later that the oxidation potential is not the only deciding factor. (iv) Whenever a compound is reduced with a relatively high yield by ci-rays it is also reduced by H202 under the same chemical conditions (CeIV, CrV1, MnVrx). To explain the preponderance of oxidation reactions one must examine various factors which are essential for radiation chemistry in solution : (i) The different behaviour of ci- and X-rays leads one to consider the effect of the initial geometrical distribution of free radicals near the ionization tracks.It is believed that for ci-rays the OH radicals are concentrated in a cylindrical space of 15-20A radius around the track, while the H atoms occupy initially a much larger coaxial cylinder of about 150A radius. If the OH radicals are responsible for oxidation and the H atoms for reduction, the latter would be more frequent, since the H atoms should be more efficient due to their more homogeneous distribution. The greater rate of diffusior, of H atoms should also favour their interactions with solutes. This behaviour was predicted by various workers but is in disagreement with the experimental results. (ii) One could assume that the reducing poher of the H atoms is smaller than the oxidizing power of the OH radicals, i.e.the dccrease of free energy AC of a reaction such as is greater than AG of the reaction This might perhaps explain why some aromatic compounds are oxidized or hy- droxylated by X-rays, but no reduction is observed.12.13 The argument is, however, invalid in most inorganic systems. It can be calculated that, for example, The equilibria are rare for X-rays. Me(" - 1) I- -4- EI+ 4- OH = Me" i -+ H20 Me" + + H = Me(" - I)+ + H+. Fez+ + OH + H + = Fe3+ + H2O ; AG =- - 35 kcal/mole, Fe3+ + H = Fe2 i- + Hi- ; AG = - 65 kcal/mole. It is also difficult to understand why very unstable compounds such as pernitric acid are not decomposed by H atoms, as appears from Allen's experiments.14 (iii) It was also suggested 33 15 that radical ions such as ]HZ or OH i- may act as oxidants.We have no further data on the existencc of these ions in solution. The formation of H2f should be strongly favoured by the decreasc of pH (H + H+ = H2+). Radiochemical oxidations should then be more frequent in acid solutions. The same behaviour is to be expected if the determining oxida- tion agent is the OH+ ion, since in alkaline solution such ion would disappear rapidly by the reaction OH+ + OH-. These predictions are in disagreement with the experimental evidence. Ce3-, for example, is not oxidized in acid solution by cc- or X-rays, but Ce4+ is quantitatively reduced.7 The same is true for the Crvr/CrX1' couple.16 The oxidation rate of As02- increases 17 with the pH. The reduction of methylene blue increases with acidity 18 in the absence of 0 2 .On the other hand, the oxidation of KI diminishes2 with increase of pH; it is the same for the oxidation of phosphites."J Thus it seems that the effects of theM . HAISSINSKY 135 acidity are determined more probably by the chemical conditions created in the solution by its components than by the conditions of radical primary formation. We shall see on the other hand, that any hypothesis which invokes only the simple ionic origin of the free radicals is not sufficient to interpret the high ionic yields obtained for some radiochemical reactions. (iv) We only mention the very recent suggestion of Dee and Richards 20 that radiochemical transformations by a-rays are due to a secondary emission of photons. This hypothesis is based on experiments which as yet are not clear and not very convincing.Moreover, this would imply a completely different mechanism in the action of a-rays in solution and in gases, and also a different mechanism for the a- and X-rays. At the present moment, this does not seem justified. 2. REDUCTION BY OH RADrcALs.-Starting from the above considerations, we have suggested, with Lefort,6 that OH radicals cause not only oxidation, but also reduction. This latter could be carried out either directly when the substance has a high oxidation potential or indirectly by H atoms produced in the secondary reaction : H2 + OH = H20 4- H. We assumed at the same time that the inertia of the primary H atoms should be attributed to their rapid combination to molecular H2. We shall see that our suggestion may be modified on this last point.The study of the reduction of Ce(SO& in 0-8 N H2SO4 enabled 11s to 7 9 21 assume the reactions : (1) Ce4+ 4- OH + H20 L- Ce3t + H202 + H+ (2) 2Ce3t + H202 = 2Ce3+ -I- 0 2 + 2H+ ( 3 ) (The ions of CeIV in sulphuric acid solution are probably in the form of Ce(S04)~ and Ce(SO4)32-, but this complex formation 22 has an effect on the rate of reduc- tion, not on its mechanism). The reaction (2) is the reverse of that proposed by Haber and Weiss for the oxidation of ferrous ions, but the intermediate formation of H202 is not essential for our hypothesis. This formation became plausible for cc-rays from the kinetic data and for X-rays from the detection of small quantities of H202 under conditions where it is not produced in pure water. The mechanism of reduction by OH radicals was applied later by Evans and Uri 23 to the photo- oxidation of water in the presence of ceric ions.The reaction of ceric sulphate with X-rays (when 0 2 is initially absent) is accompanied by the formation of relatively high quantities of 0 2 (in units of number of molecules transformed by 100 eV, the yield G is 4.25) and still more of H2 (G = 6-5). If the reduction (G = 5.3) were due to H atoms the total number of H20 molecules decomposed by the radiation would then be too high (2 x 6.5 + 5.3 = 18.3) and probably energetically impossible, even if any recombination of the radicals is neglected. On the contrary the results can be explained if it is assumed that the reduction is produced by reactions (2) and ( 3 ) and is accompanied by a partial re-oxidation by OH radicals.For x-rays, H202, which is produced even in pure water, can account for the greater part of the reduction following ( 3 ) . A less detailed study of the reduction of Cr2072-, including gas analysis, led to similar conclusions : 16 (i) formation of great quantities of H2 when irradiated with X-rays ; or (ii) reduction by OH following the reactions : Cr2O72- i- 6 OH 4- 8 H' 2 Cr3+ 4 6 H202 + HzO Cr2O72 -t 6 0 H 1 - 8 H " - 2Cr3- , - 3 0 2 1 7H20136 TRANSFORMATIONS I N AQUEOUS SOLUTIONS (G) the reduction is accompanied by a partial reoxidation ; (iv) the possibility of accounting for the greater part of the results obtained (6) with x-rays by the primary formation of H202 : The kinetic relationships are much more complicated 17 for the system AS0a3- + As02-.With X-rays a radiochemical equilibrium is established, shifting in the direction of arsenate with an increase of pH ; no reduction occurs with x-rays. The yields of the reactions and of evolved gases are small. H202 is produced in small quanti- ties during the oxidation and the reduction (even in the absence of 0 2 ) . The primary production of H202 by cc-rays has still an essential role in the oxidation and probably prevents the reverse reaction. With X-rays the reduction is at- tributed to the following processes : (7) (8) (9) (10) (1 1) (1 2) It should be noted that we have assumed in this mechanism the intervention of AsIV compounds which have not so far been observed in aqueous solution. The formation of such unstable compounds of intermediate valency is justified, in particular by the necessity of avoiding elementary reactions of higher order when the difference between the initial and the final valencies is more than one.Thus in the reduction of Crvl to C P , the intermediate formation of CrV and C P has to be postulated. Other considerations led us 24 to suppose the formation of the radical €3-OH, (or B), during H202 production by x-rays in boric acid solution. The intervention of these intermediate unstable compounds and the reductions and oxidations accompanied by reverse reactions produce what one can call the catalytic destruction of OH and H02 radicals, a phenomenon analogous to the catalytic decomposition of H202. According to the chemical nature of the dissolved substance, the final result of these transformations would be an oxida- tion (AslIr), a reduction (Crvr, CeXv) or no change at all (H3B03).One can see that only accurate kinetic studies including the analysis of all the products of the solvent (H202, H2,02) can reveal the nature of the transformations and eventually establish the mechanism of the total reaction. Unfortunately the majority of radiochemists confine their research to the examination of only the chemical changes of the solutes ; correct interpretations are then difficult, especially with increasing complexity of irradiated substances. The whole of radiochemical and radiobiological work nevertheless confirms the conclusion concerning the predominant role of oxygenated radicaIs in radiolysis.Thus in the foIIowing organic and biochemical reactions only oxidative processes are reported without indication of reverse reactions or of intervention of H atoms : Cr2072- + 3 H202 + 8 Hf = 3 Cr3+ + 3 0 2 + 7 H20. AS043- + OH + H' = H202 -+ A~032- (JAsO~) As02 $- H202 = AsOZ- + H02 + Hf H0z + OH = H2O + 0 2 AsO2- + HO2 = As02 + H02- The last reaction is probably competing with other processes : As02 4- HOz == AsO2- + 0 2 f H' and especially, in acid solution, with the reaction : H2 + HO2 = H20 + OH. /O' /Om \OH N O (i) oxidation of various aliphatic aldehydes and acids by X-rays,25M. HAISSINSKY 137 (ii) oxidation and hydroxylation of aromatic compounds 129 13 and of sterols 26 (iii) transformations of carbohydrates such as starch, glucose, fructose, etc., (iv) depolymerization of nucleic acids by X-rays ; 29 (v) inactivation of sulphydryl enzymes 30 or oxidation of thiols by a-, p-, (vi) inactivation of ribonuclease by X-rays.323 33 The oxidations and the oxygenations in all these cases were proved either by direct determination of the reaction products or by carrying out the same reactions with OH radicals produced chemically or photochemically, or finally by a study of the protection effects.Only in a very limited number of radiobiological reactions do workers assume a reduction by H atoms. Thus Stein and Weiss,34 and Scholes and Weiss 35 consider that the de-amination of some amino acids, and the liberation of phos- phates of some nucleic acids irradiated by X-rays could be brought about also by H atoms. This conclusion is based on the variation of the ionic yield, if the gaseous environment in which the irradiation is performed, is changed (H2, 0 2 or vacuum).It is, however, necessary to note that in some cases interpretations due only to protection effects are not very certain, as competition can occur not only between the irradiated substances, but also between these and the intermediate and final products. Again, the conclusions are even more doubtful for complex biochemical molecules. Thus Forssberg explained 36 the protection effects exerted by various organic substances in the inactivation of catalase by X-rays through the reducing action of H atoms. In particular, the protection effect of pyruvic acid was attributed to its transformation into lactic acid. However, Bella has since then shown 37 that pyruvic acid does not undergo this transformation with X-rays, but that lactic acid is oxidized to pyruvic acid.In any case it should be emphasized that we do not exclude reduction by H atoms (see 5 3,4), we only point out the relatively small efficiency of these atoms. 3. DIRECT FORMATION OF H2; GENERAL FEATURES OF RADIOCHEMICAL TRANS- FoRMA-rIoNs.-We now need to investigate the causes of this apparent inertia of the H atoms. This can still be understood with a-rays where the production of the two species of radicals is very localized leading to rapid combinations into Hz and H202. Even if the localization of H atoms is less pronounced than that of OH radicals, the reaction H + H occurs very probably without activation energy, while for the combination OH J- OH an activation energy is necessary, which is evaluated as about 5 kcal/mole.On the other hand, once H atoms have combined to form H2, their chemical activity towards redox systems is prac- tically suppressed, while the main product of the OH combination, H202, is still an oxidant and, towards compounds of high oxidation potentials, a reductant. It is more difficult to understand the asymmetry H/OH for X- and y-rays, where the ionization is nearly homogeneous. The probability of collision between a H atom and a solute molecule should not be very different from the probability of collision between an OH radical and the solute. There is another serious difficulty when the ionic yields are examined. The primary formation of the free radicals in water is generally attributed to the decomposition of the ions of the solvent : (1 3) H20 7- e + (H20-) + H + OH-.(14) The energy necessary to produce by X-rays a pair of ions in air is 32.5 eV, a quan- tity which is presumably not very different from that necessary to ionize liquid by X-rays ; oxidation of hydroquinone ; 27 by a-rays ;23 y- or X-rays ; 31 H20f --> Hr + OH E138 TRANSFORMATIONS I N AQUEOUS SOLUTIONS water. The ionization potentials of water and of its components are of the order of 15 eV. There remains 15-20 eV for excitation processes. The latter should lead to a supplementary formation of free radicals, which is however frequently neglected by consideration of the " cage effect " of Franck and Rabinowitsch. Dale has, however, noted38 several cases where the yields correspond to 3 or 4 active radicals per ion pair.These cases are de-aminations of complex mole- cules where one could invoke the simultaneous utilization of W and OH radicals (so doubling the theoretical yield) or perhaps even a chain reaction. But this does not hold for the high yields of molecular hydrogen that we obtained 7 during the reduction by x-rays of CeIV and Crvl and the oxidation of Fe3+. Thus the contribution of the excitation phenomena appears excessive even if it is assumed that the kinetic energy possessed by the products of the decomposition of the solvent permits a partial escape from the cage. On the other hand, the differences between the radiochemical effects of corpuscular and electromagnetic radiations are usually treated by considering only ionizations.This implies that the spatial distribution is not very different for the radicals originating either by ionization or excitation. These considerations led Magat and myself to suggest 39 that molecular hydrogen is not produced by the combination €I -t H, but chiefly in a direct step after the electronic capture of the H2O molecule : H20 + c -+ H2 + 0-. (15) In spite of the rearrangement of the H20 molecule implied in this dissociation, the reaction would be favoured by the high electronic affinity of the 0 atom, the hydration energy of 0- and perhaps also by the particular conditions of the radiation field. The 0- will react rapidly with another molecule of the solvent : 0- + H20 = OH + OH-. (1 6) Keeping in mind the dissociation of the HzOf ion (reaction (13)) one can see that for each pair of ions we have now two OH radicals instead of one and one H2 molecule instead of one H atom.A direct formation of H2 could also be effected by the processes : (1 7) (1 8) (19) The final result of these reactions is the same as for the mechanism (15)-(lQ, but step (17) requires much more energy. Lefort proposed 40 a third possibility for H2 formation : H+ + H20 = H30t (20) H3O+ + e = H2 f OH. (21) Such a mechanism appears reasonable in strong acid solutions but is unlikely in neutral or alkaline medium. Whatever the exact mechanism of the direct H2 formation might be, thjs explains immediately the preponderance of OH reactions and accounts, at least in part, for the high yields of €41. It also enables one to understand the apparent inertia of radiochemically produced hydrogen, for example, towards the very un- stable pernitric acid presumably obtained by irradiation of HNO3.As Lefort noted,'+o our hypothesis accounts for the fact that Ha production is independent of the various factors which influence the decomposition of H202, when its aqueous solutions are irradiated by X-rays. As this author and, independently, Johnson have shown,41 the yield of this gas does not vary with the intensity of radiation, initial presence of 0 2 , addition of KT, etc., factors on which the decomposition depends. If H atoms were primarily produced, it would be difficult to understand This is not evident apriori. H20 + c + H- + OH H- + H20 = H2 + OH-, or in acid solution, H- + H+ = H2.M .HAISSINSKY 139 why they do not intervene in the chain of decomposition or why their effects are independent of the concentration of H202. While further experiments in radiation chemistry and probably in mass spectro- graphy are still necessary before one could estimate the relative importance of the various mechanisms, one can now have the following general scheme for radiochemical reactions produced in water and in dilute aqueous solutions. We will not discuss the contribution of excitation processes, as this question as yet is not at all clear. (A) %-RAYS; (i) pure wafer.-Strongly localized formation of H2 and of OH radicals. The greater part of the latter combine rapidly in situ to form H202. A certain number of radicals, probably not very great, escape from the tracks and react with H2 giving H atoms which produce a partial decomposition of H202.Another part is destroyed by OH following, for instance, the chain of Haber and Weisc : H202 -1 OH H2O -t ]KO2 (22) ti202 -t HO2 0 2 $- H20 + OH (23) OH + H02+ = H2O + 0 2 (termination). (24) It is not excluded that a part of OH radicals reacts to give directly 0 2 , i.e., OH + OH = H20 + 0 followed by 0 + 0 = 0 2 . The final measured quantity of H202 is thus less than equivalent number of decomposed H20 molecules.42 The gaseous phase contains H2 and 0 2 not in stoichiometric proportions but the amounts of H2 is equivalent to that of H202 + 1/2 0 2 . (ii) Aqueous solutions of apparently indiferent substances.-A great number of compounds do not undergo any observable transformations carried about by OH, H02 or H202 (at least under certain conditions).They can nevertheless react with them to yield unstable products, notably peroxides and peracids, break- ing down more or less rapidly into the initial compounds with evolution of 0 2 . These transformations influence the yield of H202 and of the gaseous products. We actually observed that the formation of hydrogen peroxide by x-rays of radon or by the corpuscular radiations of the nuclear reaction (B, n) and (Li, n) produced in the pile depends on the presence of " inert " compounds : e.g. H3B03, Li2SO4 and MgS04 diminish the yield. The influence of boric acid is attributed to re- actions of the types : H3BO3 + OH = H20 + H2BO3 (radical) H2BO3 + OH = H3B04 H3B04 -t- OH = H2BO3 + H20 + 0 (25) (26) (27) (28) (29) On the contrary the presence of HClO4 increases the yield; this case will be examined later.(iii) Aqueous solutions of reductants and oxidants.-The fate of these depends mostly, but not exclusively, on their behaviour with H202 : following the mode of action of H202 under the given conditions, the substance could be oxidized, reduced or unaltered, Radiochemical equilibria are not excluded but as yet not reported with certainty. Perhaps a steady state could be expected for CrV1/CI1I at some intermediate pH. The OH radicals can contribute to these transforma- tions depending on their peculiar nature, being kinetically more reactive and having a redox potential different from that of H202.It is necessary, furthermore, to keep in mind that the actions of these radicals on the solute are in competition with the reactions (l), (22) and (24). According to the measurements made up to date, it seems that these effects are of small importance. Thus the amount of gaseous H2 produced during the reduction of CeIV or the oxidation of As02- is not very different from that measured in the irradiated pure water. The amount H2BO3 $- H202 = H3B03 t HO2 H2BO3 + H2BO3 4- H20 = 2H3B03 + 0, etc.140 TRANSFORMATTONS I N AQUEOUS SOLUTIONS of 0 2 depends strongly on the nature of the transformations and of the path. In the kinetic analysis of the results one must also consider the reverse reactions with the catalytic decomposition of the oxygenated derivatives and the resulting consequences mentioned above.(B) X- AND 7-RAYS (FAST ELECTRONS) ; (i) pure water.-The primary radiolytic products are again Hz and OH, but the distribution is more or less uniform, the degree of uniformity depending on the energy of the incident photon. The reaction OH + OH : H202 js almost completely suppressed. But it is as yet not clear if H202 is produced at all (Weiss) 44 or if it is destroyed by a back reaction (Allen).14 The experimental evidence is rather in favour of the last supposition. In any case with strong intensities of radiation very small amounts of H202 can be detected.45~ 46 (ii) Dissolved " indiferent " substances.-The situation is analogous to that occurring with u-rays. The decomposition of unstable compounds resulting from the interaction of OH radicals leads to the formation of 0 2 . An equivalent amount of H2 is thus protected against its consumption in a back reaction.14 The pro- duction of H202 could also be increased.This could occur through a similar mechanism to that proposed above for the reduction of ceric salts with a back reaction giving the initial compound, but to date no example is known. How- ever, in preliminary experiments we observed the formation of small quantities of H202 in perchloric acid solution irradiated by X-rays ; a redox mechanism is here hardly likely (see below). Some radiochemists claim that KI is not oxidized by X-rays which only increases the yield of H202. Actually the main transformation here is the production of I2 47 and so KI is not an " indifferent " solute.(iii) Oxidants and rehctanfs.-The great availability of OH radicals here considerably increases the probability of the formation of H atoms (H2 + OH = H20 + H). Thus the situation can become very different, sometimes quali- tatively, from that with corpuscular radiations. Some reductions, which cannot be observed with the latter, are produced with X- or y-rays. If the probability of the reverse reaction is comparable with that of reduction, one will have a " radio- chemical equilibrium ". The kinetic study of the system As043-/AsO2- shows moreover that such an equilibrium can also result from the competition for the OH radicals between the oxidized and reduced forms. One can suppose that these radicals have from this point of view a less defined, more mobile, behaviour even than H202, which under given conditions more frequently orientates the whole reaction in one direction only.This peculiar feature of OH radicals probably plays an important role in the biological actions of ionizing radiations. There are indeed close similarities between the effects of radiations and of radiomimetic and carcinogenic substances, and it is known that hydroxylation is an essential step in the metabolism of these substances.48 It is, furthermore, to be borne in mind that the radiochemical equilibria depend not only on the reactivities towards the radicals of the oxidized and reduced forms, but also on the presence of H2 competing for OH radicals and of 0 2 if present, It is known that in the latter case H202 is produced.The action of the X-rays on another solute is then not much different than that of a-rays, keeping in mind the different competition conditions (H2 + OH ; 0 2 + H, etc.). On the other hand, the protection effects on oxidized substances can be exerted not only by other oxidizing agents competing for H atoms but also by reducible compounds competing with H2 for OH radicals. In some energetically favoured cases, one can have an increased formation of H202 even in the absence of 0 2 . This can, for instance, follow a mechanism similar to that described above for the reduction of CeIV and AsV. H202 can also be produced during oxidations either by simple electron transfer,M . HAISSINSKY 141 or following the mechanism suggested by Dainton : 49 1 + OH = HI0 (31) It is interesting to note that reaction (32) implies the idea of reduction by OH radicals, but this escaped notice up to now.4. KINETIC REMARKS.-we have seen that in spite of the identical primary production of the radicals with x- and X-rays, the radiochemical behaviour of the two species of radiation presents considerable differences. It seems there- fore difficult to agree with the assumption of Dainton and Collinson 3 that " activ- ated water " could have approximately the same " equivalent redox potential " (+ 0.95 V) for both species. The possibility of reduction of I 2 (Eh = 0.53) and of AsO$ (Eh = 0-56) by X-rays and not by x-rays demonstrates an essential differ- ence. Moreover, the concept of equivalent redox potential itself should be used very carefully.The reactivities of several redox systems such as AsO&/As02-, N03-/HN02, Cr2072-/Cr3+, etc., even near an electrode do not correspond at all to their reversible electrochemical potentials, as calculated by thermodynamical data and used in the evaluation of the equivalent potential. But even for re- versible potentials, these do not always express the chemical reactivity and still less in radiation chemistry, where the kinetics are determined by special competitive conditions. The following example, mentioned briefly by Dainton and Collinson, will illustrate our statement. Many dyestuffs are considered to be reduced by ionizing radiations in spite of their low redox potential. It is possible that in some cases where bleaching was taken as the only criterion for reduction, oxidation had really taken place.This could be true for example for methylene blue irradiated by X-rays, the bleaching being irreversible.50 Recently Collinson established 18 that this dye is oxidized in air by X-rays, while in the absence of 0 2 both oxida- tion and reduction occur, the latter increasing with acidity. Waterman and Limburg51 and Loisleur52 have also observed a negative shift of the potential by irradiating aqueous solution of this compound with X-rays (see also Day and Stein53). Furthermore, Seitz has irradiated57 a great number of dyes having potentials between i-0.217 and -0.525 V and he observed that, with the exception of the compounds of very low potential all the others were bleached more or less rapidly and regained their colour on passing air through the solution.No leuco- base has been modified by irradiation. The interpretation of these reactions is in any case a delicate matter, but one can think that the oxidation of these complex molecules by OH radicals requires high activation energy. The H2 molecules win therefore in the competition for these radicals giving H atoms. Those having very small size form an activated complex with the dye with small activation energy. This is one of the most interesting cases in radiation chemistry, the clarification of which is most urgently required. We may draw attention to some other obscure points in this field. We recall that Dale pointed out 38 certain anomalies in the trend of concentration-yield curves for de-aminations and other radio- biochemical reactions.The phenomenon of " changing quotient '' in protection effects, discovered by this author,55 is still also not quite understood. The high yields obtained by Dale and his co-workers for de-aminations and the high yields of H2 found in the above redox reactions are now clearer if our point of view on the direct formation of H2 and of two OH radicals per ion pair is adopted. But even if the yields are calculated following this hypothesis, it is necessary to consider an important additional formation of radicals by processes other than ionization in spite of the fact that at present radiation chemistry is based on the ionic origin of the free radicals.142 TRANSFORMATIONS I N AQUEOUS SOLUTIONS Another difficulty is encountered in the study of the oxidation of iodides and The kinetics follows the equation of a unimolecular reaction limited arsenites.by the reverse reaction : However, in contrast to the usual significance, the rate " constants " kl and k2 vary regularly with the initial concentration of the oxidized substance. An equa- tion of the same type was also established by Dale, Meredith and Tweedie 54 for the inactivation of carboxypeptidase and by Lefort 2 for the oxidation of fluor- escein. But in these cases one could theoretically account for the equation by supposing that the probability for the reaction product to react with radicals is the same as that for the initial substance (without yielding again this substance). Such a hypothesis for the oxidation of I- or AsO2- would lead to equal con- centrations of the oxidized and the reduced forms when in equilibrium, which is contrary to experimental evidence. It seems that the solutes can react with the radicals produced by radiolysis not only chemically but also in some other manner, e.g.by some deactivation process. The increase of the ionic yield of H202 in the presence of HC104 leads to the same supposition. In effect, it is very unlikely that this acid reacts with OH radicals by a reduction- oxidation mechanism or through formation of an intermediate per-perchloric acid with final production of H202. Rather it is necessary to assume that HC104 intervenes as a protector of H202 in the reactions involving the destruction of the latter, perhaps as a moderator of hot radicals which ensure the propagation of the decomposition chain? Finally, another interesting peculiarity was noted in the study 24 of H202 in boric acid solution when irradiating in the nuclear reactor of Chiltillon with a-rays (produced in the nuclear reaction B, n).Under some conditions of pH the amount of H202 is approximately proportional to the absorbed energy if the variation of this is brought about by changing the radiation time. But if the intensity is increased (by increase of the neutron flux or of the boron concentration) a limiting value is reached, which does not increase any further with intensity. Furthermore, initially added H202 is without apparent effect on the amount produced by radiation. The full understanding of these phenomena is far from satisfactory, but it seems that in the kinetic treatment of the results one cannot apply the usual stationary state method but one must consider, at least forsome. of the components participating in the reactions, the local concentrations around the ionization tracks and not the mean concentration in the bulk of the solutions These remarks are perhaps not closely inter-related but indicate nevertheles that the reactions brought about by ionizing radiations present peculiarities which do not follow immediately from the usual radiochemical conceptions. It is necessary, however, to recognize that it is these very conceptions which have given rise to the present development of radiation chemistry.The difficulties encountered appear thus to be only symptomatic of growing pains. 1 Dainton, Ann.Reports, 1948, 45, 5. 2 Lefort, J. Chim. Physique, 1950, 47, 624, 776. 3 Dainton and Collhson, Ann. Rev. Physic. Chcm., 1951, 2, 99. 4 Burton, J. Chem. Educ., 1951, 404. 5 Haissinsky and Lefort, Conpt. rend., 1949, 228, 314. 6 Haissinsky and Lefort, Compt. rend., 1950, 230, 1156. 7 Haissinsky, Lefort and Le Bail, J. Chirn. Physique, 195 I , 48, 209. 8 c f . Amphlett, Nature, 1950,165, 977. 9cf. Miller, J. Chem. Physics, 1950, 18, 79. 10 Nurnberger, Z. physik. Chem., 1934, 38,47. 11 Todd and Whitcher, A.E.C.U., 1949, no. 458. 12 Kailan, Monatsh., 1920,41, 312.143 M . HAISSINSKY Stein and Weiss, J. Cliewr. Soc., 1949, 3245 ; 1951, 3265. 14 Allen, J. Physic. Chem., 1948, 52, 479. I 5 Weiss, Nature, 1950, 165, 728. 16 Lefort, Radhakrishna and Haissinsky, J. Chim. Physique, 1951, 48, 488. 17 Haissinsky and Lefort, J. Chim. Physique, 1951, 48, 429. 18 Personal communication of Dr. Collinson. 19 Mrs. Vermeil, Cottin and Haissinsky, unpublished. 20 Dee and Richards, Nature, 1951, 168, 736. 21 Lefort and Haissinsky, J. Chim, Physique, 1951, 48, 368. 22 cf. Hardwick and Robertson, Can. J. Chem., 1951, 29, 828. 23 Evans and Uri, Nature, 1950, 166, 602. 24 Pucheault, Lefort and Maissinsky, J. Cliim. Physique (in press). 25 Lsisleur and Latarjet and Crovisier, Conipt. rend. biol., 1942, 136, 57. 26 Keller and Weiss, J . Chern. Soc., 1950, 2704. 27 Loisleur, Latarjet, Conipt. rend. biol., 1941, 135, 1534. 28 Khenoch, J . Gen. Ckm. Rim.. 1950, 20, 1560. 29 Lirnperos and Mosher, Amer. J . Roentg., 1950, 63, 681. 30 Barron, Dickinan, Muntz and Singer, J. Gen. Physiol., 1949, 32, 537. 31 Barron and Dickman, J. Gen. Pl?ysiol., 1949, 32, 595. 32 Holmes, Nature, 1950, 165, 266. 33 Collinson, Dainton and Holmes, Nafure, 1950, 165, 267. 34 Stein and Weiss, J . Cliem. Soc., 1949, 3256. 35 Scholes and Weiss, Nature, 1950, 166, 640. 36 Forssberg, Nature, 1947, 159, 309. 37 Bella, Bull. Soc. Ital. Bid. Speriin., 1949, 25, 1268. 38 Dale, J . Cliim. Physique, 1951, 48, 245. 39 Magat, Compt. rend., 1951, 233, 954. 40 Lefort, Compt. rend., 1951, 233, 1194. 41 Johnson, J. Cliein. Physics, 1951, 18, 1204. 42 cf. Lefort, J. Chim. Physique, 1951, 48, 339. 43 Haissinsky and Pucheault, J . Chim. Physique (in press). 44 Wejss, Nature, 1944, 153, 748. 45 Allen, M.D.D.C., 1947, no. 1056. 46 Toulis, 1950, U.C. R.L., 583. 37 (5. ref. (5). 48 cf. Boyland, J. Clrim. Physique, 1950, 47, 942. 49 Dainton, 1949, C.R.C., 304 (Chalk River). 50 Shakhtman, Krasnovski and Veretshinski, Compt. w i d . U.R.S.S., 1950, 74, 757. 51 Waterman and Limburg. Biochem. J., 1933, 263, 400. 52 Loisleur, Bull. SOC. Biol., 1943, 25, 22. 53 Day and Stein, Nucleonics, 1951, 8, 34. 54 Seitz, Strahlentherap., 1938, 61, 140. 55 Dale, Davies and Meredith, Brit. J. Cancer, 1949, 3, 31. 56 Dale, Meredith and Tweedie, Nature, 1943, 151, 281. 57 rf. Hamill, Williams, Schwarz and Voiland, Hot Radical Reactions (Univ. Notre Dame, 1951).

ISSN:0366-9033    

DOI:10.1039/DF9521200133

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

THE RADIATION CHEMISTRY OF FERROUS-FERRIC SYSTEMS PART 1. REACTIONS IN AIR-EQUILIBRATED SOLUTIONS BY C . B. AMPHLETT Chemistry Division, A. E. R. E., Harwell, Berks. Received 14th February, 1952 The oxidation of ferrous ions by X- and y-radiation has been studied in dilute solu- tions in &SO4 at different acidities. The kinetic scheme proposed to explain the resuits in 0.8N H2S04 has been found inadequate to explain the variation in initial oxidation yield with pH ; it is suggested that the primary act is more complex than is usually assumed. From a study of the steady-state condition at different acidities, values have been ob- tained for the “ equivalent redox potential ” of irradiated water with respect to the ferrous- ferric system. The effect of addition of other ions and complexing agents is briefly described.The oxidation of dilute aqueous solutions of ferrous ion in acid solution was first studied by Fricke and his co-workers,l who suggested it as a means of measur- ing the integral energy absorption in irradiated solutions. In 0-8 N H2S04 oxidation is virtually complete and ferric ion hydrolysis negligible, while the high yield makes it possible to measure moderately low doses with the sensitive analytical met hods available .2 The following mechanism has been proposed3 to describe the oxidation in the initial stages in sufficiently acid media (> 10-1 N). H20 y2- H + OH Fe2f + OH --f Fe3+ + OH- Fe2f + €402 --f Fe3+ + H02- H02- + H + Z-f H202 Fez+ + H202 -+ Fe3f + OH- +- OH HA- + OH- ~2 H2O. The maximum theoretical ratio between the oxidation yields in aerated and evacuated solutions is Ii = 4/1, whereas in practice a figure of about 2.7 is ob- tained.(The value of 4 reported by Krenz and Dewhurst 3 has since been traced to contamination of the evacuated samples, possibly with mercury vapour, leading to a lowered evacuated yield (Dewhurst, private communication).) Hart has suggested 4 that a second “ primary ” step be added, viz., This represents the increased probability of pairwise combination of like radicals in regions of dense ionization (e.g. along a-particle tracks, or near the end of /3-particle tracks), where the local radical concentrations are high. It had pre- viously been postulated by Allen 5 to explain production of H202 in “ hot spots ” in irradiated water.It should not be confused with similar combination re- actions in the bulk of the solution ; the results of ferrous 4- H2Oz kinetic studies 144C . B. AMPHLETT 145 at niuch higher rates of radical production than those in this work show that the OH -t OH reaction can be neglected in comparison with step (l), while the linearity of yield up to complete oxygen depletion in irradiation experiments shows that the H + H reaction is negligible compared with step (2). By adding (0’) to the above sequence, and assigning to the two steps probabilities x and (1 - x), the aerated yield is decreased without affecting the evacuated yield, thus reducing R to less than four. By considering the ratios experimentally obtained, Hart finds that the importance of (0’) increases with increasing ion-density (e.g.from Co y-radiation to T3 P-radiation), in accordance with expectation. It is found, however, that the ratio of x/(l - x) appears to depend upon the solute, being appreciably different for dilute formic acid solutions.6 This would suggest that Hart’s treatment is an approximation, which it is convenient to employ in our present state of knowledge concerning the primary act. In 0.8 N H2S04 the oxidation yield is found to be independent of initial ferrous ion concentration over a wide range, viz., 10-4 to 10-1 M.7 The decrease in yield below 10-4 M is attributed to the relative slowness of step (4) under these con- ditions. If sufficient time is allowed between irradiation and analysis, normal yields are obtained,g and it has been shown 9 that the rate constant for the increase in yield on standing is in agreement with that for step (4).In 0.8 N H2S04 the yield is constant almost to complete oxidation, and oxidation is virtually complete, but as the pH is increased the initial yield eventu- ally decreases and evidence of a back reaction appears; oxidation is no longer complete, and the oxidation curve bends over to a steady state with respect to ferrous and ferric ion. In addition, reduction of aqueous ferric solutions, which is barely detectable in strong acid, is observed to an increasing extent. These effects may be attributed to the reduction of ferric ions, which are known to com- pete for HOa radicals in the pH-dependent step Baxendale and his co-workers have shown10 in a study of the Fez+ -I- Hz02 reaction, which has many features in common with the present system, that the ratio kS/k3 (k4/k3 in their terminology) is pH dependent, and that the reacting partners in step ( 5 ) are probably Fe3f and 02-.The possibility of ferric ion being reduced by the step Fe3+ +- HO2- + Fez+ + H02, i.e. the reverse of (3), is ruled out by its extremely low rate compared with the ferrous oxidation steps, so that it would only become appreciable at very high ratios of ferric to ferrous ion. Little work has hitherto been reported on the irradiation reaction at low acidities,ll possibly because of the difficulties introduced by hydrolysis of ferric ion; it has been shown 12 that as the pH is increased a progressive lowering of the steady-state [Fe3+]/[Fe2+] ratio is obtained.It is the purpose of this paper to discuss the effect of pH in air-equilibrated solutions, and also to detail some of the effects observed in the presence of complexing agents and added ions. Fe3’ . j HOz -tFe2+ -i- Hi 4- 0 2 . ; . ( 5 ) EXPERIMENTAL MATERIALS.-h the original experiments 12 perchlorates and perchloric acid were used to avoid complexing other than that arising from hydrolysis. It was found, how- ever, that below pH 2 the yields were appreciably higher than in sulphate solutions ; in addition chlorate (but not chloride) was formed. Weiss has since found13 a similar phenomenon in air-free solutions, together with the appearance of chloride. It appears that perchlorate ion must participate in the reaction, and sulphate solutions have since been used, despite their great complexing power.Nitrate ion was excluded, because of possible complications arising out of its reduction to nitrite ; 14 chloride ion is known to compete for OH radicals.15 A.R. ferrous ammonium sulphate was twice recrystallized from 0.8 N HzSO4 and dried in air ; all other materials were A.R., and were not purified further. Distilled water was redistilled before use, once from 0.01 M KMn04 + 0-01 M KOH, once from 0.01 M KHS04 and finally alone. All distillations were from silica apparatus. and the final146 FERROUS-FERRIC SYSTEMS water was stored in stoppered silica flasks. This water was used for all irradiated sohi- tions, and for the final rinsing of apparatus. (i) 400 mc Ra y-source in Pt capsule. IRRADIATIONS.-Three sources have been used, viz.: Samples were irradiated in Pyrex vessels mounted coaxially with respect to the source. Dose rate 33 r/min (1.9 x 1015 eV/cm3 min). Samples irradiated either in polystyrene cell containing 10 ml solution (208 r/min ; 1.22 x 1016 eV/cm3 min) or in 2 cm3 glass cells mounted in a polystyrene block (338 r/niin ; 1.99 x 1016 eV/cm3 min). 15 cm3 samples were irradiated in thin-windowed Pyrex cells mounted close to the shutter. Dose rate 1420 r/min (8.4 x 1016 eV/cm3 min). All dose rates were measured by dosimetry with 0-8 N HzSO4 solutions of ferrous ion (2 x 10-4 M to 2 x 10-3 M), using a value of G = 20.6 ions oxidized per 100 eV.2 ANALYsIS.-After irradiation the solutions were analyzed for Fez+ by complexing with a-phenanthroline at pH 4-5 in an acetate buffer,16 measuring the optical density at 510 mp with a Hilger Uvispeck spectrophotometer.The calibration curve for ferrous ion gave a value of eM = 10,660 for Feph$+ under these conditions (ph E a-phehanthrol- ine). This is low compared with the accepted value of 11,000-11,100, but was consistently obtained with this instrument and the particular cells used. The optical density of a ferric solution (free from Fez+) containing o-phenanthroline and acetate corresponded to eM -20, so that the contribution of ferric ion species towards the optical density is negligible. pH values were measured with a glass electrode and Beckman pH meter. (ii) 3.25 curie Co y-source in stainless steel container. (iii) 240 kVp X-rays, unfiltered, at 8 mA, from a Victor Maximar set.1. THE INITIAL OXIDATION YIELD AS A FUNCTION OF ACIDJTY RESULTS Table 1 shows the results obtained in H2SO4 solutions from 0-80 N to 1.6 x '10-4 N in H+ ion; table 2 gives for comparison some earlier results in perchlorate solutions. The absolute accuracy of the sulphate figures is probably :t 5 % ; the perchlorate results are less accurate, as they were obtained before conditions had been thoroughly standardized. TABLE ~.--JNITIAL OXIDATION YIELD OF FERROUS ION IN H2SO4, IRRADIATED WITH [Fe2+]o = 1-5-2.0 x 10-4 M. Dose rate 33-1420 rjmin =1-9 x 101s - 8.4 x 1016 eV/cmJ min. x- AND y-RADIATION [H2S04], moles/l. PH [H+l moIes/I. Go = no. of ions oxidized per 100 eV 0.4 (by titration) 0.39 (calc.) 0.41 20-6 0.05 ,, 1.21 6.0 X 10-2 20.6 -2.0 x 10-2 1.60 2.51 x 10-2 20.6 -1.0 x 10-3 2.65 2-24 x 10-3 16-2 -3.6 x 10-4 3-10 7.9 x 10-4 13.9 -1.3 x 10-4 3.518 2.6 x 10-4 10.3 -9.0 x 10-5 3.80 1.6 x 10-4 8.4 1.0 TABLE 2.-INITIAL OXIDATION YIELDS IN PERCHLORATE SOLUTIONS 150-180 kVp X-rays; dose rate 270-370 r/min =- 1.59-2.18 :: 1016 eV/cm3 min.PH [H+], moles/l. Gn 0.1 (calc.) 0.80 41-2 1 -05 9-8 x 10-2 29.0 2-06 8.7 x 10-3 18.1 2-20 6.3 18.1 3-00 1.0 14-5 3-10 8.0 x 10-4 12.5 3-64 2.3 12-5C . B . AMPHLETT 147 DISCUSSION The variation in Go with pH is shown in fig. 1, which also shows results ob- tained by other workers. It is not proposed to consider the perchlorate results beyond remarking that above pH 2 there appears to be no difference between perchlorate and sulphate solutions ; it may reasonably be conjectured that in this region we are dealing principally with uncomplexed ferrous ion and the ion FeOH2+ in both cases.The kinetics outlined in the introduction give a qualitative explanation of the oxidation process, and in strongly acid solutions conform to the observed stoichio- metry. If we assume that the distribution of atoms and radicals arising from y-irradiation is homogeneous, we may apply steady state kinetics to obtain an FIG. 1.-Variation in initial oxidation yield with pH. 0 SO$- solutions (this work) 6 ,, ,, (Krenz and Dewhurst, unpublished results) A 3 , ,, ; 10-3 N H2S04, 0.3 M Na2S04 El c104- ,, expression for the yield. From the calculated steady state concentrations of H, OH, H02 and H202, the overall expression for the yield of ferrous ions oxidized per 100 eV may be calculated to be where D is the dose rate in units of 100 eV per unit time, ko is the net radical pair yield per 100 eV (allowing for the back-reaction in the primary step) and x is the fraction of radicals available for reaction with substrate in the bulk of the solution.The first term in parentheses represents the stoichiometry 1 OH = 1 Fe2+, while the second represents competition between ferrous and148 FERROUS-FERRIC SYSTEMS ferric ion for the H02 radicals, each of which is capable of oxidizing three ferrous ions. In the initial stages of the reaction, when ks[Fe34] < k3[Fe2+], this reduces to the form (2) In evacuated solutions Go' = ko, and hence x may be evaluated. Using a value of R = Go/Go' = 2-7-2-86, we obtain x = 0.57-0.62, say 0.6.This implies a net radical pair production of slightly more than seven per 100 eV, corres- ponding to an energy requirement of about 13 eV per radical pair in liquid water. Although the initial portion of the oxidation curve remains linear as the pH is increased, the initial yield decreases above pH 2. The form of the curve re- lating Go to pH is different from that given by Fricke and Hart,ll the reason for this discrepancy not being obvious. On the basis of the above mechanism, the decrease in yield must be due to a decrease in concentration of one or more of the reactants concerned, viz. Fez+, OH, HO;! and H202. Since the decrease is apparent from the start of the irradiation, and since the curvature due to back reactions involving ferric ion is not apparent until much later, we cannot invoke ferric ion effects to explain the decrease in initial yield.The hydrolysis of ferrous ion is negligible at these acidities ( K h = 1.2 x 10-6 17), and it would appear that sulphate complexing has no effect; the effect of adding excess neutral sulphate to a solution at pH 3 is to depress the yield to a value corresponding to the increase in pH. It appears unlikely that the variation in initial yield can be attributed to differences in reactivity between hydrated ferrous ions and possible complexes. The results of other workers on the ferrous + H202 reaction 10 show that kq, deter- mined in HC104 and H2SO4 solutions under conditions where steps (1) and (4) are predominant, does not vary over the range N/2 acid to pH 2.65 ; this confirms the conclusions drawn from the present work.The oxidizing power of the irradiated solvent must therefore vary, and there are several alternatives. The dissociation constant of H202 is so low (Kd = 1-8 x 10-12 18) that we can eliminate the possibility of step (4) being pH dependent at moderate acidities. Similarly, the dissociation of OH into 0- and H+, followed by 0- + H --f OH- would appear unlikely; 19 both these steps are, of course, eliminated also by the results of the Fe2f + H202 studies. Weiss has suggested 20 that in air-free solutions the following pH dependent steps may operate : Go =: ko(1 + 3 4 . H + H+ z-f H2+ H2f + Fez+ -+ H2 + Fe3+. In the presence of 0 2 the equilibrium would be expected to be well on the side of H atoms, and such a mechanism in aerated solutions would actually lead to a decrease in yield with increasing acidity. The participation of OH+ ions,21 and their removal by OH+ + OH- -+ H202 will have no effect in the ferrous system, since both OH+ and H202 are equivalent to two Fe2+ ions.This leaves only H02 as a source of pH dependence. It has been pointed out 10 that whereas H02 behaves as an oxidizing agent towards ferrous ions, its anion 0 2 - will reduce ferric ions, so that with decreasing acidity the net oxidizing power should be reduced. K d for the dissociation HO2 z-f H+ + 0 2 - has been estimated as 10-3, with a possible error of about one power of ten in the value.22 If we modify eqn. (2) to allow for the fraction of H02 radicals dissociated, and therefore whenceC .B . AMPHLETT 149 Gmax, the maximum yield when Kh < [H+J will be given by (2), so that Substituting for x, Kd and Gmax, we obtain a curve relating Go and pH of the form shown in fig. 1. It will be seen that, although the decrease in yield is found to be in the correct acidity range, the forms of the theoretical and experimental curves are in fact appreciably different. Thus, while the dissociation of HO2 is probably contributory, it cannot be the only cause. The experimental values between pH 2 and 4 were found to fit the empirical relationship which is also plotted in fig. 1 ; below pH 2 the values deviate appreciably from this curve. The complexity of this expression suggests a complicated overall pH dependence, with possibly several factors operating.It should be remembered that ko and x may not be independent of acidity, as has been assumed in this treatment, and that the primary act of radical production may itself be much more complex than has hitherto been assumed. Lefort has suggested23 that an addi- tional source of primary radicals may arise from the pH dependent sequence H30f + e --f H2 4- OH (6) H2 -k OH -+ HzO + H. (7) If these steps be added to the primary step (0), it is no longer necessary to assume the production of equal numbers of H atoms and OH radicals, and the production of both species is pH dependent. However, such a step will introduce a pH dependence in opposite direction from that found in ferrous solutions, since increasing acidity will favour capture of the slow electrons in step (6) rather than by the process H20 + e -+ H20- -+ H + OH-, thus leading (in aerated solu- tions) to a decrease in oxidizing power.It seems therefore that a more complete understanding of the primary act in irradiated water is necessary for interpretation of the overall pH dependence ; further work is proceeding with evacuated solutions in an attempt to obtain more information. 2. THE STEADY STATE AND THE EQUIVALENT REDOX POTENTIAL OF IRRADIATED WATER On the above mechanism the value of G at any point on the oxidation curve is given by the expression (1). From the study of the Fe2f 4- H202 reaction, values of k5/k3 may be calculated ; 24 they range from 1-25 at pH 4 to about 0.02 in 0.8 N acid at 25" C. It is found in irradiation experiments that the point at which the yield begins to deviate from its initial value Go approaches more closely to the origin as the pH is increased, due to steadily increasing values of k5/k3.In 0.8 N H2SO4 it is so near to 100 % oxidation as to be frequently missed ; if, however, a large excess of ferric ion is initially present, thus increasing [Fe3+]/[Fe2+], the deviation from linearity is easily observed, even though the initial yield is unchanged.8 An attempt to apply expression (I) quantitatively fails for two reasons : (i) we should expect a continuously decreasing yield as oxidation proceeds, at acidities where k5/k3 is appreciable. As fig. 2 shows, the initial yield is constant for an appreciable part of the oxidation and this is still true at pH 3.8 ; (ii) the expression (1) will never yield a steady state value of G = 0, except at almost complete oxidation, for positive values of kS/k3 and of x.This suggests the introduction of some other reducing step such as (8) Fe3+ -I- H -+ Fez+ + H 1 .150 FERROUS-FERRIC SYSTEMS This will itself presumably be pH dependent, and over practically the whole range of acidities used in these experiments the contribution of FeOH2+ cannot be neglected. The substitution of step (8) into the original kinetic scheme does not, however, solve the difficulty, since we obtain negative values of ka[Fe3+]/k2[021 on substituting for kS/k3 and [Fe3+]/[Fe2'] when G = 0. It is clear, therefore, that a fundamental inadequacy of the kinetics has been exposed. Although a complete elucidation of the kinetics is not possible at present, the results obtained for the steady state ratio of ferric to ferrous ion may be used to calculate the " equivalent redox potential " (e.r.p.) of irradiated water.The concept of e.r.p. has been fully discussed by Collinson and Dainton,25 who con- clude from a survey of results on the irradiation of aqueous solutions of oxidizing and reducing agents that for X- and y-irradiation e.r.p. (H20) is somewhat morc positive than - I-OV (U.S. convention). A few examples of steady state con- &ions attained in irradiations are known, e.g. the systems I NO. --NO3 ,2b FIG 2.-Oxidation of ferroug solutions to steady state conditions at pH 391. 240 kVp X-rays, 1400 r/min Co y-rays (1.2 mV) 208 r/min Ra 'lV mean)7 33 r/min x 1-46 x 10-4 M Fe2+ 0 1.46 X 10-4 M Fez i- 0 1-48 ,, 6 1-54 ,, 7, t 1.52 ,, 6 1.66 ,, 7 , If 1-72 ,, Q 1.72 y y Y 7 0 1-95 ,, 7 7 -0- 1.82 ,, , I 0 2.04 ), 7 , The figures to the right of the curves refer to the percentage oxidation at the steady state.The horizontal scale has been considerably compressed above 10 x 1017 eV/cm3. T2-I-,26 arsenite--arsenate,29 and Fe*+-Fe3+ ; in these cases the steady state may be approached from either side, and no effect is observed if one irradiates a solution of composition corresponding to the steady state. From the steady state value of [N03-]/[N02-] = 1.5 quoted by Lefort26 and from the known EO value for this system (- 0.94 V 27 in acid solution, corrected to - 0-84 V at pH 6), we may calculate a value of - 0.49 V for the e.r.p. of X-irradiated water assuming the overall reaction to be H20 + NO2- ,f NO3- -I- 2HI + 2e.2-04 x 10-4 M Fe2tC. B . A‘MHPLETT 151 RESULTS We have measured steady state values in the Fez+ + Fe3 1- system in H2SO4 at different acidities, both for ferrous oxidation and for ferric reduction. Fig. 2 shows the results at pH 3.1 for X- and y-radiation over a forty-fold variation in dose rate. No variation either in initial yield or in steady state conditions was observed over this range. In order to calculate values of e.r.p. (H20) we have measured Eo’ values for the ferrous + ferric system as a function of pH from 0 to 3-5, using a Cambridge millivoltmeter and a saturated calomel electrode. The variation is complex, and is being studied further. At high acidity our results are in agreement with other workers,28 but no other results are known to us for lower acidities. Using these values of Eo’ we have calculated e.r.p.(HzO) from the steady state ratios by means of the Nernst formula e.r.p. (H20) = Eo’ - 0.059 loglo [Fe3 ‘]/[Fe2+]. We use the symbol Eo’ to denote the observed redox potential relative to the hydrogen electrode, under the conditions of measurement, when [FelI~l = [Fe“] ; this should not be confused with the standard thermodynamic potential EO- Table 3 summarizes the results so obtained ; the absolute values of e.r.p. (H20) are probably correct to 1 5 mV. TABLE 3 .----EQUIVALENT REDOX Po rEmiAL OF IRRADIATED WATER IN 7 HE system PI1 Fez+ 2.12- oxidation 2.1 5 247 3.10- 3.1 5 Fe3+ 2.12 reduction 3.22 Fe3 + reduction 2.08 0 3 2 - saturated) radiation Y Y Y X x Y X FERROUS-FERRIC SYSTEM [ Fe3+]/ [Fe2+] at steady state Eo’ 6.20 --0*704 V 6.30 5.30 3.28 -0.708 3-10 3-1 1 3.67 -0.703 3.25 4-14 3-06 4.22 3-76 7.1-9.3 -0.704 3.28 -- 0.702 4.94 -0.703 e.r.p.(H20) mean e.r.p. (H20) - 0.750 V 0.75 1 - 0.749 &0*002 V 0.747 - 0.738 0.737 0.737 -0*737-J:0.001 - 0.736 0.733 0.739 - 0.736 1 0.004 0.732 0.740 0-737 -0.754 to -0*757+0.003 - 0.760 -0.732 - 0.732 - 0.744 -- 0.744 J l L % C J L S $ u T r J As the pH is increased, the value of e.r.p. (H20) becomes more positive, in conformity with the general trend of redox potentials with respect to increasing alkalinity; the change in oxidizing power thus parallels the change in initial oxidation yield. The presence of H2, which is assumed to remove OH radicals by the reaction H2 + OH -+ H20 + H, produces only a slight change in potential a t pH 2. We can calculate values of e.r.p.(H2O) for 0.8 N and 0.1 N acid, putting [Fe3+]/[Fe2+] ’v 100 ; we then obtain values of N - 0-806 V in 0.8 N acid, and N - - 0.799 in 0.1 N acid, showing an appreciable decrease in e.r.p. to more negative values below pH 2. These values differ appreciably from that calculated from the results of Lefort’s work on the system N02--N03-; this may in part be due to differences in pH,152 FERROUS-FERRIC SYSTEMS but may also be due to the different substrates present in the two cases. The measured e.r.p. of water relative to any dissolved material will depend upon the competition for the oxidizing and reducing species formed between the substrate (in both oxidized and reduced forms), the solvent and the radicals themselves; in the presence of dissolved oxygen the situation will be still more complicated.It is therefore conceivable that the precise value of e.r.p. will depend upon the particular substrate chosen, as well as upon other external conditions, such as acidity, temperature, presence of oxygen, hydrogen, etc. Thus, although in principle we can conceive the idea of irradiated water possessing a redox potential, it may not be possible to measure it precisely by experiments of this kind. All the possible reactions involved in the oxidation and reduction of ferrous and ferric ions in aqueous solution are appreciably exothermic, with one ex- ception ; their AH values, calculated from cycles involving the entities concerned, are given below: --f Fez+ + H-+ ; - 63.6 ,, (iii) - 16.0 ,, (vi> The reduction of hydrated ferric ions by hydrogen atoms is energetically favourable, but in view of the pH dependence it seems more likely that reduction involves the hydrolyzed species FeOI-l2+.Work on the photo-decomposition of ceric solutions 30 and on the radiation induced reduction of acid ceric solutions 31 suggests that the principal reduction step is CeOH3t t OH + Ce3+ + H202. This is energetically more favourable (AH = - 20.1 kcal30) than (v) above ; in addition, the peroxide produced reduces more ceric ion, whereas in the present system it would oxidize ferrous ion and so produce no net reduction. The precise function of H atoms in producing reduction in aqueous solutions seems at present to be doubtful; although it appears likely that they participate in the ferric system, recombination to molecular hydrogen seems to predominate in the reduc- tion of ceric solutions.32 It is hoped to obtain further information by studying the H202 yield and gas production in evacuated systems.The apparent anomaly in the effects of acidity upon the FeZ’-Fe31- and Ce3+-Ce4+ systems disappears when viewed in the light of the potentials in- voked. In acid solutions, oxidation of ferrous solutions is predominant, while oxidation of cerous solutions is only observed in alkaline solutions.31 The above measurements show that e.r.p. (H2O) becomes slightly more positive with in- creasing pH, while the potential of the CelI1/CetV couple changes from - 1.44 in acid solution to + 0.77 V over the range of conditions employed by Lefort.31 Consequently the small change in e.r.p.is overtaken by the change in Eo’. There are no results for ceric solutions over the range of acidity used in the present work, owing to difficulties introduced by the hydrolysis of ceric ion. Under the con- ditions employed for cerous oxidation, viz. in solution in 3 M K2CO3, the cerous ion will be largely present as anionic complexes and under similar conditions it is probable that ferrous solutions would again be oxidized ( E l = + 0.56 V). It appears from the present work that although oxidation of hydrated ferrous ions is easily affected, reduction is greatly favoured by hydrolysis of ferric ion; this may be due to a greater ease of electron-transfer between the reducing species and the central ion when the hydration shell is broken, either by hydrolysis or by suitable complexing.Silverman and Dodson 33 have reported that decrease in acidity favours isotopic exchange between ferrous and ferric ions in perchlorate solutions, and attribute this to a greater facility for electron-transfer in the inter-C. B . AMPHLETT 153 mediate state Fe . . . OH- . . . Fe compared with that in Fe . . . H20 . . . Fe. The experiments on I‘ optical interaction absorption ” in mixed ferrou?+ ferric systems 34 are also taken to indicate increased probability of electron-transfer when the hydration shells are broken. It is known that ceric solutions are much more strongly hydrolyzed than the corresponding ferric solutions (92 % 30 in N HC104, compared with 0.5 % 3 9 , while a comparison of the respective EO values in HC104 and H2SO4 solutions27928 suggests that ceric ions are much more strongly complexed. Further information should be sought on the effects of pH on the irradiation of redox systems, since it may lead to knowledge con- cerning the reactivities of ions and complexes in solution.3. THE EFFECT OF ADDED MATERIALS ON THE YIELD AND STEADY STATE VALUES FLUORIDE ioN.-Fluoride (0.01 M) added to solutions with a final pH of 2.8 and 4.2, produced practically complete oxidation, with no evidence of a ferric back-reaction ; complete oxidation was still observed in 0.8 N HzSO4 solutions. This agrees with the change in redox potential when fluoride is added, e.g.addition of F- to a 0.1 N HC1 solution displaces Eo’ from - 0.771 V to -- 0.553 V at N/70 F- ion concentration.36 Addition of fluoride also affects the initial yield in a com- plex manner. Thiocyanate, which also complexes strongly with ferric ion, produces similar results to fluoride at high pH. CHLORIDE ioN.-Chloride ion had little effect on the steady state ratios. Contrary to the results reported by Dewhurst,37 a marked effect was found upon the initial yield in low chloride concentrations, the diminution in yield increasing with increasing C1- ion concentration ; values obtained from several runs were Go = 17.2 (lO--3 M Cl-), 15.2 (10-2 M Cl-), 9.1 (M Cl-), all in 0.8 N H2SO4, and 8.3 (10-2 M and 1 M) at pH 3.1. The reason for the discrepancy is not obvious, but it is proposed to examine it further in air-free solutions. A~RYLoNITR~LE.-A ddition of many reactive organic compounds has been shown to increase the ferrous ion consumption both in the irradiation reaction 37 and the Fenton reaction.38 On the other hand, addition of acrylonitrile (1 M) to a 0.8 N H2SO4 produced a decrease in yield (Go -x 13.8) without affecting the overall oxidation ; the latter point is to be expected, since only a 1 mV change in Eo’ is produced.39 Acetonitrile (1 M) produced a similar effect (Go = 12.2).In these cases the effect may be due to competitive attack by the radicals upon the unsaturated structures present, presumably producing organic radicals which are not capable of peroxidizing and so causing an induced oxidation of ferrous ion.PHENANTHROLINE AND DIPYRIDYL-The most striking effect was observed in the case of the o-phenanthroline and dipyridyl complexes, where complete re- duction of the ferric complex was observed from pH 1 to 4, with no oxidation of the ferrous complex.12 These form a reversible redox system (EO = - 1-10 V). At low pH the initial yield is quite high (Go _N 12 at pH 1.0), suggesting great ease of electron transfer for such bulky ions; this is supported by the results of radioactive exchange experiments.40 At pH 2 to 3, GO = 4 to 5, similar to the value found at pH 6 in the absence of air ; 31 we also find that the yield in this region is unaffected by 0 2 , N2 or H2. It was found during the original investiga- tions that the ferric complex was also reduced by H202; 41 the kinetics of this reduction, as well as those of the irradiation reaction, are being further studied.The author wishes to thank Sir John Cockcroft, F.R.S., for permission to publish this paper; his grateful thanks are due to Mrs. M. 0. Small for con- siderable assistance with the experimental work, to Dr. W. Wild for discussion of the results, and to the M.R.C. unit at Harwell for the loan of facilities for X-irradiation.154 PEKKOUS-FERRIC SYSTEMS NOTE ADDED IN PRooF.-Since writing my paper, further evidence has appeared concernjng the relative contributions of steps (0) and (0‘) in the mechanism dis- cussed. Measurement of the moiecular Hz yield in aerated ferrous solutions in 0.8N H2SO4 suggests 42 that x := 0.78, and not 0.62 as was estimated from Hart’s earlier figures.This assumes that thc yield of H202 in the “hot spots ” equals the molecular H2 yield, which is supported by the measurements of Sutton 43 on the post-irradiation reaction in very dilute (< 10-4 M) deaerated ferrous solu- tions. On substituting the above value of x the expression (4) is altered slightly, the factor 3x/(1 + 3x) changing from 0.645 to 0.70 ; this will not, however, greatly affect the position of this curbe in fig. 1. The discrepancy between this value of x and the earlier one derived by Hart is readily understood if we consider the mechan- ism for the oxidation of deaeratcd solutions of ferrous ion.44 The participation of H2-‘ in the latter case, but not in aerated solutions, implies that calculations of x based solely on air/vacuum yield ratios may be considerably in error, par- ticularly since the ratio is dependent upon ferrous ion concentration.1 Fricke, Physic. Review, 1928, 31, 1117. 2 Miller, J. Chem. Physics, 1950, 18, 79 ; also this Discussion. 3 Krenz and Dewhurst, J . Chem. Physics, 1949, 17, 1337. 4Hart, J. Amer. Chem. SOC., 1951, 73, 1891. 5 Allen, J . Physic. Chem., 1947, 52, 489. 6 Hart, ANL-4434 (U.S. Atomic Energy Commission, declassified document). 7 Miller, ref. (2). Todd and Whitcher AECU-458 (US. Atomic Energy Commission, declassified document). 8 Krenz, unpublished work. 9 Dainton and Sutton, unpublished work. Fricke and Morse, Phil. Mag., 1929, 7, 129. 10 Barb, Baxendale, George and Hargrave, Trans. Furatlay Soc., 195 1, 47, 462. 11 Fricke and Hart, J. Chem. Physics, 1935, 3, 60. 12 Amphlett, Nature, 1950, 165, 977. 13 Weiss, private communication. 14 Clark and Pickett, J. Amer. Chem. SOC., 1930, 52, 465. 15 Taube and Bray, J. Amer. Chem. Soc., 1940, 62, 3357. 16 Fortune and Mellon, Ind. Errg. Chem. (Anal.), 1938, 10, 60. 17 Lindstrand, Svensk Kem. Ticlskr., 1944, 56, 282. 18 Evans and Uri, Trans. Faraday Soc., 1949, 45, 224. 19 Gordon, Hart and Walsh, AECU-1534 (U.S. Atomic Encrgy Commission, de- 20 Weiss, Nature, 1950, 165, 728. 21 Collinson and Dainton, Anir. Rev. Physic. Clzem., 1951, 2, 112. 22 Evans, Hush and Uri, Quart. Rev. (in course of publication). 23 Lefort, Compt. reiid., 1951, 233, 1194. 24 ref. (lo), p. 485. 25 ref. (21), pp. 107-114. 26 Lefort, J . Chim. Phys., 1950, 47, 776. 27 Latimer, Oxidation States of the Elements niid their Potentials in Aqueous Solution (Prentice-Hall, New York, 1938). 28Smith, Anal. Chem., 1951, 23, 925. 29 Haissinsky and Lefort, J. Chim. Phys., 195 1 , 48, 429. 30 Evans and Uri, Nature, 1950, 166, 602. 31 Haissinsky, Lefort and Lebail, J. Chirn. Phys., 1951, 48, 208. 32 Lefort and Haissinsky, J. Chim. Phys., 1951, 48, 368. 33 Silverman and Dodson, BNL-82 (U.S. Atomic Energy Commission, declassified 34 McConnell and Davidson, J. Amer. Cllem. Soc., 1950, 72, 5557. 35 Rabinowitch and Stockmayer, J . Ameu. Chem. Soc., 1942, 64, 335. 36 Igarischew and Turkowskaja, 2. physik. Chem. A, 1929, 140, 227. 37 Dewhurst, J . Chem. Physics, 1951, 19, 1329. 38 Kolthoff and Medalia, J. Arner. Chem. Sac., 1949, 71, 3784. 39 Dainton, private communication. classified document). document), p. 54.C . B. AMPHLETT 155 40 Eimer and Medalia, BNL-117 (U.S. Atomic Energy Cdmmission, declassified 41 see also Barb, Baxendale, George and Hargrave, Trans. Faraday Soc., 1951, 47, 608. 42 Hardwick, this Discussion, ref. (14). 43 Sutton, this Discussion. 44 Rigg, Stein and Weiss, Proc. Roy. SOC. A , 1952, 211, 375. document), p. 33.

ISSN:0366-9033    

DOI:10.1039/DF9521200144

出版商:RSC    

年代:1952

数据来源: RSC

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C . B. AMPHLETT 155 RADIATION CHEMISTRY OF AQUEOUS SOLUTIONS CONTAINING BOTH FERROUS ION AND CARBON DIOXIDE * BY WARREN M. GARRISON AND G. K. ROLLEFSON Crocker Laboratory, Radiation Laboratory and Department of Chemistry, University of California, Berkeley, California Received 5th February, 1952 The radiation chemistry of acid ferrous sulphate solutions containing dissolved carbon dioxide has been studied using high energy helium ions from the cyclotron as the radi- ation source. The principal products formed are ferric ion and hydrogen in a proportion very close to two moles of ferrous ion to one mole of hydrogen. A small fraction of the hydrogen is used to form reduction products of carbon dioxide. The most important of these is formic acid, but there are also much smaller amounts of formaldehyde and oxalate produced.The formate concentration approaches a definite limit asymptotically and is proportional to the concentration of dissolved carbon dioxide. A proposed mechanism gives the rate law which fits the observed data over the entire range of observation. The study of the radiation chemistry of aqueous solutions containing two solutes offers two general possibilities. The first is that the solutes may react with all of the H and OH radicals produced by the action of the radiation on the water and thus eliminate any back-reaction. The yield obtained under such conditions would represent the maximum yield of the primary process. The second possi- bility is that one solute will take all of the OH or all of the H and thus serves as an internal standard for the amount of radiation introduced and permits the study of the reactions of the other radical with various substances.For a substance to serve as such an internal standard it is sufficient that the amount of this standard reacting should not depend on the concentration of the substance added. Another condition which must be fulfilled in studies of this type is that the oxidation- reduction couples present in the solution must not interact in any manner inde- pendent of the radiation. The results which we are presenting in this paper are for experiments of this second type. We have used ferrous ion as a means of removing all OH radicals formed in the solution and have studied the reactions of hydrogen atoms upon dissolved carbon dioxide which as recently observed 1 can be reduced to formic acid and formaldehyde by high energy radiation.* The work described in this paper was sponsored by the United States Atomic Energy Commission.156 FERROUS ION AND CARBON DIOXIDE EXPERIMENTAL Solutions * containing one molar ferrous sulphate, one-tenth normal sulphuric acid, and carbon dioxide, were bombarded with 45 MeV helium ions in an all-glass target cell of the type shown in fig. 4. The beam (0.1 pA in all cases) entered through a thin glass window A. Solutions were de-aerated prior to the addition of carbon dioxide by evacuation through B. The C14 labelled carbon dioxide was added to the cell by either of two methods. In the first, C1402, prepared 2 by the action of perchloric acid on BaC1403, was added by means of a Toepler pump, through B which was then sealed off.In the second method, a known amount of NaC1403 was placed in the side arm C prior to evacu- ation. In this case, the C1402 was generated, after the target cell was sealed at D, by tilting the cell so that acid solution came in contact with the Na2C1403. Since only a few milligrams of Na2C1403 were required, no significant change in the acid concentration of the solution resulted. During bombardment, the target cell was supported horizontally and rotated in the target assembly shown in fig. 5. The irradiation period could be ac- curately determined by use of the shutter arrangement 5 4 56. The helium ion beam was de-limited by the water-cooled aperture, 2. The target cell, 8, could be rotated in the assembly, 7, which in turn was fastened to the shutter assembly at 6.After irradiation, the cells were attached at E through a glass seal to a vacuum system. The break-off F was opened and the gases were removed for mass spectrometric analysis and for C1402 recovery by means of a Toepler pump. The cell was then cut from the line and flushed with carbon dioxide followed by nitrogen to remove traces of C14O2 activity. An aliquot of this solution was titrated 3 with standardized titanous chloride solution to determine the amount of ferric ion produced. The remainder of the solution was then distilled in vacuo. The amounts of HC1400H, HC14H0, C14H30H and (C1400H)2 produced in the irradiation were determined by procedures based on the addition of milligram amounts of stable carrier for each of the irradiation products.These were separated as indicated below, using modifications of previously described techniques. To the distillate were added formic acid, formaldehyde and methyl alcohol carriers in amounts to give 100 mg of the isolated product, i.e. barium formate, methone derivative of formaldehyde, and barium carbonate prepared from carbon dioxide formed on oxidation of the methyl alcohol fraction. After addition of carrier, the distillate was again flushed with carbon dioxide, adjusted to pH 7 with sodium hydroxide, and redistilled in vacuo. The formalde- hyde and methyl alcohol were obtained in the distillate. The residue which contained the formic acid fraction was acidified to pH 1 and redistilled. The formic acid distillate (- 5 cm3) was flushed with carbon dioxide, neutralized to pH 8 with barium hydroxide, centrifuged to remove barium carbonate, and was then added to 75 cm3 of warm absolute ethyl alcohol which precipitated crystalline barium formate.The barium formate was dissolved in a minimum volume of water and recrystallized by the addition of ethyl alcohol. This process was repeated three times following which the barium formate was mounted on an aluminium plate to a uniform thickness and assayed for C14 activity using standard techniques2 The fraction containing the formaldehyde and methyl alcohol was mixed with methone solution in 50 % excess and acidified to pH 1 . The methone-formaldehyde derivative which precipitates under these conditions was separated from the methyl alcohol by distillation in VQCUO.The methyl alcohol fraction from one of the bombardments (number 2-2) was wet oxidized with chromium trioxide-sulphuric acid mixture containing potassium iodate, and the evolved carbon dioxide after precipitation as barium carbonate was found to be inactive. This observation agrees with the findings previously reported.1 The methone-formaldehyde precipitate was filtered off, washed, redissolved in sodium hydroxide solution, and reprecipitated by the addition of acid. This cycle was repeated three times. The methone derivative was then recrystallized twice from acetone+ water and assayed for C14 activity. The first residue containing the ferrous sulphate and (C14OOH)2 was dissolved in 75 ml of water and to this solution was added oxalic acid in an amount equivalent to 100 mg of lanthanum oxalate.Sodium hydroxide was added to precipitate ferrous hydroxide and to bring the solution to pH 10. After centrifugation, the supernatant, *The ferrous sulphate used in these experiments was Baker and Adams Reagent Grade without further purification. Water, from a Barnstead still, redistilled in glass from alkaline permanganate was used in preparing the solution. The C1402 as BaC1403, was obtained from the United States Atomic Energy Commission, Oak Ridge. Tennessee.-25 /+ -20 3 i ”/‘ 5 0 \ 2! : /o % -5 - ‘9 1: /+ + + Fe “f/ons x /O DISCUSSION It is apparent from inspec- tionof the results presented in table 1 that the principal pro- ducts formed in this system are ferric ion and hydrogen, in a proportion very close to two moles of ferric ion to one mole of hydrogen.This is brought out particularly well by the plot shown in fig. 1 and 1 ~ . It is to be noted that these results are not in agree- ment with Nurnberger4 who worked with alpha particles of considerably lower energy and found a rather wide range of values for this ratio. It is also shown in table 1 that a small fraction of the hydrogen atoms is used to form reduction pro- ducts from carbon dioxide. The most important of these is formic acid, but there are also much smaller amounts of formaldehyde and oxalate pro- duced. Since most of the hydrogen atoms combine to FIG. 1.4. form molecular hydrogen, the effective concentration of these atoms in the solution is determined by the rate of formation from the water and the rate of the158 FERROUS ION AND CARBON DIOXIDE combination to form the molecules.In this connection we wish to call attention to the distinction between what might be called the local concentration of atoms along the particle path and an overall concentration in the solution. This has been discussed by Magee.4 The local Concentration of hydrogen atoms will be essentially constant for a given radiation. bombardment number vol. of solution (cm3) gas space (cm3) C1402 added (mc) C14 in CO2 (%) H2 molecules ( x 10-19) Fe3+ ions ( x 10-19) Fe3 + ions/Hz molecules (Fe3+) x lo3 (co2) x 103 (HCOOH) x 106 (HCHO) x 108 3-2 20.0 7.20 1 -00 8.14 0.16 2.61 16.2 2-18 5.57 5-25 1-95 TABLE 1 2-2 2-3 13.2 13.2 6-85 7.0 1.00 1.00 7.39 8-14 4.06 3.73 4.10 6.87 1.01 1.84 5-18 8.54 7.46 7.35 31.1 42-2 7.20 9-85 4-2 21.0 7.05 1 -00 8-14 8.35 16.4 1-97 5.47 4.76 13.1 39.3 5- 1 22.8 7.08 1.00 7.39 16-0 29-1 21.2 37.4 1.82 4-94 5-24 4 3 21.0 6.30 1 -00 7.39 28.6 58.9 2.06 46.8 5.15 5.34 40.7 8- 1 26.0 6-40 0.10 8.14 10.2 21-8 2-10 14.0 0.48 3-5 - 6-2 29.0 8.05 none I 23.0 43.4 24.9 1 *90 - - - [(HCOOH)/(C02)] x 103 0.94 4.17 5-74 7.20 7-57 7.92 0.73 - [(HCHO)/(C02)] x 105 0.35 0.97 1.34 0.87 1.06 1.02 - - The results indicate that the amount of formate formed from the carbon dioxide is proportional to the amount of dissolved carbon dioxide.Therefore, in comparing the results of different experiments we have used the ratio of the con- centration of formic acid to that of carbon dioxide. Fig. 2 shows a plot of this - 80 - 70 - 60 t J0 .1 0 Y --.F e J f ( m ~ h / L , m ~ ) / D 20 30 40 FIG. 2.-Dependence of (HCOOH)/(C02) ratio of ferric ion concentration. quantity against the concentration of ferric ion produced. This is essentially a plot of the ratio against time, since the amount of ferric ion produced under constant radiation conditions is a linear function of time. It is apparent that the formate rises and approaches a quite definite limit asymptotically. The amounts of formaldehyde and oxalate formed are always quite small, compared to the formate (less than 5%). This approach of the formate concentration to a definite limit must be caused by the existence of compensating reactions, one forming the sub- stance, the other destroying it, so that in the limit the rates of the two processes are equal.WARREN M .GARRISON AND G. K. ROLLEFSON 159 The following set of reactions is suggested as a mechanism to account for the H2O = H + OH (1) formation and destruction of formic acid by means of hydrogen atoms : Reaction 1 shows the dissociation of water into hydrogen atoms and hydroxyl radicals. Reaction 6 indicates the removal of hydroxyl radicals by ferrous ion. It is immaterial for our discussion whether ferrous ion reacts with hydroxyl directly or with hydrogen peroxide, formed by the combination of two hydroxyls, as long as neither hydroxyl nor hydrogen peroxide react in any other way in the system. Reaction 7 shows the combination of the hydrogen atoms. We are assuming the reactions 1 and 7 together determine what might be called a steady state concen- tration of hydrogen in the system.Reactions 2, 3, 4 and 5 show the formation and decomposition of the formate. It is assumed that those reactions occur to such a small extent that they have a negligible effect on the steady rate concentration of hydrogen atoms. It is probable that all of these reactions occur without any significant activation energy. If that is true, in order to reach a steady concen- tration of formic acid as low as that observed, the reactive species of dissolved carbon dioxide in the solution must be one present in the concentration comparable to that of formic acid at the steady state. It is for this reason that we have assumed the reactive species in reaction 2 to be carbonic acid. Also, carbon dioxide and hydrogen atoms do not react readily in the gas phase.If we apply the usual procedures for deriving a rate law from a mechanism, we obtain the following equation : On the basis of the assumptions already stated, we may replace d(HCOOH)/dt by d(WCOOH)/d(Fe3+), and if we also consider that in our experiments the carbonic acid and hydrogen atom concentrations remain essentially constant, the equation can be shown in the form : HCOOH) d[(HCOOH)i(H2CO3>1 -~ = a - ( -__ d(Fe3f) (H2C03) Integration of this differential equation shows that a plot of the logarithm of a constant minus the ratio of the concentration of formic acid to that of carbon dioxide against the ferric ion concentration should be a straight line. Such a plot is shown in fig. 3. It is apparent that the data fit this over the entire range of observations.If we assume that all of these reactions proceed without activation energy, it is possible to draw some conclusions concerning the effective concentration of hydrogen atoms in the solution. The reaction of a hydrogen atom with another hydrogen atom to form a molecule, proceeds approximately 100 times as fast as the reaction with either carbonic acid or formic acid. Since the latter two sub- stances were present in concentrations of about 10-5 My this means that in our experiments the effective concentration of hydrogen atoms was about 10-3 M. Our mechanism suggests a possible means of accounting for the formation of160 FERROUS ION A N D CARBON D l O X l D b oxalate * by having two HC02 radicals come together in the solution. Since these radicals were present in rather low concentration, it is not surprising that FIG.3.-log [ a - (Hpl;)] ____ as a function of ferric ion concentration. 8 FIG. 4.-Design of target cell. FIG. 5.-Target assembly. * Experimental results, recently obtained by Garrison, Morrison and Hamilton, soon to be published, indicate that oxalate is one of the principal products in the irradiation of hydrogen saturated solutions of formic acid with 35 MeV helium ions.WARREN M . GARRISON A N D G. K . ROLLEFSON 161 the amount of oxalate formed was not large. Formaldehyde is probably produced by some reduction of the formic acid by hydrogen atoms. The amount formed in our experiments is too small to warrant any attempt to set up a detailed mechanism. We hope to study the possible reduction of formate in detail in Iater experiments. We wish to thank Professor Joseph G. Hamilton, Director of Crocker Labor- atory, for his interest in this problem and for many helpful discussions, Dr. Amos Newton for the mass spectrometric data, Mr. D. C. Morrison and Mr. Boyd Weeks for help in some of the experimental procedures, Mr. Herman R. Haymond for several suggestions during the course of the investigation, and the staff of the 60-in. Cyclotron at Crocker Laboratory for bombardments and assistance in target design. 1 Garrison, Morrison, Hamilton. Benson and Calvin, Science, 1951, 114, 416. 2 Calvin, Heidelberger, Reid, Tolbert and Yankowich, Isotopic Carbon (John Wiley 3 Treadwell and Hall, Analytical Chemistry, voI. 2 (John Wiley and Sons, New York, 4 Nurnberger, J . Pizysic. Chem., 1934, 38, 47. 5 Magee, J . Anter. Ciiem. Suc., 1951, 73, 3270. 6 Taylor and Marshall, J . Physic. Chem., 1925, 29, 1140. and Sons, New York, 1949). 1942).

ISSN:0366-9033    

DOI:10.1039/DF9521200155

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

WARREN M . GARRISON A N D G. K . ROLLEFSON 161 THE DECOMPOSITION OF HYDROGEN PEROXIDE BY IONIZING RADIATIONS AND RELATED PROBLEMS BY JOSEPH WEISS University of Durham, King’s College, Newcastle-upon-Tyne, 1 Received 13th February, 1952 The mechanism of the photochemical decomposition of hydrogen peroxide has been discussed particularly in relation to the decomposition by ionizing radiations. The latter has been treated with reference to the reaction in and between the tracks produced by the fast particles in the solution. Finally, the formation of molecular hydrogen in the irradiation of aqueous hydrogen peroxide solutions, and in related cases, has been referred to. It is well known that hydrogen peroxide in aqueous solutions is decomposed by ionizing radiations. AIthough this decomposition shouId be different in many respects from the photochemical decomposition, these two processes are bound to have certain features in common as the free radicals, OH and HO2, which occur in the photochemical reaction, play also an important part in the radiation chemistry of aqueous solutions in general.It may be of some interest, therefore, to summarize briefly the present state of the theory of the photochemical decomposition of hydrogen peroxide in aqueous solutions. Urey, Dawsey and Rice 1 the photochemical primary process corresponding to light absorption in the ultra-violet ( A < 3100A) is represented by THE PHOTOCHEMICAL DECOMPOSITION OF HYDROGEN PEROXIDE.-According to H202 + h~ -> 20H, (1)162 DECOMPOSITION OF HYDROGEN PEROXIDE I, being the number of quanta absorbed per unit of time; this process is com- bined with the chain reactions 2,- = with H202 + OH + H20 + H02 k2 (2) H202 1- 0 2 - + 0 2 f OH-- + OH k3 (3) HOz + H+ + 0 2 (diss.const. KHo2), with a suitable reaction for the disappearance of the free radicals (chain breaking) e.g. 20H -+ H2O 4- 0 k4 (4) 2 0 + 0 2 . (4a) Eqn. (3) appears here in a modified form,Zc where the anion 0 2 - replaces H02, which can account for various pH effects, but it is important to note that in re- action (3) the 02- reacts with the (undissociated) H202 molecule. The above reactions lead to the following equations for the stationary state : d(OH)/dt == 0 == 21, - k2[H202][0tI] +- k3[H202][02-] - 2k4[OH]2 (la) d(H02)ldt = 0 =-= k2[H202][OH] - k3[H202][02-] (1b) giving I, - k4[0H]z = 0.(W For the rate of disappearance of hydrogen peroxide one obtains and for the quantum yield y (under conditions when [H202] can be regarded as practical] y constant), The dependence on the square root of the light intensity is in general agreement with the experimental results of a number of authors.3~4-5 From eqn. (Ha) (and the equations in table 1) it follows that the light intensity enters also as a linear term, which should assert itself, particularly at higher light intensities, and also under conditions when the chain-breaking process (6) is predominant. There is also considerable evidence that the rate of the photolysis is proportional to the H202 concentration.3~ 5 ~ 6 9 7 In more concentrated solutions of hydrogen per- oxide the situation appears to be somewhat complex, but there are no reliable data available for this region.The simple mechanism given above leads to a quantum yield independent of the hydrogen ion concentration. This is appar- ently in accordance with certain results reported in the literature, although there is no general agreement on this point. Recently it has been shown8 that the dependence of the quantum yield on the light intensity and pH, is closely connected with the nature of the chain- breaking process. In the case discussed above the lower limit of the quantum yield is 3/min - 1 (cf. eqn. (ITb)). In table 1 a few other possible chain-breaking processes have been listed with the corresponding expressions for the rates and the quantum yields. In certain cases a first order disappearance of the radicals may have to be considered according to reaction (6) in table 1, where X represents reactive " impurities " or a " wall ".Even under the most extreme conditions so far employed the quantum efficiency does not fall below unity and, generally, is not less than ca. 1.5. This result suggests strongly that there is no appreciable recombination of OH radicals to reform hydrogen peroxide (eqn. (8), table l), which is supported by the work of Bonhoeffer and Pearson 9 in the gaseous phase and by some work of Lea 5 in solu- tions. Theoretical considerations 10 concerning the interaction of OH radicals also favour reaction (4) rather than reaction (8). However, it cannot be ruled out completely that those cases, where the quantum yield lies between 1 and 2JOSEPH WEISS 163 might arise from a combination of reactions (8), with, e.g., reactions (5) or (7), although this is rather improbable.According to the above considerations a dependence of the quantum yield on the pH is to be expected unly under conditions when the quantum yield is appreciably greater than unity. There is, therefore, no obvious discrepancy between the earlier results of Kornfeld obtained under conditions of y > 2 and those of Heidt 11 and Lea,5 who worked in the region of low quantum yields. Most of Kornfeld's results 3 seem to be compatible with a [H+]-+ dependence of the quantum yield (see equations in table 1). chain breaking process TABLE 1 Y quantum yield Ymin (6) OH+X+ chain bc., k6 -tO (8)20H --f H202, k8 -f This particular approximation was chosen here because it gives the correct expression for long chains and also leads to the correct minimum value of the quantum yield.For a study of the pH effect over a wider range, the concentration of the [H202] in the rate equations (IIb) and in table 1 has to be expressed in terms of the total (analytical) concentration [HzOa] of the hydrogen peroxide and its dissociation constant KI1202 according to A number of photochemical experiments have been carried out using the sector technique : 49 5 there are some unexplained features in this work although nearly all the experiments lead to relatively long life times of the active radicals of the order of 0.5 to 1 sec. This point is of some importance in connection with the decomposition of hydrogen peroxide by ionizing radiations discussed below.tively few quantitative data are available in the literature concerning the decom- position of hydrogen peroxide by ionizing radiations and most of these are con- cerned with the decomposition by X-rays in dilute aqueous solutions.12.13 In the latter case one has to consider primarily the action of the radiation on the solvent, water, which is split by the processes of excitation and ionization, according to the net process : 14 As was pointed out already previously 14 the H atoms can then react according to H202 + H -+ H20 + OH + 65 kcal klo and the OH radicals can enter the chain reactions (2) and (3) and into the chain breaking processes discussed above. Although one may also have to consider, under certain conditions, other processes such as H + 0 2 -+ HO2, THE DECOMPOSITION OF HYDROGEN PEROXIDE BY IONIZING RADIATIONS.-Rela- H20 Z S ~ S + H 4- OH kg (9) (10) (1 1)1 64 DECOMPOSITION OF HYDROGEN PEROXIDE from a general point of view, the decomposition by ionizing radiations is clearly very similar to the photochemical decomposition. However, whereas in the photochemical reaction the radicals are produced more or less uniformly throughout the irradiated solution, with ionizing radi- ations they are formed along the tracks of the fast particles.The resulting inter- play of diffusion processes and of chemical reactions is difficult to describe exactly. The theoretical treatment of Lea 15 and others are based on Jaffe’s theory which, although generally valid in gaseous systems, is often not a good approximation for liquid systems.It is possible to distinguish broadly the following two limiting cases : (i) One may be able to confine oneself to the consideration of the chemical reaction in the individual tracks 16 and, as a first approximation, neglect diffusion processes. This is generally justified when the yield is independent of the dose rate, which is the case, e.g. when the reactions in the tracks are independent of each other. This will hold when the dose rate is not too high, so that the mean distance between the tracks is sufficiently great to prevent any appreciable interaction during their lifetime, which is evidently favoured when the lifetime of the radicals in the tracks is sufficiently short.In these cases, diffusion processes will still in- fluence, to a greater or lesser extent, the concentration of the radicals within the tracks and this will have to be taken into account if one is concerned with the determination of the actual rate constants. (ii) There is the other extreme case to be considered when, either on account of the high dose rates and/or due to the long lifetime of the radicals, practically complete intermixing of the tracks occurs during the early stages of the reaction. Under these conditions the situation is very similar to the photochemical case, so that the above considerations concerning the latter can be applied directly. The sector experiments carried out on the photochemical decomposition of hydrogen peroxide indicate that the radicals possess a relatively very long life- time ; 43 5 it would, therefore, not be unexpected to find an appreciable “ inter- track ” reaction in the decomposition by ionizing radiations, even at compara- tively low dose rates.Actually, Fricke,l3 who studied the decomposition by X-rays, found a dependence on the dose rate and claimed to have established the relation : rate of decomposition oc (dose rate)$ [H202]4. This result has been confirmed recently by Hart et aI.17 and to some extent also by Johnson .18 This dependence on dose rate is analogous to the dependence on the square root of the (absorbed) light intensity in the photochemical reaction. It is clear, however, that the photochemical approach can only represent a relatively rough approximation for the radiation chemistry, and this is also shown by the fact that while the decomposition by ionizing radiations is proportional [H202]* the photochemical reaction, in the same concentration range, shows at linear depend- ence on the hydrogen peroxide concentration.It is, therefore, of interest to approach the radiation chemistry from a different point of view, viz. by starting with the reaction in the tracks and investigating the interaction between the tracks. is also described by the radical mechanism outlined above except that, in general, the process cannot be treated as a stationary problem, although the method of stationary concentrations can be applied in radiation chemistry in certain special cases with some success. In general, however, a treatment of the non-stationary state is required.Therefore, in the present case, instead of eqn. (Ia) and (Tb) we shall consider the following equations : THE DECOMPOSITION OF HYDROGEN PEROXIDE WITHIN A SINGLE TRACK.-This d(OH)/dt - k2[H202][OH] -: k3[H202][02 ] -- 2k4[OH12, ( 1 1 1 ~ ) d(H02)/dt =_ k2[H202][0H] -- k3(H202][02-]. (I1 Ih)JOSEPH WEISS 165 The term corresponding to the first term in eqn. (Ia) can be omitted, because the initial production of the radicals, following the passage of the fast particles, is rapid compared with the lifetime of the radicals. An estimate of the time necessary for the initial production of the radicals can be obtained, e.g. from the pre-dissociation of excited water molecules, viz : H20* --f H + OH, which should require about 10-13 sec, and from the Debye relaxation time of water dipoles in liquid water, which is of the order of 10-11 sec, and which should represent a reasonable estimate of the lifetime of the unstable ions, e.g., according to the reaction : The non-linear differential eqn.(IIIa) and (IIIb) cannot be integrated in closed form and it is necessary to introduce certain approximations. The following calculations are confined to a more or less symbolical treatment. Thus, one may estimate the concentration n of the radicals from the differential equation : ([S] denotes the concentration of the acceptor, e.g. [ S ] = [H2021) (W which has a linear term (rate constant k,) characteristic of the reaction with, e.g., hydrogen peroxide and a quadratic term (rate constant ks) to account for chain- breaking processes.In the hydrogen peroxide decomposition, using eqn. (Ta) and (Tb), this approximation is equivalent to the assumption that - dn/dt = k,[S]n 1- 2kgn2, [H021 prop. [OH], but with a proportionality factor different from the one corresponding to the stationary state. Integration of eqn. (IV) gives, for [S] YJ const. : where no (the concentration of the radicals at t == 0) is given by no = Nolnb2, where b represents the (mean) initial diameter of the track, and No denotes the initial linear density of the radicals. The concept of a simple “ track ” of uniform linear density is, of course, also an approximation. The rate of decomposition of hydrogen peroxide is given by an expression of the form : corresponding to eqn.(Ha) given above. decomposed (A[S]) during the total lifetime of the track is then given by - d[S]/dt N const. [Slit (Via> The amount of hydrogen peroxide const. k,[S] In (1 + ’*o>. (vI~) 2kB kJS1 Denoting by Jlr the total number of tracks produced per unit of time and volume and by 1 the average length of a track, the total amount of hydrogen peroxide (fA[S]) decomposed in all the tracks produced per unit of time and volume is given by For the yield I‘, (number of molecules of hydrogen peroxide decomposed per radical pair produced) one obtains166 DECOMPOSITION OF HYDROGEN PEROXIDE The value of G (moles decomposed per 100eV) only differs from this by a proportionality factor and is given by G = Fr(l0O/W,-), where W, represents the energy (in eV) required to produce a radical pair in the solution.Eqn. (VII) shows that, with the assumption that the tracks do not influence each other, the yield is, of course, independent of the dose rate, but that it de- creases, in general, with increasing linear density of the radicals (No), e.g. the yield of the decomposition should be smaller for alpha particles than for gamma rays. The dependence on the hydrogen peroxide concentrations can be almost a linear one and may practically disappear according to the magnitude of the second term under the logarithm in eqn. (VII). and/or for sufficiently long lived radicals the situation is different, as discussed above, and the total reaction can be conveniently divided into (i) the reaction in the individual tracks, up to the time when they begin to show any appreciable overlapping, followed by (ii) the general “ inter-track ” reaction which proceeds more or less uniformly in the whole volume.A more exact theoretical treatment would be difficult and laborious and the simple calculations given below are based on a greatly simplified model which, nevertheless, may reveal some features of more general interest. THE REACTION DUE TO THE INTERACTION OF THE TRACKS.-FOr higher dose rates We start from the differential equation : 3n/3t LL- DQ% - k,[S]n - 2k@, (VIII) which is the same as eqn. (IV) except for the diffusion term (Dv22n). estimated. If eqn. (VIII) is written as The initial contributions of the different terms in eqn. (VIII) can be easily 3n/3t =; [d] - pa] - [rp] (VIIIa) at t = 0, the order of magnitudes of these terms are NO and b have the same significance as above, all concentrations, including that of the acceptor [el, are given in molecules cm-3.Z,, Z b and Pa, Pg denote the collision numbers and Boltzmann factors respec- tively of the two reactions. For water at room temperature, D N 10-5 cm2 sec-1, b - 10-6 cm and taking 2, - Z p 1: 10-10 cm3 sec-1, one obtains for the initial ratios where [S] is now expressed in moles/l. It will be seen that the diffusion term is always important initially if P,[S] < 10-4, and/or PBNo < 105. Nevertheless, for a semiquantitative treatment the diffusion term can be neglected, as pointed out above, unless the diffusion leads to a greater or lesser extent of interaction of the tracks ; this case is considered below.For this treatment of eqn. (VIII) we assume for n a Gaussian distribution, viz. : where r denotes the distance from the centre of the columnar tracks and N the linear density of the radicals. This equation will be used here only in the ap- proximation indicated above, where 6 represents a suitable mean value for the167 JOSEPH WEISS radius of the columns lying between the value of b at t = 0 and the value of r which is given by (XI where represents the mean distance between coexisting tracks (XJ. This is obtained from - R N (4Dtjj + b2)P - R - NS-* (XJ) (and for the 3-dimensional case, x’ - Ms-g), where Ns represents the number of coexisting tracks per unit of volume and can be expressed by the relation where JV (the total number of tracks produced per unit of time and volume) is proportional to the dose rate, and T denotes the mean “ lifetime ” of the tracks, which, under certain conditions, can be identified with the mean lifetime or the radicals in the tracks.It is clear that eqn. (XII) for hPs is only valid if y e t , where r represents the time of irradiation; this condition is generally fulfilled. If, on the other hand T >> t, which may be true, e.g. under certain conditions of pulsed radiations, then the number of coexisting tracks is approximated by the total number of tracks (Mt) produced during the (relatively short) time t. On the other hand t i , as de.fined by eqn. (X) is the time when the mean distance of the (coexisting) tracks is equal to the mean diameter of the columns. To obtain an estimate of the ‘‘ inter-track ” reaction we have to calculate the number of radicals which escape from the individual tracks by diffusion and thus contribute to the general volume reaction.This can be done by the following approximate method. The number of radicals N, disappearing by the chemical process per unit of length of a columnar track until the time when it has reached the radius R is given by M s = M T , (XW - dN,/dt = (k,[S]n 4- 2kp2)2.rrrdr = AN + 2BN2, (XITIa) k nb2 _ - J: where A = k,[S](l - exp (- R2/b2)), B = A(l - exp (- 232ib2)). (XIIIb) The number of radicals NO disappearing by diffusion from unit length of a - dND/dt ;= 2D%~(3rz/3r), = = ON, (XIVa) track of radius k is given by the equation : where (XIVb) The total change of N is then given by the differential equation : - dN/dt -- d(N, 4- ND)/dt ET ( A + O)N + 2BN2, (XVa) which on integration gives where No is again the linear density of the radicals at time t = 0.The fraction K of the total number of radicals disappearing by diffusion from a columnar track of radius R, per unit of length, is then given from eqn. (XIVa) and (XVb) by168 DECOMPOSITION OF HYDROGEN PEROXIDE The " inter-track " reaction itself can be treated again as a stationary problem as it is very similar to a photochemical (volume) reaction. The total number of radicals produced per unit of time and volume which contribute to the intertrack reaction is given by ( N O ~ N K ) where the symbols have the same significance as above. The stationary concentration ns of the radicals follows from an equation corresponding to eqn.(V) and is given by (XVIIn) and with eqn. (XVI) NOINK - 2kpn,2 = 0, (XVITb) For the rate of decomposition one thus obtains which shows that the rate is proportional to the square root of the dose rate and also depends on NO. The dependence on the concentration of hydrogen peroxide, after substituting from eqn. (XIIIb) (for [S] = [H202]), is of the form, which can vary from a practically linear relationship to a proportionality with (H202)*, the latter being the case if &[H202] :* 1 > K1. Under the conditions discussed above the reaction in the tracks is also modified, as the (total) number of radicals available per unit length of track is now of the order of NO (1 - K), but the reaction within each track is on& defined as such during the time interval from 0 to ti?, after which the tracks lose their independent existence.The above calculations represent only rough approximations but a more exact theoretical treatment does not appear warranted until more detailed experimental data have become available. composition of hydrogen peroxide by X-rays is accompanied by the formation of small quantities of molecular hydrogen. Although this may not be very im- portant for the general course of the decomposition reaction, a special interest attaches to it because the formation of small amounts of molecular hydrogen, even in the presence of certain oxidizing agents, appears to be not uncommon in the radiation chemistry of aqueous solutions. At first sight this may be rather unexpected because most oxidizing agents are good acceptors for hydrogen atoms.This seemingly paradoxical result can be explained. by the intervention of the hydrogen molecular ion H2+ in solution, which can be formed, under suitable conditions, according to the equilibrium : 19 THE FORMATION OF MOLECULAR HYDROGEN.-According to Risse '2 the dt- H + Hf ,-f H2f and which can no longer act as a reducing agent but only as an oxidizing agent.17 For the hydrogen peroxide reaction one may have to consider the reactions : H2f + HO2- -+ H2 + HO2 (1 3) H2+ + 0 2 - + H2 + 0 2 (14) and particularly for the formation of molecular hydrogen. These correspond to the non-ionic processes : H -t H 2 0 2 + H2 t - HO2 H -1- H02 -> PI2 + 0 2 .JOSEPH WEISS 169 In general, every suitable electron donor X- could, in the presence of H2+, lead to the formation of some molecular hydrogen according to 20 H2f + X- -+ H2 + X. Apart from this, a certain amount of molecular hydrogen may be formed on account of the primary recombination of some of the hydrogen atoms produced at rela- tively high local concentrations in the tracks of the 8-rays.21 This will depend on the nature of the radiation 14 and, to some extent, on the solute present. 1 Urey, Dawsey and Rice, J. Amer. Chem. SOC., 1929, 51, 1374. 2 (a) Haber and Willstatter, Ber., 1931, 64, 2844 ; (6) Haber and Weiss, Proc. Roy. 3 Kornfeld, Z. physik. Chem. B, 1935, 29, 205. 4 Allmand and Style, J. Chem. SOC., 1930, 596, 606. 5 Lea, Trans. Faraday SOC., 1949, 45, 81. 6 Henri and Wurmser, Compt. rend., 1913, 157, 654. 7 Taylor and Anderson, J. Amer. Chem. SOC., 1923, 45, 65. 8 cf. Weiss, Advances in Catalysis, 1952, 4 (in press). 9 Bonhoeffer and Pearson, Z. physik. Chem. By 1931, 14, 1. 10 Weiss, Trans. Faraday SOC., 1940, 36, 856. 11 Heidt, J. Amer. Chem. SOC., 1932, 54, 2840. 12 Risse, Z. physik. Chem. A, 1928, 140, 133. 13 Fricke, J. Chem. Physics, 3, 1935, 364. 14 Weiss, Nature, 1944, 153, 748. 15 Lea, Actions of Radiations on Living Cells (Cambridge, 1946). 16 Weiss, Trans. Faraday SOC., 1947, 43, 314. 17 Gordon, Hart and Walsh, AECU, 1534 (Dec. 1950). 18 Johnson, J. Chem. Physics, 1951, 19, 1204. 19 Weiss, Nature, 1950, 165, 728. 20 Rigg, Stein and Weiss, Proc. Roy. SOC. A, 1952 (in press). (University of Durham, 1951). 2lcf. also Allen, J. Physic. Chem., 52, 1948, 479. SOC. A, 1934, 147, 332; (c) Weiss, Naturwiss., 1935, 23, 64. Rigg, MSc. Thesis

ISSN:0366-9033    

DOI:10.1039/DF9521200161

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

JOSEPH WEISS 169 MECHANISM AND RATE CONSTANTS OF THE y-RAY INDUCED DECOMPOSITION OF HYDROGEN PEROXIDE IN AQUEOUS SOLUTIONS BY EDWIN J. HART AND MAX S. MATHESON Argonne National Laboratory Chicago 80 Illinois U.S.A. Received 28th January 1952 A mechanism for the y-ray initiated decomposition of dilute aqueous solutions of hydrogen peroxide has been deduced from data showing a dependence of decomposition yield on the square root of hydrogen peroxide concentration and inverse square root of dosage rate. A novel feature of this mechanism is that termination occurs through a termolecular reaction involving two hydroperoxy radicals and a hydrogen peroxide mole- cule. The propagation and termination rate constants have been measured in inter- mittent radiation experiments which were carried out in paraffin coated cells containing 0.1 M hydrogen peroxide under irradiation conditions yielding 30 molecules of oxygen per initiating radical.Evidence is presented showing that HOz + H202 = H20+02+OH and 2H02 + H202 = 2H202 + 0 2 are the rate determining propagation and termination steps. kD is found to be 530 1. mole-1 sec-1 and 2kt equals 5-3 x 1010 1.2 mole-2 sec-1. The chain character of the photochemical and X-ray initiated decomposition of hydrogen peroxide has been discussed by Allmand and Style,l Lea,2 Fricke 3 F DECOMPOSITION OF HYDROGEN PEROXIDE 170 and Johnson.4 Fricke reported a dependence of yield on the reciprocal square root of the dosage rate for X-ray irradiations. This result is evidence that two chain carriers are involved in the termination step.Allmand and Style applied the intermittent light technique to measure the lifetime of the kinetic chain. It was postulated that the hydroxyl free radical was the catalyst present and the chain was found to have a mean life of 1 sec under the conditions these authors employed. The termination step proposed was OH + OH = H202 (or M20 -t i 0 2 ) . However no proof was offered for this step. More recently Lea 2 and Volman 5 have employed bimolecular termination of two hydroperoxy radicals. However the radiolysis of hydrogen peroxide solutions has been shown by Fricke and Johnson to be dependent on the square root of hydrogen peroxide concentration. In agreement with this Allmand and Style find the relative quantum yield to be constant below one M H202 with 3650A light of constant intensity.This result is expected if the rate is proportional to (H202)i (labs)* and Iabs is proportional to (H202). These results are not explained by any of the mechanisms proposed. The purpose of the present paper is to report work in which the rotating sector was applied to the decomposition of hydrogen peroxide by y-rays in order to measure rate constants for the propagation and termination steps. A further object is to clarify if possible the mechanism of this reaction. EXPERIMENTAL PURIFICATION OF MATERIALS.-The methods previously described 6 were used for the preparation of the distilled water and for evacuation of solutions to be irradiated. Ex- haustive irradiation of such air-free triply-distilled water samples produced of the order of one micromolar carbon dioxide.This small carbon dioxide yield was the criterion used for testing the purity of the water the cleanliness of the irradiation cells and the general technique of removing air from the solutions. As shown earlier less than one micromole of oxygen per litre of water remains in the solutions after the evacuation procedure. While the complete removal of oxygen is not critical in this hydrogen per- oxide work careful attention must be taken to insure the removal of impurities from the water and the cell surfaces. Merck reagent grade 30 % hydrogen peroxide and Buffalo Electrochemical 90 % hydrogen peroxide were used in the present work. The former was purified by steam distillation and the latter used without further purification.The results appearing in fig. 2 and 4 were obtained using the steam distilled hydrogen peroxide. All of the work in paraffin coated cells at or near 0.100 M hydrogen peroxide was carried out on the 90 % product. In these cases a stock solution of 15 M hydrogen peroxide was added to freshly distilled water in the evacuation chamber which had previously been furnace heated to 550" C. Subsequent degassing of the solution and filling of the cells was usually carried out within 15 min. ANALYSES.-The Van SIyke micro-gas analysis apparatus was used for measuring the amount of oxygen formed in the irradiation cells. For concentrations of one millimolar no problem with respect to the decomposition of hydrogen peroxide during analysis was encountered.However the regular method of analysis could not be used for 0.10M concentrations. At these concentrations attack on the mercury in the Van Slyke by the hydrogen peroxide is rapid unless the solution is made strongly acid. In order to circum- vent this difficulty and to condition the surface of the mercury the Van Slyke was evacu- ated with 1 ml of 0.1 M hydrogen peroxide in 18 N sulphuric acid in the chamber. After evacuation 0.5 ml of the gas-free solution was left in the chamber. The mercury and this solution were then lowered to the 50 ml line and the hydrogen peroxide from the irradiation cell introduced. For 0.1 M hydrogen peroxide the gas analyzed consisted of 99.5 % oxygen the remainder presumably being hydrogen.Samples of 0-1 M hydrogen peroxide analyzed 10 min after evacuation showed 4 to 5 pM oxygen. This small amount of gas is believed to be due to the decomposition of hydrogen peroxide during analysis ; however it could be due to an initially high catalytic decomposition by the walls of the irradiation cells. EDWIN J . HART AND MAX S . MATHESON 171 EFFECTS OF CELL PREPARATION AND WATER PuRITY.-The rotating sector experi- ments were carried out using 0.1 M hydrogen peroxide in order to obtain the long chains assumed in the mechanism. At this concentration considerable difficulty was en- countered in obtaining consistent and reproducible results. After carrying out the series of experiments appearing in table 1 it was concluded that the irregularities in results were due to impurities in the water probably originating in the glass surface.The best solution found to this problem was to coat the inside of the irradiation cells with paraffin a procedure also used by Allmand and Style in some experiments. A non-crystalline paraffin (Bioloid Embedding Paraffin m.p. 60-62" C Will Corp. Rochester New York) was found satisfactory. Two sets of eight cells each were used in alternate experiments in the sector work. The inner walls of the cells were coated with paraffin as follows. The furnace-treated cells were warmed to 100" C in an oven and then about 0.5 ml melted paraffin drawn by suction through the standard taper capillary tip into the cell. The cell was then rotated until the entire surface except for that in the thin capillary tip was covered.The excess paraffin was then drained through the standard taper tip. After cooling to room temperature these paraffined cells were ready for use. After use the paraffined cells were recleaned by alternate treatments with methanol and benzene until the paraffin was dissolved. The cells were then further freed from organic impurities by reheating to 550" C . The effect of treatment of the cells and purity of the water on the decomposition of 0.10 M hydrogen peroxide by y-rays is given in table 1. The treatment that the cells for a particular experiment were given appears in column 2. In general eight cells were filled at one time and six of these cells irradiated for a total dosage of 3.4 x 1020 eV/l.The two unirradiated cells were used as thermal controls and analyzed for oxygen at the end of the run. Column 4 gives the averages for the irradiated samples corrected by sub- traction of the thermal reaction. The standard deviation giving a measure of the vari- ability of the reaction within an experiment appears in column 5. The experiments were designed to test the effects of several types of cleaning procedures and other variables leading to non-uniform results. Expt. 1 and 2 of table 1 show the effect of triple and regular distilled waters. The latter produces 4.9 times as much decomposition during irradiation as does the triple distilled water. Packing cells with glass helices so as to increase the surface 60-fold possibly reduces the net over-all reaction but has no marked effect.The effects of con- taminating the solution with potassium dichromate are given in expt. 4 6 and 8. Al- though the 0.1 pM dichromate does not affect the results the effect of 10 pM dichromate is impressive and hereafter the use of cleaning solution was avoided in this hydrogen peroxide work (compare expt. 6 and 8). Expt. 9-14 inclusive show the results of pains- taking attempts to obtain reproducible rates of decomposition within a run and between runs. This proved to be futile and from the work of irradiations 11 to 14 it was con- cluded that impurities from the cell wall play a predominant role. The more detailed data of irradiations 11 to 14 on which this conclusion is based are given in table 2. Irradiations 11 and 13 were performed with one set of cells and irradiations 12 and 14 with another set.In all four irradiations there is wide variability in the results for each irradiation that is much greater than the 2 to 4 % error in measuring the oxygen content of the cells (see table 2). The column headed R shows that the oxygen yields in a set of cells change by the same factor for each cell when the set is given different treatments in two successive experiments and the constancy of this factor holds even when the variability between cells within the same experiment is large. The constancy of R suggests that variation in oxygen yields is to be associated with the individual cells prob- ably with contaminations originating in the cell walls. In support of this hypothesis microscopic examination of the inner glass surfaces of the cells disclosed many irregular- ities and discoloured spots.Treatment with hydrofluoric acid failed to remove all con- tamination or to provide uniform results (see expt. 15 and 16 table 1). Further support is furnished by expt. 20 and 21 where the use of paraffin coated cells eliminated variation between cells within an experiment and reduced the variation between experiments. Paraffin coated cells were used in the rotating sector experiments. Although the impurities are postulated to originate in the cell walls the impurities that affect the reaction must presumably become dispersed throughout the solution very early in the reaction. This conclusion is based on a calculation using the usual random walk displacement theory.7 The lifetime of a kinetic chain was measured as 0.60 sec in the sector experiments to be described the jump frequency of a diffusing radical was taken as 2 X lOlo/sec,g and the jump distance from one equilibrium position to the next I72 irrad.no. treatment of cells * 2 I 3 4 5 6 7 8 9 10 11 WSF WHFF WSF NSDF 3XF 3XF HFF HFF 12 13 14 15 16 17 DECOMPOSITION OF HYDROGEN PEROXIDE TABLE NAL EFFECT OF TREATMENT OF CELLS AND PURITY OF WATER ON THE DECOMPOSITION OF 0*10M HYDROGEN PEROXIDE BY GAMMA RAYS WCSF WCSF SF WF WSF WCSF WFCSDF F WFCSDF WSDF WNSDF WNDF WNSDF no. of G l - 25 141 24 106 cells 6 6 6 6 4 2 4 4 6 6 6 6 6 6 5 3 6 6 6 4 6 av.pmoles O?/l. 4 h 266 1293 174 1433 237 41808 279 28 1 196 167 143 178 81 156 257 815 307 297 182 464 435 - 61 37 31 47 29 16 32 10 24 70 20 76 103 12 21 27 3X distilled water regular distilled water cell packed with helices cells contaminated with cleaning solution 3X distilled water 10 micromolar K2Cr207 3X distilled water 0.1 micromolar K2Cr207 3X distilled water 3X distilled water effect of HF treatment 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 0 /a v.0.0940 0,1090 0.1375 0.0740 0.2574 I 0.1326 0-1 103 0.2398 0.1737 0.1 119 0.1798 0,1235 0.1538 0.2724 0.0245 0.2476 0.3468 0.0659 0.045 3 0.0621 3XF WF W3XF paraffin 18 19 20 21 paraffin * Meaning of codes C = heated in cleaning solution ; S = steamed 1 h; F = heated in furnace to 600" C ; D = dried in oven at 110" C ; N = heated in hot concentrated nitric acid for 30 min ; W = rinsed out several times with regular distilled water ; 3X = rinsed out several times with 3X distilled water ; HF = treated with HF for 20 min ; paraffin = inner walls of furnace-treated cells coated with paraffin.-f c = standard deviation. $ Both the blank and the irradiated cell contained this same total amount of oxygen. was taken as 3.1 x 10-8 cm from the diameter of a water molecule so that the calculated total distance diffused by the radicals involved in one kinetic chain is 0.03 mm. There- fore about 0.3 % of the chains would reach the walls of the 1.8 cm diameter cells used in the sector experiments. The rates for paraffin cells in the sector experiments are in good agreement with the paraffin cell rates of table 1 when corrected to the 3.7-fold higher intensity used for table 1 experiments. This shows that in the table 1 experiments even fewer chains reached the walls since the lifetimes are somewhat shorter due to the higher rates.APPARATUS.-The rotating sector method of Briers Chapman and Walters 9 has in recent years been applied extensively to oil phase polymerization systems 103 11 and to liquid phase hydrocarbon oxidations.12.13 In the work cited ultra-violet light was used to initiate the free radical reactions studied. In the present work on the decomposition of dilute aqueous hydrogen peroxide the y-radiation from C060 (1-1-1 -3 meV) was used to initiate chains and the penetrating nature of this radiation introduced experi- mental complications not encountered with ultra-violet light. The nature of the corn- plications will be apparent in the discussion of the apparatus. remarks EDWIN J. HART AND MAX S . MATHESON TABLE 2.-A COMPARISON OF THE DECOMPOSITION OF 0.10M HYDROGEN PEROXIDE BY GAMMA RAYS IN UNCOATED CELLS cell no.itrad. no. 26 82 85 11 11 11 1 1 11 11 87 89 90 12 12 12 30 76 78 79 83 86 axis of the apparatus. S 12 12 12 Fig. 1 shows a vertical cross-section through what may be designated as the optical The cylinder S is a C060 source similar to that described earlier 6 FIG. 1 .-Vertical cross-section of intermittent y-ray device. but of 80curie strength. A is an aperture in a portion of the lead shield surrounding the source. At this point the lead shield is 25-4 cm from the source and is 16.2 cm thick. Because of the mass of lead required to stop the y-rays rotation of a disc containing the aperture is not a very practical method of obtaining intermittent radiation particularly for high speeds of rotation.Therefore the cells C were mounted on the aluminium wheel W (radius = 12.0 cm to centres of cells which were 5-1 cm long) close behind the lead shield and aperture and the rotation of this wheel caused the cells to be irradiated intermittently as they rotated past the opening in the lead shield. Holders for twelve cells were spaced equally around the wheel. The aluminium wheel could be rotated at constant speeds ranging from 600 rev/min to 0-0938 rev/min by means of a synchronous motor and a gear box. The aperture was designed to provide a cycle in which the cells were irradiated one-fourth of the time in each cycle when the wheel W was rotating at constant speed. To accomplish this the side surfaces of the aperture were generated according to the following scheme.The aluminium wheel is stopped so that a vertical line through the axis of the wheel bisects the angle between two adjacent cells. The angle between cells being 30" this means that two other cells are located so that their centre lines make angles of 45" with the vertical. The centre of the upper end of one of these 45" cells is connected with the top centre of the source by a straight line and similarly the centre of the lower end of the cell and the bottom centre T 173 &Fotg irrad. no. &$'Ph R(13/11) 119 163 0.55 0.57 0.55 13 13 13 13 131 147 66 93 72 85 82 86 13 13 151 146 0.58 0.54 0.59 R( 14/ 12) 197 201 0.89 0.89 0.88 0.90 0.9 1 165 190 165 128 175 179 456 170 150 115 0.90 14 14 14 14 14 14 DECOMPOSITION OF HYDROGEN PEROXIDE 10 rev/min sector-(492 485) av.488.5 pmoles 02/1. after 15.67 h irradiation. Sector blank-84.0 pmoles 02/1. in 22.25 h after preparation. Sector yield corrected for blank-404.5 pmoles 02/1. in 15.67 h irradiation. Sector rate-25.85 pmoles 02/1. h. Steady irradiation-(448 464) av. 456 pmoles 02/1. after 6.0 h irradiation. Steady irradiation blank-85.0 pmoles 02/1. in 28-85 h after preparation. Steady irradiation yield corrected-371-0 pmoles 02/1. in 6 ~ 0 h irradiation. Steady irradiation rate-61.83 pmoles 02/1. h. 174 of the source are likewise connected by a straight line.Then a side surface of the aperture is cut through the shield by a line moving from the cell toward the source the line during this motion remaining parallel to the plane of the aluminium wheel and intersecting both guide lines drawn from cell to source. The sides of such an aperture will form a slightly curved surface so that a point in the cell will not begin to receive radiation simultaneously from the full length of the cobalt rod. (With a point source such curvature would be eliminated.) Due to this factor and to the scattering of radiation in the aperture it follows that when the wheel is rotating the radiation intensity at a given point in the cell does not change instantaneously from zero to full intensity as is assumed in the theory.14 The correction for this is considered later.Further with this experimental arrangement the sector scans the cell so that one part of the cell may be receiving radiation when another part is not. This however will not be important unless radicals diffuse rapidly throughout the cell. With regard to certain other assumptions of the theory the use of the C060 source is advantageous. For example the emitted radiation is constant in time and also across the aperture except at the edges. Further because the radiation is penetrating and because the mean distance from the source to the cells is large (45.4 cm) the ab- sorption of energy is nearly constant from front to back in the cell. SECTOR EXPERIMENTAL PROCEDURE.-FrOm each experiment the sector rate ratio (ratio of rate in intermittent radiation to rate in steady radiation) for one sector speed was ob- tained.Eight cells filled simultaneously were used in each experiment as follows. Two cells were exposed while mounted on the aluminium wheel which was rotated at a known speed. The starting exposure time was controlled exactly by lowering the source from a shielding lead turret to the position shown in fig. 1. Raising the source similarly ended the exposure. After the cells exposed to the intermittent irradiation were removed two other cells were located on the now stationary aluminium wheel and centred sym- metrically behind the aperture in the lead shield. These were exposed to steady irradi- ation for a measured time. Each set of cells was analyzed immediately after it had been irradiated.The other four cells stored in the dark were used as blanks to determine the thermal rate of decomposition. Generally two blank cells were analyzed at the same time as the cells from interrupted irradiation and the other two at the same time as the cells receiving steady irradiation. Occasionally one of the blank cells was analyzed at the start of the sectored irradiation. The oxygen yields of the blanks were plotted against time elapsed after filling of the cells. The plots indicated a small but relatively rapid thermal reaction (in addition to the analytical blank of 4 to 5 pmoles/l.) in the first hour or less followed by a much slower very slowly decreasing rate of evolution of oxygen. The thermal reaction being considered to be ionic the oxygen yields in irradiated cells were corrected by subtracting the amounts of thermal reaction corresponding to the times at which the irradiated cells were analyzed.(Allmand and Style also corrected for the thermal rate by simple subtraction.) These amounts were read from the above plots of thermal oxygen yield against time after cell preparation. From the corrected oxygen yields and irradiation times the rates of oxygen production for intermittent and steady irradiation were calculated. Calculation of a typical experi- ment follows Sector rate ratio-4-418. An experiment to test the assumption that the thermal blank is not a free radical reaction was carried out using hydroquinone. With 3.55-fold the intensity used in the sector experiments 0-1 mM hydroquinone in 0.1 M H202 reduced the rate of oxygen production to one-fifth that obtained with no inhibitor.The retarded rate was linear as far as followed and was estimated to correspond to a chain length of 3. On the other hand for the same concentrations the hydroquinone had little if any effect on the initial rapid thermal blank reaction occurring in the first 4 h but reduced the thermal rates about 28 % jn the period extending from 20 to 200 h. If the small thermal blank rate were all free radical the estimated chain length would be about 1000 so that the hydro- EDWIN J . HART AND MAX S . MATHESON 175 quinone by cutting this chain length to 3 would reduce the blank to a negligible amount. Thus it is concluded that not more than 28 % of the thermal blank is a free radical re- action.For the sample calculation above this would correspond to a dark reaction less than 1 % of the steady irradiation rate and from fig. 1 of ref. (14) it can be seen that such a dark rate has Iittle effect on the lifetime measurement. TREATMENT OF SECTOR DATA.-Frorn eqn. (14) derived in a later section from our assumed mechanism T ~ the kinetic chain life is inversely proportional to Roy the rate of oxygen production in steady irradiation. In essence this inverse dependence of R and T ~ assumes that at constant cell temperatures the variations of R are due to variations in the rate of initiation. Since the temperature difference between experiments varied 5°C as a maximum with most experiments within a 2" C range and since the temperature coefficient for such reactions is small the observed large variations in the steady irradiation rate of oxygen production obtained in the sector experiments cannot be accounted for through temper- ature effects alone.Further no correlation between R and temperature was observed. Experiments already described have shown that even with paraffined cells simultaneously irradiated with the quite constant C060 source the standard deviation is about 5 %. Thus it is possible that the steady rate variations between sector experiments may be due at least partially to impurity effects rather than to changes in the rate of initiation. Never- theless it seemed worth while to use the inverse relation of eqn.(14) to correct sector data to a common R, since this procedure gave a moderate decrease in the standard deviation of the results and had little effect on final rate constant calculations. Consequently all experiments were adjusted to a common R (60.0pmoles 02/1. h) by changing the 11's (11 = duration of intermittent radiation pulse) and R,'s inversely with respect to each other the corresponding sector rate ratios being left unchanged.14 The experimental sector rate ratios especially those for low speeds of rotation of the sector were somewhat high indicating either that an appreciable " dark " reaction was occurring in the " dark " part of the cycle or that the effective angle of the aperture was greater than the quarter circle intended. To test this point two experiments were carried out.First with the aluminium wheel stationary and with cells 1 and 2 centred behind the aperture ten cells of the ferrous sulphate-formic acid dosimeter 8 500 pN FeS04 + 0.1 M HCOOH + 0.8 N HzSO4 + 1-2 mM 0 2 were irradiated. Under the conditions employed the reaction in this dosimeter is nearly linear with dose.15 The following results show that in most of the shielded portion of the cycle negligible dark reaction would be induced in the hydrogen peroxide reaction by scattered radiation. 2 6 5 7 3 4 11 12 cell position 1 pequ i v. 9 10 117 115 61 1.1 0.15 0.0 0.15 - 0.0 - 0.6 54.6 Fe3+/l. h The cells on the edge of the aperture (3 and 12) received half as much dose as cells 1 and 2 as was expected. In the second experiment two cells of the ferrous sulphate dosimeter using 500 pN FeS04 -t 0.8 N HzS04 (air saturated) gave 5.72 pequiv.Fe+++/l. h when exposed to steady irradiation while centred behind the aperture. When the experiment was repeated with the wheel rotating the rate was smaller by the factor 0.261. From these experiments it was concluded that the effective angle of the aperture was 0.261 of 27r. Therefore the ratios measured for an irradiation period 0.261 of a cycle were corrected to the values they would have for the same R and sector speed but for an irradiation period 0.250 of a cycle. This essentially amounted to subtracting 0.022 from the experi- mental values of two x sector rate ratio since the theoretical curves of two x sector rate ratio against b (b = t l / ~ ) differ by approximately this amount for the above two sector apertures for all values of b.A further correction was required because as already noted the transition from zero to fuil intensity at the edge of the aperture was not instantaneous. An X-ray film was exposed just in front of the wheel. A Knorr-Albers microphotometer was used to measure the density of the developed film. The plot of film density against the angular position on the wheel showed that the radiation pulse could be fairly well approximated by a trapezoid. This type of approximation has been suggested elsewhere in sector work involving the use of visible or ultra-violet light.16 Burns and Dainton showed that for a trapezoidal rather than rectangular light pulse the correction to the theory is appreci- able only for slow sector speeds.For our case where the radiation period duration DECOMPOSITION OF HYDROGEN PEROXIDE 176 (“sector width”) was close to five times the period of increasing (or of decreasing) intensity (“ beam width ”) the limiting fast sector rate ratio was naturally unchanged while the limiting slow sector rate ratio was calculated to be raised to 0.27 for the trape- zoidal radiation pulse as against 0-25 for the corresponding rectangular pulse lasting one- fourth of a cycle. This increase at slow sector speeds is due to the dependency of the rate on the square root of the dosage rate. An approximate theoretical curve of two x sector rate ratio against b was drawn for the trapezoidal case of our experimental con- ditions.Then an experimental sector rate ratio (previously corrected to an effective sector aperture 0-250 of a cycle) was fitted to this curve the value of b noted and then the corrected rate ratio taken at the same value of b from the theoretical curve for a one- quarter cycle rectangular radiation pulse. All sector experiments were carried out in the temperature range from 25 to 30” C . RESULTS AND DlSCUSSION The y-ray initiated decomposition of hydrogen peroxide is a chain reaction in which water oxygen and a small amount of hydrogen are the products. The chain length of this reaction is altered by variables such as purity of the water and contamination by cell surfaces and in addition is dependent on dosage rate and concentration of hydrogen peroxide.In the present work the C060 y-ray dosage rate has been varied in the range (0.0038 to 0.63) x 1020 eV absorbed/]. min and the concentration from 0.008 to 200 mM hydrogen peroxide. In calculating energy absorbed Hochanadel’s 17 value for the FeS04 air-saturated dosimeter of 15.5 Fe2+ oxidized/100 eV was used. Except for minor changes the present work is in agreement with that of Fricke who reported that the oxygen yield is inversely proportional to the square root of the dosage rate and varies directly as the square root of concentration. A major difference exists however in the quantitative value of the yield under certain specified conditions. EFFECT OF DOSAGE RATE.-Owing to the existence of a non-chain primary decomposition of water yielding hydrogen peroxide and to the occurrence of a reaction of oxygen (the decomposition product of hydrogen peroxide) with the hydrogen atom the kinetics of this reaction become complex at low hydrogen peroxide concentrations.Dosage curves (oxygen production against time) are non-linear in the concentration range up to above 10 mM. Above this con- centration linear dosage curves result. Fig. 2 gives the results obtained on air-free unbuffered 32-8 mM hydrogen peroxide in the dosage rate range (0-0177 to 0.63) x 1020 eV/1. min. The yield plotted against the reciprocal of the square root of dosage rate gives a straight line which it is important to note does not extra- polate to the zero decomposition indicated by Fricke. Instead an extrapolation to 2.6 molecules 02/100 eV is obtained for very high dosage rates.Since 2.73 pairs of H and OH free radicals appear per 100 eV absorbed it is apparent that a chain reaction no longer exists at high dosage rates. Fig. 3 shows the results obtained on 100 mM hydrogen peroxide solutions (air-free and unbuffered) irradiated in paraffin coated cells. The dosage rate was varied in the range (0.0038 to 0.63) x 1020 eV/1. min. The slope in fig. 3 is greater than in fig. 2 since at this concentration the chain lengths of the decom- position reaction are much longer than for the experiments of fig. 2. Because of the steep slope the uncertainty in the intercept is as large as the expected value so that the line has simply been extrapolated to the origin. This curve is similar to fig.l b of Fricke. Johnson reports that the decomposition of hydrogen per- oxide at a concentration of 0.39 mM is independent of dosage rate in the range from (0.12 to 1.56) x 1020 eV/l. min. This is not in marked disagreement with the above results since Johnson with a yield of 4-3 molecules hydrogen peroxide decomposed/100 eV is clearly not in the concentration range where a chain reaction is obtained. The expected limit at high intensities is 4-8 molecules hydrogen peroxide/100 eV as can be calculated from the constant terms in eqn. (7) and (9) and using our extrapolated value of 2.6 molecules oxygen/100 eV EDWJN J . HART AND M A X S . MATHESON FIG. 2.-Effect of C060 y-ray dosage rate on the decomposition of unbuf- fered 32-8 mM hydrogen peroxide.abscissa (1018 eV/l. min)-’ ordinate molecules oxygen/ 100 eV. 100 eV. Likewise Hochanadel working at concentrations of 0.1 and 4.0 mM hydrogen peroxide obtained low yields indicating chains of the order of unity. A com- parison of the yields of several investigators with certain of ours is given in table 3. It is to be noted that at low concentrations there seems to be general agreement in the yields. However important differences appear at concentrations where the chain attains several units in length. At 10 mM hydrogen peroxide after adjustment is made for the effect of intensity and concentration it is found that at a dosage rate of 0.63 x 1020 eV/1. min 426 29.3 and 7-3 molecules hydrogen peroxide/100 eV are decomposed in Fricke’s Johnson’s and in our experiments 177 FIG.3.-Effect of C060 y-ray dosage rate on the decomposition of un- buffered 0.1 M hydrogen peroxide in paraffin coated cells. abscissa (1018 evil. min)-a. ordinate molecules oxygen/ DECOMPOSITION OF HYDROGEN PEROXIDE 178 respectively. Since measurements of chemical yield and dosage are far more accurate than this it is concluded that trace impurities are the main contributing factors to this divergence. At present we are not in a position to state whether the high yields or the low yields represent the impurity free reaction. conc. (H202)0 G(H202) molecules/ 100 eV mM 0.10 4.0 0.027 0*100 1.3 2.1 0.25 10.0 10.0 0.40 0.39 12.4 2.5 53.5 156.0 5.4 3.4 36-6 4.0 0.286 0.326 1.14 2-13 TABLE 3.-YIELDS IN THE DECOMPOSITION OF AQUEOUS HYDROGEN PEROXIDE BY IONIZING RADIATIONS dosage rate 3.7 3.7 Hochanadel 17 Hochanadel17 Toulis 18 Toulis 18 Toulis 18 Cob0 y-rays Co60 y-rays X-rays X-rays X-rays X-rays X-rays X-rays X-rays X-rays Co60 y-rays Co60 y-rays C060 y-rays Co60 y-rays Co60 y - r a y s 3.3 4.1 7.8 9.0 10.5 EFFECT OF CONCENTRATION.-The effect of hydrogen peroxide concentration was studied in air-free unbuffered solutions at concentrations in the range from 0.008 to 200 mM.At concentrations below 0.02 mM the dependence changes but above this concentration the rate of decomposition depends on the square root of the concentration.The points on the curve of fig. 4 were obtained at a 0.63 /OO 300 200 dosage rate of 0.63 x 1020 eV/l. min. However enough points were obtained at a dosage rate of 0.0038 x 1020 eV/1. min to show that this Sam dependence holds here too. It is believed that this dependence will hold at any dosage rate where an average chain of even a few units is obtained. This follows from the fact that the average chain is only 3.7 units in length at the highest point in fig. 4. (1020eVil. min) 0.0054-2.47 0.0054-2.47 0.34 0.34 0.34 0.38 0.38 0.89 0.117-1.68 0.89 0.0 I76 0.63 0.0176 0.63 400 radiation author Fricke 3 Fricke 3 Johnson 4 Johnson 4 Johnson 4 this work this work this work this work this work EDWIN J.HART A N D MAX S . MATHESON 179 This dependence of yield on the square root of hydrogen peroxide concentration was also found in the previously cited work of Fricke and Johnson. There is definite evidence in the present work that the form of the concentration curve changes below 0.02 mM hydrogen peroxide. However all of this work on hydrogen peroxide concentration was carried out using standard glass irradi- ation cells. Therefore the effects of impurities from the glass wall would have a particularly pronounced effect at these low concentrations. Further work should be carried out in paraffin coated cells. MECHANISM.-Previous investigators have provided suitable mechanisms for initiation and chain propagation processes.A bimolecular termination step correctly giving the observed relation between yield and the inverse square root of intensity has also been proposed. However an over-all mechanism must also explain the dependence of yield on square root of concentration as found in the photochemical work of Allmand and Style in the X-ray work of Fricke and Johnson and in our y-ray work. An explanation must also be provided for the extrapolated value of 2.6 oxygen molecules produced/100 eV at high dosage rates as is shown in fig. 2. These facts can be explained by the following mechanism rate (la) Ha for decomposition H20 = H + OH H2O = 1/2 Hz + lj2 H202 (Ib) K’L] of H2O H 4- H202 H2O -+- OH OH -1- H702 = HO2 + H20 HOz I- P I 2 0 2 2HO2 -t Hz02 -= 2H202 -/ 0 2 of the K02.(2) Kla H20 + OH -4- 0 2 (4) kp(H02)(H202) (3) kp’(OH)(H>Olj (5) 2kt(M02)?( HzO?)-for disappear- ance of HO2. This mechanism postulates that reaction (4) is much slower than (3) since termination is by hydroperoxy radicals and not by hydroxyl radicals. Evidence for this has been discussed by Agar and Dainton.19 These authors have also pointed out that inhibition of the photolysis of H202 by alkali indicates that H02 and not 02- is involved in the propagation reaction. (Using the KHoz and KH20z assumed by these authors the concentration of H02 would be about three times the 02- concentration in our sector experiments.) Lea2 also found the quantum yield independent of pH in the pH range 1-6 suggesting that 02- or H02- is not involved.Further evidence that HO;! + H202 is relatively slow is found in the work of Barb Baxendale George and Hargrave,Zo who find that the H02 i- H202 reaction does not take place in the presence of ferrous or ferric ions. The results of Bray 21 and George 22 on KO2 dissolved in water can be accounted for by the above mechanism if it is assumed that the solution and re- action of KO2 produces HOz radicals. The ratio of reaction (5) to reaction (4) is Zkt(HOz)/k independent of H202 concentration and is equal to 1.0 x lOs(H02) from our rate constant measurements. Thus if the HO2 concentration is 10-9 molar or higher near the surface of the dissolving KO2 particles reaction (4) will occur to an extent which is 1 % or less than that of reaction (5) and in reaction (5) the net effect is the production of one H202 and one 0 2 molecule from two H02 radicals.This is also the observed stoichiometry for KO2 dissolving in water. If the assumption that KO2 reacts with water to yield H02 is correct con- centrations of H02 in excess of 10-6 M might well be expected near the surface Ry assuming that a steady state of intermediate free radicals is readily estab- lished in the above mechanism eqn. (6) is derived Ro = dOz/dt = kp(KI,(Ha02)/kt)3 t KI,. (6) 180 (7) - Y(H202) = - K’/2K + 2 + 2kp((H202)/ktKTa)*. According to Hochanadel the ratio K’IK is 0.34 for the hydrogen peroxide (8) (9) = 1.83 + 2kp((H202)/krKIn)h. Primary steps (la) and (1b) have been discussed recently in the papers of Allen H-0 O/o-H (4’) (5 ’> DECOMPOSITION OF HYDROGEN PEROXIDE After dividing both sides by KIu in order to express the results in terms of yield per unit energy required to produce one radical pair in (la) one obtains Y(02 ) = (d02/dt)(l /K&) = kp((H202)/ktKLJ) + 1.The expression for hydrogen peroxide disappearance becomes system. Thus the above equation reduces to O H 0 4 \O-H - Y(,,) and co-workers,23 Hochanadel,l7 and Hart.24 The present work provides addi- tional evidence for the existence of step (16). Hydrogen gas is always found in the radiolysis of hydrogen peroxide solutions. The yield is substantially inde- pendent of dosage rate and of hydrogen peroxide concentration in agreement with the above mechanism. The Haber-Willstatter 25 propagation steps (3) and (4) account for the chain decomposition and have been widely applied in thermal and photochemical mechanisms for this reaction.Our termination reaction (5) appears as a termolecular one and differs from previous ones by including hydrogen peroxide in this step. The inclusion of hydrogen peroxide in the termination reaction introduces the correct dependence between rate and square root of hydrogen peroxide concentration. Although present work offers no proof we suggest the formation through hydrogen bonding of an intermediate complex between hydrogen peroxide and the hydroperoxy radical as follows This complex intermediate may decompose as in reaction (4) or if formed in the presence of another hydroperoxy radicaI could yield reaction (5) through a larger intermediate.In either reaction a simple regrouping of electrons and moderate adjustment of interatomic distances is all that is necessary for reaction as in (4’) and (5’). H-O/ 0 //O H\o H-O/ /0-H 0 O/o-H H-0 /o-H H-0 0-H -+ ;/O+ (J 1) (12) I If reaction (3) were the slow step in propagation so that the termination reaction corresponded to eqn. (11) then eqn. (12) would be obtained. OH + OH + H202 = 2H202 (2kr’(OH)2CH202)) Y(02) = kp’((H202)/wm*. The intercept in fig. 2 predicts that hydrogen peroxide would be decomposed at a rate yielding one oxygen molecule per radical pair at infinitely high dosage rates since the data of Hart 24 show that 2-73 water dissociations occur in reaction (la) per 100 eV of C060 y-ray energy absorbed in water.This is in accord with eqn. (7) but not with (12). Therefore our data definitely favour termination step (5) and not (11). This experimental evidence provides an additional argu- ment for claiming that the rate constants measured in our rotating sector experi- ments are kp and kt and not k,’ and kt’. Reaction (11) is not written as 20H + H202 == 2H20 + 0 2 181 (H202) corr. to corr. to aperture rectangular EDWIN J . HART AND MAX S . MATHESON since 0 2 is not produced in the “ hot spot ” or “ end of the track ” reaction but H202 is presumably from OH radicals. RATE CONSTANTS.-The experimental results and corrections of the sector experiments are listed in table 4. Twice the sector rate ratios corrected to a rectangular radiation pulse lasting one-fourth of a cycle are plotted in fig.5 against the Iogarithms of the radiation pulse times. The theoretical curve shown in the same figure is plotted against log b. The ratio tl/b corresponds to a life- time of 0.60 sec. This value was calculated by averaging the logarithms of the lifetimes of the last column of table 4 weighting the less reliable starred results one-half as much as the others. The lifetime obtained from the uncorrected ratios (column 6 table 4) is 0.83 sec. TABLE 4.-LIFETIMES OF KINETIC CHAINS 2 x sector rate ratio RO corr. to Ro = 60.0 .umoles/l. h 50-3 62.0 59.0 84.1 61.83 expt. 0.025 7.5 7.5 0.625 1.50 7.5 0.025 0.025 3.00 1.50 159.9 lifetime zs sec.- 1-83 0-406 0.559 1-39 - - 0.339 0.646 pulse 1-016 0.557 0.742 0.833 0.796 0-713 0.978 1-022 0.642 expt. 22 23 24* 25 26 27* 28 29 31 32 33 mM 85.6 92.0 86.4 92.2 95.0 98.4 93.6 100.4 94.2 100 100 expt. 0.754 1 so00 1.044 0.696 0-846 0.624 0.250 1.016 0.598 0.758 0.850 0.814 0.732 0.978 1.022 0.674 0.824 0.602 pmoles/l. h 0.02095 1.038 0.620 0.780 0-872 0.836 7.75 7-37 0.876 1.547 6.94 0.0169 0.0248 3.25 1-659 142.2 55.5 40.56 59-6 65.0 66.3 53-35 0.807 0.562 *Because of experimental difficulties the results in these experiments are not con- sidered as reliable as for the other experiments.FIG. 5.-Rate of oxygen production in intermittent radiation for different radiation pulse durations tl ; curve theory for rectangular radiation pulse lasting one-fourth of a cycle and for zero dark reaction. abscissa b = (duration of radiation pulse)/(kinetic chain life) = f l / T S . ordinate 2 x sector rate ratio. @ our points on single cells. 0 our points on duplicate cells. Lea’s points (ref. (2)) corrected to our rates and for non-chain portion of reaction. DECOMPOSITION OF HYDROGEN PEROXIDE 182 CALCULATION OF RATE CONSTANTS.-The fOl1OWing eqn. (13) and (14) show that two sets of experimental data suffice to obtain the individual rate constants R$/Ri(H202) == kp2/2kt where Ri -- 2KI = rate of initiation of chains and R o ~ s = kp/2kt.Eqn. (13) is obtained by rearranging eqn. (6) and assuming long chains and eqn. (14) may be derived from eqn. (13) using the following relations conc. of free radical chains = (RiT,) conc. of free radical chains N_ (HOz) if H02 + H202 slow propagation step (16) (1 7) Ro = kp(H02)(H202). occurs with an effective primary quantum yield f. Then d(H202)/dt = 2kp(H202)'(2f~a~,/2kt)' + 2 fIabs. (1 5 ) Considering first the data needed for eqn. (13) the dose rate for steady ir- radiation of cells on the aluminium wheel centred behind the aperture corresponds to 5.72 pequiv. Fe3+/l. h produced in a 500,uN FeS04 + 0.8 N H2S04 (air saturated) dosimeter.This dose rate decomposes 1 *006 pmoles H20/1. h to radical pairs to give Rj = 5.59 x 10-10 mole radicals/l. sec. This radical pair yield for a given dose rate is estimated from the data of Hart24 on the inhibition by 0 2 of the HCOOH + H202 reaction. The rate of production of radical pairs is taken as equal to the rate of 0 2 consumption in this inhibition. In this work 0.153 mmole water dissociations/l. h to yield hydrogen atoms and hydroxyl radicals were obtained under irradiation conditions where 0.870 mequiv. Fez+/]. h were oxidized. Corresponding to the value of Ri above average values (@)+ = 1.665 x 10-8 mole/]. sec and (H202) -~ 93-26 x 10-3 mole/l. were obtained from the data of table 4. These data and eqn. (13) give kp2/2kt = 5.32 x 10-6 sec-1.Further eqn. (14) and the average Ro value found above for the sector experiments with the kinetic chain lifetime of 0.60 sec found in the same experi- ments gives kp/2kt = 530 1. mole-1 sec-1 and 2kt = 5-3 x 1010 1.2 mole-2 sec-1. Tn the sector experiments in steady irradiation (Ro - 1.66 x 10-8 mole/l. sec) the concentration of H02 is 3-4 : 10 10 mole/l. and the chain length is 30 mole- cules of 0 2 per initiating radical. The only results in the literature which may be compared with ours are those of Lea.2 The comparison can be made by assuming steps (3) (4) and (5) of our mechanism and by assuming that in photochemical initiation Lea's reaction (l) From Lea's results at high intensities 2f= 1.39 = the limiting quantum yield and from his results at low intensities usingf= 0-7 kp2/2kt = 11.1 x 10-6 sec-1.I f f = 1-0 then kp"2kt is found to be 7-8 x 10-6 sec-1 which does not differ appreciably from our value of 5.3 x 10-6 sec-1. In fig. 5 are plotted the results of Lea's sector work corrected to the same steady rate of reaction as used by us. These ratios have also been corrected for the non-chain portion of the H202 decomposition by subtracting 2JTabs (f= 0.7) from the rate in steady illumination and (fIabs)/2 from the rates in sectored light. The authors wish to acknowledge their indebtedness to Messrs R. A. Blomgren and L. S. Markheim for the design and installation of the rotating sector and to Miss P. D. Walsh for technical assistance. EDWIN J . HART AND MAX S . MATHESON 183 APPENDIX Received 10th June 1952 NEW EXPERIMENTAL msuLTs.-Hydrogen peroxide or its complex with the hydroperoxy radical has been postulated to participate in the termination step of the 7-ray induced decomposition of hydrogen peroxide.This conclusion has been drawn from the dependence of the rate of decomposition on the square root of hydrogen peroxide concentration. A first order dependence would support the customary bimolecular termination previously employed.25 The X-ray work of Fricke 3 and Johnson 4 and part of the photochemical work of Allmand and Style 9 and Kornfeld 26 are consistent with the square root dependence. While the X-ray and y-ray work clearly shows the square root dependence a first order reaction is frequently reported in photochemical work.2927 Impurities are sus- pected to be the cause of this discrepancy since they are known to cause important changes in the kinetics of hydrogen peroxide formation and decomposition.Since our sector work was carried out under conditions where less than 0.5 % decomposition occurred it was felt desirable to re-investigate the order of the reaction under condilions of more exhaustive purification and greater degree of reaction. EXPERIMENTAL PURIFICATION OF HYDROGEN PERomE.-Our previous purification consisted of a steam distillation from acid solution of Merck reagent grade 30 % hydrogen peroxide. Since this product was identical in radiolysis behaviour to inhibitor free Buffalo Electro-Chemical 90 % hydrogen peroxide and gave the square root dependence observed by Fricke and Johnson it was thought that our hydrogen peroxide was of sufficient purity for the sector FIG.1.-Apparatus for the distillation of 90 % hydrogen peroxide. experiments. Tn order to provide further information on this point a purification pro- cedure recommended by Shanley 28 of the Buffalo Electro-Chemical Co. was carried out on B.E.C. 90 % hydrogen peroxide. Sodium stannate was added to the still to neutralize nitric acid believed to be the only volatile material present besides water and hydrogen peroxide. A one-plate distillation was carried out on 175 g of hydrogen peroxide in a still maintained at 35" C (see fig. 1). The system was evacuated through F by a Hypervac pump with dry ice on traps C and D. E is a medium fritted glass filter.After one-third 184 of the hydrogen peroxide had distilled into C dry ice was placed on trap B also and the middle fraction collected and used in the radiolysis experiments. Trap A was added to catch any hydrogen peroxide carried over by entrainment of droplets. However no bumping of the liquid was observed and the distillation proceeded very smoothly over a period of about 6 h. The hydrogen peroxide (42.5 g) collected in B was diluted to 100 ml with triply distilled water in a heat treated volumetric flask and this stock solution of 13.5 M hydrogen peroxide was used in the radiation experiments. On the basis that an inhibitor might still be present in the vacuum distilled hydrogen peroxide 10 ml of the 13.5 M hydrogen peroxide was irradiated until 6-23 % decomposi- tion occurred.This pre-irradiated hydrogen peroxide was then used in some of the irradiations run at 0.1 M. ANALYSES.-TWO methods of analyses were used to measure the decomposition of hydrogen peroxide by y-rays. In one the Van Slyke apparatus was employed to measure evolved oxygen where the total decomposition was a fraction of 1 %. The second method employed ceric sulphate using ferroin as an indicator. In this method 0.100 to 2-00 ml of hydrogen peroxide was added to 10 ml of 0.8 N sulphuric acid and titrated directly with 54.64 mN ceric sulphate which had been standardized with pure iron wire. Reproducibility by this method is better than 0.5 %. PREPARATION OF CELLS AND IRRADIATION OF soLuT1oNs.-For Use when Oxygen eVOlU- tion was measured cells of the type employed for the sector experiments were coated with paraffin in the manner previously described.In a separate experiment as a further treatment designed to neutralize the alkaline surface of the glass a set of cells was given a fuming sulphuric acid treatment suggested by Dr. Shanley. In this method 2 to 3 ml of fuming sulphuric acid were drawn into furnace heat-treated cells and the surface wetted by this material. After standing 10 min the cells were rinsed several times with triply distilled water in order to remove the sulphuric acid. In the irradiations carried out to nearly compIete decomposition (followed by H202 analysis) 80 ml paraffin coated cells were used. During the irradiation there was no indication that the wax surfaces had ruptured.The analyses were performed on 2 ml samples withdrawn from the irradiated 80 ml solution. Under the conditions of irradi- ation up to 70 % of the solution could be withdrawn without appreciably affecting the dosage rate delivered to the solution irradiated. DECOMPOSITION OF HYDROGEN PEROXIDE one half order (moles/l.>?n-' init. conc. moles/l. 8-57 x 10-2 RESULTS ORDER OF THE REACTION.-The order of the y-ray induced decomposition was investig- ated using redistilled 90 % hydrogen peroxide at concentrations of 0.0314 0.1002 0.306 and 1-008 M. These solutions were decomposed by irradiation in 80 ml paraffin coated cells until a series of overlapping concentration ranges was obtained. The forms of the individual curves demonstrate clearly that the experimental data are consistent with a dependence of rate of decomposition on (H202)0-5 instead of (H202)l.O (see fig.2 and 3). In general deviation from the square root dependence is not observed until 80 % reaction has occurred whereas deviation from first power kinetics is observed at 20 "/ decomposition. A comparison of the one-half and first order rate constants is shown in table 1 for the initial rates of decomposition for the data of fig. 2 and 3. A maximum of 19 % deviation from the mean occurs for the one-half order rate constants whereas the first order rate constants show a five-fold increase as the concentration of hydrogen peroxide is lowered from 1-008 M to 0,0314 M. In no case was any sign of an induction period indicative of inhibitors found.Eqn. (9) of our paper for the rate of decomposition of hydrogen peroxide indicates that for this chain reaction there is a small constant term (1.83 x rate of production of TABLE 1 .-INITIAL RATE CONSTANTS FOR H202 DECOMPOSlTION AT CONSTANT INTENSITY (CONSTANT RATE OF INITIATION) first order I?-' 0.246 x 10-2 0.339 x 10-2 0.781 x 10-2 6-05 x 10-2 8.82 x 10-2 1.008 0-306 0.10025 0.03 142 1-273 x 10-2 6-52 x 10-2 7.43 x 10-2 av. kp2/2kt = 10.76 x 10-6 sec-1 from av. half order constant. kp2/2kt = 5.32 x 10-6 sec-1 from rates in sector experiments. radical pairs) in addition to the (Hz02)+ term. If the complete rate expression is integrated the slopes of fig. 2 may then be corrected for an additional logarithmic term.This term is most important for the lowest concentration where the chains are the shortest. Here the correction decreased the slope by 10 :,/ without however affecting the linearity. 185 Y I 0 1.008 M. EDWIN J . HART AND MAX S . MATHESON 80 40 1 80 240 200 200 0 Q 0.306 M. 40 FIG. 2.-Dependence of the rate of hydrogen peroxide decomposition on (H202)* at a dosage rate of 3.98 1020 eV/1. hr. 0 0.10025 M. The redistilled hydrogen peroxide showed a dependence of rate of decomposition on (H202)0.5 in irradiations carried out to 0.4 % reaction in paraffined cells of the type used in the sector experiments. The Same dependence was obtained on distilled 90 % hydrogen FIG. 3.-Dedendence of the rate of hydrogen peroxide decomposition on (H202)l.O at a dosage rate of 3-98 x 1020 eV/I.hr. 0 0.306 M. 160 TIME IN HOURS 0 0*10025 M. 160 I 2 0 TIME IN HOURS I I 2 0 0.5 - < 0.3 - 3 0 0.4 - %I L I 0 01 D( 5 0.2 - 0.1 0 1.008 M. peroxide which had been irradiated to 6-23 % decomposition and then diluted. The data for 0.032 and 0.10 M appear in fig. 4. This result again demonstrates that trace impurities removable by irradiation are not present in our solutions. c) 0.03142 M H202. I \ I I 1 280 0.03142 M H202. 240 280 186 Cells pre-treated with fuming sulphuric acid demonstrated the same erratic behaviour previously noted for freshly heat-treated cells. In general the amount of hydrogen per- oxide decomposed on irradiation is about one-half that found for the paraffin-treated cells.The results obtained were too erratic to distinguish between one-half and first order reactions. EFFECT OF WATER.-standard triply distilled water Barnstead still water and laboratory distilled water were used in the preparation and irradiation of 0.10 M hydrogen peroxide. The Barnstead water is prepared from tap water by a single distillation and is used as the first stage of our triply distilled water. The laboratory distilled water is prepared by a double distillation in Barnstead stills but in addition is delivered to the laboratory benches in aluminium tubing. Table 2 gives the rate constants obtained by irradiation at 3.98 x 1020 eV/1. h.The kinetics appear to conform to (H702)0.5 although a complete investigation was not made on the Barnstead watcr not- the laboratory distilled water. The Barnstead water shows a moderate increase in rate whereas the rate of the laboratory water is nearly 8 times that of our triply distilled water. DECOMPOSITION OF HYDROGEN PEROXIDE 0 distilled H202. TABLE 2.-EFFECT 0.5 order rate constant (m/M)+ h-' 8.82 x 10-2 FIG. 4. FIG. 4.-Effect of conceiitration on the decomposition of hydrogen peroxide at dosage rate of 7.6 x 1019 eV/1. hr. distilled and pre-irradiated H202. OF WATER ON THE ONL-HALE ORDER KATE CONSTANTS IK ?HE DECOMPOSITION OF HYDROGEN PEROXIDE W 2 0 2 ) o 0.10025 0-0976 type of water triply distilled Barnstead laboratory distilled 11.6 x 10-2 46-4 >< 10-2 0.0972 A dark reaction that is small compared to the y-ray reaction also takes place in each of the hydrogen peroxide solutions.However with the laboratory distilled water 85 % of the hydrogen peroxide is decomposed in 300 h standing at room temperature. This high rate of decomposition is indicative of an impurity catalyzed decomposition and the decomposition rate is first order with respect to hydrogen peroxide concentration. This behaviour is similar to that observed by Barb Baxendale George and Hargrave 20 in experiments at high ratios of peroxide to ferric ion concentrations. Thus it is possible that certain types of impurities when present in water lead to first order kinetics. In- sufficient work was done to determine whether this behaviour is carried over into the y-ray irradiated solutions.EDWIN J . HART AND MAX S . MATHESON 187 DISCUSSION THIRD ORDER TERMINATION IN HYDROGEN PERoxmE.-Since the experimental evidence clearly favours a rate of decomposition proportional to the square root of absorbed radiation and to the square root of hydrogen peroxide concentration any mechanism adopted must account for these facts. In our mechanism the correct dependence on radiation intensity and ( H 2 0 2 ) is obtained through the termination reaction HO2 +- H 0 2 t H202 -+ 2H202 -t 0 2 . HOz -+ H202 -+ [ H 0 2 - H202J (5) Since only a few termolecular reactions are known it is pertinent to consider the proposed reaction in more detail. First it may be noted that the work of Swain 29 indicates that third order reactions are not as rare as has been generally believed.He has shown that in the enolization of acetone in water solution and in the dis- placement reactions of organic halides in benzene solution three reactants are involved in the rate determining steps. Secondly it is not necessary that termination in our mechanism occur by a true termolecular reaction but only that the termination be effectively third order. This third order effect can be derived if the formation of complexes such as those in equations (4') and (5') of our paper on hydrogen peroxide is postulated. Thus steps (4) and ( 5 ) of the proposed mechanism may be replaced by the following (44 (4b) (44 (18) (19) [HOz - H2021 + HO2 + H202 kc(H02)(H202) [HOz - H 2 0 2 ] - + H 2 0 + 0 2 + OH k4c([H02 - H2021) kd(LH02 - H2021) [ H 0 2 - H 2 0 2 ] $- HO2 + 2H202 -t 0 2 k5a([H@ - H2021)(H02) (5a) The steady state assumption and certain others may be applied to the rnechan- ism as now modified.Three such cases will be considered here. Case 1. Steps 4a and 4b are taken to be much faster than 4c or 5a i.e. the complex attains an equilibrium concentration with respect to HO2 and H202 Then the rate of 0 2 evolution is dOz/dt = k4c(kc/kd)%(Kl,/ksu)~(H202)% + Kla as required by the experimental facts. If in addition kc(H202) < kd then the kp we have measured is k4kc/kd and 2kt = 2ksakc/kd. This is the situation if only a small portion of the H02 is in the equilibrium concentralion of complex.In the limiting case of short-lived unstable complex formation kc would become the rate constant for formation of activated complex and k4c/kd would become the transmission coefficient for reaction. Case 2. If kd = 0 so that irreversible formation of complex occurs the rate is given by d02/dt == k,tlt~c(Kr,/k5a)'(H202)' + Kla. Next let kc(H202) > kqc so that rapid formation of complex is followed by slow decomposition. Then k p = k4c/(H202) 2kr = 2k5ak4c/[kc(H20#] and the rate constants as measured would decrease with increasing H202 concentration. Case 3 . If kd = 0 and k4c > kc(H202) so that slow irreversible formation of complex is followed by rapid decomposition then kp = k and 2kt = 2 k ~ ~ k ~ / k 4 ~ .From the above discussion it appears that a termolecular termination real or apparent is reasonable. It remains then to consider the magnitude found for the termination rate constant. Possibly the highest termolecular rate constant for the gas phase is the 1010 1.2 moles-2 sec-1 measured by Smallwood30 for the number of effective collisions of three hydrogen atoms. In comparison with this value our k r == 2.7 X 1010 1.2 moles-2 sec-1 is somewhat higher. However this is not an impossibly high value even for a true termolecular reaction as can 188 DECOMPOSITION OF HYDROGEN PEROXIDE be seen from the following argument. If two HO2 radicals collide in aqueous solution they will have 10 or more molecules as neighbours in the cage around them. Since pure H202 is about 40 M this would mean that in a 0.1 M H202 solution one in every 40 pairs of colliding H02 radicals would have an H202 molecule as a neighbour.Further the rate constant for bimolecular collisions is about 3 x 1011 1. moles-1 sec-1 so that the rate at which H02 pairs would collide in the presence of Hz02 (for 0.1 M H202) is 3 x 1011 x (1/40) x (H02)2 = 0.75 x 1010 x (H02)2 moles 1.-1 sec-1. From our results kt(H202)2 = 2.7 x 1010 x 0.1 x (H02)2 = 0-27 x 1010 moles L-1 sec-1 indicating that about one in three collisions would lead to termination which is a rather high efficiency. On the other hand complex formation may be present to account for such high efficiency. 1 Allmand and Style J. Chem. Soc. 1930 596. 2 Lea Trans. Faraday SOC.1949 45 81. 3 Fricke J . Chem. Physics 1935 3 364. 4 Johnson J. Chem. Physics 1951 19 1204. 5 Volman J . Chem. Physics 1949 17 947. 6 Hart J . Amer. Chem. SOC. 1951 73 68. 7 Lindsay Introduction lo Physical Statistics (John Wiley and Sons Inc. New York N.Y. 1941) p. 15. 8 Orr and Butler J. Chem. Soc. 1935 1273. 9 Briers Chapman and Walters J . Chem. SOC. 1926 562. 10 Matheson Auer Bevilacqua and Hart J. Amer. Chem. Sue. 1951 73 5395. 11 Grassie and Melville Proc. Roy. SOC. A 1951 207 285. 12Bateman and Gee Pruc. Roy. SOC. A 1948 195 391. 13 Bamford and Dewar Proc. Roy. SOC. A 1949 198,252. 14 Matheson Auer Bevilacqua and Hart J . Amer. Chem. Sue. 1949 71 497. 15 Hart in preparation. 16 Burns and Dainton Trans. Faraday Soc. 1950,46,411. 17 Hochanadel paper presented at 119th Amer.Chem. SOC. Meeting (Cleveland Ohio 1951). 18 Toulis University of California Radiation Laboratory Report No. 583 Feb. 10 1950. 19 Agar and Dainton Faraday SOC. Discussion 1947 2 218. 20 Barb Baxendale George and Hargrave Trans. Faraday Sue. 1951 47 462. 21 Bray J. Amer. Chem. SOC. 1938 60 82. 22 George Faraday SOC. Discussion 1947 2 196. 23 Allen Hochanadel Ghormley and Davis paper presented at 119th Amer. Chem. SOC. Meeting (Cleveland Ohio 195 1). 24 Hart paper presented at 119th Amer. Chem. SOC. Meeting (Cleveland Ohio 1951). 25 Haber and Willstatter Ber. 1931 64 2844. 26 Kornfeld 2. tviss. Phot. 1921 21 66. 27 J. Weiss private communication. 28 E. S. Shanley private communication. 29 Swain J .Amer. Chem. Soc. 1952 72 2794 4578. 30 Smallwood J . Amer. Chem. Suc. 1934 56 1542. JOSEPH WEISS 169 MECHANISM AND RATE CONSTANTS OF THE y-RAY INDUCED DECOMPOSITION OF HYDROGEN PEROXIDE IN AQUEOUS SOLUTIONS BY EDWIN J. HART AND MAX S. MATHESON Argonne National Laboratory Chicago 80 Illinois U.S.A. Received 28th January 1952 A mechanism for the y-ray initiated decomposition of dilute aqueous solutions of hydrogen peroxide has been deduced from data showing a dependence of decomposition yield on the square root of hydrogen peroxide concentration and inverse square root of dosage rate. A novel feature of this mechanism is that termination occurs through a termolecular reaction involving two hydroperoxy radicals and a hydrogen peroxide mole-cule. The propagation and termination rate constants have been measured in inter-mittent radiation experiments which were carried out in paraffin coated cells containing 0.1 M hydrogen peroxide under irradiation conditions yielding 30 molecules of oxygen per initiating radical.Evidence is presented showing that HOz + H202 = H20+02+OH and 2H02 + H202 = 2H202 + 0 2 are the rate determining propagation and termination steps. kD is found to be 530 1. mole-1 sec-1 and 2kt equals 5-3 x 1010 1.2 mole-2 sec-1. The chain character of the photochemical and X-ray initiated decomposition of hydrogen peroxide has been discussed by Allmand and Style,l Lea,2 Fricke 3 170 DECOMPOSITION OF HYDROGEN PEROXIDE and Johnson.4 Fricke reported a dependence of yield on the reciprocal square root of the dosage rate for X-ray irradiations.This result is evidence that two chain carriers are involved in the termination step. Allmand and Style applied the intermittent light technique to measure the lifetime of the kinetic chain. It was postulated that the hydroxyl free radical was the catalyst present and the chain was found to have a mean life of 1 sec under the conditions these authors employed. The termination step proposed was OH + OH = H202 (or M20 -t i 0 2 ) . However no proof was offered for this step. More recently Lea 2 and Volman 5 have employed bimolecular termination of two hydroperoxy radicals. However, the radiolysis of hydrogen peroxide solutions has been shown by Fricke and Johnson to be dependent on the square root of hydrogen peroxide concentration.In agreement with this Allmand and Style find the relative quantum yield to be constant below one M H202 with 3650A light of constant intensity. This result is expected if the rate is proportional to (H202)i (labs)* and Iabs is proportional to (H202). These results are not explained by any of the mechanisms proposed. The purpose of the present paper is to report work in which the rotating sector was applied to the decomposition of hydrogen peroxide by y-rays in order to measure rate constants for the propagation and termination steps. A further object is to clarify if possible the mechanism of this reaction. EXPERIMENTAL PURIFICATION OF MATERIALS.-The methods previously described 6 were used for the preparation of the distilled water and for evacuation of solutions to be irradiated.Ex-haustive irradiation of such air-free triply-distilled water samples produced of the order of one micromolar carbon dioxide. This small carbon dioxide yield was the criterion used for testing the purity of the water the cleanliness of the irradiation cells and the general technique of removing air from the solutions. As shown earlier less than one micromole of oxygen per litre of water remains in the solutions after the evacuation procedure. While the complete removal of oxygen is not critical in this hydrogen per-oxide work careful attention must be taken to insure the removal of impurities from the water and the cell surfaces. Merck reagent grade 30 % hydrogen peroxide and Buffalo Electrochemical 90 % hydrogen peroxide were used in the present work.The former was purified by steam distillation and the latter used without further purification. The results appearing in fig. 2 and 4 were obtained using the steam distilled hydrogen peroxide. All of the work in paraffin coated cells at or near 0.100 M hydrogen peroxide was carried out on the 90 % product. In these cases a stock solution of 15 M hydrogen peroxide was added to freshly distilled water in the evacuation chamber which had previously been furnace heated to 550" C. Subsequent degassing of the solution and filling of the cells was usually carried out within 15 min. ANALYSES.-The Van SIyke micro-gas analysis apparatus was used for measuring the amount of oxygen formed in the irradiation cells. For concentrations of one millimolar, no problem with respect to the decomposition of hydrogen peroxide during analysis was encountered.However the regular method of analysis could not be used for 0.10M concentrations. At these concentrations attack on the mercury in the Van Slyke by the hydrogen peroxide is rapid unless the solution is made strongly acid. In order to circum-vent this difficulty and to condition the surface of the mercury the Van Slyke was evacu-ated with 1 ml of 0.1 M hydrogen peroxide in 18 N sulphuric acid in the chamber. After evacuation 0.5 ml of the gas-free solution was left in the chamber. The mercury and this solution were then lowered to the 50 ml line and the hydrogen peroxide from the irradiation cell introduced. For 0.1 M hydrogen peroxide the gas analyzed consisted of 99.5 % oxygen the remainder presumably being hydrogen.Samples of 0-1 M hydrogen peroxide analyzed 10 min after evacuation showed 4 to 5 pM oxygen. This small amount of gas is believed to be due to the decomposition of hydrogen peroxide during analysis ; however it could be due to an initially high catalytic decomposition by the walls of the irradiation cells EDWIN J . HART AND MAX S . MATHESON 171 ments were carried out using 0.1 M hydrogen peroxide in order to obtain the long chains assumed in the mechanism. At this concentration considerable difficulty was en-countered in obtaining consistent and reproducible results. After carrying out the series of experiments appearing in table 1 it was concluded that the irregularities in results were due to impurities in the water probably originating in the glass surface.The best solution found to this problem was to coat the inside of the irradiation cells with paraffin, a procedure also used by Allmand and Style in some experiments. A non-crystalline paraffin (Bioloid Embedding Paraffin m.p. 60-62" C Will Corp. Rochester New York) was found satisfactory. Two sets of eight cells each were used in alternate experiments in the sector work. The inner walls of the cells were coated with paraffin as follows. The furnace-treated cells were warmed to 100" C in an oven and then about 0.5 ml melted paraffin drawn by suction through the standard taper capillary tip into the cell. The cell was then rotated until the entire surface except for that in the thin capillary tip was covered. The excess paraffin was then drained through the standard taper tip.After cooling to room temperature these paraffined cells were ready for use. After use the paraffined cells were recleaned by alternate treatments with methanol and benzene until the paraffin was dissolved. The cells were then further freed from organic impurities by reheating to 550" C . The effect of treatment of the cells and purity of the water on the decomposition of 0.10 M hydrogen peroxide by y-rays is given in table 1. The treatment that the cells for a particular experiment were given appears in column 2. In general eight cells were filled at one time and six of these cells irradiated for a total dosage of 3.4 x 1020 eV/l. The two unirradiated cells were used as thermal controls and analyzed for oxygen at the end of the run.Column 4 gives the averages for the irradiated samples corrected by sub-traction of the thermal reaction. The standard deviation giving a measure of the vari-ability of the reaction within an experiment appears in column 5. The experiments were designed to test the effects of several types of cleaning procedures and other variables leading to non-uniform results. Expt. 1 and 2 of table 1 show the effect of triple and regular distilled waters. The latter produces 4.9 times as much decomposition during irradiation as does the triple distilled water. Packing cells with glass helices so as to increase the surface 60-fold possibly reduces the net over-all reaction but has no marked effect. The effects of con-taminating the solution with potassium dichromate are given in expt.4 6 and 8. Al-though the 0.1 pM dichromate does not affect the results the effect of 10 pM dichromate is impressive and hereafter the use of cleaning solution was avoided in this hydrogen peroxide work (compare expt. 6 and 8). Expt. 9-14 inclusive show the results of pains-taking attempts to obtain reproducible rates of decomposition within a run and between runs. This proved to be futile and from the work of irradiations 11 to 14 it was con-cluded that impurities from the cell wall play a predominant role. The more detailed data of irradiations 11 to 14 on which this conclusion is based are given in table 2. Irradiations 11 and 13 were performed with one set of cells and irradiations 12 and 14 with another set. In all four irradiations there is wide variability in the results for each irradiation that is much greater than the 2 to 4 % error in measuring the oxygen content of the cells (see table 2).The column headed R shows that the oxygen yields in a set of cells change by the same factor for each cell when the set is given different treatments in two successive experiments and the constancy of this factor holds even when the variability between cells within the same experiment is large. The constancy of R suggests that variation in oxygen yields is to be associated with the individual cells prob-ably with contaminations originating in the cell walls. In support of this hypothesis, microscopic examination of the inner glass surfaces of the cells disclosed many irregular-ities and discoloured spots.Treatment with hydrofluoric acid failed to remove all con-tamination or to provide uniform results (see expt. 15 and 16 table 1). Further support is furnished by expt. 20 and 21 where the use of paraffin coated cells eliminated variation between cells within an experiment and reduced the variation between experiments. Paraffin coated cells were used in the rotating sector experiments. Although the impurities are postulated to originate in the cell walls the impurities that affect the reaction must presumably become dispersed throughout the solution very early in the reaction. This conclusion is based on a calculation using the usual random walk displacement theory.7 The lifetime of a kinetic chain was measured as 0.60 sec in the sector experiments to be described the jump frequency of a diffusing radical was taken as 2 X lOlo/sec,g and the jump distance from one equilibrium position to the next EFFECTS OF CELL PREPARATION AND WATER PuRITY.-The rotating sector experi I72 DECOMPOSITION OF HYDROGEN PEROXIDE TABLE NAL EFFECT OF TREATMENT OF CELLS AND PURITY OF WATER ON THE DECOMPOSITION OF 0*10M HYDROGEN PEROXIDE BY GAMMA RAYS irrad.no. I 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 treatment of cells * WCSF WCSF SF WF WSF WCSF WFCSDF F WFCSDF WSDF WNSDF WNDF WNSDF WSF WHFF WSF NSDF 3XF 3XF HFF HFF 3XF WF W3XF paraffin paraffin * Meaning of codes : no. of cells 6 6 6 6 4 2 4 4 6 6 6 6 6 6 5 3 6 6 6 4 6 av.pmoles O?/l. 4 h 266 1293 174 1433 237 41808 279 28 1 196 167 143 178 81 156 257 815 307 297 182 464 435 G l -25 141 24 106 61 -37 31 47 29 16 32 10 24 70 20 76 103 12 21 27 0 /a v. 0.0940 0,1090 0.1375 0.0740 0.2574 I 0.1326 0-1 103 0.2398 0.1737 0.1 119 0.1798 0,1235 0.1538 0.2724 0.0245 0.2476 0.3468 0.0659 0.045 3 0.0621 remarks 3X distilled water regular distilled water cell packed with helices cells contaminated with cleaning solution 3X distilled water 10 micromolar 3X distilled water 0.1 micromolar 3X distilled water 3X distilled water effect of HF treatment 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water 3X distilled water K2Cr207 K2Cr207 C = heated in cleaning solution ; S = steamed 1 h; F = heated in furnace to 600" C ; D = dried in oven at 110" C ; N = heated in hot concentrated nitric acid for 30 min ; W = rinsed out several times with regular distilled water ; 3X = rinsed out several times with 3X distilled water ; HF = treated with HF for 20 min ; paraffin = inner walls of furnace-treated cells coated with paraffin.-f c = standard deviation. $ Both the blank and the irradiated cell contained this same total amount of oxygen. was taken as 3.1 x 10-8 cm from the diameter of a water molecule so that the calculated total distance diffused by the radicals involved in one kinetic chain is 0.03 mm.There-fore about 0.3 % of the chains would reach the walls of the 1.8 cm diameter cells used in the sector experiments. The rates for paraffin cells in the sector experiments are in good agreement with the paraffin cell rates of table 1 when corrected to the 3.7-fold higher intensity used for table 1 experiments. This shows that in the table 1 experiments even fewer chains reached the walls since the lifetimes are somewhat shorter due to the higher rates. APPARATUS.-The rotating sector method of Briers Chapman and Walters 9 has in recent years been applied extensively to oil phase polymerization systems 103 11 and to liquid phase hydrocarbon oxidations.12.13 In the work cited ultra-violet light was used to initiate the free radical reactions studied.In the present work on the decomposition of dilute aqueous hydrogen peroxide the y-radiation from C060 (1-1-1 -3 meV) was used to initiate chains and the penetrating nature of this radiation introduced experi-mental complications not encountered with ultra-violet light. The nature of the corn-plications will be apparent in the discussion of the apparatus 173 EDWIN J. HART AND MAX S . MATHESON TABLE 2.-A COMPARISON OF THE DECOMPOSITION OF 0.10M HYDROGEN PEROXIDE BY GAMMA RAYS IN UNCOATED CELLS cell no. itrad. no. &Fotg irrad. no. &$'Ph R(13/11) 26 11 119 13 66 0.55 82 11 163 13 93 0.57 85 11 131 13 72 0.55 87 1 1 147 13 85 0.58 89 11 151 13 82 0.54 90 11 146 13 86 0.59 R( 14/ 12) 30 12 197 14 175 0.89 76 12 201 14 179 0.89 78 12 165 14 456 0.88 79 12 190 14 170 0.90 83 12 165 14 150 0.9 1 86 12 128 14 115 0.90 Fig.1 shows a vertical cross-section through what may be designated as the optical The cylinder S is a C060 source similar to that described earlier 6 axis of the apparatus. S T FIG. 1 .-Vertical cross-section of intermittent y-ray device. but of 80curie strength. A is an aperture in a portion of the lead shield surrounding the source. At this point the lead shield is 25-4 cm from the source and is 16.2 cm thick. Because of the mass of lead required to stop the y-rays rotation of a disc containing the aperture is not a very practical method of obtaining intermittent radiation particularly for high speeds of rotation.Therefore the cells C were mounted on the aluminium wheel W (radius = 12.0 cm to centres of cells which were 5-1 cm long) close behind the lead shield and aperture and the rotation of this wheel caused the cells to be irradiated intermittently as they rotated past the opening in the lead shield. Holders for twelve cells were spaced equally around the wheel. The aluminium wheel could be rotated at constant speeds ranging from 600 rev/min to 0-0938 rev/min by means of a synchronous motor and a gear box. The aperture was designed to provide a cycle in which the cells were irradiated one-fourth of the time in each cycle when the wheel W was rotating at constant speed. To accomplish this the side surfaces of the aperture were generated according to the following scheme.The aluminium wheel is stopped so that a vertical line through the axis of the wheel bisects the angle between two adjacent cells. The angle between cells being 30" this means that two other cells are located so that their centre lines make angles of 45" with the vertical. The centre of the upper end of one of these 45" cells is connected with the top centre of the source by a straight line and similarly the centre of the lower end of the cell and the bottom centr 174 DECOMPOSITION OF HYDROGEN PEROXIDE of the source are likewise connected by a straight line. Then a side surface of the aperture is cut through the shield by a line moving from the cell toward the source the line during this motion remaining parallel to the plane of the aluminium wheel and intersecting both guide lines drawn from cell to source.The sides of such an aperture will form a slightly curved surface so that a point in the cell will not begin to receive radiation simultaneously from the full length of the cobalt rod. (With a point source such curvature would be eliminated.) Due to this factor and to the scattering of radiation in the aperture it follows that when the wheel is rotating the radiation intensity at a given point in the cell does not change instantaneously from zero to full intensity as is assumed in the theory.14 The correction for this is considered later. Further with this experimental arrangement the sector scans the cell so that one part of the cell may be receiving radiation when another part is not.This however will not be important unless radicals diffuse rapidly throughout the cell. With regard to certain other assumptions of the theory the use of the C060 source is advantageous. For example the emitted radiation is constant in time and also across the aperture except at the edges. Further because the radiation is penetrating and because the mean distance from the source to the cells is large (45.4 cm) the ab-sorption of energy is nearly constant from front to back in the cell. SECTOR EXPERIMENTAL PROCEDURE.-FrOm each experiment the sector rate ratio (ratio of rate in intermittent radiation to rate in steady radiation) for one sector speed was ob-tained. Eight cells filled simultaneously were used in each experiment as follows. Two cells were exposed while mounted on the aluminium wheel which was rotated at a known speed.The starting exposure time was controlled exactly by lowering the source from a shielding lead turret to the position shown in fig. 1. Raising the source similarly ended the exposure. After the cells exposed to the intermittent irradiation were removed two other cells were located on the now stationary aluminium wheel and centred sym-metrically behind the aperture in the lead shield. These were exposed to steady irradi-ation for a measured time. Each set of cells was analyzed immediately after it had been irradiated. The other four cells stored in the dark were used as blanks to determine the thermal rate of decomposition. Generally two blank cells were analyzed at the same time as the cells from interrupted irradiation and the other two at the same time as the cells receiving steady irradiation.Occasionally one of the blank cells was analyzed at the start of the sectored irradiation. The oxygen yields of the blanks were plotted against time elapsed after filling of the cells. The plots indicated a small but relatively rapid thermal reaction (in addition to the analytical blank of 4 to 5 pmoles/l.) in the first hour or less followed by a much slower very slowly decreasing rate of evolution of oxygen. The thermal reaction being considered to be ionic the oxygen yields in irradiated cells were corrected by subtracting the amounts of thermal reaction corresponding to the times at which the irradiated cells were analyzed. (Allmand and Style also corrected for the thermal rate by simple subtraction.) These amounts were read from the above plots of thermal oxygen yield against time after cell preparation.From the corrected oxygen yields and irradiation times the rates of oxygen production for intermittent and steady irradiation were calculated. Calculation of a typical experi-ment follows : 10 rev/min sector-(492 485) av. 488.5 pmoles 02/1. after 15.67 h irradiation. Sector blank-84.0 pmoles 02/1. in 22.25 h after preparation. Sector yield corrected for blank-404.5 pmoles 02/1. in 15.67 h irradiation. Sector rate-25.85 pmoles 02/1. h. Steady irradiation-(448 464) av. 456 pmoles 02/1. after 6.0 h irradiation. Steady irradiation blank-85.0 pmoles 02/1. in 28-85 h after preparation. Steady irradiation yield corrected-371-0 pmoles 02/1.in 6 ~ 0 h irradiation. Steady irradiation rate-61.83 pmoles 02/1. h. Sector rate ratio-4-418. An experiment to test the assumption that the thermal blank is not a free radical reaction was carried out using hydroquinone. With 3.55-fold the intensity used in the sector experiments 0-1 mM hydroquinone in 0.1 M H202 reduced the rate of oxygen production to one-fifth that obtained with no inhibitor. The retarded rate was linear as far as followed and was estimated to correspond to a chain length of 3. On the other hand for the same concentrations the hydroquinone had little if any effect on the initial rapid thermal blank reaction occurring in the first 4 h but reduced the thermal rates about 28 % jn the period extending from 20 to 200 h. If the small thermal blank rate were all free radical the estimated chain length would be about 1000 so that the hydro EDWIN J .HART AND MAX S . MATHESON 175 quinone by cutting this chain length to 3 would reduce the blank to a negligible amount. Thus it is concluded that not more than 28 % of the thermal blank is a free radical re-action. For the sample calculation above this would correspond to a dark reaction less than 1 % of the steady irradiation rate and from fig. 1 of ref. (14) it can be seen that such a dark rate has Iittle effect on the lifetime measurement. TREATMENT OF SECTOR DATA.-Frorn eqn. (14) derived in a later section from our assumed mechanism T ~ the kinetic chain life is inversely proportional to Roy the rate of oxygen production in steady irradiation. In essence this inverse dependence of R, and T ~ assumes that at constant cell temperatures the variations of R are due to variations in the rate of initiation.Since the temperature difference between experiments varied 5°C as a maximum with most experiments within a 2" C range and since the temperature coefficient for such reactions is small the observed large variations in the steady irradiation rate of oxygen production obtained in the sector experiments cannot be accounted for through temper-ature effects alone. Further no correlation between R and temperature was observed. Experiments already described have shown that even with paraffined cells simultaneously irradiated with the quite constant C060 source the standard deviation is about 5 %. Thus it is possible that the steady rate variations between sector experiments may be due at least partially to impurity effects rather than to changes in the rate of initiation.Never-theless it seemed worth while to use the inverse relation of eqn. (14) to correct sector data to a common R, since this procedure gave a moderate decrease in the standard deviation of the results and had little effect on final rate constant calculations. Consequently, all experiments were adjusted to a common R (60.0pmoles 02/1. h) by changing the 11's (11 = duration of intermittent radiation pulse) and R,'s inversely with respect to each other the corresponding sector rate ratios being left unchanged.14 The experimental sector rate ratios especially those for low speeds of rotation of the sector were somewhat high indicating either that an appreciable " dark " reaction was occurring in the " dark " part of the cycle or that the effective angle of the aperture was greater than the quarter circle intended.To test this point two experiments were carried out. First with the aluminium wheel stationary and with cells 1 and 2 centred behind the aperture ten cells of the ferrous sulphate-formic acid dosimeter 500 pN FeS04 + 0.1 M HCOOH + 0.8 N HzSO4 + 1-2 mM 0 2 were irradiated. Under the conditions employed the reaction in this dosimeter is nearly linear with dose.15 The following results show that in most of the shielded portion of the cycle negligible dark reaction would be induced in the hydrogen peroxide reaction by scattered radiation. pequ i v. cell position 1 2 3 4 5 6 7 8 9 10 11 12 Fe3+/l.h 117 115 61 1.1 0.15 0.0 0.15 - 0.0 - 0.6 54.6 The cells on the edge of the aperture (3 and 12) received half as much dose as cells 1 and 2 as was expected. In the second experiment two cells of the ferrous sulphate dosimeter using 500 pN FeS04 -t 0.8 N HzS04 (air saturated) gave 5.72 pequiv. Fe+++/l. h when exposed to steady irradiation while centred behind the aperture. When the experiment was repeated with the wheel rotating the rate was smaller by the factor 0.261. From these experiments it was concluded that the effective angle of the aperture was 0.261 of 27r. Therefore the ratios measured for an irradiation period 0.261 of a cycle were corrected to the values they would have for the same R and sector speed but for an irradiation period 0.250 of a cycle.This essentially amounted to subtracting 0.022 from the experi-mental values of two x sector rate ratio since the theoretical curves of two x sector rate ratio against b (b = t l / ~ ) differ by approximately this amount for the above two sector apertures for all values of b. A further correction was required because as already noted the transition from zero to fuil intensity at the edge of the aperture was not instantaneous. An X-ray film was exposed just in front of the wheel. A Knorr-Albers microphotometer was used to measure the density of the developed film. The plot of film density against the angular position on the wheel showed that the radiation pulse could be fairly well approximated by a trapezoid. This type of approximation has been suggested elsewhere in sector work involving the use of visible or ultra-violet light.16 Burns and Dainton showed that for a trapezoidal rather than rectangular light pulse the correction to the theory is appreci-able only for slow sector speeds.For our case where the radiation period duratio 176 DECOMPOSITION OF HYDROGEN PEROXIDE (“sector width”) was close to five times the period of increasing (or of decreasing) intensity (“ beam width ”) the limiting fast sector rate ratio was naturally unchanged, while the limiting slow sector rate ratio was calculated to be raised to 0.27 for the trape-zoidal radiation pulse as against 0-25 for the corresponding rectangular pulse lasting one-fourth of a cycle. This increase at slow sector speeds is due to the dependency of the rate on the square root of the dosage rate.An approximate theoretical curve of two x sector rate ratio against b was drawn for the trapezoidal case of our experimental con-ditions. Then an experimental sector rate ratio (previously corrected to an effective sector aperture 0-250 of a cycle) was fitted to this curve the value of b noted and then the corrected rate ratio taken at the same value of b from the theoretical curve for a one-quarter cycle rectangular radiation pulse. All sector experiments were carried out in the temperature range from 25 to 30” C . RESULTS AND DlSCUSSION The y-ray initiated decomposition of hydrogen peroxide is a chain reaction in which water oxygen and a small amount of hydrogen are the products. The chain length of this reaction is altered by variables such as purity of the water and contamination by cell surfaces and in addition is dependent on dosage rate and concentration of hydrogen peroxide.In the present work the C060 y-ray dosage rate has been varied in the range (0.0038 to 0.63) x 1020 eV absorbed/]. min and the concentration from 0.008 to 200 mM hydrogen peroxide. In calculating energy absorbed Hochanadel’s 17 value for the FeS04 air-saturated dosimeter of 15.5 Fe2+ oxidized/100 eV was used. Except for minor changes the present work is in agreement with that of Fricke who reported that the oxygen yield is inversely proportional to the square root of the dosage rate and varies directly as the square root of concentration. A major difference exists however in the quantitative value of the yield under certain specified conditions.EFFECT OF DOSAGE RATE.-Owing to the existence of a non-chain primary decomposition of water yielding hydrogen peroxide and to the occurrence of a reaction of oxygen (the decomposition product of hydrogen peroxide) with the hydrogen atom the kinetics of this reaction become complex at low hydrogen peroxide concentrations. Dosage curves (oxygen production against time) are non-linear in the concentration range up to above 10 mM. Above this con-centration linear dosage curves result. Fig. 2 gives the results obtained on air-free unbuffered 32-8 mM hydrogen peroxide in the dosage rate range (0-0177 to 0.63) x 1020 eV/1. min. The yield plotted against the reciprocal of the square root of dosage rate gives a straight line which it is important to note does not extra-polate to the zero decomposition indicated by Fricke.Instead an extrapolation to 2.6 molecules 02/100 eV is obtained for very high dosage rates. Since 2.73 pairs of H and OH free radicals appear per 100 eV absorbed it is apparent that a chain reaction no longer exists at high dosage rates. Fig. 3 shows the results obtained on 100 mM hydrogen peroxide solutions (air-free and unbuffered) irradiated in paraffin coated cells. The dosage rate was varied in the range (0.0038 to 0.63) x 1020 eV/1. min. The slope in fig. 3 is greater than in fig. 2 since at this concentration the chain lengths of the decom-position reaction are much longer than for the experiments of fig. 2. Because of the steep slope the uncertainty in the intercept is as large as the expected value so that the line has simply been extrapolated to the origin.This curve is similar to fig. l b of Fricke. Johnson reports that the decomposition of hydrogen per-oxide at a concentration of 0.39 mM is independent of dosage rate in the range from (0.12 to 1.56) x 1020 eV/l. min. This is not in marked disagreement with the above results since Johnson with a yield of 4-3 molecules hydrogen peroxide decomposed/100 eV is clearly not in the concentration range where a chain reaction is obtained. The expected limit at high intensities is 4-8 molecules hydrogen peroxide/100 eV as can be calculated from the constant terms in eqn. (7) and (9) and using our extrapolated value of 2.6 molecules oxygen/100 eV EDWJN J .HART AND M A X S . MATHESON 177 Likewise Hochanadel working at concentrations of 0.1 and 4.0 mM hydrogen peroxide obtained low yields indicating chains of the order of unity. A com-FIG. 2.-Effect of C060 y-ray dosage rate on the decomposition of unbuf-fered 32-8 mM hydrogen peroxide. abscissa (1018 eV/l. min)-’ ordinate molecules oxygen/ 100 eV. parison of the yields of several investigators with certain of ours is given in table 3. It is to be noted that at low concentrations there seems to be general agreement FIG. 3.-Effect of C060 y-ray dosage rate on the decomposition of un-buffered 0.1 M hydrogen peroxide in paraffin coated cells. abscissa (1018 evil. min)-a. ordinate molecules oxygen/ 100 eV. in the yields. However important differences appear at concentrations where the chain attains several units in length.At 10 mM hydrogen peroxide after adjustment is made for the effect of intensity and concentration it is found that at a dosage rate of 0.63 x 1020 eV/1. min 426 29.3 and 7-3 molecules hydrogen peroxide/100 eV are decomposed in Fricke’s Johnson’s and in our experiments 178 DECOMPOSITION OF HYDROGEN PEROXIDE respectively. Since measurements of chemical yield and dosage are far more accurate than this it is concluded that trace impurities are the main contributing factors to this divergence. At present we are not in a position to state whether the high yields or the low yields represent the impurity free reaction. TABLE 3.-YIELDS IN THE DECOMPOSITION OF AQUEOUS HYDROGEN PEROXIDE conc.(H202)0 mM 0.10 4.0 0.027 0*100 0.25 10.0 10.0 0.40 0.39 0.286 0.326 1.14 2-13 12.4 10.5 G(H202) molecules/ 100 eV 3.7 3.7 1.3 2.1 2.5 53.5 156.0 5.4 3.4 36-6 4.0 3.3 4.1 7.8 9.0 BY IONIZING RADIATIONS author dosage rate (1020eVil. min) 0.0054-2.47 Hochanadel 17 0.0054-2.47 Hochanadel17 0.34 Toulis 18 0.34 Toulis 18 0.34 Toulis 18 0.38 Fricke 3 0.38 Fricke 3 0.89 Johnson 4 0.117-1.68 Johnson 4 0.89 Johnson 4 0.0 I76 this work 0.63 this work 0.0176 this work 0.63 this work 0.63 this work radiation Cob0 y-rays Co60 y-rays X-rays X-rays X-rays X-rays X-rays X-rays X-rays X-rays Co60 y-rays Co60 y-rays C060 y-rays Co60 y-rays Co60 y - r a y s EFFECT OF CONCENTRATION.-The effect of hydrogen peroxide concentration was studied in air-free unbuffered solutions at concentrations in the range from 0.008 to 200 mM.At concentrations below 0.02 mM the dependence changes but above this concentration the rate of decomposition depends on the square root of the concentration. The points on the curve of fig. 4 were obtained at a /OO 200 300 400 dosage rate of 0.63 x 1020 eV/l. min. However enough points were obtained at a dosage rate of 0.0038 x 1020 eV/1. min to show that this Sam dependence holds here too. It is believed that this dependence will hold at any dosage rate where an average chain of even a few units is obtained. This follows from the fact that the average chain is only 3.7 units in length at the highest point in fig.4 EDWIN J. HART A N D MAX S . MATHESON 179 This dependence of yield on the square root of hydrogen peroxide concentration was also found in the previously cited work of Fricke and Johnson. There is definite evidence in the present work that the form of the concentration curve changes below 0.02 mM hydrogen peroxide. However all of this work on hydrogen peroxide concentration was carried out using standard glass irradi-ation cells. Therefore the effects of impurities from the glass wall would have a particularly pronounced effect at these low concentrations. Further work should be carried out in paraffin coated cells. MECHANISM.-Previous investigators have provided suitable mechanisms for initiation and chain propagation processes. A bimolecular termination step correctly giving the observed relation between yield and the inverse square root of intensity has also been proposed.However an over-all mechanism must also explain the dependence of yield on square root of concentration as found in the photochemical work of Allmand and Style in the X-ray work of Fricke and Johnson and in our y-ray work. An explanation must also be provided for the extrapolated value of 2.6 oxygen molecules produced/100 eV at high dosage rates as is shown in fig. 2. These facts can be explained by the following mechanism : rate H20 = H + OH (la) Ha for decomposition H2O = 1/2 Hz + lj2 H202 (Ib) K’L] of H2O H 4- H202 H2O -+- OH (2) Kla OH -1- H702 = HO2 + H20 (3) kp’(OH)(H>Olj HOz I- P I 2 0 2 H20 + OH -4- 0 2 (4) kp(H02)(H202) 2HO2 -t Hz02 -= 2H202 -/ 0 2 (5) 2kt(M02)?( HzO?)-for disappear-ance of HO2.This mechanism postulates that reaction (4) is much slower than (3) since termination is by hydroperoxy radicals and not by hydroxyl radicals. Evidence for this has been discussed by Agar and Dainton.19 These authors have also pointed out that inhibition of the photolysis of H202 by alkali indicates that H02 and not 02- is involved in the propagation reaction. (Using the KHoz and KH20z assumed by these authors the concentration of H02 would be about three times the 02- concentration in our sector experiments.) Lea2 also found the quantum yield independent of pH in the pH range 1-6 suggesting that 02- or H02- is not involved. Further evidence that HO;! + H202 is relatively slow is found in the work of Barb Baxendale George and Hargrave,Zo who find that the H02 i- H202 reaction does not take place in the presence of ferrous or ferric ions.The results of Bray 21 and George 22 on KO2 dissolved in water can be accounted for by the above mechanism if it is assumed that the solution and re-action of KO2 produces HOz radicals. The ratio of reaction (5) to reaction (4) is Zkt(HOz)/k independent of H202 concentration and is equal to 1.0 x lOs(H02) from our rate constant measurements. Thus if the HO2 concentration is 10-9 molar or higher near the surface of the dissolving KO2 particles reaction (4) will occur to an extent which is 1 % or less than that of reaction (5) and in reaction (5) the net effect is the production of one H202 and one 0 2 molecule from two H02 radicals.This is also the observed stoichiometry for KO2 dissolving in water. If the assumption that KO2 reacts with water to yield H02 is correct con-centrations of H02 in excess of 10-6 M might well be expected near the surface of the K02. Ry assuming that a steady state of intermediate free radicals is readily estab-lished in the above mechanism eqn. (6) is derived, (6) Ro = dOz/dt = kp(KI,(Ha02)/kt)3 t KI, 180 DECOMPOSITION OF HYDROGEN PEROXIDE After dividing both sides by KIu in order to express the results in terms of yield (7) (8) According to Hochanadel the ratio K’IK is 0.34 for the hydrogen peroxide (9) Primary steps (la) and (1b) have been discussed recently in the papers of Allen and co-workers,23 Hochanadel,l7 and Hart.24 The present work provides addi-tional evidence for the existence of step (16).Hydrogen gas is always found in the radiolysis of hydrogen peroxide solutions. The yield is substantially inde-pendent of dosage rate and of hydrogen peroxide concentration in agreement with the above mechanism. The Haber-Willstatter 25 propagation steps (3) and (4) account for the chain decomposition and have been widely applied in thermal and photochemical mechanisms for this reaction. Our termination reaction (5) appears as a termolecular one and differs from previous ones by including hydrogen peroxide in this step. The inclusion of hydrogen peroxide in the termination reaction introduces the correct dependence between rate and square root of hydrogen peroxide concentration. Although present work offers no proof we suggest the formation through hydrogen bonding of an intermediate complex between hydrogen peroxide and the hydroperoxy radical as follows : per unit energy required to produce one radical pair in (la) one obtains The expression for hydrogen peroxide disappearance becomes Y(02 ) = (d02/dt)(l /K&) = kp((H202)/ktKLJ) + 1.Y(H202) = - K’/2K + 2 + 2kp((H202)/ktKTa)*. - Y(,,) = 1.83 + 2kp((H202)/krKIn)h. -system. Thus the above equation reduces to This complex intermediate may decompose as in reaction (4) or if formed in the presence of another hydroperoxy radicaI could yield reaction (5) through a larger intermediate. In either reaction a simple regrouping of electrons and moderate adjustment of interatomic distances is all that is necessary for reaction as in (4’) and (5’).O H (4’) /0-H 0 4 \O-H 0 0-H -+ H-O/ H-0 O/o-H I 0 //O H\o (5 ’> ;/O+ H-O/ H-0 /o-H O/o-H H-0 If reaction (3) were the slow step in propagation so that the termination reaction corresponded to eqn. (11) then eqn. (12) would be obtained. OH + OH + H202 = 2H202 (2kr’(OH)2CH202)) (J 1) Y(02) = kp’((H202)/wm*. (12) The intercept in fig. 2 predicts that hydrogen peroxide would be decomposed at a rate yielding one oxygen molecule per radical pair at infinitely high dosage rates since the data of Hart 24 show that 2-73 water dissociations occur in reaction (la) per 100 eV of C060 y-ray energy absorbed in water. This is in accord with eqn. (7) but not with (12). Therefore our data definitely favour termination step (5) and not (11).This experimental evidence provides an additional argu-ment for claiming that the rate constants measured in our rotating sector experi-ments are kp and kt and not k,’ and kt’. Reaction (11) is not written as 20H + H202 == 2H20 + 0 2 EDWIN J . HART AND MAX S . MATHESON 181 since 0 2 is not produced in the “ hot spot ” or “ end of the track ” reaction but H202 is presumably from OH radicals. RATE CONSTANTS.-The experimental results and corrections of the sector experiments are listed in table 4. Twice the sector rate ratios corrected to a rectangular radiation pulse lasting one-fourth of a cycle are plotted in fig. 5 against the Iogarithms of the radiation pulse times. The theoretical curve shown in the same figure is plotted against log b.The ratio tl/b corresponds to a life-time of 0.60 sec. This value was calculated by averaging the logarithms of the lifetimes of the last column of table 4 weighting the less reliable starred results one-half as much as the others. The lifetime obtained from the uncorrected ratios (column 6 table 4) is 0.83 sec. expt. 22 23 24* 25 26 27* 28 29 31 32 33 (H202) mM 85.6 92.0 86.4 92.2 95.0 98.4 93.6 100.4 94.2 100 100 TABLE 4.-LIFETIMES OF KINETIC CHAINS RO .umoles/l. h expt. 50-3 62.0 59.0 84.1 61.83 55.5 40.56 59-6 65.0 66.3 53-35 0.025 7.5 7.5 0.625 1.50 7.5 0.025 0.025 3.00 1.50 159.9 2 x sector rate ratio corr. to Ro = 60.0 pmoles/l. h 0.02095 7.75 7-37 0.876 1.547 6.94 0.0169 0.0248 3.25 1-659 142.2 expt.1.038 0.620 0.780 0-872 0.836 0.754 1 so00 1.044 0.696 0-846 0.624 corr. to aperture 0.250 1.016 0.598 0.758 0.850 0.814 0.732 0.978 1.022 0.674 0.824 0.602 corr. to rectangular pulse 1-016 0.557 0.742 0.833 0.796 0-713 0.978 1-022 0.642 0.807 0.562 lifetime zs sec. -1-83 0-406 0.559 1-39 --0.339 0.646 *Because of experimental difficulties the results in these experiments are not con-sidered as reliable as for the other experiments. FIG. 5.-Rate of oxygen production in intermittent radiation for different radiation pulse durations tl ; curve theory for rectangular radiation pulse lasting one-fourth of a cycle and for zero dark reaction.abscissa b = (duration of radiation pulse)/(kinetic chain life) = f l / T S . ordinate 2 x sector rate ratio. 0 our points on duplicate cells. Lea’s points (ref. (2)) corrected to our rates and for non-chain portion of reaction. @ our points on single cells 182 DECOMPOSITION OF HYDROGEN PEROXIDE CALCULATION OF RATE CONSTANTS.-The fOl1OWing eqn. (13) and (14) show that two sets of experimental data suffice to obtain the individual rate constants : R$/Ri(H202) == kp2/2kt, where Ri -- 2KI = rate of initiation of chains and R o ~ s = kp/2kt. Eqn. (13) is obtained by rearranging eqn. (6) and assuming long chains and eqn. (14) may be derived from eqn. (13) using the following relations : (1 5 ) conc. of free radical chains N_ (HOz) if H02 + H202 slow propagation step (16) Ro = kp(H02)(H202).(1 7) conc. of free radical chains = (RiT,), Considering first the data needed for eqn. (13) the dose rate for steady ir-radiation of cells on the aluminium wheel centred behind the aperture corresponds to 5.72 pequiv. Fe3+/l. h produced in a 500,uN FeS04 + 0.8 N H2S04 (air saturated) dosimeter. This dose rate decomposes 1 *006 pmoles H20/1. h to radical pairs to give Rj = 5.59 x 10-10 mole radicals/l. sec. This radical pair yield for a given dose rate is estimated from the data of Hart24 on the inhibition by 0 2 of the HCOOH + H202 reaction. The rate of production of radical pairs is taken as equal to the rate of 0 2 consumption in this inhibition. In this work 0.153 mmole water dissociations/l.h to yield hydrogen atoms and hydroxyl radicals were obtained under irradiation conditions where 0.870 mequiv. Fez+/]. h were oxidized. Corresponding to the value of Ri above average values (@)+ = 1.665 x 10-8 mole/]. sec and (H202) -~ 93-26 x 10-3 mole/l. were obtained from the data of table 4. These data and eqn. (13) give kp2/2kt = 5.32 x 10-6 sec-1. Further eqn. (14) and the average Ro value found above for the sector experiments with the kinetic chain lifetime of 0.60 sec found in the same experi-ments gives kp/2kt = 530 1. mole-1 sec-1 and 2kt = 5-3 x 1010 1.2 mole-2 sec-1. Tn the sector experiments in steady irradiation (Ro - 1.66 x 10-8 mole/l. sec) the concentration of H02 is 3-4 : 10 10 mole/l. and the chain length is 30 mole-cules of 0 2 per initiating radical.The only results in the literature which may be compared with ours are those of Lea.2 The comparison can be made by assuming steps (3) (4) and (5) of our mechanism and by assuming that in photochemical initiation Lea's reaction (l), occurs with an effective primary quantum yield f. Then, d(H202)/dt = 2kp(H202)'(2f~a~,/2kt)' + 2 fIabs. From Lea's results at high intensities 2f= 1.39 = the limiting quantum yield, and from his results at low intensities usingf= 0-7 kp2/2kt = 11.1 x 10-6 sec-1. I f f = 1-0 then kp"2kt is found to be 7-8 x 10-6 sec-1 which does not differ appreciably from our value of 5.3 x 10-6 sec-1. In fig. 5 are plotted the results of Lea's sector work corrected to the same steady rate of reaction as used by us. These ratios have also been corrected for the non-chain portion of the H202 decomposition by subtracting 2JTabs (f= 0.7) from the rate in steady illumination and (fIabs)/2 from the rates in sectored light.The authors wish to acknowledge their indebtedness to Messrs R. A. Blomgren and L. S. Markheim for the design and installation of the rotating sector and to Miss P. D. Walsh for technical assistance EDWIN J . HART AND MAX S . MATHESON 183 APPENDIX Received 10th June 1952 NEW EXPERIMENTAL msuLTs.-Hydrogen peroxide or its complex with the hydroperoxy radical has been postulated to participate in the termination step of the 7-ray induced decomposition of hydrogen peroxide. This conclusion has been drawn from the dependence of the rate of decomposition on the square root of hydrogen peroxide concentration.A first order dependence would support the customary bimolecular termination previously employed.25 The X-ray work of Fricke 3 and Johnson 4 and part of the photochemical work of Allmand and Style 9 and Kornfeld 26 are consistent with the square root dependence. While the X-ray and y-ray work clearly shows the square root dependence a first order reaction is frequently reported in photochemical work.2927 Impurities are sus-pected to be the cause of this discrepancy since they are known to cause important changes in the kinetics of hydrogen peroxide formation and decomposition. Since our sector work was carried out under conditions where less than 0.5 % decomposition occurred it was felt desirable to re-investigate the order of the reaction under condilions of more exhaustive purification and greater degree of reaction.EXPERIMENTAL PURIFICATION OF HYDROGEN PERomE.-Our previous purification consisted of a steam distillation from acid solution of Merck reagent grade 30 % hydrogen peroxide. Since this product was identical in radiolysis behaviour to inhibitor free Buffalo Electro-Chemical 90 % hydrogen peroxide and gave the square root dependence observed by Fricke and Johnson it was thought that our hydrogen peroxide was of sufficient purity for the sector FIG. 1.-Apparatus for the distillation of 90 % hydrogen peroxide. experiments. Tn order to provide further information on this point a purification pro-cedure recommended by Shanley 28 of the Buffalo Electro-Chemical Co. was carried out on B.E.C.90 % hydrogen peroxide. Sodium stannate was added to the still to neutralize nitric acid believed to be the only volatile material present besides water and hydrogen peroxide. A one-plate distillation was carried out on 175 g of hydrogen peroxide in a still maintained at 35" C (see fig. 1). The system was evacuated through F by a Hypervac pump with dry ice on traps C and D. E is a medium fritted glass filter. After one-thir 184 DECOMPOSITION OF HYDROGEN PEROXIDE of the hydrogen peroxide had distilled into C dry ice was placed on trap B also and the middle fraction collected and used in the radiolysis experiments. Trap A was added to catch any hydrogen peroxide carried over by entrainment of droplets. However no bumping of the liquid was observed and the distillation proceeded very smoothly over a period of about 6 h.The hydrogen peroxide (42.5 g) collected in B was diluted to 100 ml with triply distilled water in a heat treated volumetric flask and this stock solution of 13.5 M hydrogen peroxide was used in the radiation experiments. On the basis that an inhibitor might still be present in the vacuum distilled hydrogen peroxide 10 ml of the 13.5 M hydrogen peroxide was irradiated until 6-23 % decomposi-tion occurred. This pre-irradiated hydrogen peroxide was then used in some of the irradiations run at 0.1 M. ANALYSES.-TWO methods of analyses were used to measure the decomposition of hydrogen peroxide by y-rays. In one the Van Slyke apparatus was employed to measure evolved oxygen where the total decomposition was a fraction of 1 %.The second method employed ceric sulphate using ferroin as an indicator. In this method 0.100 to 2-00 ml of hydrogen peroxide was added to 10 ml of 0.8 N sulphuric acid and titrated directly with 54.64 mN ceric sulphate which had been standardized with pure iron wire. Reproducibility by this method is better than 0.5 %. tion was measured cells of the type employed for the sector experiments were coated with paraffin in the manner previously described. In a separate experiment as a further treatment designed to neutralize the alkaline surface of the glass a set of cells was given a fuming sulphuric acid treatment suggested by Dr. Shanley. In this method 2 to 3 ml of fuming sulphuric acid were drawn into furnace heat-treated cells and the surface wetted by this material.After standing 10 min the cells were rinsed several times with triply distilled water in order to remove the sulphuric acid. In the irradiations carried out to nearly compIete decomposition (followed by H202 analysis) 80 ml paraffin coated cells were used. During the irradiation there was no indication that the wax surfaces had ruptured. The analyses were performed on 2 ml samples withdrawn from the irradiated 80 ml solution. Under the conditions of irradi-ation up to 70 % of the solution could be withdrawn without appreciably affecting the dosage rate delivered to the solution irradiated. PREPARATION OF CELLS AND IRRADIATION OF soLuT1oNs.-For Use when Oxygen eVOlU-RESULTS ORDER OF THE REACTION.-The order of the y-ray induced decomposition was investig-ated using redistilled 90 % hydrogen peroxide at concentrations of 0.0314 0.1002 0.306 and 1-008 M.These solutions were decomposed by irradiation in 80 ml paraffin coated cells until a series of overlapping concentration ranges was obtained. The forms of the individual curves demonstrate clearly that the experimental data are consistent with a dependence of rate of decomposition on (H202)0-5 instead of (H202)l.O (see fig. 2 and 3). In general deviation from the square root dependence is not observed until 80 % reaction has occurred whereas deviation from first power kinetics is observed at 20 "/, decomposition. A comparison of the one-half and first order rate constants is shown in table 1 for the initial rates of decomposition for the data of fig.2 and 3. A maximum of 19 % deviation from the mean occurs for the one-half order rate constants whereas the first order rate constants show a five-fold increase as the concentration of hydrogen peroxide is lowered from 1-008 M to 0,0314 M. In no case was any sign of an induction period indicative of inhibitors found. Eqn. (9) of our paper for the rate of decomposition of hydrogen peroxide indicates that for this chain reaction there is a small constant term (1.83 x rate of production of TABLE 1 .-INITIAL RATE CONSTANTS FOR H202 DECOMPOSlTION AT CONSTANT INTENSITY (CONSTANT RATE OF INITIATION) first order one half order moles/l. I?-' (moles/l.>?n-' 1.008 0.246 x 10-2 8-57 x 10-2 0-306 0.339 x 10-2 6-05 x 10-2 0.10025 0.781 x 10-2 8.82 x 10-2 0.03 142 1-273 x 10-2 6-52 x 10-2 init.conc. 7.43 x 10-2 av. kp2/2kt = 10.76 x 10-6 sec-1 from av. half order constant. kp2/2kt = 5.32 x 10-6 sec-1 from rates in sector experiments EDWIN J . HART AND MAX S . MATHESON 185 radical pairs) in addition to the (Hz02)+ term. If the complete rate expression is integrated the slopes of fig. 2 may then be corrected for an additional logarithmic term. This term is most important for the lowest concentration where the chains are the shortest. Here the correction decreased the slope by 10 :,/ without however affecting the linearity. Y I 0 40 80 I 2 0 160 200 240 280 TIME IN HOURS FIG. 2.-Dependence of the rate of hydrogen peroxide decomposition on (H202)* at a dosage rate of 3.98 1020 eV/1. hr. 0 1.008 M.Q 0.306 M. 0 0.10025 M. 0.03142 M H202. The redistilled hydrogen peroxide showed a dependence of rate of decomposition on (H202)0.5 in irradiations carried out to 0.4 % reaction in paraffined cells of the type used in the sector experiments. The Same dependence was obtained on distilled 90 % hydrogen 0.5 -3 0.4 - 0 %I < 0.3 -0 L I 01 D( 5 0.2 -1 I I \ I I 1 40 80 I 2 0 160 200 240 280 0.1 TIME IN HOURS FIG. 3.-Dedendence of the rate of hydrogen peroxide decomposition on (H202)l.O at a dosage rate of 3-98 x 1020 eV/I. hr. 0 1.008 M. 0 0.306 M. 0 0*10025 M. c) 0.03142 M H202. peroxide which had been irradiated to 6-23 % decomposition and then diluted. The data for 0.032 and 0.10 M appear in fig. 4. This result again demonstrates that trace impurities removable by irradiation are not present in our solutions 186 DECOMPOSITION OF HYDROGEN PEROXIDE Cells pre-treated with fuming sulphuric acid demonstrated the same erratic behaviour previously noted for freshly heat-treated cells.In general the amount of hydrogen per-oxide decomposed on irradiation is about one-half that found for the paraffin-treated cells. The results obtained were too erratic to distinguish between one-half and first order reactions. EFFECT OF WATER.-standard triply distilled water Barnstead still water and laboratory distilled water were used in the preparation and irradiation of 0.10 M hydrogen peroxide. The Barnstead water is prepared from tap water by a single distillation and is used as the first stage of our triply distilled water.The laboratory distilled water is prepared by a double distillation in Barnstead stills but in addition is delivered to the laboratory benches in aluminium tubing. Table 2 gives the rate constants obtained by irradiation at 3.98 x 1020 eV/1. h. The kinetics appear to conform to (H702)0.5 although a complete investigation was not made on the Barnstead watcr not- the laboratory distilled water. The Barnstead water shows a moderate increase in rate whereas the rate of the laboratory water is nearly 8 times that of our triply distilled water. FIG. 4. FIG. 4.-Effect of conceiitration on the decomposition of hydrogen peroxide at dosage rate of 7.6 x 1019 eV/1. hr. distilled and pre-irradiated H202. 0 distilled H202. TABLE 2.-EFFECT OF WATER ON THE ONL-HALE ORDER KATE CONSTANTS IK ?HE DECOMPOSITION OF HYDROGEN PEROXIDE 0.5 order rate constant W 2 0 2 ) o (m/M)+ h-' type of water triply distilled 0.10025 8.82 x 10-2 Barnstead 0-0976 11.6 x 10-2 laboratory distilled 0.0972 46-4 >< 10-2 A dark reaction that is small compared to the y-ray reaction also takes place in each of the hydrogen peroxide solutions.However with the laboratory distilled water 85 % of the hydrogen peroxide is decomposed in 300 h standing at room temperature. This high rate of decomposition is indicative of an impurity catalyzed decomposition and the decomposition rate is first order with respect to hydrogen peroxide concentration. This behaviour is similar to that observed by Barb Baxendale George and Hargrave 20 in experiments at high ratios of peroxide to ferric ion concentrations.Thus it is possible that certain types of impurities when present in water lead to first order kinetics. In-sufficient work was done to determine whether this behaviour is carried over into the y-ray irradiated solutions EDWIN J . HART AND MAX S . MATHESON 187 DISCUSSION THIRD ORDER TERMINATION IN HYDROGEN PERoxmE.-Since the experimental evidence clearly favours a rate of decomposition proportional to the square root of absorbed radiation and to the square root of hydrogen peroxide concentration, any mechanism adopted must account for these facts. In our mechanism the correct dependence on radiation intensity and ( H 2 0 2 ) is obtained through the termination reaction HO2 +- H 0 2 t H202 -+ 2H202 -t 0 2 . (5) Since only a few termolecular reactions are known it is pertinent to consider the proposed reaction in more detail.First it may be noted that the work of Swain 29 indicates that third order reactions are not as rare as has been generally believed. He has shown that in the enolization of acetone in water solution and in the dis-placement reactions of organic halides in benzene solution three reactants are involved in the rate determining steps. Secondly it is not necessary that termination in our mechanism occur by a true termolecular reaction but only that the termination be effectively third order. This third order effect can be derived if the formation of complexes such as those in equations (4') and (5') of our paper on hydrogen peroxide is postulated. Thus steps (4) and ( 5 ) of the proposed mechanism may be replaced by the following : HOz -+ H202 -+ [ H 0 2 - H202J kc(H02)(H202) (44 (4b) (44 [HOz - H2021 + HO2 + H202 kd(LH02 - H2021) [HOz - H 2 0 2 ] - + H 2 0 + 0 2 + OH k4c([H02 - H2021) [ H 0 2 - H 2 0 2 ] $- HO2 + 2H202 -t 0 2 k5a([H@ - H2021)(H02) (5a) The steady state assumption and certain others may be applied to the rnechan-ism as now modified.Three such cases will be considered here. Case 1. Steps 4a and 4b are taken to be much faster than 4c or 5a i.e. the complex attains an equilibrium concentration with respect to HO2 and H202, Then the rate of 0 2 evolution is dOz/dt = k4c(kc/kd)%(Kl,/ksu)~(H202)% + Kla (18) as required by the experimental facts. If in addition kc(H202) < kd then the kp we have measured is k4kc/kd and 2kt = 2ksakc/kd.This is the situation if only a small portion of the H02 is in the equilibrium concentralion of complex. In the limiting case of short-lived unstable complex formation kc would become the rate constant for formation of activated complex and k4c/kd would become the transmission coefficient for reaction. Case 2. If kd = 0 so that irreversible formation of complex occurs the rate is given by Next let kc(H202) > kqc so that rapid formation of complex is followed by slow decomposition. Then k p = k4c/(H202) 2kr = 2k5ak4c/[kc(H20#] and the rate constants as measured would decrease with increasing H202 concentration. Case 3 . If kd = 0 and k4c > kc(H202) so that slow irreversible formation of complex is followed by rapid decomposition then kp = k and 2kt = 2 k ~ ~ k ~ / k 4 ~ . From the above discussion it appears that a termolecular termination real or apparent is reasonable. It remains then to consider the magnitude found for the termination rate constant. Possibly the highest termolecular rate constant for the gas phase is the 1010 1.2 moles-2 sec-1 measured by Smallwood30 for the number of effective collisions of three hydrogen atoms. In comparison with this value our k r == 2.7 X 1010 1.2 moles-2 sec-1 is somewhat higher. However, this is not an impossibly high value even for a true termolecular reaction as can d02/dt == k,tlt~c(Kr,/k5a)'(H202)' + Kla. (19 188 DECOMPOSITION OF HYDROGEN PEROXIDE be seen from the following argument. If two HO2 radicals collide in aqueous solution they will have 10 or more molecules as neighbours in the cage around them. Since pure H202 is about 40 M this would mean that in a 0.1 M H202 solution one in every 40 pairs of colliding H02 radicals would have an H202 molecule as a neighbour. Further the rate constant for bimolecular collisions is about 3 x 1011 1. moles-1 sec-1 so that the rate at which H02 pairs would collide in the presence of Hz02 (for 0.1 M H202) is 3 x 1011 x (1/40) x (H02)2 = 0.75 x 1010 x (H02)2 moles 1.-1 sec-1. From our results kt(H202)2 = 2.7 x 1010 x 0.1 x (H02)2 = 0-27 x 1010 moles L-1 sec-1 indicating that about one in three collisions would lead to termination which is a rather high efficiency. On the other hand complex formation may be present to account for such high efficiency. 1 Allmand and Style J. Chem. Soc. 1930 596. 2 Lea Trans. Faraday SOC. 1949 45 81. 3 Fricke J . Chem. Physics 1935 3 364. 4 Johnson J. Chem. Physics 1951 19 1204. 5 Volman J . Chem. Physics 1949 17 947. 6 Hart J . Amer. Chem. SOC. 1951 73 68. 7 Lindsay Introduction lo Physical Statistics (John Wiley and Sons Inc. New York, 8 Orr and Butler J. Chem. Soc. 1935 1273. 9 Briers Chapman and Walters J . Chem. SOC. 1926 562. 10 Matheson Auer Bevilacqua and Hart J. Amer. Chem. Sue. 1951 73 5395. 11 Grassie and Melville Proc. Roy. SOC. A 1951 207 285. 12Bateman and Gee Pruc. Roy. SOC. A 1948 195 391. 13 Bamford and Dewar Proc. Roy. SOC. A 1949 198,252. 14 Matheson Auer Bevilacqua and Hart J . Amer. Chem. Sue. 1949 71 497. 15 Hart in preparation. 16 Burns and Dainton Trans. Faraday Soc. 1950,46,411. 17 Hochanadel paper presented at 119th Amer. Chem. SOC. Meeting (Cleveland Ohio, 18 Toulis University of California Radiation Laboratory Report No. 583 Feb. 10 1950. 19 Agar and Dainton Faraday SOC. Discussion 1947 2 218. 20 Barb Baxendale George and Hargrave Trans. Faraday Sue. 1951 47 462. 21 Bray J. Amer. Chem. SOC. 1938 60 82. 22 George Faraday SOC. Discussion 1947 2 196. 23 Allen Hochanadel Ghormley and Davis paper presented at 119th Amer. Chem. 24 Hart paper presented at 119th Amer. Chem. SOC. Meeting (Cleveland Ohio 1951). 25 Haber and Willstatter Ber. 1931 64 2844. 26 Kornfeld 2. tviss. Phot. 1921 21 66. 27 J. Weiss private communication. 28 E. S. Shanley private communication. 29 Swain J . Amer. Chem. Soc. 1952 72 2794 4578. 30 Smallwood J . Amer. Chem. Suc. 1934 56 1542. N.Y. 1941) p. 15. 1951). SOC. Meeting (Cleveland Ohio 195 1)

ISSN:0366-9033    

DOI:10.1039/DF9521200169

出版商:RSC    

年代:1952

数据来源: RSC

摘要:

THE FORMATION AND DECOMPOSITION OF H202 IN AQUEOUS SOLUTIONS BY THE ACTION OF HIGH ENERGY ELECTRONS AND X-RADIATION BY M. EBERT AND J. W. BOAG Radiotherapeutic Research Unit, Medical Research Council, Hammersmith Hospital, London Received 22nd February, 1952 Experiments are reported on the decomposition of H20 and H202 by 1 MeV electrons and by 1.2 MeV and 200 kV X-radiation. The dose rate ranged from 300 ergs/g sec to 107 ergs/g sec and the radiation dose from about 105 up to 1010 ergs/g. The decomposition of neutral water showed no dependence on the radiation intensity, but in acidified water the equilibrium concentration of H202 produced by large doses was dependent on intensity. The experiments reveal a marked difference between 1 MeV electron or X-radiation on the one hand and 200 kV or lower energy X-radiation on the other.The experimental arrangements, and the methods of dosimetry used for the various types of radiation are described. EXPERIMENTAL APPARATUS USED AND IRRADIATION TECHNIQUE.-The irradiations described in this paper were carried out with several different radiation sources. A Van de Graaff gener- ator was used for the fast electron irradiations and the 1.2 MeV X-rays, and a 200 kV Philips Metalix X-ray tube for the other X-ray exposures. ELECTRON RADrATIoN.-The experimental arrangement for the electron irradiations is shown in fig. 1. The electron beam, accelerated to 1-2 MeV, emerged from the vacuum tube through a 0.002 in. thick A1 diaphragm and then passed through a monitoring ionization chamber of the type described elsewhere.1 Immediately below this was an aluminium shutter, electrically controlled, by means of which the beam could be stopped, or allowed to pass into the irradiation apparatus.The latter contained an ionization chamber similar to the monitoring chamber, below which was a movable platform sup- porting the glass vessel containing the water sample, which could thus be brought up close to the ionization chamber. The current from the ionization chamber was taken to a galvanometer or microammeter giving a continuous indication of dose rate and, for short irradiations, also to an integrating instrument. The spacing between the foils of the ionization chamber was 1 mm and the stop defining the beam was 3.9 cm diam. The volume of air irradiated between the foils was thus 2.385 cm3, and the ionization current at 10s e.s.u./cm3 sec was 79.5 PA. A study of the saturation curve on a chamber of this kind2 showed, however, that even at this high ionization intensity a potential of 200V was sufficient to collect 98 % of the theoretical saturation current.The radial variation in intensity over the surface of the liquid sample was investigated by means of a small cylindrical ionization chamber with a thin aluminium wall, which was relatively transparent to electrons. The curve obtained when this chamber was moved across the beam is shown in fig. 2. The full line in this figure is a Gaussian curve of standard deviation 2.53 cm. The ionization intensity produced by the electron beam at various depths in the water was determined by means of the apparatus illustrated in fig.3. In this, the electron beam was defined by a brass cone and then passed through a known thickness of water in a cup having a thin aluminium foil base. The cup was so designed that the meniscus was not irradiated, so the absorption and scatter occurred in a parallel slab of water. The depth of water corresponding to a known volume pipetted into the cup was measured by focusing a microscope first on the foil surface and then on the 189190 water surface. The difference was corrected for refractive index. The depth dose curve obtained as described is shown in fig. 4. FORMATION AND DECOMPOSITION OF H202 0.7- PO /5 /n 5 F~G. 1 .-Sectional view of the arrangement for electron irradiations. D i s h , <$ OXIS ( m ~ ) 5 10 I5 PO As an additionaI qualitative check on the dose distribution under electron radiation some strips of bromide paper were cut, of width 3-2 mm (i.e.equal to the depth of water normally used in the glass dish). These were supported upright across a diameter of the 0.8 t B dish containing the usual 3.5 ml of water and were irradiated. The blackening was a maximum at a depth of between 1 and 2 mm below the surface. It was clear that the meniscus was shading the liquid in the outer corner of the dish but the photopaper did not permit quantitative measurements to be made.M. EBERT AND J . W. BOAG 191 The isodose curves drawn in fig. 5 are based on the depth dose measurements of fig. 44 and the radial intensity variation of fig. 2. It can be seen from curve b in fig.4 that, for a liquid depth of 3.5 mm the mean dose throughout the depth of the liquid would be just equal to the surface dose. The volume of liquid normally used was 3.5 ml and, but for the meniscus, this would have given a depth of 3.5 mm in the glass dish. The actual Thin m7f 1 \ /bullZOhibul c hnm,bgr/ WUjM FIG. 3.-Apparatus for determining electron depth dose in water. depth at the centre was 3.2 mm. The 0.3 ml of water in the corner of the dish, which was screened almost completely by the meniscus would, if it had occuped a thin layer at the bottom of the dish, have received about one-third of the average intensity, and thus contributed only some 3 % to the total energy absorption. It will be sufficiently accurate to assume that the mean throughout the liquid was the same as it would have ‘x’mrn dep/h of Walter FIG.4.-Depth dose in water for the 1 MeV electron beam. been for a parallel slab of water 3.5 mm in thickness, and this, according to fig. 4b is equal to the superficial intensity. The ionization chamber measured the average ionization intensity in e.s.u./cm3 sec over the liquid surface. This figure, converted to ergs/g by multiplying by 93, is the mean dose rate quoted in the present paper, and it is thought to be within 5 % of the true mean energy absorption in the irradiated sample. ELECTRON SHuTTER.-In some of the electron irradiations it was desired to give a small dose at very high dose rate and this involved exposures of less than 1 sec. The electron shutter normally used was not accurate for times less than a few seconds and so a rotary shutter consisting of two aluminium discs geared together was constructed. Each disc had a 30 degree sector cut out and these coincided once in 10 complete revolutions of the faster disc.The discs were driven by a variable speed electric motor and the time1 92 FORMATION A N D DECOMPOSITION OF H202 during which the electron beam could pass was variable from about 1 to $6 sec. The procedure adopted was to adjust the rotary shutter to the correct speed and leave it rotating, switch on the electron beam with the slow speed shutter closed, set the beam to the desired intensity, and then open the slow speed shutter. As soon as the current pulse from the ionization chamber above the sample was recorded on the integrating meter, the slow speed shutter was closed. The whole procedure took about 10 to 15 sec and, with the l/lOth sec exposures, the X-rays due to the electrons striking the upper aluminium shutter gave a dose to the sample which was about 10 % of the electron dose.1.2 MeV X-RADIATIoN.-The X-rays were produced by a 1.2 MeV electron beam striking a lead target 2 mm thick; with additional filtration amounting to 1 mm of copper and 3 mm of steel. For this work a special jig was constructed in which the glass dishes containing the liquid samples were supported at different distances from the target and irradiated simultaneously. The first dish was at 5 cm from the target spot and the farthest at 32 cm. The arrangement for the irradiations is illustrated in fig.6. The actual dose rate at each position was measured by means of a small Perspex ionization chamber with graphited collecting surfaces. This chamber was calibrated against four bakelite-graphite condenser ionization chambers which had been standardized at the N.P.L. For comparison with the electron irradiations, all dose values quoted have been FIG. 5.-Isodose curves for the electron irradiations. expressed in ergs/g, the conversion factor being taken as 93 erg/g rontgen for the 1.2 MeV X-radiation. 200 kV RADIATION.-A Philips Metalix 200 kV X-ray tube was used for this work. The tube was run at 195 kV and at currents up to 6 mA with no filtration except the in- herent filtration of the tube window. In order to obtain a high intensity the sample had t o be placed close t o the X-ray shutter.The dose rate at this point was measured by means of condenser ionization chambers. The capacity of these chambers was not large enough to permit the accurate measurement of the intensity at full tube current so two chambers were placed close to the shutter, and two others at a distance of about 20 cm from the target. The latter two chambers were charged to a lower potential and their loss of charge measured on a more sensitive scale of the electrometer. The positions of the upper and lower chambers were standardized by using a jig and the ratio of dose rate a t the top to dose rate at the bottom could therefore be determined accurately. The absolute intensity at the position of the lower chambers was then determined by charging them t o a higher voltage and exposing for about 20 sec.This method of measurement, which circum- stances compelled, was not as accurate as would have been desired and measurements made at different times showed discrepancies of the order of 4 %. An independent check on the intensity using a Victoreen 250 r chamber agreed within 3 %. Owing to the short target-to-sample distances and the soft radiation used, there was an appreciable variation of intensity throughout the solution. The depth dose of the radiation in water was measured and the figures quoted are mean values throughout the solution.M. EBERT AND J . W. BOAG 193 TEMPERATURE RISE DURING IRRADIATION.-The values of radiation dose delivered in this investigation were often large enough to cause an appreciable temperature rise if delivered instantaneously.The irradiation vessel, however, was in close contact with a inass of metal many times that of the water it contained, and the final temperature of the water at the end of the irradiation depended greatly on the amount of heat removed by conduction. The rate of rise of temperature during irradiation was measured in a special series of experiments, using a fine copper-constantan thermocouple immersed in the water. In a typical run at a dose rate of 5.1 x 106 ergs/g sec the initial rate of rise was 0.11" C/sec, the final rise after a dose of 3.25 x 108 ergs/g was 4.4" C, and the equilibrium tempera- ture for this dose rate when conduction balanced input, appeared to be about 8" C . These thermal measurements, incidentally, pro- vided a rough check on the ionization measure- ments of energy absorption. In the foregoing example the initial rate of rise to be expected from the ionization measurements was 0.12" C/sec.The temperature rise in the irradiation experi- ments reported would have been limited in most cases to a few degrees or fractions of a degree. Even in the few cases in which the largest doses were given at high dose rate the rise would not have exceeded 10" C. It is shown later that in acidified solution a temperature difference of 40" C between samples changed the H202 equili- brium value by only 30 %. It is unlikely, therefore, that the temperature rise due to the irradiation had an important effect on the HzOz concentrations observed. PREPARATION OF THE WATER FOR IRRADIATION. -Distilled water was re-distilled from alkaline potassium permanganate in an all-Pyrex still.Purity was checked by measuring the conductivity. If the specific resistance was less than 400 kQ at 15" C the water was discarded. The pH of water accepted for use was between 6-1 and 6-9, this small acidity being due presumably to dissolved carbon dioxide. Aeration was carried out by sucking a vigorous stream of air through the water for not less than 20 min. This air had been bubbled through 10 cm of distilled water and had passed three wet sintered glass filters to remove all dust. Redistillation of such water from alkaline manganous salts in an all-quartz still or repetition of the distillation in an all-Pyrex still with a 2 ft high fractionating column, did not alter the results of the irradiation experiments.On aerating the water without the three sintered glass filter plates, the curve of H202 , \ I 1 1 FIG. 6.-Experimental arrangement for the 1.2 MeV X-ray experiments. formation against dose showed a maximum (at approximately 2 x lo7 ergs/g) which was higher than the equilibrium value ultimately reached, but was not reproducible. We believe this was mainly due to dust or other impurities from the air used for aeration. These impurities may react irreversibly with the active radicals primarily formed and thus prevent the back reaction. Once they are used up an equilibrium value of H202 is reached which is lower than the maximum. ESTIMATION OF H202.-The estimation of hydrogen peroxide was carried out colori- metrically using the colour formed with titanium sulphate reagent.3 The colorimeter used for most of the work was a Hilger Spekker with violet filter.This was compared with a Beckmann spectrophotometer at a wavelength of 410 mp and the agreement was satisfactory. For the measurements of optical density in the work with 1.2 MeV X-rays, and in the comparison of X-ray and electron effects shown in fig. 13 the Beckmann instru- ment was used. All pH measurements were made with a Cambridge pH meter with glass electrodes. The conductivity was measured, using commercially available platinum194 electrodes, in a 50 cycle a.c. bridge circuit. Calibration of the conductivity cells was carried out using solutions of known conductivity. All reagents used (H2S04, H3P04, €3202) were of A.R.or of M.A.R. standard. IRRADIAmoNs.-Unless otherwise stated, 3.5 ml of the solution prepared for the ir- radiation was pipetted into a cylindrical glass cell 3.5 cm diam. with an optically flat bottom. The solution was usually open to the atmosphere during the irradiation, but when 0 2 , N2 or air was bubbled through the solution, a special arrangement was used which is shown in fig. 1 (inset). After the end of the irradiation, the solution was re- moved from the vessel as soon as possible by pipetting it into acid titanium sulphate reagent. The colour thus obtained remained constant for many days in the dark. GAS BUBBLING TECHNIQUE.-FOr these experiments a 4 ml sample of water or H202 was placed in a sintered glass filter S3 through which the desired gas was passed continu- ously.The general arrangement is illustrated in fig. 1 (inset). The dish was covered with very thin Pliofilm held in position by an aluminium ring. The gas escaped through the creases in the film. The loss of solution due to the gas bubbling was about 10-20 % for a run of 10 min. The gas pressure was adjusted to give a vigorous gas flow throughout the run. A series of experiments was carried out to establish that bubbling the various gases through the solution without irradiation did not give rise to any decomposition of H202. It was found necessary to keep above a minimum gas flow in order to obtain reproducible results. pared as usual were air saturated at 0" C, 20" C and 40" C respectively in separate flasks, then transferred without change of temperature to glass vessels which were kept in covered glass containers at the appropriate temperature. The glass vessels containing the solution were then inserted in turn into the aluminium cup shown in fig.1 at the same temperature, and were irradiated without delay. The irradiation usually started about 1 min after the vessel had been removed from the covered container. The radiation intensity was 1.4 x 106 ergs/g sec and the total dose was 1.6 x 108 ergs/g. FORMATION AND DECOMPOSITION O F H202 TECHNIQUE FOR STUDYING THE INFLUENCE OF TEMPERATURE.-saMpleS Of Water pre- H202 FORMATION IN NEUTRAL peroxide in purified and aerated hydrogen peroxide concentration RESULTS M'A'rER.-Fig. 7 shows the formation of hydrogen water when irradiated with 1 MeV electrons.The approached equilibrium at about 4 x 107 ergs/g. i 0 0 :252 0 0 : FIG. 7.-The formation of H202 in neutral water irradiated by 1 MeV electrons. There was no systematic increase over this equilibrium value in test runs up to a dose as high as 1010 ergs/g. No dose rate dependence was apparent for the initial part of the curve but the equilibrium value appeared to be slightly lower at higher dose rates as may be seen from the distribution of points in fig. 7. The observations are not sufficiently constant, however, for this to be considered certain.M . EBERT AND J. W . BOAG 195 TabIe 1 gives the separate observations of H202 concentration in pmolesll. made over a period of some 12 months for a range of intensities and total doses.The scatter of the points is considerably larger than can be accounted for by the accumuIated errors in the several techniques employed namely ionization current measurements, pipetting, estimations of optical density. Moreover, observations taken on any one day showed greater consistency. It appears, therefore, that some variable factor operated from day to day. This factor may possibly have been impurities in the air used for aeration, which was filtered to remove dust but not otherwise purified. Differences in oxygen tension due to variation in room temperature and barometric pressure, are another possible source of the discrepancies. The matter is still under investigation. TABLE I.-FORMATION OF H202 IN NEUTRAL WATER IRRADIATED 41 7.3 12.1 24-2 48.5 73.0 146.0 146.0 146.0 292.0 292.0 292-0 154 2.3 4.6 9.2 9.2 18.5 27.7 46- 1 92-5 240 3.6 7.2 14.4 14.4 43-2 72.0 173.0 288.0 433.0 /imole!l 61 61 77 141 138 168 162 138 171 182, 174 39 56 95 92 I15 122 11s 120 47 68 94 100 138 147 154 168 162 IN THE dose rate ergsig sec x 1 0 - 3 410 845 4,100 8,840 1 MeV ELECTRON BEAM total dose ergs/g x 10-6 41.5 41-5 41.5 115.0 249.0 249-0 249.0 243.0 1,390.0 41.6 42.2 42.7 240.0 236.0 241.0 41.3 40.5 43.6 61.6 60.5 61.0 10.5 11.7 11.3 12.5 23.7 26.6 pmole/l 127 118 118 154 150 144 156 153 165 111 111 103 146 130 130 123 112 107 144 122 145 75 85 79 81 109 129 dose rate total dose ergs/g sec ergs/g6 pmolejl.~ 1 0 - 3 X I O - 8,840 con t inued 23.8 24.1 29-9 37.1 37.1 36.9 121.0 119.0 1 16.0 144.0 246.0 3 17.0 3 19.0 10,250 41.0 41.8 41-0 247-0 248.0 248.0 252.0 2524 252.0 2534 256.0 110 114 115 124 128 1 20 142 147 143 142 143 141 135 130 121 127 141 135 132 I50 127 138 135 132 Bearing Fig.7 has been drawn in two parts, with a change of scale between them. this in mind, the simple exponential curve which is drawn, provides a reasonable fit to the experimental points. The equation is y = 140 (1 - e-kx) with y in ,umoles/l, x in ergs/g, and k = 8-5 x 10-8. The initial ionic yield calculated from this curve is 0.38. Although the points in Fig. 7 would be consistent with values between 0.3 and 0.4, subsequent work, to be reported elsewhere, gives a value of 0.34. with sulphuric acid to a pH of 2 to 3, was irradiated, no equilibrium was reached in test runs to a dose as high as 109 ergs/g.The hydrogen peroxide concentration reached was of the order of 1,000 pmoles/I. for 109 ergs/g. Table 2 gives the hydrogen peroxide con- centration in pmoles/l. for a range of intensities and total doses. H202 FORMATION IN ACIDIFIED WATER.-When Water, purified, aerated and acidified196 FORMATION AND DECOMPOSITION OF H202 In this experiment the pH of the water samples varied from 2-0 to 2.8 and this may explain in part the irregularities in the table. In spite of this, however, the observations reveal a marked dose rate dependence for hydrogen peroxide formation in acid solution at pH values below 3. Dilute phosphoric acid solutions were irradiated at a dose rate of 1.4 x 106 ergs/g sec and showed the same effect, i.e. no equilibrium concentration of hydrogen peroxide was reached even with a dose as high as 109 ergs/g and the highest H202 concentration obtained was 1,030 pmoles/l.TABLE 2.-FORMATION OF H202 IN ACIDIFIED WATER BY 1 MeV ELECTRONS dose rate ergs!g sec 24 x 103 49 9 , 95 Y, 190 9 , 384 9 9 770 9 , 3,840 ,, 7,690 ,, Hz02 concentration in Icniolesll. for a total dose of 39 x 106 ergsjg 417 417 382 332 385 365 353 348 78 x I 0 6 ergs,'g 155 x 648 616 595 515 504 473 500 404 1 0 6 ergs/g 836 765 68 1 570 565 5 10 486 463 PEROXIDE FORMATION IN WATER CONTAINING Ti(SO& REAGENT.-cUrVe 2 in fig. 8 shows the effect of 1 MeV electrons on the titanium sulphate reagent. The reagent is 0-8 N in sulphuric acid and contains 0.16 g Ti/L The optical density of the irradiated solution was measured, and expressed as the amount of hydrogen peroxide necessary to produce the same density.Curve 2 crosses curve 1 at a level of approximately 5 x 108 ergs/g where the irradiated acid HzO contains 880 pmoles/l. hydrogen peroxide. At this level the radiolysis of H202 is effective in the acid solution, whereas the titanium peroxide complex is much more stable and increases until more than 50 % of the Ti 4 / 2 3 10 20 50 /00 200 500 Id00 2OyO FIG. &-The formation of H202 in the irradiation of titanium sulphate reagent. present is oxidized. The titanium sulphate solution seems, therefore, suitable for chemical dosimetry at these very high dose levels, especially as it does not show a dose rate dependence for intensities between 9,000 ergs/g sec and 2 x 106 ergs/g sec. If the titanium concentration is kept constant at approximately 0.1 6 g titanium/l.and the sulphuric acid concentration changed, the optical density shows a flat maximum between 1 N and 0.05 N sulphuric acid. On increasing the sulphuric acid concentration the colour formation drops rapidly. BREAKDOWN OF HYDROGEN PEROXIDE.-A number of aerated dilute hydrogen peroxide solutions (5 ml) were given a dose of 1 1 x 108 ergs/g at an intensity of 9 x 106 ergs!g sec. The oxygen liberated by the breakdown of H202 cannot diffuse out of the solution during the relatively short irradiation times used (about 2 min). The initial H202 concentration may thus be regarded as an additional reservoir of oxygen. The slope of the curve in fig. 9 is, therefore, determined mainly by the time taken for the irradiation and by the surface to volume ratio of the solution.This result suggests that the temperature at which the water is air saturated will play a part in determining the final equilibrium value. The solubility of 0 2 in water varies from 1 3 . 2 ~ 02/1. at 0' C to 6 . 5 ~ 0 2 at 40" C . A few pilot experiments showed that theM. EBERT AND J. W . BOAG 197 H202 yield for aqueous solutions of pH 2.5 dropped by approximately 30 % between 0" C and 40" C. It is planned to investigate this temperature effect under better controlled conditions. These experiments indicated that when H202 is decomposed by radiation, the diffusion of 0 2 out of the solution may exert a controlling influence on the back re- action, and thus on the H202 level attained. Irradiations of solutions of H202 were carried out with different gases bubbling solution.Fig. 10 shows a typical series of experiments. Curve 1 taken gas bubbling, flattens out quickly, which we take to indicate an increased through the without any /I./_.lj'O't?s_!%- / I I r e 200 300 400 I /po , 1 FIG. 9.-The amount of H202 formed by the dose of 7 x 108 ergs/g as a function of the initial concentration of H2Oz. A I W / 2 4 5 I 1 1 I FIG. 10.-Breakdown of H202 1, with no gas bubbling through the solution ; 2, with air bubbling ; 3, with nitrogen bubbling. H202 production due to the high 0 2 tension arising from breakdown of H202. Curve 2 shows the breakdown of H202 when air is bubbled through the solution during irradi- ation. Part of the excess 0 2 , formed during the irradiation, is carried away by the air and the curve therefore falls to a lower level than curve 1.Curve 3 shows the effect of bubbling nitrogen through the solution. The H202 concentration decreases very rapidly and no change in the gradient of the decomposition curve is noticeable until a concentra- tion of about 100 pmoles/l. is reached. After 4 x 1Gs ergs/g, the remaining H202 was no longer detectable by the titanium method. If, however, H202 acidified with sulphuric acid was used under the same conditions, the concentration fell much less rapidly and even after 9 x 109 ergs/g the H202 concentration was 235 pmoles/l. and was still declining.198 FORMATION AND DECOMPOSITION OF H202 Irradiating neutral water at 9 x 106 ergs/g sec to a dose of 5.5 x 108 ergs/g with 0 2 bubbling through the water, led to an equilibrium value for H202 of 557 pmoles/l.4- N& u/’ra/ SO/U,f/ 0 r7 0 Ac/-d . z J l o L , d r ~ 1 -I x 10 a ergs Y-;% - 50 I-? I I I with oxygen bubbling through neutral solution ; / 2 3 4 5 FIG. 1 l.-H202 formation and decomposition. L, with oxygen bubbling through 0.01 N sulphuric acid solution ; 3: breakdown of H202 with oxygen bubbling through. FIG. 12.-The decomposition of H202 with nitrogen bubbling through the solution in (a) acid and (b) neutral solution. (curve 1, fig. 11). If, however, the water was acidified to 0.01 N sulphuric acid, the equilibrium value reached was 1,290 pmoles/l. (curve 2, fig. 11). When hydrogen per- oxide at 11,700 pmoles/I. was irradiated with 0 2 bubbling through, the rate of breakdownM.EBERT AND J . W. BOAG 199 was the same for both 0-01 N sulphuric acid and neutral solutions. A dose of 6-6 x 108 ergs/g (curve 3, fig. 12) brought the H202 concentration down to 1,620 pmoles/l., the relation between H202 concentration and dose being the same for neutral as for acid solutions. 1 MeV electrons was different from that published by Lefort 4 for 30 kV X-ray experi- ments. In collaboration with T. Alper, M. Lefort and H. C. Sutton, a series of experi- ments was, therefore, carried out, some in the 1 MeV electron beam and some with 200 kV X-rays. These experiments, which will be reported in the discussion, confirmed both the results reported by Lefort and those reported in this paper. They showed that, when H202 formation is plotted against energy absorption per g in water there are striking differences between 1 MeV electrons and either 200 kV or 30 kV X-rays.A further experiment has since been carried out which confirms these differences, in a particularly direct way. A sample of neutral aerated water (no. 1, fig. 13) was irradiated to 7.5 x lo7 COMPARISON BETWEEN ELECTRONS AND X-RADIATION.-The formation Of H202 Using 77he scale //me scale FIG. 13.-&02 level in a solution irradiated alternately with equal doses of electrons and X-rays. ergs/g in the electron beam at an intensity of 7,000 ergs/g sec. This took 3 h and on completion a second sample (no. 2) was treated in the same way. This dose was sufficient to bring the H202 concentration very close to the equilibrium value found in earlier experiments.In the meantime, two similar samples (no. 3 and 4) had been irradiated to the same dose level at the same intensity in the unfiltered 200 kV X-ray beam. Samples no. 1 and 2 were then mixed and their H202 concentration was found to be 141 pmoles/l. 3.5 ml of this mixture were taken (sample no. 5 ) and irradiated with 200 kV X-rays in exactly the same manner as samples 3 and 4 had been. At the end of ir- radiation the H202 concentration was 221 pmoles/l. 2.5 ml of sample no. 5 were again irradiated in the electron beam (sample no. 6) to a total dose of 1.9 x 108 ergs/g at an intensity of 106 ergs/g sec. The H202 concentration reached was 144 pmolesll. The formation and breakdown of Hz02 in this experiment is indicated graphically in fig. 13a.In the same way samples 3 and 4 were mixed and their H202 concentration was found to be 221 pmoles/l. From this mixture, sample no. 7 was taken, and irradiated in the electron beam under exactly the same conditions as for sample no. 1 and 2. The resulting H202 concentration was 156 pmoles/l. The course of this experiment is shown in fig. 13b. 1.2 MeV X-RAY IRRADIATION.-TO investigate further the effect of radiation quality, a series of experiments was carried out with 1.2 MeV X-radiation. Intensities from 300 to 15,000 ergs/g sec were used, and in neutral solution an equilibrium value of 210 pmoles/l. hydrogen peroxide was reached. The observations are given in table 3. In fig. 14 (curve a), an exponential curve has been fitted to the experimental points. The equation of this curve is y = 210 (1 - exp (- kx)) with k = 10-7 and the initial ionic yield calculated from this equation is 0.66 molecules/ion pair.The scatter of the ob- servations is such, however, that the true value may lie anywhere between 0-5 and 0.9.200 dose rate ergs/g sec 0.308 0.321 0.351 A 10-3 0,617 0.604 0.642 0.604 0.684 1.175 1.280 1.175 1.223 1.150 1.190 1.15 1-30 1-175 2.53 2.53 2-76 2.29 2.53 2-63 2-48 2.32 2-60 FORMATION AND DECOMPOSITION OF H202 TABLE 3.-FORMATION OF H202 IN NEUTRAL WATER IRRADIATED BY 1-2 MeV X-RAYS 0.370 6.2 0.578 2-4 1-68 19.0 0.370 5-3 0.724 17.4 1-13 17.4 1.45 27-0 3.28 51.0 0.352 5.4 0,500 3-2 0.704 18-2 1.100 15-0 1.380 36.4 2-13 39.4 2.76 60.2 5.61 110.0 6-23 91.0 0-380 10.3 0.758 18.2 1-08 36.2 1-37 36.2 1.51 42-3 2-37 46-5 2.97 72.4 4-18 82.0 4.61 86.0 2-48 2.46 2.53 2.80 2-29 3.67 3.73 3.96 3.12 3.50 3.67 5.34 5.34 5-83 5-34 5.50 5.24 5-41 5-24 5.33 5.89 6.63 6-73 7-13 6-30 6.63 5.95 8.12 12.1 13-4 25.8 2.20 6.72 13.1 14-4 23.0 41.5 0.80 1 -60 2-27 3.20 4.98 6-28 9.70 12.6 25.6 28-4 3.98 12.10 23.5 41.5 75.2 H202 vmoles/l 109.0 98.0 155.0 133.0 194.0 52.0 100.0 141.0 185-0 192.0 2 10.0 17.4 39.7 37.0 67.0 77.0 106.0 125.0 185-0 199.0 174-0 67.0 140-0 187.0 210.0 216.0 H202 dose rate total dose eyj$:y ~ y d ~ _ ~ pmoles/l.10.7 1.61 39.8 10.7 3.21 72.6 11.7 4.56 704 10.7 6-42 118.0 11-1 10.0 117.0 10.5 12.6 137.0 10.9 19.5 174.0 10.5 25.2 185-0 10.1 51.4 199.0 11.9 56.9 195.0 14.5 4.39 37.0 14.5 8-71 100.0 14.7 26.5 176.0 15.7 51.7 212.0 13.8 91.1 221-0 14.5 164.0 212-0 In acid solution (PH 2) the H202 concentration was still rising almost linearIy after a dose of 5 x lO7ergs/g, when the H202 concentration was already 1,140 pmoles/l.FIG. 14. (fig. 14, curve b). pair. The observations for acid solution are given in table 4. The initial ionic yield determined from this curve is 0.8 moleculesJonM. EBERT AND J. W. BOAG 201 TABLE 4.-FORMATION OF H202 IN ACIDIFIED WATER IRRADIATED BY 1.2 MeV X-RAYS dose rate total dose dose rate total dose ergsjg sec ergs/$ H202 pmoIe/l. ergs/g sec ergs/g6 H202 pmole/l. x 10-3 x 10- x 10-3 x 10- 0.30 0.3 1 0.3 1 0.58 0.62 0.59 0.62 1.16 1.12 1.18 1.15 1-18 0.18 0.74 1-57 0.35 0.74 1-45 2.97 0.35 0.67 1-42 2-75 5-65 5.6 21.5 43.5 11.5 26.4 43.8 101.0 11.5 19.7 51.2 83.1 182.0 5.3 5.3 5.1 5.4 5.2 5.4 10.7 10-7 10.2 10.8 10.5 10.8 0.83 1.59 3-04 6.45 12.5 25.7 1.59 3-18 6.1 1 12.9 25.1 51-6 21.5 41.7 85.8 174.0 323.0 634-0 41-2 81.0 147.0 321-0 578.0 1,140-0 2.5 1 0.75 21-5 2.41 1.45 42.9 2.54 3-04 84.0 2.48 5-93 165-0 2.54 12.1 355.0 DISCUSSION The equilibrium value of hydrogen peroxide formed in the electron beam will clearly depend on the velocity of the overall back reaction.A possible reaction scheme for the decomposition of hydrogen peroxide which has been discussed by Haber and Weiss 5 and others, is the following : H202 + OH -+ H20 + HO2 H202 + H02 -+ H20 + OH + 0 2 H202 + 02- -+ OH + OH- + 0 2 [H202 + O* -+ H20 + 0 2 €302 -1- OH1 20H --+ H202 20H --+ H20 + O* 2H02 + H202 + 0 2 OH -+ H02 + H20 + 0 2 The results reported above for the decomposition of hydrogen peroxide in acid or neutral solutions when N2 or 0 2 is bubbled through during the exposure are summarized in table 5.TABLE 5.-EFFECT OF IRRADIATION WITH GAS BUBBLING THROUGH SOLUTION no. gas used acid solution neutral solution 1 N2 reaction velocity falls ; reaction velocity maintained ; 235 pmole/l. H202 left after no hydrogen peroxide detect- an exposure of 9 x 109 ergs/g able after 4 x 107 ergs/g reaction velocity identical for both solutions and of the same order as for N2 bubbling in acid solution 2 G 0 2202 FORMATION AND DECOMPOSITION OF H202 Bubbling a gas through the solution removes all dissolved gases except itself. Nitrogen removes 0 2 , and oxygen leaves 0 2 . In acid solution the reaction proceeds in the same way for both gases, so, in this case, 0 2 plays no part.Hydrogen could play some part, however, and it is planned to investigate this by bubbling hydrogen through. In neutral solution the H202 concentration is reduced more rapidly with N2 bubbling than with 0 2 , so we conclude that in this case the oxygen slows down the decomposition or increases the back reaction. According to the reaction scheme suggested, the only molecules of gas formed are 0 2 molecules. Line 1 in table 5 therefore indicates that the neutral solution contains more removable 0 2 molecules than the acid solution. It appears that 0 2 is not formed at the same rate in the acid solution, but instead, the H202 level is maintained by continuous synthesis from radicals which, in the neutral solution, might be responsible for the formation of 0 2 .When oxygen is bubbled through the solution (table 5, line 2) the only gas molecules present in abundance in both acid and neutral solution are those of 0 2 . In this case there is no difference in the behaviour of the two solutions. It seems probable, therefore, that the hydrogen ion concentration influences one or more steps in the reaction by which 0 2 is formed from either H20 or H202. On the other hand, in the presence of excess oxygen, the decomposition of H202 proceeds in the same way in neutral and acid solutions (table 5, line 2) and this suggests that the utilization of 0 2 for the formation of H202 is not influenced much by pH. The important reactions in the scheme suggested which lead to the formation of 0 2 are eqn.(2), ( 5 ) and ( 6 ) and ihese all involve H02. Reaction (2b) depends on O*, which is derived from (4), and this latter reaction will have only a small probability at the dose rates used in this work. A high hydrogen ion concentration therefore seems either to stabilize H02 or at least to reduce the probability of reaction (2). This would interfere with an essential chain propagat- ing step for the breakdown reaction. Volman,6 studying the kinetics of H 2 0 2 + 0 3 * , concludes that the absence of a large temperature coefficient for this gas reaction, renders the reaction H202 -I- HO;! improbable. He also concludes that (2b) is unlikely in the gas phase, but one would, of course, expect it to take place more readily than (2c) in the liquid phase. In the reaction scheme quoted, the only species likely to be affected by a change in hydrogen ion concentration is the HO2 radical.As pointed out by Weiss,7 and by many other workers, the HO2 radical is dissociated into H+ and 02. We conclude, therefore, that HO2 itself does not react easily with H202 but that 02- is the more reactive species. It is interesting to note that Barb et al.8 found an analogous result for the reaction of H02 with ferric or cupric ions and that Bray 9 expressed similar views based on the reaction of KO2 + H20. Another relevant fact is that even 9 x 109 ergslg of electron irradiation to an acid solution of HzOz with N2 bubbling continuously only brought the Hz02 concentration down to 235pmoles/l. We can conclude that, in this case, any reaction forming H202, such as (5), must be favoured. If reaction ( 5 ) is to be important, then the product of concentration and reaction probability must be higher for the HOz radical than for the other radicals present. A high H02 concentration would arise if the view expressed above on the relative reactivities of HO2 and 0 2 - is correct. High ion density leads to a high radical concentration in the tracks and thus favours radical-radical reactions. It has already been suggested that reaction (5) is mainly responsible for the formation of H 2 0 2 near the equilibrium, and that (1) and, in neutral solutions, (2a) may be responsible for the decomposition of H202 under equilibrium conditions. Reaction ( 5 ) should be favoured by high ion density to a greater extent than (I), thus leading to a higher equilibrium con- centration of H202, in agreement with the observations. Summing up, the amount of H202 formed by a given radiation dose is dependent on both ion density of the radiation and pH of the solution. For any ion densityM. EBERT AND J. W. BOAG 203 the amount of Hz02 formed is larger for acid than for neutral solutions. Similarly for any pH the amount of H202 formed is larger for high ion density than for low ion density radiations. We are greatly indebted to Dr. L. H. Gray for his constant advice and en- couragement during the progress of the work. Many of the earlier irradiations on the Van de Graaff generator were carried out by Mr. T. Wilson, and most of the later ones by Mr. D. Moore, to both of whom we extend our thanks. 1 Boag, Pilling and Wilson, Brit. J. Rad., 1951, 24, 341. 2 Boag and Wilson, J. Appl. Physics, 1952, in press. 3 Eisenberg, Ind. Eng. Chenz., 1943, 15, 327. 4 Lefort, Tlzbes (Paris, 1950). 5 Haber and Weiss, Proc. Roy. SOC. A, 1934, 147, 333. 6 Volman, J. Amer. Chem. SOC., 1951, 73, 1018. 7 Weiss, Trans. Faraday Soc., 1935, 31, 668. 8 Barb, Baxendale, George and Hargrave, Trans. Faraday SOC., 1951, 47,462. 9 Bray, J. Amer. Chem. SOC., 1938, 62, 3357.

ISSN:0366-9033    

DOI:10.1039/DF9521200189

出版商:RSC    

年代:1952

数据来源: RSC