J. Phys. Chem. 1983, 87, 4686-4689
4688
charged with reactants as soon as possible to minimize extraneous leakage of oxygen into or out of the solutions. Visible absorption changes vs. time were monitored at 736 nm. The same procedure was followed in studying the kinetics of oxidation of cuprous acetate in pyridine, except that visible absorption changes vs. time were monitored at 632 nm.
Acknowledgment. We thank Drs. Stylianos Sifniades, Robert S. Cooke, William J. Boyle, and Willis B. Hammond for their continued help and advice during the course of this study. We especially thank Dr. George S. Hammond for his critical insights and valuable contributions during the preparation of this manuscript. Registry NO.CuCl, 775889-6; CUOAC, 59854-9; 02,7782-44-7.
Mechanistic Detalls of the Oxidation of Oxalate by Manganese( I I I ) Timothy J. Jones and Richard M. Noyes' Department of Chemlstty, University of Oregon, Eugene, Oregon 97403 (Received: December 1, 1982)
-
Oxalate complexes of manganese(II1)decompose by a net process that can be generalized as 2Mn(III) + C204'2Mn(II) + 2C02. The rate-determiningstep is a unimolecular reaction of a single Mn(II1) complex to form Mn(I1) and a radical intermediate assigned the formula .COZ-. That radical may then react either with COZor with Mn(II1). Adding acrylic acid as a radical scavenger has little effect on the rate of Mn(II1) consumption, and we reach a somewhat qualified conclusion that 2.COZ- C2042-is a more important radical consumption reaction than is COz- + Mn(II1) COz + Mn(I1) even though [Mn(III)] >> [-CO;]. In the presence of air, the decomposition of manganic oxalate induces the reaction O2 + C2Od2-+ 2H+ 2C02 + H202. However, as Kolthoff showed almost 60 years ago, this induced reaction does not affect the quantitative stoichiometry of the overall reaction between permanganate and oxalate.
-
-
Introduction Oxalate is oxidized by permanganate. Under certain conditions, the stoichiometry of process T is obeyed so 2Mn04-
+ 5C2042-+ 16H+
-
2Mn2++ loco2+ 8H20
(T)
precisely that the reaction serves as an excellent standardization procedure for quantitative volumetric analysis. Although kinetic study of the reaction was initiated well over a century ago,' some aspects of the mechanism are still unclear. If sufficient manganous ion and oxalate are both present initially, the stoichiometry of the fast initial process is well described by process A. Intermediate Mn(VI1) + 4Mn(II) 5Mn(III) (A)
-
oxidation states of manganese are certainly involved, and oxalate can compete with manganese(I1) for the reduction of some of them. The dynamics of this complex overall process are not of concern to the present paper. The manganese(II1) produced in process A is present as complexes that can be generalized as Mn(Cz04)n3-2n where n may be 1,2, or 3. The subsequent behavior obeys the stoichiometry of process B. The overall stoichiometry is then generated by (T) 2(A) + 5(B). 2Mn(C204),,-2n 2Mn2++ (2n - l)Cz042-+ 2C02 (B) -+
Kinetic studies, including those by Taube2 and by Adler and no ye^,^ show that process B is cleanly first order in manganese(II1); therefore, only one manganese is present in the transition state of the rate-determining step. There is general agreement that this step initiates process B by producing a radical intermediate such as COz- or .C204(1) Harcourt, A.
V.;Esson, W. Philos. Trans. R . SOC.London 1866,
156, 193. (2) Taube, H. J. Am Chem. SOC.1948, 70, 1216-20. (3) Adler, S. J.; Noyes, R. M. J. Am. Chem. SOC.1955, 77, 2036-42.
OO22-3654/03/2O07-4686$01.5O/O
-
derived by a 1-equiv oxidation of oxalate. The rate is slower the larger the value of n and the more compressed the transition state. We interpret this trend to indicate that the C-C bond is broken in the rate-determining step B1. This assignment of the structure of the radical in-
-
Mn(Cz04)n3-2n Mn2++ ( n - 1)C20d2+ COP+ .COP(B1) termediate is not essential to the argument that follows. The radical intermediate must react further and could conceivably do so either by step B2a or by step B2@. Mn(II1) + C02- Mn2+ + C 0 2 (B2a)
-
2.C02- .+ Cz042-
(B2P)
If the mechanism involves step B2a, (B) = (Bl) + (B2a) and the experimentally observable rate of process B is equal to the rate of step B1. If the mechanism involves step B28, (B) = 2(B1) + (Bl@)and the observed rate of process B is equal to half the rate of step B1. A mechanistic decision between steps B2a and B2g could be made by introducing a scavenger S so that process C competed efficiently with the appropriate step B2. If co2- + s P (C)
-
the mechanism in the absence of scavenger involves step B2a, introduction of S will reduce the observed -d In [Mn(III)]/dt by a factor of up to 2. If the mechanism involves step B28, introduction of S will have no effect on -d In [Mn(III)]/dt. The scavenger species S must satisfy several requirements. I t must not react directly with Mn(II1) species, it must be sufficiently reactive with COz- that process C can compete with step B2a, and any further reactions of species P must not involve reduction of manganese(II1). In the experiments reported here, we have attempted to use a polymerizable vinyl monomer for species S. 0 1983 American Chemical Society
Oxidation of Oxalate by Manganese(1I I)
The Journal of Physical Chemistry, Vol. 87, No. 23, 1983 4687
TABLE 1: Effects of Acrylic Acid on Rates of Decay at 25 "C of Solutions Containing 2.0 x M KMnO, and 5.0 x M MnSO, [Na,C,O,I/M 0.0401 0.0802 0.0401 0.0802 [ HClO, l/M 0.0199 0.0199 0.0028 0.0028 [C,H3C0,H I /M keX/min-' 0 0.005 0.010 0.078 0..148 0.294 0.588
0.0393 0.0348 0.0346 0.0347 0.0372 0.0379 0,0440
0.0342 0.0307 0.0304 0.0286 0.0278 0.0280
0.0307 0.0336 0.0347 0.0342
0.0280 0.0306 0.0316 0.0296
0.0344 0.0282
Experiment Procedures A stock solution of potassium permanganate was prepared by boiling for 15 min, cooling, filtering, and standardizing by titration of primary standard sodium oxalate by conventional analytical procedures. Other solutions were prepared from reagent-grade perchloric acid and mariganous sulfate and from dried primary standard sodium oxalate. For some runs, reagent-grade sodium perchlorate was used to maintain constant ionic strength. Triply distilled water was used for all solutions. Scavenger solutions were prepared with commercial monomers such as acrylic acid and acrylamide. For each run the MnSO,, Na2C204,and HClO, were first mixed and then thermostated at 25.00 f 0.03 "C. The KMn0, was added by pipet at time zero 80 that in all cases the resulting solution was 2.0 X lo4 M in KMnO, and 5.0 X M in MnSO,. As nearly as could be told by visual observation, formation of manganese(II1) complexes was instantaneous. Any monomer was then added, some of the solution was transferred to a silica cell with 2-cm light path, and the decomposition of Mn(II1)complexes was followed at from 450 to 520 nm with a Beckman DU spectrophotometer. Water from a thermostat was continuously pumped through the jacket of the cell compartment. However, in some experiments with large concentrations of monomer the temperature in the cell at the end of a run was as high as 26.5 "C. When the monomer concentration was no more than 0.1 M, the temperature at the end of a run was never more than 26.0 "C. In many runs, there was an "induction period" of slow reaction which is discussed below. However, every run then exhibited an extended period during which the logarithm of the absorbance was a linear function of time. The rate of change during that period was used to calculate the first-order rate constants, It,,, reported here. At long times, the absorbance became indistinguishable from that of the water used in the reference cell. Experimental Results Experiments with Acrylic Acid. Most of the scavenging experiments were carried out with acrylic acid, C H 4 H C02H. Table I shows the results of four series employing different amounts of scavenger. Most entries in the table involve an average of at least two runs whose total spread was never more than 3% and was usually less than 2%. In each of these four series, the principal species was expected to be the trioxalate complex, Mn(C204)33-,which reacts the slowest of the three. The fraction of this dominant complex should, if anything, have increased on going from left to right along the table, and the rates without scavenger are entirely consistent with this expectation. Poly(acrylic acid) is water soluble and transparent at the wavelength used. However, two pieces of evidence indicate that the monomer was indeed trapping radicals and polymerizing. First, at the larger concentrations of monomer
TABLE 11: Effects of Acrylamide on Rates of Decay at 25 "C of Solutions like Those of Table I Containing [Na,C,O,] = 0.0401 M, [HClO,] = 0.0028 M
0 0.005 0.010
0.0306 0.0328 0.0293
0.040 0.060
0.0238 0.0213
the temperature of the cell increased up to 1.5 "C, while in the absence of monomer it never increased more than 0.3 "C. Separate experiments demonstrated that these temperature changes were not associated with heats of mixing. Furthermore, a few experiments showed that the viscosity of a solution increased monotonically by about 3% during reaction. We also examined the effect of acrylic acid on solutions prepared with [Na2C204]= 0.005 M, [HClO,] = 0.020 M. Such solutions should have contained mostly the dioxalate complex, Mn(C20J2-,with perhaps a kinetic contribution from the rapidly reacting monooxalate complex, MnC204+. Additions of 0.15-0.6 M acrylic acid increased the rate of process B by up to 40%, but we subsequently found that the rate of this mixture was sensitive to ionic strength. When the ionic strength was kept constant at 1.0 M with NaClO,, additions of up to 0.075 M acrylic acid caused the rate constant to go through a shallow minimum only 7% below that with no addition of monomer. Experiments with Acrylamide. Table I1 reports the result of a series of runs of the composition of the third column in Table I except that acrylamide, CH2=CHCONH2was used as monomer. These runs show a significant decrease in rate when scavenger is added, although it does not approach the factor of 2 anticipated if step B2a were dominant in the absence of scavenger. Other Experiments. Both acrylic acid and acrylamide form water-soluble polymers but cause ambiguities in interpretation as discussed below. We therefore attempted a few experiments with the monomer acrylonitrile, CH2=CHCN. Radicals were obviously being scavenged because the solutions became turbid with polymer. That turbidity prevented kinetic measurements by spectrophotometry. We also attempted to follow process B potentiometrically in both the presence and absence of acrylamide, but we abandoned these efforts because estimation of extent of reaction involved ambiguities that were absent from the spectrophotometric measurements. Induction Periods. As was mentioned in the section on Experimental Procedures, our experiments were often plagued by initial periods of slow reaction before good kinetic data were obtained. The problem seemed to be more severe in solutions that contained larger concentrations of free oxalate ion and which hence reacted more slowly. It also seemed to be more severe in those solutions of lower pH which contained more perchloric and acrylic acids relative to the available oxalate. We had already made most of our experiments before we rediscovered the excellent paper by Launer: who discovered this effect and explained it half a century ago. Figure 1 provides a particularly dramatic illustration of the effect and of its virtual elimination by purging the solution with nitrogen before the start of reaction. The implications are discussed below. Discussion Mechanism of Process B. As was indicated in the Introduction, this study was undertaken to determine (4)Launer, H.F.J. Am. Chem. SOC.1933,55, 865-74.
4668
1.6 r 1.5
a c
Jones and Noyes
The Journal of Physical Chemistty, Vol. 87, No. 23, 1983
kinetic results would have been obtained if acrylamide competed with oxalate as a complexing agent for manganese(II1). We believe that on balance the evidence is most consistent with an interpretation that Cop-radicals are more likely to disappear by reaction with each other (B2g) rather than by electron transfer to manganese(II1) (B2cu). Because Taube2 found some peroxide product even in deoxygenated solutions, protonated C02-species may occasionally form performic acid
/-
'
0
11
0
/I I1 HCOOCH
09
06 0
I
I
I
I
I
I
5
IO
15
20
25
30
t i m e , min Flgure 1. Logarithmic plots of absorbance at 520 nm and 25 OC for three different runs each initially containing 1.9 X lo-' M KIMnO,, 3.85 X lo-' M Na,C,O,, 9.5 X lo4 M MnSO,, 1.91 X lo-* M HCIO,, and 0.9 M NaCIO,. Runs from upper to lower were saturated before the start with NP,air, and O,, respecthrely. The limiting slopes in min-' were 3.24 X lo-' for N, and 3.19 X lo-' for air.
whether step B2a or B28 provides a better description of the second step of process B. The section on Experimental Results presents thermal, viscosimetric, and turbidimetric evidence that the added monomers behaved as intended and scavenged organic radicals with initiation of polymerization reactions. We would not expect the monomers themselves to react directly with manganese(II1) species, and we see no evidence that they did so react. We are less certain that the polymeric radicals do not reduce manganese(II1). However, if .COP-radicals react so slowly with Mn(II1) that addition to monomer competes with step B2a, then a polymeric vinyl radical should be even less reactive and should preferentially engage in radical termination steps such as combination and disproportionation. For the data in Table I, acrylic acid monomer never decreased -d In [Mn(III)]/dt by more than 18%, and the effect was usually much less than this or even in a direction opposite to any allowed by the argument in the Introduction. At least superficially, such results indicate that step B2cv is unimportant relative to step B2P. Such a conclusion is not entirely unequivocal. Acrylic acid is almost as strong an acid as is HC204-,and at large concentrations of monomer equilibrium reaction D might C2H3COPH + (22042- ~ iC2H3C02. + HC204- (D) Mn(C204)33-F? Mn(C204)2-+ (E) significantly deplete [C2042-]and shift equilibrium E to the right. Such a shift would increase the rate of manganese(II1) consumption and thus might mask a rate decrease that would otherwise occur. However, our best estimates of magnitudes indicate that such an effect could not have generated the observed data if step B2a were making a major contribution to the mechanism. The data with acrylamide in Table I1 are still more equivocal. A monomer concentration of 0.06 M did indeed deplete the rate of process B by 30% consistent with considerable contribution from step B2a. However, similar
.
rather than the isomeric oxalic acid. Probably it is not surprising that we failed to reach a certain mechanistic conclusion for a reaction whose kinetics have been investigated actively for well over a century. The final resolution probably will come not from a single definitive experiment but from a consistent convergence of several suggestive pieces of evidence like that obtained here. Any stable organic compound that does not react directly with manganese(II1) will almost certainly be in a singlet ground state and will create a potentially reactive radical when it reacts with .CO2-. Perhaps more definitive conclusions can be obtained if the scavenger is an inorganic 1-equiv oxidant like Hg2+or Ce4+. Implications for Mechanisms of 1-equiv Oxidations. Process B is only one example of a large number of reactions in which organic compounds are oxidized by strong 1-equiv oxidants such as Mn(III), Ce(IV), Co(III), Fe(III), Cu(II), Ag(I), etc. These reactions are often initiated by steps like B1 in which the oxidation involves intramolecular electron transfer within a complex of the oxidant ion. One product of that electron transfer will be an organic radical which is invariably a stronger reductant than the original compound. However, in spite of the very negative AGO of step B2a, the rate constant is so small that step B2P is evidently faster even though [Mn(III)]>> [COP-]. Ganapathisubramanian and Noyes5have recently reported a very similar behavior during the oxidation of malonic acid by cerium(1V). They also have interpreted their results to indicate that the organic radicals reacted ultimately with each other rather than with Ce(1V). However, oxidation of organic radicals may occur under certain conditions. The same authors5 report that a bromide complex of cerium(1V) can oxidize malonyl radicals by process F. Apparently species like water, oxalate,
+
CeBr3+ R-
-
Ce3++ RBr
(F)
sulfate, etc., in the first solvation sphere act as "insulators" to inhibit electron transfer from organic radicals while bromine atom transfer becomes a facile way to reduce the oxidation state of an inorganic ion. It is difficult to obtain unequivocal data about reactions of organic radicals with inorganic 1-equiv oxidants, but there is need for a systematization of the type of thermodynamically allowed reactions which may or may not occur with facility. Zmplications for the Mn(VIT)-to-Mn(IIT) Transition. Unless Mn(I1) is present in considerable stoichiometric excess, during reduction of permanganate by process A some oxalate is also oxidized by manganese in oxidation states greater than 3+. Adler and Noyes3argued that such oxidations took place by 2-equiv processes because otherwise C 0 2 - . radicals would have reacted with manganese(1II) to generate branched-chain autocatalysis contrary (5) Ganapathisubramanian, N.; Noyes, R. M. J.Phys. Chem. 1982,86, 5158-62.
The Journal of Physical Chemistry, Vol. 87, No. 23, 1983
Oxidation of Oxalate by Manganese(II1)
to observation. However, that argument assumed that step B2a was dominant. The finding of this paper that step B2p is the most important invalidates the argument of Adler and Noyes. I t appears now that higher oxidation states of manganese may oxidize oxalate by either 1- or 2-equiv steps. Explanation of Oxygen Effects. The induction periods were a definite complication for our work, although a specific run in a spectrophotometer cell would eventually exhibit behavior consistent with good first-order decay. We could have avoided those problems if we had been more conversant with the prior literature. Launer4 observed the same effect long ago and explained it by the sequence of steps G1 and G2. The sequence (Bl) + (G1)
+ 02 e ozcoo-. (GI) Mn2+ + O2CO0-. + 2H+ Mn(II1) + C02 + HzOz c 0 2 -
+
-
(G2) (G2) generates the stoichiometry of process G, which is Cz042-+ 02 + 2H+
4
2C02 + Hz02
(G)
a manganese(II1)-catalyzed oxidation of oxalate by oxygen going at a rate no greater than that of step B1. When one of our solutions entered the spectrophotometer, the concentration of dissolved oxygen was about 2.6 X M. Once it had been consumed by process G, we were able to follow process B satisfactorily. This interpretation is consistent with our observation that the induction period was longest in those solutions in which step B1 was slowest. Addition of acrylic acid sometimes caused an induction period when none was observed in the absence of monomer; presumably a vinyl radical adds oxygen to produce an oxidizing peroxy radical more easily than C02-does. Finally, our observations at different acidities suggest that HOC-0 reads with oxygen more readily than O=C=O- does; such a conclusion is consistent with delocalization of the odd electron available for bonding. An alternative explanation of this pH effect is to replace steps G l and G2 by steps G3 and G4. The net effect of HC02. + 0 2 -* C02 + HOz. ((33) Mn2++ H02. + H+
-+
Mn(II1)
+ Hz02
(G4)
either pair of steps is the same. The need for hydrogen atom transfer could then explain the more severe oxygen effects in the more acidic solutions. Implications of Oxygen Effects for Analytical Chemistry. When oxalate is titrated with permanganate in a well-stirred solution, process G should certainly compete with process B for the consumption of oxalate. Nevertheless, the quantitative stoichiometry of process T is well established as a reliable analytical procedure. In 1924 KolthofP made a careful study which addressed this apparent paradox. His very farsighted discussion is directly relevant to the present paper. Kolthoff demonstrated that hydrogen peroxide did indeed form during a titration but was then destroyed by (6) Kolthoff, I . M . Z.Anal. Chem. 1924, 66, 185-211.
-
4689
process H. The stoichiometry of process T is generated
+ 5H2O2+ 6H+ precisely by 5(G) + (HI. 2Mn0,
2Mn2++ 502
+ 8H20
(H)
The existence of such precise net stoichiometry puts severe restrictions on permissible reactions in this system. Thus, hydrogen peroxide does not act as an oxidant during any subprocess, and there is no catalysis of the energetically favored process I. 2H202 2Hz0 + 0 2 (1) These restrictions are probably equivalent to saying that hydroxyl radicals, HO., do not exist as intermediates along any important reaction paths. One might reasonably expect COT and perhaps some intermediate oxidation states of manganese to act as l-equiv reductanh by processes that might be generalized by process J. For whatever the X H202 XOH HO. (J) reasons, the rates of J processes seem to be too slow to be important. Concluding Remarks. The development of this paper illustrates the frustrating nature of oxidation-reduction reactions in aqueous solution. Net processes like T, B, G, H, and I were for many years the backbone of any freshman course in general chemistry. Yet we are still ignorant about the detailed mechanisms by which many of these reactions occur. Furthermore, the technology presently available is inadequate to observe directly many of the transient intermediates that we know must be present. Techniques for their quantitative study may become available some day. Until the day comes, we must rely on the sort of laborious logical analysis employed here.
-
+
-
+
Acknowledgment. This study was undertaken as an undergraduate research project by T.J.J. and was continued during a postgraduation summer. It was supported in part by a Grant from the National Science Foundation to the University of Oregon. R.M.N. is grateful to the Max-Planck-Institut fur Biophysikalische Chemie in Goettingen for hospitality during the preparation of the manuscript. Note Added in Proof. Since this manuscript was accepted for publication, Professor Henry Taube of Stanford University has pointed out to us that manganese(II1) catalyzes the oxidation of oxalate by the halogens ClZand Br,.' That catalysis is interpreted to involve a sequence initiated by step B1 in which the organic radical is scavenged by the molecular halogen and the resulting halogen atom reoxidizes the manganese(I1). Professor Taube has referred to some of his rate measurements which seem to support the conclusions reached in the present paper. Such scavenging of organic radicals by molecular halogens should permit a still more definitive determination of the relative importance of steps B2a and B2P. Registry NO. Mn, 7439-96-5;C2044 33870-5; CHZ+HCO,H, 79-10-7; CH2=CHCONH, 79-06-1; CH,=CHCN, 107-13-1. (7) Taube, H.J . Am. Chem. SOC.1946,68,611-7; 1947,69,1418-28; 1948, 70, 3928-35.
