摘要:

THE OXIDATION AND REDUCTION OF FREE RADICALS BY METAL IONS IN AQUEOUS SOLUTION BY E. COLLINSON, F. S. DAINTON, I>. R. SMITH, G. J. TRUDEL AND (IN PART) S. TAZUKB Dept. of Physical Chemistry, The University, Leeds, 2 Received 8th February, 1960 If Rp is the rate of the radiation-induced polymerisation of acrylamide in aqueous sohtions containing metallic ions, M, which terminate the growth of polymer radicals by oxidizing or reducing them, the graph of R$ against Rp [MI always has a linear portion, of which the negative slope is proportional to the rate constant of this termination reaction. Tn addition, from certain of the graphs may be obtained (i) values of GH/COH or GH~o,/(GH+ GOH) in good agreement with established values, and (ii) relative rate constants for the reactions of metal ions with hydrogen atoms or hydroxyl radicals.The A and E factors for the oxidation of polymer radicals by (i) Fe3+ and Cu2+ in H20 and D20 and (ii) CeOH3+, TP+, Hg2+ and Ag+ in H20 are all very small, whereas for their reduction by Ti3+ in H20 and D20, A = 1012 1 . mole-1 sec-1 and E = 14.6 kcal mole-1. The oxidation of the polyacrylamide radical is considered to be an electron transfer in which the completeness or otherwise of the d-shell is the dominant influence, whereas the oxidation of H-atoms, which probably proceeds through the formation of M-H, is controlled primarily by the d+s orbital energy difference. The reductions of OH and polymer radicals by Ti3+ necessarily involve 0-H bond fission. It has long been known that simple inorganic free radicals such as He, HO., Hop and C1.can enter into oxidation-reduction reactions with inorganic ions in aqueous solution. Although relative values of the rate constants of the reactions of metal ions with a particular radical are known in a very limited number of instances, reliable absolute values of rate constants are difficult to obtain. Recently, and especially in the last decade, it has become clear that the cations of metals of variable valency can be oxidized or reduced by organic free radicals. For a free radical denoted by R. and a metal ion of oxidation number +x, these reactions may be written formally as eqn. (1) and (2) : R. + M ~ + -+M(~- l)+ + R+, (1) R. +M(X-~)+-+M~+ +R-. (2) It was pointed out by Mackinnon and Waters1 that free radicals and their anions (R-) and cations (R+) may be regarded as two redox pairs R.+R++e-, E,,,), (3) R--R.+e-, E(++ (4) having potentials E(-e) and E(+e) as indicated, and that consequentIy whether reaction (1) or (2) will occur, will depend on the relative magnitudes of E(-e), E(+e) and the redox potential EM(^) of the M(x-l)-+/MX+ couple.Thus if E(-e) is more positive than Ep~l(~) ( U . S . convention) then the oxidation reaction (1) will take place. In practice, for many radicals E(-e> and IT(+,) are not measurable under thermodynamically reversible conditions and these quantities can only be evaluated approximately by determining whether reaction (1) or reaction (2), or their anionic counterparts, take place between the same radical and different ions. By this 188COLLINSON, DAINTON, SMITH, TRUDEL, TAZUKB 189 means Haines and Waters2 have assigned an approximate redox potential E(-e)N- 0.4 V to the radicals oC(CH~)(CN)CH~CH~COOH and *c(cN) (CH2CH2COOH)2.Similarly, polyacrylonitrile radicals [R(CH2CH(CN)>,*] are reduced by Ti(III), Mo(III), U(III), W(III), CrUI), V(II), and Eu(I1) ions3 swg- gesting that E(+e) for this radical is more negative than N-0.1 V. This approach to the problem of the assessment of the oxidizability or reduci- bility of free radicals may be objected to on the grounds that, although the actual value of AG? or AGZ may be large and negative, the value of the rate constant ki or k2 may be such that under the experimental conditions chosen the particular reaction (1) or (2) may be undetectable and an erroneous conclusion about E(+=) or will therefore be drawn.The reason for this is that the actual course of the reaction written formally as (1) or (2) may be different for different ions. Thus the oxidation of a free radical may conceivably involve simple electron transfer ( q n (111, hydroxyl radical addition (eqn. (5)), Re + MXf + OH-(or HZO)-,ROH + M(X- ')+(or + H'), Re + MX++R minus H + M(X-')' + H', ( 5 ) (6) or hydrogen-atom abstraction either as a proton (eqn. (6)), or even as a hydrogen molecule, in a reaction in which the radical and the metal ion are both oxidized, Ra+MX++H+-, R minus H+H2+M(X+')+. (7) Furthermore, water molecules of the inner or outer hydration spheres of Mx+ or hydrolyzed or oxygenated or even polymeric forms of MX+ may be involved in the rate-determining step.For a better understanding of these reactions it is necessary to know the kinetic parameters of a series of closely related reactions. As a first step we aim to determine the rate constants of the reactions of a particular radical towards different cations and the effects on them of changes in temperature, ionic strength, nature of the anion, isotopic composition of the solvent, etc. To achieve this it is necessary to know the concentration of R* whilst reaction (1) or (2) is occurring. If R* is the growing polyradical chain carrier mi*, of the polymerization of a vinyl monomer (ml), the propagation step of which may be written the concentration of Re = [mj.], when high polymer is formed, is at all times given by Rp/kp[ml], where Rp is the rate of polymerization. Hence measurement of Rp allows [mi*] to be calculated, provided kp is known, and the larger kp the more sensitive the method.The polymerization of acrylamide (CH2 : CH . CO . NH2) is a very suitable reaction for this purpose, since kp is large (= 1.8 X 104 1 . mole-1 sec-1 at 25°C) and the mechanism of the reaction well understood.4~ 5 This paper is concerned with the determination of the rate constants of the reaction of polyacrylamide radicals with ferric, cupric, ceric, mercuric, thallic and argentous ions in perchloric acid media and titanous ion in sulphuric acid media and the measurement of the relative rate constants of the reactions of hydrogen atoms with argentous and mercuric ions. THE PRINCIPLE OF THE METHOD When a deaerated aqueous solution of acrylamide is irradiated with prays from cobalt-60 the reactions which ensue are :190 REACTIONS BUTWEEN RADICALS A N D CATIONS kio H.+ml-+ml.HQ. +m,-+m,. k i 1 INITIATION PROPAGATION mi. +ml%mj+ 1. (8) TERMINATION 2mj.-+dead kt polymer, (12) where GH and GOH denote the numbers of hydrogen atoms and hydroxyl radicals respectively, which are formed per 100 eV of pray energy absorbed. The steady- state rate of polymerization is then given by eqn. (13). = ([GI%+ Go~I]~abs)*kp[ml]/k~ mole I.-' sec-', (13) where Iabs is the dose rate in units of 100 eV 1 .-I sec-1. An oxidizing cation Mn+, could conceivably oxidize hydrogen atoms and polyradicals according to the formal eqn. (14) and (1 5) : H.+M*f?H+ +M("-')+ (14) inj. + M"+Adead polymer + M("- 'If.I<' (15) Increasing amounts of MX+, by diminishing the extent to which initiation by reaction (10) can occur and by augmenting the termination rate through reaction (15), will cause a progressive decrease in the rate of polymerization which is expressed in eqn. (16) : Rp" = Ri,O(GOH+ GHklO[ml]/(klO[ml] f k14[MX+]))/(GH+ - R,CMX '1 kp[mllk:/kt, (1 6) where Rp, 0 is the rate of polymerization when no metal ion is present (eqn. (13)). Two extreme cases may be distinguished depending on whether reaction (14) is negligible at all the concentrations of MX+ used. If reaction (14) is negligible the graph obtained by plotting RZ against Xp[MX+] will be a straight line of negative slope = k,[ml]/c:/k, and positive intercept on the ordinate = I?$,*.This is illustrated in fig. 1. Since k, and kt are known Ic; can be determined. Moreover at high concentrations of [Mn+] where almost all the chains are terminated by reaction (15), RZ is negligible in comparison with the other terms in eqn. (16) and hence Rp = (GH + G o d l a b s kp Lm 1 CM"+ 1 (17) and (GI,+ GOH)labs = -d[M"+]/dl* (1 8) If, on the other hand, Mx+ competes very effectively with acrylamide for hydrogen atoms, the curve of R: against Rp[Mx+] will show an initial rapid fall and then become a straight line of negative slope = k,[ml]kf//~, as before. This is illustrated in fig. 5. If the straight line portion is extrapolated to Rp[Mx+] = 0, the intercept will be R&o GoH/(GH+ GoH). Eqn. (16) may be rearranged to give eqn.(l9), (19) the whole of the left hand side of which can be evaluated from the straight line portion of the graph, its intercept and the observed values of R,,o and R,. A plot of 0 against [Mx+J/[ml] allows k14//~10 to be determined.COLLINSON, DAINTON, SMITH, TRUDEL, TAZUKE 192 If, instead of using oxidizing ions, reducing ions NJ”- capable of entering into reactions (20) and (21) are used, another set of equations : Ic ; mjo + NY i- -+dead polymer + N‘y ’ I)’, similar in form, and corresponding to (16), (17), (18), and (19) but with the sub- stitutions GH for GOH, kll for klo, k2o for k14, ky for /ti and NU+ for Mx+ are obtained. Clearly, the study of the effect of various metal ions on the rate of polymerization of acrylamide induced by ionizing radiation is of great potential utility for the determination of absolute values of ki, kr, GI* and GOH and relative values of k14 and k20.In addition, if measurements of gaseous hydrogen yields are made and the end-groups in polacrylamide are identified, then, as discussed later, valuable information concerning the precise mechanisms of reaction (lo), (ll), (14), (15), (20) and (21) may, under favourable circumstances, be obtained. EXPERIMENTAL The source of y-rays was 8 Curies of cobalt 60 and the source of X-rays was a 220 kVp G.E.C. Maximar X-ray set. The method of preparing dsaerated solutions for irradiation and the dosimetry wcre similar to those already dcscribed.4 RECORDING DILATOMETER.-The polymerization was followcd dilatometrically, the change in height of the level of the meniscus in the capillary tube being measured by the change in resistance of a 10 % w/w aqueous solution of potassium chloride separated from the irradiated solution by a column of mercury of constant length.The apparatus was essentially that shown in fig. 1 of ref. (4), exccpt that the change in resistance was ‘recorded by making the conducting column of potassium chloride solution part of an A.C. bridge circuit, the out-of-balance e.m.f. of which was alternately fed into a Honeywell- Brown 1 mV strip chart recorder with a full-scalc traverse-time of about 1 sec. At maximum sensitivity a change of 1 0 cm in height of the dilatometer meniscus corresponded to the full width (11 in.) of the chart. PREPARATION OF REAGENm-Water was twice distilled, the second distillation being from alkaline potassium permanganate in a Pyrex apparatus.The deuterium oxide (99-78 % D2O ex Norsk Hydro Elektrisk Akt.) was similarly treated. Acrylaniide as supplied by the American Cyanamid Co. Ltd., was recrystallized once from chloroform, the resultant crystals were washed several times with A.R. benzene and final traces of this substance were removed by evaporation in vucuu. A.R. perchloric was distilled twice under vacuum at 100°C. Ferric perchlorate was made by dissolving spectroscopically pure (British Bureau of Analysed Samples Ltd.) iron wire in perchloric acid and the resultant ferrous perchlorate solution diluted to 0.1 M concentration and oxidized by passage through it of ozonized oxygen. Excess ozone was subsequently displaced by a stream of pure oxygen.During oxidation a faint purple colour of permanganate ion developed which was due to manganese impurity in the iron to the extent of 0.023 %. This was reduced by the addition of an equivalent of ferrous perchlorate solution. In other experiments, G. Frederick Smith reagent-grade ferric perchlorate dissolved in perchloric acid without further purification was used. No difference was detectable between the effects of these ferric perchlorate solutions on the polymerization. Cupric, mercuric and silver perchlorate (G. Frederick Smith reagent grade) were used without further treatment. Thallous perchlorate prepared by dissolution of thallous carbonate in an excess of hot perchloric acid was thrice re- crystallized from doubly distilled water. Cerous perchlorate (G.Frederick Smith reagent grade) and thallous perchlorate, each dissolved in 2 M perchloric acid were oxidized to the ceric and thallic states in the anode compartment of an electrolytic cell of the type described by Biedermann.6 Since titanous ions reduce perchloric acid it was necessary to use sulphuric acid solution as the solvent for trivalent titanium. Approximately 10-2 M titanyl sulphate solution in 0.8 N sulphuric acid was prepared by gentle boiling of potassium titanyl oxalate in con- centrated sulphuric acid, followed by dilution with water. After further dilution with 0.8 N sulphuric acid, the titanyl sulphate solution was de-areated and then shaken under192 REACTIONS BETWEEN RADICALS AND CATIONS vacuum with liquid zinc amalgam from which it was separated by filtration through a sintered-glass disc of porosity no.3. I n order to ensure as high a concentration of deuterium oxide as possible in the heavy water solution, the appropriate acid and salt were added from an Agla micrometer syringe. ANALYTICAL METHODS.-The concentration of the ferric iron in the stock solution was determined gravimetrically as ferric oxide and in dilute solutions containing 0.8 N &SO4 it was deduced from the optical density at 304 mp assuming an extinction coefficient of 2150 1 . mole-1 cm-1 at 20°C. Cupric ion concentrations in stock solutions were deter- mined iodometrically and mercuric ion concentrations by titration with potassium thiocyanate. Any mercurous ion in the latter was estimated from the optical density at 236.5 mp using Higginson's value 7 of the extinction coefficient.Ceric ion concentrations in the stock solution were determined by titration against standardized ferrous sulphate solution, and in dilute solutions the ferric ion produced on addition of excess ferrous sulphate was estimated from the optical density at 304 mp. For the measurement of cerous ion concentration the method of Willard and Young 8 was employed. Thallic ion concentration was measured by reduction to the thallous state by passage of sulphur dioxide through the solution. After removal of the excess sulphur dioxide by boiling the solution was acidified with concentrated hydrochloric acid and titrated potentiometrically against a standard potassium bromate solution. Argentous ion concentration in the stock solution was obtained by titration against potassium chloride solution.The titanium content of the stock titanyl sulphate solution was determined gravimetrically as Ti02 obtained by ignition of the filtered precipitate produced by addition of cupferron. In all these soIutions the hydrion concentration was determined by some appropriate standard method. EXPERIMENTAL PROCEDURE.-The great sensitivity of the acrylamide polymerization to traces of impurity required strict adherence to the rigorous cleaning and de-aeration procedures which are in frequent use in this laboratory and which have often been described. The preparation of the de-aerated solution and their transfer to the dilatometer were by methods already described.4 RESULTS POLYMERIZATION IN THE ABSENCE OF METAL IONS No thermal rates of polymerization or inhibition periods were observed.At an acrylamide concentration of 0.4 M in 0.1 N perchloric acid solution R,, 0 was found to be proportional to I& over the dose rate range 1 . 3 2 ~ 1014 to 2.68 x 1016 eV 1-1 sec-1. Taking the value of GH and GOH in 0.4 M acrylamide solution as equal to 7.5 and to be independent of temperature,g the application of eqn. (13) to the measured values of R,, 0 at 18.6", 25.0", 32.5" and 40.3"C gave APIA; = 45 1.4 mole4 sec-) and Ep-(EJ2) = 1.38 f0.25 kcal mole-1. The value of kp/kj at 25°C was 4.4 I.& mole4 sec-* in good agreement with the value of 4.7 previously reported,4 and the values of kp = 1 . 8 ~ 104 and kt = 1 . 4 5 ~ 107 1 . mole-1 sec-1 given by Dainton and Tordoff.6 An increase in perchloric acid concentration from 0.1 to 1.0 M had no effect on Rp, 0, but replacement of perchloric acid by 0.8 N sulphuric acid halved the value of R,, 0 and replacement of the perchloric acid by 0.8 N sulphuric acid and 0.5 M sodium sulphate decreased Rp, 0 by 30 %.THE EFFECTS OF ADDED FERRIC AND CUPIUC PERCHLORATES The results obtained for 0.1 N perchloric acid solution are given in fig. 1 and the linearity of the graphs indicates that over the whole range of concentrations used, the metal ions present react only with polymer radicals and not with H- or -OH, both of which are captured by acrylamide. Applying eqn. (16) after putting k14 = 0, and substituting for kp and kt with the values given above, the values of k,' for ferric and cupric perchlorate at 25°C in 0.1 N perchloric acid are: k1(Fe(C104),) = 4.1 x 103 and k,'(Cu(ClO&) = 1 .2 ~ 103 1 . mole-1 sec-1. TheCOLLINSON, DAINTON, SMITH, TRUDEL, TAZUKB 193 value for the sum of these rate constants obtained from measurements made on the equimolar mixture is 5.1 x 103 1. mole-1 sec-1. Rp[M(Z+l)+] x l o 9 mole2 1.-2 sec-1 FIG. 1.-The effect of added ferric (-0-) and cupric (-O-) perchlorates and equimolar mixtures of the two (--@-) on the rate of polymerization of 0.4 M acrylamide solution in 0.1 N perchloric acid solution at 25°C. Dose rate = 2.68 x 1011 eV 1.-1 sec-1 for 0 and 0 , slightly less for 0, It has been shown elsewhere that k,' for hydrolyzed ferric ion is larger than that for unhydrolyzed ferric ion.+ s.To determine E: for the latter it is necessary to ensure that this species comprises < 99 % of the total ferric iron at all temperatures at which k;(Fe(ClO&) is measured. This was achieved by raising the perchloric acid concentration to 1 M. The Arrhenius plots for the ferric and cupric per- chlorate systems at this acidity in H20 and D2O are shown in fig. 2. THE EFFECT OF ADDED CERIC PERCHLORATE Ceric salts will initiate the polymerization of water-soluble vinyl compounds even in the dark.10 This thermal polymerization is at least second order with respect to monomer concentration 11 whereas the radiation-induced reaction is first order. Consequently in 0.2 M acrylamide solution the thermal rate can be reduced to a very small fraction of that of the radiation-induced polymerization.However, the radiation-induced reaction in the presence of ceric ions differed from all other systems investigated in that the steady-state was often not attained until several hours had elapsed. During this period, the duration of which increased with increased initial concentration of ceric ions, the polymerization slowly accelerated and the ceric ion concentration decreased initially at a rate corresponding to a large G(- Ce(1V)) value ( ~ 2 0 ) . At the end of this induction period both the concentration of ceric ion and the rate of polymerization became constant and the graph of the square of this steady rate against the product of this rate and the steady-state ceric ion concentration gave the straight line shown in fig. 3. The G1 94 REACTIONS BETWEEN RADICALS AND CATIONS value of kt(Ce(C104)4) obtained from this graph asuming that eqn.(16) with klo = 0, is applicable, is 3.4f0.3 1. mole-1 sec-1. It seems likely that the long induction period is due to reaction of ceric ion with hydrogen peroxide formed in l/T°K x 103 perchloric acid solution in H20 and D20. FIG. 2.-Log10(kpk,'/kt) plotted against T-1 for 0.4 M acrylamide solution in 1.0 N T 1- 0 Fe 011) in H20 I I Fe 011) in D20 1- - I 0 Cu (11) in H20 I I B Cu (11) in D20 1_ the primary act. The oxygen liberated adds to growing poly radicals forming mi02. radicals which either combine with other poly radicals or reduce ceric ions. The polymer peroxides are then oxidized in reactions with ceric ions which involve free radical intermediates and ultimately lead to polymers containing carbonyl groups.Many feasible reactions participate in this mechanism, the net result of which is slow reduction in the oxygen concentration and a rapid conversion of ceric ions to the cerous state. Cerous ions play no part in this process, since addition of cerous ion initially has no effect. The maximum permissible value of G(- Ce IV) based on the known radical and molecular yields is 35 and the observed value is less than this. When the concentration of ceric ions falls below a certain level the rate of regeneration of oxygen from the peroxidic products also diminishes and the concentration of oxygen becomes very low. Polymerization then proceeds at the steady rate (given by eqn. (17)) and the ceric ion concentration decreases at an undetectably small rate.The graphs of loglo (k,k:/kt) against T-* at two different acrylamide concen- trations are curved, as shown in fig. 4. This curvature is undoubtedly due to the increased hydrolysis of ccric ions as the temperature is raised, and to differentCOLLINSON, DAINTON, SMITH, TRUDEL, TAZUKb 195 I I I 5 10 R,, f [Ce LV]fx 1010 mole2 1.-2 set-1 FIG. 3 . T h e relation between the steady rate of polymerization of 0.2 M acrylamide solution in 0.1 N perchloric acid solution and the steady-state ceric perchlorate concen- tration at 25°C. Dose-rate = 2 X 1016 eV 1.-1 sec-1. l/T°K x 103 FIG. 4.---Effect of temperature on k,k,'/kt for ceric perchlorate in 1.0 N HC104. Acryl- :iniidc conccntratiori. 0 0 4 M ; 0 0 2 M ; ordinates separated for clarity.196 REACTIONS BETWEEN RADICALS A N D CATIONS values of E,' for Ce4+ and CeOH3+.In the temperature range used [Ce4+] decreases almost 6-fold whereas [CeOH3+] increases by a quarter and it is possible that the steeper slope at the lower temperatures is due to E,'(C&+)>E:(Ce-OH3+) and that the contribution which Ce4+ makes to linear termination is much smaller at the higher temperatures, in which case the limiting slope at higher temperatures will provide an upper limit for E,'(CeOH3+). THE EFFECT OF ADDED MERCURIC, THALLIC, AND SILVER PERCHLORATES The effect of each of these salts is much less marked than that of any of the salts mentioned above. In addition, the graphs of I?; against RP[Mx+], which are shown in fig. 5, have two distinct sections.The sections corresponding to the 9 0 6 0 3 0 4 a 12 Rp[M2+] x 106 mole* 1.-2 sec-1 FIG. 5.-The effect of added mercuric (-O-), thallic (-a- )and silver (-0-) per- chlorates on the rate of polymerization of 0.4 M acrylamide solution in 0.1 N perchloric acid solution at 25°C. Dose rate = 1 . 8 6 ~ 1016 eV 1.-1 sec-1. higher values of [Mx+] are linear but at low concentrations the plots are curved and of steeper slope. For mercuric and silver perchlorates the results are con- sistent with eqn. (16) when k149k10. Back extrapolation of the linear portions to intersect the ordinate gives values of (GH+GoH)/GoH = 1.93 (HgZ+) and 2-11 (A&) which may be compared with the value of 2.18 obtained for the acryl- amide water system by another method.9 The slopes of the linear portions indicated that at 25"C, k,'(AgCIOd) = 0 and k:(Hg(ClO&) = 1.05 I.mole-1 sec-1 and the plots of 0 (for definition see eqn. (19)) against [Mxf]/[M1] give straight lines corresponding to k14 (Ag+)/klo = 28 and k14 (Hg2+)/klo = 510 at 25°C. Thallic perchlorate is a less effective chain terminator than mercuric per- chlorate and from the slope of the linear part of the curve in fig. 5 we deduce that k,'(Tl(ClO4)3) = 0.34 1. mole-* sec-1 at 25°C. The rapid initial drop in rate of polymerization may be ascribed to a liberation of oxygen from the molecular hydrogen peroxide by reaction with thallic ions (eqn. (22)) T13 .t + H2O2+Tl+ + 2H' + 0, (22) followed by reaction of the oxygen with hydrogen atoms to form HO2 radicals.COLLINSON, DAINTON, SMITH, TRUDBL, TAZUKB 197 Hence, for each 100 eV energy absorbed G H ~ o ~ polymerization chains would be prevented from starting and the ratio of the intercept on the ordinate to Rg, should be (GH+ GOH- GH~o*)/(GH+ GOH) for which the predicted value 9 is 0.935 and the value obtained from the data depicted in fig.5 is 0.92. I 3.2 3.3 3.4 I 1 1/TK x 103 FIG. &-Log (kpkl/kt) plotted against T-1 for 0.4 M acrylamide solutions in 1.0N per- chloric acid : -0- added Hg(C104)~ ; -0- added TI(C104)3. In fig. 6 are shown the Arrhenius plots for the mercuric- and thallic-perchlorate terminated polymerization. THE EFFECT OF ADDED TITANOUS AND SODIUM SULPHATES Titanous sulphate solutions in 0-8N sulphuric acid have a broad, weak, ab- sorption band centred at 525 mp at which wavelength the extinction coefficient is about five.Between 4-00 mp and 280 mp, which is the long-wave absorption edge of the electron transfer band, the solutions are virtually transparent. When acrylamide is added to these solutions a new absorption band with A, = 325 mp, is produced. Measurements of the optical density at 325 mp of solutions con- taining a large, variable excess of acrylamide and a small, fixed, concentration of titanous sulphate indicated that the new absorption band is due to a 1 : 1 complex of acrylamide and titanous sulphate. In a solution containing 0.8 N sulphuric acid and 0.5 M sodium sulphate the equilibrium constant (= [complex]/r]ri 1111 [ml]) is G0.13 at room temperature and the extinction coefficient of the complex is N 285 1.mole-1 cm-1 at 325 mp. At the concentrations of acrylamide used in the polymerization experiments described below, not more than 1.3 % of the total titanous ion exists as the complex. The effects of increasing titanous sulphate concentration on the rate of poly- merization of 0.1 M acrylamide solution (a) in 0.8 N sulphuric acid solution and (b) in 0-8 N sulphuric acid solution also containing 0.5 M sodium sulphate, are shown in fig. 7. The ionizing radiation used was 220 kVp X-rays. This radiation rather than 6ocobalt y-rays was used because a higher dose-rate was desirable to compensate for the decrease in rate due to the lower monomer concentration necessary to minimize complex formation. No thermal rates or inhibition or198 REACTIONS BETWEEN RADICALS AND CATIONS induction periods were observed, and the presence of 1 .2 ~ 10-4 M titanyl sulphate had no effect on the rate of polymerization. The graphs of Ri plotted against Hp [Ti 1111 wcre similar to the corresponding graph for mercuric perchlorate (see fig. 5). The obvious explanation of these graphs is that titaiious ions compete with acrylamide for hydroxyl radicals and also terminate polymer radicals. The values of (GH+GoH)/GH obtained from the ratios of Rg,O to the intcrcept on the ordinate when the linear portions of the graphs are back-extrapolated are (a) in 0.8 N sulphuric acid solution, 1.72 and (b) in the mixed sulphuric acid+sodium sulphate solution, 1.82. This ratio has not previously been determined for 0.1 M acrylamide solution in these media for 220 kVp X-rays.In 0.04 M perchloric acid solution containing 0.1 M acrylamide 9 the ratio for cobalt-60 y-rays is 1.82, and it would be expected that the ratio for 220 kVp X-rays in 0.8 N solution would be close to, but would not exceed this value. I R, [Ti(lII)] x 108 mole2 1.-2 sec-1 FIG. 7.The effect of added titanous sulphate on the rate of polymerization of 0.1 M acrylamide solution at 25°C. Dose rate = 9.6 X 1017 eV 1.-1 sec-1. 0 denotes solvent is 0.8 N H2SO4 in H20 ; -a- denotes solvent is 0.8 N H2SO4+05 M Na2SO4 in H20 ; Qo denotes solvent is 0.8 N H2SO4 in D20. The data for solutions in D20 may be interpreted in a similar way and indicate that (GD+G~D)/GD = 1.87 which may be compared with the value of 1.86 for 0.1 M acrylamide in 0.04 N perchloric acid solution in D20 using y-rays.9 Assuming GH 3 3.95, GOH = 3.21, GD = 4.25 and GOD = 3.6 the values of k,/l/kt for 0.8 M H2S04 solution in H20, 0.8 N H2SO4 in D20 and 0.8 N H2S04+0'5 M Na2S04 are 2.36, 2.45 and 2.96 1.4 mole-* sec-) respectively.These results indicate that substitution of D for H has no effect on k,/kt but that substitution of 0-8 N H2S04 for 0.1 N HClO4 diminishes this ratio by a factor of two and that the addition of 0.5 M Na2SO4 increases k,,/kt by 50 %. It is not yetCOLLINSON, DAINTON, SMITH, TRUDEL, TAZUKk 199 known whether changes in this ratio are due to changes in kp or kt, or both. In the only remotely comparable situation which exists, namely, the polymerization of acrylonitrile in NN1-dimethyl-formamide,l2 the addition of 0.294 M lithium chloride increased the propagation rate constant without affecting the termination constant.Although final values of k," (Ti 111) must await the determination of the values of k, and kt in each of these media, it is instructive to compare the values which kl' (Ti 111) would have assuming that the change in k,/@ is due solely either to (a) a change in k,, or (b) a change in kt These values are tabulated below. TABLE 1 medium 0.8 N . H S04/Hz0 0.8 N . H2S04/D~0 0.8 N HzS04/0.5 M Na2S041HzO kP1kt 2.36 2-45 2.96 Assumption a b a b a b k11 CriIII) 1.14X I03 2.3X 103 1.59X 103 3.27X 103 1 . 1 0 ~ 103 1 . 7 9 ~ 103 The Arrhenius plot of loglo (k,(ml)k:l (Ti IIT)/kt) against 2'-1 gives a value of E,-J?Z~+E)~ (Ti 111) = 16*2 kcal mole-1 and AllAp/At = 8x 1010 1.mole-1 sec-1. Unpublished work in this laboratory 13 suggests that Et = 0 A0.5 kcal mole-1 and hence Ep = 1.38f0.25 kcal mole-1. From these data we calculate that Etl p i III) = 14.652.75 kcal mole-1 and that A i l (Ti 111) lies between 5 x 1010 and 5 x 1014 1. mole-* sec-1 in 0.8 N sulphuric acid. The quantity 8 in eqn. (19) can be evaluated for each titanous sulphate concen- tration from the data given in fig. 7 and hence we may obtain the ratio of the rate constants for the reactions of OH (or OD) with acrylamide and titanous sulphate in the various media. The values of k20 (Ti III)/kll at 25°C are : (a) in 0.8 N H2SO4 in H 2 0 , 650, (6) in 0.8 N H2SO4 in D20, 1600 ; and (c) in 0.8 N H2SO4, 0.5 N Na2S04 in H20, 860. DISCUSSION The results presented in the previous section indicate that the study of the effects of metallic salts on the rate of polymerization of acrylamide initiated by ionizing radiation afford a powerful means of determining (i) the absolute values of the Arrhenius parameters of oxidation-reduction reactions between various cations and the same free radical, namely, the polyacrylamide radical, (ii) relative values of the Arrhenius parameters of the reactions of hydrogen atoms or hydroxyl radicals with various cations and (iii), incidentally, where (ii) is possible, values of GH and GOH.Although it is our intention to study many more of these reactions and final conclusions concerning reaction mechanisms and the influences which control the magnitudes of the Arrhenius parameters must be deferred until this survey is complete, some correlation and tentative conclusions can be reached on the basis of the results obtained for the seven ions with which reactions have already been investigated.Most of the results are given in tables 1, 2 and 3, and the broad conclusions to be drawn from them are : (i) The values derived for the radiation-chemical quantities are in good agree- ment with those derived by an entirely independent method 9 for aqueous solutions of acrylamide of the same concentration. (ii) For the reactions between the radical and ferric, cupric or titanous ions the difference between the values of the rate constants in D20 and H20 are within the experimental error, but there is a genuine solvent isotope effect for the reaction between hydroxyl radicals and titanous ion which is contrary to that usually observed for oxidation-reduction reactions.TABLE 2 REDUCTION OF CATIONS (Mx+) BY THE POLY-ACRYLAMIDE RADICAL (mj*)AND BY A HYDROGEN ATOM (Ha) derived -radiation chemical quantity Fe3+ - 0.77 1 - 17.8 - 9.6 + 26.5 H20,O.l N HClO4 4.1 X lo3 .... .................. too small H20, 1.0 N HClO4 2.8x 103 1 . 4 5 ~ 10s 2.35k0.6 to Nil .................. measure D20, 1.0 N HClO4 2.6 X lo3 1.57X 10s 244kO-3 -3.5 uncertain + 1.7? H20, 1.0 N HC104 1.17%' 103 1.1 X lo7 5.4h1.3 too small .................. to D20, 1.0 N HC104 1.4 x I03 1.0 x 107 5351.3 measure CeOH3+ -1.7 - 39.0 (in 1.0 N HC10,) , HzO, 0.1 N HClO4 3.4 X lo3 .. . . too small .................. to H20, 1.0 N HClO4 3.2 X 103 2 X 105 p2.45 measure Nil Nil Hg2+ (-0.9 forHg;+/Hgz+) (-;I) .. . . HzO. 0.1 N HClO4 1.05 4.2 X 104 6*2&1*0 510 (GH+GoH)/GoH= 1.93 unknown for Hg+/HgZ+ . TI3+ Eo for TP+/TP+ Between . . . . H20.0.1 N HC104 0-34 21 2.5 -1-0.4 too small (GH+ GoH)/(GH+ GOH- GHZO~) unknown but - 18 and to - 1-09 -0*79> E"> - 1*7(e) -40 measure Ag+ +l*8 (f) +41.4 +434 +lS HzO, 1.0N HCIOd 0 .. .. 28 (GH+ GoH)I(GoH = 2-1 1 NOTES (a) Since k,/kt and E,,-(Et/2) have same values in H2O as DzO, A,, A t , Ep and Et are assumed to be the same in each of these solvents. (6) taken from Latimer, Oxidation Potentials. New York. 1952. (c) Bureau of Standards values; heat of formation of H& and H2 taken as zero. ( d ) based on convention SR+ = 0 and taking si&) = 31.2 cal deg. mole-1.(e) From the facts that Eo (TI+/Tl3+) = - 1.24 and both Tl3+ and TP+ oxidize Fe2f.17 ( f ) refers to ~H2(s)+fAga'-Ha's+Ag(6). Ei calculated from AGO. c( 0COLLINSON, DAINTON, SMITH, TRUDEL, TAZUKi! 201 (iii) Whereas the entropy of activation for the oxidation of titanous ion by the radical is nearly zero, the entropies of activation for the reduction of ferric, cupric, ceric, mercuric and thallic ions are all surprisingly large and negative, having the values - 35, - 25, - 3 1, - 38 and - 53 cal deg.-l mole-1 respectively. Moreover, these values have no obvious relation to AS" for the reduction of the ions by hydrogen (see column five of table 2). (iv) The energies of activation for the reduction of cations by the polymer radical are all much smaller than that for the oxidation of titanous ion.(v) The rate constants for reduction by polymer radicals of ferric, cupric and ceric ions are all between lo3 and 104 times larger than the rate constants for reduction of mercuric and thallic ions and bear no relation to AGO for the reduction of these ions by hydrogen. The stability of argentous ions towards the radical may be associated with the very large positive value of AGO for the reduction of this ion to atomic silver by hydrogen. (vi) It is broadly true that those ions which react rapidly with hydrogen atoms react slowly with the polymer radical, and vice versa. Thus argentous and mercuric ions react rapidly with the atom and very slowly or not at all with the radicals, and ferric ions, which oxidize the radical more rapidly than cupric ions, oxidize hydrogen atoms less rapidly.14 TABLE 3.-oXIDATION OF TITANOUS ION (TI3'') IN 0.8 N H2S04 IN H20 AND D20 BY THE POLYACRYLAMIDE RADICAL (mj .) AND BY THE HYDROXYL RADICAL OH) k2l at 25°C A21 E21 k20/kll at 250c derived radiation medium 1.mole-1 sec-1 1. mole-1 sec-1 kcal chemical quantity (a) (6) H20 1.14x lo3 2.3 x l o 3 1012*2 14*6&2*75 N 600 (GH+GoH)/GH = 1.72 D20 1-59 x 103 3.3 x 103 .. .. N 1600 (GD+GoD)/GD = 1-87 RELATED THERMODYNAMIC DATA Reaction AGO OH+TP+-+TiOz++ H+ - 62 Hf HzO+TiJ++H++H2fTi02+- 46 (a) and (b) refer to the assumptions that sulphuric acid alters (a) k, only (b) kt only. (vii) Previous studies 15@) in these laboratories have shown that k2o/lcll for ferrous ions at 25°C is 0.02 and therefore kzo(Ti3+)/~20(Fe2$)-"3x 104 which is the order expected for these ions, since AGO for the oxidation of Ti3+ by OH is e l 4 kcal more negative than that for the oxidation of Fe2+.However, the solvent isotope effect for the Ti3++ OH reaction is opposite to that which probably holds for the Fe2++ OH rea~tion.ls(~) REDUCTION OF CATIONS All attempts to correlate either the rate constants, or the entropies or energies of activation, of these reactions with thermodynamic properties of the redox system including relative free energy, enthalpy and entropy changes, or with changes in solvation heats or entropies, or ionization potentials are unsuccessful, and this suggests that entirely non-thermodynamic influences dominate these reactions. It is very striking that ions with incomplete d orfshells react rapidly with the radical, whereas the ions with complete d shells either react very slowly (Hg2+ and TP+) or not at all (Agf) with the radical and may react (Hg2+ and Ag+) more rapidly with hydrogen atoms than transition element ions such as Fe3-+ and Cu2f.This suggests that the oxidation of the radicaI involves the utilization of one of the vacant d orbitals of the cation. We may envisage the approach of the radical to the cation so as to achieve sufficient overlap of the p-orbital of the202 REACTIONS BETWEEN RADICALS AND CATIONS carbon and a d-orbital of the cation in the transition state for rapid electron transfer to take place. The release of an electron by the radical enabIes it to expel a proton. On this hypothesis the overall reaction might be represented in the case of cupric ion by eqn.(23) : Cu2+ + . C H ( X ) C H ~ ~ ~ - ~ + ( C U . . . CH(x)---CH-mj)2++Cu+ + 3d9 i 3di0 H +CH(x)=CHmj_, (23) H+ transition state (where x = CONHz), and would lead to the predictions : (a) that the energy of activation would have the low value typical of true electron-transfer processes, (6) that ions with completed d shells would not react, (c) that no marked solvent isotope effect should exist since no 0-H bond breakage occurs in the rate determining step, and (d) that the polymer should contain a terminal vinylene group. (a), (b) and (c) have been shown in this paper to be in accord with the facts and (d) will be tested. This model does not indicate that the observed negative entropy of activation is caused by stronger solvation of the transition state.Possibly electron " tunnelling " transfer is involved and the negative AS+ values are due to low permeability of the barrier. For a reaction of the type shown in eqn. (23) to be exothermic, almost complete solvation of the proton is necessary. A long-lived transition state will facilitate the development of this solvation, and the incipient bond formation through an unoccupied d-orbital, which is implicit in the above mechanism, is clearly ad- vantageous. The question then arises as to whether a similar mechanism applies to the reduction of cations by hydrogen atoms. Although in principle this might be the case, it should be remembered that hydride formation involving d-orbitals is unknown, but that a($ hydride bonds involving metals are common.If incipient formation of such a bond is a necessary prerequisite in the reduction of a cation by hydrogen atoms, the magnitude of the energy necessary to promote an electron in the cation to the lowest available s orbital, may be the controlling influence on the rate, in the sense that the larger this energy the smaller k14. This d-s energy interval will be smaller (a) the smaller the charge on the ion and (6) the larger the principal quantum numbers involved, and hence also the greater the atomic number. On this basis it is immediately apparent why the ions Hg2+(5dlO) and Ag+(MlO), having low charge and large principal quantum number should be reduced more rapidly than Fe3+(3d5) and Gi2$-(3d9) which have either high charge or low principal quantum number or both.* For thallic and Eric ions the charge and quantum number influences are in opposition, but further experiment is necessary to establish whether these two ions occupy an intermediate position between A& and Cuz+.Arguments somewhat similar to these have been advanced by Halpern 16 to account for the effectiveness of certain cations as catalysts for the activation of molecular hydrogen. These reactions involve the transfer of either an H. atom or the hydride (H-) ion to the catalyst cation and optimum catalytic activity is observed when the d-shell is filled or nearly filled and the ion can make available an empty d-orbital by promotion. The hydride intermediates proposed include ASH+, HgH+ and CuHf, which will certainly utilize bonding orbitals.OXIDATION OF TITANOUS IONS Titanium (IV) exists in aqueous solution as the Ti02+ ion. Consequently the oxidation of the titanous ion by the OH radical may perhaps be regarded as the * Results of radiation-chemical experiments indicate that H atoms do reduce Fe3+, T13+, Cu2+ and CeSO:+ ions, and that, as would be expected, kl4(Cu2+)>k,4(Fe3+).14COLLINSON, DAINTON, SMITH, TRUDEL, TAZUKE 203 transfer of charge by transfer of 0- from the OH radical and expulsion of the proton. The reaction would be expected to be more rapid than the oxidation of ferrous ion by the OH radical because -AGO is much larger, and the entropy of activation should be more positive since, unliko the ferrous+ OH reaction, no additional charge is developed in forming the transition state.If the reaction mechanism is correctly represented by the equation given in table 3, it might be expected that a change of radical from OH to OD would retard the reaction, whereas the present indications are that the reverse is true. This may be evidence of some net desolvation in forming the activated complex, associated with the decrease in charge on the titanium. Since D20 solvates cations more strongly than H20 this could lead to greater entropy gain in the former solvent than in the latter, an effect which might be sufficiently strong to predominate and to invert the normal specific solvent effect in oxidation-reduction reactions. The oxidation of Ti(II1) by mp necessarily involves the breaking of both the OH bonds of one of the bound water molecules.The reaction can be one of attachment of a hydrogen atom to (eqn. (24)), Ti~OH~++~CH(x)~CH2mj~1+TiOz++H++HCH(x)CH2mj~l, (24) or detachment of a hydrogen atom from (eqn. (25)), the polymer radical. We have some evidence, which will be presented elsewhere, that eqn. (25) is the more correct, but whcther this involves electron transfer from the radical with proton expulsion from it (eqn. (25(a))) Ti3+*CH(x)CH.mj- 1-+Ti2+ + CH(x): CHmj- 0 /OX\ H + H+ 1 H H H2 in a manner which recalIs eqn. (23), or whether there is no charge flow in the radical (eqn. (25(b))), Ti3 +.CH(x), -+Ti2+CH(x) 7 x = CONH, is not yet determined. If the former is true, d-orbital influence may be important, but only experiments with other ions can establish or disprove this. If the latter is correct the hydrated Ti(II1) ion may be regarded as a hydrogen atom capable of entering into a disproportionation reaction with the radical. The zero entropy of activation and the high energy of activation suggest that this may be the case. We are grateful to the General Electric Research Laboratory for some financial aid to S.T. and G.J.T. 1 Mackinnon and Waters, J. Chem. Soc., 1953, 323. 2 Haines and Waters, J. Chem. SOC., 1955, 4256. 3 Dainton and James, Trans. Faraday SOC., 1958, 54, 649. 4 Collinson, Dainton and McNaughton, Trans. Faraday SOC., 1957, 53, 476, 489. 5 Dainton and Tordoff, Trans. Faraday SOC., 1957, 53,499. 6 Biederman, Arkiv. Kemi, 1953, 5, 441. 7 Higginson, J. Chem. SOC., 1951, 1438. 8 Willard and Young, J. Amer. Chem. Soc., 1928, 50, 1379. 9 Armstrong, Collinson and Dainton, Trans. Faraday SOC., 1959, 55, 1375. JOBacon, Trans. Faraday Soc., 1946, 42, 140; Saldick, J. Polymer Sci., 1956, 19, 73.204 REACTIONS BETWEEN RADICALS AND CATIONS 11 Second-order with respect to monomer concentration would be expected if the initiation reaction involved a bimoIecuIar reaction between acrylamide and ceric ions and termination were exclusively by reaction (16). The recent work of Mino, Kaizerman and Rasmussen, (J. Polymer Sci., 1959, 38, 393) suggests that this mechanism is probable. 12 Bamford, Jenkins and Johnston, J. Polymer Sci., 1958, 29, 255. 13 S. A. Zahir, Thesis (Leeds), 1960. 14 Riesz and Hart, J. Physic. Chem., 1959, 63, 858. 15 (a) Dainton and Hardwick, unpublished data. (b) Bunn, Dainton and Salmon, unpublished data. 16 Halpern, J. Physic. Chem., 1959, 63, 398. 17 Higginson and Ashurst, J. Chem. SOC., 1953, 3044.

ISSN:0366-9033    

DOI:10.1039/DF9602900188

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

INDUCED OXIDATIONS OF ORGANIC COMPOUNDS BY PERMANGANATE IN ALKALINE SOLUTIONS BY K. A. K. LOTT AND M. C. R. SYMONS Dept. of Chemistry, The University, Southampton Received 20th January, 1960 Certain organic compounds, RH, unaffected by neutral solutions of permanganate, are readily oxidized at room temperature by permanganate in concentrated alkaline solution. Solutions of permanganate in aqueous alkali decompose to give oxygen and manganate : 4Mn0, +4OH-+4Mn0~-+2H2O+ O,, but when a large excess of RH is added, oxygen evolution is suppressed. Formulating the first stages of the induced oxidation by the steps ki MnO; + OH- + MnOi- + OH, k- 1 RH+ . O H ~ R . +H,o, and applying stationary state kinetics, gives the equation for the rate of the reaction. This equation is obeyed closely over a wide range of con- centrations, and the value deduced for kl is close to that derived from studies of initial rates in the absence of organic substrate.MODES OF REACTION OF PERMANGANATE Mechanisms of oxidation by permanganate in alkaline solution have been studied extensively, generally by means of rate measurements (see, for example, ref. (1) and (2)). This work has recently been summarized in two reviews.394 In dilute alkali, permanganate is reduced to manganese dioxide, which is pre- cipitated in a hydrated and reactive form, and the possibility of surface-catalyzed reactions creates a major difficulty in interpretation. Reduction to manganese dioxide involves a valency change of three. This is probably never a one-step reduction, and, at least initially, manganese dioxide is thought to be formed by a sequence of reactions such as the following : (a) reduction to manganate (MnOi-) or hypomanganate (MnOi-) ; (b) protonation followed by condensation and disproportionation ; to give, eventually, hydrated manganese dioxide and permanganate.Subsequent stages may be more complex. Protonation is a necessary precursor to con- densation, and does not occur appreciably for manganate in 0.02 M sodium hydroxide and for hypomanganate in 2-0 M sodium hydroxide.6 Condensation almost certainly precedes disproportionation because reaction (1)206 INDUCED OXIDATIONS BY PERMANGANATE is very rapid, and totally favours manganate formati~n,~ and all attempts to pre- pare a monomeric form of MnIV in concentrated aqueous alkali have failed, the disproportionation 2MnIV+MnV+ Mn"' (2) being rapid and ~omplete.~ Many of the complex processes involved in such reactions can be avoided by the use of more concentrated alkali, such that neither manganate nor hypo- manganate are protonated.Under these conditions the oxidizing power of per- manganate lies essentially in its ability to accept either one or two electrons into a vacant, doubly degenerate, antibonding n-type molecular orbital, which may, for a first approximation, be thought of as the dZ2 ; d x ~ y 2 level of the manganese atom. It is possible that permanganate can also act as an oxygen atom donor, being thereby converted into MnO:, but no direct evidence for this intermediate has yet been obtained. It is also possible that permanganate can react by initial formation of a complex, analogous, for example, to the formation of alkyl chrom- ates in the interaction between acid chromate and alcohols.8 The low basicity of permanganate (pKe- 2.25) compared with that of chromate (pK~6.5),6 must strongly disfavour such a step, especially in alkaline solution. However, inter- action with olefins almost certainly proceeds via the formation of a cyclic complex which may be thought of as a di-ester of hypomanganate? Oxidations involving direct attack by permanganate for which the breaking of a C-H bond is the most difficult stage, seem, generally, to result in the transfer of hydride ion.A clear-cut example is the oxidation of benzhydrol, which follows reaction (3), the hypomanganate being then oxidized to manganate : 10 Ph2CHO- +MnOB+Ph2CO+HMnOz-.(3) INDUCED OXIDATIONS In concentrated aqueous alkali, permanganate decomposes to manganate and oxygen, 4Mn0, +40H--+4Mn0~-+2H20+ 02. (4) The rate of loss of permangatate has been followed over a wide range of concen- trations, and was found to be complex.11 We have repeated this work using highly purified reagents, with substantially the same results.12 Initial rates have also been estimated, and although extrapolation to zero time proved to be somewhat uncertain, the results show that the rate at zero time can be described by eqn. (1) : -d[MrlO,]/dt = k1[M~1O,][OH-]. The value for kl is given in table 2. All these results, together with the observa- tion that all the oxygen comes entirely from the solvent,l3 can be explained in terms of a reaction sequence involving only proton and one-electron transfers, the intermediates *OH, -0-, HOT and -02 being treated as by the stationary-state approximation.11 Many alternative reaction schemes have been tested, but none has been found to fulfil all these requirements.The first stage in the proposed mechanism involves the formation of hydroxyl radicals, k t k-i MnO, + OH- + MnOi- + OH, (5) and it has been postulated that several organic compounds, unaffected by cold permanganate in dilute aqueous aIkaIine solution, are oxidized by these radicals when the concentration of alkali is Iarge.l4K . A . K. LOTT A N D M. C. R . SYMONS 207 The main argument in support of this contention was that when optically active carboxylate ions of structure RIR~CH(CH~)&OZ, where R is alkyl or aryl and n = 0 or 2, were oxidized, racemic hydroxyacids, RIR~C(OH)(CH~)~CO~H, were formed in good yield.In dilute alkali, these ions were unaffected by per- manganate at room temperature. Consideration of the possible modes of bond breakage for the tertiary C-H bond 15 leads to the conclusion that radical-ions of structure R1R2iI(CH2)nC02 are intermediates, whereas, had permanganate interacted directly, hydride-ion transfer would have been a far more probable course. Also catalysis by alkali would then be inexplicable since optically-active carboxylate ions were unaffected by concentrated alkali. All these considerations point to the conclusion that there is an induced re- action, represented by (9, followed by (6), .which is rapid and irreversible. Preferential attack on tertiary hydrogen would be expected if hydroxyl radicals are involved, and subsequent reactions of R- would be rapid and should not affect the rate of oxidation. To gain further insight into such induced oxidations, we have measured the rate of oxidation of sodium isobutyrate in concentrated alkaline solution, and compared the results with those for oxygen evolution in the absence of organic material. Many other organic compounds, not attacked by permanganate in dilute alkali, are also oxidized under these conditions, and we selected tert.-butanol for study because it is quite unaffected by most oxidizing agents, but is rapidly attacked by Fentcin’s reagent.16 Other reactions which may also proceed by way of reaction (5) are the conversion of sodium p-isopropyl benzoate into the cor- responding tert.-hydroxy compound,17 and the conversion of benzene into diphenyl.18 EXPERIMENTAL Water was doubly distilled from alkaline pennanganate.Stock solutions of per- manganate were filtered after refluxing and used directly, after appropriate dilution. TABLE 1 .-SUMMARY OF KINETIC DATA [OH-ll 0.997 0-997 0.997 0.997 0.997 0.997 0.610 1 -220 1.520 1.5 3.0 mole 1. [But OHIX 10 mole L-1 2.55 2-55 2-55 1.066 1 -599 2.665 2.132 2.132 2.1 32 [MezCHCO-21 X 10 1.0 0.6 [MnOd-] x lo3 KlU K2a kl x lO3b mole 1 . 4 1.-1 mole min L-1 mole min 1. mole-1 sec-1 1.355 2-710 5.420 2.710 2.710 2-7 10 2.7 10 2.710 2.7 10 6.5 6-3 6-5 6.3 6.4 7.0 6.1 6.6 6.5 1500 - lo00 2.95 1100 2.85 1 600 2-22 1360 3-05 1260 2.83 1280 2.95 1400 2.80 1330 - 1 -64 5.2 341 3.03 1 *64 5.6 332 2.76 average values of K1 and KZ were used in conjunction with eqn.(3) to construct b kl estimated from (-d[Mn04-]/dt)c+O and eqn. (Q. “ theoretical ” curves, all of which were close to the experimental curves. Conccntrations were estimated spectrophotometrically, using the absorption maximum at 526 mp, and also volumetrically. Solutions of sodium hydroxide were prepared by de- cantation from 12 M solutions of Judex Hatch Analyzed pcllcts which had been in contact205 INDUCED OXIDATIONS BY PERMANGANATE with pure alumina for several days. Fresh solutions were prepared for each series of experiments, their concentration being estimated volumetrically after suitable dilution.tert.-Butanol was purified by distillation followed by crystallization and partial melting. Sodium isobutyrate was recrystallized from dilute aqueous alkali. Visible spectra of solutions were measured by means of a Unicam SP. 600 spectro- photometer and, by comparison with the spectrum of pure manganate,s it was established that in all cases conversion to manganate was quantitative. Glass or polythene reaction vessels were completely immersed in a water bath at 25-0f0*05"CY the reactants having been allowed to come to thermal equilibrium before being mixed. The rate was followed by measurement of the optical density of aliquots withdrawn at suitable intervals. Meas- urements were made within 15 sec of removal and it was established that the temperature change during transference and measurement could be neglected.Optical densities were obtained with a Hilger-Spekker spectrophotometer fitted with no. 1 purple filters, and concentrations were then estimated from suitable calibration curves. Results were repro- ducible and independent of the nature of the reaction vessel, and are summarized in table 1, RESULTS In the absence of organic material, decomposition of permanganate in concentrated alkaline solution is approximately third order in permanganate.11 In the presence of a large excess of organic material, oxygen evolution was totally suppressed, and the overall rate of loss of permanganate greatly increased, being, then, very approximately first order in permanganate. However, there were marked initial and finaI deviations from this law, the initial rates being comparable with those in the absence of substrate.Further deviations from any simple law were the marked dependence of the approximate first-order rate constants upon the initial concentration of permanganate and upon the concentration of added manganate. These " first-order " rate constants were also strongly dependent upon the concentration of the organic component, which had to be in large excess to suppress oxygcn evolution. This was limited by the relatively low solubilities of the organic com- pounds in aqueous alkali, and in solutions saturated with sodium hydroxide, it was not possible to suppress oxygen evolution completely. The data have bcen treated in two ways. Initial rates, estimated by the method recom- mended by Livingston 19 were independent of the concentration of the organic component and fitted eqn.(1) closely. Values for kl, calculated by the use of this equation, are given in table 1. There are three alternatives if the kinetically important stages in the overall oxidation are reactions (5) and (6). (i) If k2 were large compared with k-1 the reaction rate would be given by eqn. (I) being independent of the concentration of the organic component. This alternative does not fit the results, and is inherently improbable, since the reversc of reaction (5) must be extremely fast, in order to compete effectively with the other stages postulated for the mechanism of oxygen formation.11 The other extreme would be that stage (6) is slow and rate-determining, ( 5 ) being a rapid equilibrium.Again, this neither agrecs with the results, nor is it to be expected, since oxygen evolution is totally suppressed. The steps leading to oxygen formation in the absence of organic material must all have large rate constantq11 and if (5) is to compete successfully, k2 cannot be small. Accordingly, we have applied the stationary-state approximation to (5) and (a, treat- ing the hydroxyl radical as an unstable intermediate. This gives for the rate of reaction. tends to eqn. (1). On integration (2) becomes This is in accord with the initial rate studies, since as t-20, [Mn042-]->O, and eqn. (2) t = - where K1 = l / k l and K2 = k-llklkz. Values for Kl and K2 were obtained by solving sets of simultaneous equations and were found to bc ressonably constant ovcr the range 0 to 80 Mean values for each r u n are given in table 1 and the mean of these reaction.K .A. K . LOTT AND M. C. R . SYMONS 209 results for tert.-butanol and sodium isobutyrate have been used to calculate kl and k2lk-1. These results are given in table 2, together with values for kl obtained from initial rate s tudies.12 TABLE 2 . 4 U M M A R Y OF RATE CONSTANTS experiment kl x 103 1. mole-1 sec-1 initial rates : no organic component : 12 2.1 10.5 ,, ,, +sodium isobutyratc : 2-89L05 ,, ,, + tert.-butanol : 2.95 104 from K1+ sodium isobutyrate : 3.08 k0.5 16.0 ,, ,, + tert-butanol : 2-57 f0.4 4-9 total average : 2.72 RESULTS FOR kl In the presence of organic material, kl varies within the limits 2.87f0.5 x 10-3 1.mole-1 sec-1. These results overlap with those obtained from studies in the absence of organic compounds, namely, 2.1 f0.5 x 10-3 1. mole-1 sec-1,12 although on average, they are ap- preciably higher. We have tested a variety of alternative mechanisms for these oxidations but have bccn unable to discover any which can accommodate the highly characteristic kinctic data. In view of this and of the difficulty experienced in extrapolating to zero time in the absence of organic material,l2 we are inclined to the view that the discrepancy is caused by experimental errors. The results show that the forward stage of reaction (9, oxidation of hydroxide ion, is very slow compared with such direct reactions as attack on benzhydro1(3),10 or hydrogen.20 Hence, induced reactions of this type would not normally compete if direct interaction is possible.Although the overall course of the reaction with sodium isobutyrate is known,l4 attack being confined to tertiary hydrogen, that with tert.-butanol is not. The overall stoichio- metry has not yet been established, but the first step is almost certainly attack on 13-hydrogen to give *CH2C(Me)20H. Electron-spin resonance results established that these radicals were formed when solid tert.-butanol containing traces of hydrogen peroxide was irradi- ated with light from a high-pressure mercury arc,21 and fair yields of the dimer, (.CH2C(Me)20H)2, were isolated after reaction with Fenton's reagent22 Coffman et a1.22 found that when isobutyric acid reacted with Fenton's reagent in dilute aqueous sulphuric acid, 13-, rather than a-hydrogen atoms were attacked.They interprcted this and other similar results by postulating that hydroxyl radicals are electro- philic and that the induction effect of the carboxyl group is more effective than the fact that tertiary C-H bonds are weaker than primary. However, the carboxylate group will have an opposite inductive effect, and should facilitate attack on a-hydrogen. It is possiblc that tert.-butoxide is formed in low concentration and that the initial reaction involves electron transfer to give manganate and the radical Me3CO.. However, tert.-butoxide is a very strong base, relative to hydroxide, and the great similarity between this reaction and oxidation of sodium isobutyrate suggests that they proceed by similar mechanisms.RESULTS FOR k2 Since k-1 is a comnion factor, relative results, k2 (tert.-butanol)/k2(isobutyrate) canlbe estimated. The result, that isobutyrate is attacked about four times more rapidly than tert.-butanol, is in accord with the observation that isobutyrate is attacked specifically on thc tertiary hydrogen. No value for k-1 is available so an absolute figure for k2 cannot be given. However, k-1 is undoubtedly very large, so k2 must also be large. This is in accord with calculations made from the data pertaining to the decomposition of per- manganate in the absence of organic rnaterial,*z which suggest that k2 for attack on iso- butyrate and k2' for the second stage in the mechanism leading to the formation of oxygen, k2' OH+ OH- +* 0- + H20 (7) do not cliffer by more than a factor of 102.210 INDUCED OXIDATIONS BY PERMANGANATE Acknowledgement is made to the University of Southampton for a main- tenance grant for one of us (K. A. K. L.). 1 Tompkins, Trans. Farday SOC., 1942,38, 128, 131 ; 1943,39, 280. 2 Wiberg and Stewart, J. Amer. Chem. SOC., 1955,77, 1786; 1956, 78, 1214. 3 Ladbury and Cullis, Chem. Rev., 1958, 58, 403. 4 Waters, Quart. Rev., 1958, 12,277. 5 Carrington and Symons, J. Chem. Soc., 1956, 3373. 6 Bailey, Carrington, Lott and Symons, J. Chem. Sac., 1960, 290. 7 Lott and Symons, J. Chem. SOC., 1959, 829. 8 Westheimer, Chem. Rev., 1949,45,419. 9 Wiberg and Saegebarth, J. Amer. Chem. SOC., 1957,79, 2822. 10 Stewart, J. Amer. Chem. SOC., 1957, 79, 3057. 11 Symons, J. Chem. SOC., 1953,3956. 12 Carrington, Lott and Symons, unpublished work. 13 Symons, J. Chem. SOC., 1954, 3676. 14 Kenyon and Symons, J. Chem. SOC., 1953,2129, 3580. 15 Lewis and Symons, Quart. Rev., 1958, 12, 230. 16 Merz and Waters, J, Chem. SOC., 1949, S . 15. 17 Bryce-Smith, J. Chem. SOC., 1954, 1079. 18 Jezowska-Trzebiatowska, Nawojska arid Wronska, Bull. h a d . yolon. Sci., 1954, 2, 19 Livingston, Rates and Mechanism of Reaction, ed. Weissberger (Interscience Inc., New 20 Webster and Halpcrn, Trans. Faraday SOC., 1957, 53, 51. 21 Gibson, Symons and Townsend, J. Chem. Soc., 1959,269. 22 Coffman, Jenner and Lipscomb, J. Amer. Chem. Soc,, 1958, 80, 286.4. 447. York, 1953).

ISSN:0366-9033    

DOI:10.1039/DF9602900205

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

THE MECHANISM OF THE PERMANGANATE OXIDATION OF FLUORO ALCOHOLS IN AQUEOUS SOLUTION BY Ross STEWART AND R. VAN DER LINDEN Dept. of Chemistry, University of British Columbia, Vancouver, Canada Received 1 st February, 1960 The mechanism of the aqueous permanganate oxidation of a series of aromatic fluoro alcohols has been studied. The reaction rate is proportional to the concentration of alkoxide ion and this, together with the salt effect and the activation entropies, supports the idea that the reaction takes place between alkoxide and permanganate ions. A very large reduction in rate occurs for the oxidation of the 1-deutero compounds (16 : 1). The effect of aromatic ring substituents on the rate of oxidation is slight, assuming the reaction proceeds via the alkoxide ion. All the evidence except the effect of substituents on the rate is in agreement with a hydride transfer mechanism fram alkoxide ion to permanganate ion.Various other possible mechanisms are considered. Permanganate is a powerful oxidizing agent which, not surprisingly, shows a variety of reaction paths in its reactions with organic and inorganic compounds.l.2 Previous work in this laboratory 3 has shown that one of these, the oxidation of the secondary alcohol benzhydrol, PhzCHOH, to benzophenone by aqueous per manganate, proceeds cleanly with the following kinetics : V = k[Ph,CHOH][MnO,][OH-]. The following additional facts are known: a kinetic isotope effect of 6.6 exists for the oxidation of Ph2CDOH; the entropy of activation is large and negative; a positive kinetic salt effect exists, and there is no transfer of oxygen from per- manganate to the organic substrate during the reaction.These facts all suggest that the reaction path is the following with the rate-controlling step being a transfer of hydride ion from alkoxide ion to permanganate ion. Ph2CHOH + OH' +Ph,CHO- + H20 fast, HMnOi-+MnO; +OH--+2Mn0~-+H20 fast. The acidity of this alcohol is such, however, that it can be only slightly ionized under the reaction conditions. We have accordingly turned to fluorinated com- pounds in order to obtain alcohols of sufficiently high acidity that we can examine in more detail the first two steps above if, indeed, the validity of the overall mechanism appears to be c o n w e d by this further work. A series of aryl trifluoromethyl carbinols proved to be satisfactory since they are highly ionized in 0.1 N alkali and are oxidized cleanly to the corresponding ketones by permanganate : ArCHOHCF3+2Mn0, +20H-+ArCOCF3+2Mn0:- +2H20.Ph2CHO- + MnO,'+Ph,CO+HMnO~- slow, EXPERIMENTAL The preparation of the Auoro alcohols and ketones has been reported elsewhere.4.5 The reaction was followed by an iodometric procedure similar to that previously reported.% 6 Thc only important difference was that the precision of the method was improved by 21 1212 PERMANGANATE OXIDATION OF ALCOHOLS using separate solutions for each point on the rate plot rather than by removing aliquots from a single solution. For reactions in solutions more basic than pH 11.7 a 2 : 1 ratio of permanganate to alcohol was used.For more acidic solutions a 2 : 3 ratio was used in accordance with the altered stoichiometry. Phosphate buffers were used and potassium nitrate and potassium sulphate used to vary the ionic strength. The second-order rate constants k2 were determined using the following equation when a 2 : 1 ratio of reactants was used : where VO and V, are the volumes of thiosulphate required at time zero and at time t, and [alc]~ is the initial concentration of alcohol. When the 2 : 3 ratio is used the final term becomes ( VO - V{)/( Vt - {- VO). An examination of the ultra-violet spectra of the solution after oxidation of C~HSCHOHCF~ in 0.1 N NaOH under the same conditions as were used for the kinetic experiments showed that the ketone C6HsCOCF2 was present in the solution in greater than 95 % yield. The ketone was found to be resistant to further oxidation and to hydrolysis under the reaction conditions.A few experiments which were performed with the primary alcohol, HCFzCFzCHzOH, gave results very similar to those reported herein. The rate in 0.18 N NaOH was 2.2 1. mole-1 sec-1; AH* = 10.2 kcal/mole ; AS+ = -22 cal/mole deg. ; alcohol pK = 12.1. RESULTS AND DISCUSSION KINETICS At constant hydroxyl ion concentration the reaction is second order as can be seen in fig. 1, the second-order rate constant k2 being insensitive to variations in reactant concentration. Changing the hydroxyl ion concentration changes the rate in the manner shown in fig. 2. The pH values at the midpoints of these curves all correspond quite closely to the known pK values of the alcohols.4 These range from 12.2 for the p-CH30 compound to 11.2 for the m-NO2 compound. A positive salt effect is observed for the oxidation of m-N02Cfl&HOHCF3 in 0.01 N NaOH at 25".The rate increases from a value of 4.54 1. mole-1 sec-1 at ionic strength 0.0115 to a value of 5.90 1. mole-1 sec-1 at ionic strength 0-122. The above results are all consistent with the following reaction path : ROH+OH-~RO-+H,O, (2) (3) -d[MnO,]/dt = k[RO-][MnO;]. (4) -d[MnO,]/dt = k,[alc][MnO;], ( 5 ) kz[alc] = k[RO'], ( 6 ) k RO- + MnO; +products. The rate law corresponding to this mechanism is Since eqn. (1) was derived from the expression, where [alc] = [ROH]+ [RO-1, it follows that and under conditions of almost complete ionization of the alcohol the experimental rate constant k2 becomes identical with k, the rate constant of the rate-controlling step.It is interesting that the permanganate-formate reaction kinetics are exactly analogous.7-9I R . STEWART A N D R . VAN DER LINDEN i 21 3 4 8 12 t , min FIG. 1.-Second-order rate plots, [alc]~ = 5.1 x 10-4, [OH-] = 0.2, T = 25.2" ; 0 GH5CHOHCF3 ; c) C6HsCDOHCF3 (multiply units on both axes by four for latter). [OH-] FIG. 2.Variation of rate with basicity ; CsH5CHOHCF3 ; 0 m-NOz ; CD m-Br; A p-CH3; 63 C~H~CDOHCFJ.214 PERMANG ANATE OXIDATION OF ALCO I-IOLS A termolecular step involving ROH, MnOz and OH- is also consistent with the kinetics since the rate law - d[ Mn O,]/dt = I&[ ROH][ OH'] [MnO,]/[ H, O] (7) is equivalent to the one above. This possibility will be discussed in a later section. THERMODYNAMIC FUNCTIONS Fig.3 shows a plot of logl&/T) against l/Tfor the reaction in 0.1 N NaOH. At this basicity, k 2 ~ k and one can calculate the activation parameters for what we have assumed to be the rate-controlling step, i.e., the reaction between alkoxide ion, RO-, and permanganate ion. These values are listed in table 1 for the parent compound, its deutero analogue, and its m-nitro derivative. TABLE 1 .-HEATS AND ENTROPIES OF ACTIVATION 0.1 N NaOH, [alc]~ = 5.1 x 10-4 compound AH4, kcal/mole AS', cal/mole deg. CsH5CHOHCF3 9.1 & 0.3 -24.3f1 C~HSCDOHCF~ 11.3 - 22.4 rn-NO2C6H4CHOHCF3 9.6 -23.8 The heats of activation are similar to those observed in other permanganate oxidations 3 ~ 7 - 9 and the entropies of activation are reasonable for a reaction between two anions.10 1 I I I I 1 3.2 33 3.4 3.5 1000/T"K FIG.3.Variation of rate with temperature ; 0 C6HsCHOHCF3, y = 3 ; CsH5CDOHCF3, y = 4. ISOTOPE EFFECT A large kinetic isotope effect (16.1) is observed for the oxidation of the deuter- ated compounds, ArCHOHCF3. This immediately disposes of an electron transfer step from the alkoxide ion to permanganate ion. The following mechanism (hydride transfer) requires that there be an isotope effect but the values listed in table 2 are unusually large :R . STEWART AND R . VAN DER LINDEN 215 0 2 - ll +Ar-C-CF, I n I .4r-C-H CF3 transition state CF3 +HMnOi- (8) TABLE 2.-DEUTEKIUM ISOTOPE EFFECTS AT 25" kz, 1. mole sec-1 kH kD kHlkD P H C6H5CDOHCF3 13.3 13.0 12-0 11.0 7.0 * 1.0 * 13.0 13.0 m-BrC6H4CDOHCF3 13-3 P-CH~C~H~CDOHCF~ 13-3 7.5 7.6 3.8 0.18 -00394 40096 11.6 t 2.85 $ 7.6 8.4 0.47 *47 -23 -01 53 -000242 -00055 a88 j- 0155 $.-47 -52 16.0 16.2 16.5 12.0 16.0 1.8 13.2.1. 18.4 $ 16.2 16.1 * The reactant concentrations were considerably higher in these experiments. j- T = 38.0". T = 12-65", A high isotope effect is also observed in neutral solution even though the reaction is now very slow. In acid solution the reaction is still very slow but a low isotope effect is here observed. Normal deuterium isotope effects, which result from the loss of the C-H stretching mode in the transition state, are less than half this size at the same temperature.11 Four possible explanations for the anomalous effect present them- selves.(a) The C-H and C-D stretching frequencies might themselves be anomalous. This possibility can be rejected by an examination of the infra-red spectra of these compounds. For the three pairs of compounds studied CC-H is at 2895 cm-1 and i c - ~ is at 2150 cm-1, both normal values for carbon-hydrogen bonds. (b) Some unusual consecutive process involving possibly chain branching might produce a cumulative isotope effect. This, however, is difficult to reconcile with the " clean " kinetics and the similarity of the heats and entropies of activ- ation to those of other permanganate oxidations. (c) The reaction may involve loss of the bending modes in the transition state in addition to the loss of the stretching mode. This might be caused by the reactant molecules being widely separated in the transition state as a consequence, in the present case, of electro- static repulsion.The benzhydrol+ permanganate reaction 3 and the formate ion+ permanganate reaction 79 9 which resemble this reaction in many ways have, however, normal isotope effects, (d) Quantum-mechanical tunnelling which is more probable for protium than for deuterium may be occurring.13 Further work is required to see whether mechanistic or structural factors (presence of a CF3 group, for example) are responsible for the large effect reported here. Re- gardless of the cause, it is clear that the carbon-hydrogen bond is being broken in the rate-controlling step. It is of interest that the manganate oxidation of these compounds, although slaw, also gives large effects.216 PERMANGANATB OXJDATION OF ALCOHOLS The effect of deuterium substitution on the ionization of the alcohols is slight.The pK of C~HSCHOHCF~ and that of its deutero derivative are within 0.1 pK units of one another. EFFECT 01' KING SUBSTITUENTS The principal advantage of using fluoro alcohols to study the permanganate. alcohol reaction is that the first two steps of the suspected mechanism can be studied separately. The effect of ring substitution on the ionization step is similar to that for the benzoic acid ionization and shows a satisfactory Hammett linear free-energy relation 12 with a p value of 1.01.4 In 0.2 N NaOH the alcohols are almost completely ionized and k 2 z k A hydride transfer from the alkoxide ion to permanganate should result in a negativep value, i.e.a p-methoxy substituent should accelerate the process and a p-nitro group retard it. Fig. 2 shows that the reaction rate is, surprisingly, only slightly affected by nuclear substitution and that the small variations in rate which do occur (curve A) appear to be not linearly related to the Hammett substituent constant 0. (These rates are corrected for incomplete ionization of the alcohols in 0.2 N NaOH.) Several explanations for this situation can be considered. First, a simple hydride transfer may in fact be occurring with the electronic effect of a distant group in the molecule being unimportant. Secondly, two different processes with different electronic requirements may be occurring, one with a positive, the other with a negative, p.This would account for the curved plot but it is not certain that the small curvature is significant and this possibility will not be pursued here. These mechanisms can be written as follows : Thirdly, one of several termolecular reactions may be taking place. 0. (8) 1 Ar e n l n I H,O @-C-O- Mn0,-+H30++Ar-C-CF3+MnO:- - CF3 0. 1 n I A Ar-C-CF, + MnO,*'Sproducts. OH I Ar HO- @-C-OH MnO~-+H,O+Ar-C-CF, +MnO:- (5) I * * I . CF3 OH Ar-C-CF, + MnO,*%products, 0 Ar Ar f-31 n I HO- C--H MiiO,--+HO-C + HMnO?-. /\ HO CF3 /\ HO CF3 In reaction (8) the permanganate acts as an electron abstractor with a water moIecuIe acting simultaneously as a proton abstractor. This mechanism is inR . STEWART A N D R . VAN DER LINDEN 217 agreemcnt with the kinetic and isotopic evidence and since the alkoxide ion is losing both an electron and a proton in the transition state the effect of substituents should be slight.Reaction (9) is kinetically equivalent to (8) but hydroxyl ion removes the proton as the permanganate abstracts the electron from the hydroxyl group of the neutral substrate. (The permanganate cannot be removing a hydrogen atom or a hydride ion from the hydroxyl group since there is no appreciable change in rate when a 60 % D20+40 % H20 solvent is used.) It would be difficult to predict the p value for this reaction. It is of interest, however, that a reasonably linear relation is obtained when is now plotted against the logarithm of the rate constant (fig. 4, curve B). This treatment requires that the rate in 0.2 N NaOH be corrected for partial ionization rather than for incomplete ionization as before.I I 1 I I 1 -0.4 -0.2 0 0 . 2 0.4 0.6 0.8 U FIG. 4.-Hammett plot, T = 25*2", [OH-] = 0-2 ; 0 ki = k2[alcJ/[RO-] ; 0 ki = k2[alcl/lROH]. Eqn. (10) shows hydride transfer from the neutral alcohol being aided by dis- placement by hydroxyl ion to produce the stable hydrated ketone.4 Of the three termolecular mechanisms considered here, (8) does not require a termolecular solute collision and, in addition, it accommodates satisfactorily the rest of the evidence accumulated about this reaction. In summary, there are two unusual features of the permanganate oxidation of aryl trifluoromethyl carbinols, viz., the size of the deuterium isotope effect and the effect of aromatic substituents on the rate. If a simple hydride transfer from alkoxide ion to permanganate ion is rejected on the basis of the negligible effect of substituents one is forced to conclude that a termolecular process is occurring. The most attractive of these alternatives is that in which a permanganate ion and a water molecule remove an electron and a proton respectively from the alkoxide ion. The financial support of the National Research Council of Canada is gratefully acknowledged.21 8 PERMANCANATE OXIDATION OF ALCOHOLS 1 Ladbury and Cullis, Chem. Rev., 1958, 58, 403. 2 Waters, Quart. Rev., 1958, 277. 3 Stewart, J. Amer. Chem. SOC., 1957, 79, 3057. 4 Stewart and Van der Linden, Can. J. Chern., 1960, 38,399. 5 Stewart and Van der Linden, Tetrahedron Letters, 1960, OOOO. 6 Wiberg and Stewart, J. Amer. Chem. SOC., 1955, 77, 1786. 7 Wiberg and Stewart, J. Amer. Chem. SOC., 1956, 78, 1214. 8 Hill and Tompkins, Trans. Roy. SOC. S. Africa, 1943, 30, 59. 9 Taylor and Halpern, J. Amer. Chem. SOC., 1959, 81, 2933. 10 Frost and Pearson, Kinetics and Mechanism (Wiley, New York, 1953), p. 132. 11 Wiberg, Chem. Rev., 1955,55,713. 12 Hammett, Physical Organic Chemistry (McGraw-Hill, NewjYork, 1940). 13 Bell, Fendley and Hulett, Proc. Roy. SOC. .4.,!1956, 235, 453.

ISSN:0366-9033    

DOI:10.1039/DF9602900211

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

THE EFFECT OF ACID AND ALKALI ON HYDROGEN TRANSFER FROM ISOPROPANOL AND DIISOPROPYE ETHER IN DILUTE AQUEOUS SOLUTION BY C. F. WELLS* British Rayon Research Association, Heald Green Laboratories, Wythenshawe, Manchester, 22 Received 12th February, 1960 In the photosensitized autoxidation of alcohols and ethers the oxidative reactivities of aqueous solutions of isopropanol and diisopropyl ether are reduced by the addition of acid and that of isopropanol by the addition of alkali. A study of the product yields and the kinetics under these conditions suggests that the retardation in acidic solutions is caused by the formation of protonated substrate molecules which are unreactive to the photo- excited sensitizer, and that the retardation in alkaline solutions is caused by the alcoholate ion being unreactive to the photo-excited sensitizer.Values are obtained in dilute aqueous solution for the equilibrium constants for the association of a proton with isopropanol, cyclohexanol and diisopropyl ether, and for the equilibrium constant for the dissociation of a proton from isopropanol. These equilibria are discussed in relation to solvation It has been shown 1-3 that the reaction between alcohols RH and the photo- excited anthraquinone-2-sulphonate ion Q* in neutral aqueous solution follows the equation, where the rate constants are designated as follows : Q+hvAQ* Q * ~ Q Q* + RH%QH + R. (1) Reaction (1) has been shown 1-3 to be a direct hydrogen-atom transference from a C-H bond a to an OH group and reaction (0) is a non-radiative, adiabatic transfer from the excited state to the ground state of the quinone.le2 The radicals -QH and R* react rapidly with oxygen to give HOy and ROy radicals which interact and decompose to give the stable products.4~ 5 The quantum yield has been shown to be unity for ethanol 4 and for isopropanol.1 Addition of acid and alkali retards the rate of oxygen absorption.Fig. 1 shows that plots of oxygen uptake against time in solutions containing HzSO4 and NaOH are similar to those observed in neutral solution 1 where the deviation from linearity is caused by light absorption of the coloured products formed in a side reaction of the sensitizer.1 Similar retardations are produced by the addition of other acids. The kinetics and product yields of the oxidation at 0°C of iso- propanol and diisopropyl ether in acidic and alkaline solutions and of cyclohexanol in acidic solution have now been studied in some detail, using sodium anthra- quinone-2-sulphonate as a photosensitizer.219 * present address : Chemistry Dept., The University, Birmingham, 15.220 EFFECT OF ACID AND ALKALI O N HYDROGEN TRANSFER time (min) FIG. 1,-Plot of oxygen absorption against time for aqueous solutions of isopropanol at 0°C. [Sensitizer] = 3 x 10-4 M, [isopropanol] = 021 M : 0, neutral ; El, 0.886 N H2SO4 ; 8,0.266 N H2SO4 ; X ,0.266 N H2SO4,0.4 M o\IH4)2so4 ; V ,0-03 N NaOH. EXPERIMENTAL MATERIALS Isopropanol was purified by refluxing with 2,4dinitrophenylhydrazine and hydro- chloric acid and subsequent fractional distillation.1 Diisopropyl ether was first shaken with solutions of sodium bisulphite and ferrous sulphate 6 and fractionally distilled under nitrogen : the ether was stored under nitrogen, but soon decomposed giving acetone and hydrogen peroxide as products and had to be repurified; aqueous solutions prepared from this ether and stored under nitrogen were stable for only very short periods, and it was found necessary to prepare a fresh solution each day to obtain reproducible results.Cyclohexanol was given the same treatment as isopropanol, followed by crystallization from the pure liquid. Sodium anthraquinone-2-sulphonate was purified as described elsewhere.1 Anthra- quinone-2-sulphonic acid was prepared in solution by passing a 10-2 M solution of the sodium salt down a column of suitabIe length packed with IR .120(H) ion-exchange resin. Picric acid was recrystallized from water. The hydrochloric acid used in the calibration of the method of determining total hydrogen ion conccntration was A.R. grade. Water was distilled once : for the work in alkaline solution all solutions were prepared in distilled watcr which had been boiled out and cooled down and stored under nitrogen, and all vessels were flushed out with nitrogen before putting solutions into them. Sul- phuric acid, potassium bisulphate and ammonium sulphate were all of A.R. grade. Solu- tions of sodium hydroxide were prepared by dilution of B.D.H. " carbonate-free " standard solution with boiled-out distilled water. PROCEDURE Solutions contained in silica vessels immersed in a thermostat were irradiated by light from a monitored Hanovia lamp which passed through a Chance O.X.1 filter : a constant light intensity was used in the experiments described in this paper; rates of oxygenC. F.WELLS 22 1 absorption were determined manometrically. The apparatus and procedure are described in detail elsewhere; 1 for the experiments in alkaline solution, oxygen admitted to the apparatus was first passed through lime water and the manometric apparatus was kept filled with CO2-free oxygen. 0°C was obtained by using melting ice 1 and 25°C by using a Hg+CC14 regulator, 250 W heater and a Sunvic E.A.3 control.1 ANALYTICAL METHODS Hydrogen peroxide was determined by titration with standard ceric solutions at 0°C. Acetone and cyclohexanone were determined spectrophotometrically using 2,4dinitro- phenylhydrazine ; 1, 7 all volatile material was first removcd from the sensitizer by dis- tillation under high vacuum.ESTIMATION OF TOTAL HYDROGEN-ION CONCENTRATION This was determined spectrophotometrically using picric acid. The picrate ion is coloured yellow and has a maximum absorption at 360 mp which obeys Beer's law : the undissociated acid is colourless in the visible region. When acid is added to solutions of picric acid, the hydrogen-ion concentration is given by Do-D D [H'] = Kp---, where D and DO are the optical densities with and without added acid and Kp is a calibration constant. Measurements were made at 0°C by filling the cell compartment of the Unicam SP 500 with dry air before each measurement to prevent condensation on optical faces of the cell.The method was calibrated at 0°C using HCl : Kp = 1.19 at 0°C. RESULTS AND DISCUSSION PRODUCT YIELDS The yields of acetone and hydrogen peroxide from isopropanol are both 100 % in neutral solution, except for slight deviations which occur at low alcohol con- centrations and long extents of reaction 1 (percentage yield = no. of moles of isopropanoll M 0.314 0.314 0.167 0.167 0.167 1.20 0-314 1 -20 0.167 0.167 0.167 0.314 0.314 1 *20 1 -20 1 -20 1 -20 1 *20 1 -20 1 -20 1 -20 1 -20 TABLE 1. CONCENTRATION OF SENSITIZER = 3 X 10-4 M ['acid] or 0 2 absorption, moles % yield of [alkali] per 10 mlx 106 HzOz neutral 21-7 96.5 Y Y 63-4 94 Y9 6 100 Y Y 20 92 Y, 63 89 64 96.5 0.625 GHzS04 65.1 96.0 0.625 N HzSO4 66-8 98.5 0.01 M OH- 5 90 0.01 M OH- 12 80 0.01 M OH- 44 73 0.01 M OH- 6 91 0.01 M OH- 58 82 0.01 M OH- 17 97 0.01 M OH- 63 94 0005 M OH- 20 98 0.125 M OH- 16 99 0.015 M OH- 13 98 002 M OH- 10 103 0.03 M OH- 10 96 0.04 M OH- 9 91 0.08 M OH- 6 92.5 % yield of acetone 98.5 95 100 94 91 99.5 95.5 97 95 84 73 80 98 95 98 101 102 102 94 97 - - product x 100/no.of moles of 0 2 absorbed). Table 1 shows that the yields of ace- tone and hydrogen peroxide are also both 100 % in acidic solutions. In alkaline222 EFFECT OF ACID A N D ALKALI O N HYDROGEN TRANSFER solution, yields of acetone and hydrogen peroxide are both 100 % at short extents of reaction, but deviations occur with increasing extents and increasing concen- tration of alkali although the yields of acetone and hydrogen peroxide always remain equal : examples of these yields are given in table 1 and compared with yields in neutral solution.The following observations can be made. (a) With constant sensitizer concentration and at any particular concentration of alkali the yields remain equal but decrease from 100 % with decreasing concentration of isopropanol. (6) At any particular isopropanol concentration and alkali concentration yields remain equal but decrease from 100 % with increasing extent of reaction. (c) With constant sensitizer concentration at constant isopropanol concen- tration the yields remain equal but decrease from 100 % with increasing alkali concentration at constant extent of reaction. Observations (a) and (b) closely parallel effects observed under extreme con- ditions of long extents of reaction and high [Q]/[O2] ratios at low concentrations of isopropanol in neutral solution,l where, from the manner of the variations, the deviations have been attributed to the effect of a side reaction of the excited sensitizer which gives coloured products from the sensitizer and probably absorbs oxygen.It would appear from the above observations that this side reaction is accelerated in alkaline solution. Thus the effect of alkali and acid on the rate of oxidation of isopropanol does not operate on the radical reactions and must be on the excitation of Q or on (0) and (1). The product yields from cyclohexanol are 100 % cyclohexanone and 100 % hydrogen peroxide, both in neutral and acidic solutions. Table 2 shows that the yields of acetone and of hydrogen peroxide are both 100 % from diisopropyl ether in neutral and acidic solutions : short extents of TABLE 2 [SENSITIZER] = 3 X lop4 M ; [DIISOPKOPYL ETHER] = 0.12 M [acid] or [alkali] neutral 0.9 N H2S04 10-5 M OH- 5 x 10-5 M OH- 10-4 M OH- 10-3 M OH- 10-2 M OH- 2x 10-2 M OH- 4 x 10-2 M OH - % yield of H202 92 93 95 92 88 74 75 71 76 % yield of acetone 88 97 92 70 69 67 70 66 66 reaction are used to minimize the effect of the side reaction of the excited sensitizer as the concentration of ether is so low because of its low solubility.These yields decrease sharply in alkaline solutions (table 2); the yield of acetone decreases to 70 % abruptly between 10-5 M OH- and 5 x 10-5 M OH-, and that of hydrogen peroxide to between 70-75 % between 10-4 M OH- and 10-3 M OH-.These effects are quite different from those observed for isopropanol, and the sharpness of the decreases suggests effects of acid-base equilibria on the reactions of the radicals formed in reaction (1) and subsequent reactions. REACTION IN ACIDIC SOLUTION Fig. 1 shows that the rate of oxidation of isopropanol is unaffected by the addition of (NI44)2S04 in acidic solutions, so it is concluded that there is no eKectC . F . WELLS 223 l/[isopropanol] (moles-1 1.) FIG. 2.-Plots of (-d[02]/dt)-l against l/[isopropanol] in acidic solutions at 0°C. [Sen- sitizer] = 3 x 10-4 M : X , neutral ; A, 0.133 N H2SO4 ; El, 0-266 N HzSO4 ; 7 , 0 * 6 2 3 N HzSO4; 0, 0.886N H2SO4; #, 1.315N H2S04; 8, 1.615N HzS04; a, 0.60M KHS04; A, 1.00 M KHS04. Ol 0 S 10 I S 2 0 l/[cyclohcxanol] (moles1 1.) FIG.3.-Plots of (-d[Oz]/dt)-' against l/[cycIohexanol] in acidic solutions at 0°C. [Sen- sitizer] == 3 ?: 10 4 M : 0, neutral ; a, 0.624 N HzSO4; A, 1.30 N H$304; v, 2.00 N J42SO4 ; X , 2.88 N H2SO4.224 EFFECT OF ACID AND ALKALI ON HYDROGEN TRANSFER of ionic strength on the reaction in acidic solution and no further account is taken l/[diisopropyl ether] (moles-1 1.) FIG. 4.-Plots of (-d[O,]/dt)-l against l/[diisopropyl ether] in acidic solutions at 0°C. [Sensitizer] = 3 x 10-4 M : M , 4 . 9 N &So4 ; of it. Plots of (-d[O~]/dt)-l against l/[substrate] all gave straight lines for isopropanol, cyclohexanol and diisopropyl ether (fig. 2, 3 and 4), all intercepting A, neutral ; 0, 4 - 4 5 N H2SO4 ; #, -1.2 N H2SO4; D, -1.8 N H2SO4; n, -2.5 N H2S04.I I 0 1.0 1.5 2 . 0 I O L . I .. - __ I 0 5 [H+]T (moles/l.) FIG. 5.-Plots of SIT against [H+-]T for isopropanol, cyclohexanol and diisopropyl ether at 0°C. [Sensitizer] = 3 x 10-4 M : a, isopropanol with H2SO4 ; A, isopropanol with KHS04; 0, cyclohexanol with H2SO.4; 8, diisopropyl ether with H2SO4. the ordinate at the same point as obtained in neutral solutions with these substratesC. F . WELLS 225 and a wide range of other substrates at this light intensity: 1 the variation with acidity operates on the second term of eqn. (a). Plots of the ratio SIT of the slopes S and intercepts T obtained from fig. 2, 3 and 4 against the total hydrogen ion concentration, [H+]T, in fig. 5 are all straight lines intercepting the ordinate at their respective values of SIT in neutral solution.The total hydrogen ion concen- tration was obtained from the concentrations of sulphuric acid in the absence of added substrate by the picric acid method discussed above. From fig. 5, where B is a constant for each substrate. acid. The spectrum of the sensitizer was found to be unaltered by added sulphuric POSSIBLE MECHANISMS As the retardation is not caused by an effect on the reactions of the free radicals produced in the hydrogen transfer (l), the following possibilities exist : (i) an acceleration by H i 0 ions of the first-order energy degradation (0) ; (ii) an acid-base equilibrium involving the unexcited sensitizer ; (iii) an acid-base equilibrium involving the photo-excited sensitizer ; (iv) an acid-base equilibrium involving the substrate alcohol or ether.(i) is unlikely as H20 is known to be a poor deactivator of excited states, and moreover, fig. 1 shows that added NHf ions, which have a structural resemblance to H i 0 ions, have no effect on the reaction. (ii) is unlikely as eqn. (a) shows that it would involve a change of intercept of the plots of (-d[02]/dt)-l against l/[substrate] with acidity, as well as a change of slope, which should give SIT invariant with acidity, if the sensitizer associated with the proton is not a photo- sensitizer ; and if the latter is a photo-sensitizer, straight lines are unlikely to be obtained for plots of (- d[Oz]/dt)-l against l/[substrate]. (iii) is unlikely as this would also involve a change of intercept withacidity whether this operated without an intermediate complex, [QH"l Q* 1 CH + I Q* + H+ +QH+ K = K Q* + R H k QH + R.or through a complex C, ki k' Q* + RH+C Kh c + H+ +CH+ C k QH + R. Applying stationary-state conditions, we have without the complex, H226 EFFECT OF ACID A N D ALKALI ON HYDROGEN TRANSFER and with the complex, C, where k = k'+k,, and in both cases [H+] = [H+]T. To check whether the (-d[O2]/dr)-l intercept is indeed invariant with acidity some experiments were done at high acidities. Fig. 6 shows that the plots of (-d[O2]/dt)-l against l/[isopro- pan011 at high acidities intercept the ordinate at the same point as obtained in neutral solution, within experimental error. In particular, no increase in the intercept with m+] of the order expected by the kinetic equations for (iii) is possible : therefore possibility (iii) is eliminated.-_ 1 i i i i ' i b ; C l/[isopropanol J (moles-1 1.) FIG. 6.-Plots of (-d[O2]/dr)-l against l/[isopropanol] at high [HzSO~] at 0°C. [Sen- sitizer] = 3 x 10-4 M : 0, neutral ; A, 276 N H2SO4; El, 5.52 N H2SO4. (iv) would involve the formation of an oxonium ion of the substrate which is less reactive to the excited sensitizer than the substrate itself. Kb f ROH + H t O+ROH, + H,O + ki Q* + ROH2+ QH + H + +radical, + with k; < k,. and assumingC . F . WELLS 227 over the range of concentrations used here, where f represents activity coefficient, Kb will be constant under these conditions. Applying the stationary-state conditions, The total concentration of added alcohol [ROHIT is given by + [ROHIT = [ROHI + [ROH,] = [ROH](l +(Kb[H~O]/[HzO])) and substituting (d) into (c), and intercept is invariant with acidity, as observed.From (e), The total concentration of added hydrogen ions [H+]T is given by and substituting for [Hi01 in (f) from:(g), (Kb[R6H]/[H20]) = ([ROHd/[HfO]) and must be small for (- d[Oz]/dt)'l against ~/[ROH]T to be linear, so that + and Comparing eqn. (h) with eqn. (b), and values of Kb/[H20] = [R6H21/[ROHl[H~O] = KL can be computed from the slopes and intercepts of fig. 5. KL is 2.2 for isopropanol, 2.0 for cyclohexanol and 3.2 for diisopropyl ether. The linearity of the plots in fig. 5 confirms that in all cases k:<kl. There is the possibilitfthat the equilibrium, 4- HSO, + ROH+SOi-+ ROH,,228 EFFECT OF ACID AND ALKALI ON HYDROGEN TRANSFER exists under the conditions of these experiments.To test this values of SIT were determined for isopropanol in acidities derived entirely from potassium bisulphate, with [H+]T measured spectrophotometrically using picric acid as described above. Fig. 2 shows that the plots of (-d[02]/dt)-l against l/[isopropanol] are linear and intercept the ordinate at the same point as obtained in the absence of KHS04. Fig. 5 shows that the values of S/Tlie within experimental error on the line obtained using sulphuric acid. Therefore, the equilibrium above involving HSO, ions has no effect on the values of Ki. TEMPERATURE VARIATION OF & The variation with temperature of this retardation in acidic solutions has been examined for isopropanol with added sulphuric acid.In fig. 7 a plot of (-d[02]/dt)-l against l/[isopropanol] in 0.9 N H2SO4 at 25°C is compared with a similar plot at 0°C : the plot at 25°C has the same intercept on the ordinate as the plot at 0°C. It can be easily shown from eqn. (h), that for the same acidity at 0" and 25", and that for isopropanol under these conditions, KL25 S2 - 0.08 -- - KLo So-0.08 0 C, i i j i i i i i i i o i i 2 1 j i ~ l/[isopropanol] (moles-1 1.) FIG. 7.-Temperature variation of (- d[02]/dt)-l against l/[isopropanol] in 0.9 N HzSO4. [Sensitizer] = 3 x 10-4 M : El, 0°C; 0, 255°C. From fig. 7, KL 25 = 0-76 KLO, and as KLO = 2.2, KL 25 = 1.7. From these values of KL at 0°C and 25"C, AG and A H can easily be calculated for 4. H+ + (CH3)2CHOH+(CH3)2CHOH2 in water, from AG = -RT In KL, AS = -d(AG)/dt, and AH = AG+TAS.C.F . WELLS 229 Thus - AG := 429 cal/mole at 0°C and 315 cal/mole at 25.5"C; from which, - AS == 4-48 cal/mole and - AH = 1.6 kcal/mole. DETERMINATION OF THE DISSOCIATION CONSTANT FOR ANTHRAQUINONE-2-SULPHOMC ACID The solution of anthraquinone-2-sulphonic acid prepared as indicated above was concentrated by slow evaporation under reduced pressure until the saturation point was nearly obtained : the concentration was estimated by potentiometric titration with a 10-2 N NaOH solution. From the measurement of the pH of a suitable concentration of the acid at 0°C an approximate value for the dissociation constant was determined. From a knowledge of the approximate value of the constant several solutions were prepared with the ionic strength adjusted ap- proximately to 0.1 by addition of a 1 M sodium nitrate solution.The pHs of these solutions were measured at 0°C using a Pye direct-reading instrument which was standardized at 0°C. All operations of preparation, standarization and pH measurement involving the sensitizer acid were all carried out in the dark. Al- though the range of concentrations which can be used for these determinations is narrow at O"C, due to the restriction of solubility at one end and full dissociation at the other, reasonable agreement was obtained for the three solutions used for K, = [sulphonate ion][H+]/[undiss. acid]. An average value of K, = 0.08, giving a value for the basicity constant of the sulphonate ion KL = 12.5. It is unlikely that there is much variation between unexcited and excited sen- sitizer in this respect and it is seen that KL for the sulphonate ion is quite different from the values of KL determined for the substrates from the kinetics.l/[isopropanol] (moles-1 1.) FIG. 8.-Plots of (-d[OJdt)-l against l/[isopropanol] in alkaline solutions at 0°C. [Sen- sitizer] = 3 x 10-4 M : X neutral ; A, 0.005 N NaOH ; Ely 0.01 N NaOH ; 0, 0.02 N NaOH ; V 0.03 N NaOH. REACTION IN ALKALINE SOLUTION ISOPROPANOL Plots of (- d[O2]/dt)-l against l/[isopropanol] are given in fig. 8. All intercept the ordinate at the same point as obtained in neutral solution for isopropanol and230 EFFECT OF A C I D AND ALKALI ON IIYDKOGEN ‘I’RANSFER a wide range of other substrates,l but deviations from linearity are observed at low isopropanol concentrations, the magnitude of the deviation increasing with increasing [OH-].It is evident from this that the deviations are caused by the side reaction of the excited sensitizer discussed above. The criterion used for an accurate rate measurement is that the product yield>95 % in the first five minutes of reaction, which is regarded as the minimum period for an accurate initial linear portion of the rate curve to be obtained. This assessment was used in drawing the straight lines in fig. 8. The slopes S can be calcuIated from these linear portions in fig. 8 and a plot of SIT against [OH-IT, the total added concentration of alkali, is linear (fig. 9). The equation to this line is where A is a constant. FIG. 9.-Plots of SIT against [OH-IT for isopropanol at 0°C.[Sensitizer] = 3 x 10-4 M. The constant intercept of fig. 8 again eliminates any possibility of an acid-base effect on the excited or unexcited sensitizer as being responsible for the retardation ; and again any effect such as Q* + OH--, Q + OH- is unlikely. An acid-base equilibrium of the isopropanol seems the most likely explanation for the retardation, K , OH- + isoPr OH$ H,O + isoPrO - Q* +isoPrO-%QH+radical, with k ; t k l . From a similar treatment to that used in acid solution with K , = [isoPr 0 -1 [ H2 01 /[OH -3 [isoPr OH], and assuming f i s o p r o - fH,o/ f o H - fisoPrOH = a constant over the range of concentrations used here, then K, will be a constant under these conditions. At the stationary state,C .P. WELLS 23 1 If /c;<kl, SIT for plots of (-d[O;?]/dt)-l against l/[isopropanol] is given by and substituting for the total added alkali concentration, K,[isoPr OH]), [OH-], = [OH-] 0) -++- S k K a T kl [H,O] 1 [ H A as 1 + Ku[isoPrOH]/~20]-l, for (- d[O2]/dt)-l against l/[isoPrOH]~ to be linear. Comparing (i) with ( j ) , A = Kuko/[H20]kl and from the slope and inter- cept of fig. 9 [isoPr 0 -1 = KI, = 78. KU -1 = [OH-][isoPrOH] = K Z , = 9.0 x 10-14 [isoPr 0-1 [H '1 [isoPr OH] Therefore, KY = at O'C, where Kw = 0.1 15 x 10-14 at 0"C.S The linearity of fig. 9 shows that k;<kl. DIISOPROPYL ETHER . The rate of oxygen absorption of an aqueous solution of diisopropyl ether containing sodium anthraquinone-2-sulphonate as a photo-sensitizer decreases slightly with addition of alkali, but table 3 shows that, unlike the effect with isopropanol, the rate soon attains a constant minimum value.This must be connected with the sharp changes in product yields described above. It can be concluded, therefore, that there is no effect of alkali on the interaction of the photo-excited quinone with the ether (reaction (1)). TABLE 3 [SENSITIZER] = 3 X 10-4 M [DIISOPROPYL ETHER] = 0.12 M [OH-] added initial rate, moles/l. secx 106 neutral 1 -95 0.005 M 1 *63 0.01 M 1.23 0.02 M 1.1 1 004 M 0.95 THE BASICITY OF ALCOHOLS AND ETHERS IN AQUEOUS SOLUTION The equilibrium constants Kb for + H~0+(CH,)2CHOH+(CH,),CHOH2+H20 and for cyclohexanol and diisopropyl ether have been determined in dilute solu- tions of the alcohol or ether in water : no other values for alcohols or ethers appear to be known under these conditions.Using the hydrogen halides, values of the basicity constant have been determined at 25oC in alcohols with added water by conductiometric methods for methanol and ethanol by Goldschmidt and his co-workers 99 1% 13@) (r = [Rd.H~][H20]/~:0] = 0.235 (MeOH), = 0.06 (EtOH))232 EFFECT OF ACID AND ALKALI ON HYDROGEN TRANSFER and by Thomas and Marum 11 (r = 0.25 (MeOH), = 0.06 (EtOH)) and for ethanol by Bezman and Verhoek 12 (r = 0.05) ; by catalytic methods for methanol, ethanol and isobutanol by Goldschmidt and co-workers 9(Q 13 (r = 0.21 (MeOH)), at high [acid] = 0.15 (EtOH), at low [acid] = 0.08 (EtOH), = 0.088 (isoBuOH)) and for ethanol by Deyrup 14 (r = 0-15) ; and by indicator methods for ethanol by Braude and Stern 15 (Y = 0.077) and by Deyrup 14 (r = 0-1 8).Because of the different solvation conditions existing in dilute solutions of alcohols in water and dilute solutions of water in alcohols, such values will be expected to bear no relation to the values determined in this paper. Braude and Stern 15 have found in HzO+EtOH mixtures that the acidity function Ho first decreases on addition of EtOH to water and attains a minimum which is followed by a steep rise in HO to absolute alcohol. A similar minimum is observed in the catalytic activity of H+ ions in H20+EtOH mixtures 16 and in the conductivity of solutions of acids in H20+MeOH, H20+EtOH and H20+isoBuOH mixture~.g(~)# 109 12, 13(b)(c) The effects on adding water to ethanol are accounted for by the basicity of H20 being greater than that of EtOH in the equilibrium, + EtOH,+H,O+EtOH+H:O.+ Braude and Stern calculate their value of r = [EtOH2][H2O]/[H$O] for ethanol from HO measurements on the assumption that no EtOH2 ions exist at the minimum : Braude attributes the decrease in catalytic effect on adding EtOH to water containing acid to a dielectric effect 16 and the decrease in HO to added EtOH breaking down the quasi-tetrahedral structure of the water causing the proton affinity of the H20 molecules to increase.15 Mitchell and Wynne-Jones 17 have shown that the variation of l/Ho, which represents the basicity of the solvent, with the composition of H20+EtOH mixtures is paralleled by the variation of the excess thermodynamic functions of mixing of these solvents.In particular, they point out that the variation of the thermodynamic functions with composition of HzO+EtOH and H20+ MeOH mixtures indicates that there is a strong interaction between water and alcohol molecules at low concentrations of the alcohols, but that the variation of the maximum densities with temperature 17.18 shows that low concentrations of the alcohols have no structure-breaking effect on the water. In the method outlined in this paper, dielectric effects are ruled out by the linearity of the plots in fig. 2, 3 and 4. The values of KL depend upon the pro- tonated alcohol or ether being unattacked by the photo-excited sensitizer, which leads to the direct estimation of the concentration of the unprotonated alcohol or ether via the oxygen absorption.From Frank and Evans' theory of solvation in aqueous solution,lg where alcoholic solutes stabilize " ice-berg-like " clusters of H20 molecules, these basicity constants in dilute aqueous solution represent equilibria such as + ( H, O),H + + isoPr OH (H, 0),, +( H, O), + [ is0 Pr OH (H, O),] H -+ , where rn, n, x and y are not necessarily equal and are variable throughout the solution. Here protons are partitioned between agglomerates of water molecules some of which contain alcohol or ether molecules and these are more basic than those not containing alcohol or ether molecules. Thus in dilute aqueous solutions of isopropanol, cyclohexanol and diisopropyl ether the organic solute appears more basic than water. Similarly, the basicity constants (r-values) of Goldschmidt, etc., in methanol, ethanol and isobutanol with added water will not be represented by equilibria as simple as + ROH,+H,O+ROH+H;O, but solvation conditions will have to be taken into account in an analogous way to the above for aqueous solutions : heremethanol, ethanol and isobutanol appear less basicC.F . WELLS 23 3 than water. As Bell has pointed out,2o general conclusions 21 regarding basicities of solvent molecules cannot be drawn from a comparison of the equilibrium constants for reactions like : + RNH,+H20+RNH2+H~0 RNH, + R’OH+RNH, + R‘OH, obtained in the pure solvents,14,21 as solvation effects are completely neglected. The values of KL determined in this paper appear to be little affected by varying inductive effects of the alkyl groups, which might be expected to exert an influence on the basicities, as observed for the basicities in the pure liquids,22 and especially considering the marked influence of these effects on the oxidative reactivities.1-3 This insensitivity of the KL values in aqueous solution to inductive effects must be due to the predominating influence of solvation effects on the equilibria, as de- scribed above.Table 4 shows that a similar comparative insensitivity to inductive + + TABLE 4. KI: B-tH++BH+ B KLx 10-10 20°C 40°C 6.12 1 -45 6-46 1 -05 5.13 1.15 6.03 1-29 8.36 2.25 12.6 3.16 effects is found in the KL values of amines in dilute aqueous solution calculated from the data of Evans and Hamann 23 and Everett and Wynne-Jones.24 THE ACIDITY OF ISOPROPANOL IN AQUEOUS SOLUTION The equilibrium constant KL for OH- +(CH,),CHOW+(CH,),CHO- +H,O has been determined in dilute aqueous solutions by a method which determines directly the concentration of the uncharged alcohol, the alcoholate ion being unreactive to the photo-excited sensitizer.No other values for KL appear to exist for isopropanol in water ; but Hine and Hine 25 have found for solutions in iso- propanol, Ke = [OH-]/[H20][isoPrO-] = 1-2. Similar considerations regarding solvation as discussed above for the basicities might be expected to apply in a1 kaline solution. However, aqueous alkaline solutions of ethanol and methanol appear to behave in different ways. Thus, Caldin and Long,26 using an indicator method, have found the equilibrium constant K’ = [oH-]a~:o~/[OEt-]a~,o to vary in a regular way with composition : log K’ varies linearly with the inverse of the dielectric constant, from which by extrapolation K’ = 6.3 in water and K‘ = 0.5 in EtOH at 25°C.Unmack 27 has found for alkaline methanol+water mixtures that values of K’ at 18°C derived from conductivities pass through a maximum (K’ = 0.06 at high [MeOH] and K’ = 0-26 at low [MeOH]); catalytic constants and electromotive force measurements also show a maximum in their variation with composition.234 EFFECT OF ACID A N D ALKALI ON HYDROGEN TRANSFER VALUES OF ki AND k; Both ki and k; for k\ Q* + protonated substrate +* QH + H+ +radical Q* +isoPrO-% QH + radical are <kl for the uncharged substrate. As electron-reIeasing inductive effects have been found to increase the reactivities of C--H bonds to hydrogen tra+nsfer,l-3 it would be expected that an electron-attracting substituent such as a - OH2 group will reduce reactivity. However, a substituent -0- group might be expected to increase reactivity, but in this case the repulsive forces between the alcoholate ion and the very large -SO, group of the sensitizer must outweigh this electron releasing effect. This work forms part of a programme of fundamental research undertaken by the Council of the British Rayon Research Association.1 Wells, in course of publication. 2 Wells, Rksumis 16th Int. Congr. Pure AppZ. Chem., 1957, Division n, p. 11. 3 Wells, Nature, 1956, 177, 483. 4 Bolland and Cooper, Proc. Roy. SOC. A, 1954, 225,405. 5 Cooper and Talbot, unpublished work. 6 Vogel, Practical Organic Chemistry (Longmans, London, 1956), p. 163. 7 Wells, to be published. 8 Harned and Hamer, J. Amer. Chem. Soc., 1933,55, 2194. 9 Goldschmidt and Dahll, (a) Z. physik. Chem., 1924, 108, 121 ; (b) 2. physik. Chem., 10 Goldschmidt, (a) 2. Elektrochenz., 1914, 20, 473 ; (b) 2. physik. Chem., 1915, 89, 11 Thomas and Marum, 2. physik. Chem., 1929, 143, 191. 12 Beman and Verhoek, J. Arner. Chem. SOC., 1945, 67, 1330. 13 (a) Goldschmidt and Udby, 2. physik. Chem., 1907, 60, 728 ; (b) Goldschmidt, 2. Elektrochem., 1909, 15, 4 ; (c) Goldschmidt and Tlieusen, 2. physik. Chem., 1912, 81, 30; (d) Goldschmidt and Melbye, Z. physik. Chem., 1929, 143, 139; ( e ) Goldschmidt, Haaland, Melbye, Z. physik. Chem., 1929, 143, 278. 1925, 114, 1. 129 ; (c) 2. physik. Chem., 1916, 91, 46. 14 Deyrup, J, Amer. Chem. SOC., 1934,56, 60. 15 Braude and Stem, J. Chem. SOC., 1948, 1976. 16 Braude, J. Chem. SOC., 1944,443. 17 Mitchell and Wynne-Jones, Faraday SOC. Discussions, 1953, 15, 161. 18 McHutchison, J. Chem. SOC., 1926, 1898. 19 Frank and Evans, J. Chem. Physics, 1945, 13, 507 ; Frank and Wen, Faraday SOC. 20 Bell, Acids and Bases (Methuen, London, 1952), p. 28. 21 Hammett, PhysicaZ Organic Chemistry (McGraw-Hill, New York, 1940), p. 260. 22Braude, J. Chem. Soc., 1948, 1971. 23 Evans and Hamann, Trans. Faraday Soc., 1951,47, 34. 24 Everett and Wynne-Jones, Proc. Roy. SOC. A , 1941, 177,499. 25 Hine andlHine, J. Amer. Chem. SOC., 1952, 74, 5266. 26 Caldin and Long, J. Chem. SOC., 1954, 3737. 27 Unmack, 2. physik. Chem., 1927,129,349; 1928,131.371 ; 1928,133, 45, Discussions, 1957, 24, 133.

ISSN:0366-9033    

DOI:10.1039/DF9602900219

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

OXIDATION AND REDUCTION BY ATOMIC HYDROGEN IN AQUEOUS SOLUTIONS BY GABRIEL STEIN Dept. of Physical Chemistry, Hebrew University, Jerusalem, Israel Received, 4th February, 1960 The oxidizing action of hydrogen atoms (produced by an electrodeless high-frequency discharge) on aqueous solutions of iodide and ferrous ions is examined. It is shown that tho assumption that Hias ions may be the actual oxidizing agents agrees well with the results. Comparison with results in solutions containing ferro-ferri sulphate and ferro-ferri cyanide permit an evaluation of competitive oxidation-reduction reactions due to H atoms. The results are critically compared with evidence from radiation, photo- and electro-chemistry. Atoms and free radicals may, in general, act as either oxidizing or reducing agents, according to whether they acquire an additional electron from a donor or lose their odd electron to an acceptor.For the simplest case, that of a hydrogen atom, this fact is well known. They may act as reducing agents by electron transfer or H atom addition. They may also act as dehydrogenating, i.e. oxidizing agents on molecules from which H may be abstracted whilst the molecular H-H bond is gained. As to electron transfer to H, the electron affinity of H in the gas phase is only of the order of 0.5 eV. In aqueous solution the reaction H-+aq+H,+OH,< (1) would have to occur in addition to provide the maximum gain in energy. This sequence of reactions may be given a pH-dependent form. In acid solutions another possibility arises. The formation of the hydrogen molecule ion H+H+-+H,+ (2) is exoergic by 61 kcal.In aqueous solution the hydration energy of H+ is ap- proximately 260 kcal. If, then, the hydration energy of HZ is around 200 kcal, the species HZas could exist in acid solutions. It would be a strong electron acceptor. The ionization potential of H2 being 359 kcal, the process would be exoergic by some 150-160 kcal. Coulson 1 examined the problem theoretically. The reducing action of hydrogen atoms (produced by electric discharge, and introduced into the aqueous solution) was investigated in great detail. The possible oxidizing action of atomic hydrogen was not considered seriously until recently. Ethier and Haber2 were the first to report briefly experiments which could be best explained by assuming that hydrogen atoms act as oxidizing agents in homogeneous aqueous solutions.The effect of increasing acidity was observed and a termolecular reaction, M,+,+H+H,+,-+M,~,+ +H, was postulated. 235 (4)236 OXIDATION AND REDUCTION B Y ATOMIC HYDROGEN THE EVIDENCE FROM RADIATION CHEMISTRY There was no further work in this field, until the development of radiation chemistry of aqueous solutions. The action of ionizing radiations, e.g. y-rays, on dilute aqueous solutions is best understood according to Weiss,3 if one assumes that the main radical intermediates formed are hydrogen atoms and hydroxyl radicals : HOH--s\w+H+ OH. ( 5 ) In the radiation chemistry of aqueous solutions of Fe2+ ions, in the absence of oxygen, the experimental results could best be explained4 by assuming that H atoms can oxidize ferrous to ferric ions, in acid solution, according to Fe:: +H:aq+Fe,3,$ +H,.The quantitative results of Rigg, Stein and Weiss 5 were confirmed by Barr and King6 and the work was further extended by Rothschild and Allen.7 Im- portant additional evidence was recently obtained by Shubin and Dolin.8 They carried out reaction ( 5 ) in ferrous ion solution in the presence of high pressures of H2. OH radicals are replaced by H atoms according to OH + H2 -+ H20 + H, (7) but the oxidizing equivalents remained practically constant in acid solution. Thus, evidence from radiation chemistry conclusively shows that hydrogen atoms can oxidize ferrous ions to ferric in acid aqueous solution. Regarding the nature of the actual oxidizing species, radiation chemistry could not provide conclusive evidence for or against Weiss’ assumption that Hzaq is the active intermediate.In addition to Haber’s three-body-collision mechanism (4), Uri 9 suggested the possibility of the reaction Fe2+HOH+H--+Fe3+OH- + H,. (8) In radiation chemistry the possibility of the thermalized electrons, eaq, and H atoms, or alternatively H atoms and HZfaq ions being the pair of species dominant at different pH values, was pointed out by Barr and Allen.10 Isotopic experiments using deuterium gave results,ll which did not absolutely contradict the possibility of HZaq being the actual species, but could be reconciled with it only by making special ad hoc assumptions.12 Similarly, the investigation of competition reaction between, e.g.Fez+ and 0 2 , for the reacting species 79 13 was difficult to reconcile with other evidence.12 Thus radiation chemistry could onIy suggest, but not prove, the possible mechanism of oxidation by H atoms in aqueous solution. THE EVIDENCE FROM PHOTOCHEMISTRY The detailed mechanism of the photochemistry of aqueous solutions of ions, e.g. Fez+ or I-, which on u.-v. irradiations yield molecular hydrogen, is still quite obscure. Indeed, even the basic question whether the true initial yields are at all pH-dependent or not, has not yet been conclusively decided. Rigg and Weiss investigated the photochemistry of ferrous 18 and iodide 19 ions. They found both reactions to be pH-dependent and proposed mechanisms, in which the pH- dependence was attributed to the role of Hiaq in the mechanism.However, Lefort and Douzou20 claimed that the ferrous ion reaction at least is not pH- dependent at all. We have therefore recently re-investigated the photochemical oxidation of ferrous 21 and iodide 22 ions, with particular attention to the determination of the pH-dependence of the initial yield. We found that in both cases the initial yield indeed depends on the pH. However, it was proved, as envisaged by Farkas andG . STEIN 237 Farkas 23 and by Platzman and Franck,24 that the primary photochemical formation of the H atoms was itself largely responsible for the pH-dependence observed. Our photochemical results are consistent with comparable results from radiation chemistry, and yield for comparable reactions practically identical rate constants.They prove that oxidation by H atoms occurs. However, whilst the results are entirely consistent with the assumption that H atoms act as oxidants in the actual form of HZ ions, they cannot serve as conclusive proof in this respect. THE EVIDENCE FROM ELECTROCHEMISTRY In investigations of electrode processes resulting in the evolution of H2, it was suggested by Kobosew and Nekrassow,lo that the HZ ion may be the decisive intermediate on high overvoltage electrodes. Horiuti and his co-workers 15 developed in great detail the convincing evidence in favour of this reaction step. Poltorak 16 showed that this active intermediate may penetrate the aqueous solu- tion. Recently, Ives 17 confirmed these results and summed up the electrochemical evidence in favour of the formation of H2f ions on the electrode surface and probably evaporating from there, in the body of the solution. REACTIONS OF ATOMIC HYDROGEN To avoid the ambiguities inherent in photo and radiation chemistry, it is de- sirable that the question whether H atoms can act as oxidizing agents should be investigated using H atoms produced as such and introduced into the solution.Indeed, simultaneously two groups 2 5 ~ 2 6 reported results in agreement on the point. Davis, Gordon and Hart26 produced H atoms by dissociating H2 gas through an internal discharge or by a heated tungsten wire. Czapski and Stein 25 used an electrodeless ring discharge. Both groups reported that H atoms indeed oxidize ferrous ions in acid solution. Their accuracy did not permit the deter- mination of pH-dependence.Using increased accuracy, it was shown recently 12 that the oxidation of ferrous ions by H atoms is pH-dependent. This excludes mechanism (8) as the onZy one operating. However, with ferrous ions the practic- able p H range is limited, The results could not differentiate between the pH- dependent mechanisms (4) or (6), occurring by themselves, or in addition to another mechanism, e.g. (8). The situation is much more favourable with I- ions. Here, the oxidation of I- by H atoms was proved.27 It could be shown that varying the pH from 0.4 to 7.0 and the iodide concentration between 0.006 and 0.6 M, the yield Y of iodine was given by 2A (9) where a = k$/2k, (k2, being the velocity constant of the formation of Hiaq in aqueous solution, and k,, the velocity constant of the recombination reaction, H+H-+H2) ; p = [I-]/([k2,/krJ+ D-]), where kI- is the velocity constant for the oxidation reaction of by H2-aq ; and A g atoms 1.-1 sec-1 is the rate of intro- duction of H atoms into the solution.It was shown that the results disagree with both the pH-independent mechanism and the three-body-collision process (4). They were found to be in excellent agreement with the mechanism involving the specific formation of HZaq which acts as the actual oxidizing species. THE ENERGETICS OF THE H i a q ION From the fact that H;aq can oxidize Iaq in solution an upper limit for SH+aq the energy of hydration of Hiaq can be obtained for the process238 OXIDATION A N D REDUCTION BY ATOMIC HYDROGEN Using the value of (E,+S,-) == 147 kcal, and EH+ = 365 kcal, an upper limit of approximately SHi-= 21 8 kcal is obtained, negIecting energies of hydration of the neutral entities.In our work on the photochemistry of I- soIutions,Zz resuIts were obtained which indicated that the nature of the reducing species changes at about pH = 2. Interpreting this as indicating that the pK of Hzaq is between say 1 and 3, and as- suming the energy of hydration of H+ to be 240 kcal, the energy of hydration of Hiaq has a value of approximately 204 kcal. We thus come to the tentative conclusion that SH; is probably 210 f10 kcal. 2 THE COMPETITION BETWEEN OXIDATION AND REDUCTION PROCESSES BY H ATOMS From the work on the oxidation of I- ions by H i , the velocity constant for the formation of Hiaq from H and H& could be calculated.27 k2 was found to have the value of -102 1.mole-1 sec-1. Thus the formation of HZaq is a relatively slow process. Indeed, it was found that as the pH increases, the reduction of, e.g., ferric to ferrous predominated, since H atoms were captured by the oxidizing agent. The results also indicated that the velocity of reduction by H atoms of the hexahydro Feiz ion was less than, e.g., of the Fe3+0H- complex.12 The role of the complexing group in determining the reduction velocity has been investigated in radiation chemistry22 for the groups OH-, SO:-, C1-, (C1-)2, PO:+, F- and (CN-)6 with Fe3+. The results indicated that in the reduction process, which may in different cases proceed by either group transfer, according to or by conducted electron transfer through the bridging complex, according to the nature of the complexing group is decisive for the reduction velocity. It is of great interest to investigate the detailed mechanism of such processes using H atoms.We investigated first, the reduction of ferricyanide, which prob- ably proceeds by conducted electron transfer. It was shown that when Fe(CN)i- was used,29 the velocity constant of the reaction, H + Fe(CN)i - +Fe(CN)i- + H + , was of the order of 106 1. mole-1 sec-1. The formation of Hi,, was too SIOW ta compete with this process above pH-1, and reduction predominated. Using the large protein containing molecule of ferricytochrome-c, at about pH-7, similarly high velocity constants of the order of 106 1. mole-1 sec-1 were obtained.30 Thus in any particular system, the relation between oxidation and reduction appears to be determined by the pH, oxidation processes by H atoms being favoured in acid slution.The quantitative analysis of the results excludes a purely pH- independent process, such as (8) as the sole reaction path, and excludes the ter- molecular mechanism (4), which cannot account for the quantitative pH-de- pendence. The results are compatible with the assumption of H i being the actual oxidizing species, but some uncertainties remain. It is not known whether a process like (1) may not on detailed analysis yield similar results. The use of H atoms can also yield important results towards our understanding of the detailed mechanism of reduction in ionic solutions, particularly the role of the complexing groups in determining the actual reaction path and the specific velocity constants.Added in Proof.-Since the writing of this paper we have fully investigated the implications of the hydride mechanism suggested in the paper. With Dr. G. Czapski and Dr. J. Jortner we found that it is possible to account for all theG . STElN 239 observed phenomena if we assume that two alternatye pathways of oxidation by H atoms exist. One involves the reaction of H atoms with Hiq to form HZaq. The velocity constant of this reaction, KH=H+ =; 102 1. mole-1 sec-1. H&q formed in this relatively show reaction oxidizes acceptors, c.g., I- as discussed in the paper. This slower reaction is the one also operating in pure aqueous solutions, its slow- ness accounting for the absence of isotope exchange as observed by Friedman and Zeltman. If, however, other acceptors, e.g., metal cations, are present, this may directly form intermediate complexes with H in a reeaction which may have faster rate constants.For example, k g e z + + ~ = 6x 104 1. mole-1 sec-1. When this faster hydride complex pathway is available the slower H i mechanism may not be of kinetic importance. The existence of this faster pathway explains the results of competition reactions between ferrous ions and other acceptors, e.g., 0 2 or alcohols. We could show that the specific expressions derived from the two mechanisms yield different quantitative dependence on the concentration of H+ and the acceptor. We could show that in the case of the iodide ion indeed the HZ mechanism is the only one that fits the results, whilst the hydride mechanism does not agree with it.This latter does, however, account very satisfactorily for the quantitative results in the case of the oxidation of ferrous ions by H atoms. 1 Coulson, J. Chem. SOC., 1956,778. 2 Ethier and Haber, Naturwiss., 1930, 18, 266. 3 Weiss, Nature, 1944, 153, 748. 4 Weiss, Nature, 1950, 165, 728. 5 Rigg, Stein and Weiss, Proc. Roy. SOC. A, 1952, 211, 379. 6 Barr and King, J. Amer. Chem. Soc., 1954, 76, 5565. 7 Rothschild and Allen, Rad. Research, 1958, 8, 101. 8 Shubin and Dolin, Doklady Akad. Nauk S.S.S.R., 1959,125, 1298. 9 Uri, Chem. Rev., 1952, 50, 376. 10 Barr and Allen, J. Physic. Chem., 1959, 63, 928. 11 (a) Gordon and Hart, J. Amer. Chem. Soc., 1955, 77, 3981. (b) Friedman and 12 Czapski and Stein, J. Physic. Chem., 1959, 63, 850. 13 Baxendale and Hughes, 2. physik. Chem., 1958,14, 323. 14Kobosew and Nekrassow, 2. Elektrochern., 1930, 36, 529. Kobosew, Zhur. Fiz. Khfm., 1952,26, 112. 15 Horiuti and Okamoto, Sci. Papers Inst. Physic. Chem. Res. Tokyo, 1936, 28, 321. Horiuti, Keii and Hirota, J. Res. Inst. Cataly8is5 Hokkaido, 1951, 2, 1. Horiuti, Z.physik. Chem., 1958, 15, 162. 16 Poltorak, Zhur. Fiz. Khim., 1953,27, 599. 17 Ives, Can. J. Chem., 1959, 37,231. 18 Rigg and Weiss, J. Chem. Physics, 1952, 20, 1194. 19 Rigg and Weiss, J. Chem. SOC., 1952,4198. 20 Lefort and DOUZOU, J. Chim. Physique, 1956, 53, 536. 21 Jortner and Stein, to be published. 22 Jortner, Levine and Stein, to be published, 23 Farkas and Farkas, Trans. Faraday Soc., 1938,34,1113. 24 Platzrnan and Franck, Farkas Memorial Volume (Jerusalem, 1952), p. 21. 25 Czapski and Stein, Nature, 1958, 182, 598. 26Davis, Gordon and Hart, J. Amer. Chem. SOL, 1958,80,4487. 27 Czapski, Jortner and Stein, J. Physic. Chem., 1959, 63, 1769. 28 Schwartz, J. Amer. Chern. Suc., 1957, 79, 534. Schwartz and Nritz, J. Amer. Chem, 29 Czapski and Stein, J, Physic. Chem., 1460, in press. 30 Czapski, Frohwirth and Stein, to be published. Zeltmann, J. Chem. Physics, 1958, 28, 878. Soc., 11933, 80, 5636.

ISSN:0366-9033    

DOI:10.1039/DF9602900235

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

THE PROPERTIES OF SOLUTIONS OF COPPER SALTS IN PYRIDm AND QUINOLINE BY M. PARRIS AND R. J. P. WILLIAMS Inorganic Chemistry Dept., South Parks Road, Oxford Received 5th January, 1960 Cuprous salts in pyridine and quinoline catalyze homogeneously the reduction of cupric ions by molecular hydrogen. The rate of reduction depends upon the cupric ion concentration, the anion present, and substitution in the pyridine or quinoline molecule. An overall kinetic scheme is proposed from which the rate of reaction under particular circumstances can be deduced. The redox potentials of the cuprous cupric couples and the absorption spectra of the complexes have been examined in an attempt to understand the reactions. In 1938, Calvin 1 examined the reduction of cupric ions by molecular hydrogen in quinoline and observed catalysis by cuprous ions.He found the rate equation d[H2]/dt = k[Cd]. Later 293 the equation was modified in the specific case of the reduction of cupric acetate catalyzed by cuprous acetate and the second-order rate equation was given d[H2]/dt = k[Cu1]2, after due allowance had been made for inhibition by acetic acid. The nature of the inhibition was not made clear. The earlier first-order dependence upon cuprous ion in the reduction of, for for example, cupric salicylaldehyde was not restudied. In contrast to the results in quinoline, the rate of the reduction of cupric acetate in pyridine 4 has been shown to be first order in cuprous concentration and in this solvent no correction for the acetic-acid concentration appears to have been required.Various other anions and solvent support the cuprous catalysis but none of these systems has been examined in much detail. In an attempt to understand the different rate expres- sions found under different conditions, we have made a detailed study of the catalysis of homogeneous hydrogenation in these solvents. EXPERIMENTAL RATE EXPERIMENTS The rate of hydrogenation or hydrogen absorption was measured in a conventional Warburg apparatus. The cupric salt solution in the Warburg vessel was first degassed at room temperature for some hours. Little reduction ( ( 5 %) took place at this temper- ature. A total volume of 5 ml of cupric solution was used and the dead space of hydrogen gas in the vessel and connected manometer was about 25ml. The hydrogen pressure was 1 atm.The Warburg vessel was then transferred to a thermostat at 100°C and measurement of hydrogen absorption followed after allowing 10 min for temperature equilibrium to be established. The cupric solution was stirred very vigorously with a magnetic stirrer and it was shown that the rate of hydrogen absorption was independent of stirring speed. Blank experiments were performed to ensure that the hydrogen-uptake measurements were not affected by the sudden change in temperature of the liquid. No significant effect was ever measured in the blanks. At the end of an experiment the solu- tion was examined for precipitation of copper. None was observed. With cupric salts of o-hydroxy-acetophenone, phenylacetic acid and cyanoacetic acid. entirely anomalous results, e.g. very rapid reduction, were observed.We do not understand these reactions. 240M. PARRIS AND R. J . P. WILLIAMS 241 MATERIALS Anhydrous cupric acetate was prepared by the procedure of Spath, and other cupric salts by the general methods of Pfeiffer,g i.e. normally by crystallization from alcoholic solutions of the acid and cupric acetate. The salts were stored over P205 to obtain minimum hydration. Analyses were made of all the compounds to ascertain their water content. The results showed that the cupric benzoate, the three nitro-benzoates, and the two bromo-benzoates contained two molecules of water while all other salts were anhydrous. A.R. pyridine was distilled, dried with KOH for a fortnight, and then fractionated at atmospheric pressure.A.R. quinoline was treated with benzoyl chloride and NaOH for several days, then washed successively with aqueous NaOH and water, After distillation and drying over KOH for a fortnight, it was fractionated, first at atmospheric pressure, and then at 20 mm, the portion distilling at 102-103°C being collected. Cyanopyridines were recrystallized from ligroin, other solvents being fractionated by distillation at reduced pressure. Oxidation-reduction potentials were measured in a specially designed cell in the ab- sence of oxygen. The apparatus was in all essentials similar to that described earlier.10 The reference electrode was a silver electrode made according to Gupta’s 7 procedure. The ascorbic acid was added through a micrometer syringe and the solution was vigorously stirred throughout the titration.The potentials were measured on a Cambridge pH meter. Spectrophotometric measurements were made with a Unicam SP 600 spectrophoto- meter, for which we wish to thank the Royal Society, and with a Beckmann DU spectro- photometer which was lent to us by Dr. H. Irving. time, min FIG. 1.-Rate of absorption of molecular hydrogen by solutions of cupric acetate in (a) pyridine, (b) quinoline. The curves are labelled with the total number of moles of cupric salt ( x lO+4) dissolved in 5 ml of solvent. RESULTS AND DISCUSSION Fig. 1 illustrates the general pattern of the uptake of hydrogen by a cupric salt in pyridine and quinoline. The rates we observe with cupric acetate in quinol- ine are slightly greater than those observed by Calvin and Wilmarth3 under the same conditions.The experimental methods are rather different and the general agreement is satisfactory. In pyridine, the rates of hydrogen absorption by cupric acetate are almost identical with those found by Wright, Weller and Mills.4 However, we have worked for the most part at much lower cuprous and cupric ion concentrations than the above authors and observe with no correction for acetic acid the rather different rate expression242 COPPBR SALTS IN PYRlDINE AND QUINOLINE d[HJ k,[Cu"][Cu'] -- - dr k2 4- k,[Cu"] At high cupric concentrations the equation agrees with that found in pyridine.4 In an attempt to understand the differences between ourselves and other authors for the rate expression in quinoline, we have also followed the rate of hydrogenation of quinone as catalyzed by cuprous ions.Unfortunately, there is a difficulty in the interpretation of the rates of this reaction not recognized in previous work. The cupric acetate/cuprous acetate redox potential in quinoline is so close to that of quinone/hydroquinone that complete conversion of cupric ions to cuprous ions cannot be brought about by adding hydroquinone. The equilibrium constant €or the reaction HQS. CU~~-+CU~+ Q has been determined by spectrophotometric and potentiometric methods as 4-0fO.5 at 25°C. It is probably somewhat larger at higher temperatures. As such an equilibrium exists, the addition of successive quantities of HQ to cupric solutions will not give an increase linearly related to the amount of cuprous ion in solution.Similarly, at constant total hydroquinone additha of cupric ions will not give a linear increase in the cuprous ion concentra- tion. As in our experience the rate of hydrogenation depends upon both the con- centration of cupric and cuprous ions, it will be appreciated that the rate of hydrogenation is related in a complex manner to the amounts of hydroquinone in the solutions. However, assuming that the equilibrium constant obtained at 25°C applies at lOO"C, we can calculate the cuprous and cupric ion concentrations at any total hydroquinone and total copper and compare the observed rates of hydrogenation and the calculated ones. The results are shown in table 1. The agreement is reasonable. If the assumption had been made that all the hydro- quinone was oxidized by the cupric ions there would have been an apparent rough relationship between rate of hydrogenation and [Cur]2 ignoring any cupric de- pendence or any dependence on acetic acid concentration. TABLE RA RATE OF HYDROGEN ABSORPTION IN PRESENCE OF HYDROQUINONE hydroquinone added rate (mmlmin) moles x 104/5 ml 0.2 0.4 0.8 1.2 1.6 2.0 (a) calc.0.7 1.0 1.7 2.0 2.1 {Oh. 0.4 0.8 1.5 1.9 2-0 2.4 0.9 1.5 3.0 4.0 4.9 5.5 5.8 1.0 1.5 3.4 4.5 5.4 5.8 6.0 (b) (a) total copper 1 X 10-4 mole/5 d ; (b) total copper 2 X 10-4 mok/5 ml. Note: 1*0mm/min = 0.83 x 10-7 moles hydrogen/min. We have not yet undertaken a study of the inhibition by acetic acid of hydrogen uptake in quinoline. The inhibition in pyridine is shown in fig. 2, The rate of hydrogen uptake was examined where there was little or no cupric dependence and it was found that although acetic acid inhibits the reaction the rate is well expressed by a fist-order dependence on cuprous ion at all acetic acid concentra- tions studied and at the extrapolated condition of zero acgtic acid concentration. We extrapolated linearly as we have not been able to determine a satisfactory relationship between rate and acetic acid concentration.At high acetic acid concentration there is little further inhibition on adding more acetic acid and the linear extrapolation was only employed as it appeared to fit the experimental results at lower concentrations quite well, The inhibition appears to be different in quinoline.3 The rate of hydrogen absorption by other cupric salts in pyridine has a sirnilat dependence on cuprous and cupric ion concentration, With cupric salicylaldehydeM.PARRIS AND R. J . P, WILLIAMS 2 3 the kinetics are very similar to those of the acetate. In the preseii& of excess salicylaldehyde there is somewhat more inhibition than with equimolar quantities 0 2 0 40 6 0 8 0 too yo Cu reduced FIG. 2.-Rate of hydrogen absorption by cupric acetate in pyridine (2 x 10-4 mole ib 5 ml) plotted against concentration of cuprous salt in the presence of different numbers of moles of added acetic acid. The labelling of the lines refers to number of moles of added acid x lo+% The dotted line is that obtained at the theoretical zero acid concentration. Some experimental points {circles) are given to illustrate the agreement with first-order kinetics.of acetic acid but the rate of reaction reihains Arst order in cuprous ion cori- centration. The rates of reduction of some other cupric salts are listed in table 2 togethet with the dependence of the rates on total mpric concentration initially present. The dependence on cupric concentration is found over different con- centration ranges for the different salts, The rates iri table 2, coli~mn (a), canoat be directly compared. A comparison between these results and those of Calvin and Wilmarth3 shows that the rates and the anion dependence of the rates in pyridine are very like those in quinoline. It is difficult to understand how this similarity can be explained by the differences in mechanism which have been proposed 3* 4 in the two solvents. Inhibition of the hydrogenation can also be brought about by the addition of phenanthrolines.Virtually complete inhibition is achieved if the ratio of phen- anthroline to copper exceeds 2 : 1. At lower concentrations the inhibition is not great until the ratio exceeds 1 : 1. This suggests that complexes of the formula Cul(phenan)l and CulI(phenan)l are reactive but that complexes in which two phenanthroline molecules are bound to the cations are inactive. With the cuprous salts the binding of the second phenanthroline molecule probably displaces the anion from the co-ordination sphere of the cation. Change of solvent also Affects the kinetics of the reaction. This was shown by Wright, Weller and Mills,4 who studied briefly the effect of a large number of solvents. We haw observed that th? reducthi of cupric acetate by hydrogen does not occur in 4-methyl quinoline; it occurs in 2-methyl pyridine at the same rate as in pyridine itself but in 3- or 4-cyano pyridine the rate is faster than in244 COPPER SALTS IN PYRIDINE AND QUINOLINE TABLE 2.-D"LATIVE RATES OF HYDROGEN ABSORPTION BY DIFFERENT CUPROUS SALTS M PYRIDME anion chloride acetate stearate sebaca te chloracetate benzoate o-NO2 97 m-02 Y9 P-NO2 99 o-NH2 99 m-NH2 99 P-NH2 99 3 -methyl Y9 5-methyl 9 9 5-bromo 99 5-chloro 9 9 3-nitro 9 ) 5-nitro 9 9 3,5-dinitro 99 2-hydroxy-1 -naphthaldehyde salicylaldeh yde PK, < 0.0 4.76 484 4.92 2.86 420 2.1 8 3.49 3.42 4.98 4.79 492 1035 11-50 10-55 9.20 9.35 6.70 6.55 2.65 9.30 - E" 40 350 360 390 170 307 305 190 180 305 345 385 3 30 375 400 205 240 100 65 15 250 (4 770 0 655 4.7 650 4.3 660 4.0 667 0 0 650 0 6 660 0.5 660 0.35 660 0.35 660 0-6 647 1.2 655 2.2 725 4 7 720 4 5 750 7.2 685 0 0 0 0 685 0.0 675 0.0 695 0.0 680 0.0 (6) 0 4.7 4.3 4.0 0.0 1.1 1.0 0.55 0.65 1.4 1.8 3.8 4.7 10.0 0.0 00 0-0 0.0 0 0 0.0 pK, values are acid dissociation constants in water except for the salicylaldehydes which are values in 75 % aqueous dioxan.E" values are quoted with reference to the silver electrode described in the experimental section.7 They refer to the cuprouslcupric couple in pyridine. The position of the absorption maximum is that observed for the cupric salt in pyridine. Rate (a) is that found in the presence of 05x 10-4 mole of cuprous and 0.5 x 10-4 mole of cupric salt in 5 ml pyridine.Rate (b) is similar to rate (a) but in the presence of three times the quantity of cupric salt. Rate units are as in table 1. pyridine and is second order in cuprous salt, if no account is taken of inhibition by the acetic acid produced (fig. 3). The following scheme can be used to account for all the above observations : k i k2 cu' 3. H,=CuXE12, Cu'H, + Cu"ku'H + Cu'+ HI, Cu'H, + Cu1%2Cu'H, Cu'H+ Cu"%2Cu'+H1. Assuming steady-state concentrations in CdH and Cu1H2, k5 [ Cu 'H] [CU I 'I = k4[ CU 'HJ [ CU '1 + k3 [ CU 'HJ [ CU "J , kl[C~'][H2] = [Cu1H2](k, + k,[Cu"] +kq[CUg}-M. PARRIS AND R. J. P . WILLIAMS 0 245 6.0 I2 ' j c C 0 0 . 5 1.0 Pyrid inc Pyridinc moles of cuprous x 104/5 ml Fia. 3.-Second-order kinetics of hydrogen absorption in cyano-pyridines ; rate in same units as fig.2. Combining these equations, The following cases are of interest Eqn. (3) also holds for ortho/para hydrogen conversion in the absence of cupric ions if the back reaction k-4 is assumed rate determining : Conditions (6) and (c) might be expected to obtain either at high cuprous con- centration or where k4 is large compared with k3. Condition (b) will hold rather than (c) depending only on the relative size of k2 and k4[Cu1]. This could be different in different solvents and we must suppose that k2 is relatively large in cyano-pyridines and small in pyridine in order to explain the kinetics. The differences between our kinetics and those of Calvin and Wilmarth in quinoline may depend upon the fact that they worked at much higher cuprous concentra- tions where condition (6) might well hold.Condition (a) is valid throughout the lower ranges of copper concentration in quinoline and pyridine but there is a transition to condition (d) at the highest concentrations. Eqn. (1) expresses this246 COPPER SALTS IN PYRIDINE AND QUINOLINE transition. The transition is not independent of anion and the ratio k3/k2 appears to decrease with the change from simple aliphatic acid salts to salicylaldehyde salts to benzoic acid salts. Although within each of the three groups there is a de- pendence of rate on the basicity of the anion this cannot be associated generally with an effect upon kl, for the rate which we observe for the benzoates is related to klk3/k2. In this case, it is not possible to say that the dependence of rate on basicity is any guide to the mechanism of the initial step in the hydrogen uptake. The reaction scheme we propose is essentially the same, independent of solvent and anion.The initial step, kl, is of primary interest. In order to understand this step further, we have made a detailed study into the nature of the complexes in the solution by two methods, the determination of the oxidation reduction potentials of the cuprous/cupric couples in pyridine and the examination of the absorption spectra of the complexes. The oxidation reduction potentials were measured 7 (see experimental section) under the same conditions as the hydrogenation reactions were studied except for a change of temperature from 100°C in the latter to 25°C in the former.Mid- point potentials were obtained by following the titration of the cupric salt with ascorbic acid. The potentials obtained indicated that both species, cuprous and cupric, were polymerized to the same degree. From the absorption spectra it seems highly improbable that there is any appreciable polymerization of the cupric species so that polymerization of the cuprous complexes can probably be ruled out. The mid-point potentials refer to the potential in the presence of half an equivalent of the acid, HX, derived from the cupric salt CuX2, per equivalent of the cupric and cuprous ions. As we have shown that the potentials are sensitive to excess acid to a degree dependent on the acid, the mid-point potentials are not exactly comparable.However, the changes in potential due to additions of acid are much smaller than the changes due to changes in the acid (see table 2 and fig. 4). > E r;" I 2 5 0 * I 3 0 0 , 3 5 0 0 ACID 3 7 5 f 10 moles of acid/20 ml x 104 FIG. 4.-The effect of additions of acetic acid (lower curve) and salicylaldehyde (upper curve) on the redox potentials. We shall discuss the dependence of the potential on the anion and the acid in more detail elsewhere. Here, it is sufficient to note that the rate of reduction of a cupric salt by hydrogen does follow the oxidation reduction potential quite closely. The oxidation reduction potential is related within each of the three groups of anions to the dissociation constant of the acid. Ligands which increase the potential toM. I'ARRIS AND R .J . P. WILLIAMS 247 more than about - 450 mV with respect to the reference eleotrode,7 e.g., 8-hydroxy- quinoline and bis-salicylaldehyde ethylenedi-imine, do not form catalytically active hydrogenation catalysts either in pyridine or in quinoline.3 It may be that the potential of these couples is so high, thereby favouring high cupric concentrations, that it is not thermodynamically possible for the reduction of cupric to cuprous to take place by the action of molecular hydrogen to any marked degree. The situation would appear very similar to that in water where it has been observed that hydrogen will not reduce cupric salts in the presence of ethylenediamine and ethylenediamine tetra-acetic acid.6 Both these chelating agents lower the Cul/CuII potential far below the Hz/HI potential.5 On the other hand, the effect of phen- anthroline on the copper acetate potential is very slight and its inhibitory effect must be ascribed to a totally different condition.The absorption spectra of the complexes indicates that this is the displacement of the acetate ion from the co- ordination sphere of the cations. Excess acetic acid and excess salicylaldehyde raise, i.e. make more positive, the oxidation reduction potentials in the same way, The degree to which these reagents lower the rate of reduction of the cupric ion by hydrogen roughly parallels the expected fall in rate calculated from a plot of rate of reduction against oxidation reduction potential. We shall discuss the absorption spectra of the complexes in detail elsewhere. Our interpretation is consistent with the formulation of the cuprous complexes as CuIXpy3 and with that of the cupric complexes as C~~~X2py4.We have come across no evidence for dimers of either species in any solvent or with any of the anions we have studied. INTERPRETATION OF RESULTS In those cases where the rate equation takes the form (4) or (S), we have no grounds for assuming that the initial step is different from the reaction between a monomeric cuprous-anion species and a hydrogen molecule. As there is a rough dependence of this rate on the basicity of the anion bound to the cuprous ion, and as the catalysis is prevented by reagents, phenanthrolines, which remove the anion from the neighbourhood of the cuprous ion, we suppose the initial step is not very different from heterolytic fission.11 The hydrogenated complex is then H- H+ The H- is associated more with the co-ordinated pyridine and the Hf with the co-ordinated anion, X. Where there is a clear-cut demonstration by us of a de- pendence of rate on [Cu1]2 (cyano-pyridines as solvent), we imagine that because it is easier for the cuprous ion than the cupric ion to bind to the cyan0 group the magnitude of k4[Cu1] exceeds k3[CUI1]. In other cases under our experimental conditions, k3[CuI1] appears to be greater than k&ul]. One of us (M. P.) wishes to express his thanks to The British Rubber Producers Research Association for a maintenance grant. I Calvin, Trans. Faraday Sue., 1938,34, 1181. 2 Calvin, J. Amer. Chem. Soc., 1939, 61, 2230. 3 Calvin and Wilmarth, J. Amer. Chem. SOC., 1956,78, 1301. 4 Wright, Weller and Mills, J. Physic. Chem., 1955, 59, 1060. 5 James and Williams, unpublished observations. 6 Halpern, Quart. Rev., 1956, 10, 463. 7Gupta, J. Chem. SOC., 1952, 3473. 8 Spath, Monatsh., 1911, 33, 235. 9 Pfeiffer, Breith, Liibbe and Tsumaki, Annalen, 1933,503, 84. 10 Tomkinson and Williams, J. Chem. SOC., 1958, 1153. 11 Halpern, J. Physic. Chem., 1959, 63, 598.

ISSN:0366-9033    

DOI:10.1039/DF9602900240

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

GENERAL DISCUSSION Dr. C. F. Wells (British Rayon Res. Assoc.) (partly communicated): Whilst I appreciate Prof. Halpern's point about the composition of the activated state, it is important to identify the species taking part in the reactions. It is often diffi- cult in the reactions of Co(II1) ions with organic substrates to decide whether the the Hf-dependence is on the Co(II1) or on the substrate, particularly in the re- action of cobaltic sulphate with formic acid1 which he quotes, and also in the reactions of cobaltic sulphate with alcohols 2 and with formaldehyde 2 9 3 and in the reaction of cobaltic perchlorate with formaldehyde.3 However, in our study 4 of the reaction of cobaltic perchlorate with benzene, it is clear that the H+-de- pendence is on the Co(II1) and that both Co(II1) and Co(1II)OH- oxidize benzene with [Co(III)OH-]< [Co(III)].5 This suggests, for instance, that the H+- dependence in the cobaltic perchlorate 3- formaldehyde reaction in a similar range of [HClO4] might be due, at least in part, to Co(1II)OH- as well as Co(1II) re- acting with formaldehyde : indeed, when the published data 3 are replotted they conform quite well to the relation k = kl+(k2/[H+]).It seems possible in the reactions of cobaltic sulphate mentioned above that the H+-dependence can involve two Co(1II) species, A and B, presumably sulphate-containing complexes, related by A+B+H+, reacting with the substrate. However, the difference in the kinetic behaviour of the reactions of cobaltic sulphate with crotonic acid and with the olefines 6 suggests that in the reaction of Co(II1) sulphate with formic acid the H+-dependence is largely on the formic acid.I wonder if there is any evidence that the hydrolytic species of the cations used by Halpern and Taylor in solutions of perchloric acid are capable of oxidizing organic substrates ? Dr. L. H. Sutcliffe (Liverpool University) said: The fact that Bawn and White 7 used reaction mixtures containing sulphuric acid in the concentration range 1.3 to 5 M would lead one to expect sulphate complexes of Co(II1) to take over from hydrolyzed species.8 The oxidation of Ce(II1) by Co(II1) in perchloric media has been studied in the presence of a variety of simple anions 99 10 including sulphate.11 When the anions are present at moderate concentrations then CoOH2f appears to be main reactive form of Co(II1).However, concentrations of bisulphate ions less than about 10-2 M introduce a competitive step involving sulphate complexes of Co(III), therefore it seems fair to assume that the experimental conditions adopted by Bawn and White would exclude reactions in which CoOH2+ participates. Dr. W. A. Waters (Oxford University) said: Prof. Halpern favours electron- pair transfer processes for reactions of formic acid. His list can be extended by noting that even the powerful 1-electron transfer reagents Mn3+ and VO', scarcely oxidize formic acid. In contrast, oxalic acid is much more easily oxidized by I-electron transfer than by 2-electron transfer (pure acid permanganate does not * Barn and White, J . Chem. SOC., 1951, 339.2 Bawn and White, J. Chem. Sac., 1951, 347. 3 Hargreaves and Sutcliffe, Trans. Faruduy SOC., 1955, 51, 786. 4 Baxendale and Wells, to be published ; for summary see Furaday SOC. Di.wmion.7, 5 Baxendale and Wells, Truns. Furuday Soc., 1957, 53, 800. 6 Bawn and Sharp, J. Chem. SOC., 1957, 1854. 7Bawn and White, J. Chem. SOC., 1951, 339. 8 Hargreavts and Sutcliffe, Trans. Furuday SOC., 1955, 51, 786. 9 Sutcliffe and Weber, Trans. Furuduy SOC., 1956, 52, 1225. 10 Sutcliffe and Weber, Truns. Furaday SOC., 1959,55, 1892. 11 Sutcliffe and Weber, to be published. 1953, 14, 239. 248GENERAL DISCUSSION 249 oxidize it) : thus the differing reactions of formic and oxalic acid can be used for diagnosis of oxidation mechanisms. When electron-pair oxidation by oxy-anions, e.g.MnOi, MnOi‘, HCr0; takes place the mechanism involves oxygen atom transfer. The steric release of internal compression in the closely packed, tetrahedral, MO4 groups by conversion to more planar MO3 structures may be of considerable energetic importance in favouring these oxidations, even though MO3 ions promptly break down to still simpler structures such as Mn02 or CrOz, the formation of the latter having been demonstrated in oxidations effected by acid chromate. Electron transfer with concurrent change of ionic structure has not yet been considered by the theorists present at this discussion, but the application of this type of conformational analysis to inorganic reactions may be very important. In reply to Dr. Symons, though simple electron transfer reactions of ions MO4 to structurally similar products, e.g., MnOi‘, CrOi- has been conclusively demonstrated, this is not sufficient evidence for refusal to concede the transient existence of MO3 ions, as found in sulphites or selenites.The reduced MO4 ions form only under highly alkaline conditions, or in glasses, and these conditions are not relevant for the majority of oxidations of organic compounds. Mr. R. P. Bell (Oxford University) (communicatcd) : The dissociation constants reported in Halpern and Taylor’s paper for HCOOH and DCOOH ( 5 . 6 ~ 10-4 and 8 . 0 ~ 10-4) call for some comment. A secondary isotope effect of 1.4 is unusually high, since although the isotopic substitution is close to the dissociating group, only one hydrogen is substituted. Since the molecules involved are simple, it should be possible to calculate the anticipated isotope effect from the vibrational frequencies of the species HCOOH and HCOO-.Although the assignment of the observed frequencies is not altogether certain,l calculations made in con- junction with Dr. R. E. Weston (Brookhaven National Laboratory) indicated that it was not possible to account for an isotope effect of more than a few per cent on the basis of any reasonable assignment. These calculations refer to a hypoth- etical gas-phase reaction, and the large isotope effect observed would imply that the modes of motion of at least one of the species HCOOH and HCOO- are greatly modified by the presence of the solvent. Dr. N. S. Hush (Bristul University) (comnzunicated): It is not easy to estimate the solvation energies of large hydrocarbon ions, and it is correspondingly difficult to try to distinguish between different possible detailed reaction mechanisms for hydrocarbon/hydrocarbon- exchange reactions.However, both solution and electrode kinetic data obtained by Hoijtink and others suggest that the free energy of activation is not, in general, very dependent on molecular size, so that solvation effects are probably of less importance here than in metal ion exchange reactions. It is worth noting that an abiabatic mechanism will require a slight deformation of the carbon skeleton in the transition state to accommodate the electron which is being transferred into an antibonding orbital. This should not vary very much from one molecule to another and may account for the greater part of the small observed activation energies.Prof. R. A. Marcus (Polytechnic Inst. of Brooklyn) said : I should like to com- ment on Weiss’ and Laidler’s interesting treatments of electron-transfer reactions, referred to in the paper by Aten et al. Weiss’ treatment provided a helpful first examination of various factors influencing electron-transfer rates, while Laidler examined the role of dielectric saturation in affecting coulombic repulsion b e tween the reactants. If the distance between the hydrated cations in the activated complex is the sum of the hydrated radii, then Laidler’s calculations indicate that dielectric saturation is negligible. (To be sure, it must be admitted, this fortunate concIusion is based on a long theoretical extrapolation of Malsch’s data on saturation.) 1 cf.Miyazawa and Pitzer, J. Chem. Physics, 1959, 30, 1076.250 CENEKAL DISCUSSION There is one point, however, which is referred to in the present paper and to which I would like to take exception here. Both Laidler and Weiss do not con- sider any re-organization of the solvation atmospheres about the reactants prior to electron transfer. Since the equilibrium solvation atmosphere of any species as a reactant and as a product is very different, because the charge on the species is different, and since a re-organization of solvation atmospheres during electron transfer was not considered, the authors presumably assume, tacitly or otherwise, that it occurs after the electron transfer.Indeed, this inference was drawn by Hoijtink and his collaborators. Let us examine the consequences of this assumption. (i) Referring to fig. 1 of my paper, we see that if no re-organization of solvation atmospheres occurred prior to the transfer, the electron transfer would have to involve a vertical motion on that graph, from the potential-energy curve of the reactants to that of the products. However, such a process can only occur as a result of the absorption of a light quantum. Since the authors' mechanism is supposed to be a thermal one, it is clear that there is a contradiction and that the original assumption is incorrect. (ii) This contradiction reveals itself in another way as well-an interesting violation of the principle of microscopic reversi bility. Consider, for example, an electron transfer between Fe+2 and Fe+3 and suppose that solvation re- organization occurred after but not before electron transfer. Let us denote by brackets, ions which are temporarily present in foreign solvation environments : >.Fe'2+Fe+3+[Fe+3] + [Fe+2] (electron transfer only) [Fe+3] + [Fe+2]-+Fe'3 +FeS2 (solvation reorganization) By the principle of microscopic reversibility, the reverse process would have to occur at the same net rate. That is, lB Fe+3+Fe+2+[Fef3]+[Fe+2] [Fe+ '1 + [Fe+2] -+Fe+2 + Fe+3 Thus, it follows from microscopic reversibility that there also occurs, and with equal rate, a process in which solvation re-organization precedes electron transfer. Furthermore, the chance that the reaction sequence A has exactly the same rate as B, for each possible microscopic process, is quite negligible. So we conclude that some compromise situation prevails-partial re-organization of solvation atmosphere both before and after electron transfer.Such a path which satisfies the principle of microscopic reversibility is given in fig. 1 of my paper. Dr. D. R. Stranks (University of Leeds) (communicated) : Aten, Dieleman and Hoijtink 1 suggest that for electron transfer between an aromatic molecule and its ion, the rate constant k would be given by k = ke exp (dS*/R), where ke is the " encounter rate constant ". They have placed considerable stress on the finding by Ward and Weissman 2 that electron transfer between naphthalene and naphthalene negative ion proceeds with " zero activation enthalpy ".How- ever, k, itself must be diffusion-controlled and would normally give rise to an apparent activation enthalpy of from 3 to 5 kcal mole-1. Ward and Weissman, on the other hand, estimate that the upper limit for the activation energy for naphthalene-naphthalene negative ion in tetrahydrofuran is 2.6 kcal mole-1. (No temperature study was reported for dimethoxymethane.) In connection with this apparent anomaly, two possibilities which might be considered are: (a) that diffusion of naphthalene and its ion in tetrahydrofuran proceeds with an unusually 1 Aten, Dielemann and Hoijtink, this Discussion. 2 Ward and Weissman, J. Amer. Chem. SOC., 1957,79, 2086.GENERAL DISCUSSION 25 1 low enthalpy of activation, and (b) that the eiithalpy of ion-association between cations and the naphthalene ion (mentioned by Ward and Weissman) partly cancels the true activation enthalpy for electron transfer.Since it is difficult to visualise how a reaction in solution cannot be limited ultimately by diffusion, further investigation of these systems would seem to be desirable. In the analogous case of the ferrocene-ferricinium electron exchange, the evidence suggests an activation enthalpy exceeding 4 kcal mole-1.1 Prof. R. A. Marcus (Polytechnic Inst. of Brooklyn) (communicated): It was interesting to see that both homogeneous and electrochemical electron transfers are being studied by using the same compound. Since quantitative studies of fast electrochemical electron transfers are relatively new, and since electrochemistry and homogeneous chemical kinetics have conventionally constituted rather different research areas, there has been only a little correlation attempted between electron- transfer mechanisms in the two types of systems.Some correlation would be expected when salt effects are appropriately considered.2 A greater interchange of information would seem desirable. Certainly electrochemists could benefit from a knowledge of the careful mechanistic studies made in isotopic exchanges, while chemical kineticists might be able ultimately to draw inferences about electron as contrast atom transfers based on such comparisons. Prof. F. S. Dainton (Leeds University) said: One of the aims of our work was to determine the Arrhenius parameters of the reactions of a wide range of aquated and complexed cations with a common uncharged reactant, thereby minimizing the Coulombic interactions between the reactants and perhaps enabling better tests of existing theories. If the uncharged reactant is to be capable of oxidation and reduction it must be capable of accepting or losing either electrons or electron- carriers such as H, OH, Cl, etc.An organic free radical has just these properties but since preliminary experiments indicated that the bimolecular rate constants of their reaction with transition metal ions may be very large, perhaps as high as 107 1. mole-1 sec-1, the radical concentration must be kept very low, say, in the range 10-5 to 10-9 M. One of the easiest and most sensitive ways of measuring low radical concentrations is to make use of the fact that they add vinyl monomers very readily at a rate Rp = kp[ml][R.], where kp is of the order 104 1.mole-1 sec-1. Moreover, it is easy to measure rates of polymerization of the order of 10-4 M secl, especially if a recording dilatometer is used. Our paper is primarily concerned with the determination of (i) the absolute rate constants of polyacrylamide radicals with metal ions and (ii) the relative rate constants of reactions of metal ions with, on the one hand, hydrogen atoms and on the other, hydroxyl radicals. The striking difference between these two groups is discussed in the paper. Since the paper was in the proof stage the follow- ing new information has been obtained. First, infra-red examination of the polymers has shown that the polymer produced by oxidation of the radical contains a terminal vinylene group and therefore that the reaction involves hydrogen-atom abstraction as depicted in eqn.(23). This result is consistent with the absence of a solvent isotope effect in these reactions (see fig. 2) which indicates that 0 - H bond Assion is not involved. Secondly, when ferric ion is the oxidant the presence of 0.1 M chloride concentration increases k’, 30-fold, so that it would appear that chloride ion has an analogous catalytic effect to that which it has on ferrous-ferric exchange. A test will shortly be made to see whether this catalysis also involves chlorine atom addition to the radical. Thirdly, the reaction with titanous ion does not involve hydrogen formation and therefore eqn. (25a) and (256) do not describe the course of this reaction. Dr.W. A. Waters (Oxford University) said: Prof. Dainton has distinguished between the two groups of metallic cations which can easily, or slowly, reduce 1 Stranks, this Discussion. 2e.g. Marcus, Can. J. Chem., 1959, 37, 155. cf. Randles, Trans. Faraday Soc., 1952, 48, 828. Vlcek, COIL Czech. Chem. Comm., 1955,20, 894.252 GENERAL DISCUSSION organic radicals, pointing out that reference to AG is not adequate. An alter- native distinction is that only the metals of the latter (slowly reducing) group form stable covalent organometallic compounds. Dr. J. Halpern (University of British Columbia) said: The higher rates of reduction of Ag+ and Hg2+ by H atoms, found by Dainton et al., compared to those for Fe3+, Cu2-t- and CeOH3+, probably reflect a difference in mechanism. The last three ions are efficient 1-electron oxidants and it is likely that their re- duction proceeds by an electron transfer mechanism of the type Fe3-'-X-+H-t[Fe3-t .. , X- . . . H]->Fe2+XHY (1) foIIowed by either Fe2+XH+ H2O+Fe2+-X-+ HgOf, ( 2 4 or Fe2+XH+ H20+Fe2+0H2+ HX (or H++ X-). (2b) Step (1) is effectively the transfer of an electron from H to the metal ion through a bridging ligand and is thus formally analogous to the bridged electron-transfer reactions between metal ions (e.g., the oxidation of Cr2+ by (H3N)&o3-t-X-) described by Taube. Whether this step is followed by (2a) or (2b), resulting in an overall reaction which corresponds formally to " electron transfer " or " ligand abstraction ',, respectively,l will depend on the substitution lability of Fe2+X.Thus (2a) is likely to be favoured for Fe(CN)a- and (2b) for FeC12f. With FeOH2+, step (1) alone completes the reduction process. This mechanism ac- counts adequately for the higher rates of reduction of FeOH2+, FeCP-:, etc., compared to Fei:. An actual ligand abstraction mechanism involving a co- ordinatively unsaturated intermediate seems rather unlikely. For As+, homogeneous reduction by electron transfer (i.e., to an Ag atom) is much less favourable and it seems likely that this reaction proceeds by a different mechanism in which the first step is the formation of AgH+.2 This and similar hydrides (AgH, CuHf, etc.) are known to be formed as intermediates in the reduction of metal ions by molecular hydrogen in aqueous solution2 The observed 4 activation energy of 15 kcal/mole for the reaction 2 A&+ H2-22 AgH+ implies considerable stability for AgH+ (with respect to dissociation into either AS++ H or Ag+ H+) and suggests both that the reaction A&+ H-kAgH' will be very rapid and that it will occur in preference to electron transfer.The rcduction of Hg2+ may proceed through a similar mechanism. Prof. F. S. Dainton (Leeds University) said: In reply to Dr. Halpern, I would not deny the possibility that the reduction of certain metal cations by hydrogen atoms might involve the formation of a metal hydride cation, e.g., AgHf, as the rate-determining step. Even in this case, however, d-+s promotion energies might exert a major influence. We shall be in a better position to comment when a larger number of ions have been ranked in order of reactivity towards He.Someone has raised the question as to whether we are sure that hydrogen atoms are the actual reducing species, to which the answer must be negative. In all our solutions the pH was less than 2 and, as may be seen from the figures shown in my comment on Dr. Stein's paper, under these conditions hydrogen atoms are very rapidly converted to H+2 ions in the reversible reaction. H+ H++Hl. 1 Stein, this Discussion. 2 Halpern, Czapski, Jortner and Stein, Nature, in press. 3 Halpern, Advances in Cutalvsis, 1957, 9, 302; 1959, 11, 301. 4 Webster and Halpern, J. Physic. Chem., 1957, 61, 1239.GENERAL DISCUSSION 253 It could well be that H+, is the actual reductant and that eqn.(14) should be written Hi + W++2 H++ M(x-l)f and that the order of reactivity is determined by the electrostatic repulsion of H; for MX+ which will increase as x increases or z=+ decreases. Further experi- ments are needed before a decision on this point is possible. Dr. M. C. R. Symons (Southampton University) said: With respect to the decomposition of permanganate in aqueous alkali to give manganate and oxygen, some recent experiments on the addition of hydrogen peroxide are relevant. When manganate is in large excess, added peroxide reduces Zess than an equivalent of permanganate. Also, if water containing excess H$80 is used the evolved oxygen contains considerably more than the natural amount of 180. These results are rationalized if the mechanism originally postulated for this decomposition is accepted : MnO;+ OH-+MnOi-+ *OH (1) Mn0;+ 0- + OH- +MnOi- + HOi (3) MnO;+ HO;+MnO:- + HO2 (4) HO,+ OH-+02+H20 ( 5 ) MnO, + 0, ->MnOz - + 02.(6) The major feature of this scheme which renders it improbable is the adverse free energy change for step (1).1 We would particularly welcome comments on this aspect of the problem. Dr. W. A. Waters (Oxford University) said: We shall shortly be publishing from the Dyson Perrins laboratory evidence which indirectly supports Dr. Stewart’s view that the permanganate oxidation of PhCH(OH)CF3 must involve electron- pair transfer. Some years ago, Mosher discovered the occurrence of C-C bond fission in chromic acid oxidations of related alcohols, ArCH(OH)CMe3, and Westheimer has provided evidence that this fission is not due to Crvl.As an alternative the 1-electron fission, ArCH(O-)CMe3z%ArCH=O+ CMe3, which gives two fairly stable products, has been suggested for the C-C rupture process2 and Mr. J. R. Jones has now found that with the highly reactive (V(OH)3}2+ ion, which reduces only to VIV, such fission occurs to about 80 % in the direct oxidation of PhCH(OH).CMe3. In view of the stability of the CF3 radical, one could anticipate that C-C fission would occur extensively in l-electron oxidation of Dr. Stewart’s alcohol. Mr. R. P. Bell (Oxford University) said: The very large hydrogen isotope effects reported for the reactions studied by Stewart and Linden are of great interest, and the suggested explanation in terms of the tunnel effect seems the most likely one.It receives some support from the difference observed between the entropies of activation (or pre-exponential factors) for the two isotopes. In terms of the latter we have approximatelyAD/AH = 3 for the oxidation of C~HSCHOHCF~. A treatment not invoking the tunnel effect would predict values of AD/AH much closer to unity, and certainly not exceeding 24. Similar results for proton transfer reactions have been accounted for quantitatively in terms of tunnelling 3 and the 1 Symons, J. Chem. Soc., 1953, 3956. 2 Waters, Quart. Rev., 1958, 12, 286. 3 Bell, Fendley and Hulett, Proc. Roy. SOC. A , 1956, 235, 453. Bell, Trans. Faraday Soc., 1959, 55, 1.254 GENERAL DISCUSSION same treatment is applicable to reactions involving the transfer of a hydrogen atom or hydride ion.Mr. R. P. Bell (Oxford University) said: If the pK of isopropanol in water is as low as 13 (cp. Well's paper) it should be possible to measure it directly by the method described by Ballinger and Long,l depending upon its effect on the conductivity of sodium hydroxide solutions, and thus to obtain an independent check of the value derived from the kinetic measurements. Dr. C. F. Wells (British Rayon Res. Assoc.) said: In reply to Mr. Bell, it is indeed possible to determine the acidic dissociation constants of alcohols in water by conductivity measurements. Since submitting my paper for this Discussion, a papcr by Ballinger and Long 2 has appeared in which the acidity constants of some simple alcohols measured at 25°C by use of a conductivity method are given.Unfortunately, isopropanol is not one of the alcohols used, and the difference between the constants for ethanol and methanol at 25°C determined by con- ductivity and that for isopropanol at 0°C in my paper is probably largely due to the differences in the solvation properties of the solutions between 0°C and 25"C, especially considering the changes in structure of aqueous solutions which occur at about 0°C. I am indeed interested in the statement by Prof. Wynne-Jones about their results on the equilibrium : isoPrOH+ OH-zisoPrO-+ H20 over the whole range of alcohol-water composition. However, in the case discussed in my paper, the isopropanol concentration does not exceed 15 % or 2 M (see fig. 8), i.e.the discussion is restricted to dilute solutions of isopropanol in water. In the kinetic analysis of the results it is assumed that ~ & , R O - ~ H ~ O ~ ~ ~ H & P ~ O H = a constant (say F&), over the range of concentrations used, where F'H may be -1 for such an equilibrium involving no total change in ionic charge, as assumed by Ballinger and Long in their conductivity work discussed above. The term [isoPrO-][H2O]/[OH-] [isoPrOH] is equated to the equilibrium con- stant Ka and this concentration product is used throughout as a simplification. The thermodynamic equilibrium constant KOH = (aisoPrO-aH,O)/(a~oP~HaOH->, where a = activity, and Kb = KoH/F'H. Any deviation from constancy of FOH will be seen in the slopes of fig. 8. These are, in fact, seen to be variable with l/[isopropanol].In the paper this deviation from linearity is accounted for by the side-reaction of the photo-excited sensitizer, and this effect must indeed be largely responsible for this deviation. However, in view of Prof. Wynne-Jones' results, it would appear likely that deviation of FOH from constancy may be r e sponsible for some of the curvature. Nevertheless, the linearity of the plot in fig. 9 indicates the essential correctness of the treatment in alkaline solutions and the approximate constancy of FOH over the linear partiom in fig. 8. In reply to Prof. Stewart, the accuracy of the determination of K , for anthra- quinone-Zsulphonic acid is, of course, restricted by the narrow range of con* centrations over which measurement was possible.However, the value is suf- ficiently accurate to distinguish the KA value for the sulphonate anion from the values determined from the kinetic measurements. There is one point in the paper which is perhaps not clear. In the temperature variation of KL, the relationship used requires kl/ko to be invariant with tem- perature. This is irideed the case and the experimental evidence for this will be presented elsewhere. 1 J. Amer. Chem. Soc., 1959, 81, 2347. 2 Ballinger and Long, J. Amer. Chem. Soc., 1960, 82, 795.GENERAL DISCUSSION 255 Dr. Ross Stewart (University of British Columbia) said : The method used by Dr. Wells to determine the total hydrogen ion concentration, i.e., spectrophoto- metric measurements with picric acid, actually measures h- rather than [ H f l (assuming that the pK of picric acid is tied into the dilute aqueous system).l The acidity function required to measure the basicity of isopropyl alcohol is ho.This may account for the very much higher basicity reported here than that previously determined, PKBH+ = -3.2,2 although the discrepancy seems rather to be com- pletely accounted for in this way. The acidity of anthraquinone-2-sulphonic acid reported here is certainly lower than that expected for a substituted benzene sulphonic acid 3 and the method used to determine it is of questionable validity for an acid of this strength. Dr. C. F. Wells (British Rayon Res. Assoc.) (communicated) : Concerning the point raised by Prof. Ross Stewart about the nature of [H+] derived from the measurements with picric acid, under these conditions this method gives total molar concentration of H-f added in absence of organic substrate by reference to the calibration over the same [€-I+]-range against molar concentration of Hi- from HC1 added to the picric acid.The pK of picric acid is eliminated in the calibration. The difference between the valuc of Kbt for isopropanol derived in my paper and that derived by Bartlett and McCollum 1 is probably due to the difference in media, as described abovc. Dr, M. C. R. Symons (Southampton University) said: I would like to mention some experimental evidence which supports the suggestion of Dr. Wells-that, other things being equal, one might expect RO- to be more reactive than ROH towards radical attack on a-hydrogen. Electron spin resonan& studies of the radicals eH20H and dH20- trapped in rigid media show that the unpaired electron is more delocalized on the oxygen in dH20- than eH2OH.Thus, if there is an appreciable amount of bond breaking in the transition state, the anion should react more rapidly than the alcohol. Dr. C. F. Wells (Bristol Rayon Res. Assoc.) (communicated): For the other type of equilibrium discussed in my paper, ROH+ H $ O ~ R & H ~ + H~O, the concentration product [Ri;H21[HzOl/[ROHI[HfOl = Kb is also used for simplicity throughout the kinetic treatment, and as stated, the activity coefficient term is assumed to be a constant (say FH) over the range of concentrations used. FH may be -1. The thermodynamic equilibrium constant KH = ~ R ~ H ~ ~ H z O / ~ R O H ~ H ~ + O , where a = activity, and K’ = KHIF!, Therefore, any variation in FH over the conditions that I used would be reflected in the slopes of fig.2, 3 and 4 : these plots are all good straight lines indicating that indeed FH is a constant under these conditions. I have now had an opportunity of considering the paper on Hydride Transfer Mechanisms in Strongly Acid Media by Bartlett and McCollum4 which Prof. Stewart mentioned to me privately. From their work they deduce KL = 6x 10-4 1 Paul and Long, Chem. Rev., 1957, 1. 2 Bartlett and McCollum, J. Amer. Chem. Soc., 1956, 78, 1441. 3 Hammett, Physical Organic Chemistry (McGraw-Hill, 1940), p. 261. 4 Rartlett and McCollum, J. Amer. Chem. Soc,, 1956, 78, 1441.256 GENERAL DISCUSSION for the basicity of isopropanol. Assuming that their proposed mechanism for the reaction they studied is correct, I can only presume that the difference between their value of KL and that deduced in my paper is due to the difference in media in which they were determined. Thus, Bartlett and McCollum worked in 22-75 % H2S04 at 25°C and my work with isopropanol was done at 0°C in solutions of (1.62 N H2S04 or 8 % w/v H2SO4. These are really two entirely different media : the one, H2S04+H20 mixtures, the other, aqueous solution ; the solvation properties, which appear largely to control the basicities, will be entirely different in the two media.This is borne out by the general agreement in order between my values for isopropanol, cyclohexanol and di-isopropyl ether at 0°C in aqueous solution and those of Goldschmidt and other authors quoted in my paper for methanol, ethanol and isobutanol at 25°C in alcohol-rich solutions, although, as stated in my paper, no exact agreement can be expected between water-rich and alcohol-rich solutions. As I mentioned during the discussion, the acidic dissociation of picric acid does not appear to be as simple as represented by picrate ion+ H++picric acid, at least, when the concentration of the picrate ion is determined spectrophoto- metrically.The spectrum of the picrate ion has a maximum absorption at 360 mp which obeys Beer’s law. Since the work on my paper was completed I have found that Amax. for the picrate ion is shifted by the addition of H+ ions but not by other simple ions. This is illustrated in the table.All the maxima are fairly broad and Amax. is the value at the centre in each case. TABLE 1.-5.45 x 10-5 M PLCRIC ACID AT 20°C additive [additive] L a x . (mp) &ax, - c 360 0.780 NaCl 0.95 N 360 0.775 H2S04 0.95 N 348 0.565 HCl 095 N 348 0.480 HC104 -1 N 348 0.440 However, the calibration plot of [H+] against (Do-D)/D using HCl gave a good straight line with all measurements of D at 360 mp ; but the value of Kp = 1.19 determined from the slope must be regarded as a calibration constant rather than a true value of the dissociation constant of picric acid. All measurements of [H+]r using sulphuric acid were done at 360 mp and compared with the calibration curve, so that the values of [H+]T in my paper are accurate. This shift in Amax. does not detract from the value of picric acid for determining high values of [H+].Prof. G. Stein (Hebrew University, Jerusalem) (communicated) : Since the writing of my paper we have fully investigated the implications of the hydride mechanism suggested in the paper. With Dr. G. Czapski and Dr. J. Jortner we found that it is possible to account for all the observed phenomena if we assume that two alternative pathways of oxidation by H atoms exist. One involves the reaction of H atoms with H& to form Hi,,. The velocity constant of this reaction, k(H = H+) = 102 1. mole-1 sec-1. H;aq formed in this relatively slow reaction oxidizes acceptors, e.g., I- as discussed in the paper. This slower reaction is the one also operating in pure aqueous solutions, its slowness accounting for the absence of isotope exchange as observed by Friedman and Zeltman.If, however, other acceptors, e.g metal cations, are present this may directly form intermediate complexes with H in a reaction which may have faster rate constants. For example, k(Fe2++ H) = 6 x 104 1. mole-1 sec-1. When this faster hydride complex pathway is available, the slower Hi mechanism may not be of kinetic importance. The existence of this faster pathway explains the resultsG EN F, K A I, 1) I S C U S 15 I 0 N 257 of competition reactions between ferrous ions and other acceptors, e.g. 0 2 or alcohols. We could show that the specific expressions derived from the two mech- anisms yield different quantitative dependence on the concentration of H+ and the acceptor. We could also show that for the iodide ion the H i mechanism is the only one that fits the results, whilst the hydride mechanism does not agree with it.This latter does, however, account very satisfactorily for the quantitative results in the oxidation of ferrous ions by H atoms. Prof. J. Weiss (University of Durhanz) said: Recent investigations on the photo- chemistry of ferrous ions in aqueous solution carried out by Dr. Hayon about two years ago, have also fully confirmed the pH dependence of the initial yields in agreement with the earlier work of Rigg. A detailed paper is now in the course o f publication. It is a matter of considerable interest that Stein has now been able to confirm the formation and the reactions of H,’ with hydrogen atoms intro- duced directly into a solution.With regard to his suggestion that in addition to the H i mechanism there is an alternative pathway at higher ferrous ion concentrations, it would be par- ticularly important to establish with some certainty the relative rate constants of the different reaction steps involved. Also some additional evidence for the hydride complex, not derived entirely from kinetic evidence, would be very desir- able before a final decision can be made with regard to the importance of the hydride formation. We have recently also found some indications in certain systems of some change in the mechanism at relatively high ferrous salt con- centrations. We have tentatively attributed this to a reduction of the ferrous ions to a monovalent state although we have not been able, as yet, to arrive at any definite conclusion in this matter. Prof.F. S. Dainton (Leecis University) said: During the past year, Dr. D. Peterson and Mr. Sills have obtained some rather striking results which indicate that nitrous oxide is an extremely good reagent for discriminating between the ionized and unionized forms of (i) hydrogen atoms and (ii) hydroxyl radicals. PH 10 FIG. 1.-G(H2), G(02) and G(N2) as a function of pH for 7-irradiated solutions of N20 in pure water. Full details wiII be published elsewhere but the most important results are dis- played in fig. 1 and 2 and are explicable on the basis that N20 is inert to the radicals I., HOB and HO2. but reacts much more rapidly with hydrogen atoms than H i ions and much more rapidly with 0- ions than with OH, each of these reactions resulting in nitrogen formation according to the equations, I258 GENERAL DISCUSSION H+N20+N2+ OH, O-+N20--+N2+0- 2 .At pH 2.4, hydrogen ions begin to compete with NzO for hyclrogen atonis to form H; and at pH 2.0 this competition is overwhelmingly in favour of H l formation, so that in this pH range G(N2) and $(N2) change very rapidly with pH. In the FIG. 2.-4@2), 4(H2) and 4(N2) a s ? function of pH for u.-v. irradiated solutions of potassium iodide and N20 in pure water. photochemical system, 4(N2) falls to zero whereas in the radiation chemical system, G(N2) does not become zero at low pH, a result which is ascribed either to the presence of some solvated electrons or excited water molecules which can react with N20 by eiq+N20->N2+OH+OH-, H20*+N20->N2+20H.Thus there seems little doubt about the possible existence in water of the H: molecule ion with a pK, of about 2.3 and of hydroxyl radicals with a pK, of 11.3. There are other reasons for believing that at high pH hydrogen atoms may be converted to solvated electrons by reaction with hydroxide ions, i.e., H+ OH-+H20eTq. Prof. J. Weiss (Univevsify of Durham) said: In support of Prof. Dainton's remarks, I can say that we have also now quite definitely come to the conclusion that in water irradiated with ionizing radiations one has three species related to hydrogen atoms. This was already discusscd previously in some detail in a paper by Dr. Hayon and myself presented at the Geneva Conference in 1958. The situation appears to be as follows: the absorption of, c.g., y-radiation leads to ionization and excitation. In the ionization process, the electrons removed from the water appear in the first instance as self-trapped electrons in the polar medium (" polarons ").These polarons which for the sake of simplicity may be denoted by (H20)- are then capable of undergoing reactions with other solutes in these systems, In the presence of hydrogen ions, one of these reactions is the capture of the polaron to give the hydrogen atoms according to (H20)-+ H+->H+ HzO.GENERAL DISCUSSION 259 I n sufficiently acid solutions, the hydrogen atoms can then lead to the formation of H t ions in the way discussed previously. A point of interest is that, in addition to hydrogen atoms produced from polarons, some hydrogen atoms are apparently also formed more directly.There are reasons to believe that the latter arise from excited water molecules according to The G-value of this latter process which is independent of pH was found to be about G-0.6. The evidence for these directly produced hydrogen atoms has been particularly clearly established by the recent work of Scholes and Allan on the pH dependence of thc radiolysis of solutions of different alcohols. With regard to the species related to the OH radical, I may mention briefly that it appears that the first species produced again corrcsponds to a self-trapped hole (H20)-1- (positive polaron), which may then lead further to the formation of OH radicals. It seems that it is rather more difficult to distinguish between the reactions of the two species (H20)i- and OH.However, wc have now obtained some definite indications for the reactions of the (H20)’ ions by studying the stereo- specificity of somc radiation-induced hydroxylation reactions. Dr. J. K. ‘lhomas (U.K.A.E.A., Wantage) said: The evidence both for and against the cxistencc of Ht that has been derived from radiolysis studies is con- flicting. Prof. Stein has mentioned several relevant cases and I would like to draw attention to further work in this line. By the study of the X-radiolysis of dilute aqueous solutions of acetic acid+ hydrogen peroxide mixtures it is possible to find the ratio of the ratc constants of the reactions of the postulated H atom pt-oduccd, with acetic acid and hydrogen peroxide, is., I c ~ / / c ~ can be measured.HaO*+H+OH. X-radiation H20--->H+ OH, H + CH~COOH- +Er2+ CH~COOH, kl I-I+ H202-tH20 + OH. k2 I t was found that this ratio /cI//c~ was constant over a pH range from 1.2 to 5.4, and would seem to indicate that the species H+2, the formation of which is said to be pH-dependent, is not present in the radiolysis of this aqueous system in the However, the pH effect noted by Prof. Stein in the radiolysis of aqueous metal pH I - U I ~ C 1.2-5.4. ion systems can be explained in terms of the following equations : X-radiation H20w--+H20 t+ e, H++ e--zH, k 3 k 4 e t (metal ion)’*+-->(metal ion)(n--l) 1 , the ratio / c j / k q then being pH-dependent. 1 wonder whether this sort of scheme could explain the results that Prof. Dainton has found in the N@+water systcm, and whether the list of relative rates of reaction of metal ions with H atoms that he indicated in his paper, are valid in view of the uncertain nature of the H atom produced in the radiolysis of water.Dr. W. A. Waters (Oxford Uizivcrsity) said: Thc fact that the CuI1 oxidation o f molecular hydrogen requires catalysis by somc complex of Cur is not a unique example amongst reactions of Cu“. The oxidation of glucose and other simple sugars by Fchling’s solution is also catalyzed by a Cur intermediate and Mr. B. A. Marshall and 1 will shortly be publishing kinetic analyscs of the oxidations of acetoin and benzoin by alkaline cupric citrak and malate complexes. In thesc oxidations there is an autocatalytic period in which CuT accumulatcs and eventually precipitates as CuzO.If oxygen is present the Cut- is immediately re-oxidized and260 GENERAL DISCUSSION no net reduction of CulI can be observed though the acyloin slowly autoxidizes. W e ascribe the main reaction, which is of zero order with respect to [Cu*I], to the oxidation by Cu2+ of a Cur chelate complex formed by the di-enol of the acyloin. I understand that Dr. J. H. Baxendale in Manchester found, a few years ago, that the Cu" oxidation of acetylenes had similar characteristics, and required the prior formation of an autoxidizable Cul acetylide. Oxidations dircctly effected by C~i2-t~ other than direct elcctron transfers with other simple ions, scem to be very rare. I would be glad to lcarn of othcr examples of oxidations of covalent molecules by metallic ions in which thc highcr valcncy state becomes reactive onIy in the presence of the lowcr valency. Mr.D. A. Dowden (Z.C.Z. Ltd., Billingltnnt Division) said: Dr. N. Mackcrizic and Dr. M. S. Spencer, working with me on thc descriptive chcmistry of the homo- geneous activation of molecular hydrogen, have made a number of observations worthy of note. While confirming that thc catalytic activity of cuprous itcctate- quiiioline complexes is not due to colloidal metal, it was found that warm quinoline passed through freshly prepared Raney coppcr took up mctal (with decrcasing readiness as the copper aged) to give a green solution; the solution absorbctl hydrogen in an autocatalytic manner with a colour change to rcd and the appcar ance of the characteristic catalytic properties. It seemed that simple anions might not be necessary to the catalytic species. However, Raney coppcr is hardly a suitable material with which to test this hypothesis. Thereforc evaporated coppei- films were made from spectroscopically purc copper in ;I high vacuum and con- tacted with carefully purified quinoline (comparable to that used by Parris and Williams) with rigorous exclusion of air. On warming, part of thc copper film dissolved to form a red solution which catalyzed quinoiie reduction with hydrogen and turned green when exposed to dry air. Although the quinoline sample might have contained other bases it could not have contained appreciable halogen or other anions unless there was a gross error. Perhaps this result can be checked by workers still interested in this topic. The rough outlines of an activity pattern, similar to that found among hydrogcl~ reactions catalyzed by transitional metals and oxides, begins to appear. Ovcr these solids, catalytic activity increases sharply as the first long period is traversed to the left from copper. The copper complexes in solution affect only a very few simple catalyses (c.g. hydrogenation of quinone), aqueous nickel (I) cyanides will hydrogenate acetylene to ethylene although not truly catalytically and cobaltous cyanide complexcs will hydrogenate conjugated di-olefincs to olefincs, sometimes with intercsting specificity. Thus again moving to the left from coppcr brings about considerablc increase in activity as in heterogeneous catalysis. In reply to Prof. Halpern, I must make it yuitc clear that we do not attribute the activity of the copper solutions to colloidal copper. Cupric hydrides, cuprous oxide and cupric oxide gave no colour effects with quinolinc and appear to be quite insoluble. Copper hydride (from cuprous iodides with lithium aluminium hydride) dissolves to give a red-brown solution which does not hydrogenatc quinone; it oxidizes to a green solution which cannot bc reduced again with hydrogen. In reply to Dr. Williams, although solids do form in some aqueous solutions of' cobaltous cyanide complexes, we havc observcd hydrogenation of dienes in solutions which do not appear to contain solids.

ISSN:0366-9033    

DOI:10.1039/DF9602900248

出版商:RSC    

年代:1960

数据来源: RSC

摘要:

A U T HO I< INDEX * Adamson, A. MI., 120, 125, 163, 173. Atcn, A. C., 182. I)ASOlO, I:., 80, 133. Baughan, c., 135. Hell, R. ID., 92, 119, 249. 253. 254. Benson, I)., 60. Carpenter, I,. G., 92. <:his, Y., 109. Collinson, ti., 188. Csinyi, I,. ,I., 146, 171. Lhrintos, 1.'. S., 121, 125, 171, 188, 251, 252. I)iclcniao, J., 182. lhdson, It. \V., 92. 134. Ihivdcn, D. A., 260. Iliick, 1;. W.. 169. I;ord-Sniith, 11. I-I., 92, 114. (;ilks, S. \\'., 102. 757. G r a y , I'., 160. I-Iaggett, RI. l,., lS3. Hallwrn, J., 7, 32, 119, 174, 252. liigginson, \V. C. I<., 49, 122, 123, 136, 1.15. lioijtink, G . .I., 182. .Jones, l'., JS3, 173. King, 1.1. I,., 109, 130. v;in dcr Linden, H., 21 I . Marcus, K. A., 21, I IS, 119, 120. 124, 120, Morris, M. L., 80. Orgel, I,. l . , 7, 32, 1 10.I rllsil, N. s., I 13, I 16. 13.7, 240. im, K . A. K., 205. 249, 25 1 . Parris, M., 240. Pearson, H. G., 80, 133. l'oc, A. .I., 133. Proll, 1'. J., 60, 129. Ihscinsky, I). R., 49, 128. Sliimi, I . A. W., 122. Smith, I). It., J88. Stead, .1. H., 49, 126. Stcia, G . , 235, 256. Stcivart, H., 21 1, 255. Strmks, D. H., 73, I 16, 131, 250. Sutcliffe, 1.. H., 00, 128, 129. 248. Sutin. N., 134. Sykcs, A. G., 40. Symons, hl. C. It., 1'70, 172, 17.1, 205, 252, 'lbul)~, I-I., 42, 132. 255. 'I'aylor, S. hl., 174. l'azuki., S., 188. 'l'homas, J . I<., 250. 'I'rutlel, G. .J., 188. l's:lO, M., 137. Vlcek, A. A., 114. Waind, G . M., 102, 127, 134. \V;ilklcy, .I., 60. Watci's, M'. A., 170. 23s. 251, 25.3, ?SO. Wciss, J . , 118, 125, 128, 170, 257, 258. \Vclls, <:. F., 128, 171. 219, 248. 354, 255. \\'iIlianw, It. J. ID., 12 I , 132, 177, 240. Wilmarth, W. K., 137, 100. Wynne-Jones, W. 1;. K., 123, 172. %wickel, A. M., 42. 261

ISSN:0366-9033    

DOI:10.1039/DF9602900261

出版商:RSC    

年代:1960

数据来源: RSC