摘要:

SALT EFFECTS ON THE RATE OF EXCHANGE OF THALLOIJS AND TH-FAELIC IONS IN WATER AND HEAVY WATER BY S . W. GILKS AND GWYNETH M. WAIND Chemistry Department, Queen Mary College, University of London, Mile End Road, London, E.l Received 28th January, 1960 At 25°C in perchlorate solutions the thallous-thallic exchange rate in water is 1.5 times as fast as in heavy water. There is a linear decrease with increase in perchlorate ion con- centration from 1.0 to 5.0 M; in this range replacing protons by sodium or barium ions has little effect on the exchange rate. At an ionic strength of 3.0 the thallic-ferrous re- action is catalyzed by hydroxyl ions ; the thallous-thallic exchange is not. Both reactions are catalyzed by platinum metal. The thallous-thallic reaction is one of a group for which the hydrogen-ion inhibition is known to change markedly with increase in ionic strength.l.2 Such inhibition is usual for oxidation-reduction reactions of metal salts and has been interpreted as due to the faster reaction of hydrolyzed as compared with unhydro- lyzed metal aquo ions.3-9 This is of particular interest for the ferrous-ferric electron exchange because a mechanism of hydrogen-atom transfer has been based on the much greater catalytic effect of hydroxyl ions (as compared with other small anions) together with the size and existence of the isotope effect when the solvent (water) is replaced by heavy water.loy11 This is an analogous study of the exchange reaction between thallous and thallic ions in perchlorate solutions of ionic strengths greater than one.At these concentrations equilibrium properties show that ionic atmosphere effects are small compared with specific anion-cation interactions. The thallous-thallic exchange involves a formal two-electron transfer and is catalyzed by platinum metal ;2 we have therefore compared our results with those for the thallic-ferrous oxidation where there is evidence for a two-stage process,l2 and in order to make this comparison more complete, have studied the effect of platinum metal on the rate of this latter reaction. EXPERIMENTAL MATERIALS Conductivity water was used throughout. Deuterium oxide (99h0.5 %) was obtained from I.C.I. in sealed glass ampoules. Thallous perchlorate was prepared by evaporation of a perchloric acid solution of thallous nitrate (obtained from Johnson, Matthey & Co.) and analyzed by titration with potassium iodate solution.13 Thallic perchlorate was prepared in perchloric acid solu- tions from thallous sulphate.12 The thallic concentration was determined gravimetric- ally as oxide 14 and volumetrically with potassium iodide and sodium thiosulphate. The total thallium concentration of these solutions was determined in the same way after reduction with sulphur dioxide, and the formal hydrogen ion concentration by titration with sodium hydroxide after treatment with hydrogen peroxide.This was confirmed by potentiometric titration with standard alkali, correcting for precipitated thallic oxide. Sodium perchlorate was obtained from B.D.H. and was chlaride-free. Initial diff- culties encountered in attempts to prepare and recrystallize sodium perchlorate are de- scribed in detail elsewhere.15 A.R.60 % perchloric acid was standardized by titration. 102S . W. GILKS AND G. M. WAXNL) 103 Barium perchlorate solutions were prepared from a recrystallized B.D.H. product and were analyzed for barium gravimetrically as the chromate and as the sulphate. All the other materials were A.R. grade and were used without further purification. Solvent mixtures of deuterium oxide and water were analyzed using an Abbe refracto- meter. Stock solutions of radioactive thallous and thallic perchlorate were made by adding a small quantity of 204Tl perchlorate obtained from the Radiochemistry Centre, Amersham, Bucks. RATES OF BXCHANGE These were followed at 25 &O*Ol"C by precipitating thallous ion as chloroplatinate.16 All samples were counted as solids using a G.E.C. G.M.4-type end window counter. Since samples of the same weight were always compared no absorption or scattering corrections were made. The average deviation between the measured activities of duplicate samples was always less than 1 %. All measurements were made in duplicate. THALLIC-FERROUS OXIDATION RATES All solutions were made and standardized exactly as described in ref. (12) but at room temperature. One aliquot was removed from 100ml of reaction mixture in a beaker containing the platinum metal and titrated as in ref. (12). [ClOil, M FIG. 1. 0 HC104, NaCIO4, X Ba(C104)2 k = 0.5489-0.08794 [ClO,] RESULTS The bimolecular rate constants were calcuIated in the usual way 17 and are given in fig.1 and tables 1 and 2. All measurements were made at 25°C. The line1 04 EXCHANGE OF THALLOUS A N D THALLIC IONS drawn in fig. 1 is the least-squares slope (excluding data below 1.0 and above 5.0 M perchlorate) and corresponds to the equation, k = 0.5489 - 0.0879[C10~]total. ( l a ) The corresponding equation for 80 % D20 solutions calculated from the data in tables 1 and 2 is k = 0.392 - O~0628[C10~]tota~. (1 b) TABLE 1 EI(C104)3] 040770 M, [TIClO4] 0.020 M, ionic strength 1 *20 0.336 0.388 0.377 0279 0.564 0.422 0.690 0.300 0.906 0.441 1 a00 0300 1.134 0.456 [HC10]4, M itHzo, mole-1 1. h-1 [DClO4]*, M k*,,, mole-1 1. h-1 *total D/H = 80 % throughout TI(C104)3 0.00773 M, [TIC1041 0.020 M, ionic strength 3.0 [HC104], M kHz0, mole-1 1.h-1 DClO4]*, M k*D20, mole-1 1. h-1 0.42 0268 0.40 0.90 0.271 0.70 1-37 0.271 1.03 1.92 0268 1.38 2.39 0.265 1 -75 2.93 0.268 213 [Tl(C104)3] 0.00770 M, [TIC1041 0.020 M, LHC1041, M kHzo, mole-I 1. h-1 DClO41*, M 0.365 0.0827 0.283 0.936 0.0762 0.832 2.01 0.0687 1.75 3.08 0.0558 2.66 4.5 1 0.0485 5.93 0.0439 0.204 0.202 0.198 0.1 99 0.200 0.203 ionic strength 6.0 k*,,,, mole-1 1. h-1 0.0667 0-0563 0.0502 0.0471 TABLE 2 [TlC104] = 0.020 M, solvent 100 % H20 TI(C104)3] M = 0.001345 0,00269 0.00538 0*00807 p = 1.20 HC104, M = 1.11 k = 0.474 0,473 0459 0.437 mole-1 1. h-1 k = 0.271 0.270 0.263 0.257 k = 0.275 0.280 0.273 0.265 k = 0.0907 0.0888 0.0843 0.08 17 k = 0.0627 0.0610 0.0635 0.0561 p = 3.0 p = 3.0 p = 6.00 HClO4, M = 2.499 HClO4, M = 2.910 HClO4, M = 0.473 HClO4, M = 4.28 p = 6.00 Tables 1 and 2 are for solutions at fixed ionic strengths ( p = 1-2, 3.0 and 6.0) and contain the reactants together with a large excess of sodium and hydrogenS .W . GILKS AND G. M. WAlND 105 perchlorate only. The ionic strength here is therefore effectively the perchlorate concentration. In water at p = 1.2, the rate constant increases from 0-388 to 0-456 ; at p = 3-0 there is no change ; at p = 6-0 the rate constant decreases from 0.083 to 0.044, when in each case the formal hydrogen ion concentration is increased from about 0.4 M to p. Fig. 2 shows that at p = 6.0 the rate is proportional to the molar concentrations of perchloric acid and sodium perchlorate. I I 1 1 0 I 2 3 [H+l/INa+l FIG. 2. ionic strength 6.0; 0, present work (H20); e, Harbottle and Dodson kNa+ = 0.470 kHi = 0.517 The curvature of the second-order plot for the thallic-ferrous oxidation observed by Higginson12 after about 60 % reaction and attributed by him to the back reaction FeITr+ TII1=FeI1+ T P is eliminated by platinum metal (fig.3). DISCUSSION The linear decrease in rate with increasing perchlorate concentration (fig. 1) of a reaction between two positive ions which is known to be catalyzed not only by nitrate 2 but also by sulphate ions 14 seems unusual. It has, however, been pre- viously shown that catalysis by sulphate ions can be accurately described 18 and also qualitatively understood 19 if the reaction intermediates are TI+SO2--Tl3+ and SO~-Tl+SO~-T13+SO~--. In these species the reactants niay :or may not, also be separated by water molecules and the reaction therefore [may be either the outer-sphere or the inner-sphere type of Ta~be.20~9 20b It seems likely that the nitrate catalysis is due to Tl3+NO3Tl+, as a structure in which the nitrogen must interact with the thallic ion not only explains the kinetic data but also incor- porates the explanation given by Gutowsky21 of the equal dependence of the chemical shifts of thallous and thallic nuclei on increasing ratio [nitrate]/[thallium]106 EXCHANGE OF THALLOUS AND THALLIC tONS t (min) FIG.3. - - - - no pt. 1 in. square 0.1 mm thick [FeII] 0.003884 M bright pt. fnIII] 0.005202 M [H+] 0567 M p - 0.9 with NaClO4 room temperature in concentrated nitrate solutions. Although similar n.m.r. (2OsTl resonance) studies show that concentrated thallic perchloric solutions (which cannot be prepared free from thallous ion) give a very broad absorption band ;22 the variation in position and in band width with change in perchlorate concentration and with change in thallous-thallic ratio have not yet been measured.Many properties of concentrated solutions containing tetrahedral oxy-anions have been attributed to either (i) change in solvent structure,23 or (ii) cation-water-anion complexes.24 For the solutions used here there is no information which makes it possible to distinguish between these effects; either or both of which might be expected to cause the observed linear decrease in the exchange rate with increasing perchlorate ion concentration. HYDROGEN-ION DEPENDENCE The change in rate with increasing perchloric acid concentration is very different from that of the thallic-ferrous electron transfer where there is a marked acid catalysis above 2 M.25 At constant perchlorate concentration hydrogen ion inhibition is large only in the 6 M solutions.As recent studies of the hydration of protons 26 and of triply charged cations 27 suggest that this may reflect changes in the degree of hydration as well as of the hydrolysis of the thallic ion, and as studies of the ultra-violet absorption of these solutions confirm this,28 we have only calculated specific rate constants for TW and TlOH2+ in the more dilute solutions.S. W. GILKS AND G . M. WAIND 107 In 3 M perchlorate solution, two independent determinations of the degree of hydrolysis of the thallic ion are in complete agreement.28929 The hydrolysis constant at 25°C is 0.073 and is the same in water and in mixtures of water and heavy water.28 Application of eqn.(2) and (3) to the data in tables 1 and 2 gives : where (3) w3 +I ; cIo=- [T10H2 '][H'] [TI3+] [TI"']' K = ko = 0.267, kl = 0.301 mole-1 1. h-1; and to the data of Higginson 12 for the thallic-ferrous reaction, ko = 0.44 mole-1 1. min-1, Icl = 15.1 mole-1 1. min-1. All these values differ from those quoted in the literature.30~ 31 Although the presence or absence of hydroxyl-ion catalysis can only give definite information about the composition of the transition state, i.e. it either does or does not contain a hydroxyl ion; in view of the much discussed importance of bridging groups in electron transfer reactions and the ability of the hydroxyl ion to penetrate the hydration sheaths of ions, it is interesting that while dimeric ferric hydroxyl complexes are well-known, thallic and mercuric ions are the only cations for which this behaviour has not been postulated.32, 33 D/H ISOTOPE EFFECT This is independent of ionic strength (table 1) and corresponds to ~ H ~ o / ~ D ~ o = 1.5 for 100 % deuterium oxide.The significance of such an effect has recently been discussed.*os 31~34 The importance of water as a constituent in the activated complex for electron transfer reactions has also recently been stressed.35 1 Harbottle and Dodson, J. Amer. Chem. SOC., 1951,73,2442. 2 Prestwood and Wahl, J. Amer. Chem. SOC., 1949, 71, 3137.3 Anderson and Bonner, J. Amer. Chem. SOC., 1954,76, 3826. 4 Silverman and Dodson, J. Amer. Chem. SOC., 1952, 56, 846. 5 Meier and Garner, J. Physic. Chem., 1952, 56, 852. 6 Sherril, King and Spooner, J. Amer. Chem. Soc., 1943, 65, 170. 7 Keenan, J. Amer. Chem. Soc., 1956,78,2339. 8 Furman and Garner, J. Amer. Chem. SOC., 1952,74,2333. 9 Sutcliffe and Weber, Trans. Faraday Soc., 1956,52, 1225. 10 Davidson and Dodson, J. Physic. Chem., 1952,56, 866. 11 Reynolds and Lumry, J. Chem. Physics, 1955,23,2460. 12 Ashurst and Higginson, J. Chem. Soc., 1953, 3044. 13 Andrews, J. Amer. Chem. SOC., 1903,25,766. 14 Brubaker and Mickel, J. Inurg. Nuclear Chem., 1957,4, 55. 15 S. W. Gilks, Thesis (London, 1960). 16 Challenger and Masters, J. Amev. Chem. Sac., 1956, 78, 3012. 17 McKay, Nature, 1938, 142, 997. 18 Brubaker et al., J. Amer. Chem. Soc., 1957, 79,4641. 19Libby, J. Amer. Chem. SOC., 1940, 62, 1930. 20 Henry Taube, (a) Int. Cunf: Co-ordination Chemistry (Spec. Publ. Chem SOC., London, 1959), p. 57 ; (b) Advances in Inorganic Chemistry and Radiochemistry, vol. 1 (Academic Press, 1959). 21 Gutowsky and McGary, Physic. Rev., 1953, 91, 81. 22 B. N. Figgis, private communication. 23 Robinson and Stokes, Electrolyte Sulutiuns, 2nd ed. (Butterworths, 1959), chap. 1. 24 Sykes, J. Chem. SOC., 1959, 2473. 25 Forscheimer and Epple, J. Amer. Chenl. Soc., 1952, 74, 5773.108 EXCHANGE OF THALLOUS AND TI-IALLIC IONS 26 Bell and Bascombe, Faraday SOC. Discussions, 1957, 24, 158. 27 Glueckauf, Trans. Faraday Soc., 1955, 51, 1235. 28 Waind and Rogers, to be published elsewhere. 29 Biederman, Arkiv. Kenzi., 1953, 5, 441. 30 Rossotti, J. Inorg. Ncrcl. Chem., 1955, 1, 159. 31 Basolo and Pearson, Mechanisms of Inorganic Reactions (John Wiley and Sons Inc 32 Gluskova, Russ. J. Inorg. Chem., 1959, 7. 33 Sillen, Acta Chem. Scand., 1954, 8, 299, 1607. 34 Marcus, J. Chem. Physics, 1956, 24, 966. 35 Maddock, Trans. Faraday Sac., 1959,55,1268. 1958), chap. 7.

ISSN:0366-9033    

DOI:10.1039/DF9602900102

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

THE REACTIONS OF CHROMIUM(II) AND THE ISOMERIC DIFLUOROC€€ROMIUM(rII) IONS BY YUAN-TSAN CHIA AND EDWARD L. KING Dept. of Chemistry, University of Wisconsin, Madison, Wisconsin Received 26th January, 1960 The transition state for the electron-transfer reaction of cis-difluorotetra-aquochromium (111) ion with chromium(1I) ion which involves two fluoride ions acting as bridging groups is not detected. Reaction by this pathway is very much slower than by the pathway involving a single fluoride ion acting as the bridging group. The values of AH* and AS* for the reactions of chromium(II) ion with cis-difluorotetra-aquochromium(II1) ion and with monofluoropenta-aquochromium(II1) ion are very similar. Bridged transition-states have been demonstrated to provide the predominant reaction pathway for many oxidation-reduction reactions involving metal ions.1 What of the possibility of a transition-state in which two anions act as a bridge between the reacting metal ions ? Although transition-states containing two anions are known for such reactions, e.g.the transition-states (Fe2C12+)* 2 and (Fe2F;+)+ 3 for the iron(I1)-iron(II1) exchange, it cannot be assumed, as has been suggested,4 that these reactions proceed by transition-states in which both anions act as bridging groups. The present study was designed to investigate this question for the reaction of chromium(1I) ion and cis difluorotetra-aquochromium(II1) ion ; because of the inertness of chromium(II1) species, this system is one which can be expected to provide unequivocal information.That a fluoride-bridged transition- state does provide the easiest pathway for the " electron-transfer " between chrom- ium(I1) and monofluoropenta-aquochromium(I1I) has been shown by the occurrence of chromium exchange between these two species unaccompanied by a q ~ a t i o n . ~ Chromium(I1) can be expected to catalyze the aquation of each isomer of difluoro- tetra-aquochromium(II1) ion by a path which involves a transition-state analogous to that for the exchange of chromium(I1) and monofluoropenta-aquochromium(II1) ion with a single fluoride ion acting as the bridging group, / \ + FCr(OH& '. \ / F(H20),CrF++Cr2+-+ Only the cis-isomer of difluorotetra-aquochromium(II1) ion can possibly form a transition-state with chromium(I1) ion in which two fluoride ions bridge between the two chromium atoms, and this transition-state for the exchange reaction can be studied only if it has a stability greater than or comparable to that of the monofluoro-bridged transition- state which accomplishes the net aquation of cis-difluorotetra-aquochromiurn(II1)- ion.109110 CHROMI UM(II) + DIELUOROCHROMIUM(III) REACTIONS EXP E RI MENTAL The isomeric difluorotetra-aquochromium(II1) ions, shown by analysis to have a ff uoridel chromium ratio of 2-02&0+09, have been prepared and separated from one another by an ion-exchange procedure 6 patterned after that which has proved successful for the preparation of the isomeric dithiocyanatotetra-aquochromium(II1) ions 7 and the isomeric dichlorotetra-aquochromium(II1) ions.8 As in these previously studied systems, the more easily eluted isomer is assumed to be the trans isomer, an assignment of configura- tion which is supported by the spectra of the two species." A solution of chromium(I1) perchlorate in perchloric acid was prepared electrolytically in the same manner employed previously.5 Chromium-51 was obtained from the Oak Ridge National Laboratory in the form of a very dilute solution of chromium(II1) chloride in hydrochloric acid.In the preparation of solutions containing tagged chromium(II), the radioactive chromium(II1) was either added to a perchloric acid solution containing a very large excess of chromium(II1) per- chlorate and heated prior to the electrolysis or added to the chromium(II), present in very large excess, and allowed to exchange ; 9 the chloride ion concentration of the resulting chromium(I1) solutions was less than 10-2 M.In the preparation of solutions contain- ing tagged difluorotetra-aquochromium(III) ion, the radioactive chromium(II1) was added to a large excess of chromium(II1) perchlorate prior to equilibration with fluoride ion; a comparison of the radioactivity and the difluorotetra-aquochromium(II1) ion content of the appropriate eluent portions demonstrated that the radioactivity of the eluent portions was due to chromium in the form of difluorotetra-aquochromium(II1) ion. The reaction mixture containing one of the isomeric difluorotetra-aquochromium(II1) ions and perchloric acid contained in a Pyrex glass vessel was rid of oxygen by bubbling carbon dioxide through the solution prior to the addition of chromium(I1).At various times, measured portions of reaction mixture were forced by the pressure of carbon dioxide into hydrogen peroxide solution which predominantly converts chromium(l1) into hexa- aquochromium(II1) ion.10 The difluorotetra-aquochromium(1LI) ion of charge + 1 was separated from the species of higher charge, monofluoropenta-aquochromium(II1) ion, hexa-aquochromium(II1) ion, and any polymeric chromium(II1) species by an ion exchange procedure. The separated difluorotetra-aquochromium(II1) ion was analyzed for radio- activity by use of a well-type gamma ray scintillation detector (Atomic Instrument Co., model 810) and for chromium content by a measurement of the absorbance at 372 mp, after conversion to chromate ion with alkaline hydrogen peroxide.11 In one run, the total chromium content of the difiuorotetra-aquochromium(II1) ion fractions was determined both by the chromate procedure and by a procedure using 1,5-diphenylcarbohydrazide ; 12 the results were in good agreement. In most experiments the chromium(II) Concentration of the reaction mixture was checked several times during the course of an experiment by using tri-iodide ion as the quenching agent; the excess tri-iodide ion was determined by titration with thiosulphate ion, The chromium0I) concentration decreased by no more than 2 % in the experiments in which it was followed.For experiments involving the cis isomer plots of In [CrFZ] against time were linear up to 65 %-SO % aquation and then started to level off; presumably this was due to the approach to equilibrium in the aquation reaction.The approach to bquilibrium was not studied carefully; since the reaction was carried out in glass vessels, the exact concentration of fluoride ion was not known. For the trans isomer, curvature in the plots of In [CrFtl against time occurred much earlier, and only experimental points during the first 20 %-25 % reaction were used in the evalu- ation of the rate constant. The cause of this behaviour of the trans isomer is not known. It cannot be due simply to the establishment of the aquation equilibrium; the trans isomer is the less stable isomer 6 and will, therefore, aquate more completely. The curvature is consistent with the aquation reaction being higher than first order in trans- difluorotetra-aquochromium(II1) ion, but too few experiments have been performed to confirm this possibility.(In none of the experiments at a particular temperature was the initial concentration of the trans isomer varied appreciably.) t The observed values of the second-order rate constant k2 = [Cr2+]-1 d In [CrF,+]/dt, are given in table 1. The aquation is shown by these data to be first order in the catalyst, chroniium(I1) ion. * A tabulation of the absorbance indices of the two isomeric difluorotetra-aquo- chromium(III) ions as well as monofluoropenta-aquochromium(II1) ion over the wavelength range 210 mp to 400 mp will be furnished upon request. Further experiments on the reaction of chromiuin(I1) ion and traat3-diflbiorotetra- aquochromium(ZT1) ion are in progress and will Ire reported at a later date.Y .CHIA AND E. L. KING 111 In certain experiments, chromium(i1) was tagged and the radioactivity of the separated difluorotetra-aquochromium(II1) ion was determined. The radioactivity of the samples taken at the shortest reaction times showed an apparent extent of exchange of -1 %; samples taken at a later time when 6 % to 15 % aquation had occurred showed no more exchange than those taken when 1 % to 2 % aquation had occurred. Whatever the cause of this small apparent amount of exchange, it is not due to the direct second-order exchange reaction. Considering the various experimental uncertainties, it can be stated that the second-order rate constant for the direct exchange reaction of cis-difluorotetra- aquochromium(II1) ion and chromium(I1) ion is definitely not greater than 1 % of the second-order rate constant for the chromium(I1)-catalyzed aquation of cis-difluorotetra- aquochromium(1II) ion.25" 34.6" TABLE 1.-vALUES OF THE RATE CONSTANTS FOR THE AQUATION REACTIONS cis- or ~ ~ ~ ~ s - ( H ~ O ) ~ C ~ F ~ + H ~ O + ~ ~ ( € - I ~ O ~ ) C ~ F ~ + + H F in 1 M perchloric acid k2 = [Cr2+]-1 d In [CrF$]/dt cis isomer trans isomer [Cr2+]~ 102 k 2 ( ~ 103 mole L-1 sec) [CrZ+]x 102 kz( x 103 mole 1.4 sec) temp. - 0" 2.83 1.18 1.68 020 4.1 3 1.1 1 3.29 017 calc.0 1.2 1.87 10.1 1.73 1.5 ca1c.a 10 0.95 17.4 1.53 8.9 1-85 22-5 1 *88 6 9 204 21.5 2.26 19.4 2.54 17.6 ca1c.a 19 45.5" ' 0.859 20 1-83 19 (a) calc. using the parameters given in table 2. A more accurate limit of the value of the rate constant for exchange should be derivable from experiments in which the aquation proceeds to a greater extent.In such experi- nients, of course, there will be interference from the reversibility of the aquation reaction, which provides a pathway for exchange. The reverse reaction HF+(H20),CrF2+%3++(H20)4CrF~ +H20 can be suppressed by thorium(IV) ion which forms a very stable fluoride complex13 but it was found that thorium(1V) ion also accelerates the aquation of difluorotetra-aquo- chromium(II1) ion, presumably by the direct reaction H, 0 + (M, O),CrF; + Th4+ +(H,0),CrF2 + +ThF3 +. Two experiments involving chromium(I1) and thorium(IV) ion at a concentration equal to that of the tagged difluorotetra-aquochromium(II1) ion did not reveal appreciable chromium exchange.The amount of exchange observed in experiments with trans-difluorotetra-aquo- chromium(II1) ion was also negligible. With this isomer, of course, direct exchange is not possible ; had direct exchange been observed, it would have indicated that the assign- ment of configuration to the isomers was incorrect. DISCUS SION Although the exchange of chromium atoms between chromium(I1) ion and cis-difluorotetra-aquochromium(UI) ion via the symmetrical difluoro-bridged112 c H R o M I u M (I I) + D I F L u o R o c H R o MI u M (I I I) RE A c TI o N s transition state seemed a reasonable possibility in view of the occurrence of such a four-atom-ring configuration in known compounds (e.g. the aluminium halides14~ 19, the reaction path was not detected in the present study.The transition state involving two bridging groups has also been shown to be unimportant relative to that with a single bridging group in the reactions of chromium(I1) ion with cis- diaquotetranaminecobalt(II1) ion or cis-diaquobisethylenediaminecobalt(II1) ion.16 In the absence of any direct experimental proof of the operation of transition states for metal ion oxidation-reduction reactions involving two bridging groups, and for most systems such proof seems inaccessible, it is reasonable to assume that the transition state involves, at most, one bridging group. TABLE 2.-vALUES OF THE PARAMETERS ASSOCIATED WIl'H THE RATE LAW rate = ~ [ k T / k ] exp (AS*/R) exp (- AH*/RT)[Cr2"][CrF~3-~)+l k2 at 34.6' a AS+ AH+ ( x mole 1.-1 sec) (X cal-1 mole deg.) (X kcal-1 mole) cr"* species (H20)sCrF2+ 4.9 x 10-2 -20 5 13.7 5 cis-(H20)&rF f 1.9 x 10-2 - 24 13 (a) kz = rate/[Crz+][CrFi3-')+] ; values presented are calculated using values of AS+ (b) calc.assuming K = 1.00. and AH+. A summary of the values of AH+ and AS+ for the reactions of chromium(I1) and various fluorcsaquochromium(II1) species via fluoridebridged transition-states is given in table 2. Before values such as these are compared with one another, a correction for the symmetry number factor should be made.17 The values are uncertain enough, however, that these small corrections will not be made. The similarity of the values of AH+ and AS* for the reactions of chromiun(I1) ion with cis-difluorotetra-aquochromium(II1) ion and monofluoropenta-aquochrom- ium(II1) ion is of interest for in each of these chromium(II1) species a water mole- cule is located trans to the fluoride ion which is bonded to chromium(I1) in the transition state.It is the trans group which is believed to exert the dominant influence upon the rate at which reactions of this type occur.18 We wish to thank the United States Atomic Energy Commission for financial support (Contract AT(l1-1)-64, Project No. 3) and Sister M. J. M. Woods., O.P., for help in performing the experiments. 1 Taube, Advances in Inorganic Chemistry and Radiochemistry led. Emeleus and Sharp) (Academic Press, New York), vol. 1, 1959, chap. I. 2 Silverman and Dodson, J. Physic. Chem., 1952,56, 846. 3 Hudis and Wahl, J. Amer. Chem. SOC., 1953, 75, 4153. 4 Zwolinski, Marcus and Eyring, Chem. Rev., 1955,55, 157. 5 Ball and King, J. Amer. Chem. SOC., 1958, 80, 1091. 6 Chia and King, to be published. 7 Hougen, Schug and King, J. Amer. Chem. SOC., 1957,79,519. 8 King, Woods and Gates, J. Amer. Chem. SOC., 1958, 80, 5015. 9 Anderson and Bonner, J. Amer. Chem. SOC., 1954,76, 3826. 10 Ardon and Plane, J. Amer. Chem. SOC., 1959, 81, 3197. 11 Haupt, J. Res. Nut. Bur. Stand., 1952, 48, 414. 12 Allen, Anal. Chem., 1958, 30,447. 13 Dodgen and Rollefson, J. Amer. Chem. SOC., 1949,71,2600. 14 Palmer and Elliot, J. Amer. Chem. SOC., 1938, 60, 1852. 15 Harris, Wood and Ritter, J. Amer. Chem. SOC., 1951, 73, 3151. 16 Kruse and Taube, J. Amer. Chem. Sac., 1960,82,526. 17 Benson, J. Amer. Chem. SOC., 1958,80,5151. Orgel, Report Tenth Solvay Conference (Brussels, 1956), p. 289.

ISSN:0366-9033    

DOI:10.1039/DF9602900109

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

GENERAL DISCUSStON Dr. N. S. Hush (Bristol University) said The activation energies of " outer-sphere " electron-exchange processes should depend strongly on the energy of re-orientation of the ligands in the first co-ordination shells of the ions to the transition-state configuration In the thermoneutral M3+/M2+ isotopic exchange reactions of transition-group ions in aqueous solution this re-orientation energy should vary sharply with atomic number as a result of the crystal-field stabiliz-ation of the aquo comp1exes.S In the table below some results are given of cal-culations of this term for the Fe3+/Fe2+ and V3+/V2+ exchanges assuming these to be adiabatic outer-sphere reactions. The calculated terms for the corresponding electrode processes i.e., M3++ e (electrode)+M2+ are also shown.These results have been obtained by a calculation similar to, but less approximate than that described in ref. (7). The V3+/V2+ exchange both in solution and at an electrode is predicted to have a smaller activation energy than the Fe3+/Fe2+ exchange because the re-orientation energy of the first shell of water molecules around the ions should be smaller for the vanadium reactions. This is connected with the fact that the large crystal-field stabilization energy of the V(H20)2+ ion is accompanied by a con-traction of the V . . . 0 bond distance to a value which will be very close to that in the V(H20)$+ ion so that an unusually small expansion of the (H2O)6 shell is needed to accommodate the transferring electron. The expansion is much larger for the Fe(H20)%+ ion-probably about 0.16w.In agreement with this the activation energy for the V3+/V2+ exchange at an electrode is lower than that for the Fe3+/Fe2+ exchange. For the homogeneous reaction the activation energy for the Fe3+/Fe2+ exchange is reasonably well predicted. The calculated value for the V3+/V2+ homogeneous exchange however is much lower than the value suggested by Krishnamurty and Wahl4 on the basis of an analysis of the overall V1I1/VI1 exchange rate in 1 M HC104 solution. This high activation energy is difficult to understand and it would be interesting to see whether it can be con-firmed by an independent measurement. calculated inner-shell re- total activation expt. activation orientation energy energy energy ref. system elec- Fe3+/Fe2+ 2.5 10.1 9 (1) homo- Fe3+/Fe2+ 5.0 108 9.9 (3) (energies in kcal/g ion) trode{ V3+/V2+ 1.0 7.9 '7-8 (2) geneous{ v3+/v2+ 20 7.5 13.2 (4) 1 Randles and Somerton Trans.Faraday Soc. 1952 48 937. 2 average of results obtained by Randles and independently at a mercury electrode (private communication 1960). 3 Silverman and Dodson J. Physic. Chem. 1952 56 846. 4 Krishnamurty and Wahl J. Amer. Chem. SOC. 1958,80 592 5 Orgel J. Chem. Soc. 1952,4756. 6 Hush Trans. Faraday Soc. submitted for publication. 7 Hush J. Chem. Physics 1958,28,962. 113 py Parsons and Mehl 114 GENERAL DISCUSSION Dr. A. A. VEek (Polarographic Institute Prague) (communicated) The free energy AF* of activation of a redox reaction involving electron transfer can be separated into two terms where AFZ denotes the part of the free energy of activation due to all the electro-static effects operating in the course of the transition state formation and AF: denotes the part of the free activation energy connected with the structural changes in the reacting particles accompanying the transition state formation.The value of AF $ for most reactions is smaller than the '$ electrostatic " free energy of activation and reactions of simple ions possessing a not very different structure, are usually relatively fast. On the other hand there is a large group of reactions, the free energy of activation of which is relatively high and much higher than the calculated AF $ value. In such reactions the structural changes-i.e. the changes in interatomic distances or electronic states-operate and the value of AFZ is of great importance.In these reactions however the electrostatic contribution to the free energy of activation plays also an important role. To establish the importance of AFZ in reactions accompanied with large struc-tural changes we have followed the rate of electron transfer between the mercury electrode and a series of cobaltic complexes. It is however not possible to separate the experimental free activation energy into the parts AFZ and AF but it is possible to measure the change of AFZ with the change of the environment when the assumption is made that the composition of the solution influences the value of AFF only to a very small extent. Under this assumption the observed difference in the free energy of activation [A(AF+)] of the reaction in two different solutions is equal to the difference of the " electrostatic " free energy of activation [A(AF$)].The change of AFF is caused by the change of the potentid at the electrode-solution interface and tli'e values of A(AF$) can be supposed to represent a rough measure of AF$. AF* = AFZ+AF:, TABLE 1 ion A(AFz) (kcal) [Co(N&)d3 + - 3.0 [Co(NH3)5N0212+ - 1.00 [CO(NH3)5N03I2+ - 0.90 [CO(NH~)SAC]~ + - 1-10 trans [Co(NH3)4(N0&If -0.10 [cO(NH3)3m(?2)31 w +0.1 [COCNH~)~(NO~)~I- w + 1.0 cis [Co(NH3)4(N02)21+ -0.30 The last two values are only approximate due to complications of electrode process by convection. Tables 1 and 2 summarize some characteristic values. A(AF+) is defined as where (AF+)l is measured in a solution of 0.1 M HC104 0.9 M NaC104 and (AF4)2 in 1 M NaHS04 solution.It has to be noted that all the reactions proceed on the positive side of the electrocapillary curve and that the reactions studied are highly irreversible. The results in table 1 show a general agreement with the electrostatic theory : the electrostatic energy of activation is high for the reaction of highly charged particles with the electrode carrying the same charge (repulsion) it decreases with decreasing charge of the reacting particle and becomes even negative (attraction) when the reacting particle has the opposite charge with respect to the charge of the electrode. A(AF') = (AF*)l- (AFs)2 GENERAL DISCUSSION 115 A new point of view is brought by the results summarized in table 2.The values of A(AF+) for complexes having the same total charge depend on the struc-ture of the complex i.e. on the distribution of the charge in the complex particle. The ions cannot thus be regarded as spherical conductors but the inner charge distribution has to be taken into account in calculations of AFt. TABLE 2 ion A(AF$) (kcal) [CO(NH3)6I3 ' - 3.0 [Co(NH3)5H2013+ - 1 -70 cis [Co(NH3)4(H20)2]3 + - 065 trans [Co(NH3)4(NO2)]+ -0.10 [CO(NH3)5S04]+ - 0.80 Let us suppose for simplicity that the electrostatic free energy of activation arises from the necessity to bring an ion of total charge z from the bulk of the solution where the value of electric field is conventionally set to be equal to zero, into an inhomogeneous field at the interface the strength of which at a given point is g(r).The work necessary to make such a transport is trans [Co(NH3)4(H20)# -0.55 where 60 is the potential of the field due to all charges and dipoles outside the ion considered at the centre of the ion; m, my and m are the components of the di-pole and 411 412 the components of the quadrupole of the ion. To the electro-static free energy of activation further components contribute (such as the first and second Stark effect in the inhomogeneous field the change of the polarization of the medium the van der Waals forces) which we shall not consider for the present time. TO simplify the considerations the ion can be regarded as an assembly of discrete charges ei and in such a case the work W will be given by where 4j denotes the potential of the field at the site of the charge ej due to all other charges and dipoles.Making some assumptions concerning the dependence of # j on the position of the charge el and the values of charges el a rough in-formative calculation of W can be carried out. Such calculations show that one of the most important factors is the orientation of the reacting particle with respect to the electric field which it is entering. The comparison with the experiment shows that the orientation is not always such as to make Wa minimum. This result can be interpreted on the basis of the fact that the particle is oriented in such a way as to minimize the total free energy of activation. The particle is strongly polarized in the field (for complexes of transition ions also the first Stark effect plays an important role) and the polarization of the particle tends to decrease the value of AFZ and the orientation of the particle will be such as to decrease the total activation energy as much as possible even at the expense of an increase in the electrostatic work.The detailed study of entropies and enthalpies of activation shows that the process is more complicated but the results are in agreement with the simple picture described. The conclusions obtained from the study of electrode reactions can be easily generalized for electron-transfer reactions in the solution in the considerations of which the importance of the non-uniform charge distribution in the ions and hence of the relative orientations of the two reacting particles has also to be taken into account 116 GENERAL DISCUSSION Dr.D. R. Stranks (University of Leeds) (communicated) Marcus 1 has derived expressions for the contribution to the overall free energy of activation from reorganization of the primary co-ordination spheres of reacting molecules. For many polyatomic ions in water the force constants and bond lengths not only of the transition state but also of the reactants are unknown. A rigid calculation is then virtually impossible. For electron exchange between the hexammines of Co(P1) and Co(I11) an estimate can at least be made in the following manner. The two C O N stretching force constants may be taken as k a ~ = 2x 105 dyne cm-l2 and k11 = 0 . 8 ~ 105 dyne cm-1.3 The ColII-N distance rIII is 2-05f 0.02 whereas the CoI1-N distance rI1 has been widely quoted as 2.5 A.5 The latter distance based on some exploratory X-ray measurements,6 seems abnormally large.A value of 2-39A is consistent with the occurrence of the electron-transfer band of the Co(NH,);+ ion at 2000 A ; this value will be adopted here. For thermal electron transfer the two co-ordination spheres of the reactants must first re-arrange to an intermediate bond distance r+ such that the energy required for this re-arrangement is minimal. To a first approximation this re-arrangement corresponds on the valence-force model to the symmetrical dl 3 breathing frequency of an octahedrally-co-ordinated molecule. Orgel 7 has quoted this energy as but this expression assumes that I"+ = -&II-rIII) which only holds for the special case ICII = k111.The more general expression for the re-arrangement energy E is where For electron exchange between the hexammines of Co(I1) and Co(1II) this " energy barrier '' amounts to 32 kcal mole-1. This forces the electron transfer act to pro-ceed via a bridged transition state.8 This may well be a general circumstance. Dr. N. S. Hush (Bristul Urziversity) said If an adiabatic mechanism is assumed, it is possible to discuss the activation barrier for outer-sphere electron-exchange reactions between simple ions in an ionizing solvent in terms of a simple physical model. These reactions are of the type : Ep = 3kIII(r* -r111)2 + 3k11(rII-r*)25 r* = (k&+ k$IIT-~J(k~I+ &). azi + bZ2+aZt-1+ bZ2+1. The reactant ions are assumed to come together from the initial state i to a distance cr to form a loose collision complex with the first solvation sheaths intact and in contact.This is state p just prior to electron transfer. The free energy of activ-ation for the exchange is then written as 9 In this equation the energy change in passing from state p to the transition state t is partitioned into the change of electronic binding energy of the transferring electron (Pd'Gbhdhg) and the change of environmental or solvation energy of the a . . . (e) . . . b complex ('A'GenvkonmenQl). If the probability density of the trans-ferring electron associated with the acceptor nucleus a is A+ in the transition state, AG' = 'APG+PAfGbinding+PA1G,nvi,onmental. (1) 1 Marcus this Discussion.2 Block Trans. Faraday Soc. 1959,55 867. 3 Powell and Sheppard J. Chem. SOC. 1956,3108. 4 Watanabe Atoji and Okazaki Acta. Cryst. 1950,3 405. SBrown J. Physic. Chem. 1952,56 868. 6 Biltz 2. anorg. Chem. 1927,164,246. 7 Orgel Report 10th SoZvay Con$ 1956 329. 8 Stranks this Discussion. 9Hush 2. Elektrochem. 1957 61 734 GENERAL DISCUSSION 117 and the overlap integral between the one-electron wave functions t,ba and $b is small enough to be neglected it is easily shown that where 'AfGbinding is the overall change in electronic energy (i.e. the difference of ionization potentials of the ions bzz+l and aZ1). Hence substituting in eqn. (l) the activation free energy is 1 (3) The bracketed expression on the r.h.s. of eqn. (3) can be shown21 3 to be of the form A*(l-A*)O where 8 is a function of the dielectric constant of the solvent and the ion radii.For a thermoneutral reaction ('AfG = 0) AG+ is a maximum where A* is 3 i.e. the electron charge density will be symmetrically distributed between the two nuclei in the transition state. For an endothermic transfer AG* will have a maximum at A*>+ i.e. the transition-state electron density dis-tribution will tend toward that of the product ions. The reverse will be the case for exothermic reactions. The activation free energy is connected with the overall free energy through the last term ;l*fAfG of eqn. (3). Similarly the entropy of activation will contain the term A+;AfS which connects it with the overall reaction entropy. The experimental heats and entropies of activation for thermoneutral exchange processes in aqueous solution are reasonably consistent with the values predicted from a simple interpretation of eqn.(3). This lends some support to the view that these are electron-exchange reactions of the outer-sphere type. When we examine the data for the larger class of reactions with finite overall free energy of reaction it appears that there are extremely few reactions for which both kinetic and thermodynamic data are accurately known. One exception is the PuV1/Pu3f exchange, for which the overall thermodynamics are known.4 A calculation of the rate parameters based on eqn. (3) assuming an outer-sphere mechanism gives 3 AG' = iAPG + {PA*G,n,ironmental- ;l+iAfGenvironmental) + 2"'AfG. puo; + + Pu3 + +PUOz' + Pu4 + AH+ 4.3 AS* -45 AG' 17.7.(Energies in kcal/mole ; entropy in cal/deg. mole.) The experimental values 5 are : (25°C) AH* 4.8 AS* -40.4 AG+ 16.9, in the same units. The agreement here is exceptionally good. For other reactions, no exact calculations can yet be made. The approximate correlations between AS+ and iAfS for a number of different reactions of the same charge type indicated by Dr. Higginson is qualitatively in agreement with eqn. (3) assuming an average value of A+ of about 3. Similarly the trend of AH* with 'AfH for exchange re-actions of U Np and Pu shown by Newton and Rabineau 6 also conforms quali-tatively to this model. However for the majority of these reactions the entropy of activation appears to be more positive (by perhaps 20-38 cal/deg.mole on average) than the outer-sphere model would suggest using approximate values for the overall reaction entropies. This suggests that the electrostatic entropy of repulsion which dominates the entropy of activation for thermoneutral reactions, 1 Hush 2. Elektrochem. 1957 61 734. 2 Hush J. Chem. Physics 1958 28 962. 3 Hush Trans. Faraday Soc. 1960 submitted for publication. 4 Katz and Seaborg The Chemistry of the Actinide EZements (Methuen London 1957). 5 Rabideau and Kline J. Physic. Chem. 1958 62 617. 6 Newton and Rabideau J. Physic. Chem. 1959 63 365 118 GENERAL DISCUSSION may be smaller in these exchanges. This would be so if the reactions proceeded through a more compact transition state e.g. one in which one molecule of water was squeezed out from the first hydration sheath of one reactant ion with the ions separated by only one water molecule in the transition state.The interionic repulsion entropy would be much smaller in absolute magnitude for such a com-plex owing to dielectric saturation effects which are of small importance in the more open type of outer-sphere complex. It may be that closer approach of the ions is generally necessary when the two nuclei are dissimilar as a result of small orbital overlap. Prof. R. A. Marcus (Polyteclznic Inst. of Brooklyn) (partly communicated) : The bridging group in an electron-transfer reaction may function in several ways, one of which has been treated in detail by Halpern and Orgel in their interesting paper. I should like to discuss these various ways making use of fig.1 in my paper. When the electronic interaction between the reactants is very weak (a non-adiabatic reaction therefore) the bridging group may enhance this electronic coupling. This effect increases the probability of electron transfer per passage through the intersection region in fig. 1 and is the one considered by Halpern and Orgel. It clearly has an upper limit namely when this probability becomes unity. Or if this probability were originally unity for the uncatalyzed reaction i.e. if the reaction were originally adiabatic this effect could of course not occur. However even in such cases further enhancement of electronic coupling in-creases the rate. A large electronic coupling results in a large “ splitting ” of the intersection of the potential energy curves in fig.1 in the usual way thereby lower-ing the activation energy. (In the first effect above only the frequency factor in the rate constant was affected other things being equal.) This lowering can be computed from the energy terms estimated by Halpern and Orgel when the splitting is not too large and when therefore their perturbation approach is appropriate. Again the bridging group may alter the relative heights of the two curves in fig. 1 and could increase the reaction rate if it thereby decreased the standard free energy of reaction. Evidently all of these effects should be considered in any comparison of theory with experiment. Finally in the atom-transfer mechanism some or all of these effects would occur. In fig. 1 the profile co-ordinate then largely involves motion of the atom being transferred.An application of Halpern and Orgel% considerations to this case would assume that the original electronic coupling is very weak (i.e. that there is a very small splitting in fig. 1). In other types of atom transfer the coupling is normally strong. Prof. J. Weiiss (University of Durham) said The dependence of transition probability on the square of the time (see Halpern and Orgel’s paper) is obtained if the transition takes place between discrete levels. If however transitions were taking place from a discrete ground state to an upper (narrow) continuous region, one would obtain a transition probability which is linear with time. In solution, it is likely that the levels will be broadened due to different interactions with the surrounding molecules.Would there be any difficulty with such an assumption? Prof. R. A. Marcus (Polytechnic Inst. of Brooklyn) (communicated) Halpern and Orgel have discussed the time-dependence of their electronic-transition proba-bility and have raised some questions about the significance of its form. It must also be recalled too that the role played by the atomic motion in effecting the electron transfer should also be explicitly considered in any dynamical treatment of a non-adiabatic process. Simultaneous electronic and atomic motions have been treated previously by Landau,ls 3 and by Zener 2 ~ 3 for the case in which the atomic motion may be assumed quasi-classical. While their equation 2 was 1 Landau Physik. 2. Sowjetunion 1932 2 46.2 Zener Proc. Roy. Soc. A 1932 137 696 ; cf. 1933 140 660. 3 cf. Kauzmann Quantum Chemistry (Academic Press New Yotk 1957) pp. 536-544. Landau and Lifschitz Quantum Mechanics (Pergamon Press London 1958) 9 87 GENERAL DISCUSSION 119 deduced for a one-dimensional case the results obtained by Halpern and Qrgel could be introduced into it as a first approximation. To complete this quantitative estimate for the transition probability per passage through the intersection region (cf. fig. 1 of my paper) some estimate of typical slopes of the two potential-energy curves at the intersection is needed and can be made. Mr. R. P. Be111 (Oxford University) said The mechanism discussed in Halpern and Orgel's paper involves electron conduction through the conjugated bridge system.It would be interesting to know whether the same kind of treatment can be applied to solids such as graphite or metals which have measurable macroscopic conductivities and which are known to catalyze electron transfer reactions such as the t hallous-t hallic exchange. Ips. J. Halpern and Dr. 1;. E. Orgel (Cambridge University) (communicated): In reply to Bell electron transport of the type we discuss occurs between donor impurity centres e.g. phosphorus atoms in semiconductors such as germanium even when the distance between phosphorus atoms is too large to permit direct orbital overlap. An even more closely related phenomenon is the electrical conductivity of mixed-valence metal oxides. In reply to Weiss we believe that the linear dependence of the rate on time is probably the correct one.However we do not understand the detailed dependence of the rate law on the distribution of upper energy states. Prof. R. A. Marcus (Polytechnic Inst. of Brooklyn) (communicated) Mr. Bell has raised the interesting question of whether the observed catalysis of an electron exchange between two ions by a solid such as platinum is an example of the bridge model discussed by Dr. Halpern and Dr. Orgel the solid serving as the bridge. It is known from other experimental results that another mechanism occurs when this metal behaves as an electrode in an electrochemical system. The usual data on electrochemical reaction orders imply transition states associated with separate collisions of each redox species with the electrode.Such an electrode in equilibrium with the solution is presumably the same as the solid immersed in a solution. Both are in electrochemical equilibrium with the latter. This equilibrium being a dynamic one the solid continually accepts electrons from colliding reductants and donates them in collisionally independent acts to colliding oxidants. It thereby would automatically cause isotopic exchange in the Tl+CTl+3 reaction. These remarks do not of course eliminate the solid-bridge mechanism for the solid catalyzed isotopic exchange (it would not be observable electrochemically) but do show that an alternative well-established path is available one whose applicability to isotopic exchange can perhaps be tested experiment ally. The apparent reaction order in such a catalytic mechanism would be some-what complex because of the potential difference existing between metal and solution.This potential difference established by the dynamic electron-exchange equilibrium between the redox species and the metal affects the rate constants for forward and reverse reactions in such a way as to make the corresponding rates equal at equilibrium. As Randles has shown from such considerations, certain simple-electron transfers A+ metal %B involving collisions between A and metal and separately B and metal have rates proportional to ~ C A C B when they are not diffusion controlled. Orders in more complex electron transfers could be inferred from electrochemical reaction orders and electrochemical transfer coefficients in such systems.When CA and CB are equal this electrochemical mechanism is seen to be first order in CA a result reflecting the first order nature of the individual electron transfer acts. Accordingly for simple electron transfers at least the electrochemical mechan-ism for catalysis would increase in relative importance over the homogeneous exchange with decreasing concentration since the homogeneous reaction is second order. This behaviour differentiates the electrochemical mechanism from th 120 GENERAL DISCUSSION solid-bridge one which would presumably be second order if the fraction of ad-sorbed species is small. At least for macroscopic metals the electrochemical mechanism would be expected to prevail over the solid-bridge one for I believe. nothing would be gained by a simultaneous collision of the two redox species the same environmental re-organization about each being necessary in both cases.Prof. R. A. Marcus (Polytechnic Znst. of Brooklyn) said The role of ionic size in the kinetic salt effect described by Zwickel and Taube is an interesting one. They attribute the greater effectiveness of the larger ions to the corresponding smaller re-arrangement energy required to modify the ionic atmospheres prior to electron transfer. While this effect is as they say in agreement with the ideas of Libby and of myself' about re-organization of the systems undergoing electron transfer I was rather surprised that they observed it experimentally. Some approximate unpublished calculations which I made based on eqn. (4.3.3) of my paper suggest that this effect would probably be swamped by specific effects associ-ated with the much larger salt effect of reduction of the coulombic repulsion of the reactants by the ionic atmosphere.For this reason it was of interest to search the literature to see if the effect described by Zwickel and Taube occurs in other more complex reactions. In a preliminary survey Dr. Norman Peterson of this laboratory found two examples for which suitable data were available 192 2BrO+BrO; + Br- (followed by other reactions) S 2 0 g +21-+2S84+12 (first order in I- and also in S,Oi-) (2) In both cases a greater ionic size of the counter ions in the atmosphere caused a greater increase in reaction rate as found by Zwickel and Taube in their simple electron transfer and by Sheppard and Wahl3 in theirs.Accordingly a different reason for the effect might be present. For example, the highly charged activated complex in a reaction between ions of like sign is, in a free-energy sense considerably more solvated than the pair of isolated re-actants.4 Highly solvated ions in the ionic atmosphere will compete with the activated complex for this solvation and so will be less effective in promoting the formation of this complex than will the less solvated ones. Smaller ions being more solvated than the larger ones this effect is seen to be consistent with the data. (Other specific effects would occur too,l and could be more important.) While this solvation effect and the re-organization effect are in the same direction for reactions between ions of like sign and so cannot be easily differ-entiated they are in the opposite direction for reactions between ions of opposite sign and a study of the latter would be of interest.In each case of course the effect of the different ionic media on the " standard " free energy of reaction in the prevailing electrolyte should also be discussed. Any effects on it would also affect the rate (cf. eqn. (4.3.3) and (4.4.1) of my paper. Prof. A. W. Adamson (University of Southern California) said I t should be noted that with Cr(dip)$+ there apparently exists a rapid equilibrium with Cr(1) and Cr(II1) species.5 Although in the cases discussed by the authors the possi-bility is apparently excluded by the simple second-order kinetics observed in general 1 Perlmutter-Hayman and Stein J.Physic. Chem. 1959 63 734. 2 von Kiss and Bruckner 2. Physik. Chem. 1927 128 71. 3 Sheppard and Wahl J. Amer. Chem. Soc. 1957 79 1020. 4To compute this additional free energy of solvation term we subtract from the coulombic term ele2/Du the term due to direct interaction of the charges ele2/r and so obtain -qeZ(l-l/D)/r. 5 Herzog et al. Z. anorg. Chem. 1952 267 337 ; 1958 297 323 ; Naturwiss. 1956, 43 35. ' GENERAL DISCUSSION 121 it is not unlikely that an oxidation of Cr(dip):' might proceed through the above disproportionation followed by oxidation of the Cr(1) species. Assuming that the oxidation with Co(en)i + did follow the simple bimolecular path a point of interest is the intimacy of contact required for the two ions to undergo a charge transfer.As a means of testing this question Mr. S . Spees in this laboratory carried out some experiments whereby Cr(dip) + in fairly dilute solution was oxidized by a deficiency of L-Co(en);+ and the product solution checked for optical activity. None could be observed. Unfortunately the experiment is not decisive since the failure to observe activity could have been due either to a lack of stereospecificity in the transition state for electron transfer or to too rapid a racemization of the product Cr(dip)i+. Further experiments of this type on other systems should be worth carrying out. Prof. F. S . Dainton (Leeds University) said If I understand Prof. Adamson's question correctly he asks whether the actual entity which reduces cobaltic ions might be the chromous tris bipyridyl ion.If the mechanism were in fact represented by the equations 2 Cr(bip)$ ++Cr(bip); + Cr(bip)z + the rate law would be - d[Cr(bip):+]/dT = k,k3[Co1I1][Cr(bip)$+]/{kZ[Cr(bip)~ +I+ kJ [CoIrr]). I do not of course know whether the authors specifically studied the effects on the rate of (i) added chromic tris bipyridyl ion or (ii) varying the initial cobaltic ion concentration but it does seem to me to be very unlikely that they would have failed to notice the retardation by products implied by this mechanism and which should have been readily detectable in kinetic runs carried to any substantial degree of conversion. Dr. R. J. P. Williams (Oxford University) (partly communicated) The complex Cr(bip) described by Zwickel and Taube has several interesting properties.First, its stability is unusually greater than the complex with ethylenediamine. This can be understood as the bipyridyl complex is a low-spin complex while the ethylene-diamine complex is in a high-spin state. In much the same way ferrous tris bi-pyridyl is very much more stable than ferrous tris ethylenediamine. The parallel between the ferrous and the chromous complexes is even closer for it appears that in both cases the third-step stability constant must be larger than the first two in the formation of the tris bipyridyl complexes. If this were not the case it is diffi-cult to understand how Zwickel and Taube were able to obtain such a satisfactory stability constant for the chromous complex by the method they applied (their table 1).The complex Fe(bip)$ has an intense charge-transfer band at 510mp and this intense band is absent in the ferrous complexes with a lower ratio of bi-pyridyl to iron. The absorption spectrum of Cr(bip) is also highly distinctive. No other chromous complex has such an intense band in the visible. It seems likely that this band has the same origin as that of ferrous tris bipyridyl. In the ferrous case the stability of the tris complex is so much greater than that of the bis complex that it is impossible to obtain a high concentration of the latter in the presence of the former. Even in acid solution at a pH where the mono bi-pyridyl complex of ferrous begins to form there is a measurable amount of the tris complex. Is this also true of the series of chromous complexes? If so then the catalysis of the reactions of chromous bipyridyl complexes by acid may not be due to the formation of bis or mono complexes as suggested by Zwickel and Taube, but may arise from the association of the proton with the tris complex.The complexes of other metal cations with bipyridyl have properties which are very sensitive to hydrogen ion concentration although they do not break up 122 GENERAL DISCUSSION The exchange reactions of Go(en)% + and Co(NH,); ' with the corrcsponding cobaltic complexes have about the same rate. The cobaltic ammine complex is reduced much more rapidly than the ethylenediamine complex by Cr(bip)g +. Do Zwickel and Taube eliminate the possibility that the two sets of reactions go by different outer sphere mechanisms ? For Cr(bip)z + is the reaction intermediate Cr(bip)g + H ' NIP Co(NH3) 2 + excluded ? In the comparison between the reactions of Cr(H20)g and Cr(bip)g Zwickel and Taube have noted the importance of orbital symmetry factors in electron exchange.There is another difference between the two complexes. In low-spin Cr(bip) + the bond distance Cr-N will be little different from that in the chromic complexes. In the hydrate a considerable bond-shortening is to be expected on going to the higher valence state. Prof. H. Taube (University of Chicago) (communicated) We have no direct evidence for the existence at high acid of Cr(I1)-bipyridyl complexes containing less than 3 molecules of the ligand per Cr(I1). Dr. Williams points out correctly that explanations alternative to ours can be given for the rate effect we observe at high acid.The issue could be resolved by a more complete rate study-if bi-pyridyl is actually lost as we have chosen to infer the rate should be sensitive to its concentration in the acidic solutions-but we have not done the indicated experiments. The issue is important enough to merit further investigation. We must also admit that even when Cr(bipy) + is the reactant we cannot be entirely certain of its constitution in the activated complex. The ion is very labile and although we know that bipyridyl is not lost when the complex ion is oxidized one end of the ligand may be dissociated. The ambiguity in interpretation would be reduced with Cr(o-phen) + as reactant but this ion seems to be difficult to prepare.I am inclined to believe that the formulation Cr(bipy)~+H+NH,Co(NH,)~ + for the activated complex in the range of " normal " rate behaviour is excluded by the results obtained in replacing N H 3 by ND3. ~ Thus for the enolization of ketones, the isotope effects observed kH/kD are in the range 1 of a factor of 4-7 and much the same isotope effect might be expected for dissociation of H from NH3. I agree with Dr. Williams' final comment. The question of the distortion of the reactant complexes required prior to electron transfer as well as the distribution of the electrons in the molecules is important in determining the rate of reaction. It is difficult to evaluate the relative importance of the two factors because in general the very ligands that lead to formation of the kind of complexes which do not undergo much change in size in the redox process are also those which help to distribute the electrons over the molecules.Dr. I. A. W. Shimi (A'in Shams University Cairo) and Dr. W. C. E. Higginson (Manchester University) (communicated) Some time ago we investigated the rates of oxidation of various substrates including hydroquinone S20i - Fe(CN)$ -and SnC12 by the sexadentate complex ion ColIIY- (where Y4- = ethylenediamine-tetra-acetate) and by the related quinquedentate complexes C O ~ ~ ~ Y H ~ O - and CoIWf3r2-. For a given reducing agent reactivities were in the order CoYBr2->CoYHzO- >COY-, the average of the relative rates for all reducing agents being CoYBr2- ca. 103; CoYH20- ca. 20 ; COY- 1. The potentials of solutions containing equal con-centrations of oxidizing agent and of CoDY2- were measured at 25" and pH = 4.7 against a saturated KCl calomel reference electrode with the following results : CoYBrz- 0.266 V ; COYH~O- 0-184 V ; COY- 0.161 V.We concluded that the differences in reactivity among these oxidizing agents were partly due to the differences in their electron affinity as shown by the potentials observed. In their paper Zwickel and Taube report that Cr(dipy)g+ is oxidized ca. 100 times more rapidly by Co(NH3)5H203+ than by CO(NH~)~+ and have ascribed 1 Wiberg Chem. Rev. 1955 55 713 GENERAL DISCUSSION 123 this to the easier conduction of ail electron through the ligand l i z 0 compared with NH3. It is probable that Co(NH3)$&03+ is a stronger oxidizing agent from a thermodynamic viewpoint than Co(NH,);+ and in the light of our conclusions we wonder whether this is partly the cause of the greater reactivity of the former reagent.Particularly in reactions which proceed via an outer-sphere transition complex thermodynamic factors may be of importance in comparisons of rates in systems of the same charge type. An interpretation largely in terms of ligand conductivities seems to us an unwarranted over-simplification ; indeed in outer-sphere type systems ligand conductivity may not be the dominant factor re-sponsible for differences in rate of reaction. Dr. W. C. E. Higginson (Manchester University) said We have given a tentative interpretation of the positive entropies of activation observed for several bi-molecular oxidation-reduction reactions between metal cations.There is of course difficulty in drawing conclusions about entropies of activation of reactions occurring at high ionic strengths since effects which may be expected when the ionic strength is low are likely to be opposed to some extent by ionic-atmosphere effects. Also the rather imprecise entropies of reaction quoted in table 2 hold at zero ionic strength. However the form of equation we have used to compare the entropies of reaction and activation may partly counteract this discrepancy, although the values of a are likely to lose their significance. reactants A Cr(H20)5F2++ Cr2+ Cr(NH3) 5F2+ + Cr2 + Cr(NH3)5CP++ Cr2+ Cr(NH&Br2++ Cr2+ C0(m3)58H2++ Cr2-t B Co(NH3)5H203++ Cr2+ Co(NH3) 2 + + Cr2+ Co(NH3) + + Cr (dipy) + C FeOH2++ Fez+ FeF2++Fe2+ FeCP++ Fez+ D Fe3 + + Fez + v3 + + v2+ TABLE 1 I AS* mole/l.cal mole-1 deg.-l 1.0 - 20 1.1 - 30 1.1 - 23 1.1 -33 1.2 - 18 1.2 - 52 0.4 - 30 001-02 e- 12 0.55 - 18 0.5 -21 0.55 - 24 0.55 - 25 2.0 - 25 bridged ? Y e s Yes Yes Yes Yes Yes no no ? ? ? ? ? ref. 1 2 2 2 3 334 5 5 6 7 6 6 8 If our interpretation is acceptable it leads to the conclusion that for most of the reactions cited the transition complex is extended i.e. more than one molecule of water separates the two cations. It is difficult to see how a positive entropy of reaction could greatly modify the negative entropy of activation anticipated 1 Ball and King J.Amer. Chem. SOC. 1958 80 1091. 2 Ogard and Taube J. Amer. Chem. SOC. 1958 80,1084. 3 Zwickel and Taube J. Amer. Chem. SOC. 1959,81 1288. 4 Kruse and Taube J. Amer. Chem. Soc. 1960 82 526. 5 Zwickel and Taube this Discussion. 6 Silverman and Dodson J. Physic. Chem. 1952 56 846. 7 Hudis and Wahl J. Amer. Chew. Sac. 1953,75,4153. 8 Krishnamurty and Wahl J. Amer. Chem. Soc. 1958 80 5921 124 GENERAL DISCUSSION for the formation of a compact transition complex. Thus we do not expect the relation proposed between AS and AS+- to be of general validity for reactions be-tween two ions; indeed it may well be restricted to electron-transfer reactions of the non-bridged type. If it is correct to regard reactions showing an inverse dependence upon the hydrogen-ion concentration as proceeding by a bimolecular reaction between partly-hydrolyzed ions the entropies of activation for this process are almost always negative.Here much more compact possibly bridged transition com-plexes seem probable. In support of this contention Ball and King 1 have noted that the values of AS+ for certain reactions involving iron(II1) and iron(I1) are similar to those for reactions of the same charge type between' chromium(II1) and chromium(I1); compare groups C and A in the annexed table. On this basis they have suggested that the former group of reactions may also proceed via bridged transition complexes. On the other hand a similar comparison between the reactions in group D with those in group B of this table suggests that the transition complexes for the Fe3++ Fe2+ exchange and the V3++ V2+ exchange are probably of the non-bridged (extended) type.Unfortunately the experimental evidence is scanty although the difference in AS+ for the reactions Co(NH&H203 + + Cr2 + and Co(NH& + + Cr2 + is notable. Prof. R. A. Marcus (Polytechnic Inst. of Brooklyn) said As Higginson et al. point out one would expect the entropy of activation between ions of similar sign to depend largely on two factors the standard entropy of reaction and some entropy change present in an isotopic exchange reaction (for which AS" is zero) of the same charge type. It is of interest to see what form the equations of my electron-transfer theory take for this entropy of activation AS*. We consider here only the cases where there is no change in the number of particles in the redox step.Differentiating eqn. (4.2.3) (4.3.3) and (4.4.1) for AF? AF$ and rn with respect to Tin my paper we obtain dwP d W - -rnAS"'+rn-(l+rn)-aAF* A s * = - -dT dT dT' on neglecting a minor term involving dA/dTand assuming that the fractional change in any force constant Ks is small (in which case AS?% -mAS&). AS"' dwp/dT and dwfdT denote the " standard " entropy of reaction of the elementary redox step the temperature dependence of the work required to bring the products to-gether from iniinity and the corresponding term for the reactants all terms being evaluated at the prevailing electrolyte concentration. In this equation m is a factor reflecting the nature of the activated complex.As may be shown from eqn. (3.1.3) or (3.3.3) m is zero when the activated complex resembles the reactants (which would be the case for reactions having a very negative AF"). It is - 1 when the complex resembles the products (i.e. when AF" is very positive). It is - 112 when the complex is as much like the reactants as it is like the products (e.g. in an isotopic exchange reaction where AF" = 0). Finally, it takes on intermediate values for intermediate situations. Bearing this significance of rn in mind the dependence of AS+ on the various quantities in the above equation becomes clearer. Comparing this equation with that proposed by these authors, we see that their a becomes -m here (whose theoretical value is now given by eqn. (4.4.1)). Their AS& becomes a related term but one in which both the charges of the products and those of the reactants appear weighted by factors -(l+m) and m factors which are equal only for the symmetrical case (AFO' = 0 etc.).Only when the charges of the two reactants are the same as those of the products 1 Ball and King J. Amer. Chem. Suc. 1958 80 1091 GENERAL DISCUSSION 125 (but interchanged) is there no ambiguity in these authors' definition of '' isotopic exchange reactions of the same charge type ". For this case wp = w and the expression for AS+ becomes aAS"'-dw/dT. It is easily shown that -dwldT is the entropy of activation of this isotropic exchange reaction of the same charge type and we then obtain their expression. Above all we note that AP' refers to the given elementary redox step restricted here to one which involves no change in number of particles, and not to the AS"' of the overall reaction.Prof. F. S . Dainton (Leeds University) said The activation entropies for almost all isotopic exchange reactions between cations are negative presumably because the charge density and solvation of the transition state exceeds the sum of the corresponding quantities for the reactants. This is clearly evident for ferrous-ferric exchange where the bridging ligand is a halogen or pseudo-halogen (see table). bridging ligand F- CI- Br- OH- -SCN N3 AS* (cal mole-1 deg.-1) -21 - 24 - 25 -18 -27 1-68 The azide ion is a notably exceptional bridging ligand in that AS+ is p0sitive.l TWO factors may account for this. First the azide ion is linear and both terminal nitrogen atoms are negatively charged so that the formation of the transition state may be regarded as the approach to the ferrous ion of a negatively charged nitrogen atom.The polarity and solvation of the system consequently diminish as the transition [(EP~O)SF~~~~N N N-]2++ Fe(H20)6-+[(H20)5Fe N N N Fe(QH2)5I4++ Hzo state is formed. Indeed it may be that as written in the equation a water molecule is displaced from the ferrous hydration sphere. That the solvent is somehow involved in activated complex formation is indicated by the fact that between 4" and 17°C the exchange rate is 50 % larger in H20 than D2Q. The second factor may be that the azide ion is an especially efficient bridge possibly by virtue of its highly symmetrical structure and the possibility of overlap of its TC orbitals with the d orbitals of the iron atoms.At present we have no experimental means of discovering the relative importance of these two factors. Prof. J. Webs (University of Durham) said The entropies of activation have been derived by Higginson from the second-order constants. To what extent can it be excluded that these second-order constants are not in fact rather more complex expressions of several rate constants due to a pre-established equilibrium SO that these second-order constants would more nearly correspond to a bi-molecular rate constant multiplied by the equilibrium constant of the pre-established equilibrium. Under such conditions the entropies of the activation calculated froin the second-order constants would of course not themselves be of significance for the rate-determining process.Prof. A. W. Adamson (University of Southern California) said Mercury (11) will slowly oxidize Fe(I1) according to the equation and in some unpublished work Mr. B. J. Wood in this department found that the forward reaction was second order in 1 M perchloric acid with specific rate constants of 1-22 x 10-4 M-1 min-1 and 2.25 x 10-4 M-1 mh-1 at 80" and 9O"C, respectively. The corresponding activation quantities are 15.5 kcal for AE* and - 42 cal/mole deg. also for AS". The observed rate law does not permit a choice between two alternative mechanisms : 1 D. Bunn F. S . Dainton and S . Duckworth to be published 126 GENERAL DISCUSSION (1) Fe2++ I-Ig2+%Fe3++ Hgf (slow), 2 Hg++Hg2+ (fast).(A) Fez+ + Hg+-tFe3+ + Mg" (fast) (B) Hg" + Hgz+-+Hg;+ (fast). (a (2) First step same as above followed by Judging froin their paper Higginson et al. would prefer mechanism (2) since nowhere in their paper do they appear to consider the possibility of reaction (A) being important in their systems. For example from what detail is given for the Ag(I) catalyzed oxidation of [Hg(I)]2 by Ce(IV) it would appear that a clear-cut determination was not made as to whether the concentration of Hg(I) in the denominator of eqn. (3) is present as the first power (as shown) or as the square root. In the latter case the mechanism would presumably be : Ce(IV) + Ag(I) 3 Ce(II1) + Ag(II), k - 8 K [Hg(I)12 + 2 HgQ) (fast), Ag(II)SHg(I) 5 Ag(I)+Hg(II).I would think that equilibrium (A) should be quite rapid and that it would be reasonable to suppose that it is present during the reaction in the two systems discussed above. Returning to the reaction between Fe(I1) and Hg(1I) it is interesting that if the authors' eqn. (6) is applied using a probably reasonable value of + 15 cal/mole deg. for the entropy of Hg+ then an a value of unity results. Thus in spite of the large negative entropy of activation observed the reaction does fit in well with the systems summarized in table 2. Dr. J. B. Stead and Dr. W. C. E. Higginson (Mmchester University) said: In reply to Prof. Adamson eqn. (3) of our paper can be rewritten in the form - {d log, [Ce"]/dt}-' = A[Ce"']~~(]Hg'),~ +B, where A and B are constants provided [Ag'Iinitial is constant.If the concentration of (Hg1)2 in the denominator of eqn. (3) is present as the square-root the equation corresponding to that given above is similar but with 2/[(Hg1)2] present in the denominator. We plotted 10g&e~~] against time for each of three kinetic experi-ments in which the initial concentration of AgI was the same. From each of the resulting graphs we obtained several values of - d log&eIV]/dt by drawing gradients. The reciprocals of these values were then plotted against (i) corres-ponding values of [CeI1*]/[Hg1)2] (ii) corresponding values of [CeI*I]/d [(Hg1)21. In case (i) the points from the three experiments lay on the same straight line; in case (ii) the points lay on three separate curves. The highest concentration oi (Hg1)2 was eight times that of the lowest.The reaction between CoI'I and (Hg1)2 in our opinion a more likely system for the occurrence of the type of mechanism suggested by Prof. Adamson is also first order in (Hg1)2. We can offer no suggestion concerning the relative proba-bility of the mechanisms he has given for the reaction between Wgrr and Fe". We presume that the rate of this reaction has been shown to be independent of the hydrogen-ion concentration otherwise the value of AS* quoted would not be that of the bimolecular reaction between Hg2+ and Fe2+. Prof. Weiss has asked whether the second-order constants we have quoted can be regarded as true constants or whether they may include equilibrium constants for pre-established equilibria. The only type of pre-equilibrium we have no GENERAL DISCUSSION 127 taken into account is association between perchlorate ions and the reactant cations, and there is independent evidence for certain cations that such association does occur.To this extent many of the velocity constants we have quoted may be complex quantities. For our purposes it is therefore necessary to enquire whether the corresponding entropies of activation are greatly different from the true en-tropies of activation which would hold at zero or very low concentrations of perchlorate ions. The only convincing answer would be based on experiment and unfortunately data at sufficiently low concentrations of perchlorate ions for satisfactory extrapolation to zero concentration are lacking. We know of only one reaction of this type for which AS* has been quoted at zero ionic strength, this is the TW+ T1+ exchange reaction referred to by Dr.Waind in this Discussion. The effect of changes of perchlorate ion concentration upon equilibria in which two positive ions associate seems relevant to this question in that the kinetic and equilibrium processes should be similar as far as interaction with solvent species is concerned. Here relatively small changes in the entropy of association are usually found e.g.,l FeOH2++H++Fe3'+H,8; AS (cal/rnole deg) = - 25 ( I = 0) - 21.6 ( I = 1-0). AS (cal/mole deg.) = - 18 (I = O) - 14 (I = 1.0). 2FeOH2' +(Fe2(QH),)4+ ; Thus there are some experimental indications that association between perchlorate ions and metal cations is unlikely to modify greatly the true entropy of activation for bimolecular reactions between metal cations.This does not seem very surpris-ing for it is probable that such association does not involve penetration of the first co-ordination sphere of the cation at any rate in solutions of perchlorate up to ca. 5 M. In this connection Sykes concluded that in dilute solutions two water molecules separate Fe3f and ClO in the association complex.2 Further the effects of such a loose association are likely to be similar for reactions of the same charge type Our use of the apparent entropies of activation for comparative purposes therefore seems admissible. We accept however that uncertainty about the effects of perchlorate ions remains and that this is a weak point in the hypothesis advanced in our paper.Finally we refer to a major difference in interpretation between Dr. Mush and ourselves. In his second contribution to the Discussion Hush regards abnormally positive values of the entropy of activation in bimolecular electron-transfer re-actions between aquated cations in which the nuclei are dissimilar as evidence in favour of a compact transition state whereas we believe that the transition state is extended. It does not seem easy to decide between these opposing views. Reactions between ions of like charge in which compact transition states are known to be formed almost invariably show highly negative values of AS+ (see e.g., ref. (20) of our paper and table 1 of our second contribution to the Discussion). Although these examples do not include electron-transfer reactions between two aquated cations we find it difficult to suppose that if a compact transition state occurs in such a reaction there would be a substantially different (less negative) contribution to AS+ compared with these other reactions on account of a much smaller increase in solvent orientation when the transition state is formed.Dr. Gwyneth M. Waind (Queen Mary College Londun) said The data at 25°C (given in Halpern and Orgel's paper) for the Tl+-Tl3+ exchange reaction can be combined with those at 50°C 3 to give the following values for the Arrhenius 1 Milburn J. Anzer. Chem. Soc. 1957 79 537. 2 Sykes J. Chem. Soc. 1959,2473. 3 nodson J. Artier. Chem. Soc. 1953,75 1795 128 GENERAL DISCUSSION energy of activation Ea and for AS+.The bimolecular rate constants were ob-tained from those measured in mixtures of sodium perchlorate+ perchloric acic by linear extrapolation to all sodium perchlorate. Ic mole-1 1. h-1 25°C 50°C Ea kcal mole-1 AS+ cal mole-1 deg.-1 NaC104 hl 6.0 0.0885 0.78 1 16.6 - 26.5 3.0 0.270 2.174 16.0 -266 1.2 0.385 2.788 15.3 - 28.2 0 0459 3-45 14.5 - 30f 1) Higginson and co-workers (this Discussion table 1) report entropies of activ-ation for reactions between metal cations in " dilute perchloric acid solutions " which range from 0.55 to 4.5 M. The above figures show that for this reaction at least such large changes in perchlorate concentration lead only to small changes in AS*. Dr. D. R. Rossdnsky (University of the Witwatersrand South Africa) (com-municated) Part of the contribution to the positive entropies of activation in reactions involving Coly reported by Higginson and collaborators may arise from the preliminary equilibrium if C0~~I(d,)4(d,)2 is more reactive than Co111(d,)6.Magnetic susceptibility measure-ments show that at least a small fraction of Cot? exists in the high-spin state.1 A change of radius of about 0.1 to 0.2A would be expected causing a positive entropy change of 10 to 20 cal/mole deg. according to the Latimer-Powell entropy expression. This point could be cleared up by measurements on the system FeII where partial hydrolysis in the transition state is improbable. The rate is rather high for conventional measurements &>lo4 M-1 min-1 in 6 M and 3 M perchloric acid at - 7"C).2 Prof.J. Weiss (University of Durham) said I am not aware that Shaffer's principle of two-electron equivalence change has any physical basis. In solutions, most electron transfer reactions are adequately described by one-electron transfer processes. Two-electron transfers are of course not excluded and experimental evidence for such processes would be of considerable interest. I am referring to electron transfer reactions as such which would normally proceed as non-adiabatic processes. obtained with concentrations of cobaltic perchlorate < 3 x 10-3 M for the existence of dimeric Co(III) and the spectrophotometric data would suggest that a large proportion of the Co(II1) exists in this form. I would like to ask Dr. Higginson whether he has considered the possibility that dimeric Co(II1) is one of the reacting species in the systems he considers? Dr.L. H. Sutcliffe (Liverpool University) said Recently Dr. J. R. Weber and I 4 have attempted to study the formation of polynuclear species of Co(II1) in perchloric acid solutions using a spectrophotometric method. Polynuclear forms do seem to be present and they are favoured by high Co(II1) concentra-tions and low acidities. Baxendale and Wells 5 have postulated that dimeric Co(II1) is a reactive entity in the decomposition of cobaltic perchlorate by water. The experimental conditions they used were such that a high proportion of polymeric species would be expected from our findings. Dr. C. F. Wells (British Rayun Res. Assoc.) said There is kinetic evidence 1 Friedman Hunt Plane and Taube J.Amer. Chem. Soc. 1951,73,4028. ' 2 Rosseinsky Thesis (Manchester University 1958). 1 3 Baxendale and Wells Trans. Faraday Soc. 1957,53 800. j 4 Sutcliffe and Weber J. Inorg. Nucl. Chern. 1960,12,281. 5 Baxendale and Wells Trans. Faraday SOC. 1957,53 800 GENERAL DISCUSSION 129 Mr. P. J. Proll and Dr. L. H. Sutcliffe (Liverpool University) (communicated): Incorporation of the equilibrium Co(I1) + Co(II)+Co(I) + Co(III) suggested by Longuet-Higgins in place of the dimerization equilibrium does not lead to a rate law corresponding to that observed. Several deficiencies become apparent : (i) the order with respect to Pb(IV) would be predicted to vary whereas it was found to be unity under all conditions ; (ii) either the reaction would not go to completion or added Co(II1) would retard the reaction neither of these effects was observed ; (iii) the highest power predicted for the dependence on the concentration of cobaltous acetate would be 2 but a value as high as 2-3 was obtained.This fact led to the postulation of the dimeric form of Co(I1). Prof. R A. Marcus (Polytechnic Inst. of Brooklyn) said In his treatment of electron transfer one of Laidler’s1 assumptions was that the reaction rate equals the rate of diffusion (as computed by Debye 2) of the reactants towards each other under the influence of their interionic forces multiplied by some factor which includes an electron tunnelling term. This ad hoc assumption is it will be shown below incorrect. The correct multiplicative factor will include a diffusion con-stant which in Laidler’s case cancels the one in his diffusion rate.The final expression does not depend on the diffusion constant when the reaction is inefficient. Examples where the role of diffusion in bimolecular reaction rates has been incorrectly treated in force-free cases have been discussed recently.3 Using related arguments we deduce here an expression applicable to systems where inter-molecular forces influence the diffusion. The reaction scheme normally consists of a diffusion step (1) to form some A . . . B pair which in turn can either diffuse apart (reverse process in step 1) or react : Let c(r) distance r. k i A+B+A . . . B (1) (2) k - 1 A . . . B+products. k2 denote the concentration of B with respect to some A at an A-B (Ref.(3) gives a precise interpretation of c.) The net flux of diffusion of B towards an A J(R) at r = R mustequal the reaction rate, where c(R) is the concentration of A . . . B’s and R is a typical A . . . B distance. The diffusion equation to be solved is J(r) = 4nr2[D(dc/dr) + vc] (4) where v the relative velocity caused by the forces equals a velocity per unit force, u multiplied by the force dw/dr w(r) being the work required to bring A and B from infinity to r. D is the sum4 of the diffusion constants of A and B. The term u is readily shown to be D/kT. (At equilibrium J = 0 and c = co exp (- w(r)/kT) co being the value of c at r = 00. Substituting these results this value of u follows at once.) One has therefore dc Dcdw J(r) = 4nr2 D-+- .( dr k T d r ) (5) 1 Laidler Can. J. Chem. 1959 37 138. 2 Debye Trans. Electrochem. SOC. 1942 82 265. 3 Collins and Kimball Ind. Eng. Chem. 1949 41 2651. 4 Smoluchowski 2. physik. Chem. 1917 92 129 130 GENERAL DlSCUSSION Tn the steady-state J(r) is independent of r. Solving (5) subject to boundary condition (3) and setting c = co at r = 00 we find for the overall rate constant, k which is the flux per unit CO J/CO : 1 exp w(R)/kT . 1 + - . . - = k k2 4nD/I exp (w]kT)dr]r2 r = R Substituting Laidler’s numbers into (6) one finds that the second term in the r.h.s. is negligible. Or more simply a comparison of the known bimolecular rate constant k with the diffusion term in (6) leads to the same conclusion. Thus when a reaction is sufficiently inefficient its rate constant does not depend on D.An alternative method of deducing (6) which is perhaps somewhat more illus-trative of the physical processes involved and which shows more clearly the relatioiiship with Debye’s equation for a purely diffusion-controlled reaction rate, is the following. In (1) and (2) we can set dcA . . B / d t = 0 in the steady state and obtain for the overall rate constant kl is the rate constant computed by Debye by integrating (5) subject to the Smoluchowski boundary condition,l c = 0 at r = R. (This boundary condition prevents back diffusion.) kl is the flux per unit co and was found in this way to be2 kl = 4nD/j * exp (w/icT)dr/r2. (8) R We see incidentally from (7) that k indeed equals Debye’s term kl multiplied by some efficiency factor but not the one assumed 3 previously.Similarly k-1 is the flux when c(R) = 1 and c(c0) = 0. We find upon in-tegrating (3, co k = 4 n ~ exp [ w ( ~ ) ] k ~ ] / J exp (w/kT)dr/r2. R (9) Combining (7) (8) and (9) (6) is once again obtained. In (6) incidentally it is convenient to define a rate constant k which would be the value of k if c(R) always had its equilibrium (Boltzmann) value which it does when k2<k+ This value of ke it is seen from (6) or (7) equals k2 exp [- w(R)/kT]. When k is very small relative to the diffusion term we see that from (6) that k = k and the reaction rate is “ activation controlled ”. When the converse is true k equals Debye’s expression 4nD exp(w/kir) r/r2d and the reaction is diffusion controlled. Prof E.L. King (University of Wisconsin) said In the paper by Dr. Stranks, a value of the second-order rate constant for the reaction of Co(NH,);+ and the ion-pair Co(NH3);’ . C1’ is presented. To obtain this from the experimental data it is necessary to use a value of the equilibrium quotient for the reaction ic Co(NH3); + + C1’ = Co(NH3);+ . C1’. 1 Smoluchowski 2. physilc. Chenz. 1917 92 129. 2 Debye Trans. Elektrochem. Soc. 1942 82 265. 3 ‘L.nidler Con. J . C’hem. 1959 37 138 GENERAL DISCUSSION 132 In view of the large uncertainty in the value of this quantity,l it is preferable to summarize the experimental results in the third-order rate law : rate = k[Co(NH3)2 +][Cl-][Co(NH& 'I, the magnitude of the rate constant for which comes directly from the experimental data and does not depend upon a somewhat arbitrarily selected value of the equilibrium quotient for the association reaction.One cannot dismiss the amido species Co(NH3),NHi' as the participant in the exchange reaction pathway showing the first-order hydroxide-ion dependence on the basis that the value of K1 (eqn. (3.1) in Dr. Stranks' paper) obtained in the analysis of the kinetic data agrees with the value obtained in equilibrium measure-ments It has been suggested that the low rate of hydrogen exchange between solvent water and hexamminecobalt(II1) ion in acidic solution indicates the acidity of hexamminecobalt(II1) ion is due to the formation of the ion-pair CO(NH~)~+ . OH- and not to the formation of the amido species CO(NH~)~NH%+.~ This suggestion appears to be unjustified.In alkaline solution the exchange rate is high and thus does not shed light upon relative concentrations of C O ( N H ~ ) ~ N H ~ + and Co(NH3)g* . OH- and although the exchange rate is low in acidic solution, this observation is consistent with the rapid establishment of the equilibrium, Co(NH,);+ + OH-$Co(NH,),NH$ .+ + HzO, for the rate of establishment of equilibrium at constant pH is determined by the pseudo-first order rate constant (kl[OH-]+k2) but the rate of exchange is deter-mined by the pseudo-first order rate constant (kl[OH-]). If the equilibrium re-action given above is the mechanism for the exchange reaction the values of kl and k2 may be obtained from the rate of exchange 2 and the magnitude of the acid dissociation constant of hexamminecobalt(II1) i0n.3 At pH values less than -12 Ic2 >kl [OH-] and thus the rate of establishment of equilibrium is independent of pH and in addition is quite high tgg4x 10-5 sec.It must be concluded, therefore that the question of whether it is Co(NPI3);' . OH' or CO(NH~)~NH$+ which is incorporated into the transition state for the exchange pathway catalyzed by hydroxide ion is not settled ; in view of the high rate at which these two cobalt-(111) species may come to equilibrium the question will probably remain unsettled. Dr D. R. Stranks (University of Leeds) (conzmunicated) Prof. King has drawn attention to the question of the reaction of the ion-pair Co(NH3)2+Cl-and the large uncertainty attached to its association constant K,. We have found that the following analysis of the kinetic results largely overcomes this admitted uncertainty in &.We first measured the electron-exchange rates in pure per-chlorate media (rate &) in which kinetically-significant association with per-chlorate is assumed to be absent. Then a series of measurements was made in solutions of steadily increasing chloride ion concentrations (rate R for complete association of Co(NH&+ with Cl-). In these chloride solutions a fraction a of the Co(NH3)g+ will exist as Co(NH3)2+C1- and the observed exchange rate, Robs will be given by k k2 Robs = aR + (1. - a)R,. A graphical plot of (Robs- Rp) against a should be linear with a slope of (&- Rp). Since the value of a is quite uncertain we have made two estimates of a ; one 1s calculated from a value of K,<O.2 (35" I == 0.9) reported by King and his co-workers 4 and the other from an earlier value of K = 91 (35" I = 0-054) reported 1 King Espenson and Visco J.Physic. Chem. 1959 63 755. 2 Block and Gold J. Chern. Soe. 1959 966. 3 Caton and Prue J. Chem. Soc. 1956 671. 4 King Espenson and Visca J. Physic. Clletn. 1959 63 755 132 GENERAL DISCUSSION by Evans and Nancollas.1 Our own spectrophotometric measurements suggest K ~ 0 * 3 5 (35" I = 0.9) whilst extrapolation of a conductimetric result 2 suggests that ICa?4 (25" I = 0.9). We have taken King's K value as a lower limit and Evans' value as an upper limit in calculating a. In both cases the graphical plot of (Xobs-Rp) against a is linear within experimental error. This plot then yields a second-order rate constant for exchange via the Co(NH,);+Cl- ion pair of 5.1 x 10-2 1.mole-1 min-1 (lower limit of K,) and 3 . 7 ~ 10-2 1. mole-1 min-1 (upper limit of KO). I have adopted a mean value of 4 . 4 ~ 10-2 1. mole-1 min-1 in my paper. Despite the large uncertainty in Ka and a the value of the rate constant seems reasonably sure. Strictly speaking one may only summarize the experi-mental results in the third-order rate law proposed by Prof. King. However I think it is reasonable to attribute the kinetic significance of the term in chloride ion to reaction via the Co(NH3)g+Cl- ion-pair since although the chloride ion cannot penetrate the first co-ordination sphere of Co(NH3);+ it must be within the next one or two co-ordination spheres to be able to affect the electron-transfer rate.That is it must be as closeas thechloride ion is in the ion-pair Co(NH3)2+C1-. Prof. King has raised the second point concerning the possible participation of the amido species Co(NH3) ,NH$+ in the electron-exchange reaction involving Co(NH3);'. Block and Gold 3 accepted that rate-determining dissociation to the amido species is the route to deuterium exchange but suggested that the Co(NM3);+0H- ion-pair may be an effective intermediate. The same species has been invoked by Caton and Prue4 who have assigned to it an association constant of 72.7 (25" I = 0). Whilst accepting Prof. King's view that the rate of formation of the amido species would be quite rapid even in acid solutions it should be pointed out that in our electron-exchange studies these hydrolytic pre-equilibria are not rate-determining and in interpreting the exchange path involving the first-order hydroxide dependence we must consider the equilibrium concentra-tions of either Co(NH3)2+OH- or CO(NH,)~NH~+.The former species may be detected quite readily in the electron-transfer band of the Co(NH,);+ ion in basic solutions 5 but no spectral evidence for the amido species is found in the d-d bonds. It would seem then that the equilibrium concentration of the amido species is considerably less that of the hydroxide ion-pair in basic solutions. For the pre-equilibrium involving Co(NH,)g+OH- the value of KI defined in my paper is found from our kinetic analysis to be 4 x 10-12 (64.5") and 5x 10-11 (80.1"). Values calculated from the work of Caton and Prue are 7x 10-14 (29 4~ 10-12 (64.5) and 2x 10-11 (80.1") all for an ionic strength of unity.Within the large experimental error (a factor of two to five arising in the extrapolation of our results) the two sets of values are equal. Since our preliminary spectrophoto-metric studies on the electron-transfer band of the hydroxide ion-pair confirm the reaction velocity work of Caton and Prue the balance of evidence seems to favour the ion-pair rather than the amido species as the effective exchange intermediate. Dr. R. J. P. Williams (Oxford University) said In the discussion of their results Morris Basolo and Pearson use an argument based on crystal field theory. They say since ammonia has a stronger crystal field than chloride the dz2 orbital will not lose electrons as easily in a cis dichloro as in a trans dichloro isomer.This statement poses several questions. How do the authors measure crystal field strength? If they use the spectrochemical series then ignoring the fact that this series does not coincide with any thermodynamic series of stabilities in platinum chemistry do the authors suppose that kinetic observations can be correlated with the spectrochemical series? It seems that in exchange reactions a series closer to the nephelauxetic series holds. There would also appear to be some difficulty 1 Evans and Nancollas Trans. Faraday SOC. 1953 49 363. 2 Jenkins and Monk J. Chem. Soc. 1951 68. 3 Block and Gold J. Chem. Soc. 1959 966. 4 Caton and Prue J. Chem. Soc. 1956 671. 5 Pearson and Basolo J.Anier. Chenz. SOC. 1956 78 4878 GENERAL DISCUSSION 133 about the cis/trans effect. Is it not conceivable that in those cases where rate is controlled by the conductivity of ligands which I presume to be unrelated to ligand field effects as measured by the spectrochemical series the trans isomer may give a lower conduction level than the cis isomer i.e. the grouping Cl-Pt-C1-Pt conducts electrons more readily than NH3-Pt-Cl-Pt. The symmetry of the d orbitals could then be a subsidiary consideration. Prof. Ralph G. Pearson and Dr. F. Basolo (North Western Urziversity Ill.) said In Dr. Williams' remarks he appears to be thinking of planar divalent platinum complexes where it is true a crystal field series from spectra is not well established nor correlated with kinetic or thermodynamic properties.However, we were discussing the situation in an octahedral complex of tetravalent platinum. Here we made the implicit assumption that the spectrochemical series is the one usually given. For octahedral complexes there is ample evidence that kinetic observations can be correlated with spectrochemical measurements. It should also be noted from our paper that we reject the crystal field explanation for the exchange reaction. Mr. A. J. Poe (Imperial College) said Although as Basolo Morris and Pearson have pointed out in their paper there is evidence for the addition of a fifth group to some square planar complexes no such evidence seems to have been reported for Pt(1I) complexes in aqueous solution. However in a spectrophotometric examination of the relative stabilities of the various trans-dihalogeno-bisethylene-diamine Pt(1V) ions evidence for the ion-pairs Pten2Br$+ Br- and Pten,I$+ I' was obtained.1 The absorption spectrum of the trans-Pten2€€ali + ion changes considerably on addition of excess of Hal- when the halogen is bromide or iodide, but not when it is chloride.The changes can be accounted for if ion-pairing occurs to an extent governed by the formation constants G0.1 1. mole+ and -2 1. mole-* for the bromide and iodide systems respectively at an ionic strength of 0-5 M and at a room temperature of -20". A redox equilibrium such as Pten,Hal~++HaI-+Pten~* +Hal, can be rejected because of the nature of the variation of the absorbance with the concentration of free halogen ion.No similar evidence was observed for ion pairing between Ptenif and Cl' Br- or I-. The formation of ion-pairs involving Pt(1V) species suggests that in exchange or replacement reactions of the bromo-or iodo-complexes an alternative mechanism involving only the free halogen ion and the Pt(1V) complex could compete with or even replace the redox mechanism postulated for the chIoride exchange. Dr. N. S. Hush (BristoZ University) said The TP+/Tl+ isotopic exchange reaction is particularly interesting as it is one of the few exchanges between fairly simple ions involving overall transfer of two electrons. There is no direct evidence that the reaction in aqueous solution involves TP+ as intermediate although this is often assumed. However Vetter 2 has recently reported that the corresponding exchange at a platinum electrode proceeds by way of two consecutive one-electron steps with T12+ as intermediate.As there is a general parallelism between mechan-isms of solution-phase transfers and electron-transfer at electrodes this suggests that the solution-phase transfer probably also proceeds by two one-electron steps. TI+ ~13++2~12+ (1 1 (in aqueous solution) is not yet known so that no direct calculations of reaction rates can be attempted. It is interesting however to work in reverse and see what value this constant must have in order to obtain a calculated rate in agree-ment with experiment assuming two consecutive electron-transfer steps with the outer-sphere type of mechanism. The theory mentioned earlier in the discussion The equilibrium constant for the dismutation 1 A.J. Po& unpublished work. 2 Vetter 1960 private communication 134 GENERAL DISCUSSION can be fitted both to the heat and free energy of activation reported by Prestwood and Wahl 1 if the free energy of dismutation reaction (1) is - 11.5 kcal/inole.2 As the Tl3+/Tll+ potential is 1.25 V,3 this would put the TP+/Tl+ and Tl3+/Tlz+ potentials at 1-50 V and 1-00 V respectively. The calculation assumes spherical symmetry of the first hydration sheaths of all ions involved which may not be altogether satisfactory but is unlikely to influence the derived dismutation energy very much. This value is consistent with the electrochemical data ; 4 it is however, quite provisional. Dr. M. H. Ford-Smith Dr.N. Sutin and Dr. R. W. Dodson (Brookhaven National Laboratory) (communicated) Additional measurements on the rate of oxidation of T11 by Br2 have been made by a spectrophotometric method. Pre-liminary results are cited here. Continuous flow and stopped flow procedures were employed. The reaction was followed by the disappearance of bromine or by the appearance of the thallic bromide complex. The figure shows the second-order specific rate calculated on the basis of TI+ and 331-2 as a function of initial [Br-] from zero to 8 x 10-3 My at 25"C I = 0.5. In the experiments at [Br-] = 0 the initial Tlf and Br2 concen-trations ranged from 4x 10-4 to 3x 10-3 and 2x 10-5 to 1 . 5 ~ 10-3 respectively. In the experiments with added bromide the initial Tl+ and Br2 concentrations were 4X 10-4 and 2x 10-4 respectively.Under the conditions employed there is definitely a linear dependence of rate on Br-Iy as required by the interpretation advanced for the exchange results. (Br) M x 10 The redox equilibrium calculations have been re-examinedy in view of the uncertainties involved in extrapolating the Tl+-Tl3* potential and the thallic bromide complexity constants to zero ionic strength. It was concluded that more reliable estimates could be obtained by using (i) the directly measured Tl+-Tl3+ 1 Prestwood and Wahl J. Amer. Chem. SOC. 1949 71 3137. 2 Hush Trans. Faraday Suc. 1960 submitted for publication. 3 Lather Oxidation Potentials (Prentice-Hall New York 2nd ed. 1952). 4 Vetter 1960 private communication GENERAL DISCUSSION 139 potential in 0.5 M perchloric acid; (ii) the standard Br-Br2 potential corrected to I = 0.5 (the activity coefficient of Br- was taken as 06S) and (iii) the formation constants of TlBr2+ TlBr; and TlBr3 corrected to I = 0-5 from Benoit's data at I = 0-12.On this basis the equilibrium constants for the reactions TlBr; = Tl++ $1-2 and TlBr3 = Tl++Br2+Br- are 1 . 6 ~ 10-10 aud 1 . 6 ~ 10-149 respectively The predicted reverse specific rate constants are 2.1 x 107 M-1 h-1 and 1.2 x 1011 M-2 h-1 at 25°C. The corresponding values from the spectrophotometric studies are 0.7 x 107 M-1 h-1 and 1 . 4 ~ 1011 M-2 h-1. The agreement supports themechanisms proposed for the first-order terms in the exchange rate law. The discrepant results of the e.m.f. rate measurements are probably due to failure of the simple assumption that the potential of the platinum electrode would be controlled essentially by the bromide-bromine (rather than the thallous-thallic) redox system.DF. W C. E. Higginson (Manchester University) said The overall reaction TllI1+ 2FeI1-+Tl1+ 2Fe111 is believed to occur according to the mechanism : TlIII+ FeII+TlII+ FeIII Tlrl+ Felll-+Tllll+ FeII TllI+ F&I-+Tll+ FelI1. When the concentration of FeIII is low the reaction is first order in both Tlul and FeII but when FeIU accumulates sufficiently the second-order plot shows curvature owing to the back-reaction (-1). In experiments with no FelI1 present initially the gradient of the linear part of the second-order plot has been shown to give the rate constant of reaction (1).1 Gilks and Waind have now demonstrated that in the presence of a platinum surface the reaction shows simple second-order kinetics up to at least ca.80 % of completion. The gradient of the corresponding second-order plot is identical with that of the initial linear part of a plot obtained from a similar experiment in the absence of platinum. This shows that there is no catalysis of reaction (1) and hence that there can be no catalysis of reaction (- 1). Evidently only reaction (2) is catalyzed. The potential attained by the platinum catalyst in this system seems to be of major importance in interpreting these observations. We suggest that the platinum assumes a potential approximately midway between the potentials of the TllI1/Tll and FelI1/FelI couples i.e.ca. 1.0 V From the instability of T1I1 we conclude that the potentials of the thallium system are TllI1/TllI< 1-1 V TllI/Tll> 1.4 V. Thus a T P species striking the platinum is much more likely to be reduced than oxidized since the negative change in free energy for reduction is considerably greater than for oxidation. This argument should hold provided that the potential attained by the platinum does not exceed that of the TllI1/Tll couple in the system and this could be checked by experiment. Prof. C. Baughan (Shrivenham) said I would like to make a comment on the interesting results of Gilks and Waind on catalysis by platinum metal. If this effect is due to " free " electrons in the metal as seem likely there may be here an important general method for testing for free electrons in solids and there are many biologically important systems for which such a test would be useful.Dr. Gwyneth M. Waind (Queen Mary College London) said The experimental data available for the Tl+-Tl3+ exchange reaction prior to those reported by Halpern and Orgel were summarized by Dodson 2 who concluded at that time that (i) the thallic ion was largely hydrolyzed (ii) the exchange rate decreases with increasing ionic strength. We have now found that (i) not only is the thallic ion 1 Ashurst and Higginson J. Chem. Soc. 1953 3044. 2Dodson J. Amer. Chem. SOC. 1953,175,1795 136 GENERAL DISCUSSION slightly hydrolyzed but also that changes in the degree of hydrolysis are relatively unimportant as far as the exchange rate is concerned? (ii) the rate decreases with perchlorate concentration. We also suggest that our results show clearly (i) that the uptake of water is necessary for the formation of the transition state. On the simple transition-state treatment we may write the exchange reaction as Tl3&".+ Tli-q,+ aq.=w If the activity coefficient ratio reflects only changes in hydration? i.e. the removal of free water as water of hydration on the formation of the transition state, then eqn. (1) becomes rate = k[al[b] a$ (2) where 6 is the difference between the hydration number of the transition state and the reactants and a the activity of the water in the perchlorate solutions. For the perchloric acid solutions used by us a is available and the data of fig. 1 give a linear log log plot : loglo k = 1*5+ 3-48 loglo a,. (ii) The existence and size of the D/H isotope effect shows that in this solvated complex hydrogen atoms must move. Consideration of all the exchange rate constants for this reaction in 80 % water and heavy water which have been obtained in this laboratory and are listed below, suggests that the slight variations with ionic strength in the size of the kinetic isotope effect are not larger than experimental error. We do not know the exchange rate in solvent mixtures of other compositions but we do know that the change in the composition of thallic perchlorate solutions (as measured by the ultra-violet absorption) is a linear function of the volume fraction of heavy water. We have, therefore converted our data to 100 % D20 and report the kinetic isotope effect as about 1.5. measured k[H+l perchlorate concentration [H+] formal dependence of rate -k[D+1 1.2 1.0 low chloride 1-372 1.2 0.4 to 1.2 acid 1 *418 3.0 0-4 to 3.0 acid 1.338 3-0 1.0 low chloride 1.357 6 0 0.4 to 60 acid 1-306 In reply to Prof. Baughan graphite catalyzes this reaction? boron nitride does not

ISSN:0366-9033    

DOI:10.1039/DF9602900113

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

II(a). OXIDATION-REDUCTION BEACTTONS INVOLVING INORGANIC SUBSTRATES AQUEOUS CHEMISTRY OF INORGANIC F E E RADICALS PART 3.-".HE KINETICS AND MECHANISM OF THE REACTION OF PEROXYDISULPHATE ION AND HYDROGEN PEROXIDE BY MAAK-SANG TSAO AND W. K. WILMARTH Dept. of Chemistry, University of Southern California, Los Angeles 7 Received, 20th January, 1960 A kinetic study has been made of the reaction of peroxydisulphate ion and hydrogen peroxide in aqueous solution at 30°C. In three different ranges of concentration of the reactants the rate law approaches three different limiting forms : (i) first order in peroxy- disulphate and half order in hydrogen peroxide, (ii) first order in peroxydisulphate and (iii) half order in peroxydisulphate. The proposed mechanism involves a chain reaction, with the SO,-, OH, and H02 radicals carrying and terminating the reaction chain.The present paper is concerned with the kinetics of the reaction of persulphate ion and hydrogen peroxide, with the results yielding detailed information about the reaction mechanism and about various aspects of the chemistry of the H02, SO4 and OH radicals. When hydrogen peroxide, hereafter referred to as peroxide, reacts with per- sulphate ion under the conditions outlined below, the stoichiometry of the re- action does not deviate appreciably from that given by eqn. (I) : S 2 0 g - + H202+ O2 +2HSOq. (1) However, it can be shown that all of the statements and conclusions presented below would be equally valid under conditions where the persulphate-induced decomposition of hydrogen peroxide was quite appreciable.2H202+H20+ 0 2 . (2) Experiments carried out under these latter conditions will be reported in a later paper of this series. EXPERIMENTAL To obtain results reproducible to within the value of 510 %, the limit which seems to have been achieved in the present study, it was necessary to take extreme precautions to eliminate impurities from the chemical reagents and from the solvent water. The presence of impurities in a particular kinetic experiment could generally be detected by the ob- servation of a long induction period, after which the reaction increased markedly in rate. In some experiments the induction period was found to persist for a time as long as 3-6 h. After the induction period was over, the rate of the reaction usually returned to the ex- pected value, indicating the complete conversion of the impurities to kinetically inactive species.Organic impurities, likely contaminants, would presumably be oxidized to carbon dioxide and water. Ordinary distilled water was purified by two successive redistillations, with the steam from each distillation being transported by an oxygen carrier gas through a vycor tube 137138 REACTION OF PEROXYDISULPHATE ION AND HYDROGEN PEROXIDE enclosed in a furnace maintained at 800"-900"C. The first redistillation was from an alkaline persulphate medium, the second from an acidic persulphate medium, with the latter solution being approximately 0.1 M in sulphuric acid. Solid chemicals including K2S208, NaC104, and KH2P04 were recrystallized at least twice from the purified water, with the mother liquor being removed by washing the crystals with the purified water.Baker Analyzed 30 % hydrogen peroxide, free from inhibitor, was used without further treatment after the following results indicated that purification was unnecessary. Use of unpurified Mallinckrodt 30 % hydrogen peroxide produced a 30 % decrease in rate, an observation which indicated the presence of an impurity. After a single steam-distillation of the Mallinckrodt product, the steam being generated from purified water, the observed rate was in agreement with that obtained when using the Baker product obtained from any of several bottles bearing different lot numbers. This latter rate was also observed when using the Baker hydrogen peroxide which had been steam-distilled. In an early study using impure sulphuric acid at a concentration of 0.30 M, [H202] at 00045 M, and [S208-] at 0.025 M, it was found that after the induction period was complete, the rate of reaction was the same as that observed in the absence of the acid.This observation provided a means of evaluating the purity of an acid, the criterion of purity being that upon addition of the acid there should be no change in reaction rate from that observed in an unbuffered solution, with the reactants at the concentrations listed above. The purification of Baker and Adamson concentrated perchloric acid re- quired three successive vacuum distillations at about 5mm pressure in an apparatus containing a minimum number of ground joints, since the ground joints could be lubricated only with the perchloric acid itself.In each of three distillations only the middle portion of the distillate was retained. An early attempt to purify Baker and Adamson concentrated sulphuric acid by vacuum distillation was apparently successful, the rate being normal and free from induction periods in experiments where this acid was used. However, more recently, all efforts to reproduce this result have failed, with inhibition being observed even after several successive distillations. Preheating the acid for 12 h at elevated tem- peratures either with or without added potassium persulphate did not increase the efficiency of the distillation process. Attempts to prepare sulphuric acid by dissolving solid sodium sulphate in perchloric acid were also unsuccessful, presumably because of impurities present in the sodium sulphate which are inactive in unbuffered solutions at pH2, but which become active in 0-30N acid. Fortunately, the experiments requiring the presence of both hydrogen ion and sulphate ion could be carried out, since it was finally discovered that pure potassium bisulphate solutions could be prepared by heating solutions con- taining potassium persulphate and hydrogen peroxide, the stoichiometry being given by eqn.(I). The kinetic experiments were carried out in glass reaction vessels maintained at 30'. At suitable time intervals, 3.0 ml aliquots of the solution were removed for analysis. The persulphate was determined iodometrically, after the peroxide had been destroyed by addition of alkaline bromine water, followed by re-acidification and removal of the excess bromine by bubbling nitrogen through the solution for 10min or less, depending upon the concentration. Concentrations of peroxide as large as 1.0 M could be destroyed in a period of time not exceeding 10 min, with about half of this interval occurring before the re-acidification.Before proceeding with the iodometric analysis, the pH of the solution was adjusted to 5.5-7.0 using phosphate buffers. The peroxide analysis was ceratimetric, the persulphate being removed by passing the solutions through a column packed with activated alumina,l followed by elution with 1.0 N perchloric acid. It has been reported in the literature that alumina catalyzed the decomposition of hydrogen peroxide,2 but no such difficulty has been encountered in our work using Harshaw Catalyst Grade alumina.Preliminary experiments using solutions of known concentration demonstrated the validity of both of the above analytical procedures. Numerical values of -d[S*Og-]/dt were obtained from the analytical data in two different ways. In early work when the form of the rate law was still largely unknown, -d[S2Og-]/dt was evaluated from the initial slope of a plot of [S20;-] against time. This method had the possible disadvantage of emphasizing the data obtained in the initial stages of the experiment, a period when trace amounts of impurities would create a maximum perturbation. In later work it became clear that under some experimental conditions the rate law assumed the rather simple limiting forms discussed below. Under these conditions all of the analytical data acquired during an experiment could be utilized, the procedure being to evaluate the rate constant from a linear plot of the appropriate function obtained from the integrated form of the rate2aw.M.TSAO AND W. K . WILMARTH 139 During the course of our work many experiments were carried out with impure re- agents, the impurities being detected by abnormally slow rates, usually accompanied by induction periods. These experiments involved either unpurified reagents or reagents which had been accidentally contaminated. All data obtained in experiments of this sort were discarded, and the results presented below are, to the best of our knowledge, free criticism.I 6 1 I I I 1 1 I I 3 I \ 0.10 M K.,S.,Oo C L O 0 0 0 U 0 O'OZ5 K2S208 n L'." a--9 Q 0 I 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 [H202]*, moles) I.-) FIG. 1.-The rate of disappearance of persulphate ion under various experimental conditions. 0 0.001 N HCIO4, 0 0.001 N HC104-l-0.225 M NaC104, A 0.30 N H2S04, x 0.125 N HC104, 0 pH = 3, KH2P04+HC104, p 0.20. RESULTS The results to be presented in fig. 1 and table 1 will be discussed in terms of the empirical rate law given by eqn. (3), the time here being measured in seconds : 5' (3) - - [s20;-1 1 9.5 x lo6 7.6 x 107[S20i-] f7.9 x lo8 +2*9 x 10'o[S20i-] CH2021 Our choice of this particular equation was based in part upon theory to be pre- sented below, and in part upon kinetic studies at concentrations of the reactants where, in turn, each of the terms in the denominator on the right-hand side of eqn.(3) played a readily identifiable role in governing the rate of reaction. For- tunately, the relative values of the numerical constants differ considerably, so that this latter approach may be carried through successfully. In fact, under the appropriate experimental conditions either the first, the third, or the fourth dc- nominator term becomes,so predominant that it determines the form of the rate law, the overall order of the reaction under these conditions being three-halves, one, and one-haIf, respectively. In early work carried out at 0.001-0,005 M peroxide, the reaction was found to have an overall order of three-halves, being first order in persulphate ion and half order in peroxide.The results of experiments carried out under these140 REACTION OF PEROXYDISULPHATE ION AND HYDROGEN PEROXIDE conditions are assembled in table 1. In the first seventeen experiments the average deviation from constancy of the quantity (- d[S20~-]/dt([S20~-][H20216) tabulated in column four is within the limit of reproducibility of &lo %. These results also demonstrate that the same rate is observed in unbuffered solutions, in 0.3 N sulphuric acid, in 0401 and 0.125 N perchloric acid, and in phosphate buffers at pH 3. In this region of limiting behaviour, the form of the rate law indicates that the denominator term 9 . 5 ~ 106/[H202] in eqn. (3) is by far the most important one, especially at the lower peroxide concentrations. At peroxide concentrations above 0-005 My conditions employed in expt.18-31, deviations from the three- halves order rate law become very appreciable, with the rate being considerably less than that which would be predicted from the first seventeen experiments. A further consideration of these data will be undertaken after examining the results presented in fig. 1. TABLE 1.-RATE OF REACTION OF PERSULPHATE ION AND HYDROGEN PEROXIDE UNDER VARIOUS CONDITIONS expt. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 0*00104 0*00098 0*00104 0.001 11 0.001 16 0.00 122 04016 0001 6 0.0020 0.00295 0.00304 0.003 1 8 0.00332 0.00377 0.00396 0.00427 0.0048 0.0086 0.0095 0.0091 0.010 0.025 0.0238 0.025 0.0238 0.025 0.025 0.00095 0.00295 000825 0.0272 0.005 12 0.010 0-025 0.025 0.025 0.025 0.025 0.025 0.025 0.0020 0.025 0.025 0.003 1 8 0.010 0-025 0.025 0.00427 0-025 0.025 0.010 0.025 0.025 0~0010 0*0020 0.0020 0.010 0025 0.025 010 0.10 0.10 0.10 0.25 2.93 2.55 2.73 2 70 2.92 2-20 2.35 2-35 2.97 2-45 2.23 2.97 2.77 2-80 2.55 2.73 2-50 1-50 2.33 2.37 2.00 2-13 2-00 1.68 1.50 1.85 1.68 2.20 1 -78 1 -08 *75 -93 (cxpt.) MIS% 9.5 20.0 22.0 22.5 24.8 19.3 23.5 23.3 2.7 33.3 30.8 5.3 1.6 43.0 40.3 7.7 43.3 34.7 22.7 56.6 50.0 3.4 6.1 5.3 23.3 73.3 66.6 67.6 97.4 97.4 125.0 167.0 (calc.) M/sec 9.5 21-7 22-3 23.0 23.3 23.8 26.8 26.8 2.7 34.0 34.3 5.0 1.5 37.0 37.5 7.5 40.0 46.6 21.5 47.3 4 8-3 2.9 5.6 5.7 25.6 56.2 56.2 67.6 102.0 131-0 151.0 21 3.0 0.30 N H2SO4 phate buffer 0.30 N H2SO4 0-30 N H2SO4 phate buffer 0- 125 i’HC104 pH = 3, phos- phate buffer 0.001 N HClO4 0.125 N HClO4 pH = 3, phos- phate buffer 0.30 N H2SO4 pH = 3, phos- phate buffer 0.30 N H2S04 unbuffered unbuffered 0.30 N H2so4 pH = 3, phos- phate buffer 0.001 N HC104 pH = 3, phos- 0.30 N H2S04 pH = 3, phos- 0.30 N H2S04 pH = 3, phos- phate buffer unbuffered 0.001 N HC104 0.001 N HC104 0.001 N HCI04 0.001 N HC104 0.001 N HC104 pH = 3, phos- pH = 3, phos- ,? o m N ~ ~ 1 0 phate buffer phate bufferM .TSAO AND W. K . WILMARTH 141 In extending our work to peroxide concentrations beyond the range covered in table 1, it was found that above 0.025 M peroxide the reaction became zero order in peroxide, with this limiting behaviour continuing to at least 1.0 M per- oxide, the maximum concentration employed.The zero-order dependence is illustrated by the horizontal lines in fig. 1, a plot of - d[S2Og-]/dt against [H202]). Only about half of the data of table 1 is included in fig. 1, since all of these points fall in a small area near the origin. TABLE 2.-RATE OF THE REACTION IN THE PRESENCE OF ADDED SULPHURlC ACID AND PERCHLORIC ACID - d[SzO;-l 10, -- dt [SzOi-l, M WzOz1, M [H+][SO:-]* [HC1041, M 0.10 -10 010 -10 -10 *lo -10 so25 ,025 -025 -025 a025 *025 -025 -075 ,050 408 5 -10 010 -025 -025 -025 ,025 *025 1.0 1.0 -74 -54 -56 -30 ,156 1.0 1.0 *95 -55 0156 -156 -096 el96 -156 -1 56 -70 -1 56 a 5 0 -20 ~156 -156 1.0 O W 0 -0052 -01 74 -01 85 ~0052 4052 moo52 *OOOO 4052 -0052 -0052 -0052 -0052 -0052 -0052 -0052 4052 -001 3 -001 3 * m 3 4 -00034 -00034 -00034 40055 M/sec (MIC.) 16.4 15.1 13.0 11-9 14.9 13.8 12 6.3 6.1 6.1 5.9 4.9 4.9 4.3 10.1 7.86 2.1 -125 16.1 -125 15.0 -125 6.4 ,125 6.4 a125 6.3 -71 6.2 4 25 6 3 M/sec (expt.) 17.3 17.3 16-7 12.3 14-7 13.3 11.7 6.7 4.8 5.3 4.1 4.8 4.7 4.2 10.0 7.9 2.3 14.7 15.3 4.5 4.5 4.8 5.0 4.7 * calculated using an assumed value of 0.10 for the dissociation constant of bisulphate ion.In the region above 0.025 M peroxide where the reaction is zero order in this reactant, two limiting rate laws may be observed, one at high and one at low persulphate concentrations. At relatively low persulphate concentration the re- action is first order in persulphate ion, a fact which may be deduced from the spacing of the two lower horizontal lines in fig. 1. In two comparable experiments, both at 0.25 M peroxide, a five-fold change in persulphate concentration from 0-002 M to 0.01 M produced an approximate five-fold change in rate, the numerical values of - d[S2O$-]/dt being 5.9 x 10-8 mole 1.-1 sec-1 and 2.9 x 10-7 mole 1.-1 sec-1 respectively.In this interval the reaction is first order in persulphate ion. In the other extreme at high persulphate concentration, an approximate half-order de- pendence of persulphate ion is observed. At 0.10 and 0.25 M persulphate, with the peroxide now at 1.0 M, the numerical values of -d[S2O$-]/dt are 1 . 7 ~ 10-6 and 2.8 x 10-6 mole 1.-1 sec-1, rcspectively.1 4 2 REACTION OF PEROXYDISULPHATE ION AND HYDROGEN PEROXIDE In preliminary attempts at curve fitting using eqn. (3), the denominator term 7.6 x 107 [S20i-’J/[H202] was omitted.The numerical values 7.9 x 108 and 2.9 x 1010 were obtained from the intercept and slope of the best straight line drawn through the points in a plot of ([S208-]/(- d[S2Oi-]/dt))2 against [S20$-], using only data at high peroxide concentrations where the first denominator term may be neglected. The constant 9-5 x 106 in the first denominator term was then evalu- ated using the rate data from expt. 9 of table 1 and eqn. (3), again neglecting the term in [S20$-]/[H202], since it contributes to a negligible extent at this relatively small persulphate conccntration. The three constants given above are adequate except when the ratio [S2Q;-]/[H202] is very large as it is in the last five experiments in table 1. These experiments were used to evaluate the numerical constant 7 .6 ~ 107, with this term never contributing more than 40 % to the total value of the denominator of the right side of eqn. (3). In contrast to the other three con- stants which are fixed to perhaps 10 or 15 %, the value 7 . 6 ~ 107 is rather poorly defined, since it influences the predicted rate in a very limited concentration region. The agreement between theory and experiment at low peroxide concentrations may be examincd by comparing columns 5 and 6 in table 1. At higher peroxide concentrations the solid curves in fig. 1 represent the variation of -d[S~Oi-]/dt with [H202]) predicted by eqn. (1). At 0.002, 0.01 and 0.025 M persulphate, no point deviates from the appropriate theoretical line by more than A 1 0 %, the limit of reproducibility of the experiments.At 0.10 M persulphate, and to an even greater extent at 0.25 M persulphate, the points do deviate very appreciably from the theoretical curves with the discrepancies apparently decreasing with an increase in the ratio [H2021/[S20;-]. To test whether the discrepancy observed at the higher persulphate concentra- tions might arise from the increase in acidity occurring during the course of the reaction, experiments were carried out in solutions containing added sulphuric or perchloric acid, the conditions differing from those of table 1 in that the peroxide concentration was large enough so that thc reaction was zero order in peroxide. The results presented in table 2 demonstrate that both sulphuric acid and perchloric acid do decrease the rate under these conditions.However, the effect does not appear to be large enough to cxplain the discrepancies mentioned above, assuming that the decrease in rate at high persulphate concentration might have resulted from sulphuric acid generated during the course of the reaction. DISCUSSION In discussing a reaction as complex as the present one it is convenient to outline the mechanism in detail before considering alternative formulations. In the following mechanistic interpretation, our initial and primary concerns will be with the data of table 1 and fig. 1 ; only a brief consideration of the more complex results of table 2 will be presented at the end of this paper. Eqn. (4)-(8) would seem adequate to represent the overall mechanism, except for the chain termination reactions presented below : s, 0; - 2 2s0, (4)M.'I'SAO AND W. K. WILMARTH 143 Chain initiation is caused by reaction 4, the thermal decomposition of persulphate ion. At 30°, kl has the numerical value 5.3 x 10-8 sec-1.3 Reactions 5, 6, 7 and 8 are the chain-carrying steps. Hydrogen peroxide and bisulphate ion compete for the hydroxyl radical, the competitive processes being reaction 6 and the re- verse of reaction 5. In deriving the rate law, the approximation will be made that kg[H202]~k4[H+][SO~-], with the validity of the approximation to be re- considered later in the discussion. In principle, persulphate ion and peroxide compete for HO2 in reactions 7 and 8, but eqn. (8) could be omitted here since it is unimportant in the present studies where the stoichiometry is that given by eqn.(1). Of the six possible bimolecular termination reactions which H02, SO;, and OH might undergo, only reactions 9, 10, 11 and 12 appear to be of importance in the range of experimental conditions which we have covered. HO, + OH3 0, + H2Q SO, + OH~HSO; (10) HO, + SO, 2 0, + HSO, so, +so,%,o;-. (12) (9 (11) Both here and above, the rate constants over the arrows refer to disappearance of reactants ; to illustrate this for reaction 12, - d[SOJdt = /cd[SOJ2. Photolytic studies reported earlier indicate that in the presence of radicals, the HSOj pro- duced in reaction 10 would decompose rapidly to form oxygen and bisulphate i0n.4 The rate law describing the disappearance of persulphate ion may now be readily derived by making the usual steady-state approximation and neglecting the rate of chain initiation and termination steps in comparison to the rate of chain- carrying reactions, a valid approximation when the chain lengths are as large as they are in the present system : By comparing eqn.(3) and (13) it is possible to identify the quantities which have been evaluated by use of the experimental data. As the terms in the denominator of eqn. (13) indicate, the relative importance of the various termination reactions depends both upon the numerical values of the termination rate constants k,, kb, k, and kd, and upon the chain-carrying reaction constants k3, k6 and k7. By utilizing the numerical constants of eqn. (3) it can be shown that (k,kd)/(kbk,) = 9.2, a result comparable with the value of unity if reactions 9, 10, 11 and 12 all proceeded with the same efficiency.Evidently the observed termination reactions do not all proceed with equal efficiency, but the result just presented suggests that they do not differ greatly, especially when it is recognized that the quantity 9.2 might actually be reduced somewhat SihCe it depends on the poorly defined constant kb. Other results of interest include : k,/k7 = (0.0136kd)/k, = (0.12kb)/k, k3/k6 = (1.3 x 10-3kd)/kb = (1.2 x 10-2k,)/ku and144 REACTION OF PEROXYDISULPHATE ION AND HYDROGEN PEROXIDE From these latter equations it seems reasonably safe to conclude that k6>k7>k3, and that the differences here play a major and probably a dominant role in deter- mining the observed form of the rate law.The conclusions just presented imply that the concentrations of the SO, and H02 radicals are relatively large compared to the concentration of the OH radicals, and thus favour termination reactions involving the HO2 and SO;; radicals. It is of interest to inquire why reactions 14 and 15 are not observed under our experimental conditions : 20H--%H202, (14) 2 H O 2 ~ ~ O 2 + H , O 2 . (1 5 ) The equations and ~OHI/CH02l = (k,[S20; --1)/(k6[H2021) [O~lICSOil= k3/(k6[H2O2l) indicate that the OH radical should predominate at sufficiently large values of [S20$-]/[H202] and 1/[H202], a situation which should cause reaction 14 to com- pete with reactions 9 and 10 and in the limit to yield a rate law of the form of eqn. (16) : This behaviour is clearly not observed in our experiments.Further, the agreement between calculated and observed rates presented in table 1 indicates that the term k,[S20~-]/2klk~[H202]2, omitted from eqn. (1 3), never contributes more than perhaps 15 % to the denominator of that equation. The result of this assumed limit of detection of reaction 14 leads to the relationships that and In view of the discussion above, it is entirely possible that the inequality k6>k7>k3 may involve factors as large as 17 and 130. In short, our data do not provide any very firm basis for assuming that reaction 14 is an inherently inefficient process. By contrast, similar considerations suggest that reaction 15 should have been detected, if kf were not abnormally small. The inclusion of reaction 15 in the mechanism would have resulted in a term kf/2klk$[SzO$] inside the brackets in the denominator of eqn.(13). Reaction 15 should make its maximum contribution to termination at large values of the ratio [H202]/[S20;-], the conditions corres- ponding to the lower right corner of fig. 1 where reaction 11 is the important termination process. Proceeding as before, and again assuming that a 15 % discrepancy in rate would have been detected, the restriction which is obtained is of the form kc/kf2850 k3/k7. While the considerations above do suggest that k7>k3, it is very doubtful that the factor is as’large as 850. To devise a rate law of the desired form it was necessary to assume that neither reaction 17 or 18 played an important part in the mechanism. S ~ O ; - + H ~ O ~ - F H S O ~ +SO,- +HO,, (17) SO, +H2O,+HSO, +HO,.(18) If the assumption is made that chain initiation arises from reaction (17) instead of reaction 4, 161 should be replaced by I G ~ [ H ~ O ~ ] in eqn. (13) and in the termsM. TSAO AND W. K . WILMARTH 145 k,[S20~-]/2klk~[H202]2 and kf/2klk?[S20;-] omitted from eqn. (1 3). The resulting rate law would be entirely incompatible with experiment in that there are no limiting forms which are zero order in peroxide. Replacing reaction 6 by reaction 18 would also eliminate the limiting forms zero order in peroxide and half or first order in persulphate, the foremost characteristic of the present kinetic system. It is now necessary to reconsider the validity of the approximation k6[H202]> k4[H+][SO$-], especially in the presence of added sulphuric acid, the condition em- ployed in theexperiments of table 2.If the restriction under consideration werelifted, the constants k,, kb, k, and kdof eqn. (13) would be replaced by ka, Akb, Ak, and Azkd, where A- 1 = (k4[H+][SO$-])/[k6[H202]). It is important to note here that the constant ka remains unchanged, a result which implies that in the limiting three- halves order region the rate should be acid-independent, as observed. An estim- ation of the magnitude of k41k6 can be made from the result k4/k6 = 12.2kckb/kake derivable from the equations k3/kg = (1.2 x 10-2kc)/ka and k3/k4 = 9.8 x 10-4ke/kb, the latter relationship resulting from data obtained in our earlier study of the photolytic decomposition of persulphate i0n.4 The calculated values of the rate listed in column 4 of table 2 were obtained by assuming k4/k6 = 12.3, the numerical value being selected to give the best agreement between experiment and theory. The agreement is entirely satisfactory in those experiments with added sulphuric acid. Further, the constant 12.3 is small enough so that eqn. ( 1 3 ) may be used without appreciable error in any consideration of the data in table 1 or fig. 1. Unfortunately, despite this apparent success, it must be recognized that the cor- rection terms in A may have only qualitative theoretical significance, since the theory does not provide an explanation for the observed decrease in rate in per- chloric acid solution. The calculated values here, based upon a sulphate ion concentration equal to that generated at tlls, are well outside the limit of error of the experiments. To explain these results in the perchloric acid medium it seems plausible to assume that the acid-catalyzed decomposition of persulphate ion, which becomes important at this acidity,3 generates a reactive intermediate which is an efficient radical scavenger. Our unpublished kinetic studies of the hydrogen- ation of persulphate ion tend to support this conclusion. It is a pleasure to acknowledge the fmancial support for the present work provided by the Atomic Energy Commission and travel funds received from the National Science Foundation. 1 Fritz, Yamamura and Richard, Ind. Eng. Chem. (And.), 1957,29, 158. 2 Dainton and Rowbottom, Trans. Faraduy Soc., 1953,49, 1160. 3 Kolthoff and Miller, J. Amer. Chem. SOC., 1951, 73, 3055. 4 Tsao and Wilmarth, J. Physic. Chem., 1959, 63, 346.

ISSN:0366-9033    

DOI:10.1039/DF9602900137

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

ON THE INDUCED REACTIONS WITHIN THE PEROXY COMPOUNDS BY L. J. CS~NYI Institute of Inorganic and Analytical Chemistry, University of Szeged, Hungary Received 22nd January, 1960 During the oxidatioh of hydrogen peroxide in presence of peroxy compounds of other types, an induced reaction occurs because, in the reaction between hydrogen peroxide and I-equivalent oxidizing agents, HO2 radicals are formed, which readily attack the peroxy compound present. In most cases the reaction is chain-like. The extent of the induced reaction depends markedly on the experimental conditions, e.g., the concentration of the participants, the speed of titration and the dilution of the solution, the rate of stirring, the temperature, the hydrogen-ion concentration and presence of foreign substances catalyzing the induced reactions.The dependence on these conditions can be regarded as a general characteristic of the chemical induction. It had been already recognized in the fifties of the last century that in several oxidation-reduction reactions the so-called co-existence principle (the assumption that the individual processes take place independently of each other) is not always valid. The phenomenon of one reaction taking place alone at a very slow rate (if at all) and being accelerated markedly by the simultaneous occurrence of another reaction of measurable velocity, is called, after Kessler,l chemical induction. An hduced reaction may be represented by the following scheme. If the reaction occurs separately A+I-+C and A+Ac+D, then in a solution of the three substances the reaction A+Ac-tD also takes place : A+I+C, {A+,,+,.Reaction (1) is called the inducing, main or primary reaction which induces reaction (2). Substance A, taking part in both reactions (1) and (2), is called the actor, substance I the inductor, and Ac the acceptor. The extent of the in- duced reaction is conventionally expressed by the induction factor &, defined as the ratio of the equivalents of the induced reaction to those of the primary reaction. The actor may have either reducing or oxidizing properties. The chemical character of the inductor and the acceptor is, respectively, always the same and opposite to that of the actor. When the actor is an oxidizing substance then the acceptor and the inductor are reducing agents, and vice versa. Soon after the beginning of this century Luther and Schilow 2 attempted to systematize the induced reactions.They distinguished coupled and induced chain reactions. Coupled reactions are formed when the primary reaction results in the production of an intermediate which enables also the acceptor to react. The main characteristic of this reaction is that the value of the induction factor even under favourable conditions for the induced changc is low, maximally 2. The plotting of F;. as a function of thc ratio of the initial concentrations of the acceptor I46L. J. CSANYI 147 and inductor, fi = f[Ac]/[I]o, results in a curve having a saturation value. The limiting value makes it possible to describe the overall equation of the induced reaction, i.e.the stoichiometry of the change. According to the authors mentioned above, induced chain reactions take place when the inductor catalyzes the very slow reaction between the acceptor and the actor. However, since the chemical character of the inductor and ac- ceptor is the same, the actor reacts also with the inductor ; consequently, a part of it is excluded from the catalysis (catalysis with the destruction of the catalyst). The principal characteristic of reactions of this type according to Luther and Schilow is that the value of F;: largely exceeds 2; further, the function fi = f[Ac]/[I], yields an exponential curve. The chemical induction is a common phenomenon and plays an important role in many difficult problems. The importance of the induced reactions is most obvious in analytical chemistry since the overwhelming majority of our procedures is based on the co-existence principle, But the role of the induced reactions is no less important in various autoxidation, polymerization and biological processes.It is obvious that the study of induced reactions gives many useful results con- cerning the chemical properties of intermediates (radicals, radical-like ions, ions with unstable valency) that cause the induced reactions. Our aim in these investigations started with analytical interests. Previously, we dealt in detail with the problem of the simultaneous determination of peroxy compounds of different types and pointed out the shortcomings of methods gener- ally used.3 Among these, we first emphasized the inaccuracy and poor repro- ducibility.Further investigations were therefore carried out to elucidate and eliminate the sources of errors and irreproducibility. Our preliminary results showed that the majority of disturbing effects can be traced to chemical inductions. Therefore by choosing different, analytically important systems, the induced reactions occurring therein were investigated. In the present paper we deal ex- haustively only with the H2Qf H2S2O8 system. Results obtained with other systems are only briefly mentioned. When titrating hydrogen peroxide in the presence of peroxydisulphate with standardized oxidizing solution (potassium permanganate or ceric sulphate), negative errors of about the same amount appear, both in the amounts of hydrogen peroxide and peroxydisulphate.This shows that, in the reaction responsible for the errors, hydrogen peroxide and peroxydisulphate take part in a mole ratio of 1 : 1. The reaction between hydrogen peroxide and peroxydisulphate ions is very slow, the time of +rd conversion being 86.3 h (table 1); therefore during the 1-2 min titration time, observable changes may occur only when for some reason the rate of the reaction increases. TABLE THE TIME OF +rd CONVERSION OF REACTIONS (IN MIN) H2S20.9 self decompositions H202 H202+H2S208 H ~ S ~ O ~ + A S ~ O ~ H202fAszO3 catalyst 20300 57000 5180 563 134 - 7680 20000 30 6 105 CU" The rate in question is slightly influenced by manganous or cerous ions formed from the oxidizing agents, thus an induced reaction is supposed to take place. The principal characteristics of the induced reaction may be summarized as follows : (i) In every case an induced reaction takes place when hydrogen peroxide is titrated in the presence of peroxydisulphate with 1 -equivalent oxidizing agent (direct titration).(ii) The extent of the induced reactions depends on the concentration of the reaction component-the higher the concentration of the partners, the greater is the error.148 INDUCED REACTIONS (iii) On increasing the acid concentration of the solution the error decreases. The rate of the reaction between hydrogen peroxide and peroxydisulphate likewise is inversely proportional to the acid concentration of the solution.4 (iv) In direct titrations the increase of the speed of titration decreases the error, but in indirect titrations increases it.(v) On diluting the solution to be titrated, other factors being constant, the induction error increases. (vi) On raising the temperature of the solution the error increases. Increasing the rate of stirring of the solution yields similar result. (vii) During the induced reaction the polymerization of acetanilide and bi- phthalate can be observed to a small extent, as a result of which the solution becomes brownish or yellowish. In the presence of these substances the induction error markedly decreases. Acrylonitrile, although it also decreases the error during the induced reaction, is not polymerized. (viii) On adding arsenous acid to the given system the H202 error becomes practically zero; instead the As203 error appears and the H2S208 error still remains.(ix) Regarding their effects on the induction reaction, the substances fall into two main groups : (a) the induction error is increased by addition of cupric, silver and ferric ions. These substances are effective catalysts in the oxidation reactions of peroxydisulphate. Moreover, fluoride ions also cause an increasing error. (6) the induction error is decreased by addition of alkali, alkali-earth metal ions, ions of aluminium and tin groups, and by some tervalent ions (chromic, cerous, lanthanum). Halide, acetate, and uranyl ions very greatly decrease the error. It is worth noting that cupric ions do not change the course of the induced reaction but merely considerably increase the extent of the reaction. (x) During induced reactions, intermediates with strong oxidizing power are formed which, if their concentrations are sufficiently high, partly or fully destroy the very resistant ferroin indicator.(xi) When carrying out the reversed titration, i.e. when ceric sulphate is titrated with hydrogen peroxide in the presence of peroxydisulphate the induction error is practically zero. DISCUSSION The course of the induced reaction is interpreted as follows : it is generally known 5 that the oxidation of hydrogen peroxide with 1-equivalent reagents takes place in two steps : H202+Ox-+H02+H'+Red, (3) HO, + Ox-, O2 + H+ +Red. (4) The rate-determining step of the reaction6 is (3). The H02 radical formed in (3) is a strong reducing agent 7 (02+H++ e+H02 . . . +0*13 V) therefore it readily attacks the peroxydisulphate ion present : HO2+S2O:--'O2+HSOi + SOL.( 5 ) The sulphate radical-ion formed in reaction (5) is a strong oxidizing agent, therefore it reacts with hydrogen peroxide which has amphoteric properties from the redox point of view :L. J . CSANYI 149 and the radical H02 is formed again. The H02 radical attacks the peroxydi- sulphate ion again and so, without addition of the oxidizing solution, considerable amounts of hydrogen peroxide and peroxydisulphate react. The amount of hydrogen peroxide and peroxydisulphate disappearing during these reactions, apparently depends on the ratio of the rates of the competing reactions (5) and (4). The greater this ratio, w5/w4, the greater is the error. There- fore any factor, which decreases the rate of reaction (4) or increases that of re- action ( 5 ) effects an increase of the induction error.This accounts for the variation of the induction error on changing the speed of titration. At a low speed of titra- tion, only a slight amount of the oxidizing agent is added to the solution, and thus the local concentration of the measuring solution is also smaller during a more rapid titration. Consequently, reaction (4) takes place to a smaller extent and the error increases. On increasing the dilution, or raising the rate of stirring of the solution titrated, the local measuring solution concentration is similarly decreased, thus causing an increasing error. Under optimal experimental conditions for the induced reactions the value of the induction factor is 0.42. The function fi =f[Ac]/[I]o gives a curve having a saturation value.Bearing in mind the size of Fi and the shape of this curve, the reaction must be considered as a coupled reaction. However, several observations point to the occurrence of a chain process. We found that with a low speed of titration when discrete drops are formed, the induction error further increases on increasing the drop time. It was noted that, independently of the drop time, the size of the single drops is the same, therefore at a constant stirring rate the local concentration of the measuring solution in every case is the same. Now it would be expected that the induction error would not change with further increase of drop time. On the contrary, the error increases, which observation can only be interpreted if the HO2 radical formed in step (3) starts a chain reaction.The reaction having a longer drop time has a longer time at its disposal than at greater dropping frequencies, since the destruction of the radicals according to reaction (4) has less probability of taking place when the drop-time is longer. The chain-character of the reaction is supported also by the inhibiting effect of halides ; thus halide ions readily react with sulphate radical ions thereby minimiz- ing the error by breaking the chain. The fact that only a small amount of certain halide ions exerts strong inhibiting effect shows that the length of chains is fairly considerable. During halide inhibition, only minimal amounts of free halogen are formed; this can be attributed to the fact that the halogen molecule, formed by reaction of the sulphate radical ion, reacts with the HO2 newly formed in reaction (3) and halide is formed anew.It was observed that the extent of induced reaction is strongly decreased, not only by halide ions but by halogen molecules, too. This is possible because they oxidize the H02 radical formed, thus inhibiting the pro- duction of reaction chains. The occurrence of induced reaction by the presence of radicals is confirmed by the observation that during titration the polymerization of acetanilide and bi- phthalate is also perceptible. Naturally these substances by reacting with the chain carriers, strongly decrease the extent of the induced reaction. The kinetic character of the induced reactions is supported by the fact that on increasing the concentration of the participants of the reaction and raising the temperature, the induction factor increases.The fact that such a stable molecule as ferroin is partly or wholly destroyed during the induced reaction, also supports the correctness of the part played by reaction (5), as the sulphate radical ion is one of the strongest oxidizing agents. On adding arsenous acid to the solution, the induction errors are different; the H202 error practically disappears and instead the As203 error is introduced while the H2S208 error is only slightly increased. To interpret the role of arsenous acid, we suppose that instead of (6) the following reaction occurs : AS I1 I + s 0; -+AP + s 0:- . (7)150 INDUCED REACTIONS The rate of reaction (7), noting that the potential values of the AsIII/AsIV and H202/H02 couples are -0.7 and - 1.5 V, is, as expected, greater than that of reaction (6).AdV formed in reaction (7) is a fairly strong reducing agent (its normal oxidation potential 8 being about 0.3-0.4 V), so the following reaction may take place : As Iv + S2 0; - -+AsV + SO, + SO:-. Steps (3, (7) and (8) form a closed reaction chain which explains, on the one hand, the decrease of the H202 error, and on the other, the appearance of the As203 error and the persistence of the same HzS2O8 error. On the basis of the above reaction scheme, it is clear that in reversed titrations the induction error should be zero, because the oxidizing agent is present in excess until the end-point, and therefore radicals H02 formed in step (3) are mainly consumed by the competing reaction (4).The effect of foreign substances can be interpreted as follows. Silver, cupric and ferric ions, as has been previously pointed out,9 greatly weaken the -0 . 0- bond of peroxydisulphate, thus making it more easy for electron transfer to the peroxydisulphate ion to take place. It is therefore reasonable that the rate of reaction (5) in presence of these ions is markedly increased (see table 1). Alkali and alkali-earth metal ions, as may be experimentally proved, increase the rate of reactions (3) and (4), which results in a decrease of the concentration of H02 radical, and thus in a decrease of the induction error. The error-decreasing effect of halide ions and other easily oxidable and re- ducible substances has previously been interpreted by assuming that these sub- stances, by consumption of chain carriers, effect a decrease of the induced change.Our results may be summarized in the following reaction scheme : (3) (4) ox+ H202-H02+ Ox-+02 (6) + + 4 ( 5 ) (7) \I + + I' (8) SO, += S2Qg- AS'''-+ AsIV Other peroxy acids, e.g. peroxy sulphuric, peroxy acetic and peroxy phosphoric acid behave analogously. The presence of molybdenate, tungstate and osmium tetroxide causes an in- duced error in the oxidimetric titration of hydrogen peroxide because these sub- stances form with a part of the hydrogen peroxide present, a compound of the hydroperoxide type, which reacts with the radical H02 acting as an acceptor. All these induced reactions also belong to the group of closed-chain reactions.In the H202+Br2 system the induced reaction takes place similarly to that described above. Although under normal experimental conditions hydrogen peroxide is incapable of oxidizing bromide ions to bromine, in contrast to the previous systems, an open-chain reaction is formed, and not a closed-chain one, We found no induced reaction in the H202+oxalic acid system. Our investigations show that an induced reaction may proceed not only during the oxidation of hydrogen peroxide, but also that the reduction of peroxy com- pounds may also serve as an inducing step. In the HSCN+H2SQs+H202 and HSCN+ CH3CBOOH+ H202 systems, the induced disappearance of hydrogen peroxide takes place during the reaction between peroxy acids and thiocyanate ions because a strong reducing intermediate (sulphoxylic acid) is formed which, in contrast to thiocyanate ions, instantaneously reduces hydrogen peroxide.This induced reaction belongs to the group of open-chain reactions. These results, in addition to clearing up the sources of error occurring during the simultaneous determination of peroxy compounds and rendering possible theL . J . CSANYI 151 avoidance of these errors, also iqdicate that the viewpoint formerlv developed to explain induced reactions is too rigid to characterize correctly this dynamic phenomenon. Furthermore, our experimental results draw attention to the fact that only on the basis of the magnitude of the induction factor and the shape of the plot of the function F;:(=f[Ac]/[I]o) it is not possible correctly to classify in- duced reactions.In most of the systems investigated by us, the limiting value of F;: was less than 1 (except for the H202+Os04 system where it is 7.5) and the function always gave curvcs having limiting values and the experimental data in most cases supported the existence of a closed-chain reaction. Therefore, in our opinion, it is more correct to regard the induced reactions as chain reactions. The open- and closed-chain reactions (previously called coupled and induced chain reactions, respectivcly) are obtained as limiting cases and consequently the more frequently occurring transitory reaction types are not to be considered exceptions. On this basis, to characterize the induced reactions, a simplified, general reaction scheme is given, from which, depending on the rate of the individual steps, the limiting cases, and even the transitory types, can be derived.Assuming that the actor forms its stable end-products through steps A,,+A1-+ Ared and the acceptor is oxidized according to (Ac),d-.(Ac)l-+(Ac)o,, and furthermore, for simplicity, assuming that the inductor is transformed directly to the end-product according to IredgIox, which is a reversible reaction, we can write : (Ac)l j(Ac)ox+Arcd (14) According to this scheme, if the rate of reactions (12) and (13) is very slow, steps (9), (lo), (11) and (14) form an open-chain reaction, i.e. a typical coupled reaction is produced in which the coupling is made by the A1 chain carrier. The value of fi depends on the rate of competing steps (lo), (1 1) and (14).With increas- ing conccntration of (AC)red steps (1 1) and (14) become more predominant and with sufficiently high concentration, A1 will react exclusively by (11) and (14), and thus F;: reaches its limit value. When the rate constants of the reactions (10) and (ll), klo>kll, and the rate of reactions (12) and (13) (w12 and ~ 1 3 ) are very low, the tendency of Fi to reach a limiting value owing to the considerable competition of reaction (11) cannot be proved experimentally, consequently the shape of function 4, as conceived earlier, refers to an induced chain reaction. Naturally if the acceptor possesses only two diffcrent oxidation stages (e.g. in the HzOz+Br2 system where a redox change is possible only between the forms of bromine of 0 and - 1 oxidation number), then an open-chain reaction will also occur.In contrast, if ~ 1 3 or w12, or both, are sufficiently great, and klo is low enough, or at least kloSkll, then a typical induced chain reaction results. In this case, fi does not reach a limiting value even in the presence of a very great (Ac),~ concentration (e.g. in the HzOz+ FeII system). But if klo>kll, and ~ 1 3 and w12 are not too great, it may happen that function Fi = f[Ac]/[I]o gives a curve having a saturated value. Naturally in this case, the value of Fi will be very low, in general, less than 1, even with high concentrations of the acceptor. Then, although the induced reaction belongs to closed-chain systems, on the basis of the earlier classification it should be considered a coupled152 INDUCED REACTIONS reaction. The induced reaction in the H202+H&Os-system is an example of this. On the basis of the considerations briefly outlined above, the earlier data and our own results can now be uniformly interpreted and, it may be hoped, can con- tribute to the further elucidation of the complicated phenomenon of chemical induction. 1 Kessler, Pogg. Ann., 1863, 195, 218. 2 Luthcr and Schilow, 2. physik. Chem., 1903, 46, 777. 3 Csanyi and Solyniosi, Acta Chim. Hung., 1957,13,9 ; Acta Chim. Hung., 1957,13, 19. 4 Csinyi and Majzik, in press. 5 Baer and Stein, J. Chem. SOC., 1953, 3176. Ardon and Stein, J . Cheni. Soc., 1956, 6 Baxendale, J . Chem. SOC. (Spec. Publ.), 1954, 1, 40. Sigler and Masters, J . Atner. 7 Latimer, The Oxidation States of the Elenzents and their Potentials in Aqueous 8 Csanyi and Szab6, Talantu, 1958, 1, 359. 9 Galiba, Cshnyi and Szab6, Z. finorg. Chern., 1956, 287, 169. 104. Chen?. SOC.. 1957, 79, 6353. Solutions (Prentice Hall, 1952), 2nd ed., p. 48.

ISSN:0366-9033    

DOI:10.1039/DF9602900146

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

PEROXY-COMPLEXES AS INTERMEDIATES IN THE CATALYTIC DECOMPOSITION OF HYDROGEN PEROXIDE BY M. L. HAGGETT, PETER JONES AND W. F. K. WYNNE-JONES Chemistry Dept., King's College, University of Durham, Newcastle-upon-Tyne, 1 Received 28th January, 1960 Kinetic studies of the iron-salt-catalyzed decomposition of hydrogen peroxide in H202+ H20 mixtures have suggested the importance of ferric-peroxy complexes as inter- mediates in the reaction. In this paper, spectrophotometric evidence for the existence of complexes of the type proposed is described. Measurements of the kinetics of chromium(V1)-salt catalysis of the decomposition of hydrogen peroxide in H202-t.H20 mixturesare also reported. The results are shown to be consistent with the model pro- posed for iron-salt catalysis. The implication that the iron and chromium catalyses proceed by a similar mechanism is discussed and the analogy between the proposed model and an enzyme reaction is pointed out.Many investigations of the homogeneous iron-salt-catalyzed decomposition o hydrogen peroxide in dilute aqueous solution have been described. The field has been reviewed by Baxendale 1 and by Weiss.2 Although a ferric-peroxy complex (formulated as (Fe3+HO;) or (FeH02)2+) has been reported by Evans, George and Uri,3 it has been rejected 4 5 as an intermediate in the catalytic reaction. The kinetic experiments in dilute solution have led to the conclusion that the key reaction is of the form : Fe3+ +H202-+Fe2+HO,+H+, Fe3+ +HOT +Fe2+ + HO,. or There are various schools of thought as to the details of the further stages of the reaction but all models involve the radicals OH and HO2 as intermediates, following the theory of Haber and Weiss.6 All the models are amplifications of von Bertalan's suggestion7 that the catalysis occurs as a result of the alternating reduction of ferric and oxidation of ferrous according to the overall reactions : 2Fe3C+H202-+2Fe2++2H++02, 2Fe2+ +H202-+2Fe3+ +20H-.The possible mode of intervention of the ferric-proxy complex in the reaction was considered 4 according to Fe3+ +H0,+(FeH02)2++Fe2+ +HO,. It was rejected because no deviation from first-order dependence on [HzOz] was observed in the dilute aqueous hydrogen peroxide solutions employed. We have recently reported measurements of the kinetics of the reaction in H202+ H20 mixtures.* These mixtures provide unusually simple mixed solvents for kinetic work in that they do not depart greatly from ideal behaviour and are effectively isodielectric over the whole range of composition at ambient temper- atures.9 These properties enable the variation of the catalytic reaction velocity with solvent composition to be related directly to corresponding changes in the 153154 DECOMPOSITION OF HYDROGEN PEROXIDE concentrations of reactant species. The results obtained 8 were consistent with control of the reaction velocity by ferric-peroxy complexes as interrnediatcs.The model proposed for the reaction may be summarized : [ Fe( H, O),] + + H, 0:s H, 0 + [ Fe( H, 0), (H, 0, )] + 2 products + [Fe(H,0),(H,0,)]3+ +H2O$2H2O+ [Fe(H,0)4 (H,0,),]3'~products k3 products.This model reduces to the required form in dilute aqueous solution where the water concentration is constant. The dependence of the reaction velocity on [H+] is not shown above. The reaction velocity decreases with increasing [H+] and it is likely that this means that an essential preliminary to further reaction of the complexes is acid ionization of a proton from a hydrogen peroxide molecule in the complex. This would allow electron-transfer processes of the type : Fe3+HO; +(FeH0,)2f+Fe2' +HO,. A more detailed interpretation may require other protonation processes involving the whole solvent complex. Although peroxy-complexes have been rejected as catalytic intermediates for iron-salt catalysis, they have been invoked to explain the behaviour of many of the wide variety of homogeneous hydrogen peroxide catalysts, including iron com- plexes, e.g.ferricyanide,l and, notably, the haemoprotein enzymes such as catalase.5 In this paper, we first present independent evidence for the existence of the type of ferric-peroxy complex proposed. Kinetic results for chromium(V1)-salt catalysis are also described. This system has long been known for the curious kinetic form which it exhibits in dilute hydrogen-peroxide solutions. The results are shown to be consistent with the concept of peroxy-complexes as intermediates developed for iron-salt catalysis. The cxperimental arrangements have been described in detail elsewhere.8 RESULTS FERRIC-PEROXY-COMPLEXES Evans, George and Uri 3 have interpreted spectrophotometric data for solutions of ferric perchlorate in acidified hydrogen peroxide in terms of the formation of a peroxy complex, Fe3++HO; +Fe3+HO;(or [FeH0,I2').Evans et al.3 obtained results at concentrations of H202 up to 33 M but the above model refers only to low H202 concentrations. The results given in fig. 2 of their paper may be compared with the peroxy-complex model that we have proposed for iron-salt catalysis. The limiting case of our model at sufficiently high [W] corresponds to the formation of a monoperoxy complex only : [Fe3+(H,0),] + H,0,2[Fe3 '(H,O),(H,O,)] + H,O, [Fe3+(H,O),(H,O2)]%[Fe3+(H,0),(HO;)] + H'. Under conditions where neither the hexa-aquo ferric ion nor the undissociated peroxy-pentaquo-ferric ion absorb appreciably the optical density of the solutions is given by155 where X , is the iron-salt concentration, Y the mole fraction ratio of hydrogen peroxide to water Xp/X, = Xp/(l - X,), X,+ the hydrogen ion concentration, E the extinction-coefficient of the complex and I the cell-path length.M. L . HAGGETT, P . JONES AND W. F. K . WYNNE-JONES XVf/XP FIG. 1 .-Optical density data for solutions of iron perchlorate in H202-1- H20 mixtures, [Fe(C104)2] = 5 x 10-3 M, [HC104] = 0.5 M, temperature 20°C. Full circles-Evans. George and Uri ; open circles-this work. In fig. 1 we compare Evans, George and Uri results with eqn. (1). We have made confirmatory measurements using a Unicam SP 500 spectrophotometer at 430mp with a slit width of 0.04mm. The two sets of results are in good agree- ment with eqn.(1) and in fairly good agreement with one another. At higher peroxide concentration our optical density values are higher than those of Evans et al. We found that, under these conditions, quite rapid decomposition occurred and the optical density decreased with time. The values shown in fig. 1 were ob- tained as soon as possible after mixing solutions. This result provides strong support for the type of ferric-peroxy complex postulated on kinetic grounds. CATALYSIS BY CHROMIUM(VI) SALTS Addition of chromium(V1) salts to strongly acidified aqueous solutions of hydrogen peroxide results in the immediate formation of ‘‘ blue perchromic acid ”, which decomposes quickly, evolving oxygen and leaving the chromium completely reduced to Cr(II1). Subsequent decomposition of the hydrogen peroxide is very slow.At lower acid concentrations the addition of chromium(VI) salt produces a violet solution and a catalytic decomposition of the hydrogen peroxide ensues. Catalytic decomposition persists when neutral or alkaline conditions are employed, but the colour of the solution changes gradually, becoming orange-yellow in alkaline solution. The kinetics of the catalytic reaction in dilute aqueous solution of hydrogen peroxide have been studied by Spitalsky 10 and by Kobosev.11 These investiga- tions have been reviewed by Baxendale.1 The reaction is reported to be ap- proximately first-order in total chromium. The striking feature is the curious dependence on hydrogen peroxide concentration, the reaction sometimes bcing approximately first-order, but under other conditions, the rate against [H202] curve156 DECOMPOSITION OF HYDROGEN PEROXIDE shows a sharp maximum. The appearance of the solutions has led to suggestions of peroxy-complexes as intermediates, but their constitutions have not been definitely established. We have measured the velocity of the catalytic reaction under various con- ditions in H202+ H20 mixtures.In alkaline solutions the reaction is accurately first-order in total chromium concentration but in acid solutions a small additional second-order component is observed. The variation of reaction rate (- dX,/dt) with the mole fraction of hydrogen peroxide X , is shown for an added potassium dichromate mole fraction Xc = 1 . 6 ~ 10-5 and various acidities in fig. 2, 3 and 4.Fig. 2 shows the change in form when the acid concentration is decreased from XHt=8.05x 10-5 to " neutral " conditions. At the highest acid concentra- tion the curve is monotonic (curve (a)) but, as the acid concentration diminishes, a maximum appears and increases in magnitude while the rate at the peroxide-rich end of the curve diminishes (curves (b), (c), (d)). Under alkaline conditions (fig. 3) the maximum at first increases as XOH- (or more strictly XHO;) increases and the rate at the peroxide-rich end now starts to increase also (curves (e) and (f)). At higher XOH- the maximum eventually decreases and the peroxide-rich limb has now increased so that the curve passes through both a minimum and a maximum (curve ( g ) ) . In curve (h) the maximum and minimum have disappeared and in curve ( j ) this situation persists although the rates are considerably lower.At even higher XOH- a new maximum appears and curve (k) shows both a maximum and a minimum. In curve ( I ) the maximum has developed and the peroxide-rich limb now shows a monotonic decrease of rate with increasing X,. I1 / fa),' 0.2 0.4 0.6 XP FIG. 2. FIG. 2, 3,4.-Variation of rate of decomposition with mole fraction of hydrogen peroxide. X, (added as potassium dichromate) = 1.61 x 10-5. Temperature 25°C. Added HClO4 mole fraction x 105 : (a) 8.05, (b) 4.03, (c) 0-81, (d) no added acid or alkali. Added NaOH mole fraction x 105 : (e) 0.81, ( f ) 8.05, (g) 20.9, (h) 40.2, ( j ) 80.4, ( k ) 201, (I) 402.HAGGETT, P . JONES AND W . F . K. WYNNE-JONES 157 0.2 0.4 0.6 XP FIG.3 XP FIG. 4.158 DECOMPOSITION OF HYDROGEN PEKOXIDE This striking collection of kinetic data is immediately explicable in terms of the peroxy-complex intermediate model with few additional assumptions. The limiting case of the model at high acid concentrations is the formation of a mono- peroxy complex only. The proton-acid behaviour of the complexes will not be written into the model explicitly at the moment. For this case KI k i A + H202 +H20 + B ; B-+ products, where A is the initial form of the catalyst and B the mono-peroxy complex inter- mediate. The reaction velocity is given by or rearranging, (3) where Y = Xp/Xw = Xp/(l - X,) the mole fraction ratio of [HzO2]/[HzO] and X, is the total catalyst concentration. In fig. 5 the data of curve (a) arc compared with eqn.(3) by plotting l/v against l / r = Xw/Xp. xw/x, FIG. 5.-Graph of l/v against Xw/Xp for data of curve (a). The subsequent development of the rate maximum requires the assumption of the formation of a bi-peroxy complex. For this case ki A + H , 0 t 2 H 2 0 + B ; B-products, €3 + H202$H20 + C ; C2products. where C is the bi-peroxy complex. The reaction velocity is given byM. L . HAGGET'T, 1'. JONES AND W. F . K . WYNNE-JONES 159 This equation assumes a simple form if the bi-peroxy complex is inactive, i.e. k2 = 0. The reaction velocity will first increase with r (i.e. approximately first- order if Xw-+l), pass through a maximum and finally, when KlK2r2>1+K1~ will have the form : A plot of z, against l / r should be a straight line through the origin.comparison of the data of curve (d) with eqn. (5). Fig. 6 shows a 1.c v x 104 Q ! 0 . 5 I *o 1.5 Xwi X , FIG. 6.-Graph of er against X,/X, for data of curve (d). At higher XOH- the rate maximum disappears. Fig. 7 shows a plot of rate against Y for the data of curve (h). This is consistent with the above model with k2 # 0. Eqn. (4) can then assume a simple form when l<K1r>K1&~2, where II = k,X,+ k2K2Xcr. (6) Under more alkaline conditions, the appearance of a new maximum (curves (k) and ( I ) ) may be the result of the formation of tri- and tetra-peroxy complexes. Our results are less detailed in this region. We have so far avoided giving precise formulation to the chromium(V1) peroxy complexes. The simplest type of process which would give the required result would be : CrOi- +H,O,+CrO;- +H,O, thus formulating the complexcs as the anions of various pcroxychromic acids.The way in which the kinetics vary with acidity suggests that, for example, tlic complex which wc have termed mono-peroxy can in fact exist in a variety of forms160 DECOMPOSITION OF HYDROGEN PEROXIDE which differ in degree of protonation. This is not unreasonable by analogy with known processes, such as Hf + CrOi- +HCrO,. 6 - 5 - 4 - V X 104 3 - 0 5 1.0 I. 5 XplXw FIG. 7 . 4 r a p h of v against X,/Xw for data of curve (h). Equilibria of this type may affect the whole solvent complex so that the actual number of species could be large. These species which could well be classed as “mono-peroxy ”, ‘‘ bi-peroxy ”, would presumably differ in reactivity so that both rate constants and equilibria could be complex functions of the hydrogen-ion concentration.In acid solutions a further likely complication is the formation of binuclear species. Nevertheless the results clearly demonstrate the close similarity between the kinetics of the iron and chromium catalyses. DISCUSSION Although our experiments are, as yet, incomplete, they provide strong evidence that the iron and chromium catalyses of the decomposition of hydrogen peroxide are essentially similar processes. This is an important result because it suggests a synthesis of the two theories which have been developed to account for the decomposition of hydrogen peroxide by various catalysts. The theory of com- pensating oxidation-reduction reactions has been developed for iron-salt catalysis and has also been applied to halide catalysis and copper-salt catalysis.It is based on the fact that the catalysts are capable of existing in at least two oxidation states. The idea of an intermediate complex, usually peroxidic, has been used for other catalysts, but has not been developed in a general form. These ideas are reviewed in detail by Baxendale.1M. L . HAGGETT, P . JONES AND W. F. K . WYNNE-JONES 161 Our results suggest that the two ideas may be readily combined, that is, the actual participants in a reaction proceeding by a. compensating oxidation-reduction mechanism are peroxy-complexes. The essential feature of our method is the use of mixed H202+H20 solvents which, because of their simple properties, enable the competition between the solvent components as complexing agents to be studied and so reveal the nature of the species formed.The role of co-ordinated solvent is obscured when measurements are restricted to aqueous solutions so that direct tests of the possible application of our model to other systems cannot be made at the moment. The only other system for which kinetic data in H202fH20 mixtures are available is copper-salt catalysis.12 The reaction shows a pronounced photo- chemical effect and a dependence on the ageing and intermediate treatment of the reaction vessel. Our results suggest that the reaction may not be truly homo- geneous and it is possible that traces of iron from the Pyrex container interfere in the measurements.The reciprocal promotion of iron and copper catalyses is well known. For many other systems the importance of peroxy-complex intermediates is either established or inferred. As Baxendale 1 points out an attractive feature of a peroxy-complex intermediate model is the similarity of such a scheme to an enzyme reaction. Many authors have attempted comparisons between the cata- lytic behaviour of " free " ionic iron and the haemoprotein enzymes, notably catalase. This problem is considered in some detail by George.5 Catalase- hydrogen peroxide complexes have been shown to be important catalytic inter- mediates by Chance.13 Peroxy-complexes are also formed by the catalytically active protein-free enzyme precursors such as hemin.14 George considers that the ferric-peroxy complex described by Evans, George and Uri 3 plays no distinct role in the inorganic catalysis and from this standpoint it is diacult to draw a definite analogy with the enzyme systems.Our kinetic results and reassessment of the data of Evans et aZ.3 show up a remarkable analogy between the inorganic and enzyme catalyses. The peroxy- complex model proposed is entirely analogous to the Michaelis-Menten model for the enzyme systems when the variation in water concentration is taken into account. Thus the enzymic scheme : E + S fES+products + E, ES + S +ES, (inactive), which includes inhibition by excess of substrate, exactly describes the kinetics observed for chromium catalysis under " neutral " conditions. George 15 has studied the catalase decomposition of hydrogen peroxide over a range of hydrogen peroxide concentrations up to 5 M and has observed that the velocity passes through a maximum. This result is particularly interesting in the present context but a more extended study would be desirable in order to test the form of the curve quantitatively .The influence of structure is evidently of great importance in determining the reactivity of peroxy complexes. In bi-peroxy complexes the mutual influence of the two peroxy groups which, in some cases, amounts to an inhibition of chromium catalysis, is an interesting effect. In octahedral ferric complexes this must be complicated by the possible formation of cis and trans isomers. An interesting possibility with bi-peroxy complexes is that both oxidation and reduction could occur within the complex by a flow of electrons from one peroxy group to the other via the central ion.This intramolecular process would avoid the necessity for free radicals, although with mono-peroxy complexes it is difficult to avoid the free-radical mechanism. One aspect of this problem is revealed by studies of the inhibition of the cata- lytic reaction. We have shown, for example, that the retardation of iron-salt F162 DECOMPOSIT~ JN ;?F HYDROGEN PEROXIUE catalysis by suIphate ion can be quantitatively accounted for by the formation of a ferric-sulphate complex, Fe3 + SO:-+(FtSO,)', on the assumption that the complex is inactive16 Inhibition by complete block- ing of the co-ordination sphere of the catalyst ion is readily imagined and is established for ferricyanide catalysis where an induction period is observed which involves the removal of one of the C N - groups.17 The sulphate effect must be on the energetics of breakdown of the complex.Presumably the negatively charged sulphate group increases the energy requirements of electron transfer to the central ion, an effect which seems reasonable on a simple electrostatic model. We are indebted to Messrs I.C.I. Ltd., for the award of a Fellowship to P. J., to the Council of King's College for the award of a Levin Scholarship to M. L. H., and to Laporte Chemicals Ltd., for gifts of hydrogen peroxide. 1 Baxendale, Advances in Catalysis (Academic Press, N.Y.), 1952, 4, 31-86. 2 Weiss, Advances in Catalysis (Academic Press, N.Y.), 1952, 4, 343-365. 3 Evans, George and Uri, Trans. Faraday SOC., 1949,45, 230. 4 Barb, Baxendale, George and Hargrave, Trans. Faraday SOC., 1951,47, 462, 591. 5 George, Advances in Catalysis (Academic Press, N.Y.), 1952, 4, 367-428. 6 Haber and Weiss, Proc. Roy. SOC. A , 1934, 147, 332. 7 von Bertalan, 2. physik. Chem. A, 1920,959, 328. 8 Jones, Kitching, Tobe and Wynne-Jones, Trans. Faraday SOC., 1959,55,79. 9 for a review see Schumb, Sattcrfield and Wentworth, Hydrogen Peroxide (Reinhold, New York, 1955), chap. 5. 10 Spitalsky, Z. anorg. Chem., 1907, 53, 184; 1908, 56, 72 ; 191 1, 69, 179. 11 Kobosev and Galbreich, Actaphysicochim., 1945, 20, 479. 12 Jones, Tobe and Wynne-Jones, unpublished. 13 Chance, Acta Chem. Scand., 1947, 1, 236; for a review see ref. (5). 14 Haurowitz, Enzymologia, 1937, 4, 139. Haurowitz, Brdicka and Kraus, Enzy- 15 George, Nature, 1947, 180,41; Biochem. J., 1948, 43, 287; 1949,44, 197. 16 Jones, Tobe and Wynne-Jones, Trans. Faraday Soc., 1959, 55, 91. 17 for a review, see ref. (l), pp. 67-71. mologia, 1937, 2, 9.

ISSN:0366-9033    

DOI:10.1039/DF9602900153

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

PHOTOCHEMICAL OXIDATION-REDUCTCON REACTIONS OF SOME TRANSITION METAL COMPLEXES BY ARTHUR W. ADAMSON Dept. of Chemistry, University of Southern California Received 2 1 st January, 1960 Photoredox reactions of various complex ions are discussed, with particular reference to Co(1II) compounds. It is found that the entire range of behaviour from mainly photo- aquation to mainly photoredox dccomposition occurs in the acidopentamine series, with a close correlation bctween type of photolytic behaviour and ease of oxidation of the ligand. Results on the wavelength and temperature dependence of the partial quantum yields for photoaquation and photoredox decomposition are given for Co(NH&Br2+ and Co(NH3)5(SCN)2+. There is a pronounced decrease in both +aq. and +red. as well as in the ratio '&d./&q.as wavelength increases from 370 mp to 600 mp, as well as some in- version of the temperature dependence. The findings are interpreted in terms of a homolytic fission of the metal-ligand bond as the first consequence of light absorption. The application of this type of mechanism to other transition metal complexes is also considered. The photochemistry of a number of complexes of various transition metals has been under investigation in this laboratory and the present paper summarizes some recent findings, with emphasis on the behaviour of the Co(II1) acidopentamine series. It was noted in a previous publication 1 that Co(II1) complexes show a marked preference for photochemical redox reactions as opposed to substitution processes.Thus, irradiation of Co(NH3)512+ with either visible or near ultra- violet light led exclusively, and with high quantum yield, to the formation of Co(I1) and iodine, under conditions such that the thermal reaction was solely one of aquation. Similarly, the ion Co(NH3)5(SCN)2+ showed appreciable photoredox behaviour where none at all was observed in the thermal reaction. Here, however, photochemical aquation also occurred, and in proportion which appeared to be wavelength dependent. The species Co(NH3)&12+, on the other hand, showed mainly photoaquation. These effects appeared to be ones of degree and related to the ease of oxidation of the ligand involved. The dependence of the ratio of redox to aquation reaction on the wavelength of the light used raised the additional point of whether the nature of the photochemical process was sensitive to the energy of the absorbed light quantum in some general way or whether there was some qualitative differ- ence between absorption in the wavelength region of an electron-transfer-type band and in the region of a ligand-field band.It was felt that some light would be shed on this and other questions if a series of closely related complexes were studied in more detail. The acidopentamine family, being perhaps the best known, was selected for this purpose with two members which showed an intermediate type of behaviour, namely, Co(NH&(SCN)2+ and Co(NH3)5Br2+ selected for special emphasis. EXPERIMENTAL MATERIALS [Co(NH3)s(SCN)][C10412 was prepared as described previously ; 1 similarly, aquo-, bromo-, cnrbonato-, nitro-, and fluoropentarnine nitrates were obtained by standard I63164 PHOTOCHEMICAL OXIDATION-REDUCTION REACTIONS procedures.2p3 [Co(NH3)5NJCl* was made according to Linhard and Flygare,4 and the sulphatopentamine according to Jorgensen.5 The preparations were carefully recrystal- lized and checked for purity by means of their spectra and tests for ionic impurities. Co60-labelled complexes were similarly obtained.APPARATUS The same general irradiation equipment was used as previously described,l except that a thermosiatted irradiation cell was employed so that temperatures are good to 0.1"C. Although the bolometer used provided an absolute measure of light intensity incident on it, there was considerable absorption by the multiple windows of the cell and some diver- gency of the relatively large area light beam.For this reason, absolute quantum yields are tied to calibrations by means of the uranyl oxalate 6 and ferric oxalate 7 systems, and the bolometer used for relative light intensities. PROCEDURES All solutions were acidified to a pH of ca. 4 with acetic acid (except the carbonato pentamine). The determination of the amounts of Co(I1) and of Co(NH3)5(H20)3+ formed involved the use of Wo-labelled complexes. The procedure was essentially that previously described 1 with the following variation, for the bromopentamine complex. Briefly, inactive aquopentamine complex was added to the solution of irradiated Co60(NH&BrZ+, and the latter precipitated as the chloride salt ; the precipitate was partially redissolved and then reprecipitated.Finally, the aquo complex remaining in each supernatant was thrown down as the perchlorate salt. The average of a number of trials, including some tests with labelled aquopentamine complex, showed that by the procedure 92 % of the photolytic aquopentamine was obtained, with 5 % contamination by bromopentamine. Similarly, in the other fraction, 92 % of the original bromo- pentamine was recovered, with 4.5 % contamination by aquopentamine. The amount of free bromide ion formed was determined by conductimetric titration. With Co(NH3)5NO$ +, Co(I1) was determined by direct precipitation of cobaltous hydroxide and then measuring the optical density at 660 mp of the solution in concentrated hydro- chloric acid.Aquopentamine formation was estimated from the change in optical density at 330 mp. Since Co(NH&N$+ showed no photoaquation, its photolysis was followed by the reduction in optical density at 520 mp. The widths of the optical windows of the filters used, which were either glass or inter- ference in type, are indicated in the section on results. At the longer wavelengths, double filters were sometimes employcd to determine that the very low quantum yields observed were not partly due to leakage of shorter wavelength light. RESULTS A general summary of our findings on the photolytic behaviour of the Co(II1) acido- pentamine family is given in table 1, along with the available data on thermal aquation rates. Some of the more detailed results for Co(NH3)5(SCN)2f are shown in fig.1 in which the quantum yield for SCN- ion formation is plotted against wavelength. It should be noted, however, that the amount of S C N - produced was between that of the aquopentamine formed and the sum of this and the amount of Co(I1). That is, while aquation always produced SCN- ion, redox decomposition led to a mixture of oxidation products so that part of the time SCN- ion was still liberated. With this reservation in mind, the quantum yield showed little or no temperature dependence beyond 450mp, and a small one corresponding to a few kilocal in the region 370 mp to 410 mp. Quali- tatively, the ratio of Co(1I) to aquopentamine complex produced by photolysis decreased from about 2.1 at 370 mp to 0.24 at 550 mp. The results are more complete and more accurate for Co(NH&Br2+.As shown in table 2 both the partial quantum yields and the ratio of redox to aquation reaction de- creased markedly with increasing wavelength, as with the thiocyanatopentamine. Here, however, the amount of Br- ion produced very nearly equalled the sum of aquopentamine and Co(II) formed. Evidently, any bromine formed by the redox mode of photolysis was reduced by the ammonia liberated. The solutions did show some post-irradiation oxidizing ability towards iodine, but the nature of the species involved was not identified.A . W. ADAMSON 165 350 390 430 470 510 550 A mtL FIG. 1 .-Photolysis of Co(NH3)5(SCN)2+. Upper section : absorption spectrum ; lower section : quantum yield for SCN- production (0.01 M solution, 25"C, pH 4).The temperature dependence of the partial quantum yields was mixed in nature. It was always positive for &&,x but for 4aq. it was negative around 370 mp, slightly positive around 450 mp and possibly slightly negative at the longest wavelengths. Apparent activation energies ranged up to about 10 kcal/mole. TABLE 1 .-PHOTOLYSIS OF Co(III) ACIDOPENTAMINE COMPLEXES nature of X in CO(NH,)~X"+ NH3 so:- F- Cl- BI- NO, SCN- NF I- 0.01 M complex, 25"C, pH 4 quantum yield and nature of the reaction 370 mp 550 mp (340-400) (550 plus) <2 x 10-3 <2 x 10-3 ~ a . 10-3 0.011 100 %A 1.5 x 10-3 100 %A 0.21 50 %A 1-3 x 10-3 100 % A small 9 tr. R tr. R 50 %R tr. R 1.0 35 %A 65%R 0.045 30 %A 6 . 7 ~ 10-4 80 % A 70 %R 20 %R 0-44 <1%A 0.01 1 <1 % A 100 %R l00%R 066) <1 %A 0.10 (1 %A l00%R l00%R thermal aquation rate k (min-1) 7 x (ref.(8)) 4 x 10-6 (ref. (9)) 1 x 10-4 (ref. (8)) 4 X 10-4 (ref. (8)) <1 x 10-6 2 x 10-7 (ref. (8)) < 1 x 10-6 5 x 10-4 (ref. (8)) Nox-AIand R denote aquation and redox decomposition, respectively.I 66 1'1-1 OTO C li E M 1 C A 1, 0 X I I) A T 1 0 N -I< 8 I 1 U C'I'IO N I< Is A C ' I I 0 N S Soiiie qualitative observations were niadc on the oxidizing ability of solutions undcr irradiation. Irradiation of 0.01 M Co(NII3)5So4+, Co(NH&Br2+ and Co(NH3)5N$+ in 0.1 M KI led to definite iodine production. On thc other hand, no bromine resulted if this last complex were irradiated in the presence of Br- ion, which brackets the oxidizing ability of the intermediate involved, prcstimably N3 radical.It should also be noted that it was important that the irradiated solutions remain acid ; othcrwise the ammonia liber- ated on redox decomposition Icd to precipitation of the cobalt. This, in turn, can change the entire course of the photolysis. Thus, if solutions of the azidopentamine complex are allowed to become alkaline under irradiation, Co(OH)3 and azidc ion are produced,4 rather than only nitrogen, as was observed here. wavelength ( m d 3 70 410 (405-425) 450 520 550 (530-650) 590 (340-400) (445-455) (470-580) (580-650) TABLE 2.-PHOTOLYSIS OF Co(NM3)sBrz+ 0.01 M complex, pH 4 temp. quantum yields ("C) Br- aquation CO(l1) 15 0.2 1 0.12 0.1 1 25 0.2 1 0.067 0.15 15 0.065 0.053 0.036 25 0.1 3 0.069 0.07 1 25 0.040 0.027 8.0 x 10-3 15 2.0 x 10-3 2.0 x 10-3 ca.10-4 25 2.2 x 10-3 OX 10-3 ca. 10-4 15 1.5 x 10-3 1.5 x 10-3 4 0 - 5 25 1.4 x 10-3 1.4 x 10-3 15 1.3 x 10-3 1-3 x 10-3 < 10-5 25 POX 10-3 POX 10-3 ratio redoxlaq uat. 0.95 2.2 0.67 1.0 0.30 ca. 0.05 ca. 0.05 < 0.006 < 0.0 1 NOn.-For wavelengths beyond 500 mp, the aquation yield was taken to be the same as that for Br-. DISCUSSION The acidopentamines in table 1 are listed approximately according to increasing photosensitivity, which turns out also to be the approximate order of increasing proportion of photorcdox as opposed to photoaquation reaction. The sequence, furthermore, correlates well with the ease of oxidation of the acido group and hence with the expected thermodynamic instability of the complexes towards redox decomposition. On the other hand, there is no correlation at all with the instability towards aquation, as measured by the thermal aquation rate constants.For the acidopentamines, then, the oxidation potential of the acido group plays the dominant role in determining the photolytic behaviour of a complex. Next in importance appears to be the energy of the light quantum involved, as is evident from fig. 1 and 2 as well as from the data of table 1. The operation of both these factors is consistent with the mechanism proposed earlier 1 in which it was proposed that the first chemical consequence of absorption of a light quantum is a homolytic fission of the ligand-metal bond. However, unless con- siderable excess energy is available, the net quantum yield would be decreased by failure of the nascent radicals t o escape their solvent cage and consequent reforma- tion of the original complex.Thus, high quantum yields are predicted to be favoured by use of short wavelength light and by choice of easily oxidized ligands. The negative temperature coefficient for (baq. for Co(NH3)5Br2+ indicates that aquation is competing with some other process at 370 mp, which is consistent with the further steps postulated in the above mechanism. The same is true of the observation that the ratio +rdox/+aq, decreases steadily with increasing wavelength.A . W . ADAMSON 167 Briefly, it was supposed that net redox reaction occurs only if the homolytic fission products are formed with sufficient energy to escape completely, avoiding both reformation of the original complex (which affects the total quantum yield) and back electron transfer (which leads to aquation as the net process).1 I 1 I I I 350 390 430 470 510 5 5 0 5 9 0 A, mtL FIG. 2.-Photolysis of Co(NH&Br2+. Upper section : absorption spectrum ; lower section : quantum yield for Co(I1) (curve A), for aquation (curve B), and total quantum yield (curve C). The principal qualitative features are thus all explainable in terms of an essenti- ally chemical mechanism which assigns no role to the fact that the nature of the excited state involved is different at different wavelengths. It is surely significant, however, that the rapid drop in &dox for both Co(NH3)5(SCN)2+ and Co(NH3)5Br2+ occurs in the wavelength region between the near ultra-violet electron transfer type band and the ligand field band in the visible, and that, for the latter, q5aq. levels off at long wavelengths.Finally, it might be noted that the nitropentamine complex differs from the others in that a one-electron oxidation of the acido group yields a stable radical. Preliminary values for the quantum yields seem unusually high and the system warrants further study. A more de- tailed investigation of the nature of the oxidizing intermediates formed during irradiation should also be of interest. Turning briefly to complexes of other transition metals, it is possible to account qualitatively for the variations in photosensitivity towards redox reactions by supposing that, as for Co(II1) complexes, it is necessary that the product of a homolytic bond fission be relatively stable towards back electron transfer.Thus the fact that Cr(II1) complexes invariably give photosubstitution rather than photo- redox reactions 1 7 1 0 , ~ is consistent with the much greater reducing potential of Cr(I1) as compared to Co(I1). Similarly, Fe(I1) species show only photosub- stitution 12 while Fe(II1) complexes tend to undergo photolytic redox decomposi- tion.7~ 1 3 ~ 1 4 Pt(I1) complexes such as PtCIi- and cis-Pt(NH3)2Cl;! give only168 PHOTOCHEMICAL OXIDATION-REDUCTION REACTIONS photoaquation (with high quantum yield),'s while Pt(1V) compounds such as PtC1;- and PtBri- undergo homolytic bond fission on irradiation.1, 16 Clearly, however, there is room for much more information on the photolytic behaviour of transition metal complexes.These investigations have been supported in part by the U.S. Atomic Energy Commission. The author, in addition, wishes to acknowledge the assistance of Mr. P. Simpson and Mr. D. Ben Hur on the measurements with Co(NH&Br2-k and CO(NH~)~(SCN)~-~, respectively. 1 Adamson and Sporer, J. Amer. Chem. SOC., 1958, 80, 3865 ; see also Abstr. (136th meeting of the American Chemical Society, Atlantic City, September, 1959). 2 Fernelius, Znorganic Synthesis (McGraw-Hill, New York, 1921), vol. 1, p. 186. Bailor, Inorganic Synthesis (McGraw-Hill, New York, 1921), vol. 4, p. 171. 3 Jorgensen, 2. anorg. Chem., 1894, 5, 172. 4 Linhard and Flygare, 2. anorg. Chem., 1950,262, 340. 5 Jorgensen, J. prakt. Chem., 1865 (2), 31, 268. 6Leighton and Forbes, J. Amer. Chem. SOC., 1930, 52, 3146. Forbes and Heidt, 7 Hatchard and Parker, Proc. Roy. Suc. A , 1956,235, 518. 8 Basolo and Pearson, Mechanisms 0.f Inorganic Reactions (Wiley and Sons, New 9 Values estimated from Linhard and Weigel, 2. anorg. Chem., 1951, 266, 4. 10 Adamson, J. Inorg. Nuclear Chem., in press. 11 Edelson and Plane, J. Physic. Chem., 1959, 63, 327. 12 Emschwiller and Legros, Compt. rend., 1954,239, 1941. 13 Adamson, Welker and Volpe, J. Amer. Chem. Soc., 1950,72,4030. 14 MacDiarmid and Hall, J. Amer. Chem. SOC., 1954, 76, 4222. 15 unpublished work in this laboratory by Greenstadt and by Reddi. 16 Taube, J. Amer. Chem. Suc., 1954, 76, 2609. J. Amer. Cheni. SOC., 1934, 56, 2362. York, 1958), p. 122.

ISSN:0366-9033    

DOI:10.1039/DF9602900163

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

GENERAL DISCUSSION Dr. E. W. Duck (Koninklijke fShell Laboratorium, Amsterdam) said: Experience with emulsion polymerization systems would indicate that one of the equatioiis given by Prof. Wilmarth is unlikely to occur. In emulsion polymerization systems the thermal decomposition of the persulphate ion at -20°C according to is frequently used as a source of free radicals for initiating polymerization. Since these radicals are generated in the aqueous phase and then move through this to the locus of polymerization, viz., to the micelles which contain solubilized monomer, there would appear to be ample opportunity for the reaction between the SOi radicals and water to occur. If this occurred, the SO; radicals would be consumed and little or no polymerization would ensue. Work by Kolthoff, O’Connor and Hansen,l and Meehan, Kolthoff, Tamberg and Segal2 describes the polymeriz- ation initiation by the SO; radicals and the evidence for sulphate sulphur in the resultant polymer, This offers, therefore, independent support to Prof.Wilmarth’s conclusion that the concentration of SOX radicals is relatively large compared to the con- centration of OH. radicals. Prof. W. K. Wilmarth (University of Southern California) said : In reply to Dr. Duck, it should be emphasized that studies of the thermal 3 amd photolytic 4 decomposition of persulphate both provide convincing evidence that reaction ( 5 ) SO,+H,O+OH+HSO,. (5) does occur in aqueous solution. To understand the incorporation of the SO, radical in the polymer, it must be assumed that this radical reacts with the monomer at a low concentration much more rapidly than it does with the water solvent.This result suggests that the SO, radical may show a considerable specificity in its behaviour towards scavengers, with the nature of the specificity perhaps being influenced by the delocalization of the unpaired electron in the radical. As Dr. Grapy and Dr. Symons point out, the specificity displayed in the reactions of the SO, with water and with hydrogen peroxide does not parallel the overall energy changes in these reactions. This behaviour may be somewhat unexpected, but the kinetic evidence that SO; reacts preferentially with water, as outlined in the paper, seems to us to be unusually compelling. By varying the concentration of the reactants over an unusually large range, it was possible to establish three limiting forms of the rate law, a factor which imposes very severe restrictions upon the possible reaction mechanisms.Dr. P. Gray (University of Leek) said: Tsao and Wilmarth’s mechanism for the reaction between hydrogen peroxide and the peroxydisulphate ion requires that the singly-charged sulphate radical SO, reacts with water to form the hydroxyl radical but does not react with hydrogen peroxide to form the hydroperoxyl radical : SOz+HzO->SO:- +H+ +HO, (5) SOi + H202+SO:- + H+ + HOZ. (17) This equation is SOi+HzO-tHSO,+OH.. s,o: - -+2SOT 1 J. Polymer Sci., 1955, 15, 459. 2 J. Polymer Sci., 1957, 24, 215. 3 Kolthoff and Miller, J. Amer. Chem. Soc., 1951, 73, 3055. Tsao and Wilrnarth, J. Physic.Chem., 1959, 63, 346. 169170 GENERAL DISCUSSION Although favoured by the concentration ratios (about 50 : 1 in M solution), this situation is the reverse of that expected on energetic grounds. Existing thermochemical data 1 suggest that, in aqueous solution, the energy requirements of reaction (5) are at least 20 kcal mole-1 (actually, 27.4 f 6 kcal mole-1) greater than those of reaction (17). A difference as large as this in the thermochemistry would be expected to be reflected in the rate constants, and in the activation energies in particular, of the two reactions. Conversely, if (as Wilmarth and Tsao suggest) reaction (5) is favoured, then unusually large differences in steric factor are callcd for in compensation. Dr. M. C. R. Symons (Southampton University) said: It is surprising that the reaction : is unimportant whereas reactions (5) and (8) are thought to play a major role in the mechanism proposed.Reaction (18) is similar to the step SO,+ H202-+HSOz+ HO2 (1 8) except that it is far more favourable energetically, so, in order for reaction (5) to compete so effectively with attack on hydrogen peroxide (18), it must be extremely rapid. Prof. J. Weiss (University of Durham) said: I believe that there is a good deal of experimental evidence to show that the oxidation of the As(II1) to As(V) goes by a univalent change via the tetravalent state of arsenic. This may be in- ferred also from the fact that many oxidation reactions of arsenite (similar to that of sulphite) are oxidation chain reactions of considerable chain length.This is particularly true also in alkaline solutions when the reacting species are the arsenite anions. The occurrence of chain reactions is a clear indication of univalent oxidation steps (single electron transfer) and, therefore, of the formation of the intermediate tetravalent compound of arsenic. Dr. W. A. Waters (Oxford University) (partly communicated) : The postulation by Csiyni of a transient form of AsIV merits attention, for there is clear experi- mental evidence 2 that the highly active l-electron abstracting agent Mn3+ cannot oxidize As"' to AsV under conditions in which 2-electron transfer by the very mild oxidant iodine is rapid. CeIV and probably Vv also do not appear to oxidize As111 by direct 1-electron transfer. Again (MnO4)z- seems to attack by 2- electron and not by 1-electron removal 3 though this anion can, and often does, react by both routes. Probably the tautomeric arsenite system OH OH I I I I (a) :As--OH+H-AS=O (b) OH OH oxidizes most easily by hydrogen abstraction from the favoured form (b), hetero- lysis such as c I-I+H-AsO(OH)~-+I-+ I-H++AsO(OH)~ b requiring very little activation energy.atoms out of other molecules, e.g., On the other hand, very reactive radicals that are capable of picking hydrogen H0.f H-AsO(OH)~+HO-H+ .AsO(OH)2, 1 Gray, Trans. Faraday SOC., 1959, 55, 408. 2Land and Waters, J. Chem. SOC., 1958,2129. 3 Pode and Waters, J. Chem. SOC., 1956, 717.G EN E R A L I) IS C US S I 0 N 171 could react homolytically in a way that has no exact equivalence for ions such as Mn3f which merely gain electrons, and this may be the implication of Cshnyi's results with *SO';.This route of homolytic hydrogen transfer has been estab- lished for radical attack on both phosphite and hypophosphite ions 1 but little as yet is known of the free radical chemistry of arsenic. Prof. F. S. Dainton (Leeds University) said: The HO2 radical is frequently regarded as a 1-equivalent oxidant (in its unionized form) or as a 1-equivalent reductant (in the ionized form). In principle at least, it may also be envisaged as a 2-equivalent oxidant by donation of an oxygen atom to a reductant. The oxida- tion of sulphite ion by alkaline hydrogen peroxide may in fact involve as one of the participating reactions 2-equivalent reactions of the 0-atom transfer variety such as HO~+SO;-+HO-+SO:- and HO, + SO: -+HO+ SO:-.If one admits the possibility of reactions of this kind, but with Asr1T replacing SOz-,e.g., HO~+AS"'+HO+AS~, then the effect of arsenious acid on the peroxide-persulphate system, referred to by Prof. CsAnyi may perhaps be accounted for without invoking the participation of AsIv to which Dr. Waters has raised objections. Dr. L. J. Cshyi (Szeged University) said : In reply to Prof. Dainton, in principle it may be allowed that H02 can react as a 2-equivalent oxidant : HOz->HO+ 0. However, this reaction makes it impossible to interpret the effect of arsenous acid on the induced reaction between H202 and H2S208. According to Prof. Dainton's explanation, in the presence of arsenous acid a closed chain reaction occurs, in which hydrogen peroxide and arsenous acid are consumed in a mole ratio of 1 : 2.On the contrary, we have found that in the induced reaction arsenous acid and peroxydisulphate disappeared in a ratio of 1 : 1 and the H202 error is practically zero. Therefore we persist in our explanation, i.e. arsenous acid reacts according to steps (7) and (8). Concerning Dr. Waters's objection, we consider the existence of the AsN entity to be probable on the basis of the induced reaction of chlorate ions, too ; chlorate ions are reduced by arsenous acid when it is oxidized by 1-equivalent reagents.2 Dr. C. F. Wells (British Rayon Res. Assoc.) said: In connection with the paper on the Fe(m)+H202 system it is perhaps pertinent to mention the results of our study3 of a similar system, Co(III)+H202, at O'C, 125°C and 25°C in dilute solutions.We found spectrophotometric evidence for the existence of a Co(1II)HO; complex at 0°C and 12.5"C. Although under certain conditions, over 90 % of the total Co(1II) was present as Co(III)HO,, from the kinetics we were not able to distinguish between any, or all three, of the following mechanisms : CO(III)+HO~-+CO(II)+HO~* (1) Co(III)+ HO~+CO(III)HO,-SO(II)+ HO2. (2) Co(1II)OH-+ H202-+Co(II)H20+ H02.. (3) As this reaction is extremely rapid, it is unlikely that the reaction can be followed at the higher concentrations of H202 used by Haggett, Jones and Wynne-Jones, unless a rapid reaction technique is used. 1 Nonhebel and Waters, Proc. Roy. SUC. A, 1957, 242, 16. 2 Talanta, 1958, 1, 359.3 Baxendale and Wells, Trans. F a r e Suc., 1957,53, 800.172 GENERAL DISCUSSION Dr. P. Jones and Prof. W. F. K. Wynne-Jones (King's Colloge, Newcastle-upon- Tyne) (communicated) : Kremer and Stein have recently re-investigated the Fe3++ H202 reaction in dilute aqueous solution using a combined kinetic and spectro- photometric approach.1 In their more concentrated solutions (up to 0.2 M H202), they find that the reaction velocity is controlled by the concentration of a mono- peroxy complex, essentially the same as the one we have described, and they have more recently shown that the velocity constants derived from their investigation are in excellent agreement with our own? At lower concentrations they find evidence for a spectroscopically distinct secondary mono-peroxy complex and have made suggestions as to its nature and its role in the reaction.This work is complementary to our own, since Kremer and Stein have chosen conditions under which further detail of the breakdown of at least one of the reactions intermediates can be studied. It would be of interest to test further the comparison of iron catalysis with chromate catalysis by a similar study of the latter system. Dr. M. C. R. Symons (Southampton University) said: The following observa- tions may be of interest with reference to the result of Haggett, Jones and Wynne- Jones that Cr(V1) is reduced to Cr(II1) by hydrogen peroxide in acid solutions but that in alkaline solution catalytic decomposition ensues with no observable change in the concentration of Cr(V1). When peroxide is added to sodium manganate in concentrated aqueous sodium hydroxide, hypomanganate is rapidly formed, after which catalytic decomposition of peroxide occurs.If, however, a con- centrated solution of potassium hydroxide is used as solvent, manganate is con- verted into hypomanganate at room temperature, but the reverse occurs at elevated temperatures, the reaction in each case being followed by catalytic decomposition of peroxide. The relevant equations are 3 and 2 MnOi'+HO,+ OH'+2 MnO:-+H,O+ 0, 2 MnO:-+HO;+H,O+2 MnOi-+30H'. Dr. R. J. P . Williams (Oxford University) said: Prof. Longuet-Higgins has pointed out that two light absorption processes can occur in the region of the spectrum over which Prof. Adamson is studying photochemical reactions.The spectra of the cobaltic complexes under study are of immediate interest. In the table we quote some relevant data. The overlap between the ligand field band and the charge transfer band increases in an obvious sequence. The sequence is maximum of ligand approximate wavelength X in Co(NH&X field bands where charge transfer band WPc) extinction coef. is 100 NH3 476 339 270 F- 516 355 270 so: - 515 355 270 Cl- 535 364 310 new band B r 550 ( ? ) 380 7 , NO; 460 (326) 330 Y Y SCN- 497 (360) 350 Y. 520 ( ? ) 420 Y9 583 (383) 470 7 , *i Figures in parentheses indicate that the charge transfer (new) band overlaps the second ligand field band. very similar to that of the relative importance of photo-dissociation and photo- redox reaction in table 1 of Prof.Adamson's paper. It would appear that the two photochemical processes are substantially independent in most of these com- plexes. I- 1 Kremer and Stein, Trans. Faraday SOC., 1959,55, 959. 2 Kremer and Stein, 2nd Ini. Cong. Catalysis, (Paris, July, 1960). 3 Camngton and Symons, J, Chem. SOC., 1956, 3373.GENERAL DISCUSSION 173 Prof. A. W. Adamson (University of Southern Califarnia) (communicated) : In response to Dr. R. J. P. Williams, may I say that it is perfectly possible to suppose that the processes occurring (i.e. photoaquation and photoredox decomposition) when irradiation involves an electron transfer and have a common mechanism, but one which is different from that involved when photoaquation alone, is occurring on irradiation of a purely ligand-field-type band.While this possibility was intimated in paragraph four of the Discussion section, the explicit proposal was not advanced since it seemed an as yet unnecessary complication, although not an improbable one. Current studies on the photolysis of cis and trans diacidobisethylenediamine cobalt (111) complexes may reveal differences in the nature of the photo-substitution reactions occurring when electron-transfer and ligand-field regions of the absorption spectrum are irradiated. The thought described by Prof. Longuet-Higgins was the somewhat different one that photoaquation occurred at 370 mp (e.g. with Co(NH&Br2+) due either to direct absorption into a ligand-field band hidden by the more intense electron- transfer band, or by conversion from the latter excited state to the former. Photo- aquation is thus imagined as the specific consequence of reaching a ligand-field excited state, and photoredox decomposition, of reaching (and remaining in) an electron-transfer excited state.The difficulties with this picture include the following. First, quantum yields for some one type of process tend not to vary greatly with wavelength (e.g. photo- substitution with Cr(II1) complexes and photoaquation with Co(a)51-3, where the situation is not complicated by redox decompositions). One would therefore expect that since the quantum yield for photoaquation with Co(NH&Br+2 is low at 550 mp, it should also be low at 370 mp, and, in fact, lower to the extent that conversiop from the electron-transfer excited state was incomplete. In other words, it requires assumptions contrary to existing experience if one is to account for the actually very high photoaquation quantum yield at 370 mp. Secondly, for this complex, photoaquation showed a negative temperature coefficient at 370 mp, which seems to require an activated return conversion from the ligand-field excited state to the electron-transfer one. This seems improbable in view of the short life-times of these excited states. It is difficult to rule out such a set of assumptions on any rigorous grounds, but they seem improbable and not to be adapted until and if necessary. Dr. M. C . R. Symons (Southampton University) said: It is possible that photo- lysis of some of the compounds studied by Adamson in a rigid medium at low temperature would yield informative results. Use of this technique has, for example, shown fairly conclusively that photolysis of permanganate with 3650 A light follows the equation MnO,+hv+MnO;+ 02, oxygen being " extruded " in one step.1 chromates,2 Another example of co-operation breakdown is the photolysis of monoalkyl- R'C'H ' o'CrO,+hv-+RRICO+HCrO;, R1/ \ 0 / It seems possible that other complex compounds may photolyze by similar co- operative mechanisms under some circumstances. 1 Klaning and Symons, J. Chem. SOC., 1959, 3269. 2 Klaning and Symons, J. Chem. SOC., 1960, 977.

ISSN:0366-9033    

DOI:10.1039/DF9602900169

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

II(b). OXIDATION-REDUCTION REACTIONS INVOLVING ORGANIC SUBSTRATES OXIDATION OF FORMIC ACID AND FORMATE ION IN AQUEOUS SOLUTION BY SOME INORGANIC OXIDANTS BY J. HALPERN AND SANDRA M. TAYLOR Dept. of Chemistry, University of British Columbia, Vancouver, Canada Received 8th February, 1960 Kinetic data are reported for the oxidation of formic acid by MnO;, Hg2 I., Hg$+ and TP+ in aqueous pcrchloric acid solution. In each case the predominant reaction path (identified with the reaction of formate ion with the oxidant) is first order in each of the reactants, inversely first order in H i, and exhibits a large deuterium isotope effect. For the Mn04 reaction the rate-determining step probably involves hydride ion transfer from HCOO- and for the other reactions a two-equivalent reduction of the metal ion by electron transfer from a co-ordinated HCOO- ion, coupled with proton transfer to a water molecule.With Mn04 there is also an H+-independent path corresponding to reaction of HCOOH. The absence of an isotope effect for this path suggests that it proceeds by a different mechanism. This paper describes kinetic studies of the oxidation of formic acid and formate ion in aqueous solution by several inorganic oxidants, Mn04, Hg2+, Hg$+, and TP+. Particular consideration is given to the following points which also arise in connection with many other oxidation-reduction reactions : (i) the Hf-dependence of the reaction reflecting, in this case, differences in reactivity between HCOOH and HCOO-, (ii) whether the overall oxidation of formic acid, which involves a net two-equivalent change, occurs in a single step or through successive one- equivalent steps, and (iii) the various alternative possibilities of electron and group transfer redox mechanisms.Insight into the last point has been sought par- ticularly through measurement of the deuterium kinetic isotope effects. EXPERIMENTAL A.R. formic acid was distilled twice before use. A stock solution of mercuric per- chlorate was prepared by dissolving mercuric oxide (twice reprecipitated from perchloric acid solution with sodium hydroxide) in dilute perchloric acid. Part of this solution was reduced with mercury to give a solution of mercurous perchlorate. Sodium perchlorate was recrystallized from aqueous alcohol and dried at 120". The preparations of deutero- formic acid, deutero-perchloric acid and thallic perchlorate have been previously de- scribed.lp2 All other chemicals were of reagent grade and the water was distilled twice from alkaline permanganate.The procedures used for the kinetic measurements on Mn03 and Tl3+ have been previously described.ls2 To measure the rates of oxidation by Hg2+ and Hg3+, samples of the solution were withdrawn periodically and analyzed for Hg$+ by adding an excess of iodine and back-titrating with thiosulphate. RESULTS A N D DISCUSSION GENERAL The products of oxidation of formic acid were, in each case, C02 and Hf. The stoichiometries of the reactions examined are thus : 1 74J . HALPERN A N D S . M . TAYLOR 175 2Mn0, + 3HC0OH+2H"-+2MnO2 +3C02 +4H,O, (1) 2Hg2++ HCOOH-+Hgg+ + C02 +2H+, (2) Hg$++HCOOH-+2Hg+C02+2H+, (3) T13+ + HCOOH+Tl++ C 0 2 +2H +.(4) All the kinetic measurements were made in perchloric acid solution, complica- tions due to hydrolysis and complexing of the metal ion oxidants thereby being minimized. In every case the rate showed an inverse dependence on the H+ concentration which could be accommodated by a general rate-law of the form, In acid solutions where HCOOH is largely unionized this becomes -d[HCOOH]/dt = k,[HCOOM][Ox] + k,[HCOO-][Ox]. (5) -d[HCOOH]/dl =(k,+k,KJ[H*])[HCOOH],,,COxl (6) = k'[HCOOH],o,[Ox], (7) where Ki is the ionization constant of formic acid. The measurements yielded 0 5 10 l/{H+], M-1 FIG. 1.-Permanganate oxidation of formic acid at 29*9", p = 1.1. (l), HCOOH; (2), DCOOH. linear plots of k' against l/[H+] for each system (fig.1 and 2) from whose inter- cepts and slopes kl and k2Ki were determined. For Hg2+, Hg$+ and TP, kl was immeasurably small, essentially all the reaction being due to HCOO-.* Ki is * The identification of the H+-dependent path in each case with a rate-determining step involving HCOO- is convenient for purposes of comparison but not necessarily of mechanistic significance. The rate-law defines only the composition of the activated complex and does not distinguish between reaction paths involving different reactant species in rapid equilibrium with each other, e.g. between HCOO-+TP+ and HCOOHf TIOH*+. Tie rate-constants corresponding to different choices of reactants are related through the appropriate equilibrium constants.176 OXIDATION OF FORMIC ACID not known for all the temperatures and ionic strengths and the values employed throughout (5.6 x 10-4 for HCOOH and 8.0 x 10-4 for DCOOH) were those deter- mined at 30" and p = 1.The values of k2 for the other temperatures and ionic strengths and the corresponding values of AH" and AS+' are thus subject to some error from this source but this is estimated to be small. Thus, available data 3 indicate that the correction to AH+ due to the heat of ionization of formic acid is about 1 kcaI/mole. The compensating error in AS+ is Iess than f5 cal/moIe deg. The kinetic results are summarized in table 1 which also includes data of Bawn and White 4 for the oxidation of formic acid by Co". l/[H+], M-1 FIG. 2.-Oxidation of formic acid by metal ions.(l)? Tl3* (75~2"~ p = 4.0) ; (2), Hg2+ (46-5", p = 2.0) ; (3), Hgzf (65~1"~ p 2.0). PERMANGANATE ION The results for this system have been reported in detail.1 The important (i) both HCOOH and HCOO, apparently react with MnOi but the difference in reactivity is such (k2/kl-2000 at 25") that the contribution of the former is important only at very high acidities and was undetected in various earlier investigations.5~ 6 (ii) Differences in both AH+ and ASs contribute to the higher reactivity of HCOO- and both are in the opposite direction to that expected from purely electrostatic considerations (i.e. the favoured reaction is that between ions of opposite charge). (iii) The HCOO- reaction exhibits a large deuterium isotope effect (k~c00-l ~DCOO- = 7) which is absent for the HCOOH reaction (fig.1 and table 1). A solvent isotope effect is also observed for k2 (kpO/kpO = 0.38) but not for k1.1 (iv) The reaction is catalyzed by Fe3f, the catalytic effect approaching a limit- ing value when the Fe3+ concentration exceeds about 0.02 M. features are :J . HALPERN AND S. M. TAYLOR 177 The first three of these observations strongly suggest that the HCOOH and HCOO- reactions proceed by quite different mechanisms. Earlier workers 5 s have concluded that the rate-determining step in the permanganate oxidation of formate ion involves a two-equivalent reduction of Mn0; leading to the formation of Mn(V) which disproportionates rapidly to Mn02. This is strongly favoured on thermodynamic grounds, HCOO- + MnO, +C02 + H+ + MnOZ- (AGO = - 32 kcal ; AH" = - 56 kcal), whereas the alternative possibility of an initial one-equivalent reduction, HCOO- + MnO, -+C02 +H+MnO:- (9) (AGO = +27 kcal; AH" = +22 kcal), is much less favourable and difficult to reconcile with the observed AH+ and AG+.A one-equivalent step leading to the formation of HC02*, instead of He, is unlikely to be more favourable since dissociation reactions of the type, RC02*-+R*+C02, are generally exothermic.7 The latter reflects the high stability of C02 and also suggests that a likely path for the oxidation of HCOO- is by transfer of a hydride ion, HCOO-+MnO~-+CO,+HMnO~-. ( 0) This is consistent with the large DCOO- isotope effect for this path but requires some qualification in the light of Wiberg and Stewart's 6 observation, using 0 1 8 - labelled MnOz, that an appreciable fraction of the oxygen in the C02 product ( 4 .4 atom per C02 molecule) is derived from Mn04. This may reflect an addi- tional path or, more likely, that hydride transfer occurs through an intermediate configuration of the system in which oxygen exchange between Mn0; and HCOO- is possible. Wiberg and Stewart have considered several such possibilities but the available evidence does not permit an unequivocal interpretation. The D20 iso- tope effect for this path also remains to be explained. The absence of corresponding deuterium isotope effects in the HCOOH re- action suggests that the rate-determining step in this case does not involve ap- preciable weakening of either the G H or 0-H bonds. The most plausible mechanism consistent with this would appear to involve a rate-determining attack of M n 0 ~ at the C=O position of formic acid leading to an intermediate of the type OH H ' 0-C-0-MnO, which may either decompose directly to H2CO3 and Mn03 or undergo hydrolysis to H2CO3 and MnOi-.Two factors are probably relevant to this difference in mechanism between HCOO- and HCOOH, (i) the stabilization of the activated state by the C02 structure provides a driving force for hydride elimination from HCOO- but not from HCOOH, and (ii) the negative charge of HCOO- also favours hydride elimin- ation but makes it a less favourable case than HCOOH for nucleophilic attack by Mn04. The two types of mechanisms invoked above find pertinent analogies in other permanganate oxidation reactions.Thus Stewart 8 has shown that the oxidation of benzhydrolate, (C~HS)~CHO-, to benzophenone probably proceeds by hydride ion transfer to permanganate and has drawn attention to the striking similarities178 OXIDATION OF FORMIC ACID (kinetics, isotope effect, etc.) between this reaction and the oxidation of formate ion. In this case also the negative charge and the stabilization of the activated state by the conjugation energy of the benzophenone structure are likely to favour hydride elimination. On the other hand, the acid-catalyzed permanganate oxid- ation of aromatic aldehydes 9 probably proceeds through a mechanism analogous to that proposed for HCOOH. This involves attack of MnOz at the carbon atom of the conjugate acid of the aldehyde with the formation of the intermediate OH C,H,-C-OMnO, H whose subsequent decomposition by proton transfer to a base yields C6H5COOH and Mn(V).In accord with this it was found that most of the excess oxygen in the benzoic acid product originates from permanganate rather than from the solvent. In contrast to the oxidation of formic acid, however, this reaction ex- hibits a large deuterium isotope effect suggesting that decomposition of the inter- mediate rather than its formation is rate-determining. The catalysis of the permanganate oxidation of formic acid has been interpreted 1 in terms of a mechanism involving concerted oxidation by the two species, HCOO ' + MnO, + Fe3+-+H' + CO, + MnOi- + Fe2+ (AGO = -42 kcal; AHo = -42 Itcal), leading to the formation of Mn(VI), rather than Mn(V), as the initial reduction product.This is analogous to the catalytic effect of Ag+ in the permanganate oxidation of H2. The approach of the catalytic effect to a limiting value with increasing Fe3+ concentration can be explained by complexing between Fe3+ and MnOz but direct codirmation of this is lacking. Hgzf, Hg3+ AND Tl3+ OXIDATIONS In view of their kinetic similarity and of the likelihood that they proceed by similar mechanisms these three reactions will be considered together. The oxidations of formic acid by Hg2+ and Hgz+ have previously been examined by Topham and White 11 in nitric acid solution. The present measurements were made in perchloric acid solutions and extended to higher acidities to avoid com- plications due to hydrolysis of the metal ions. Where they overlap the results are in good agreement with the earlier ones. In contrast to those for MnOa the plots of k' against l/[Ns] (fig.2) pass through the origin thus indicating the absence of an H+-independent path due to HCOOH. The rates for both the Hg2+ and Hgz+ reactions showed an inverse dependence on the ionic strength, i.e. on the concentration of NaC104 at constant HC104. Thus k' (1. mole-1 sec-1) for Hg2f at 46.5", 0.82 M HClO4, decreased progressively from 11.6 x 10-4 at 0.2 M NaCl04 to 6.5 x 10-4 at 5.3 M NaC104. The corres- ponding decrease of k' for Hgg+ (65.1", 0.5 M HC104) from 8.8 x 10-5 at 0.5 M NaC104 to 4.8 x 10-5 at 1.5 M NaC104 is somewhat more marked and may be due to complexing between Hg$+ and ClOz.12~13 The oxidation of formic acid by Tl3+ was originally investigated by Halvorson and Halpern2 although the H+-dependence of the reaction was not resolved.The kinetics are somewhat complicated both by complex formation between the reactants and by hydrolysis of TP+ and it is only at very low reactant concentra- tions and high acidities, where both these effects are negligible, that the overall kinetics assume the simple form of eqn. (6) and (7). We have extended the earlier measurements to the region of lower reactant concentrations ( 4 . 0 1 M) where to a first approximation the effects of complexing can be neglected. In this region the kinetics are approximately of first order inJ . HALPERN AND S . M. TAYLOR 179 each of the reactants and, as with Hg2+ and Hg$+, yield linear plots of k' against 1/[H+] which pass through the origin." The similarity of the deuterium isotope effects for the three reactions (k~cOO-/k~coo- = 3-4) also favours the conclusion that they proceed through similar mechanisms.In view of the favourable possibilities for co-ordination to the metal through the oxygen of formate ion, we consider the most likely mechanism for these reactions to be a two-equivalent reduction of the metal ion by electron transfer from a co- ordinated formate ion coupled with proton transfer to a water molecule, e.g. HCOO . ~1~ + +H+ + coZ +TI+. (13) The overall thermodynamics of such a process are very favourable not only for Tl3+ (AGO = -70 kcal) whcre the stable products are formed directly, but also for Hg2+ (AGO = -40 kcal) and Hg=+ (AGO = -28 kcal) where the initial reduction products are Hg atoms.Two further possible mechanisms might be considered. 1. HYDRIDE TRANSFER TO THE METAL IoN.-This is considered less likely than with MnO;i-, particularly in a structure in which the formate ion is co-ordinated to the metal through one of its oxygen atoms, but cannot be excluded. Hydride transfer has been demonstrated 16917 in the oxidation of H2 by certain metal ions, e-g., C U ~ + +H~+CUH+ +H+. (14) Both this and the previous mechanism are consistent with the deuterium isotope effects and with the absence of H+-independent paths for these reactions. OF H e , TI2+, IiTC., AND AN HCOO* RADICAL OR H ATOM.t-This has been sug- gested 11 in the case of Hg2+ and Hg$+ but, as with MnOz, seems very unlikely on thermodynamic grounds and difficult to reconcile with the observed AG* values.In this connection it is noteworthy that the reactivity of formic acid toward one-equivalent oxidants, e.g. Ag+ and Fe3-I- is much lower than toward two- equivalent ones, such as Hg2+, of comparable potential. This would be very sur- prising if reaction of the latter proceeded by one-equivalent steps. In the region of higher reactant concentrations, more extensive complexing between Tl3+ and HCOO- leads to modification of the apparent kinetics of this reaction.2 Thus, if the total TllI1 concentration is held constant at 0-02 M and the HCOOH coiicentration increased, the rate approaches a limiting value above about 0.5 M HCOOH. In this region the kinetics are of first order in TllI1 and zero order in HCOOH.Furthermore we have found that the first-order rate constant k in this limiting region is, unlike k', independent of the H+ concentration between 1 and 4M. The most obvious explanation of these observations, which makes the behaviour in the two regions kinetically equivalent, is in terms of a formate complex of Tl3+ (rather than the formic acid complex previously suggested 2 which is not readily reconciled with the different H+-dependence in the two limiting This involves a correction for the hydrolysis of TlJ+ and for this purpose we have estimated the hydrolysis constant from the known value for 25" (0.073 M at p = 3) 14 assuming a heat of hydrolysis of 10 kcal/mole ; most known heats of hydrolysis of similar metal ions are within 1-2 kcal of this value.15 $ Some evidence for the formation of free radicals is provided by the observation that addition of mercuric salts to formate solutions containing acrylonitrile causes precipita- tion of polymer.l* This may, however, reflect a minor reaction path or one involving acrylonitrile.2. A ONE-EQUIVALENT RATE-DETERMINING STEP LEADING TO THE FORMATION * k' is evaluated for a rate-law expressed in terms of the concentration of TW.180 OXIDATION OF FORMIC ACID regions) whose formation is essentially complete above 0.5 M HCOOH. Thus, the two rate-laws, -d[HCOOH]/dt = kz[HC00-][T13+] = k[HCOO. T12+] (15) are seen to be equivalent through the relation k2 = kK', (16) where K'(= [HCOO . T12+]/[HCOO-][T13+]) is the equilibrium constant of reaction (12) and k, the rate-constant of reaction (13).Values of k and k2 and of the corresponding heats and entropies of activation, derived from the measurements of Halvorson and Halpern 2 through these relations are listed in table 1. These measurements also yield a value of 9 x 104 for K' and values of A H 4 and AS-22 cal/mole deg. for reaction (12). Direct measurement of the deuterium isotope effect for k yielded a value of kHCoO-/kDCoO- = 2.7. Thus the overall AH* of the reaction and the deuterium isotope effect are largely reflected in k while the large positive overall AS+ is almost wholly accounted for by K', i.e. by the associ- ation of Ti3+ and HCOO-. While not conclusive evidence this is at least con- sistent with the suggestion that the activated complex involves a " co-ordinated " formate ion.The AS+ values for the other reactions also follow the order ex- pected for association between ions of the corresponding charge type. TABLE SU SUMMARY OF KINETIC DATA AH+ AS+ AG9 (25") kcal/mole cal/mole deg. kcal/mole oxidant temp. range, "C rate constant MnO, 15-35 Y9 15-35 Hg2+ 36-61 HgZ + 60-80 ~ 1 3 + 65-85 Y ¶ 65-85 CO"1 * 0-30 ¶ ¶ 0-30 16 12 20 21 26 26 26 21 - 19 21 - 14 16 -I- 3 19 0 21 -l- 21 20 - 1 26 + 19 21 $21 15 kH- C "D-C 1.0 (29.9") 7.0 (29.9") 3.4 (46.5") 3.9 (65.1") 3.4 (752") 2.7 (75.2") * from the data of Bawn and White.4 ONE-EQUIVALENT REDOX MECHANISMS Comparison of the potentials for the two reactions, HCOO--+H++C02+2e(Eo = 0.2 V), HCOO--+H+CO,+e(E"- -1.8 V), suggests that one-equivalent oxidation of formic acid or formate ion should be favoured only with very powerful one-electron oxidants.While unlikely for any of the oxidants we have examined this is probably the case for CoIII (C02++C03+ + e ; E" = - 1.8 V). The oxidation of formic acid by Con1 has been investigated by Bawn and White4 in sulphuric acid solution where the cobaltic ions are un- doubtedly present as sulphate complexes. The kinetics conform to eqn. (6) and while the [H+]-l-path again predominates, the contribution from the H+-inde- pendent path is measurable in this case (table 1). Because of the substitution inertness of CoIIr co-ordination of formate ion may not provide a favourable path for this reaction. An alternate possibility, also consistent with the [H+]-l-path being favoured, involves transfer of a hydrogen atom from formic acid to a water molecule or OH- ion in the co-ordination shell of the cobaltic ion,J . HALPERN A N D S . M. TAYLOR 181 We are grateful to Dr. L. E. Orgel for helpful discussion and to the NationaI Research Council of Canada for financial support. 1 Taylor and Halpern, J. Amer. Chem. SOC., 1959,81, 2933. 2 Halvorson and Halpern, J. Amer. Chem. SOC., 1956, 78, 5562. 3 Harned and Embrce, J. Amer. Chem. SOC., 1934, 56, 1042. 4 Bawn and White, J. Chem. SOC., 1951, 339. 5 Mann and Tompkins, Trans. Furday SOC., 1941, 37, 201. 6 Wiberg and Stewart, J. Amer. Chem. SOC., 1956,78, 1214. 7 Jaffk, Prosen and Szwarc, J. Chem. Physics, 1957,27,416. 8 Stewart, J. Amer. Chem. SOC., 1957, 79, 3057. 9 Wiberg and Stewart, J. Amer. Chem. SOC., 1955,77, 1786. 10 Webster and Halpern, Trans. Furaday Soc., 1957, 53, 51. 11 Topham and White, J. Chern. SOC., 1952, 105. 12 Hietanen and Sillkn, Arkiv Kemi, 1956, 10, 103. 13 Armstrong and Halpern, Can. J. Chem., 1957,35, 1020. 14 Biedermann, Arkiv Kemi, 1953,5,441. 15 Stability Constants, vol. 2, Chem. SOC. Special Publ., no. 7 (London, 1958). 16 Halpern, Macgregor and Peters, J. Physic. Chem., 1956, 60, 1455. 17 Webster and Halpern, J. Physic. Chem., 1957, 61, 1239. 18 Dainton and James, Trans. Faraduy SOC., 1958, 54, 649.

ISSN:0366-9033    

DOI:10.1039/DF9602900174

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

ELECTRON TRANSFER TO AROMATIC MOLECULES BY A. C. ATEN, J. DIELEMAN AND G. J. HOIJTINK Chemical Laboratory of the Free University, Amsterdam Received 8th February, 1960 A qualitative discussion is given of the mechanism of the electron transfer to aromatic molecules in solution and at electrode surfaces. The rate of electron transfer depends on whether the aromatic negative ions are free or associated with the positive ions to ion-pairs. In the latter case, the electron transfer is coupled with the " transfer " of the positive ion so that the overall process may be regarded as an atom transfer. The experi- mental data indicate that a re-orientation of the solvent molecules takes place after the electron transfer. During the last few years, electron-transfer processes have been studied by various investigators, particularly the isotopic exchange reactions and electron- transfer processes at electrode surfaces at the standard potential of the ox-red couple.Among these processes, the more simple ones are those in which the co-ordination of the participating ions remains unchanged. These reactions have in common that the activation entropy is markedly negative, e.g. for the isotopic exchange reaction : 1 one finds AS+ = -25 cal/mole deg. Although the various theories lead to a reasonable estimate of the rate constant and predict a large negative activation entropy, the authors differ about what is the rate-determining factor in these processes. According to Weiss,2 the electron transfer in ionizing solvents takes place by quantum-mechanical tunnelling of the electron through the compact solvent layers which separate the two ions in the collision complex.Re-orientation of the solvent molecules is assumed to occur after tunnelling, which seems to be a reason- able supposition, since the average velocity of the electron will be high compared with the frequency of tumbling of the solvent molecules. Eyring 3 and Marcus,4 on the other hand, assume that a reorientation of the solvent molecules should take place before the electron tunnels. Among the various possible conformations of the collision complex only those would be favourable for electron transfer in which the orientation of the solvent molecules around the two ions is such that the potential energy of the electron is the same at either side of the barrier.In an elegant treatment, Marcus derived general equations for the free enthalpy of such a non-equilibrium state with minimum energy. From these equations it follows that the electrostatic repulsion between the two ions gives the greatest contribution to the negative activation entropy. The electron tunnelling enters the expression for the rate constant as a trans- mission coefficient, which Marcus believes to be equal to about unity. Both Weiss and Marcus started from the simplified model in which the dielectric constant changes abruptly at the boundary between the compact solvent layer and the solution, taking the optical value for the dielectric saturated solvent layer and the static value for the surrounding solution. Recently, Laidler 5 has shown that this simplification may lead to serious errors for that part of the activation entropy due to the repulsion between the ions.Following Weiss, this author assumed that re-orientation of the solvent Fe2+ + Fe* + %Fe3 + + Fe* 2 + , (1) 182A . C. ATEN, J . DIELEMAN A N D G . J. NOIJTINK 183 molecules takes place after electron tunnelling. Contrary to the authors mentioned above, he took into account the continuous decrease of the dielectric constant with increasing field strength of the ion. In addition, the encounter rate was corrected for the electrostatic repulsion between the ions. The final results of his calculations revealed that the greatest contribution to the negative activation entropy was due to a low probability of electron tunnelling, whereas the repulsion between the ions gave rise to a small positive entropy value.In the light of this treatment, the mechanism proposed by Marcus becomes somewhat doubtful, since the negative activation entropy should be largely due to repulsion. Laidler’s results, however, do not contest the idea that re-orientation of the solvent molecules should precede electron transfer, the contribution of this orientation to the activation entropy being small in comparison with other con- t ri bu tions. Similar considerations as for the isotopic exchange reactions may be given for the electron transfer at electrodcs at the standard potential of the ox-red couple. As a consequence one is here again confronted with the alternative viewpoints, re-orientation of the solvent molecules before or after electron transfer.A completely different standpoint has been taken by Hush 6 who assumed the orientation of the solvent molecules to be in equilibrium at any stage during the electron transfer with the probability density of the electron at the ion. In order to check the above theories, it is worthwhile considering electron-transfer processes in which one of the collision partners is neutral, since then the relatively strong contribution of the electrostatic interaction between the reaction partners to the activation free enthalpy is absent. A simple type of such a process is the electron transfer between an aromatic molecule and its mono-negative ion in solution and the corresponding electrode reaction, which will be discussed in this paper.Special stress will be laid on the behaviour of aromatic hydrocarbons, since their negative ions have been studied in some detail in the past few years. Before considering the electron-transfer processes, a brief review is given of the properties of negative ions of aromatic hydrocarbons and related compounds which are of significance for the study of the electron transfer. PROPERTIES OF AROMATIC NEGATIVE! IONS Aromatic hydrocarbon mono- and di-negative ions are formed by the reaction with alkali or alkali-earth metal in ethers, amines and similar solvents. Measure- ments of the electrical conductance of these solutions indicate that under favourable conditions (high dielectric constant, small radius of the metal ion and low tem- perature) the ions are free.Thc association of hydrocarbon negative ions with alkali ions has recently been investigated by one of us (J.D.).7 The results un- equivocally show that the association leads to the formation of ion pairs in which the ions are separated by at least one layer of solvent molecules. Since the results are of particular interest for the present discussion they are briefly summarized below. (i) The solvation free enthalpy increases with increasing dielectric constant and with decreasing radii of the ions. In accordance with the negative temperature coefficient of the dielectric constant, the solvation becomes stronger when the temperature is lowered. (ii) Ionic association increases with increasing radius of the metal ion, with decreasing dielectric constant and with increasing temperature.The di-negative ions, which are not further considered in this paper, are associated with at least one, and usually with two, gegen-ions. From the above results it becomes very likely that the physical behaviour of these hydrocarbon negative ions in ionizing solvents is closely similar to that of uncomplexed inorganic ions. Accordingly, a reasonable estimate of the solvation184 ELECTRON TRANSFER T O AROMATIC MOLECULES free enthalpies of these negative ions may be obtained with the aid of Born’s equation : Ags(M-) = --?( 1 -;), in which N is Avogadro’s number, D is the static dielectric constant and Y the radius of the negative ion with one compact solvent layer. The following solvation free enthalpies have been calculated by Lyons 8 for aqueous solutions; the radii were obtained from the molar volumes of the hydrocarbons.benzene - Ag,(M-) = 46 kcal mole-1, naphthalene = 43 kcal mole-1, anthracene = 40 kcal mole-1. The solvation entropies follow from --(--.-) Ne2 8 1 . 2r d7’D (3) From eqn. (2) and (3), one finds for the solvation enthalpy : For solvents like tetrahydrofuran this equation becomes ( 0 ~ 7 9 5 , (dD/aT),w -0.03, T = 300K”)’ Ahs(M-)z 1*2Ags(M-). In comparison with the solvation energies of the mono-negative ions those of the hydrocarbon molecules may be neglected. In ion-dipoles of hydrocarbon negative ions with a relatively low polarizability, such as the ions of the peri- and cata-condensed hydrocarbons, the solvated metal ion will be located at the centre of the plane of the ion.In the ion dipoles of the strongly polarizable polyphenyl ions, on the contrary, the position of the ions may be quite different.7 ELECTRON TRANSFER IN SOLUTION Ward and Weissman 10 have investigated the rate of electron transfer for the reaction using as solvents dimethoxyethane and tetrahydrofuran. Various hydrocarbon mono-negative ions display an e.s.r. spectrum with a well-resolved hyperfine pattern. At a given concentration of the molecule, the hyperfine structure bands broaden owing to a rapid electron transfer between the mono-negative ion and the corresponding molecules. From the widths of the fine structure bands and the concentrations of the molecule and the ion at which broadening is observed, the rate constant of electron transfer could be calculated.Unfortunately, owing to the high concentration of the hydrocarbon molecule required, these experiments remained restricted to naphthalene. The results found by Ward and Weissman for sodium naphthalene are : [naphthalene]-+ naphthalene+naphthalene+ [naphthalene]-, (6) dimethoxyethane : tetrahydrofuran : k = lo9 1. mole-1 sec-1, k = lo7 1. mole-1 sec-1. These rate constants differ only very slightly with temperature, the activation enthalpy being lower than 2-6 kcal mole-1. Since the free enthalpy of diffusion will amount to a few kcal mole-1, the residual factors can contribute only to the activation entropy.A . C . ATEN, J . DIELEMAN AND G . J . HOIJTINK 185 On the basis of eqn. (5) and the solvation free enthalpies listed in table 1, the equations derived by Marcus predict an activation enthalpy for re-orientation of the solvent of about 5 kcal mole-1.This implies that re-orientation of the solvent molecules does not take place before electron transfer, so that the mechanism proposed by Marcus 4 may be disregarded. The absence of an activation enthalpy also means that the reaction partners remain solvated in the collision complex. From our investigations, it appeared that in dimethoxyethane sodium naph- thalene is completely dissociated into ions whereas in tetrahydrofuran it dissolvcs as an ion-pair. In all probability, the two different rates of electron transfer thus refer to different structures of the reacting species. The absence of an activation enthalpy for the electron transfer as such suggests the following formula for the rate constant : k = k, exp (APIR), (7) where k, stands for the encounter rate constant.Taking lc, = 1010 1. mole-1 sec-1 one obtains for the electron transfer between the molecule and its free mono- negative ion AS* = - 4.5 cal/mole deg., which is appreciably less negative than the value found for similar electron-transfer processes of inorganic ions.1 Most authors ascribe this negative activation entropy to a low probability of electron tunnelling. In our opinion, however, these authors localize the orbital of the electron at the donating ion too strongly within the crystallographic sphere of the ion. With Mulliken,ll one may say that the electron acceptance volume of the molecule or ion is much larger than the van der Waal's or crystallographic volume.This means that the electron may pass into the accepting orbital of the collision partner without an appreciable change of energy even when the two partners are separated by two compact solvent layers. The probability of such a transfer will depend on the overlap between the accepting and donating orbitals and the electrostatic interaction between the electron and the accepting molecule or ion. As for the present case, one may say that the orbitals possess more anti-bonding character and consequently will give stronger overlap according as the interaction between the electron and the neutral molecule is weaker. Hence, the two effects may compensate each other for a great deal, so that the probability of electron transfer very likely does not differ too much for different hydrocarbons.In tetrahydrofuran the mechanism is quite different. For this case one obtains in the same way AS!' x - 13.5 cal/moIe deg. The absence of an activation enthalpy for the electron transfer as such again suggests that the reacting systems remain solvated and no re-orientation of the solvent molecules takes place before electron transfer. Since the mono-negative ion of naphthalene is strongly associated with the sodium ion the collision complex will possess a sort of sandwich structure, the solvated sodium ion lying between the two hydrocarbon systems. Owing to the opposite charges the alkali ion and the molecule and molecule-ion will very likely be separated by one layer of solvent molecules. In comparison with the foregoing case, the electron transfer requires a much higher order of the collision complex so that the more negative activation entropy is not surprising.Owing to the presence of the alkali ion in the collision complex the rate of electron transfer will depend on the nature of the positive ion. as has indeed been observed by Ward and Weissman.10 An important result of this type of electron transfer is, that after the molecule has accepted the electron the sodium ion changes partner. Every successful collision thus involves the transfer of an electron and a sodium ion. As a con- sequence, electron transfer is here identical with atom transfer.I86 ELECTRON 1’KANSI;EK ‘I0 AKOMATlC MOLECULES Recent iiivestigations by Adam and Weissman 12 strongly support this mechan- ism.They investigated the e.s.r. of a solution of sodium benzophenone in di- methoxyethane. From measurements of the electrical conductivity of this solution in our laboratory, it appeared that the benzophenone ion must be strongly associated with the sodium ion. In this connection Adam and Weissman speak about a “ ketyl molecule ”, a name which unjustly might be associated with the classical description of benzophenone with the oxygen linked to a sodium atom. It should be emphasized, however, that such a “ molecule ”, or rather radical, has a much higher electron affinity than the benzophenone molecule so that it disproportionates into the corresponding radical ion and benzophenone, according to the reaction 2MNa+MNa- +M+Na+. Hence, we must conclude that the benzophenone ion forms an ion-dipole with the sodium ion.Nevertheless, Adam and Weissman observed that a finite electron density exists at the sodium ion, which in the light of the foregoing discussion means that a fairly strong overlap exists between the orbital of the benzophenone ion and that of the solvated sodium ion. In the electron-transfer complex where the sodium ion on either side is neighboured by a benzophenone system, this electron density at the sodium atom may even be higher. Adam and Weissman indeed found that at high concentrations of benzophenone the spectrum of the ion-dipole collapses, after which a fine structure remains corresponding to the interaction of the magnetic moment of the electron with that of the sodium nucleus. From this result the authors rightly concluded that an “ atom transfer ” takes place.ELECTRON TRANSFER AT ELECTRODE SURFACES The rate constants of electron transfer to various aromatic molecules in dimethyl-formamide at the dropping mercury electrode have recently been measured by one of us (A. C. A.). Following an experimental technique closely similar to that used by Randles,l3 the faradaic impedances were measured at the d.c. half-wave potentials, using tetrabutyl ammonium iodide as supporting electro- lyte. These investigations have been described in more detail elsewhere.14 Unfortunately, the rate constants of electron transfer to the aromatic molecules appeared to be too high to be measured. The only conclusion that may be drawn is that these rate constants are higher than 5 cm s e c l .If one assumes that the molecule is separated from the electrode surface by two solvent layers this corre- sponds to a lower limit for the frequency of electron transfer of about 5 x 107 sec-1. Since the diffusion of the molecules and ions has been taken into account explicitly the rate constant is given by k = (kT/h) exp (- AG*/RT). (9) Owing to the large excess of supporting electrolyte, the potential at a distance from the electrode of two solvent layers will differ only slightly from that in the bulk of the solution, so that repulsion of the negative ion by the electrode will not contribute very much to the activation enthalpy. By analogy with the electron transfer in solution, one may therefore suppose that the activation enthalpy is very small and the rate constant is practically determined by the activation entropy.This leads to a lower limit for the activation entropy of about -20 cal/mole deg. The high dielectric constant of dimethyl formamide ( D rn 37) suggests that the mono-negative ions are practically free. In that case one may expect a rate constant comparable with the homogeneous rate constant for sodium naphthalene in dimethoxyethane, which implies that the activation entropy should be roughly -4.5 caljmole deg., and consequently the rate constant should be about three orders higher than the lower experimental limit, provided the extent of overlap between the molecular orbital and electrode orbital is about the same as between two molecular orbitals.A . C . ATEN, J . DIELEMAN AND G . J . NOIJTINK 187 An important conclusion may be drawn from the measurements of the rate of electron transfer to the aromatic mononegative ions.14 The rate constants of these processes which are appreciably lower than those for the transfer of the first electron are practically independent of the structure of the aromatic compound. On the basis of these results it seems reasonable to assume that the rate constant for the first reduction step will also be of the same order of magnitude for all aromatic compounds. 1 Silverman and Dodson, J. Physic. Chem., 1952, 56, 846. 2 Weiss, Proc. Roy. SOC. A, 1954, 222, 128. 3 Zwolinski, Marcus and Eyrhg, Chem. Rev., 1955, 55, 157. 4 Marcus, Can. J . Chem., 1959, 37, 155 and related papers. 5 Laidlcr, Can. J. Chem., 1959, 37, 138. 6 Hush, J. Chem. Physics, 1956,24, 965. 7 Dieleman and Hoijtink, to be published. 8 Lyons, Nature, 1950, 166, 93. 9 Critchfield, Gibson and Hall, J. Amer. Chem. SOC., 1953, 75, 6044. 10 Ward and Weissman, J . Ainer. Cheni. SOC., 1957, 79, 2086. 11 Mulliken, Rec. trclv. chirn., 1956, 75, 845. 12 Adam and Weissman, J. Amer. Chem. SOC., 1958, 80, 1518. 13 Randles and Somerton, Tram. Furuday SOC., 1952, 48, 937, 14 Aten, Thesis (Free Univcrsity, Amsterdam, 1959). Aten and Hoijtink, Int. Congr. Poiurogruphy (Cambridge, 1959).

ISSN:0366-9033    

DOI:10.1039/DF9602900182

出版商:RSC    

年代:1960

数据来源: RSC