GILBERT GORDON AND P. H. TEWARI
200
The Kinetics of the Reaction between Vanadium(I1) and Chlorate in Aqueous Perchloric Acid
by Gilbert Gordon and P. H. Tewari Department of Chemistry, University of Maryland, College Park, Maryland
807.42 (Received August 9, 1966)
The reaction between excess vanadium(I1) and sodium chlorate has been studied in 0.30 to 1.50 M perchloric acid. The reaction was found to be first order with respect to vanadium(I1) and to chlorate. Spectrophotometric investigations suggest that only vanadium(111)is formed as a product and seem to rule out the possibility of the formation of appreciable quantities of vanadium(1V). The effect of hydrogen ion was complicated in that, when sodium perchlorate was used to maintain constant ionic strength, the rate was found to increase with increasing hydrogen ion concentration. In the case of lithium perchlorate, however, the rate was found to decrease as the hydrogen ion concentration was increased. The calculated rate constant at 0.5" and 1.60 M ionic strength maintained with sodium perchlorate is 4.85 M-l set.-'. The values of AH* and AS* under these conditions are 9.9 kcal./mole and -23 e.u., respectively.
was obtained by reduction of a suspension of vanadium Introduction pentoxide in perchloric acid with sulfur dioxide which The kinetics of some simple oxidation-reduction was generated by treating sodium sulfite with perreactions have been extensively studied by many chloric acid. The excess sulfur dioxide was removed workers recently in an attempt to elucidate the details by bubbling carbon dioxide through the reduced soluof the electron-transfer process.l-' The reaction betion. Prepurified nitrogen (1 atm.) was passed over tween vanadium(I1) and sodium chlorate has been the aqueous vanadium perchlorate solution during the studied for comparison with the corresponding reducreduction and during the storage of the freshly protion of chlorate by uranium(IV),8 chromi~m(II),~ duced vanadium(I1). The nitrogen was bubbled and iron(I1). In the present work, the experimental through separate traps of aqueous chromium(I1) and conditions have been chosen such that an excess of vanadium(I1) to minimize the oxygen content of the vanadium(I1) is always present and that the stoichioto maintain the appropriate water vapor nitrogen and metric products appear to be vanadium(II1) and chlopressure. ride as shown below Vanadium(II1) perchlorate was obtained by mixing 6V(II) C1036Hf = equimolar quantities of vanadium(I1) and vanadium6V(III) C13Hz0 (1) (IV) under 1 atm. of purified nitrogen. The vana-
+
+
+
+
Experimental Section Analyzed reagent grade chemicals were used throughout this study. Distilled water was obtained by triple distillation from a Barnsted still and was passed through an Ilco-Way Research Model exchange column. Vanadium(I1) perchlorate solutions were prepared by t,he electrolytic reduction of vanadium(1V) perchlorate at a mercury cathode with the apparatus described elsewhere.6 The vanadium(1V) perchlorate solution The Journal of Physical Chemistry
(1) T.W.Newton and F. B. Baker, J . Phys. Chem., 69, 176 (1965). (2) T. W.Newton and F. B. Baker, Imrg. Chem., 1, 368 (1962). (3) T. W.Newton and F. B. Baker, J . Phys. Chem.,67, 1425 (1963)' (4) G. Gordon and H. Taube, J . Inorg. Nud. Chem., 16, 272 (1961). (5) G.Gordon and F. Feldman, I n o r g . Chem., 3, 1728 (1964). (6)G.Gordon and A. Andrewes, ibid., 3, 1733 (1964). (7) J. G.Mason and A. D. Eowalak, ibid., 3, 1248 (1964). (8)L. A. Fedorova and E. A. Kanevskii, Kinetika i Kataliz, 3, 332 (1962). (9) R. C. Thompson and G. Gordon, t o be published.
KINETICSOF
THE
V(II)-CIOa- REACTION IN AQUEOUSHC~OI
dium concentration was determined directly by titration with potassium permanganate. The hydrogen ion concentration of the vanadium solutions was determined by passing an aliquot of the solution through a column of the acid form of Dowex-50 resin and by titrating the effluent hydrogen ion with standard alkali. The actual hydrogen ion concentration was obtained by subtracting the appropriate number of equivalents of hydrogen ion released by the aqueous vanadium species from the total cation concentration. Sodium perchlorate was prepared by neutralizing reagent grade sodium carbonate with perchloric acid. The calculated amount of the acid required for complete neutralization was added slowly to a rapidly stirred slurry of sodium carbonate in water. The solution was filtered, acidified slightly, and boiled to expel carbon dioxide. Solid sodium perchlorate was obtained by partial crystallization, followed by three recrystallizations. The acid content of this solution was determined by titrating an aliquot directly with sodium hydroxide with phenolpthalein as the indicator. The total cation concentration was determined by passing a second aliquot through a Dowex-50 ionexchange column. The concentration of sodium ion was taken as the difference between total cation concentration and the hydrogen ion concentration. Lithium, zinc, and magnesium perchlorates were prepared in a similar fashion by the neutralization of reagent grade lithium, zinc, and magnesium carbonates. The reactions were studied by means of a modified stopped-flow apparatus which has been described in detail elsewhere.'" The change in optical density as a function of time was followed by means of a Varian Model G-14 strip-chart recorder. In a typical experiment, perchloric acid, sodium perchlorate, vanadium(I1) perchlorate directly from the storage vessel, and water in the appropriate amounts were placed into a 2-cm. cylindrical spectrometer cell under 1 atm. of nitrogen to make up a solution of the desired acidity and ionic strength. The cell was capped with a rubber cap to prevent air oxidation of the vanadium(I1). The rapid-mixing syringe was filled with a solution of sodium chlorate made up to the appropriate ionic strength with sodium perchlorate to minimize Schlieren effects upon mixing at the initiation of the reaction. The solutions of sodium perchlorate, perchloric acid, and water were carefully degassed prior to mixing in an attempt to prevent bubble formation which might tend to obscure spectrophotometric observation of the reaction, Both the syringe and the sample cell were thermostated to f0.1'. The optical spectra of aqueous solutions of vanadium(II), -(III), and -(IV) were measured, and the
201
extinction coefficients for the various species are shown in Table I. Under conditions of low acidity and in particular with vanadium(I1) solution produced by the reduction of vanadium(1V) with zinc amalgam, a brown vanadium(II1) species was observed in the reaction mixture. Hydrolyzed vanadium(II1) species have been shown to be responsible for the formation of this dinuclear vanadium(II1) cation which reacts with acid to produce the monomer.11J2 The spectrum of the complex between 3500 and 7000 8. consists of a single, unsymmetrical band with the peak at ~ 4 1 0 8. 0 The extinction coefficient for this species is estimated to be greater than 6500 M-' em.-'. With solutions which contain greater than 0.25 M perchloric acid and electrolytically produced vanadium(I1) no such hydrolyzed product was observed in the reaction mixtures. Table I: Extinction Coefficient for Aqueous Vanadium Species at 25" in 0.1 M HClOl Wave length,
~[V(III) I,
e [VOs+l,
A.
M-1 o m - 1
3750 4000 4500 5600 5900 7500 7600
0.0 0.0 0.03 2.59 5.10 17.0 17.1"
M-1
e [V(II)
om. -1
M-1
6.83 8.12" 2.31 4.83 5.64 1.20 1.28
I,
cm.-1
2.05 0.82 0.70 4.37" 3.84 1.65 1.80
" Maximum value.
For the kinetic experiments, the ?ppearance of vanadium(II1) was followed at 4000 A., and the results were analyzed by means of the integrated second-order rate expression kt =
6
W(I1)l i
- 6[ClOa-]i In
X
[( [ao~-Ii) IY(II)
Ii
[V(WIi [ClOa-Ii
-X
- X/6
]
(2)
where ii is the apparent second-order rate constant, X is the increment of reaction at any time, i denotes the initial concentration, and the square brackets denote the molar concentrations. In general, second-order plots were linear for more than 90% of the reaction, The rate constants and their standard deviations were calculated from the measured optical densities aa a (IO) R. C. Thompson and G. Gordon, J . Sci .In&., 40, 408 (1964). (11) T. W. Newton and F. B. Baker, J. Phy8. C h . ,68,228 (1964) , (12) T.W. Newton and F. B. Baker,Inorg. Chem., 3, 669 (1964).
Volunze 70,Number 1 January lQ66
GILBERTGORDON AND P. H. TEWARI
202
function of time by means of a least-square program at the University of Maryland Computer Center.l8 The standard deviation was usually less than 2% for any given kinetic experiment, and the specific rate constants for identical rum differed from each other by less than 1.8%. Each data point reported is the average of at least three replicate runs.
in a monotonic fashion as the hydrogen ion concentration is varied from 0.3 to 1.5 M at an ionic strength of 1.7 M maintained by the use of sodium perchlorate. These observations are consistent with a rate law which includes rt second term with a positive hydrogen ion dependence.
Results Sugimoto14 has studied the vanadium(II1) and perchlorate ion reaction at 50". King and Garner16have extended these results by studying the reaction of perchlorate ion with vanadium(I1) at 49.9". The second-order rate constant is 2.2 X 10" M-l set.-'. As part of this study, the vanadium(II1)-chlorate reaction was also shown to be too slow to be of any consequence. Newton and Baker1' report that the rate of oxidation of vanadium(I1) by perchlorate is
