REACTION BETWEEN SULFTJROUS ACIDAND FERRIC ION: KINETICS
April, 1963
is quite close (ca. 1/22 as large) to that found by Randles and Somerton using a.c. techniques for the same system as plat,inum.*’ While me did not measure k, in this background a t carbon paste, we do check well with the platinum data in 1 M potassium chloride. It is probably safe to conclude there is little difference between lc, values at carbon paste and untreated carbon electrodes (where these can be used). It appears then that the effect of the organic liquid in the carbon paste on charge transfer rates is relatively sma,ll. (21) J. E. B. Randles and K. 9.51 (1952).
W. Somerton, Trans. F a r a d a ~Soc.,
48,
871
One useful consequence of the slower rate at carbon paste may be that such electrodes can be used to “slow down” reactions to the region where precision r.d.e. studies are possible. A detailed treatment cf rates a t the two electrodes will be given elsewhere. Acknowledgments.-We are indebted to F. O’Brien and R. Cillmore for the design and construction of the disk electrodes and to A. Rogers for technical assistance. This work was supported by the Atomic Energy Commission through contract AT(11-1)-686 and this support is gratefully acknowledged.
THE KINETICS OF THE REACTION BETWEEN SULFUROUS ACID ASD FERRIC ION1 BY
D. G. KARR.4KER
Savannah River Laboratory, E. I . d u Pont de Nemours d% Co., Aiken, South Carolina Received October 15, 1966 The rate of oxidation of sulfurous acid to sulfate by ferric ion in perchloric acid solution was found t o agree with a kinetic route that required a thionate (HSOs) free radical as a primary reaction product. The proposed mechanism requires ( 1 ) the disproportionation of the Fe(HS03)+2complex ion to form Fei2 and HSO, (2) the reverse of realetion (I), and (3) the oxidation of the HSOa radical to sulfate by Fe+3. At 25’ and in a perchlorate medium of unit ionic strength, kl is estimated t o be 7 rnin.-l and k2/k3 = 22. From measurements of the reaction rate in the presence of oxygen, it is estimated that the scavenging of HS03radicals by oxygen is about thirtyfold faster than by ferric ion.
Introduction This study is concerned with the oxidation of sulfurous acid to sulfuric acid. Previous work indicated a probable route for this reaction, but no specific studies of the reaction have been reported. The kinetic route for production of sulfate by oxidat’ion of sulfurous acid was of interest because of the probability that the mechanism and kinetics of this reaction would be pertinent Lo similar reactions of possible cominercial interest, such as the ferric-catalyzed oxidation of SO2 by 0 2 2 and the radiat,ion-catalyzed oxidation of SOZby 0 2 to produce sulfuric acid. Sulfurous acid is oxidized by ferric ion to produce either sulfate or dithionate ions, depending on the reaction conditions. The net reactions involved are H2S03
+ 2Fe+3 + H20+
+ 4H+ + S04-2 S206-2+ 4Hf + 2Fe+* 2Fe+2
+ 2Fe+3-+
2H2S03
(A) (B)
Previous investigators3**have studied the stoichiornetry of reactions A and B under different reaction conditions, and the k:inetics of the cupric-catalyzed oxidation of sulfurous acid to dithiona,te has been studied by Higginson and M a r ~ h a l l . ~The latter investigators explained their resdts by a mechanism involving the reaction of the thionate (HS03) radical. An attempt to measure the kinetics of the uncatalyzed reaction by (1) The information contained in this article was developed during the course of work under contract AT(07-2)-1 with the U. S. Atomic Energy Commission. (2) (a) D. M. Yost and €I. Russelll, Jr., “Systematic Inorganic Chemistry of the Fifth-and-Sixth-Group h’onmetallic Elements,” Prentioe-Hall, Inc., New York, N. Y.,1944; (b) F. H. Neytsell-deWilde and L. Traverner, P r o c . U . h’. Intern. Conf. Peaceful Uses A t . Energy, Bnd, Geneva, 8 , 303 ( 1 958).
(3) H. Basaett and A. J. Henry, J . Chem. Soc., 914 (1935). (4) H. Bassett and W.G . Parker, ibid., 1540 (1951). ( 5 ) W. C. E. Higginson a,nd J. W. Marshall, ibid., 447 (1957).
Pollard, et did not lead to a satisfactory agreement with the mechanism proposed by Higginson and Marshall. Experimental Reagents.-Ferric perchlorate stock solutions were prepared by dissolving analytical grade iron wire in perchloric acid, oxidizing the ferrous ion with hydrogen peroxide, and destroying the excess peroxide by boiling. I n some experiments, the ferric perchlorate was purified by crystallization from 60% perchloric acid. Ferrous perchlorate solution was prepared by dissolving iron wire in perchloric acid. Sodium perchlorate stock solution was prepared by neutralizing standardized sodium hydroxide solution with perchloric acid. Gaseous SO2 was absorbed in water to prepare a stock solution of sulfurous acid. Perchloric acid was of reagent grade and was used without further purification. Ordinary distilled water was used in most of the experiments; triple-distilled water in the remainder. Analyses.-Conventional analytical methods were used to measure the concentrations of most of the reagents in this work: ferrous iron was determined by titration with standard ceric sulfate solution with o-phenanthroline (“ferroin”) as an indicator; ferric ion was determined colorimetrically by reduction to ferrous ion and measurement of the absorbance of the ferrous o-phenanthroline complex; sulfurous acid was determined by titration with standard iodine solution; and hydrogen ion was determined by titration with standard sodium hydroxide in the presence of sodium oxalate solution to prevent interference by iron. Dithionate was detected by a modification of the method of Glasstone and Hickling,’ which involved oxidation of dithionate ion to sulfate by boiling with standard dichromate in 6 Jri HC104, reduction of the excess dichromate with a standard ferrous sulfate solution, and measurement of the excess ferrous ion by titration with ceric sulfate solution. Oxygen was measured indirectly by +ssuming that the difference betmeen the sulfate produced by the reaction between Fe+3 and HzSOa and the total sulfate (after blank correction for reagents) was due to the reaction: 2&SO3 O2 2H2SO4. Sulfate was measured turbidimetrically by comparing the turbidity developed by samples in dilute barium chloride solutions with the turbidity developed by weighed sodium sulfate standards.
+
--f
(6) F. H. Pollard, P. Hanson, and G. Nickless, J . Chromatog., a, 68 (1961). (7) S.Glasstone and A. Hickling, J . Chem. Soc., 5 (1933).
Snalytical determinations of iron, acid, etc., were estimated to be accurate to &sfh, with a precision of about 1 2 ' ' . The turbidimetric analysis for sulfate was considered to be accurate to &lo%. Run Procedure.--~olutions of Fe( CIO4)3,HClO1, and NaV104 ~ wcre mixed in tlir desired proportions t 0 produrc t h desired acidity and iron concentrations, and an ionic strength of unity; oxygen was removed from the solutions by boiling Nitrogensaturated kerosene was poured on the surface of the boiling solution to prevent readsorption of air and, after cooling, the solution was transferred by nitrogen pressure into a volumetric flask. Sufficient deaerated water was added to provide the correct volume. .4 layer of kerosene covered the surface of the deaerated solution in all subsequent steps. These measures \%ere not always successful in removing oxygen. In many experiments a small amount of HBSO~ was added 20-30 minutes before initiating the reaction, to react with and destroy any residual oxygen. I n these experiments, the initial concentration of ferrous ion was measured before starting the reaction. Portions of this solution were pipetted into the flask that served as a reaction vessel, and were allowed to reach thermal equilibrium in a constant-temperature (25.6') bath. In the experiments with air-saturated solutions, the precautions to exclude air %ere omitted. The reaction was initiated by adding and mixing a measured amount of sulfurous acid solution, Some SO2 escaped during the addition, but after this initial loss, experiments with air-tight apparatus showed identical results compared to experiments in which no effort was made to prevent SOzloss. It was concluded that after the initial loss of SOP,there was no further loss during the reaction. The progress of the reaction was followed by withdrawing samples a t timed intervals. These samples were immediately diluted with 20 ml. of 3 M HC104and the unreacted H2SO3was removed from solution by sparging with a stream of nitrogen for 8-10 minutes. The ferrous ion produced by the reaction was determined by titration with Ce(IV) with pre-neutralized "ferroin" as an indicator. This procedure was evolved after preliminary tests showed that there was essentially no reaction in 3 N HC104, and that a 5-minute sparge with nitrogen was sufficient to remove unreacted HzS03. Data obtained in this manner showed that the concentration of Fe+2increased with time t o a constant value that was taken to be equivalent to the final concentration of sulfate produced by the reaction between Fe+3 and HzS03,and thus was equivalent to the initial concentration of HzSO~. It was determined that the difference between the final Fe+2 concentration and the Fe+z concentration at any time decreased logarithmically with time. The half-life of the reaction was determined graphically for each run with an estimated precision of & l o % . The data from a typical run are shown in Table I, and a graph of log [(Fe+2)m - (Fe+Z)+]us. time for these data is a straight line. The reaction conditions were designed t o oxidize to sulfate rather than dithionate, and therefore all reactions used a large excess of Fe+3 and low concentrations of both reactants. The absence of any substantial amount of dithionate in the reaction products was verified in many of the experiments by analysis for dithionate. It was observed experimentally that the production of dithionate was also indicated by a second component in the first-order rate curve. Data for these runs were discarded; pertinent reaction data were restricted by dithionate interference to Fe+3 concentrations below 0.09 M, and H2S03concentrations below 0.006 M . TABLE
1
TYPICAL RUK DATAFOR Fe+3OXIDATION OF &So4 Initial concn. : 0.089 M Fe+a, 0.255 M HCIOa, 0.0027 $1 H&Oa Time of reaction, min.
1
3 G 11 17 25 35
50
F e 1.2, ,?.f
43 08 73 16 82 5 20 5 33 1 2 2 4 4
5.33
x x x x x x x
x
(Fe +*)
- (Fe +*)t, M
10-3
4 9 x 10-3
10-3
10-3
10-3 10-3 10-3 10-3
10-3 10-3
3 3 x 2 6 X 1 2 x o5x 1 x 1
1
10-3 10-3 10-3 1
The results of these experiments are shown in Table 11. One run in an air-eaturated system is reported.
TABLE I1 RESULTS
EXPERIMENTAL ---------Initial Fe + g l 111
run oonditions--------€12ROi, .M H+, ,U
r e 's, -2.1
Rrartion half .life, rnin.
6.53 X 3.64 X 0.255 ... 6.4 6.53 x 10-2 5 , 5 3 x 10-3 ,255 , . . 10 6.53 x 2.67 x ,255 ... 5.5 6.53 x 1 0 - 2 3 51 x 10-3 , 2 5 5 4 . 9 4 x 10-3 10.2~ 6.53 x 10-2 3.43 x 10-3 ,255 11.4 x 10-3 13.7 6.53 x 1 0 - 2 1 . 3 5 x 10-3 ,255 ... 3.3 6.53 X 4.81 X .255 .,. 7.6 8.90 x 3.06 X ,265 ... 4.9 6.53 X 10-2 1.43 X ,255 3.9 7 . 5 0 x 10-2 2.80 x 10-3 ,255 ... 5.3 4.32 X 2.02 X ,255 ... 8.5 6.53 X 10-2 2.50 X ,255 0 . 9 X 8.2 6.53 x 1 0 - 2 3.84 x 10-3 ,255 0 . 0 x 10-3 8.4 6 . 5 3 x 10-2 1 . 6 3 x 10-3 1 . 0 0 . 1 x 10-3 18 1.3 X 17 6.53 X lo-* 2 . 6 0 x lou3 0.63 1.0 x 10.6 .5l 6.53 X 2.25 X 6 . 5 5 X 10-2 3 . 2 0 X ,255 4 . 9 X 7 .5b a Compared with the air-sa,turated experiment t,o estimate Air-saturated; the rate constant for the 02-HS03 reaction. Oz concentration = 8 X lo-* M . I
.
.
Results Derivation of the Rate Expression.--.i mechanism similar to that' proposed by Higginson and Marshall for the product'ioii of dithionate was found to be in reasonable agreement with t'he experiment'al results. This mechanism depends upon the reactions of the t,hioiiate radical (HS03) and involves t'he steps
KI H$O3 J_ H + K2
+ HSG-
Fe+3
Fe(HSOa)t2
li1 Fe(HS03)+2---f
HS03 HS03
+ FeS2
+ Fe+3 + HzO
+ + HSO3
k2 --j
Fe+3
+ HSOa-
x.3 --f
SO4+
+ Fe+2 + 3H+
The rate expression for the oxidation of HzSO3 by this series of reactions is
kz(Fe+2)(HS03) (1) From the steady-state approximation
+
Replacing (Fe+2) by the identity (Fe+2)= (Fe+2)0 2(H2S03)0- 2(H2S0,)and substituting ( 2 ) into (l),the variables in the rate expression can be separated for conditions applying t o these experiments; k.,Fe+3and H + are essentially constant for each experiment. Performing the integration and evaluating the integration constant, the followiizg expression is obtained
April, 1963
R E A C T I O N BETWEEX SULFUROUS h I I ) AITD
FERRIC IOIT : KINETICS
COhlP.4RISON O F
+
[(Fe ‘4)n 2( HzS03)0]
The term 27cz(H+)[(&803) - (H?SOa)O] in the numerator is neglected; its effect is to introduce a n additive term of increasing magnitude to the first-order rate expression. This term is zero a t the start of the reaction, so its effect on the initial slope of the first-order rate is negligible. Experimentally, no appreciable deviation from a first-order rate equation was observed. Neglecting this term, the rate expression becomes
kl ’k3(Fe +3) 2t
I;,(H+) [ ( i + + 2 ) o
+ 2 ( H 2 ~ ~ 3+) 1o c~3 ( ~ e + 3 ) ( ~ +(4)
0.209 x 10-2 3.10 x 1.54 X 10-2 3.41 X 4.97 x 1 0 - 2 0 . 7 8 X 10-2 2.77 X 1 . 7 8 X 10-2 0.82 X 10-2 1 . 6 1 X 10-2 1 . 1 7 X 10-2 1.68 X 2.46 X 3.7 X 3 . 7 x 10-2 3 . 1 X 10P2
873
TABLE I11 REswrs WITH EQUATION 5
.. ki (Fe
+J)
(€1 +)
0.87 X .87 X .87 X .87 X .87 X .87 X .87 X 1.17 X 0.87 X .98 X .56 X .87 X .87 X 3.4 X 2.1 x 1.7 X
Sum of columns 1 and 2
10-2 2.96 X 3.97 X 10-2 2 . 4 1 X 4.28 X 5.84 x 1.65 X 3.64 X 10-2 2 . 9 5 X 1.69 X 2.59 X 10-2 1 . 7 3 X 10-2 2.55 X 3.33 X 10-2 7 . 1 X 10-2 5 . 8 X 10-2 4 . 8 X
lo-*
lo-% IO-%
10-2 lod2
(10)
nrv.,
(Fe . + 3 ) 2 t l j 2
‘%
- 8 2.73 X 7 4.27 X 2.34 X - 3 2 4.36 X .. 5.85 X 1 . 4 1 X 1 0 P - 10 - 9 3.34 X 27 3.87 X - 1 1.67 X 14 2.98 X - 9 1.58 X 31 3.50 X 7 3.58 X 8 7.7 X 22 7.2 X 4 . 3 X l o M 2 -11 11 Av. dev.
where kl’
=
klK1Kz
Comparison with Euperiment.-Rate data were compared with eq. 5, obtained by substituting t t / , = t, 0.693 = In [(HzS03)/(H2S03)J] and rearranging eq. 4
h(H+)
(Fe+3)2t1,,= 0.693 7 [(Fe+”O kl k , 0.693
+ 2(HzS03)0] +
- (Fe+3)(H+) (5) 11 .
‘
The constants, k,’ and kZlk3, in (5) were evaluated graphically from data at 0.0655 Fe+3,0.255 &IfH+, and varying [(Fe+2))ot 2(H2S03)o]. Equation 5 predicts that a graph of til, us. [(Fe+Z)o 2(H,S03)0]will be a straight line and this prediction was verified. The constant Icl’ was calculated to be 1.35 min.-’ from the intercept on the y-axis, k 2 / k 1 ‘ k 3 = 16.3 min. from the slope, and thus lc2/1c~ = 22. ,4 further test of the kinetic expression was made by calculating the agreement of the data at 25’ with eq. 5 ; kl’ = 1.35 m i n - l and k 2 / k 3 = 22. The results of the comparison of the left and right sides of eq. 5 are shown in Table 111.
+
Discussion The Ferric Ion-Sulfurous Acid Reaction.-The results of this work on the oxidation of sulfurous acid by ferric ion are considered to be in good agreement with the proposed mechanism in every detail. For the acidity range studied, sulfurous acid is largely unionized, but the assumption that, the reacting species is HzS03leads to a consistent deviation from the derived rate expression of - 35% at higher acidities, compared to an average deviation of 14yoif bisulfite is assumed to be the reacting species. The assumption of a ferricbisulfite complex does not affect the rate expression, but is included because the obvious color change that occurs on addition of H,SOa to ferric ion is considered to be strong evidence for the formation of a complex. Under the conditions of this work, where ferric ion is in large excess, a 1: 1 complex is considered to be the most reasonable.
The constants determined in this work, k,‘ and h / l c 3 , may be compared to those obtained from the data of Higginson and Marshall, which are IC,’ = 0.73 niin.-l and k 2 / k 3= 0.1 a t 25O, 1.1 = 2.0, and H+ = 0.08 M , in a sulfate medium. These values do not agree with the values determined in this work, kl’ = 1.35 min.-l and 1c2/k3 = 22. Under the conditions of Higginson and AIarshall’s work, ferric ion is totally complexed by sulfate, and the major difference between their constants and those of this work is ascribed to that cause. Combining kl’ with values for K1 and Kz will lead to the rate constant, kl, for the disproportionation of the ferric-sulfite complex ion. Values for Kz are unknown, and the acceptedS value of 1.72 X lo-’ for K I becomes quite uncertain when extrapolated to an ionic strength of unity. However, from Higginson and Plarshall’s report that no visual indication of a ferric-bisulfite complex was formed in their work, Kz is much less than 9zj9the complex constant constant for Fe(SOr)+. An estimate of 7 min.-l for lcl is obtained by using K1 = 0.02, and assuming that K zis of the order of 10. Effect of Oxygen.-Reactions carried out in airsaturated solutions were more rapid than those in the absence of oxygen. This effect is due to competition by the well knowii5,6ferric-catalyzed oxidation of sulfurous acid by oxygen. A reasonable mechanism for the reaction of oxygen is k4
+ + HzO +SO*-’ + 2H+ + HOz H02 + Fe+2 + HzO +Fe(OH)+2 + H202 HzOz + Fe+*---+ Fe(OH)+2+ OH HS03 + OH --+ 2H+ + S04-2
HSOQ
0 2
This mechanism, due in essence to Haber, lo embodies reactions similar to those demonstrated for the Fricke dosimeter1, and the decomposition of HzOz by ferrous (8) H. 1’. Tartar and H. H. Garretson, J. A m . Chem Soe , 63, 808 (1941). (9) R. -4.Whiteker and N. Davidson, zhzd., 76, 3081 (1953). (10) F. Haber, Naturwzss., 19, 450 (1931), quoted in ref. 2h. (11) A. 0. Allen, “The Radiation Chemistry of Water and Aqueous Solutions,” D. Van Kostrand Co., Princeton, N, $ J , 1961.
574
Vol. 67
GUNNAR W E T T E R M A R K A S D J O H S SoUSA
ion.I2 It is chosen in preference to a mechanism proposed by B a c k s t r o n ~involving ~~ peroxysulfuric acid (HZS06) because the known properties of H2S05 do not correspond to those of the active intermediate required. However, in derivation of the rate expression, the intermediates are eliminated and the same kinetic expression results for either mechanism. In the absence of any net reaction between Fe+++ and HSO3, the rate expression can be derived to be -d In H2S03- _~ k1'(Fe+3) dt (H
+>
(
we-2)) 2kd02)
This expression can be iiitegrated when Fe.f2,Fef3, and O2are constant to yield
saturated system, the rate expression can be derived as -d lii H2S03- k1'(Fe+9 ____ I
dt
(H+) k3(Fe+3) 21c4(02) (8) k2(Fe+2) k 3 ( F e 9 4 2k4(02)
+
+
From data obtained during comparable reactions both in the presence and absence of oxygen, the ratio k 4 1 ~ can be estimated. This estimate involves (a) integrating eq. 8 assuming Fe+++, Fe++, H+, and 02 are constant, (b) equating the right side of this expression with an equivalent expression for the case where O2 is absent, and (c) eliminating common terms to obtain
(
) ~ k2(Fe+2)) ( 7 ) In H2S03 - l ~ ~ ' ( F e + 31-(H+) 2724(0?) (H2SOdo The data of n'eytzell-deWilde and Taverner,2b obtained at 10' in ea. !OW2 .U H2S04, show general agreement with this expression. These investigators found the reaction betweeii 02 and SO2 to depend directly on ferric ion, inversely on the acidity, and to be essentially independent of 0 2 , under conditions where there was a twofold excess of 02. They state that the reaction is second order in H2SO3, but careful examination of their data shows that most of their data under conditions where the reaction was rapid (til, = 25 minutes or less) fits a first-order rate expression, with kl' = 0.3. During reactions that were quite slow ( t l / 2 greater than 30 minutes), their data are no longer first order in HzS03. It may be that some other mechanism is affecting the reaction, or that experimental difficulties, such as the loss of SO*, affected their measurements. Considering all the competitive reactions in an air(12) W G Barb, et al., Trans Paradail Soc , 47, 462 (1'351). (13) IT. L~ J. Bnckstrom, J . A m Chern S o r , 49, 1480 (1927).
t,,* [1~~(I;'e+~)] kz(Fei2) l ~ ~ ( I : e + (0) ~)
+
Substituting the data from an experiment in the presence of oxygen in the left of (9) and data at the approximately equal Fe+3and H2SO3 concentrations in the absence of oxygen on the right, the equation can be solved to yield kdlc3 30, and thus k2/k4 0.7. Despite the approximate nature of these estimates, the results demonstrate that oxygen reacts with the HS03 radical much more rapidly than does Fe+3,and about as rapidly as Fe+2. It is unfortunate that this inrestigation results in the determination of ratios of rate constants, rather than the rate constants themselves. It is suggested that measurement of the free radical concentration by e.s.r. techniques in this system and other systems could lead to the determination of the individual rate constants for the individual reactions. Acknowledgment.-The author is indebted to R . S. Dorsett for performing many of the analyses.
QUANTT_" YIELDS FOR THE PHOTOCHR03IISI\l OF %(5?-XITRO-4-CYAn'OBEKZYL)-PYRIDINE BY
GUiiN-4R T$'ETTERMARK'
ASD
JOHX SOUSll
Pioneering Research Division, Quartermaster Research and Engineering Center, U. S. A r m y , h'aticb, Massachusetts Received October 17, 1982 2-(2-Nitro- .l-cyanobenzyl)-pyridine has been shown to display photochromism in EPA glass a t liquid nitrogen temperature. The molar absorptivities for the photochemically produced species have been measured; 1.3 X lo4 and 1.1 X lo4 were obtained for the absorption maxima a t 450 and 590 mk, respectively. The quantum yields for the production of the colored species were found to be low (0 032 for 254 mp and 0.014 for 365/366 r n p ) .
Introduction The long-known photochromism of 2-(2,4-dinitrobenzyl)-pyridine2 recently has been the subject of many iiivestigations, 3 . 4 The color change phellomenon call (1) National Academy of Sciences-National Research Council Visiting Scientists Reriearch Associate and Guest of the Massachusetts Institute of Technology associated with Prof. L. J. Heidt of the Department of Chemistry. (2) A. E. Tschitschibabin, B. &I. Kuindshi, and S. W. Benewolenskaja, Ber., 58, 1580 (1925). (3) R. Hardnick, H. S. Mosher, and P. Passailaigue Trans. Faraday Soc., 6 6 , 44 (1960).
be observed with crystals of the compound or in solution. During the last tlvo years it has also been found that several compounds which are closely related structurally to 2-(2,4-dinitrobenzyl)-pyridine also exhibit photochromism. Mosher, et found that the ?-isomer [4-(2,4-dinitroben2;yl)-pyridine] showed photochro(4) (a) J. Sousa and J. Weinstein, J . Om. Chern., 27, 3155 (1962); (b) G. Wettermark, J. Am. Chern. Soc., 84, 3658 (1962). (5) H. S. Nosher, C. Souers, and R. Eardwick. J . Chem P h y s . , 32, 1888 (1960).
