Kinetics of the reaction of methyl iodide with sulfite and thiosulfate ions in aqueous solution1 R. A. HA STY^ AND S. L. SUTTER Pacific Northwest Laboratory, Battelle Memorial Institute, Richland, Waslrington 99352 Received June 30, 1969 The rate of reaction of methyl iodide with sulfite ion is determined. In addition, the rate of reaction of methyl iodide with thiosulfate ion is reexamined and the rate of reaction of methyl iodide with bisulfite ion is estimated. A pronounced effect of ionic strength on the reaction rate in the methyl iodide - sulfite ion system is observed, this effect does not occur in the methyl iodide - thiosulfate ion system. The second order reaction rate constant and activation energy for the reaction of methyl iodide with the respective nucleophiles are: SOSZ-, 4.4 x los2 M-' s-', 18.6 kcal mole-'; HS03-, 1 x M-' s-', 18.4 kcal mole-'; and SZ03Z-, 3.1 x M-l S-I, 19.4 kcal mole-'. Canadian Journal of Chemistry, 47, 4537 (1969) Introduction The reaction of methyl iodide with nucleophiles such as sulfite, bisulfite, sulfate, or thiosulfate ions presents an opportunity to compare and correlate the reactivity with non-kinetic properties of these nucleophiles. The reaction of methyl iodide is the simplest of the nucleophilic displacement reactions of the alkyl halides and should be the most easily understood. Excellent correlation of the reaction rate constant with the oxidative dimerization potential (1) is obtained with a number of nucleophiles (2-4). Fluoride ion is the only exception to the correlation. The reaction rate of methyl iodide with thiosulfate ion has been reported (3,4) and is reexamined in this work. The determination of the rates of reaction of sulfite, bisulfite, and thiosulfate ions with methyl iodide are reported in this paper. In addition, the effect of ionic strength on the reaction rate is determined for both the thiosulfate ion and sulfite ion reaction with methyl iodide. The comparison of the two reactions presents an extreme contrast in two reactions which in most aspects are very similar Experimental The gas chromatographic system utilized in this study has been described (5). The preparation of methyl iodide solutions and the details of the kinetic measurements are also reported. The only unusual feature of the chromatographic separation and detection of the gas which exits from the reaction vessel containing acidic solutions 'This paper is based on work performed under United States Atomic Energy Commission Contract AT(45-1)- 1830. 2Present addrcss: Department of Chemistry, Montana State University, Bozeman, Montana. (ph 4.5) from observations reported earlier is the appearance of the chromatographic peak due to SOZ. The elution of SO, from the chromatographic column occurs after the elution of 0, and before the elution of CH31. A small caustic scrubber removes SOZ from the gas stream. Methyl iodide is not removed by this scrubber. Sodium sulfite, sodium bisulfite, and sodium thiosulfate, reagent grade, are used without purification to prepare the reaction solution. The total sulfite or thiosulfate concentrations of the reaction solutions are determined by titration with potassium iodate standard solutions in the presence of carbon tetrachloride. The disappearance of iodine from the carbon tetrachloride layer is taken to be the end point of the titration. The determination of the rate of reaction of rncthyl iodide in very dilute solutions (ca. M) has been described (5). The rate of disappearance of methyl iodide from the reaction solution, k,,,,, and the total amount of methyl iodide which is sparged from the reaction vessel are determined for at least four flow rates of the sparge gas. The rate of reaction of methyl iodide in the solution is obtained from a graphical representation of (I/k,,,,) vs. the amount of methyl iodide recovered from the reaction vessel. Results Preliminary experiments at 24.70 "C, a ph of 12.05, and 0.020, 0.040, and 0.080 M sulfite ion gave0.84 x 1.65 x and 3.3 x s-' for the pseudo first order reaction of methyl iodide. This demonstrates that the rate expression is first order with respect to the sulfite ion concentration. Although these measurements are not obtained at constant ionic strength (p = 0.11, 0.17, and 0.29, respectively), subsequent measurements at constant ionic strength verify the conclusion. The lack of an ionic strength effect in these preliminary results is attributed to the small variation of ionic strength coupled with the experimental uncertainties. The rate expression is also first order with respect to the methyl iodide
CANADIAN JOURNAL OF CHEMISTRY. VOL. 47. 1969 TABLE I Rate constant of methyl iodide - sulfite and - bisulfite ion reaction at 24.73 t 0.03 "C, and ionic strength and ph dependence -- ---- Total so32- [NaCl] [N~zSO~I HSO3- kz 5 lo2 ph (M) (M) (M) N (M- s-i) 'NaCI formed by neutralization of NaOH with HCI. testirnate of rate constant only, higher HS03- concentration needed (see next row). concentration, as demonstrated by the exponential decrease of methyl iodide concentration from the reaction vessel. The effects of ph, chloride ion concentration, and ionic strength on the reaction rate are determined. The results of these measurements are given in Table I. Hydrochloric acid is added in some ex.periments to partially neutralize the sodium hydroxide and obtain solutions of a desired ph. In order to determine the effect of chloride ion on the rate of reaction, the chloride ion concentration is varied. The second order reaction rate constant is reduced ca. 30% in these experiments. Since this reduction in the observed rate constant could be due either to the increase in chloride ion concentration or to the increased ionic strength of the solution, sodium sulfate is added to adjust the ionic strength of the solutions. These results demonstrate that the rate constant is decreased as the ionic strength of the solution is increased. They are displayed in Fig. 1. The temperature dependence of the sulfite and bisulfite reactions with methyl iodide are determined from the measurement of the reaction rate of methyl iodide at 33 and 49 O C and are reported in Table 11. The rate of reaction of methyl iodide with sodium thiosulfate is determined at constant and variable ionic strength. The results of these measurements are reported in Table 111. Discussion Thiosulfate, sulfite, and bisulfite ions react with methyl iodide by a simple displacement mechanism (S,2). The kinetics of these processes are straight forward as are most displacement reac- I 0.2 0.4 0.6 0.8 Ionic strencth, M FIG. 1. Reaction rate constant for the methyl iodide - sulfite ion reaction as afunction of ionic strength, A, Na2S04 added; 0, NaCl and NaCl-Na2S0, mixtures added. I
HASTY AND SUTTER: REACTION OF METHYL IODIDE WITH SULFITE AND THIOSULFATE IONS 4539 TABLE I1 Rate constant of methyl iodide - sulfite ion reaction at a ph of 12.1 t 0.1, temperature dependence - Z[S032-] k,x103 k27102 T (M) ( s ) (M-s-') ("c) P TABLE I11 Rate constant of methyl iodide - sodium thiosulfate reaction in 0.052 N sodium hydroxide solution -- --- [SZO~~-I Ionic k, x lo3 kz x lo2 (M) strength* (s-l) (M-I s-l) Temperature, 24.70 k 0.02 "C 0.0996 0.65 3.16 3.2 0.0505 0.0254 0.65 0.65 1.51 0.76 3.O 3.O 0.0103 0.65 0.31 3.0 0.0498 0.20 1.51 3.O 0.0500 0.35 1.58 3.2 0.05047 0.15 1.70 3.4 0.03011 0.165 0.96 3.2 Temperature, 33.28 k 0.02 "C 0.0251 0.20 2.0 8.0 0.03705 0.20 2.71 7.3 Temperature, 49.04 + 0.04 "C 0.00489 0.20 1.95 40.0 0.00292 0.20 0.99 34.0 0.00739 0.20 2.96 40.0 "Ionic strength adjusted by the addition of sodium sulfate. fph of 9.0 compared to 12.2 for 0.052 N sodium hydroxide solutions. tions which involve methyl iodide (1, 3, 6). The only complication is the hydrolysis by hydroxide C2] - dt ion in the case of the slower reactions. The rate of reaction can be generalized in the form d [CH31] [CH311 = k2'[ch3i][n](x) f k2"[ch31][~](1 - X) where [N] is the total sulfite-bisulfite ion con- [I] -,, centration and x is the sulfite ion fraction of the where N represents the nucleophile. In the case of the sulfite ion as the nucleophile, the sulfite-bisulfite equilibria profoundly affects the observed rate of reaction of methyl iodide as shown in Table I. In the case of the sulfitebisulfite equilibria, the term which involves the nucleophile in eq. [I] is separated into two parts and the terz which involves hydroxide is neglected. total. In theory, the fraction of sulfite ion can be calculated from the dissociation constant (6) of bisulfite ion of 5.6 x lo-* but, in practice, the ionic strength is too high to permit reliable estimation of the activity coefficients of bisulfite or sulfite ions. Reasonableagreement between the observed decrease in reaction rate and the fraction of sulfite ion is obtained when values of the activity coefficient calculated by the formula cited by Conway (7) are employed. Although the bimolecular reaction rate constant for sulfite ion with methyl iodide can be
4540 CANADIAN JOURNAL OF CHEMISTRY. VOL. 47, 1969 determined, the bimolecular reaction rate constant for bisulfite ion with methyl iodide can only be approximated. At a ph of 4.5, the sulfite ion composes approximately 1% of the total bisulfitesulfite concentration when the equilibrium constant is considered and activity coefficients are estimated as before. Thus, the bimolecular reaction rate constant for bisulfite ion with methyl iodide of0.064 M-' min-' is only an upper limit. It should be noted that this value is close to that predicted by the Edwards correlation (1). The high ionic strength of the bisulfite solutions which is required because of the low reactivity of bisulfite ion with methyl iodide also requires that caution be used when stating this reaction rate constant. In the case of thiosulfate ion, the dissociation constant of bithiosulfate ion (6) is lo-'. The bithiosulfate ion is less than 1% of the thiosulfate concentration at ph greater than 4. The effect of ionic strength is first noted in the studies which involved the dependence of reaction rate of sulfite-bisulfite with methyl iodide upon the ph. Extension of the ionic strength studies is made for both the sulfite and thiosulfite reaction with methyl iodide at high ph to preclude the formation of other species. The two reactions, i.e., sulfite - methyl iodide and thiosulfate-methyl iodide, although quite similar (both arc reactions of dinegatively-charged ion with a dipole reaction and both have the same order-of-magnitude reaction rate constant), exhibit extremely contrasting dependence upon ionic strength. The thiosulfate - methyl iodide reaction rate constant is (3.1 k 0.1) x lo-' M-' s-' at 24.70 "C which is in excellent agreement with previously reported values (3, 4) at much higher methyl iodide concentrations (lo-' M compared to lo-' im in this work). The reaction rate constant is independent of ionic strength up to an ionic strength of 0.65. In contrast, the sulfite ion - methyl iodide reaction rate constant is very strongly dependent upon ionic strength of the solution as shown in Fig. 1. The Hiickel extension of the Debye- Hiickel treatment for activity coefficients (ref. 8, p. 229) to include a second term bp, where b is a constant which is related to the dielectric saturation of the medium and p is the ionic strength of the medium, is an attempt to explain the effect of ionic strength on ion-dipole reactions. The relationship can then be derived that [3 1 log k p = log ko + A bp where k is the reaction rate constant (at ionic strength p or zero) and A b is the difference of the constant, b, between the transition state and the reactants. For A b much less than unity theexpression can be written as [4 1 kp = ko + A bp Normally Ab is small and there is little or no dependence of the rate constant of an ion-dipole reaction on the ionic strength as in the case of the thiosulfate - methyl iodide reaction. In the case of the sulfite - methyl iodide reaction, Ab is sufficiently large to cause the decrease of reaction rate constant for the sulfite -methyl iodide reaction to be nearly as large as in ion-ion reactions (ref. 8, p. 219). An alternative quantitative treatment such as proposed by Amis (9) results in too small a predicted decrease of the reaction rate constant for the sulfite - methyl iodide reaction (10% decrease going from ionic strength of 0 to 1 as compared to ca. 40% decrease observed). Perhaps the principal reason for the poor quantitative correlation is the inability to accurately estimate values of the dielectric constant and distance of closest approach of the ion and dipole. Extrapolation of the rate constant presented in Fig. 1 in accord with eq. [4] gives the bimolecu- lar reaction rate constant at zero ionic strength and 24.7 "C to be4.4 x M-Is-'. The activation energies of the reaction are summarized in Table IV. For the thiosulfate - methyl iodide reaction the activation energy of 19.44 f 0.89 kcal mole-' is somewhat higher -- Nucleophile TABLE IV Summary of Arrhenius parameters for various nucleophiles k, at 25 "C log10 A Ea (mole 1-' s-') (mole 1-' s-') (kcal mole-') SOj2-4.4 x lo-= 12.3 18.6k0.6 HSOj- 1 x lo4 10.5-18.4 S20,,- (3.1 +. 0.1) x 12.75 19.4k0.9
HASTY AND SUTTER: REACTION OF METHYL IODIDE WITH SULFITE AND THIOSULFATE IONS 4541 but in agreement with 18.88 kcal mole-' previously reported. The activation energy of the sulfite - methyl iodide reaction is 18.6 + 0.6 kcal mole- '. The activation energy of the bisulfite - methyl iodide reaction is ca. 18.4 kcal mole-' but, because of the uncertainty of the sulfite ion contribution, this value is only an approximation. 1. J. 0. EDWARDS. J. Amer. Chem. Soc. 76, 1540 (1954). 2. E. A. MOELWYN-HUGHES. Proc. Roy. Soc. Ser. A, 220, 386 (1953). R. H. BATHGATE and E. A. MOEL- WYN-HUGHES.. J. Chem. Soc. 2647 (1938). E. A. MOELWYN-HUGHES. J. Chem. Soc. 779(1959). B. W. MARSHALL and E. A. MOELWYN-HUGHES. J. Chem. Soc. 7119 (1965). R. A. OGG. J. Amer. Chem. Soc. 60, 2000 (1938). G. C. LALOR and E. A. MOELWYN- HUGHES. J. Chem. Soc. 2201 (1965). 3. A. SLATOR. J. Chem. Soc. 85, 1286 (1904). 4. E. A. MOELWYN-HUGHES. Proc. Roy. Soc. Ser. A, 196, 540 (1949). 5. R. A. HASTY. J. Chem. Phys. 73, 317 (1969). 6. L. G. SILLEN and A. E. MARTELL. Stability constants of metal-ion complexes. Special Publication No. 17, Chem. Soc. London. 1964. 7. B. E. CONWAY. Electrochemical data. Elsevier Publishing Co., New York. 1952. p. 102. 8. K. J. LAIDLER. Chemical kinetics. 2nd ed. McGraw- Hill Book Co., Inc., New York. 1968. 9. E. S. AMIS. Kinetics of chemical change in solution. The MacMillan Company, New York. N.Y. 1949. pp. 165-174.