Reaction between Oxalate Ion and Peroxodisulfate
1265
iron(II1) on the calomel formation in the light is accounted for by the photodecomposition of oxalatoferrates which has been studied in detail by Cooper and DeGraffG Fe(II1)Ox
+
hv
-
Fe(I1)
+
COz
+ COz*-
(16)
the C02.- free radical initiating the chain reaction. In the additional presence of hydrogen peroxide the iron(I1) formed reacts as Fe(I1) + HzOz+ Ox Fe(II1)Ox + OH- + OH. (measurable rate)
-
-
the OH- reacting with C Z O ~ - ~
+
OH- + COz + COz*- (fast) (17) followed by reactions 3-5. Under our experimental conditions all hydrogen peroxide had disappeared at the end of the reaction period and the number of moles disappeared was considerably smaller than the number of moles of calomel formed, which is accounted for by reactions 16, 3, 4, and 5. From the experimental results it is clear that the rate of the photodecomposition of oxalatoferrate in the presence of peroxide determines to a great extent the amount of calomel formed. The reactions involving peroxide, iron(I1) and (111), oxalate, and oxygen in the dark and light have been discussed in a previous paper.4 The retarding effect of oxygen is easily accounted for by the reaction OH.
C204?-
The OzCOz- reacts further by a chain mechanism which yields hydrogen peroxide.2 Apparently the reaction of CO2.- with 0 2 is much faster than that with mercuric chloride (eq 3). Actually hydrogen peroxide formation in an oxygen atmosphere under conditions described previously213is not affected by the presence of mercuric chloride. The retarding effect by the chloride ion is accounted for by several reactions. Mercuric chloride with chloride ion forms complexes HgC13- and HgCl2- which may react with C02.- with a slower rate than HgC12 (eq 3). Also the rate of reaction 5 may be decreased by the complex formation. Furthermore chloride promotes the reaction HgZClz(solid) + C1- * + HgC1,which has bqen studied by Sillen, et a2.7 The solubility of calomel is increased by chloride. In a reaction mixture such as in Table I, but which was 2 M in potassium chloride the solubility a t 50" was 1 X 10-6 M . Probably the main cause of the retarding effect of chloride is that the stability of Hg(1) formed in reaction 3 greatly decreases 2Hg'
t Hgzz+
since [Hgz2+] becomes extremely small in the presence of excess chloride ion.8 (6) G. D. Cooper and B. A. DeGraff, J. Phys. Chem., 75,2897 (1971). (7) B. Lindgren, A. Jenssen, and L. G. Siiien, Acta Chem. Scand., 1, 479 (1947). (8) i. M. Kolthoff and C P. Barnum, J. Amer. Chern. SOC., 62, 3061 (1940).
Catalyzed Reaction between Oxalate Ion and Peroxodisulfate. I. Copper( I I) as Catalyst Masaru Kimura Department of Applied Chemistry, Yamagata University, Yonezawa 992 Japan (Received October 27, 1972)
According to the composition of the reaction mixture, two separate mechanisms of the reaction are presented to account for the copper-catalyzed reaction between oxalate and peroxodisulfate. (1)In an excess of oxalate the catalyzed reaction kinetics are half order in catalyst, first order in peroxodisulfate, and zero order in the oxalate concentration. The reaction is completely inhibited by the presence of 1%acrylonitrile, and remarkably inhibited by the molecular oxygen. (2) In a low concentration of oxalic acid, in the relatively high concentrations of catalyst, and in 0.1 M sulfuric acid, the catalyzed reaction is almost zero order in the catalyst, first order in peroxodisulfate, and half order in the oxalic acid concentration. The reaction is not subject to the inhibition by the molecular oxygen but is completely inhibited by 1% acrylonitrile. Iron(1II) and manganese(I1) are considerably strong inhibitors of the reaction, and the former is much more effective inhibitor than the latter. The effects of the various factors on the copper-catalyzed reaction between oxalate and peroxodisulfate are reported. Mechanisms and kinetic treatments are presented to account for these facts.
Introduction Recently, cobalt(I1) catalysis on the reaction between oxalic acid and peroxodisulfate was investigatedl. The cobalt-catalyzed reaction was not subject to inhibition by molecular oxygen, but was c o m p ~ e t e ~inhibited y by 1% acrylonitrile or 1%allyl acetate. Ben-Zvi and Allen2 have
studied the copper-catalyzed reaction of oxalate with peroxodisulfate and found that molecular oxygen is a strong inhibitor Of the catalyzed reaction. However, only one de(1) M. Kimura, J. Phys. Chern., submitted for Publication.
(2) (a) T. L. Allen, J. Amer Chem. Soc., 73, 3589 (1951); (b) E. BenZvi and T.L. Ailen, !bid, 83, 4352 (1961) The Journal of Physical Chemistry, Vol. 77, No. 10, 1973
Masaru Kimura
1266
termination in the oxygen-saturated and in the nitrogensaturated solution has been compared. No quantitative investigation of the oxygen inhibiting effect has been made. The effect of oxygen on the free radical reactions is quite general and is an important aspect. An intensive investigation of the kinetics of the reaction between oxalate and peroxodisulfate is made in the present paper. Experimental Section Chemicals. Reagent grade potassium peroxodisulfate of Kanto Kagaku Co., Inc. was recrystallized twice from the redistilled water and dried a t 40" in a vacuum desiccator. The redistilled water was prepared from the anion-cation exchange resin water first by successive distillation from the dilute alkaline permanganate solution and finally without adding any reagent in a glass still. The redistilled water was used for all the experiments and also to wash the glassware. The reaction vessel was always filled with the redistilled water for a t least one day before use. All other chemicals used were of the guaranteed reagents and were used without further purification. Technique. The reaction vessel was placed in a thermostat and oxygen and/or nitrogen gas was bubbled through the contents (90 ml) before the reaction was started. A reaction was started by adding peroxodisulfate (10 ml) which was also in a thermostat and had oxygen and/or nitrogen gas bubbled through it. The concentration of the dissolved oxygen gas was continually checked and controlled to fairly good constancy by bubbling nitrogen and/ or oxygen gas. The whole vessel was covered with black adhesive plastic tape to ensure darkness. The aliquot samples were withdrawn a t appropriate times and mixed with a cold solution of 0.5% acrylonitrile to stop the reaction. Then the concentration of peroxodisulfate was measured. A total time of about 5 min was needed from the withdrawal of the solution until completion of the analysis of peroxodisulfate. Analysis. The concentration of peroxodisulfate was determined by using polarograms a t 0.2 V us. sce at 25" in a solution of 0.5% acrylonitrile, 0.01% polyacrylamide. and 0.05 M sulfuric acid. Polyacrylamide was used as a maximum suppressor. Concentration of oxygen in the reaction mixture was determined by using a Dissolved Oxygen Meter, Model DO-IA of Toa Electronics Co., Ltd. Results. Experiments in Excess of Oxalate Decomposition of Peroxodisulfate without Catalyst or Oxalate. In the same conditions as in the experiments of the copper-catalyzed reaction, the rate of the decomposition of peroxodisulfate was checked. In the initial solution of 0.004 M potassium peroxodisulfate and 0.1 M potassium oxalate, the decomposition of peroxodisulfate was only 2% in 10 hr at 40" in the dark, in either oxygen or nitrogen saturated solution. In the same condition as in the above except 0.004 M cupric sulfate instead of 0.1 M potassium oxalate, the decomposition was also 2% in 10 hr. The 2% was practically equal to the decomposition of peroxodisulfate without any other chemical reagents. Kolthoff and Miller3 have determined the rate constants of the decomposition of peroxodisulfate. The calculated value from their data is k1 = 1.2 X 10-5 min-1 a t 40°, and the decomposition of about l%in 10 hr is calculated. Catalyzed Reaction by Copper(II). The reaction between oxalate ion and peroxodisulfate is greatly enhanced by the presence of copper(I1). A few examples are shown in Figure 1. As can be seen in Figure 1, the catalyzed The Journal of Physicai Chemistry, Vol. 77, No. 10, 1973
I
I
I
I
I
I
I
I
I
I
I
1
2
3
4
5
Reaction time in hours
Figure 1. Examples of kinetics r u n at 40". Initial solution: 4 X M in potassium peroxodisulfate: 0.1 M in potassium oxaM in cupric sulfate; pH 6.5; dark; ( l ) , N2 satulate; 4 X rated: (2), air saturated; (3),02 saturated: ( 4 ) , as in (1) but in 1% acrylonitrile. reaction is completely inhibited by 1% acrylonitrile, and the about 1%decomposition in 6 hr is actually equal to the uncatalyzed decomposition of peroxodisulfate. The reaction is considerably inhibited by the molecular oxygen; however, even in the oxygen saturated solution the reaction proceeds to a large extent. Dependence of Peroxodisulfate. During an experiment the logarithm of the persulfate concentration decreased linearly with time. This indicates that the kinetics of the reaction are first, order in persulfate and are described by the equation -d[S,O,'-]/dt
=
k,,p[Sz082-]
where kapp indicates an apparent rate constant to be a function of the copper concentration and the oxygen concentration. The above rate law was fulfilled in both the presence and absence of oxygen. Dependence of Copper(II). To determine the order of reaction with respect to catalyst, total concentration of copper(II) was varied from 1 X 10-6 to 4 X 10-3 M . The values of kapp were plotted against the square root of the copper concentrations, As can be seen in Figure 2 the plot is on a straight line and the intercept is practically zero. This indicates that the reaction rate is proportional to the square root of the catalyst concentrations. [Cu(II)] in Figure 2 indicates the total concentration of various forms of copper(II) with oxalate. In excess of oxalate they would be mostly the dioxalatocuprate(I1) complex. Dependence of the Oxygen Concentration. The catalyzed reaction rate was determined in a reaction mixture (3) I. M. Koithoff and i. K. Miller, J. Amer. Chem. Soc., 73, 3055 (1951);the kinetics are described by the equation -d[S208*-]/dt = kt[S208'-] k2[H*][S2082-]. The values of k , and k ~ respec, min-' and 3.5 X M - ' min-' at 50' at tively, are 6.0 X an Ionic strength of 0.4. The activation energy for the k , path is 33.5 kcai and for the k 2 path 26.0 kcai.
+
Reaction between Oxalate Ion and Peroxodisulfate
I
I
2
-‘ E
1.5
N
2 X
a n Y
1
I
1267 I
I
TABLE II: Effect of Oxalate Concentration on the Catalyzed Reactiona kapp x
Concn of oxalate, M
0.001 0.002 0.01 0.05 0.1
N2 sat.
02 sat.
4.0 6.3 11 23 23.24
3.5 5.0 8.3 3.3 2.0
a Conditions as In Table I except for the varied concentrations of potassium oxalate.
TABLE Ill: Effect of Sulfuric Acid Concentrationa Concentration of sulfuric acid added,
$
g
M
0 (pH 6.5) 0.01 (pH 4.2) 0.02 ( p H 3.7) 0.05 (pH 2.2) 0.1 0.2
0.5
a
kapp
x io3. min-’ O2sat.
N2 sat.
23,24 21 22 20 23 23
2 1.3 3 7 12
I
0
1
2 [CU”]’”
3
5
4
x IO2, Mtfl
Figure 2. Plot of kapp vs. [CU”]”~. Conditions as in Figure 1 , except varied concentrations of copper(l1): a,,0 2 saturated; 0, Nz saturated. TABLE I: Effect of Oxygen on the Reaction Ratea Rate constant,
kapp x 103,min-1 Concn of oxygen,
M
Obsd
0 (N2 sat.)
3x 6X 9x 1.9x 3.7 x 5.3 x 6.2 x 9.3 x 1.0 x 1.2 x 2.8 x 9.6X
io3, min-’
10-6 10-6 10-5
10-5 10-5
10-5 10-5 10-4 10-4
10-4 (02sat.)
23.0,23.7 13.3 12.3 8.8 6.8 5.8 5.3 5.0 5.0 4.3 4.0 3.3 2.0
Calcd
22.2 14.8 11.9 10.3 7.6 5.8
5.0 4.7 4.0 3.9 3.7 2.9 2.3
”
Solution 0.004 M in potassium peroxodisulfate, 0.1 M in potassium M in cupric sulfate: 40”; dark. oxalate, and 4 X
of the varied concentrations of the dissolved oxygen. The values of happ obtained are listed in Table I together with the calculated values which are discussed later. Under conditions similar to those in Table I, Ben-Zvi and Allen2b have reported k a p p to be 22.4 X 10-3 min-1 in the nitrogen- and 3.9 X min-1 in the oxygen-saturated solution at 40.5’. These values are in good agreement with the corresponding data in Table I. Ben-Zvi and Allen have taken special efforts to clean the reaction vessel and also have taken special precautions for surface effects. E f f e c t of the Oxalate Concentration. The effect of the oxalate concentration on the rate was checked over a wide
“Condition as in Table I except for the varied concentrations of sulfuric acid.
range of concentrations. (See Table II.) The rate appears to be independent of the oxalate concentration in the range 0.05 to 0.1 M . However, the rate decreases markedly a t concentrations lower than 0.05 M. It is interesting to note that the inhibiting effect of oxygen becomes extremely small when the oxalate concentrations are lower than 0.01 M . Effect of the Concentration of Sulfuric Acid. The effect of the sulfuric acid concentration was not large in the nitrogen saturated solution but it was rather large in the oxygen-saturated solutions. (See Table 111.) Experiments in Low Concentration of Oxalic Acid and Relatively High Concentration of Catalyst in 0.1 M Sulfuric Acid Solution In this section the experiments are mostly carried out in the solutions of 0.002 or 0.02 M copper(II), 0.002 or 0.005 M oxalic acid, and 0.1 M sulfuric acid. The major reason for this composition is for purpose of solving the fact that the inhibiting effect of oxygen on the copper catalyzed reaction almost disappears at very low concentrhtions of oxalate. Kinetics. The dependence of the peroxodisulfate concentration on the rate was essentially the same as in the preceding section. However, the plot of logarithms of the persulfate concentration with time deviates from a straight line after about t 1 p . This is why the reaction rate depends on the oxalic acid concentration. Accordingly, the rate constant ( k a p p ) was determined from the rate of the initial stage of reaction. Stoichiometry. To verify the stoichiometry of the reaction a kinetic run was followed up to the final stage of the reaction. Some examples are given in Figure 3. As can be seen in Figure 3, the molar concentration of oxalic acid reacted is equal to that of peroxodisulfate, so the overall reaction must be s20,2-
+ c,o,2-
-
2s042-
+
2c0,
The Journai of Physical Chemistry, Voi. 77, No. 10, 1973
1268
Masaru Kimura
I
I
0
2
I
I
I
I
I
I
I
I
I
4 6 8 io Reaction t i m e i n hours
Figure 3. Examples of kinetics run in low concentrations of oxM in potassium alic acid at 40". Initial solution: ( l ) , 4 X peroxodisulfate, 0.02 M in cupric sulfate, and 0.1 M in sulfuric acid; N2 saturated, and in the dark; (2), conditions as in (1) exM in oxalic acid; (3), conditions as in cept to contain 1 X (1) except to contain 2 X M in oxalic acid; (4), conditions M in peroxodisulfate instead of 4 X as in (3) except 3 X 10-3 M.
Dependence of the Oxalic Acid Concentration. In order to know the reaction order with respect to the oxalic acid concentration, the reaction rate was measured at varied concentrations of oxalic acid from 1 x 10-3 to 1 x 10-2 M. The values obtained are plotted us. the square root of the oxalic acid concentration in Figure 4. The plot of kapp us. [H2C204]1/2 showed a straight line, and the intercept was not zero. The value of intercept corresponds to the rate constant of decomposition of peroxodisulfate. Dependence of Copper(I0. The dependence of the catalyst concentration was examined in the wide range of the copper(II) concentration, from 2 x 10-4 to 2 x 10-2 M.It appears that the reaction rate is essentially independent of the copper concentration (Table IV). Effect of Oxygen. The catalyzed reaction rate was measured in the nitrogen-saturated and the oxygen-saturated solution. As can be seen in Table V, the rate was practically identical in both reaction atmospheres. Effect of Other Substances. The effects of a few substances on the catalyzed reaction were checked and listed in Table V together with the oxygen effect. The reaction was completely inhibited by the presence of 0.5%'acrylonitrile and considerably inhibited by iron(1II) and manganese(I1). The inhibiting effect of iron(II1) was over 10 times that of manganese(I1). Copper-Catalyzed Reaction between Hydrogen Peroxide and Peroxodisulfate. In the same composition of solution except that hydrogen peroxide is placed instead of oxalic acid, the copper-catalyzed reaction was measured. The results are given in Table V together with those in the oxalic acid reaction scheme. As can be seen in the table, there is essentially no difference in the reaction rate beThe Journal of Physical Chemistry, Vol. 77, No. 10, 1973
Figure 4. Plot of k,, vs. square root of oxalic acid concentration. Solution 0.004 fi in potassium peroxodisuifate, 0.02 M in cupric sulfate, 0.1 M in sulfuric acid, and varied concentrations of oxalic acid; 40";dark; nitrogen atmosphere. tween the hydrogen peroxide and the oxalic acid reaction scheme catalyzed by copper(I1).
Discussion In the excess amount of oxalate almost all of the copper(I1) should be dioxalatocuprate(I1) in the solution, calculated from the equilibrium constant^.^ However, the symbol CuT10x is used for all the forms of the oxalato complexes in this paper. Copper(II1) also would form one or more complexes with oxalate, and the symbol CuIIIOx is used for all the forms of the oxalato complexes. The following reaction mechanism is postulated to account for the facts under the conditions of Table I or Figure 1.
szo2-
+
CU"OX
cu"'0x
--
2s04- [very slow]
SO,-
-cu"ox
+ SO,,+ co, + coz-
CU"'OX
+ 0, OzCOz- [rapid] szo,2- + coz- -so,2+ so*- + co, S , O ~ - + O Z C ~ , --SO,'+ so4- + COz + COz-
i--)
(1) (2) (3) (4)
(5) 0 2
(6)
Reaction 1 is the initial reaction, and reactions 2 to 6 constitute a chain reaction. The most probable terminating reactions are CO,SO,+ CO, [fast] (7)
+
+
OzCOzSO,- -SO:+ COz + 0, tfastl (8) The radical which is formed in reaction 3 may be either (4) "Stability Constants of Metal-Ion Complexes." compiled by L. G . Sillen and A. E. Martel, The Chemical Society, Burlington House, London, W. 1,1964,~361.
Reaction between Oxalate Ion and Peroxodisulfate
1269
TABLE IV: Dependence of the Copper Concentrationa Concn of copper (I I),
38,
2
x 10-4 x 10-4 1 x 10-3 2 x 10-3 5 x 10-3 1 x 10-2 2 x 10-2
4.8
5
5.0 5.3 5.3 5.5
M
min-'
6.7 8.0
a Solution 0.004 M in potassium peroxodisuifate, 2 X M in oxalic acid, 0.1 M in sulfuric acid, and the varied concentrations of cupric SUIfate; 40"; nitrogen atmosphere; dark.
TABLE V: Effect of Oxygen and Other Substancesa
ihbstances added
Noneb None
N2 sat.
02 sat.
12.7b 10.8
14.0b 11.0
2 X 10-4MFe(lll) 2 X 1 0 - 3 M Fe(lll) 5X M Mn(ll) 2 X M Mn(ll) 0.5 % acrylonitrile
6.7
ca. 0 . 2
0.5 % acrylonitrile*
ca. 0 . P
1.7 4.2 3.1
ca. 0.2 ca. 0 . P
aSoiution 0.004 M in potassium peroxodisuifate, 0.005 M in oxalic acid, 0.002 M in cupric sulfate, 0.1 M In sulfuric acid, and substances added; 40"; dark. Condition as in a, except 0.005 M H202 instead Of H2C204.
COz- or Cz04-; however, the two forms do not appear to be distinguishable on the basis of kinetic or spectroscopic evidence.5 Following the consensus of the previous workers, COz- is written in the present paper. The assumption of steady states for sod-, COz-, OZCOZ-, and CuI1I0x leads to
where k , indicates the rate constant of reaction ( n ) ,and K is the equilibrium constant of reaction 4. If the chains are long, the second term on the left can be neglected. This leads to -d[Sz0,2-] / dt =
It is obvious from Figure 2 that the term of k l [ s Z o s 2 - ]is very small to be negligible, and that the chain must be long. Consequently, eq I11 is derived.
=
kapp[S2082-]
where, a = kaK/ks, b = k7/(klkZk5), c = k s K / ( k l k z k s ) . Plots of k,,, us. [ C U ( I I ) ] ~should /~ be on a straight line in which the slope corresponds to the values of [ ( l + a[Oz])/ ( b + c [ 0 ~ ] ) ] ~ /Using ~. the data of Table I, eq I11 was treated to minimize the sum of squares of the per cert de-
viations between the observed and calculated values.6 This leads to the following values and standard deviations for the parameters: a = (3.53 d= 1.18) X 103, b = (8.13 =k 1.39) X and c = (3.32 f 0.42) X 105. These parameters yield the calculated values which are listed in Table I. The root means square deviation between observed and calculated values was 11%. If it may be assumed that the value of k7 is almost equal to that of ks, K = 4 X lo5 M - l and ke,/ka = 9 x 10-3 are obtained. This may be the reason why oxygen retards greatly the reaction rate in which COzr free radical is involved. In the above reaction mechanism, it is postulated that sod- oxidizes Cu1I0x to CuIIIOx but does not oxidize CzOd2-. The oxidation of Cu(I1) and oxalate by SO4- is a competing reaction, and the former appears to be much faster than the latter. It is quite general that the complex formation leads to favorable conditions for the oxidation of metal ions. This may be presumably why the oxidation of Cu(I1) to Cu(II1) is favored over the reduction when the complex forming-oxalate exists in excess. At a low concentration of oxalate the reaction rate becomes considerably slow, and the inhibiting effect by oxygen becomes extremely small. This may indicate that the oxidation of copper to trivalent state becomes difficult, and that on the contrary the reduction to the univalent by COz- or OzCOz- occurs. There are several evidences that copper(II) is reduced by COz- and OZCOZ- radicals. For example, in an investigation of the oxidation of formic acid by X-rays in aqueous solutions including copper(II), the reaction AI has been shown7. c u 2 + + c0,- -cu+ + coz (-41) Kolthoff, et al.,s have recently studied on the effect of copper on the formation of hydrogen peroxide upon the oxidation of oxalic acid, and presumed the reactions Az and AB. Cu(I1) + OZCOZ- -CU(I) -t COZ 02 (A,)
+
Cu(I1) + HO, -Cu(I) -t H" + 0, (A3) According to Andersen and Hart,S and many subsequent workers, HCOz ( i e . , the protonated COz-) reacts with molecular oxygen HCOZ + 0, -HO, + COZ (4a) In order to examine the existence of the reactions A3 and 4a, hydrogen peroxide instead of oxalic acid was used in a copper-catalyzed reaction and compared with the reaction between oxalic acid and peroxodisulfate (see Table V). The reaction between hydrogen peroxide and peroxodisulfate was also enhanced by the presence of copper(I1) ion, and the reaction rate was comparable with that of oxalic acid with peroxodisulfate. Therefore reactions A3 and 4a would be expected. The copper-catalyzed reaction between hydrogen peroxide and peroxodisulfate may be accounted for by the following mechanism.
-
szo,2Sz0,2- + H"
2s0,SO, + HS0,-
(1) (la)
G. D. Cooper and B. A. deGraff, d. Phys. Chern., 75, 2897 (1971). The least-squares method is that described by R. H. Moore and R. K. Zeiger in Los Alamos Scientific Laboratory. ReDort No. LA-2367 . (1958rwith addenda. J. H. Baxendale and D. Smithies, 2. Phys. Chern., 7, 242 (1956). I. M. Kolthoff, E. J. Meehan, and M. Kimura, Talanta, 20, 81 (1973). A. R. Andersen and E. J. Hart, RadiationRes., 1 4 , 689 (1961)
The Journal of Physical Chemistry, Vol. 77, No. 70, 7973
Masaru Kimura
1270
-
+ SO4HOz + SO4'- + H+ H20, + SO, -HO, + SO,- + H+ Cu(I1) + HO, - Cu(1) + 02 + Hf ~ ' 0 : - + CU(I) CU(II) + SO,- + S O:H,02
(9) (9a) (-43)
-
(10) When the composition of reaction mixture was the low concentration of oxalic acid and the relatively high concentrations of copper, the catalyzed reaction did not appear to be inhibited by molecular oxygen, and the reaction rate was comparable with that between hydrogen peroxide and peroxodisulfate. These two facts would indicate that the reaction mechanism must be analogous to the mechanism of the reaction between hydrogen peroxide and peroxodisulfate. Consequently the following mechanism was presumed to account for the reaction under the conditions of low concentrations of oxalic acid, the relatively high concentrations of copper, and of an acidic medium.
s208'S2O8'-
+
C ~ O , ~+- SO,
C,O,~-
+
SO,-
-
2~0,-
-
+
H+ -SO,
-
SO,-
0,
O2CO,- + H+ -HOz C 0 2 - + 0, + H+-H02 Cu(I1) CU(II)
+
Cu(I1)
szo,2-+
CO,-
+
-
+ coz-
HS0,-
+
+ CO,- + o,co,-
-SO,'-
c0,- +
+
(1)
Cu(1)
(la)
CO,
(11)
C 0 2 (12) (4)
CO,
(44
+ C02 + co,
(-41)
+ COZ + 0, Cu(1) + 02 + H+ + so,- + so,,-
O,CO,- -CU(I)
(A,)
+
(A3)
HOZ
Cu(1) -Cu(II)
-
The most probable termination reaction will be CU(I) SO,CU(II) + SO,'-
+
The assumption of steady states for O ~ C O Z -HOz, , and Cu(1) leads to
sod, so4-,
(10)
(13) coz-,
The second term on the left will be negligible because the chain is long, so the overall rate is described as
+
where ko = k l k l a [ H + ] . Since the concentration of oxygen does not appear in eq V, the reaction rate should be independent of oxygen. Equation V is consistent with the experimental results. A plot of kapp us. [HzCz04]1/2 was a straight line as shown in Figure 4. The intercept should correspond with the rate
The Journal of Physical Chemistry, Vo/. 77, No. 10, 7973
of the uncatalyzed decomposition of peroxodisulfate. From the intercept, the value of ko = 2 x 10-4 min-1 was obtained. This is in good agreement with the value of 3.3 X min-1 which is calculated using the literature data3 for the conditions of [H+] = 0.2 M and at 40". As can be seen in Table V, iron(II1) is a strong inhibitor. This may be accounted for by the following reactions. COz- + Fe3+ COz + Fez+ [fast] (14)
SO,- + Fez+
-,.
+
Fe3+ [fast] (15) Two radicals of COz- and SO4- are removed from the reaction scheme by a chain reaction of eq 14 and 15. Therefore the catalyzed reaction would be greatly inhibited by even a small amount of iron(II1). Baxendale and Smithies7 have reported the reaction of (AI) and (14). Generally with COz- and the other organic radicals Baxendale and coworkers have pointed out that Cuz+ is more reactive than Fe3+, and especially with COP- in 0.1 N HzS04, the rate constant for Cuz+ is about 6 times as large as that for Fe3+. This is consistent with the fact that iron(II1) is a strong inhibitor and that, however, it can inhibit incompletely even when the amount of iron(II1) existed equally to that of copper(I1). It is well known that Fez+ reacts with peroxodisulfate.
-
+
SO , -; + 2Fe2+ 2Fe3+ 2S0,'(16) However, reaction 16 will hardly occur under the situation because reaction 15 will be much faster than reaction 16. The inhibiting effect of manganese(I1) was much weaker than that of iron(II1). In the case of manganese(II), the following reactions may be presumed. Mn(I1) SO,Mn(II1)Ox + SO,'(17)
+
-
+ + -
Mn(II1)Ox
Mn(I1)
C0,-
+
C02
(18)
+
Mn(II1)Ox C02Mn(I1)Ox COz (19) If reaction 18 were as fast as reaction 17 and only reactions 17 and 18 occurred, manganese(I1) could not inhibit the catalyzed reactions. So reaction 18 must be relatively slow and reaction 19 would occur to some extent. Reactions 17 and 19 will be the cause of the inhibiting effect. In case of iron(III), the internal redox reaction of oxalatoferrate(II1) as in reaction 18 does not occur in the dark but occurs in the light.5 Hydrogen ion accelerated the coppercatalyzed reaction. This may be mainly due to the promotion of initiation reaction l a by hydrogen ion. Protonations of COz- and OzCOz- also probably have an effect on the radical reaction rate. Oxalatocuprate(I1) will be Cu(Cz04),(2-2~)in which n should be changed by the hydrogen ion concentrations in solution, and the oxidation potential of copper will change accordingly. Thus a few effects of hydrogen ion on the reaction probably will occur simultaneously. Acknowledgment. The author wishes to thank the Japanese Ministry of Education for the financial support granted for this research.
