4352
EPHRAIRI BEN-ZVJ
AND
THOMAS L. ALLEN
Vol. 53
aqueous solutions. However, it seeins reasonable vations were mnde of the effect on the aiiirnonin to assume that the implication of “free” is that such spectrum of the addition of water to diethyl ether molecules are not specifically solvated but are solutions of ammonia. There were measursd the simply subjected to the forces of the general dipolar optical densities a t 2125, 2150 and 2175 A. (1 0 environment of the medium, water. If “free” has cm. cell), of ether solutions of NH3 with various this meaning, then it seems reasonable to assume concentrations of water. As may be seen in Table that the spectral behavior of such “free” ammonia 11, the presence of 5 to 10 ml. of HZO per liter of molecules would be very like that of ammonia ether causes a large reduction in the effectivc. abmolecules in diethyl ether solution, where there are sorptivity of the ammonia solutions a t these wnvc the strong general dipolar forces associated with lengths. the C-0 bonds but no possibility of specific solvaTABLE I1 tion such as is implied in the formula NHbOH. l 2 Subject to the assumptions indicated in the pre- EFFECT O F HsO O N THE SPECTRUM O F ~1~~I N DIE,TIITL ETHER vious paragraph, one can conclude that the absorpM 0.0139 0.0130 0.1::o tion spectrum of aqua ammonia, as given by Curve ? f (”I), __ Co (HzO), M 0.0000 ,0334 I1 of Fig. 1, is the sum of two quantities O.D. (1 cm.) 2125 A ( a q . NHs)
= (1
-
W )
‘lNIi8
+
1.460
aAhIIpOIi
2150 2175 K
(3)
0 . (io5
,686 332 . 1.A7
. ZY:,
.2fj
0.337 ,132 where (Y is the fraction of the dissolved ammonia in 0.454 + 0.012 0.406 i 0 . 0 0 5 the form of NHdOH, ie., specifically solvated, 3 . 4 i 0.1 2.6 * 0.1 ANH*OH is the absorptivity of NH40H a t the wave If one assumes the optical density of each of these is the absorptivity of “free” length A, and solutions is proportional to the concentration of ammonia molecules and is to be taken equal to that “free” ammonia in the solution, one can calculate of a diethyl ether solution a t A. Since OIAXH~OII is of necessity a positive quantity, we h a r e from (3) an apparent association constant for the reaction for the fraction of free ammonia molecules, 1- a , XHd H20 S H a O H ( i n ether) (:) the inequality from the expression
+
and the conclusion that less than 0.5y0of the animonia in water solutions is “free.” Similar reasoning applied to the data on tlie vapor’ and water solution spectra of water, phosphine and hydrogen sulfide leads to the conclusion that less than 0.2% H20, 30% PH3 and 4276 H,S are “free” in water solution a t about 23”. In order to ascertain that the decrease in absorptivities associated with change of solvent from ether to water in the case of ammonia is appropriately attributable to specific complex formation, obser(12) By t h e formula, XHiOH, we imply an “outer complex” in t h e sense of Fig. 2 of the definilivc paper of R. S . hlulliken, J . P h y s . Chem., 66, 801 (1952).
since the total NHs concentration in the solution is small compared with the water concentratioii. As may be seen in the last line of Table 11, tlic spectral data give -3 l./mole for the (NHa--I-IGQ) association constant in the ether solution. Such a value for the association constant would imply that considerably less than lY0 of the ammonia in water solution is free, if the association constant were independent of solvent. Clearly the results of the measurements on the effect of water on the spectrum of the ether solution of ammonia are completely compatible with the interpretation given above of the difference between the spectra of the water and the ether solutions of ammonia.
[COSTRIBU TION PROX THE DEPARTXENT OF CHEMISTRY, vsIVERSITY O F CALIFORSIA,
DAVIS,CALIFORNIA]
The Oxidation of Oxalate Ion by Peroxodisulfate. 11. The Kinetics and Mechanism of the Catalysis by Copper(11) BY EPHRAIM BEN-ZVI~ AND THOMAS L. ALLEN RECEIVED MAY20, 1961 The oxidation of oxalate ion by peroxodisulfate is strongly catalyzed by copper( 11),but the catalysis is subject to inhibition by molecular oxygen. An investigation has been made of the kinetics of the catalyzed reaction in the absence of oxygen. The rate law is first-order in peroxodisulfnte, zero-order in oxalate, and half-order in the catalyst. A free-radical chain mechanism, involving oxidation of copper to the terpositive state, is postulated for the reaction. It is shown that copper does not take part in the chain-initiating step. Reduction of copper t o the unipositive state prohably does not occur.
I n the first part of this series3 it was shown that copper(I1) is a very effective catalyst for tlic oxi-
dation of oxalate ion by the &Os= and a brief study of the kinetics of the copper-catalyzed re-
(1) This work was assisted b y a research g r a n t from tlie National Science Foundation. ( 2 ) Abstracted in part from the P h . D . Dissertation 01 Ephraim Ben-Zvi, University of California, Davis, 1900. 73, 3ZSQ (1951). ( 3 ) T. L. Allen. J. A m . Chem. SOL.,
(4) I n common usage 5 2 0 8 - is persulfate. Under t h e “ 1940 Rules” [W. 1’. Jorissen, H. Bassett, A . Damiens, F. Fichter and 1%.R i m y , ibid., 63,889 (1911)) and in Chemical Abstracts i t is designated peroxydisulfate. I n the ’‘ Definitive Rules for Nomenclature of Inorganic Chemistry,” i b i d . , 8’2, 5j23 ( l 9 0 0 ) l i t is peroxodisulfate.
Nov. 5 , 1961
COPPERCATALYSIS IN PEROXODISULFATE OXIDATIOK OF OXALATE
action was included. A detailed investigation now has been made and the results are reported below. Experimental Materials.-Redistilled water was used in all of the kinetic experiments. It was prepared from ordinary distilled water by successive distillation from dilute alkaline KMnO4 solution, dilute KHS04 solution and finally, without addition of any reagent, through a 36-inch Vigreux column. The connection between the column and the condenser was heated with a heating tape t o break the liquid film on the inner wall of the distillation system, thus preventing diffusion of impurities into the distillate.6 Potassium peroxodisulfate was recrystallized from redistilled water and dried under vacuum a t 50". The purified reagent contained an acidic substance (probably KHSOP) which lowered the PH of the reaction solution, containing five times a s much h'azCzO4 as K~S208,t o 6.0 f 0.2. ( A slightly acidic solution was desirable t o avoid precipitation of the catalyst.) T o provide an inert atmosphere Matheson prepurified NZand "bone-dry" COZ were used. All other substances were reagent grade. So that reproducible results might be obtained standard solutions were used.3 All solutions were filtered through sinteredglass filters. Apparatus.-The reaction vessels were 250-ml. Pyrex gas-washing bottles with fritted discs, modified t o permit withdrawal of samples by pipet through the gas-outlet tube. X separate bottle of the same type was used to remove 0 2 from the K2S203 solution prior t o mixing. NZ(or COZin a few experiments) saturated with water vapor was passed through flowmeters t o the gas-washing bottles. Tygon tubing was used for connections. The reaction vessels were maintained a t a constant temperature in a water thermostat. Light was not excluded, as preliminary expts. showed that the results were not significantly different when the expts. were conducted in the dark. Measurements of PH were made with a Beckman model G pH meter on samples a t room temperature. The absorption spectra were measured a t 25" with a Beckman model DU spectrophotometer using 10-cm. cells. The reaction rate and the reproducibility of the results depended t o some extent upon the method used t o clean the reaction vessels. In each case the vessels were dried after washing with one of these sequences of reagents: (1) Soap, concentrated HCl, distilled water ( 7 times) and redistilled water. (2) Hot concentrated aqua regia, distilled water (12-15 times) and redistilled water ( 3 times). (3) Redistilled water ( 3 times). ( 4 ) A 10% solution of NaOH in redistilled water and redistilled water (6-S times). (*4fter about 5 washings the wash water was colorless t o phenolphthalein indicator.) The effect of coating the glass surface with Beckman Desicote also was tested. Procedures.-Appropriate volumes of the standard solutions, except K?S208, were pipetted into the reaction vessel, which was then placed in the thermostat. The K~SZOS solution was kept a t room temperature, and NZwas bubbled through each solution a t 350 ml./min. After 40 min. the K2SlOp bottle was placed in the thermostat and 20 min. later the without disrupting the flow of Nz, K~SZOS solution having reached the thermostat temperature, 25 ml. were pipetted into the reaction vessel. The flow of NZwas maintained throughout the experiment. In the experiments with allyl acetate the solutions were deaerated by bubbling NZthrough them for 30 min. at 61.8'. They were then transferred t o glass-stoppered Erlenmeyer flasks t o prevent evaporation. Analyses.-Aliquot portions were withdrawn a t suitable intervals and analyzed for S 2 0 8 - by a slight modification of the method adopted p r e v i ~ u s l y . ~It should be noted that quenching of the oxalate reaction was accomplished by the much faster reaction of SZOS-with iodide ion and not by precipitation of CUI as reported earlier.8 (Copper(II), when complexed with oxalate, does not react with iodide.) Calculations based on the results of King and co-workers6 ( 5 ) R. Ballentine. Anal. Chcm., 26, 549 (1954). (6) E.Jette and C. V. King, J . Am. Chem. Soc., 61, 1034 (1929); C.V. King and E. Jette, ibid., 61, 1048 (1929); C.V. King and M. B. Jarohs, i b i d . , 53. 1704 (1931).
4353
have shown that, under the conditions employed, the peroxodisulfate-iodide reaction was a t least one hundred times faster than the reaction under investigation.
Results Stoichiometry.-To verify the stoichiometry of the reaction, a solution containing 0.00830 14 K2S208,0.01000 M Na2C204 and 2.0 X loF4 M CuSO4 was allowed to react for 41 hr. a t 40.2O under NS. It was then analyzed, and the final concentrations were 1 X M S20~' and 1.74 X low3 M C204-. The decrease in CrOd concentration (0.008% M ) was essentially identical with the decrease in S L O p concentration (0.00829 M ) , in accordance with the equation for the same reaction catalyzed by silver ion3,? SzOa'
+ CnOa'
=
2S04'
+ 2c02
Peroxodisulfate Dependence.-Durin g an experiment the logarithm of the SnOe' concentration decreased linearly with time, showing that the rate law for the reaction is first-order in peroxodisulfate concentration
- d(SzOa")/dl
= K'(Si0s')
Values of k' were obtained from graphs of the logarithm of the volume of thiosulfate solution (which is proportional to the concentration of S 2 0 8 ' ) versus time. As the graphs were quite straight it was possible to obtain an accurate value of k' for each experiment. However, the slopes for different experiments under the same conditions varied somewhat. (Reproducibility is discussed in detail in the section on surface effects.) The first-order dependence was also verified by changing the initial concentration of peroxodisulfate. The results are shown in Table I. TABLE I PEROXODISULFATE DEPENDENCE' (KzS?Oa), Ai'
I O 4 k', sec.-l
0 . 0 ~ 8 ~ 3.07 . 0083b 3.42 .O1lBb 3.43 0C46" 3.03 .0083" 3.73 .0332c 3.77 A f CuS04. O40.47', 0.040 iM Na2C204 and 4.0 X Reaction vessel cleaned by method 2. c Method 4.
.
Oxalate Dependence.-The rate is essentially independent of oxalate concentration (Table 11). When Na2S04was added to maintain constant ionic strength, variation of the oxalate concentration over a four-fold range did not produce any significant change in rate. When the ionic strength was allowed to vary, the rate tended to increase slightly with increasing oxalate concentration. Similar results were obtained in other experiments with (Na2C204) from 0.008 to 0.16 M , p from 0.05 to 0.51 and cleaning by method 2. (The p H of the experiments of this section varied from 5.1 to 6.5. It will be shown that this had no significant effect on the rate.) Copper Dependence.-To determine the order of the reaction with respect to the catalyst, (CuS04) was varied from 1.6 X to 40.0 X M using Nz and from 0.8 X to 50.0 X M (7) C.V. King, ibid., SO, 2059 (1928).
EPHRAIM BEN-ZVIAND THOMAS L. ALLEN
4354
TABLE I1 OXALATEDEPENDESCE~ (NanCnOa), .VI
fib
0.04
0.55 .55 .55
.08
l o 4
K ' , sec. -1
3.08 3.72 3.13 3.63 3.73 3.87 4.75
.16 .09 .02 .04 .l5 .08 .27 .12 .39 .54 4.30 .17 ~ 4 0 . 4 7 " ,0.0083 iV1 K&O8, 4.0 X J4 CuSO4 and cleaning by method 4. 6 Ionic strength adjusted with NazSOa in first 3 experiments.
using C02. (The other concentrations were 0.0083 llf KZS208and 0.040 JW NaiC204, the temperature was 40.4' and the reaction vessels were cleaned by method 1.) In Fig. 1 K' is plotted verszhs the
,
I
30
/
I
5
square root of the CuSO4 concentration. Each point represents the least-squares method average of duplicate or triplicate experiments. The best straight line determined by least squares for the experiments under N2 is shown; it fits the results with COz almost equally well. Therefore the rate is independent of the concentration of CO, (a reaction product), and it is proportional to the square root of the catalyst concentration k' = k(Cu)+/n
where (Cu), designates the total concentration of the various species of copper in solution. Surface Effects.-Table I11 summarizes the results of experiments designed to test the effect of TABLE I11 SURFACE EFFECTS~ Added substance
three experiments powdered BaSOa was added to the reaction solution. Although these experiments indicate that the extent of the surface has no appreciable effect, the method of treating the reaction vessel surface did affect the rate and its reproducibility (Table IV). The highest rate and most reproducible results were obtained using method 4. It was adopted after the experiments with Desicote, and the values of l o 4 k' (sec.-l) in the first five consecutive experiments were 2 . 8 , 2.40, 2.83, 3.42 and 3.77, respectively. Following this period of conditioning the surface and/or removing the last traces of Desicote the results were fairly consistent, as iiidicated in Table IV. TABLE 11. EFFECT OF Surface treatment
Method 1 Method 2 Method 3 Method 4 Desicote a 40.47', 0.0083 10-4 M CUSO~.
!
10 15 20 25 103 (cUso,)l!*, ~ 2 . Fig, 1.-The effect of catalyst concentration on the first-order rate constant: 0, under iY2; 0 , under Cot. 0
104 k ' , set.-'
.... ...
2.17 Pyrex glass 2.18 0.025 g. BaS04 2.38 ,227 g. BaSOu 2.37 2.27 ,737 g. BaS04 a 40.2', 0.0083 M K?S208, 0.040 114 Na?C?04, 4.0 X lo-: M CuS04and cleaning by method 1.
surface area. I n the second experiment small Pyrex glass tubes were placed in the reaction vessel, doubling the surface-to-volume ratio. In the last
Vol. SB
SCRFACE T R E A T M E S T '
10' k ' ( a t - e . ) , sec. -1
of eupts.
Iii,.
13 11 15 10 4 R~S208,0.040
Xvr. d c v
,
0
2.58 10 3.42 7 2.97 8 3.73 6 1.80 26 AI Sa2CzOL and 4.0 X
Ionic Strength Dependence.-The effect of changing ionic strength was studied using NaC104 (Table V). In the range 0.13-0.39, ionic strength has no significant effect on the rate; a t higher ionic strength k' is somewhat smaller. Similar results were obtained in other experiments using from 0.04 to 0.32 -ld SaLS04. The experiments in which ( Na2C204) was varied (see above) provide evidence that the rate is essentially independent of ionic strength in the range 0.05-0.54. TABLE V IOSICSTRESGTH DEPENDENCE* 104 ii'
(XaClOi),
.I I
...
@
0.15 .17 ,21 .%7 .30
isec.
-1)
3.42" 3.40 3.10 3.13 3.25
0.021 ,061 ,12 ,24 -2 . 75 . GO . IS 1.20 1.35 2.10 40.17', 0.0083 >lf KPSQOS,0.040 Jf Sa?CzOa, 4.0 X 1 0 . 2 5 , av. JI CuSOl and clcnning by method 2. of 14 experiments. '2
Sulfate Dependence.--Xs sulfate ion is a reaction product its concentration increases during an experiment. The lack of any significant deviation from first-order kinetics therefore indicates that the reaction rate is independent of sulfate concentration. (However, this is not true of the experiments a t low fiH discussed below.) Further evidence is provided by the first three experiments of Table I1 and by the experiments on ionic strength dependence with NaZSO4 mentioned above. In these experiments anhydrous Na2S04 was used. With the decahydrate the rate decreased with increasing sulfate concentration. Although the effect was not large, the presence of an inhibitor in the analytical reagent grade decahydrate is indicated.
Nov. 5, 1961
COPPER
CATA4LYSISIN
PEROXODISULFATE
OXIDATION O F OXALATE
pH Dependence.-The effect of PH has been investigated by the addition of NaOH or KHSO4 and by replacing part or all of the NazCz04with NaHCz04. During each experiment the pH of the solution remained constant. The results in Table VI show that the rate is essentially independent of pH in the range 3.85-6.32. At higher $H the rate decreases, probably because of precipitation of the catalyst. At low PH (2.52-3.60) the semi-logarithmic graphs are non-linear. Although the effect on the initial rate is not very pronounced, the rate decreases rather markedly as the reaction proceeds.
4355
lou4 M CuSO4; the other contained no catalyst. I n both expts. the first-order rate constant was 9.0 X sec.-l. For comparison, the rate constant for the thermal decomposition of peroxodisulfate in 0.1 M NaOH a t this temperature, calculated from the data of Kolthoff and Miller,g is 6.7 X see.-'. Influence of Other Substances.-Various other substances were added to the reaction solution, usually a t a concentration of 1.0 X M. (The temperature was 40.2', and the other concentrations were 0.0083 iM KzSlOg, 0.040 M Na:CS04 and usually 4.0 X M CuSO4.) Silver nitrate increased the rate markedly. SubTABLE VI stances which inhibited the catalyzed reaction to fiH DEPENDENCE' varying degrees included Mn++, MnOd-, NHdOH, flH 104 k', sec. - 1 S,Os' and C1-. Dust was also found to be an 7.08 2.68 effective inhibitor. Substances with no appre6.60 2.92 ciable effect on the rate included NH4+, Zn++, 4.07 6.32 Mg++, Sn++, Pb++, Fe++, Fe+++, AI+++, HzOz, 6.00 3.73* SOs=, OAc-, methanol, 1-propanol, 1-butanol 5.43 3.30 and tert-butyl alcohol. Variation of the NI flow 4.32 3.72 rate over an eight-fold range also had no effect. 3.85 3.75 Absorption Spectra.-The absorption spectra of a 40.47", 0.0083 M K&Os, 4.0 X M CuS04, 0.040 4.0 X M (NaSC204 NaHC204) and cleaning by method 4. f solutions containing 0.040 M Na2C2O4, M CuSO4 and either with or without 0.0083 M 0.22, av. of 10 expts.; PH i= 0.20. K2S2Oswere measured over the range 600-850 mp. Temperature Dependence.-The variation of The presence of peroxodisulfate had no effect on rate constant with temperature is shown in Table the spectrum. There was a broad maximum a t VII. When log k' was plotted versus 1/T the 710-720 mp with a molar absorptivity (based on experimental points fell along a straight line. From copper concentration) of 35. the slope of the line the activation energy was found Cuprous Oxalate.-To determine whether copto be 32.2 kcal./mole. per(1) could exist in an oxalate solution, a sample of cuprous chloride was dissolved in a sodium TABLE VI1 oxalate solution. As it dissolved, i t disproporTEMPERATURE DEPENDENCE^ tionated, forming dioxalatocuprate(I1) ion and 104 k', sec. Temp., '(2. metallic copper. 25.64 0.322 Comparison with Earlier Results.-In the first 32.70 1.02 article of this series3 the copper-catalyzed reaction 40.47 3.73 was studied a t 69.7' using very low concentrations 6.83 44.30 a0.0083 M K&Os, 0.040 di Ka&2204, 4.0 X M of catalyst. Oxygen was not excluded. To about 50% reaction the rate was zero-order in both CuSO4 and cleaning by method 4. SZOS' and C204=. Dependence on CuSO4 conOxygen Inhibition.-In some preliminary ex- centration was linear with a relatively large conperiments i t was found that the rate increased stant term. Concentrations were not varied when: (a) air was removed with oil-pumped over a very wide range, and reconsideration of the nitrogen; (b) prepurified nitrogen was used instead; data shows that they fit the rate law established (c) the reaction vessels were modified so that in the present work fairly well. The large constant samples could be obtained without removing the term had been interpreted as being due to a copper stoppers. However, the presence of air did not impurity of 1 X lo-' Id,but this hypothesis conchange the form of the rate law. In experiments flicted with analyses showing that the impurity using 4.0 X M CuSO4, 104k' was 0.65 sec.-l was only about 2 X M . When the data are a t 40.47' when OS was bubbled through the solu- reinterpreted using the new rate law, this paradox tion, compared to 3.73 see.-' under Nz and other- disappears. A comparison of the reaction rates wise similar conditions. Tests of the reaction obtained in the earlier work with values calculated solutions with titanium sulfate showed that the using the results of the present investigation shows presence of oxygen did not lead to the formation that the two sets of data are in fairly good agreeof hydrogen peroxide. ment. Allyl Acetate Inhibition.-Allyl acetate is an Discussion extremely effective inhibitor for this reaction. There are several reasons for believing that the Two expts. were carried out a t 61.8' and a p H of 4.7, using 0.0083 M K&Og, 0.040 M Na:Cz04, copper-catalyzed reaction of peroxodisulfate and 0.230 M allyl acetate and 0.005 M KHS04. (The oxalate is a free-radical chain reaction : (1) The very marked inhibition by allyl acetate KHSO4 was added to prevent hydrolysis of the allyl acetate.*) One solution contained 4.0 X is indicative of a chain reaction involving SOafree radicals. It has been shown by Kolthoff, ( 8 ) I. M. Kolthoff, E. J. Meehaa and E. M. Carr, J.A m . Chem. SOL.,
+
-1
7 6 , 1439 (1953).
((1)
I. hl. Kolthoff and I. K. Miller, ibid.,73, 3055 (19.51).
EPHRAIMBEN-ZVIAND THOMAS L. ALLEN
4356
Meehan and Carr that allyl acetate is a highly efficient captor of sulfate radicals.s Furthermore, the first-order rate constant in the experiments with allyl acetate was essentially the same as that of the thermal decomposition of S20s=in 0.1 NaOH, where the reaction Sz08- + 2S04- is the rate-determining step.g Therefore neither oxalate nor copper participates in the chain-initiating reaction, which is the homolytic cleavage of Sz08-. (2) Inhibition by oxygen indicates that this is a chain reaction in which COz- (or CZOd-) radicals participate. Oxygen is known to inhibit certain chain reactions involving COS- radicals, such as the oxidation of oxalate by perrnanganate1O-lz or mercuric chloride. l 3 - I 5 (3) The difficulty of obtaining reproducible results and the sensitivity of the rate to the treatment of the reaction vessel surface indicate that this is a chain reaction with very long chains. Other studies of free-radical chain reactions in aqueous solution in which similar problems have been encountered include the gamma-ray induced decomposition of hydrogen peroxidelGand the peroxodisulfate induced exchange of oxygen atoms between water and molecular oxygen.1 7 , l 8 A reasonable mechanism of this type which leads to the experimental rate law is ki
SZOS'
+2s04-
(1)
k?
+ CU111(C204)nCu"'(C?O*)?- 3_ CU"C*O4 + cot + coika CUTTCz04+ CUrT(c2O4)~ks COe- -1 +coz + Sod- + sodke Con- + sodco:!+ SOa'
so,- f CUrT(C?Oa)2' --f SO4-
(2)
k
c204-
s208'-
--f
(3) (4)
(5) (6)
Using the usual steady-state approximation and the assumption that 4k2k~(CuTr(Cz04)2=) >> klkc, it may be shown that -d(SzO,-)/dt
terpositive copper or other copper-containing reaction intermediate is relatively small. Therefore to a good approximation (Cu), = ( C U I I ( C ~ O ~ ) ~ = ) , and the experimental rate constant k = (klk2. kj/kg) '12. At 40.47' K I is 2.0 X IO-' set.-' as calculated from the data of Kolthoff and Miller.'J In this investigation k' = (klkzkb/k6) '/2(Cu1I(C204)z=)1/2varied from about 50 to over 2000 times as large as kl a t the same temperature, depending on the catalyst concentration and other conditions. Therefore the average chain length varied from about 50 to over 2000. Also, it follows that ( 2 k ' / k l ) ? = 4 k ~ K 5 ( C u r 1 ( C ~ O ~ ) ~ - 3 ) / k lo4, lk~ showing that the results are consistent with the inequality assumed in the derivation. Oxidation of copper to the f3 state is postulated in the mechanism. Compounds of terpositive copper have been prepared using several powerful oxidizing agents, including peroxodisulfate, in alkaline solution.20,21 In several investigations it has been postulated that Cu+++ is a reaction intermediate, formed by the reaction of Cu++ with atomic chlorine.22~23 Alternatively, reactions 2 and 3 may be combined in a single concerted step without affecting the form of the derived rate law. This change avoids the copper(II1) complex, but the function of copper in the mechanism is then less distinctive. Another alternative mechanism may be obtained by combining reactions 3 and 4. As in several other oxalate reactions, substitution of ( 2 2 0 4 for COz- has no effect on the rate law. Evidence concerning the relative merits of the two species is summarized by Adler and Noyes.12 The reactions between anions should be hindered by coulomb repulsions. il mechanism involving OH radicals avoids all but one of these reactions. In this alternative mechanism so4- radicals generate OH radicals, which then substitute for sod-. Besides reactions 1, 3, 4 and 5, it includes SO,-
= (kik2kS/ke)'/z(S208=)(cUrr( C204),-)'/2
This rate law has the same form as the experimental rate law, provided that the concentrations of other species of copper in solution are negligible with respect to that of dioxalatocuprate(I1) ion. Calculations based on the equilibrium constants determined by Watters19 have shown that in a typical experiment dioxalatocuprate(I1) ion comprised about 99.7% of the total copper(I1) in solution, assuming equilibrium between the various copper(I1) species. Since the absorption spectrum of a solution containing dioxalatocuprate(I1) was found to be unaffected by the presence of SzOs-, it is likely that the steady-state concentration of (10) H. F. Launer, J . A m . Citem. SOL.,6 4 , 2397 (1032); 65, SG3 (1933). (11) II.F. Launer and D. hi. Y o s t , i b i d . , 56, 2571 (1934). (12) S. J. Adler and R. M. N o y e s , i b i d . , 7 7 , 20:3D (1955). (18) W.E . Roseveare and A. R. Olson, i b i d . , 61, 171G (1929). ( I - % ) W. E. Roseveare, i b i d . , 62, 2G12 (1930). ( 1 . 5 ) R. A . Hausman and T . W.Davis, i b i d , , 7 6 , 5341 (19.5,). (11;) E, J . H a r t and hi. S. hfatheson, Disr!cssioits Pavatiiry Soc., 12, 1G9 (1952). (17) H. Taube, A n n . R P WS. ? ~ r l e aSi r i . , 6 , 2 8 (19X). (18) C. R. Giuliano, N. Schw-artz and W. K. Wilmarth, J . Pizys. Chrm.. 63, 3 3 (19.79). ( 1 9 ) J . I. \\'attrrs. J . A n . Chmi. .Tor , 81, I X O (19:9).
VOl. 83
OH
k, + 1320 + SO,' + H+ + OH
+ Cu"(C20c)n'
ks
--+ OH-
Hf + O H - + OH
kio
+ CuTT'(C~0o()?-
(7) (8)
ks H20
+ COz- --+ OH- + CO2
(9) (10)
The derived rate law is the same as the one obtained previously, except that k p and kg are replaced by K8 and &, respectively. A more complex mechanism involving dioxalatocuprate(1) ion also leads to the experimental rate law. However, when cuprous chloride is dissolved in a sodium oxalate solution it disproportionates, as described above. The stabilization of the +2 oxidation state of copper with respect to the +1 state occurs with other chelating agents, such as ethylenediamine. 2 0 , 2 4 Oxalate ( 2 0 ) J . Kleinber~."UnfamiliarOxidation States and Their Stabilization," University of Kansas Press, Lawrence, Kansas, 1850. (21) I
