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Isotopic exchange in the cobaltous-cobaltic system in perchloric acid—III influence of some ions and heavy water on exchange
J. inorg, nucl. Chem., 1967, Vol. 29, pp. 1983 to 1991. Pergamon Press Ltd. Printed in Northern Ireland
ISOTOPIC E X C H A N G E IN THE COBALTOUS-COBALTIC SYSTEM IN PERCHLORIC A C I D - - I I I INFLUENCE OF SOME IONS A N D HEAVY WATER ON EXCHANGE J. SHANKAR a n d B. C. DE S o u z A Bhabha Atomic Research Centre, Chemistry Division, Bombay, India (First received 23 December 1966; in revised form 13 March 1967)
Almtraet--Sulphate, acetate and nitrate ion have been found to enhance exchange in the cobaltouscobaltic aquo system in perchloric acid. Fluoride ion has no appreciable effect. Changing the solvent from light to heavy water reduces exchange proceeding through acid-independent and acid-dependent paths by a factor of about 2 through a reduction in the velocity constants of the two paths. Exchange through sulphate and acetate dependent paths is also reduced considerably in heavy water, but it is not possible to decide in the light of the evidence at present available whether the reduction is due to a change in the velocity constants for the exchange processes or to a change in the values of the formation constants of the complexes through which the exchange proceeds or to a change in both. The addition of small amounts of silver ion considerably increases the exchange in the system. The processes responsible for this are briefly discussed. PRONOUNCED effects due to the presence o f ions have b e e n f o u n d in a n u m b e r o f r e d o x reactions. I s o t o p i c exchange in t h e f e r r o u s - f e r r i c system for instance, is influenced b y chloride, cl) fluoride, (2) b r o m i d e , (3) sulphate, (4) azide, (5) t h i o c y a n a t e ~6) a n d o x a l a t e ions33) Evidence o b t a i n e d f r o m studies o f these effects has h e l p e d c o n s i d e r a b l y in the f o r m u l a t i o n o f theories on the m e c h a n i s m s o f the r e a c t i o n s involved. F o r a similar reason, studies o f the effect o f h e a v y w a t e r on exchange a n d o t h e r r e a c t i o n s have also been c a r r i e d out. This p a p e r describes a s t u d y o f similar effects in t h e c o b a l t o u s cobaltie system, w o r k which was carried o u t with a view to gaining f u r t h e r insight i n t o t h e n a t u r e o f exchange in the system, which has a l r e a d y received c o n s i d e r a b l e attention. ~7-u) S o m e o f the w o r k described here overlaps t h a t p u b l i s h e d recently b y HABIB et al. (~) EXPERIMENTAL Preliminary experiments showed that some anions, like chloride and thiocyanate, were quickly oxidized by cobaltic ions. Sulphate, acetate, fluoride, nitrate and argentous ions were found suitable for study. These were added to the reaction mixtures in the form of standard solutions of sulphuric acid, acetic acid, sodium fluoride, nitric acid and silver nitrate. The experiments were carried out as before~9~; the added ions did not interfere with the procedure except for fluoride ion, the presence of which prevented spectrophotometric estimation of Co(III) ~xl j. SILVERMANand R. W. DODSON,J. phys. Chem. 56, 846 (1952). ~2)j. HUDIS and A. C. WAHL,3". Am. chem. Soc. 75, 4153 (1953). ~s~R. A. HORNE,d. phys. Chem. 64, 1512 (1960). c4) K. H. LIESER and H. SCHROEI)ER,d. inorg, nucl. Chem. 14, 98 (1960); R. L. S. WILLIX, Trans. Faraday Soc. 59, 1315 (1963). (51 D. BONN, F. S. DAINTONand S. DUCKWORTH, Trans. Faraday Soc. 55, 109 (1960). ~6~G. S. LAWRENCE,Trans. Faraday Soc. 53, 1326 (1957). t~) N. A. BONNEt*.and J. P. HUNT, J. Am. chem. Soc. 74, 1866 (1952). ts~ N. A. BONNERand J. P. HUNT, J. Am. Chem. Soe. 82, 3826 (1960). c9) j. SHANKARand B. C. DE SOUZA,J. inorg nucL Chem. 24, 187 (1962). ix0) j. SHANKARand B. C. DE SOUZA,d. inorg, nucl. Chem. 24, 694 (1962). ~lx~B. C. DE SOUZA,Ph.D. Thesis, Univ. of Bombay (1962). ¢x2)H. S. HABIBand J. P. HUNT, J. Am. chem. Soc. 88, 1668 (1966). 1983
1984
J. SHANKAR and B. C. DE SOUZA
through the thiocyanate complex of Fe(III). In this case a blank mixture containing all the reactants in the same quantities with the exception of fluoride ion, was simultaneously prepared and used for the estimations. The concentrations of the different ions were such that: (a) their addition (i.e. replacement for the perchlorate ions (bulk electrolyte) by these ions) did not cause any activity coetficient changes, (b) the concentration of anionic complexes containing more than one anion in the co-ordination spheres was kept low, (c) the half-time of overall exchange lay generally between 3-25 min. Longer or shorter half-times led to errors in their determination--in the former case due to appreciable decomposition of the cobaltie species during the course of an experiment (six half-times) and in the other due to sampling time, (d) the acidity of the reaction mixtures was kept high so as to avoid complications arising from the presence of appreciable quantities of hydrolysed species (some of unknown composition) as well as the difficulty of estimating [H +] at low concentrations. Calculation of ion concentrations: Concentrations of argentous and nitrate ions were taken as being equal to the formalities of silver nitrate and nitric acid solutions respectively, as these compounds dissociate completely at the concentration used in the experiments. A knowledge of the corresponding dissociation quotients is necessary for the calculation of the actual concentrations of sulphate, acetate and fluoride ions. Sulphate ion concentration was calculated as follows. Complete dissociation of H, SO4 into HSO4- and H + was assumed. KERKERt18~has reported a value of 1.083 F at/~ = 2 and 18°C for the dissociation quotient (Q,). Using this and the value of heat dissociation which has been reported as --5.2 :k 0.5 kcal, c.4~ the Q2 value at 0°C and # = 2 was calculated to be 1.95. Knowing this, the sulphate ion concentration at any value of [H +] could be calculated in terms of the initial concentration, c, of the added sulphuric acid according to the relation: 2 " C
[SO?-]
-
-
-
Q, + [H +] •
BAKERet aL t*5~have shown that at 25°C and/z = 1, the dissociation quotient of sulphuric acid is 1"75 :k 0"17 times lower when the solvent is changed from light to 8970 heavy water. This is also the value to be expected on the basis of the rule of Rule and LaMer. Assuming that the ratio is not much affected by temperature and ionic strength, Q, was calculated as being 1.10 at 0°C and/~ = 2 in heavy water. Similar calculations were made for the determination of fluoride ion using the data calculated by HuDiS et aL ('~ at/~ = 0"5 based on the values of BRo~r,m et aL (16) In the study of the acetate ion effect, [OAc-] values were calculated using the relation: KD [OAc-] ~ ~-7]'" e where c is the initial concentration of acetic acid and K~) is its dissociation constant. From the reported value t~7) of 1-657 × 10-5 F for the dissociation constant of acetic acid at 0°C and/~ = 0, the corresponding value at 0°C and # = 1 at which these experiments were conducted, was found to be 5.27 x 10-5 F. The value of the constant B in the extended Debye--Huckel equation is 0.324a ° at 0°C, where a ° is the mean ionic radius in angstroms. Assuming a ° to be nearly equal to the radius of a molecule of acetic acid, the value taken for it was 2.03 A, based on the density of glacial acetic acid. Due allowance has been made for the larger value of a ° obtained from this calculation by applying a correction based on the calculated and experimental values of the radius of a water molecule, the assumption being made that the correction factor is similar in the two cases. The decrease in the dissociation constant of acetic acid in 100 70 heavy water has been reported c16~ to be by a factor of 3"3. Assuming linear change in this factor over a small range of deuterium content, tlS~ M. KERKER,J. Am. chem. Soc. 79, 3664 (1957). cx4~K. S. Pn'ZER, J. Am. chem. Soe. 59, 2365 (1937). (xs) F. B. BAKERand T. W. NEWTON,J. phys. Chem. 61, 381 (1957). ~xs~H. H. BROENEand T. DEVRIES,J. Am. chem. Soc. 69, 1644 (1947). cxT~H. S. HARr,~D and R. W. ErmER, d. Am. chem. Soc. 55, 652 (1933). tlS~ S. KORMANand V. K. LAMER, J. Am. chem. Soc. 58, 1396 (1936).
Isotopic exchange in the cobaltous--cobaltic system in perchloric acid--III
1985
its value was taken as 3.06 in 93 ~ heavy water, the concentration at which these experiments were carded out. The dissociation constant under these conditions was calculated to be 1.72 x 10-5 F. In the calculation of actual ion concentration, the effect of complexation on these concentrations of ions, has been considered as being negligible under the experimental conditions. The heavy water used in some of the experiments was 98 per cent in deuterium content. Experiments made for the study of the heavy water effect were conducted in the same way as the others but the volumes of the reaction mixtures were kept between 20 to 30 ml instead of the usual 50 ml. Wherever possible solutions of salts were prepared in and dilutions done with heavy water. The H20 in the final reaction mixtures came mainly from the 70 Yoperchloric acid, used for maintenance of the acidity of the mixtures without any pretreatment for replacement of its hydrogen by deuterium. RESULTS T a b l e s 1, 2, 3 a n d 4 s h o w t h a t s u l p h a t e , a c e t a t e a n d n i t r a t e i o n s e n h a n c e e x c h a n g e in the cobaltous-cobaltic aquo system while fluoride ion does not appear to have any n o t i c e a b l e effect. T a b l e s 1 a n d 2 a l s o s h o w t h a t c h a n g i n g t h e s o l v e n t f r o m l i g h t t o h e a v y w a t e r h a s a n a p p r e c i a b l e effect o n t h e s u l p h a t e a n d a c e t a t e i o n c a t a l y s e d e x c h a n g e . D a t a o n t h e a n i o n - i n d e p e n d e n t e x c h a n g e at 0 ° C a n d 10°C in h e a v y w a t e r a r e g i v e n i n T a b l e 5. F i g u r e 1 s h o w s t h e i n f l u e n c e o f a r g e n t o u s i o n o n t h e e x c h a n g e . TABLE 1.--EFFECT OF SULPHATE ION ON EXCHANGE. CONCENTRATION OF PERCHLORIC ACID 2"0 F . IONIC STRENGTH 2"0. TEMPERATURE OF RUNS 0 ° C
Medium H~O Concentration of total cobalt ,---0"9 × 10-s F
Medium 88 ~ D~O (v/v) Concentration of total cobalt ~1.3
× 10 -8 F
Initial concn of H2SO4 c( × 10~-F)
Velocity constant k(F -x min -~)
Initial conch of H2SO4 c( × l0 s F)
Velocity constant k(F -~ rain -1)
0"21 0"34 0"56 0"90 1"15 0"00
65 72 84 111 125 52
0"50 0"70 I'00 1"25 1"50 0'00
46 52 62 70 79 26
TABLE 2.--EFFECT OF ACETATEION ON EXCHANGE. CONCENTRATIONOF PERCHLORIC ACID 1"0 F. IONIC STRENGTH 1"0. TEMPERATUREOF RUNS 0°C Medium H20 Concentration of total cobalt ~1"0 × 10 -8 F
11
Medium 93 ~ D~O (v/v) Concentration of total cobalt ~ 2 . 0 × 10-a F
Initial concn of CH3COOH c(F)
Velocity constant k(F -1 min -1)
Initial concn of CH3COOH c (F)
Velocity constant k(F -1 min -1)
0"10 0'20 0'37 0"50 0"70 0"00
52 59 72 87 104 43
0-20 0"37 0"50 0"70 1-00 1-50 0.00
28 32 34 42 49 57 22
1986
J. SHANKAR and B. C. DE SOUZA
TABLE 3.~EFFECT OF NITRATE ION ON EXCHANGE. CONCENTRATION OF PERCHLORIC ACID 2"0 F. I o m c STRENGTH 2"0. TEMPERATUREor RUNS 0°C Concentration o f total cobalt
,---0.31 × 10-a F
Initial concn o f HNO3 e(F)
. Velocity constant k(F -x min -1)
0"15 0"20 0.30 0.40 0"50 0.00
59"0 61"5 66.0 69"5 75.5 52.0
TABLE 4 . - - E F F E C T OF FLUORIDE ION ON EXCHANGE. CONCENTRATION OF PERCHLORIC ACID 0"5 F . IONIC STRENGTH 0"5. TEMPERATURE OF RUNS 0 ° C
Concentration o f total cobalt Initial concn of NaF e(x 102 F)
,~1.4 × 10 - s F
[F-](x 104 F)
0-3 0"8 1.2 3"5 5-0 10-0 15"0 20"0 0.0
Velocity constant k(F -1 rain -I)
0-1 0"28 0"42 1.35 2-01 4.35 7"92 11.5 -
44 43 44 43 44 41 46 52 44
-
TABLE 5.---EFFECTOF HEAVYWATERON EXCHANGE. IONIC STRENGTH3"0. TEMPERATURE OF RUNS 0°C Medium H20 Concentration of total cobalt ,--4).5 × 10-s F
Medium 93 ~ D20 (v/v) Concentration of total cobalt ~ I . 0 x 10-8 F
Concentration of HC104 [H+](F)
Velocity constant k(F -1 min -1)
Concentration of DC104 [D+](F)
0"26 0"36 0"47 0"58 0"67
163 142 130 114 106
0'24 0"32 0"54 0"64 0"76
Velocity constant k(F -1 rain -1) 92 69 61 52 48
Isotopic exchangein the cobaltous-cobalticsystemin perchloric acid--lll
1987
i
-r -~ 12o b_
Io0 7
8¢
6¢
40
I
I
I
2
I 3
r 4
I
5
__
[Ag +] xlO 4 F
FIG. 1.--Effectof argentous ion concentrationon exchangein the cobaltous-cobaltic aquo system. Temperature= 0.0°C. Ionic strength =- 1.0. Concentrationsof perchloric acid, cobaltic and cobaltous perchlorates are 1.0 F, 0.5 × 10-s F and 1.2 × 10 -8 F.
DISCUSSION
Anion-dependent exchange The results obtained when sulphate, acetate and nitrate ions are present in the reaction mixtures can be expressed in the following empirical form: k ---- a + b[A"-],
(1)
where a and b are constants and [A"-] represents the actual concentration of the anion A"-. The first order dependence on ion concentration indicates the participation of one anion in the activated complex through which the anion-dependent exchange proceeds. The most probable mode of its incorporation into the activated complex is through its prior complexation with one of the exchanging ions. Evidence obtained from studies of similar anion complexes of other elements suggests that electrostatic forces, dependent more on the charges and sizes of the complexing ions than on their chemical nature, predominate in these type of complexes. These considerations favour the formation of cobaltic rather than cobaltous complexes with the anion. Again, the tendency of an ion to reduce its charge through complexation with anions will favour the formation of cobaltic complexes. Thus these species will be present in solution in appreciable quantities. The charge difference between complexed cobaltic species and uncomplexed cobaltous one will be less than that between the uncomplexed cobaltous and cobaltie species or uncomplexed cobaltic and complexed cobaltous ones. Thus the environmental and structural differences between the two exchanging species, being governed
J. SI-IANKARand B. C. DE SOUZA
1988
by this difference, will tend to favour the following exchange: ,Co2+ + COA3q. _ , CoA.q 3 - . + Coaq, 2+ the overall rate of which will be given by the equation: R = k2[Co~][Co3+] + kl[Co~][CoOH ~-] + k(~,_)[CoZ+][CoAa-n], where the first two terms are the contributions of the anion-independent mechanisms (previously discussed) and k(x,_ ) is the velocity constant for the anion-dependent exchange. The suffix aq has been, and will be, omitted for the sake of convenience. TABt~6.--VALUES OF kcx,,-}g[coAa-n)FOR THE DIFFERENT ANION-DEPENDENT PATHS OF EXCHANGE IN
HzO AND D~O
AT 0 ° C
Added anion
Ionic strength
SO42-
1
H20 8 0 ~ DzO
1"31 -4-0"15 × 104 0"99 + 0"10 × 104
CHaCOO-
1
HzO 93 ~o D . O
1"70 -4- 0"4 × l0 s 0"68 4- 0"06 × 106
NOB-
2
H~O
45"7 4- 4"8
Solvent
k~--~Kcoox3--j
It can easily be shown by combining the above equation with that obtained from kinetics, i.e. R = k[Co(III)][Co(III)], where k is the overall rate constant, that: k ---- k 0 + k(~,-)K(CoA~-.)[A~-],
(2)
where k 0 is the contribution from the anion-independent paths of exchange. This is constant ff the [H+] of the reaction mixtures is constant. K(coxs_,) is the formation constant of the complex CoAa-~. The above equation holds strictly only when the concentrations of species of the type CoA~-~ and COA32-2n are very small. It can thus be seen that constants a and b in Equation (1) are identifiable with k o and k(x,- ) K(cox,-, ). The values of the latter for the different anions in light and heavy water have been given in Table 6. Because of the instability of the cobaltic ion in aqueous solutions and its strongly oxidizing nature, the study of its properties has been mainly restricted to its use as a powerful oxidant for organic compounds. Thus, while its complexes, which are more stable, have been extensively studied, the properties of the cobaltic ion itself are not well known. This limits the usefulness of the data given in Table 6. It is tempting to carry out calculations based on the values of the corresponding quantities of similar complexes of the ferric ion. It is, however, doubtful whether this will be valid knowing that most of the cobaltic complexes are inner sphere in nature while a majority of those of Fe(III) are of the outer sphere type. However, qualitative comparisons may be made, because evidence suggests that a paramagnetic state lies fairly close to the usual diamagnetic one.
Isotopic exchange in the cobaltous--cobalticsystem in perchloric acid--Ill
1989
ASHURT et aL ¢1a~ have obtained a value of 22 4- 7 F -1 for K¢coso - ) at p = 2.7 and 15°C. From this and Table 6 one gets a value of ,~1.2 × 103 F -x min - t at 14° and p = 2 for k~so~2_). Our observations agree with those of HAmB et al. ~2~in that there is an appreciable sulphate-ion effect. We were however, unable to find any appreciable fluoride-ion effect even at double the concentration of sodium fluoride used by Habib et al. The large effect observed by these authors cannot be explained on the basis of higher working temperature (18.55°C vs. 0°C) and slightly higher free fluoride-ion content of the reaction mixtures used by them in the study of the fluoride-ion effect. In these experiments the authors have worked at an initial acidity of 0.2 N. Measurements made here using quinhydrone electrodes showed that on making a 0.2 N solution of perchloric acid 0.1 F with respect of sodium fluoride, the free acid concentration decreased to almost 0.1 N. Under these conditions the observed increase in exchange can be put down as almost entirely due to an increase in the acid-dependent exchange. Our observations of the presence of an acetate-ion catalysed path for exchange have not been borne out by those of HABIB and HUNT~2. These authors however, have worked at acetate-ion concentrations insufficient to cause any noticeable effect. Influence o f heavy water
The rates of exchange in the cobaltous-cobaltic system are appreciably affected by a change of solvent from light to heavy water, as can be seen from Tables I, 2 and 5. From the data in Table 5, one can calculate at 0°C and/z ---- 3, the values of klKh and k2 (where kl and k2 are the velocity constants for the acid-dependent and acidindependent paths and Kh is the hydrolysis constant of the cobaltic ion). These values come out to be 27 min -1 and 66 F -1 min -~ in light water and 13 min -~ and 32 F -~ min -1 in 93 ~o heavy water. The decrease in the value of k~Kh on changing to heavy water may be due to change in kx, Kh or both. HUDIS et al. {2°~ have shown that the hydrolysis constant of the reaction: Fe 3+ ÷ H20 ~ FeOH ~+ -k H+, is not influenced by the change. A similar observation has been made by ROGERS et al. ~2~) on the hydrolysis of the thallic ion. The change in k l K h on changing the solvent can therefore be attributed to the change in kl, the velocity constant of the acid-dependent exchange. It is not possible in the light of present knowledge to pinpoint the cause of the appreciable heavy water effects on the values of k(A~-) • K(coA3-,) observed in the cases of sulphate and acetate ions (Table 6). The heavy water effect found on the acid-dependent path is consistent with the mechanism proposed c1°) earlier for exchange through this path. However, a similar effect is also found for the acid-independent path which does not seem consistent with the mechanism suggested (1°~ for this path. Recent studies (2~) of heavy water effects have shown, however, that evidence for a particular type of mechanism obtained from t19~K. G. ASrIURTand W. C. E. I-IIGGtNSON,J. chem. Soc. 343 (1956). t20~j. HtrDISand R. W. DOBSON,J. Am. chem. Soc. 78, 911 (1956). {2a~T. E. ROOERand G. M. WAIND, Trans. Faraday Soc. 57, 1360 (1961). ~22~W. KRUSEand H. TAttoo,J. Am. chem. Soc. 82, 526 (1960); A. E. O~ARDand H. TAtmE,J. Am. chem. Soc. 80, 1084 (1958); N. StrrlN, J. K. ROWLEYand R. W. DOBSON,J.phys. Chem. 65, 1248 (1961).
1990
J. SHANKARand B. C. DE SOUZA
heavy water studies should be treated with reserve until the properties of heavy water are well understood. Argentous ion effect on e x c h a n g e
A study of the effect of silver ion on exchange in the cobaltous--cobaltic aquo system was suggested by the observations of NoYES et al. c~a~. Figure 1 shows that addition of small quantities of silver'nitrate leads to a considerable increase in exchange in the system. In the presence of argentous ions the following equilibrium will be established: (1)
Ag + + Co s+ ~- Ag ~+ + Co ~-.
Formation of Ag ~- will result in an additional path being made available for transfer of electrons from cobaltous to cobaltie species. Since the plot (Fig. 1) is not linear in the range of silver ion concentrations used here, the above process is evidently not the only one responsible for the observed silver ion dependent exchange. The following equilibrium has been previously proposed by NoYEs et al. ~ (2)
Ag e+ ~-- Ag + + Aga+.
The Aga+ so formed can facilitate exchange through the establishment of the following equilibrium: (3)
Ag a+ + Co ~+ ~---Ag ~+ + Co 3+
An approximate idea of the overall equilibrium constant K, can be obtained from the half electrode potentials t~5~for the reactions Ag + ,~ Ag e+ + e and Co 2+ ,~- Co 3+ + e. These values are respectively -- 1-84 V and -- 1.98 V at 25°C. K c o m e s out to be 4.6 × 10-s. The actual value may be somewhat higher because the value of --1.84 V has been obtained in 3 N nitric acid where appreciable complexing can occur. The magnitude of K indicates that only a very small fraction of the added silver ions exists in the higher-valent states. Thus it is evident that the velocity constants of the argentous ion dependent paths must be fairly high to account for the large increase in the overall exchange rate on the addition of the ions. From the results it appears that the cobaltous species part with their electrons with equal if not greater ease to the silver species as they do to the cobaltic ones, This is probably due to the larger radii of Ag z+ and Ag a+ and the lower charge of Ag e+ (as compared to those of Co 3+) which result in lower excess free energies (in terms of MARCUS' modeF26)). In a previous publication, tg) the effect of ionic strength on exchange rates in the cobaltous--cobaltic system was found to differ from that expected on the basis of the Debye-Huckel equation. This adverse effect had been ascribed to appreciable perchlorate complexing of cobalt ions. In the light of the present work, it is felt that over-emphasis had been placed on the implications and applicability of the equation. t2a~A. A. NOYESand T. J. DEAHL,J. Am. chem. Soc. 59, 1337 (1937). ~4~ A. A. NOVES,C. D. CORYELL,F. STITTand A. KOSSAIKOFF,.L Am. chem. Soc. 59, 1316 (1937). A. A. NoYEs, J. L. HOARD,and K. S. PrrzER,J. Am. chem. Soc. 57, 1221 (1935); A. A. NOYES, K. S. PITZERand C. L. Dtn,rN, J. Am. chem. Soc. 57, 1229 (1935). t~5~R. PARSONS,Handbook of Electrochemical Constants, p.54. Butterworths, London (1959). t~a~R. A. MARCOS,J. chem. Phys. 26, 867 (1957).
Isotopic exchangein the cobaltous-cobalticsystemin perchloricacid--III
1991
CONCLUSION The behaviour of the system under study appears to be similar to that of the ferrous-ferric one as far as can be judged from the influence of acidity, anions and change of solvent from light to heavy water. Thus the electron-transfer mechanism, through water molecule bridges, with the anions accentuating exchange by drawing the two exchanging species closer together, which has been proposed ~a~ to explain exchange in the ferrous-ferric system, is probably operative in the present system also. Thus the ionic strength may be due to reduction in the number of water-bridges available for exchange, due to displacement of some of the water molecules from the vicinity of the cobalt ions. The lack of fluoride ion effect indicates that fluoride ions do not form true complexes with cobalt ions or only form very weak ones. HUDtS et al. ~2~ have also observed that the fluoride ion effect in the ferrous-ferric system is not as large as would be expected on the basis ofits large charge-to-size ratio. SVTCLIFFEetal.t2~havefound that fluoro-complexes of cobalt play a negligible role in the fluoride ion effect on the oxidation of Ce~+ by Co 3+. ~ L. H. StlTCLIFFEand J. R. W~BER,Trans. Faraday Soc. 55, 1892 (1959).
ISOTOPIC E X C H A N G E IN THE COBALTOUS-COBALTIC SYSTEM IN PERCHLORIC A C I D - - I I I INFLUENCE OF SOME IONS A N D HEAVY WATER ON EXCHANGE J. SHANKAR a n d B. C. DE S o u z A Bhabha Atomic Research Centre, Chemistry Division, Bombay, India (First received 23 December 1966; in revised form 13 March 1967)
Almtraet--Sulphate, acetate and nitrate ion have been found to enhance exchange in the cobaltouscobaltic aquo system in perchloric acid. Fluoride ion has no appreciable effect. Changing the solvent from light to heavy water reduces exchange proceeding through acid-independent and acid-dependent paths by a factor of about 2 through a reduction in the velocity constants of the two paths. Exchange through sulphate and acetate dependent paths is also reduced considerably in heavy water, but it is not possible to decide in the light of the evidence at present available whether the reduction is due to a change in the velocity constants for the exchange processes or to a change in the values of the formation constants of the complexes through which the exchange proceeds or to a change in both. The addition of small amounts of silver ion considerably increases the exchange in the system. The processes responsible for this are briefly discussed. PRONOUNCED effects due to the presence o f ions have b e e n f o u n d in a n u m b e r o f r e d o x reactions. I s o t o p i c exchange in t h e f e r r o u s - f e r r i c system for instance, is influenced b y chloride, cl) fluoride, (2) b r o m i d e , (3) sulphate, (4) azide, (5) t h i o c y a n a t e ~6) a n d o x a l a t e ions33) Evidence o b t a i n e d f r o m studies o f these effects has h e l p e d c o n s i d e r a b l y in the f o r m u l a t i o n o f theories on the m e c h a n i s m s o f the r e a c t i o n s involved. F o r a similar reason, studies o f the effect o f h e a v y w a t e r on exchange a n d o t h e r r e a c t i o n s have also been c a r r i e d out. This p a p e r describes a s t u d y o f similar effects in t h e c o b a l t o u s cobaltie system, w o r k which was carried o u t with a view to gaining f u r t h e r insight i n t o t h e n a t u r e o f exchange in the system, which has a l r e a d y received c o n s i d e r a b l e attention. ~7-u) S o m e o f the w o r k described here overlaps t h a t p u b l i s h e d recently b y HABIB et al. (~) EXPERIMENTAL Preliminary experiments showed that some anions, like chloride and thiocyanate, were quickly oxidized by cobaltic ions. Sulphate, acetate, fluoride, nitrate and argentous ions were found suitable for study. These were added to the reaction mixtures in the form of standard solutions of sulphuric acid, acetic acid, sodium fluoride, nitric acid and silver nitrate. The experiments were carried out as before~9~; the added ions did not interfere with the procedure except for fluoride ion, the presence of which prevented spectrophotometric estimation of Co(III) ~xl j. SILVERMANand R. W. DODSON,J. phys. Chem. 56, 846 (1952). ~2)j. HUDIS and A. C. WAHL,3". Am. chem. Soc. 75, 4153 (1953). ~s~R. A. HORNE,d. phys. Chem. 64, 1512 (1960). c4) K. H. LIESER and H. SCHROEI)ER,d. inorg, nucl. Chem. 14, 98 (1960); R. L. S. WILLIX, Trans. Faraday Soc. 59, 1315 (1963). (51 D. BONN, F. S. DAINTONand S. DUCKWORTH, Trans. Faraday Soc. 55, 109 (1960). ~6~G. S. LAWRENCE,Trans. Faraday Soc. 53, 1326 (1957). t~) N. A. BONNEt*.and J. P. HUNT, J. Am. chem. Soc. 74, 1866 (1952). ts~ N. A. BONNERand J. P. HUNT, J. Am. Chem. Soe. 82, 3826 (1960). c9) j. SHANKARand B. C. DE SOUZA,J. inorg nucL Chem. 24, 187 (1962). ix0) j. SHANKARand B. C. DE SOUZA,d. inorg, nucl. Chem. 24, 694 (1962). ~lx~B. C. DE SOUZA,Ph.D. Thesis, Univ. of Bombay (1962). ¢x2)H. S. HABIBand J. P. HUNT, J. Am. chem. Soc. 88, 1668 (1966). 1983
1984
J. SHANKAR and B. C. DE SOUZA
through the thiocyanate complex of Fe(III). In this case a blank mixture containing all the reactants in the same quantities with the exception of fluoride ion, was simultaneously prepared and used for the estimations. The concentrations of the different ions were such that: (a) their addition (i.e. replacement for the perchlorate ions (bulk electrolyte) by these ions) did not cause any activity coetficient changes, (b) the concentration of anionic complexes containing more than one anion in the co-ordination spheres was kept low, (c) the half-time of overall exchange lay generally between 3-25 min. Longer or shorter half-times led to errors in their determination--in the former case due to appreciable decomposition of the cobaltie species during the course of an experiment (six half-times) and in the other due to sampling time, (d) the acidity of the reaction mixtures was kept high so as to avoid complications arising from the presence of appreciable quantities of hydrolysed species (some of unknown composition) as well as the difficulty of estimating [H +] at low concentrations. Calculation of ion concentrations: Concentrations of argentous and nitrate ions were taken as being equal to the formalities of silver nitrate and nitric acid solutions respectively, as these compounds dissociate completely at the concentration used in the experiments. A knowledge of the corresponding dissociation quotients is necessary for the calculation of the actual concentrations of sulphate, acetate and fluoride ions. Sulphate ion concentration was calculated as follows. Complete dissociation of H, SO4 into HSO4- and H + was assumed. KERKERt18~has reported a value of 1.083 F at/~ = 2 and 18°C for the dissociation quotient (Q,). Using this and the value of heat dissociation which has been reported as --5.2 :k 0.5 kcal, c.4~ the Q2 value at 0°C and # = 2 was calculated to be 1.95. Knowing this, the sulphate ion concentration at any value of [H +] could be calculated in terms of the initial concentration, c, of the added sulphuric acid according to the relation: 2 " C
[SO?-]
-
-
-
Q, + [H +] •
BAKERet aL t*5~have shown that at 25°C and/z = 1, the dissociation quotient of sulphuric acid is 1"75 :k 0"17 times lower when the solvent is changed from light to 8970 heavy water. This is also the value to be expected on the basis of the rule of Rule and LaMer. Assuming that the ratio is not much affected by temperature and ionic strength, Q, was calculated as being 1.10 at 0°C and/~ = 2 in heavy water. Similar calculations were made for the determination of fluoride ion using the data calculated by HuDiS et aL ('~ at/~ = 0"5 based on the values of BRo~r,m et aL (16) In the study of the acetate ion effect, [OAc-] values were calculated using the relation: KD [OAc-] ~ ~-7]'" e where c is the initial concentration of acetic acid and K~) is its dissociation constant. From the reported value t~7) of 1-657 × 10-5 F for the dissociation constant of acetic acid at 0°C and/~ = 0, the corresponding value at 0°C and # = 1 at which these experiments were conducted, was found to be 5.27 x 10-5 F. The value of the constant B in the extended Debye--Huckel equation is 0.324a ° at 0°C, where a ° is the mean ionic radius in angstroms. Assuming a ° to be nearly equal to the radius of a molecule of acetic acid, the value taken for it was 2.03 A, based on the density of glacial acetic acid. Due allowance has been made for the larger value of a ° obtained from this calculation by applying a correction based on the calculated and experimental values of the radius of a water molecule, the assumption being made that the correction factor is similar in the two cases. The decrease in the dissociation constant of acetic acid in 100 70 heavy water has been reported c16~ to be by a factor of 3"3. Assuming linear change in this factor over a small range of deuterium content, tlS~ M. KERKER,J. Am. chem. Soc. 79, 3664 (1957). cx4~K. S. Pn'ZER, J. Am. chem. Soe. 59, 2365 (1937). (xs) F. B. BAKERand T. W. NEWTON,J. phys. Chem. 61, 381 (1957). ~xs~H. H. BROENEand T. DEVRIES,J. Am. chem. Soc. 69, 1644 (1947). cxT~H. S. HARr,~D and R. W. ErmER, d. Am. chem. Soc. 55, 652 (1933). tlS~ S. KORMANand V. K. LAMER, J. Am. chem. Soc. 58, 1396 (1936).
Isotopic exchange in the cobaltous--cobaltic system in perchloric acid--III
1985
its value was taken as 3.06 in 93 ~ heavy water, the concentration at which these experiments were carded out. The dissociation constant under these conditions was calculated to be 1.72 x 10-5 F. In the calculation of actual ion concentration, the effect of complexation on these concentrations of ions, has been considered as being negligible under the experimental conditions. The heavy water used in some of the experiments was 98 per cent in deuterium content. Experiments made for the study of the heavy water effect were conducted in the same way as the others but the volumes of the reaction mixtures were kept between 20 to 30 ml instead of the usual 50 ml. Wherever possible solutions of salts were prepared in and dilutions done with heavy water. The H20 in the final reaction mixtures came mainly from the 70 Yoperchloric acid, used for maintenance of the acidity of the mixtures without any pretreatment for replacement of its hydrogen by deuterium. RESULTS T a b l e s 1, 2, 3 a n d 4 s h o w t h a t s u l p h a t e , a c e t a t e a n d n i t r a t e i o n s e n h a n c e e x c h a n g e in the cobaltous-cobaltic aquo system while fluoride ion does not appear to have any n o t i c e a b l e effect. T a b l e s 1 a n d 2 a l s o s h o w t h a t c h a n g i n g t h e s o l v e n t f r o m l i g h t t o h e a v y w a t e r h a s a n a p p r e c i a b l e effect o n t h e s u l p h a t e a n d a c e t a t e i o n c a t a l y s e d e x c h a n g e . D a t a o n t h e a n i o n - i n d e p e n d e n t e x c h a n g e at 0 ° C a n d 10°C in h e a v y w a t e r a r e g i v e n i n T a b l e 5. F i g u r e 1 s h o w s t h e i n f l u e n c e o f a r g e n t o u s i o n o n t h e e x c h a n g e . TABLE 1.--EFFECT OF SULPHATE ION ON EXCHANGE. CONCENTRATION OF PERCHLORIC ACID 2"0 F . IONIC STRENGTH 2"0. TEMPERATURE OF RUNS 0 ° C
Medium H~O Concentration of total cobalt ,---0"9 × 10-s F
Medium 88 ~ D~O (v/v) Concentration of total cobalt ~1.3
× 10 -8 F
Initial concn of H2SO4 c( × 10~-F)
Velocity constant k(F -x min -~)
Initial conch of H2SO4 c( × l0 s F)
Velocity constant k(F -~ rain -1)
0"21 0"34 0"56 0"90 1"15 0"00
65 72 84 111 125 52
0"50 0"70 I'00 1"25 1"50 0'00
46 52 62 70 79 26
TABLE 2.--EFFECT OF ACETATEION ON EXCHANGE. CONCENTRATIONOF PERCHLORIC ACID 1"0 F. IONIC STRENGTH 1"0. TEMPERATUREOF RUNS 0°C Medium H20 Concentration of total cobalt ~1"0 × 10 -8 F
11
Medium 93 ~ D~O (v/v) Concentration of total cobalt ~ 2 . 0 × 10-a F
Initial concn of CH3COOH c(F)
Velocity constant k(F -1 min -1)
Initial concn of CH3COOH c (F)
Velocity constant k(F -1 min -1)
0"10 0'20 0'37 0"50 0"70 0"00
52 59 72 87 104 43
0-20 0"37 0"50 0"70 1-00 1-50 0.00
28 32 34 42 49 57 22
1986
J. SHANKAR and B. C. DE SOUZA
TABLE 3.~EFFECT OF NITRATE ION ON EXCHANGE. CONCENTRATION OF PERCHLORIC ACID 2"0 F. I o m c STRENGTH 2"0. TEMPERATUREor RUNS 0°C Concentration o f total cobalt
,---0.31 × 10-a F
Initial concn o f HNO3 e(F)
. Velocity constant k(F -x min -1)
0"15 0"20 0.30 0.40 0"50 0.00
59"0 61"5 66.0 69"5 75.5 52.0
TABLE 4 . - - E F F E C T OF FLUORIDE ION ON EXCHANGE. CONCENTRATION OF PERCHLORIC ACID 0"5 F . IONIC STRENGTH 0"5. TEMPERATURE OF RUNS 0 ° C
Concentration o f total cobalt Initial concn of NaF e(x 102 F)
,~1.4 × 10 - s F
[F-](x 104 F)
0-3 0"8 1.2 3"5 5-0 10-0 15"0 20"0 0.0
Velocity constant k(F -1 rain -I)
0-1 0"28 0"42 1.35 2-01 4.35 7"92 11.5 -
44 43 44 43 44 41 46 52 44
-
TABLE 5.---EFFECTOF HEAVYWATERON EXCHANGE. IONIC STRENGTH3"0. TEMPERATURE OF RUNS 0°C Medium H20 Concentration of total cobalt ,--4).5 × 10-s F
Medium 93 ~ D20 (v/v) Concentration of total cobalt ~ I . 0 x 10-8 F
Concentration of HC104 [H+](F)
Velocity constant k(F -1 min -1)
Concentration of DC104 [D+](F)
0"26 0"36 0"47 0"58 0"67
163 142 130 114 106
0'24 0"32 0"54 0"64 0"76
Velocity constant k(F -1 rain -1) 92 69 61 52 48
Isotopic exchangein the cobaltous-cobalticsystemin perchloric acid--lll
1987
i
-r -~ 12o b_
Io0 7
8¢
6¢
40
I
I
I
2
I 3
r 4
I
5
__
[Ag +] xlO 4 F
FIG. 1.--Effectof argentous ion concentrationon exchangein the cobaltous-cobaltic aquo system. Temperature= 0.0°C. Ionic strength =- 1.0. Concentrationsof perchloric acid, cobaltic and cobaltous perchlorates are 1.0 F, 0.5 × 10-s F and 1.2 × 10 -8 F.
DISCUSSION
Anion-dependent exchange The results obtained when sulphate, acetate and nitrate ions are present in the reaction mixtures can be expressed in the following empirical form: k ---- a + b[A"-],
(1)
where a and b are constants and [A"-] represents the actual concentration of the anion A"-. The first order dependence on ion concentration indicates the participation of one anion in the activated complex through which the anion-dependent exchange proceeds. The most probable mode of its incorporation into the activated complex is through its prior complexation with one of the exchanging ions. Evidence obtained from studies of similar anion complexes of other elements suggests that electrostatic forces, dependent more on the charges and sizes of the complexing ions than on their chemical nature, predominate in these type of complexes. These considerations favour the formation of cobaltic rather than cobaltous complexes with the anion. Again, the tendency of an ion to reduce its charge through complexation with anions will favour the formation of cobaltic complexes. Thus these species will be present in solution in appreciable quantities. The charge difference between complexed cobaltic species and uncomplexed cobaltous one will be less than that between the uncomplexed cobaltous and cobaltie species or uncomplexed cobaltic and complexed cobaltous ones. Thus the environmental and structural differences between the two exchanging species, being governed
J. SI-IANKARand B. C. DE SOUZA
1988
by this difference, will tend to favour the following exchange: ,Co2+ + COA3q. _ , CoA.q 3 - . + Coaq, 2+ the overall rate of which will be given by the equation: R = k2[Co~][Co3+] + kl[Co~][CoOH ~-] + k(~,_)[CoZ+][CoAa-n], where the first two terms are the contributions of the anion-independent mechanisms (previously discussed) and k(x,_ ) is the velocity constant for the anion-dependent exchange. The suffix aq has been, and will be, omitted for the sake of convenience. TABt~6.--VALUES OF kcx,,-}g[coAa-n)FOR THE DIFFERENT ANION-DEPENDENT PATHS OF EXCHANGE IN
HzO AND D~O
AT 0 ° C
Added anion
Ionic strength
SO42-
1
H20 8 0 ~ DzO
1"31 -4-0"15 × 104 0"99 + 0"10 × 104
CHaCOO-
1
HzO 93 ~o D . O
1"70 -4- 0"4 × l0 s 0"68 4- 0"06 × 106
NOB-
2
H~O
45"7 4- 4"8
Solvent
k~--~Kcoox3--j
It can easily be shown by combining the above equation with that obtained from kinetics, i.e. R = k[Co(III)][Co(III)], where k is the overall rate constant, that: k ---- k 0 + k(~,-)K(CoA~-.)[A~-],
(2)
where k 0 is the contribution from the anion-independent paths of exchange. This is constant ff the [H+] of the reaction mixtures is constant. K(coxs_,) is the formation constant of the complex CoAa-~. The above equation holds strictly only when the concentrations of species of the type CoA~-~ and COA32-2n are very small. It can thus be seen that constants a and b in Equation (1) are identifiable with k o and k(x,- ) K(cox,-, ). The values of the latter for the different anions in light and heavy water have been given in Table 6. Because of the instability of the cobaltic ion in aqueous solutions and its strongly oxidizing nature, the study of its properties has been mainly restricted to its use as a powerful oxidant for organic compounds. Thus, while its complexes, which are more stable, have been extensively studied, the properties of the cobaltic ion itself are not well known. This limits the usefulness of the data given in Table 6. It is tempting to carry out calculations based on the values of the corresponding quantities of similar complexes of the ferric ion. It is, however, doubtful whether this will be valid knowing that most of the cobaltic complexes are inner sphere in nature while a majority of those of Fe(III) are of the outer sphere type. However, qualitative comparisons may be made, because evidence suggests that a paramagnetic state lies fairly close to the usual diamagnetic one.
Isotopic exchange in the cobaltous--cobalticsystem in perchloric acid--Ill
1989
ASHURT et aL ¢1a~ have obtained a value of 22 4- 7 F -1 for K¢coso - ) at p = 2.7 and 15°C. From this and Table 6 one gets a value of ,~1.2 × 103 F -x min - t at 14° and p = 2 for k~so~2_). Our observations agree with those of HAmB et al. ~2~in that there is an appreciable sulphate-ion effect. We were however, unable to find any appreciable fluoride-ion effect even at double the concentration of sodium fluoride used by Habib et al. The large effect observed by these authors cannot be explained on the basis of higher working temperature (18.55°C vs. 0°C) and slightly higher free fluoride-ion content of the reaction mixtures used by them in the study of the fluoride-ion effect. In these experiments the authors have worked at an initial acidity of 0.2 N. Measurements made here using quinhydrone electrodes showed that on making a 0.2 N solution of perchloric acid 0.1 F with respect of sodium fluoride, the free acid concentration decreased to almost 0.1 N. Under these conditions the observed increase in exchange can be put down as almost entirely due to an increase in the acid-dependent exchange. Our observations of the presence of an acetate-ion catalysed path for exchange have not been borne out by those of HABIB and HUNT~2. These authors however, have worked at acetate-ion concentrations insufficient to cause any noticeable effect. Influence o f heavy water
The rates of exchange in the cobaltous-cobaltic system are appreciably affected by a change of solvent from light to heavy water, as can be seen from Tables I, 2 and 5. From the data in Table 5, one can calculate at 0°C and/z ---- 3, the values of klKh and k2 (where kl and k2 are the velocity constants for the acid-dependent and acidindependent paths and Kh is the hydrolysis constant of the cobaltic ion). These values come out to be 27 min -1 and 66 F -1 min -~ in light water and 13 min -~ and 32 F -~ min -1 in 93 ~o heavy water. The decrease in the value of k~Kh on changing to heavy water may be due to change in kx, Kh or both. HUDIS et al. {2°~ have shown that the hydrolysis constant of the reaction: Fe 3+ ÷ H20 ~ FeOH ~+ -k H+, is not influenced by the change. A similar observation has been made by ROGERS et al. ~2~) on the hydrolysis of the thallic ion. The change in k l K h on changing the solvent can therefore be attributed to the change in kl, the velocity constant of the acid-dependent exchange. It is not possible in the light of present knowledge to pinpoint the cause of the appreciable heavy water effects on the values of k(A~-) • K(coA3-,) observed in the cases of sulphate and acetate ions (Table 6). The heavy water effect found on the acid-dependent path is consistent with the mechanism proposed c1°) earlier for exchange through this path. However, a similar effect is also found for the acid-independent path which does not seem consistent with the mechanism suggested (1°~ for this path. Recent studies (2~) of heavy water effects have shown, however, that evidence for a particular type of mechanism obtained from t19~K. G. ASrIURTand W. C. E. I-IIGGtNSON,J. chem. Soc. 343 (1956). t20~j. HtrDISand R. W. DOBSON,J. Am. chem. Soc. 78, 911 (1956). {2a~T. E. ROOERand G. M. WAIND, Trans. Faraday Soc. 57, 1360 (1961). ~22~W. KRUSEand H. TAttoo,J. Am. chem. Soc. 82, 526 (1960); A. E. O~ARDand H. TAtmE,J. Am. chem. Soc. 80, 1084 (1958); N. StrrlN, J. K. ROWLEYand R. W. DOBSON,J.phys. Chem. 65, 1248 (1961).
1990
J. SHANKARand B. C. DE SOUZA
heavy water studies should be treated with reserve until the properties of heavy water are well understood. Argentous ion effect on e x c h a n g e
A study of the effect of silver ion on exchange in the cobaltous--cobaltic aquo system was suggested by the observations of NoYES et al. c~a~. Figure 1 shows that addition of small quantities of silver'nitrate leads to a considerable increase in exchange in the system. In the presence of argentous ions the following equilibrium will be established: (1)
Ag + + Co s+ ~- Ag ~+ + Co ~-.
Formation of Ag ~- will result in an additional path being made available for transfer of electrons from cobaltous to cobaltie species. Since the plot (Fig. 1) is not linear in the range of silver ion concentrations used here, the above process is evidently not the only one responsible for the observed silver ion dependent exchange. The following equilibrium has been previously proposed by NoYEs et al. ~ (2)
Ag e+ ~-- Ag + + Aga+.
The Aga+ so formed can facilitate exchange through the establishment of the following equilibrium: (3)
Ag a+ + Co ~+ ~---Ag ~+ + Co 3+
An approximate idea of the overall equilibrium constant K, can be obtained from the half electrode potentials t~5~for the reactions Ag + ,~ Ag e+ + e and Co 2+ ,~- Co 3+ + e. These values are respectively -- 1-84 V and -- 1.98 V at 25°C. K c o m e s out to be 4.6 × 10-s. The actual value may be somewhat higher because the value of --1.84 V has been obtained in 3 N nitric acid where appreciable complexing can occur. The magnitude of K indicates that only a very small fraction of the added silver ions exists in the higher-valent states. Thus it is evident that the velocity constants of the argentous ion dependent paths must be fairly high to account for the large increase in the overall exchange rate on the addition of the ions. From the results it appears that the cobaltous species part with their electrons with equal if not greater ease to the silver species as they do to the cobaltic ones, This is probably due to the larger radii of Ag z+ and Ag a+ and the lower charge of Ag e+ (as compared to those of Co 3+) which result in lower excess free energies (in terms of MARCUS' modeF26)). In a previous publication, tg) the effect of ionic strength on exchange rates in the cobaltous--cobaltic system was found to differ from that expected on the basis of the Debye-Huckel equation. This adverse effect had been ascribed to appreciable perchlorate complexing of cobalt ions. In the light of the present work, it is felt that over-emphasis had been placed on the implications and applicability of the equation. t2a~A. A. NOYESand T. J. DEAHL,J. Am. chem. Soc. 59, 1337 (1937). ~4~ A. A. NOVES,C. D. CORYELL,F. STITTand A. KOSSAIKOFF,.L Am. chem. Soc. 59, 1316 (1937). A. A. NoYEs, J. L. HOARD,and K. S. PrrzER,J. Am. chem. Soc. 57, 1221 (1935); A. A. NOYES, K. S. PITZERand C. L. Dtn,rN, J. Am. chem. Soc. 57, 1229 (1935). t~5~R. PARSONS,Handbook of Electrochemical Constants, p.54. Butterworths, London (1959). t~a~R. A. MARCOS,J. chem. Phys. 26, 867 (1957).
Isotopic exchangein the cobaltous-cobalticsystemin perchloricacid--III
1991
CONCLUSION The behaviour of the system under study appears to be similar to that of the ferrous-ferric one as far as can be judged from the influence of acidity, anions and change of solvent from light to heavy water. Thus the electron-transfer mechanism, through water molecule bridges, with the anions accentuating exchange by drawing the two exchanging species closer together, which has been proposed ~a~ to explain exchange in the ferrous-ferric system, is probably operative in the present system also. Thus the ionic strength may be due to reduction in the number of water-bridges available for exchange, due to displacement of some of the water molecules from the vicinity of the cobalt ions. The lack of fluoride ion effect indicates that fluoride ions do not form true complexes with cobalt ions or only form very weak ones. HUDtS et al. ~2~ have also observed that the fluoride ion effect in the ferrous-ferric system is not as large as would be expected on the basis ofits large charge-to-size ratio. SVTCLIFFEetal.t2~havefound that fluoro-complexes of cobalt play a negligible role in the fluoride ion effect on the oxidation of Ce~+ by Co 3+. ~ L. H. StlTCLIFFEand J. R. W~BER,Trans. Faraday Soc. 55, 1892 (1959).