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CATALYSIS OF DECOMPOSITION OF CARO’S ACID
which must be broken in step 9, is twice as strong as the 0-0 bond involved in (10).I6 A recent reinvestigation of the reaction of atomic hydrogen with hydrogen peroxide has again been interpreted in favor of Geib’s mechanism.17 On the assumption that his measurements referred to reaction 10, Geib calculated the collision yield and, taking a steric factor of 0.1, he estimated the activation energy to be about 4.5 kcal. A rough calculation of Elo also may be attempted from the present results by obtaining first an estimate of k~ a t the temperature of our measurements. Lewis and von Elbe7have given an expression for kl but there are reasons to believe that it is in error by a factor of lo3 a t 520’. According to Baldwin and Walshl*a more plausible value would be em. molecules-’ sec.-l a t that temperature. WarrenlShas measured the temperature dependence of the second explosion limit, and obtained E,, = (El - Ed) = 20 1 kcal. Now, Hoare and WalshZ0have proposed for reaction 4 a negative activation energy of about -4 kcal., which makes E1 cr: 16 kcal. in good agreement with the endothermicity of that reaction, 17 kcal. a t 700°K.21 These data lead to kl N 3 X C M . molecules-’ ~ set.-' a t 447’. Taking a = 65 one gets klo = 2akl N 4 X ~ m molecules-l . ~ sec.-l a t the same temperature. Finally, using Geib’s figures for the collision diameters of H and HzO2 the collision yield may be calculated, leading to EloN 7 kca1.22with a steric factor of 0.1. Considering the
*
(16) T. L. Cottrell, “The Strengths of Chemical Bonds,’’ Butterworths, London, 1954. (17) J. S. Batzold, C. Luner and C. A . Winkler, Can. J . Chem., 31, 262 (1953). (18) R. R . Baldwin and A. D. Walsh, Disc. Faraday Soc., 17, 97 (1954). (19) D.R. Warren, Proc. Roy. SOC.( L o n d o n ) , !4llA, 96 (1952). (20) D. E. Hoare and A. D. Walsh, Trans. Faraday Soc., 53, 1102 (1957). (21) F. D. Rossini, et al., “Selected Values of Chemical Therrnodynamic Properties,” Vol. 111, Circular 500, Washington, D. C., 1952.
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entirely different approach of the two methods, the agreement with Geib’s estimate is gratifying. It is of course no proof for the value of the activation energy. The alternative interpretation of Geib’s results proposed by Lewis and von Elbe7 would thus appear doubtful. As yet no quantitative measurements have been reported on the reactions of OH radicals. The mass spectrometric investigations of Foner and Hudson15 point to a very high rate for reaction 12. From the value obtained here for constant “b,” this reaction should be faster than 3 by a factor of a t least five. I t may also be noted that reaction 12 is twice as exothermic ( A H = -32 kcal.) as reaction 3. An interesting question raised by the present findings concerns the possible influence of hydrogen peroxide on the second explosion limit under conditions of predominant second-order branching, e.g., with a clean Pyrex surface, such as used here, but at higher temperatures. I n Egerton and Warren’s4 derivation of equation I the coefficient of the ( p ~ , ) - ’ / ~term is proportionaI t o the half-power of the rate of chain initiation. This has prompted Warren19 to suggest that explosion sensitizers may expand the limits by affecting the rate of initiation. It is entirely conceivable that under appropriate conditions the increased chain-initiating effect of added hydrogen peroxide would far outweigh its chain-breaking effect through steps 10 and 12. Thus, a t a high enough temperature, hydrogen peroxide might become an explosion sensitizer instead of an inhibitor. Such a conclusion seems all the more plausible as reactions 10 and 12 very likely have low activation energies, making their rates almost temperature independent, whereas the second-order branching term in equation I involves a large activation energy, of the order of 35 k ~ a l . ~ The authors are grateful t o the National Research Council of Canada for financial assistance. (22) A value of 300 for “a” would make E M N 5 kcsl. other things being equal.
THE CATALYSIS OF THE DECOMPOSITION OF CARO’S ACID1 BY DONALD L. BALLAND JOHN 0. EDWARDS Metcalf Chemical Laboratories of Brown University, Providence I d , R. I . Received October SO, 1967
Evidence has been found for the catalysis of the decomposition of Caro’s acid (peroxymonosulfuric acid) by specific substances in aqueous phosphate buffer. Cob?lt (11)and molybdenum(V1) are especially effective, although some ot,her metal ions also act as catalysts for the decomposition. The observed catalytic decompositions are first order in the concentration of Caro’s acid with the exception of the cobalt-catalyzed decomposition which is second order in Caro’s acid. The order in catalyst concentration could not be determined except in the case of molybdate wherein an order of one-half and an induction period were found.
Introduction The decomposition of Caro’s acid (Le., peroxymonosulfuric acid, H2S06, to form oxygen and sulfuric acid) in aqueous buffer solutions is susceptible to catalysis by trace amounts of impurities.l Apparently these catalytic paths are eliminated by the (1) The Kinetics and Mechanism of the Decomposition of Caro’s Acid. 11; prior paper, J. Am. Chem. floc., 78, 1125 (1056).
addition of sinal1 amounts of ethylenediaminetetraacetic acid (EDTA). I n the absence of EDTA, the observed rate of decomposition was dependent on the sample of phosphate used in preparing the buffer solutions; therefore the observed effect has been ascribed to the nature and amounts of the cat,alytic substances present as impurities in the phosphate. The fact2 that dipicolinic acid stabilizes
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DONALD L. BALLAND JOHN 0. EDWARDS
solutions of Caro’s acid is further evidence for catalysis by trace impurities. Recently, Kuhn3 has investigated the decomposition of Caro’s acid by small amounts of nitric acid in strong sulfuric acid solutions. Additional evidence for the catalysis of the Caro’s acid decomposition was desirable. The present paper provides the results of a survey of the possible catalytic effect of various substances on the decomposition, also kinetic investigations have been made for the two most effective catalysts found. I n view of the difficulties in interpretation, etc., only a summary of our work is presented here.4 Experimental Most of the procedures used in this study have been discussed previously.’ The catalytic survey was made in solutions buffered by phosphate (total phosphate = 0.75 M) in the pH range 6.6-6.8. The initial concentrations of Caro’s acid were from 0.01 to 0.02 M. All experiments were conducted a t 25.0”. The rate of decomposition before adding the test substance was apparently first order in Caro’s acid.’ With the sample of phosphate used in this survey, the original rate of decomposition (resulting from the spontaneous decomposition plus the catalytic rate due to catalysts originally present as impurities) was not rapid (Vi = ca. 10-8 rnin.-l). I n general, the test substance was dissolved in water before its introduction into the reaction solution. The element to be tested was added as its highest oxidation state, when available. The amount added was kept small enough that a stoichiometric reaction would not be mistaken for catalysis; usually the concentration of the test substance in the reaction solution was about 1 X IOw4 M . The effect of an added substance was assumed to be demonstrated by the difference between the rate of decomposition observed after the substance was added and the rate originally observed.
Results Catalytic Survey.-Small amounts of cobalt (11), copper (&I),nickel(I1) , ruthenium( 11), iridium(III), vanadium(V), molybdenum(V1) and tungsten(V1) increased the rate of decomposition of Caro’s acid. With the exception of the studies with cobalt(II), the catalytic decompositions appeared t o be first order in Caro’s acid; plots of the logarithm of the Caro’s acid concentration versus time were linear. The effects of cobalt and molybdenum appeared t o be over ten times greater than that of any other catalyst observed. The other substances mentioned above appeared to have roughly the same small catalytic effect. No attempt to specify their relative effectiveness is justified; their concentrations were not equivalent in general and, in some instances, were not accurately known. No observable catalytic effect was displayed by Ag(I), TI(I), Cd(II), Pb(II), Hg(II), Pd(II), Zn(II), Sb(III), As(III), Bi(III), Ce(III), Fe(III), Pt(IV), Rh(IV), Th(IV), Zr(IV), I(V), Ta(V), Cr(VI) and Mn(VI1). Some of the species declared catalytically inactive might have exerted an effect a t a higher concentration. It is also possible that, in some cases, an impurity in the substance tested was responsible for the observed catalysis; a purer sample might not have shown any catalytic effect. I n considering the results of the catalytic survey, these possibilities must be borne in mind. (2) F. P. Greenspan and D. G. MaoKeller, U. S. Patent 2,663,621 (Dec. 22, 1953). (3) L. P. Kuhn, J . A m . Chem. Soc., 79, 3661 (1957). (4) Further details, especially for molybdate catalysis, may be found in the Ph.D. thesis of D. L. Ball, Brown University, 1956.
Vol. 62
The Cobalt-catalyzed Decomposition.-Cobaltous ion a t an apparent concentration of 10-8 M noticeably accelerates the decomposition of Caro’s acid. Uniquely, the cobalt-catalyzed composition appears to be second order in the concentration of Caro’s acid. Customarily observations were made over a period exceeding one “half-life”; plots of reciprocal Caro’s acid concentration versus time ’ small compared to gave straight lines when k ~ was the catalytic rate. The observed rate law is of the form -d B = kH‘ [HSOa-] dt
+ ~~’[CO(II)]~[HSO~-]*
where k ~ is’ the rate constant observed in the absence of cobalt(II), kc’ is the cobalt-catalyzed rate constant, and n is the order in the concentration of cobalt ion. In the pH range of these studies, Caro’s acid exists primarily as the ion HS06-. The data obtained are summarized in Table I. Uncertainties in the nature of the cobalt catalyst in the system make a determination of n quite impossible a t present. It is likely that the cobalt(I1) is oxidized to cobalt(II1) immediately upon its introduction. Sob015 reports a value of 3.2 X for the activity solubility product constant of Co(OH),. A Kspof this magnitude would require that essentially all the cobalt(II1) in the system was in the form of precipit,ated cobaltic hydroxide. The possibility of precipitated cobaltous and/or cobaltic phosphates cannot be ignored; these salts are known t o be very insoluble, although useful quantitative data are not available. The observed catalytic effect then could be heterogeneous in nature and the variation of rate with concentration is not inconsistent with this. Thus uncertainties regarding the appropriate solubility analysis make further consideration of the data unwarranted. TABLE I SUMMARY OF THE EXPERIMENTS ON THE COBALT-CATALYZED DECOMPOSITION OF CARO’S ACID PH
[Co(II)]
x
10s. M
Rca
6.09 6.11 6.62 6.61 6.65 6.65
5.3 1.3 23 2.5 1.4 ca. 0.28 2.8 ca. 0.45 430 6.0 770 7.1 a Ro = kc’ [Co(II)]”. The units are (moles/l.)-l min.-l.
The Molybdate-catalyzed Decomposition.M ) of molybdenum have an Small amounts effect on the rate of the decomposition of Caro’s acid. The catalyst customarily was added as a solution of NazMo04; however comparable results were obtained using (NH4)6M07024.The decomposition was first order in the concentration of Caro’s acid; sample plots are given in Fig. 1. The observed first-order catalytic rate constant is defined as kc. An induction period was always observed after the addition of the catalyst. However, the length of this induction period (usually from five to ten (5) 9. I. Sobol, Zhur. Obshchei Khim., 906 (1953); C. A . , 48, 31090 (1954).
March, 1958
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CATALYSIS OF DECOMPOSITION OF CARO’S ACID
minutes) did not appear to have any correlation with other variables, such as pH, the concentration of Caro’s acid, and the concentration of molybdate ion. All runs were conducted in phusphate buffers a t p H values ranging from 5.90 to 7.34. I n this range, the decomposition appeared t o be one-half order in the concentration of molybdenum(V1) as may be seen in Fig. 2. . Molybdate ion is known to polymerize in the pH region of the kinetics experiments.% Therefore, the amount of molybdate added represents the total amount in all forms, both monomeric and associated. The suggested rate law in the pH range investigated is - d[HS06-1 = ka’[HSO&-] dt
+ ~,“[Mo(VI)]~/~[HSO~-]
However, variations of IC,“ with pH were not entirely reproducible. The sample of phosphate used in the buffer had an apparent effect on the decomposition rate. Using phosphate(I1) (l), the observed value for IC,” was 0.66 i 0.14 (moles/l.)-‘/2 min.-’ (29 rune) in the pH range from 6.39 to 7.34. The value for k c ” observed using phosphate(II1) (1) was 0.33 f 0.12 (19 runs) in the same pH range. At lower pH values (5 runs in the range 5.90 to 6.10) the observed value for k,” (phosphate(II1)) was considerably lower (0.13 f 0.01). The data suggest that the monomer HMo04participates in the catalytic mechanism, and that the dimer Mo&- is the principal form of molybdenum(V1) in the pH range studied. The decrease in the catalytic activity in the acidic extreme of the pH range studied could be explained by the formation of higher polymolybdates. It should be noted that molybdate has no catalytic effect on the decomposition of Caro’s acid in the presence of small amounts of EDTA. It appears that a third substslnce, or “co-catalyst,” is necessary for the obeerved acceleration of the decomposition. Since a given buffer gave fairly reproducible data, whatever the source of molybdenum, it is believed that Mo(V1) is necessary for the catalytic decomposition; in other words, no impurity in the molybdate is solely responsible for the observed catalysis. Attempts to gain more reproducible data using conductivity water and recrystallized phosphate in the (6) H.J. Emeleus and J. S. Anderson, “Modern Aspects of Inorganic Chemistry.” D. Van Nostrand Co., Inc., New York, N. Y., 2nd edition, 1952, pp. 213-216.
1
I
0.5
I
I
120 180 Time in minutes.
60
240
Fig. 1.-First-order plots of the molybdate-catalyzed decomposition a t pH 6.42. The times of the addition of catalyst are indicated by the vertical bars. Runs with a total molybdate concentration of 4.8, 8.7, 41 and 68 X 10-611.f are given by curves 1, 2, 3 and 4, respectively. 1.30
1
0.50
I
,
1
II
I
I
I
I
1
0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0 2.2 2.4 Loa [Mo(VI)I 5. Fig. 2.-A plot of the logarithm of the catalytic rate constant, k, (rnin.)+, versus the logarit#hmof the concentration of total molybdate. Runs buffered a t average pH values of 6.42, 7.10 and 6.92 are given by curves 1, 2, and 3, respectively. The use of phosphates I1 and I11 is indicated by the symbols 0 and 0, respectively.
+
buffer solution were not successful. This observation compares with the failure of similar attempts to gain reproducible measurements of the rate of decomposition of Caro’s acid in unmodified (no added EDTA) phosphate buffer solutions. Acknowledgment.-We are grateful to the Office of Ordnance Research of the U. S. Army for financial support, and D. L. B. is pleased to acknowledge the aid of a Union Carbide and Carbon Fellowship.
