The kinetics of the oxidation of hydroquinone by Pu(VI) and by Pu(IV) in aqueous perchlorate solutions

J. inorg, nucl Chem., 1974, Vol. 36, pp. 639-643. Pergamon Press. Printed in Great Britain.

THE KINETICS OF THE OXIDATION OF HYDROQUINONE BY Pu(VI) AND BY Pu(IV) IN AQUEOUS PERCHLORATE SOLUTIONS* T. W. NEWTON Los Alamos Scientific Laboratory of the University of California, Los Alamos, New Mexico 87544

(Received 29 May 1973)

Abstract--The reduction of Pu(VI) and Pu(IV) has been studied in aqueous perchlorate solutions. The reactions go to completion with the reduction of two moles of plutonium for each mole of hydroquinone oxidized. The kinetics are quite complicated: the rate laws show inhibition by the reduced plutonium species, but not by quinone, and for Pu(VI) is approximated by: -d[hydroquinone]/dt = kl[Pu(VI)] [hydroquinone]/(l + a[Pu(V)]/[Pu(VI)] + b([Pu(V)]/[Pu(VI)])2). The same form, with different values for the parameters, is observed for Pu(IV). For Pu(VI), kl is 1.6 x 103 M -1 see -~ at 25° in 1 M HC1Og, depends inversely on [H+], and has an activation energy of 15.7 + 0.1 keal/mole. For Pu(IV), kt is about 9'6 x l0 t under the same conditions, is essentially independent of [H+], and has an activation energy of 12 + 1 kcal/mole. The results are compared with those for other metal ion oxidizing agents.

INTRODUCTION

EXPERIMENTAL

l ~ c ~ w r work on the oxidation ofp-hydroquinone and p-toluhydroquinone by N p O 2 + [1] has shown that the reactions are cleanly first order in each reactant, independent of the concentrations of the reaction products, and independent of the acid concentration. A purpose of the present paper is to compare these results with those for the analogous reaction between PuO 2+ and p-hydroquinone. N p O 2+ is a better oxidizing agent than PuO~ + by 0.22 V but otherwise the two ions are very similar and the same terms often appear in the rate laws for corresponding reactions[2]. However, P u O 2+ shows much more complicated kinetics. It has not been possible to formulate a complete rate law, but it is clear that the plutonium reaction is strongly inhibited by the product, Pu(V), and that the initial rates are predominantly inverse first power in the hydrogen ion concentration. The work was extended to the reduction of Pu g+ by hydroquinone and that reaction was found to be quite complicated also, The rate law is similar to that for the Fea+-hydroquinone reaction[3] in that the rates are inhibited by the metal ion product, but additional, poorly understood, terms are present also.

Reagents Hydroquinone, twice sublimed under vacuum, was provided through the courtesy of Dr. J. C. Sullivan, Argonne National Laboratory. Weighed amounts were dissolved in deoxygenated solutions 0-04 M in HCIO4. The solutions were prepared fresh each day and were standardized by spectrophotometric titrations with Co(IV) in 0.5 M H2SO4. Solutions of Pu(III) were prepared and standardized as previously described[4]. Plutonium (VI) was made using fuming HCIO414] and standardized by adding a known excess of Fe(II) and titrating to Pu(IV) with Ce(IV) in 0-5 M H2SO4. Solutions of Pu(IV) were prepared immediately prior to the runs by adding known amounts of Cr(VI) or Ce(IV) perchlorate to excess Pu(III). The observed rates were independent of the oxidant used. The distilled water and the solutions of HC104 and LiCIO4 were prepared and standardized as previously described[4]. The concentration units are moles per liter, M, at 23*. Various of the reagents were repurified without significant effect on the reaction rates, the plutonium solutions by ion exchange, the HC104 by vacuum distillation, and the hydroquinone by recrystallization from toluol. Although oxygen from the air was found not to affect the observed rates, it was routinely excluded by means of a stream of scrubbed argon.

Stoichiometry * Work supported by the U.S. Atomic Energy Commission.

The stoichiometry of the reactions was determined by spectrophotometric titrations in which increments of standard hydroquinone were added to the plutonium solutions 639

640

T.W. NEWTON

with a micrometer buret. Plots of absorbance vs ml added were linear with sharp breaks at the end-points. Reaction rates

Rate runs were made in stirred absorption cells as previously described[4, 5]. The wavelengths used were 8304, 6005, 4696, 2885, and 2450/~, where the principal absorbing species are Pu(VI), Pu(III), Pu(IV), hydroquinone, and quinone, respectively. The apparent rate constants were essentially independent of the wavelength used. Initial rates were estimated by fitting the absorbance vs time data to four-parameter power series. Instead of extrapolating to the time of mixing the slope was determined halfway between the first twO,experimental points. Various other smoothing functions, including the integrated forms of the appropriate empirical rate laws, gave essentially the same slopes. The initial apparent second order rate constants were found by dividing the slopes by the concentrations of the two reactants at the appropriate times. Extrapolation to the calculated initial absorbance readings gave times within about 1 see of the times of mixing. Thus there was no evidence for the rapid formation of intermediates.

Table 1. Effect of Pu(V) on the Pu(VI)-hydroquinone reaction* initial concentrations (M × 104) Pu(V) Pu(VI) HzQ Ratio1" (k')o:~ 0

0"76

1"12

1"94

1"06

1.12

3'94

0.74

1"13

5.12

0'96 1'04

1'13

0.11 0.11 1"96 2"18 5.56 5.60 5.14

127 126 68.4 65"6 30.0 33.7 29.0 32'5

kl§ 132'6 _+ 1"4 131.4-1-1"6 132.6+ 1'4 135"3 + 1'3 132.0+ 2"9 148.2 :t: 5"9 128"5_ 1"4 132.6 + 1'6

*Conditions: 1.0 M HCIO 4, 1"8°. i"The ratio [Pu(V)]/[Pu(VI)]at a time halfway between the firsttwo experimentalpoints. ~:The "initial"rate,determined at the time given in ~'. § The rate constant and its standard deviation from the integrated form of the rate law: -d[Pu(VI)]/dt = 2ki(I + 0.407EPu(V)]/[Pu(VI)] + 0.0368([Pu(V)]/[Pu(VI)])2)- 1.

strength. Runs were made with and without added Pu(V) so that its effectcould be determined. The concentration of Pu(V) was kept low enough so that the b in The Pu(VI)-hydroquinonereaction Eqn (2)could be ignored. By considering the initialrates The stoichiometry of the overall reaction is in accord it was found that the parameter a depends slightlyon with the equation [H+]: a--0.34+0.31[H+]. Absorbance versus time 2PufVI) + H2Q = 2Pu(V) + Q (1) data for the first75-90 per cent of the reaction were then used to determine the best values for kl in Eqn (2).The where H2Q stands for C~H,(OH)2 and Q stands for results are summarized in Table 2. C6H402. In three determinations 14 ml of 7.01 x 10 -4 Six additional runs were made in 1.00 M H C I O , at M Pu(VI) (9"82 t~-mole) required 2"095 ___0.007 ml of 25.2° using the same reactant concentrations as those 2.347 x 10 -a M H2Q (4.917 ± 0'02 #-mole) giving a listedin Table 2. The value for a was found to be 0.73, stoichiometdc ratio of 2.00 + 0.01. essentiallythe same as for 1.8°. The average value and Concentration versus time data from the various rate mean deviation for kl was 1603 + 38 M - I sec- i when runs give non-linear second order plots. For a typical the initialconcentration of Pu(V) was zero and 1606 + run in 1 M HC10, at 1.8" using 7.6 x 10 -5 M and 22 M -I sec -I when [Pu0O]0 was 4 x i0 -s M. These 1.12 x 10-* M Pu(VI) and H2Q respectively, the initial results lead to an activation energy of 15.7 ± 0.1 kcal/ slope corresponds to 116 M-1 sec-1 while that between mole. The initialapparent second order rate constants in 70 and 80 per cent completion is only 50 M-1 sec-1. Experiments showed that the product, quinone does not I M H C I O , at 1"8° reported in Table 2 are significantly affect the rate but that Pu(V) does. The data summarized larger than the values given in Table 1.The resultssumin Table 1 show how the initial apparent second order marized in Table 3 provide further evidence that these rate constants are decreased by Pu(V). The initial parameters depend somewhat on the initialconcensecond order rate constants can be approximated by trations. the empirical expression: RESULTS

(k')o = 2k1(1 + ar + br2) -1

(2)

The Pu(IV)-hydroquinone reaction

The stoichiometry of the overall reaction is in accord where r = [Pu(V)]/[Pu(VI)], a = 0.407, and b = 0"0368. The change in apparent rate constant given by this with the equation expression is in good agreement with the concentration 2Pu(IV) + H 2 Q - - 2Pu(III) + Q. (3) vs time data for the individual rate runs. Values for kl were determined for each run using the integrated form Known amounts of Pu(IV) were prepared by adding of the rate law implied by Eqn (2) and the values for a about 3/4 the stoichiometric m o u n t of standard Ce(IV)and b given above. Except for one discrepant value, the perchlorate to Pu(III) in 1 M HC10 4. Examination of others are all near 132 M -t sec-t and show no trend the solutions at 830 nm showed that about 0.5 per cent of the plutonium was oxidized to Pu(VI) under these with intial [Pu(V)], see last column of Table 1. The hydrogen ion dependence was studied in a series conditions. In three experiments the ratio of Pu(IV) of experiments at 1"8° in LiCIO, solutions of unit ionic reduced to hydroquinone added was found to be

Oxidation of hydroquinone

641

Table 2. Hydrogen ion dependence for the Pu(VI)-hydroquinone reaction* HCIO4 (M)

Initial concentrations Pu(VI) H2Q Pu(V) (M x 105) (M x 105) (M x 105) al"

1.00 0.624 0'598 0'245 0"198

9"68 5-01 5.01 9"46 5'01 9'46

2.11 1.94 1.94 2.06 1'94 2'06

0 3"88 3'88 0 3'88 0

0"65 0'65 0'53 0'52 0'42 0"40

kl kl (calc):~ (M-1 see-1) (M-1 sec-1) 169 + 172 _ 257 _ 274 + 623 + 767 +

5§ 4 5 4 10 8

171 171 258 268 618 770

* Conditions: 1'8°, # = 1'0 M (LiCIO4). ~"From a = 0'34 + 0"31[H+]. :~Calculated using k 1 = 26 + 145[H+]-1. § Mean deviation, four determinations; two determinations for other entries. 2.005 ~ 0.014. The small amounts of Pu(VI) present Were assumed to be reduced to Pu(V) at the endpoints. Quinone, prepared by oxidizing hydroquinone with Cr(VI), has no significant effect on the reaction rates. In the experiments which show this, the initial ratio of quinone to Pu(IV) ranged as high as 1.2 and the initial rates agreed within the experimental error. Pu(III), however, has a large effect, as illustrated by the data in Table 4. These data are consistent with empirical E qn (2) with kl = 1.89 x 104 M -1 sec -1, a = 2 . 7 and b = 0.65, for 1 M HCIO4 at 1.8°. Variation of the acid concentration between 0.2 and 1 M at/~ = 1~) M (LiC10,) showed that kl is essentially independent of [H +] and that a is proportional to [I-I+]; the effect on b could not be determined. The temperature dependence between 1.8° and 12.6° gives an activation energy of 12 kcal/mole for kt and 17 kcal/mole for a. Experiments using a [Pu(III)]/[Pu(W)] ratio of about 3.7 showed that initial concentrations of Pu(IV) and I'I2Q as high as 7.3 x 10 -5 M and 2.96 x 10 -5 M respectively were without significant effect on the initial rates. However, using a ratio of 1.0, a significant effect was noted. The initial rate constants decreased smoothly from 5260 M-1 sec-1 down to 3000 M-~ sec-t when Table 3. Effect of initial reactant concentrations* H2Q (M x 104)

Pu(VI) (M x 104)

klt (M-lsec-1)

0'21 0"22 0-22 0.40 0-40 0.79 1"12 1.12 1.18 1.45 1.54

0'95 2"99 4'01 1'11 1.45 0.74 0-76 1.02 1'11 0'89 0.73

169 180 191 167 167 144 132 134 141 129 135

* Conditions: 1 M HC104, 1.8°. t Calculated from the integrated form of the rate law.

the initial concentration of Pu(IV) was increased from 0'68 × 10-* to 4.2 x 10-* M. These experiments were done in 1 M HC10, at 1.8" using 3 × 10 -6 M hydroquinone. DISCUSSION The apparent second order rate constant for the oxidation of hydroquinone by Fe(HI) was found[3] to be governed by the expression: k' = 2kl/[1 + (k2[Red.]/k3[Ox.])].

(4)

This was interpreted, without regard to H +, in terms of the mechanism: kl

Ox. + I-I2Q ~ Red. + HQ

(5)

k2

k3 Ox. + HQ -~ Red. + Q

(6)

where Ox. stands for the oxidizing agent, Fe(II1) in this case, and Red. for its reduced form. The net activation processes for the formation of the two activated complexes from the initial reactants are: Ox. + H2Q = [*]a + mH+ 2 0 x . + H2Q = [*]b + nil+ + Red.

(7) (8)

where m and n indicate the hydrogen ion dependences. When Np(VI) is the oxidizing agent, no inhibition by Np(V) was observed[I]; this is consistent with the above mechanism provided k 3 >> k2, which is reasonable since Np(VI) is a much stronger oxidizing agent than Fe(III). For both Pu(VI) and Pu(IV), the empirical rate law, implied by Eqn (2), is somewhat more complicated than (4) in that an additional term involving the square of the [Red.]/[Ox.] ratio is required. This suggests that unknown additional reactions of the radical HQ also occur. It is very likely, however, that consecutive reactions like (5) and (6) are involved and the two most important activation processes are those indicated by (7) and (8). Making this interpretation, the hydrogen ion

642

T. W. NEWTON Table 4. Effect of PU(III) on the Pu(IV)-hydroquinone reaction* Initial concentrations (M x. l0 s) Pu(III) Pu(IV) H2Q 0'75 0.285 0-46 0.64 1"17 1-52 2.65 2-70 3.54 2'96 3'74

1"35 0.92 1.21 0.97 1"31 1"54 0"98 1"24 1'39 1"01 1.01

No. of det'ns

0'295 0"24 0'30 0.36 0'35 0-26 0-31 0"33 0-39 0"38 0"34

Ratio?

(k')o:~ (M- 1 sec- 1)

0"24 if31 0'38 0-66 0-895 0'99 1'49 2'18 2'55 2'94

9830 10800 + 100§ 10080 + 440 6700 5320 ___94 4000 + 110 3360 + 150 1940 + 85 1450 + 15 1254

1 2 2 1 2 2 3 4 2 1 4

* Conditions: 1 M HC104, 1'8°. ? The ratio [Pu(III)]/[Pu(IV)] at a time halfway between the first two experimental points. :~The "initial" rate, determined at the time given in ?. § The mean deviation from the mean. dependences and other rate parameters for the various oxidations of hydroquinone are compared in Table 5. It is seen that all the parameters change from oxidant to oxidant and in general depend on the reduction potential. The change in hydrogen ion dependence with driving force is the same as that observed for other actinide reactions[2]. The k2/k3 ratios show the expected trend with potential while the kl values for Pu(IV) and Np(VI) are reversed. The observed kx and k2/k3 values may be used to estimate upper and lower limits for the reduction potential for the quinone-hydroquinone couple: HQ + H + + e- = H2Q.

(9)

can be shown* that steady-state conditions will be reached rather slowly and the kinetics will be perturbed significantly. Since the kinetics in the inltial stage of reaction were as expected, it is probably safe to assume that [HQ] was never larger than 0.2[H2Q]. As an example, these considerations will be applied to the Pu(VI) data. For the reverse of Eqn (5) k2 will be 0.7 k 3 or _<_5 × 109 M - 1 sec r- 1. The equilibrium quotient, kl/k2, is greater than 1.6 x 103/5 × 109 and the reduction potential for Eqn (9) is < 1.31 V. The steady-state ratio [HQ]/[H2Q] is given by (kl/k3)/(1 + k2[Red.]/ k3[Ox.]), so kl/k3 <-0"2 and kl/k2 = Q < 0 " 2 / 0 ' 7 = 0"29, and E ° > 0"95 V. The results of similar calculations for Fe(III) and Pu(IV) are given in the last column of

The upper limit is based on the hypothesis that the larger of k2 or k 3 is diffusion controlled (kdin = 7"109 M - 1 sec- 1) and the lower limit on the validity of the steady-state approximation. If the steady-state concentration of HQ is larger than about 0.2 of that o.f H2Q, i t

* Numerical calculations using kl =0.748 and k2/k 3 = 9'09, typical values for the Fe(III)-H2Q reactions, show that about 20 see are required for the [HQ]/[H2Q] ratio to reach 0"75 of the value calculated using the steady-state approximation. Table 5. Oxidation of hydroquinone by various oxidants? n

Oxidant

(V)

Fe(III) Pu(VI)

0-74§ 0"9211

I I, 0

2 I, 2

Pu(IV) Np(VI)

0-9811 1.14¶

0 0

1 -

k1 (M- 1 see- t) 0.69?? 1.6 x 103

k2/ka 15.47? 0.7

9.6 x 10~: 0-24:~:~ 2.26 x 10"§§ <0"011111

AHI* ASI* (kcal/mole) e.m.u. 25.0 15.1

11.4 8.2

24.6 6.8

S**mp

E°(HQ)~¶ (V)

- 3.6 27.1

0-85-1.33 0.95-1.31

2.5 -32.5 -11-2§§ 1 6 - 1

0.98-1.23 ???-1.35

t 1 M HC104, 25°. Formal reduction potential. §R. E. Connick and W. H. McVey, J. Am. chem. Soc. 73, 1798 (1951). II Rabidcrau, et al., Proceedings of the Second UN Geneva Conference, Vol. 28, p. 361. United Nations, Geneva (1958). ¶ A. J. Zielen and J. C. Sullivan, J. phys. Chem. 66, 1065 (1962). ?? Data from [4], extrapolated to 1 M HC104. :[::~This work, extrapolated to 250. §§ From [2], assuming that kobs= 2kl. [I[[ [2], and the assumption that 3-46 x 10 -3 M Np(V) decreased the rate by less than 2.5 per cent. ¶¶ Limits for the reduction potential of the semiquinone radical, estimated as described in the text. ?1"? A lower limit cannot be estimated from the data.

643

Oxidation of hydroquinone Table 5. It should be pointed out that although the Pu(IV) reaction gives narrowest limits they are the least reliable because of uncertainties in the interpretation of the rate law. It is of.interest to compare these limits with values based on previous estimates of related thermodynamic quantities. Marcus[6] estimated the free energy change for the reaction H2Q + Q = 2HQ to be about 2.3 kcal/ mole. If this is combined with 0.699 V for the reduction potential of quinone[7], a reduction potential of 0.75 is obtained for Eqn (9). More recently Hale and Parsons [8] estimated AF° for the reaction H2Q + Q = 2HQ to be 14.4 kcal/mole based on the equilibrium quotient in alkaline solution[9]. This value leads to a reduction potential of I'01 V. Thus the kinetic measurements and their interpretation in terms of Eqns (5) and (6) are consistent with the second estimate of the reduction potential but not the first. The formal entropies of the activated complexes range from "32.5 e.m.u, for (Pu. H2Q4+) * up to 27 e.m.u, for (PuO2. HQ÷) *. These values are considerably more positive than those for metal ion-metal

ion reactions but the dependence on the charge is similar to that observed previously[2]. REFERENCES

1. K. Reinschmiedt, J. C. Sullivan and M. Woods, lnorg. Chem. II, 1639 (1973). I thank the authors for providing the manuscript of this paper before publication. 2. T, W. Newton and F. B. Baker, Adv. Chem. Ser. No. 71, p. 268. American Chemical Society, Washington, D.C. (1967). 3. J. H. Baxendale, H. R. Hardy and L. H. Sutcliffe, Trans. Faraday Soc. 47, 963 (1951). 4. R. B. Fulton and T. W. Newton, J. phys. Chem. 74, 1661 (1970). 5. T, W. Newton and F. B. Baker, J. phys. Chem. 67, 1425 (1963). 6. R. A. Marcus, J. chem. Phys. 26, 872 (1957). 7. W. M. Latimer, Oxidation Potentials, 2nd Edn, p. 138. Prentice-Hall, Englewood Cliffs, N.J. 0952). 8. J. M. Hale and R. Parsons, Trans. Faraday Soc. 59, 1429 (1963). 9. H. Diebler, M. Eigen and P. Matthies, Z. Naturf. 16b, 629 (1961).