Indian Journal of Chemistry Vol. 47A, September 2008, pp. 1373-1376
Kinetics and mechanism of oxidation of some thioacids by morpholinium chlorochromate N Malani, S Pohani, M Baghmar & Pradeep K Sharma* Department of Chemistry, J.N.V. University, Jodhpur, 342 005, India Email:
[email protected] Received 11 April 2008; revised 20 June 2008 Oxidation of thioglycollic, thiolactic and thiomalic acids by morpholinium chlorochromate is first order both in MCC and thioacids. The reaction is catalysed by hydrogen ions. The hydrogen ion dependence takes the form kobs=a+b [H+]. Oxidation of the thiolactic acid has been studied in nineteen different organic solvents. The solvent effect has been analysed by using Kamlet’s and Swain’s multiparametric equations. A mechanism involving the formation of a thioester and its decomposition in slow step is also proposed. 8
IPC Code: Int. Cl. C07B33/00
Halochromates have been used as mild and selective oxidizing reagents in synthetic organic chemistry1-5. Morpholinium chlorochromate (MCC) is also one of such compounds used for the oxidation of benzylic alcohols6. We have already reported7-10 the kinetics and mechanistic aspects of oxidation by complexed Cr(VI) species. There are only a few reports11 on the oxidation aspects of MCC. We report herein the kinetics of the oxidation of thioglycollic (TGA), thiolactic (TLA) and thiomalic (TMA) acids by MCC in dimethyl sulphoxide (DMSO) as solvent. The mechanistic aspects are also discussed. Experimental The thioacids (Fluka) and dithiodiglycollic acid (Evan Chemicals, USA) were commercial products and used as such. Dithiodimalic and dithiodilactic acids were prepared by the oxidation of the corresponding thiols by ferric alum12. The method results in products with nearly 99% purity. The solutions of the thioacids were freshly prepared in DMSO and standardized by titrating them against a standard solution of iodine12,13. MCC was prepared by the reported method6 and its purity was checked by an iodometric method. The solvents were purified by usual methods14. p-Toluene sulphonic acid (TsOH) was used as a source of hydrogen ions.
Stoichiometric determinations, as well as the characterization of the products, carried out polarigraphically15,16 (using an automatic Heyrovsky TP 55A). It was found that the cathode wave given by a known sample of disulphide dimmer coincided with the wave given by the final product of the oxidation. The reaction exhibited a 1:2 stoichiometry, i.e. 2 moles of the thiol are oxidized per mole of MCC reduced. Further, the reaction mixtures with an excess of MCC were allowed to go to completion and the residual MCC was determined iodometrically. These results also gave a 1:2 stoichiometry. The oxidation state of chromium in completely reduced reaction mixtures, determined by an iodometric method, is 3.95±0.15. 2 R – S – H + O2CrClOMH → R – S – S – R+ H2O + OCrClOMH
… (1)
Thus, MCC undergoes a two-electron change. This is in accordance with the earlier observations with MCC11 and other halochromates7,9. It has already been shown that both pyridinium fluorochromate (PFC)17 and pyridinium chlorochromate (PCC)18 act as two electron oxidants and are reduced to chromium (IV) species. The reactions were carried out under pseudo-firstorder conditions by keeping a large excess (×15 or greater) of the thioacids over MCC. The solvent was DMSO, unless specified otherwise. The reactions were carried out in flasks blackened from the outside to prevent any photochemical reactions and were followed up to ca. 80% conversion by monitoring the decrease in the [MCC] at 370 nm on a spectrophotometer (AIMIL, India, Model MK-II). The MCC showed a λmax of 370 nm. Further, no other species had any noticeable absorbance at 370 nm. The pseudo-first-order rate constants, kobs, were evaluated from the linear least-squares plots of log [MCC] against time. Duplicate kinetic runs showed that the rate constants are reproducible to within ±3%. The second order rate constants were evaluated from the relation k2 = kobs/[reductant]. All reactions, other than those to study the effect of [H+], were performed in the absence of TsOH.
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Results and discussion Kinetic dependence — The reactions are of first order with respect to MCC. Further, the values of kobs are independent of the initial concentration of MCC. The reaction is first order with respect to thioacid also (Table 1). Induced polymerization of acrylonitrile — The oxidation of thioacids by MCC, in an atmosphere of nitrogen, failed to induce the polymerization of acrylonitrile. Further, the addition of acrylonitrile had no effect on the rate (Table 1). Thus, one-electron oxidation giving rise to free radicals is unlikely. Effect of temperature — The rates of oxidation of three thioacids were determined at different temperatures and the activation parameters were calculated (Table 2). Effect of acidity — The reaction is catalysed by hydrogen ions (Table 3). The hydrogen ion dependence has the following form kobs = a + b [H+]. The values of a and b, for TLA, are 6.11±0.02 × 10−3 s−1 and 10.7±0.04×10−3 mol−1 dm3 s−1, respectively (r2 = 0.9953). Effect of solvents — The oxidation of thiolactic acid was studied in 19 different organic solvents. The choice of solvents was limited due to the solubility of Table 1 — Rate constants for the oxidation of thioacids by MCC at 298 K 103 [MCC] (mol dm-3)
[Thioacid] (mol dm-3)
1.0 1.0 1.0 1.0 1.0 1.0 2.0 4.0 6.0 8.0 1.0
0.10 0.20 0.40 0.60 0.80 1.00 0.40 0.40 0.40 0.40 0.20
TGA
103 kobs (s-1) TLA
TMA
1.41 2.70 5.49 8.22 11.0 13.5 6.12 5.58 6.27 5.45 2.97a
6.25 12.3 25.2 37.1 49.3 61.2 26.1 24.3 27.0 24.7 12.6a
3.15 6.22 12.2 18.8 25.0 30.6 12.6 11.7 13.0 12.0 6.75a
MCC and its reaction with primary and secondary alcohols. There was no reaction with the solvents chosen. The dependence of rate on [Thioacid] and [MCC] was studied in all the solvents. The kinetics results were similar in all the solvents. The k2 values at 298 K are recorded in Table 4. The observed hydrogen-ion dependence suggests that the reaction follows two mechanistic pathways, one is acid-independent and the other is aciddependent. The acid-catalysis may well be attributed to a protonation of MCC to yield a protonated Cr(VI) species which is a stronger oxidant and electrophile (2). + [O2CrClOMH] + H+ [HOCrOClOMH] … (2) Formation of a protonated Cr(VI) species has earlier been postulated in the reactions of structurally similar quinolinium fluorochromate (QFC)7 and 2,2′bipyridinium chlorochromate (BPCC)8. Solvent effect — The rate constants of the oxidation, k2, in eighteen solvents (CS2 was not considered, as the complete range of solvent parameters was not available) were correlated in terms of the linear solvation energy relationship (LESR) of Kamlet and Taft19 Eq. (3). log k2 = A0 + pπ* + bβ + aα
… (3)
In this equation, π* represents the solvent polarity, β the hydrogen bond acceptor basicities and α is the hydrogen bond donor acidity. A0 is the intercept term. It may be mentioned here that out of the 18 solvents, 12 have a value of zero for α. The results of Table 3 — Effect of hydrogen ion concentration on the oxidation of thioacids by MCC ([MCC] = 0.001 mol dm-3; [Thioacids]=1.0 mol dm-3; Temp.=298 K)
a
contained 0.001 M acrylonitrile
Thioacid
0.10
TGA TLA TMA
15.7 71.1 35.1
103 kobs (s-1) at [H+] (mol dm3) 0.20 0.40 0.60 0.80 18.6 84.6 42.3
23.6 105 52.2
26.1 122 61.2
33.3 144 75.6
Table 2 — Rate constants and activation parameters for the oxidation of thioacid by MCC 103 k2 (dm-3 mol-1 s-1) 298 K 308 K
318 K
∆ H* (kJ mol-1)
∆ S* (J mol-1 K-1)
∆ G* (kJ mol-1)
25.2
45.0
43.5±0.5
−154±2
89.3±0.4
61.2
104
171
37.1±0.3
−163±1
85.6±0.3
30.6
53.1
94.5
40.7±0.6
−157±2
87.3±0.5
Thioacid TA
288 K
TGA
7.38
13.5
TLA
36.0
TMA
17.1
1.00 37.8 171 85.5
NOTES
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Table 4 —Effect of solvents on the oxidation of thiolactic acid by MCC at 298 K Solvents
103 k2 (dm mol-1 s-1)
103 k2 (dm mol-1 s-1)
Solvents
-3
Chloroform 1,2-Dichloroehane Dichloromethane DMSO Acetone Dimethyl formamide Butanone Nitrobenzene Benzene Cyclohexane
-3
14.8 18.6 20.0 61.2 16.2 33.9
Toluene Acetophenone Tetrahydrofuran t-butyl alcohol 1,4-Dioxane 1,2-Dimethoxyethane
5.25 23.9 9.77 5.62 10.2 4.47
12.0 24.5 6.31 0.63
Carbon disulfide Acetic acid Ethyl acetate
2.82 2.29 6.46
correlation analyses in terms of Eq. (3), a biparametric equation involving π* and β, and seperately with π* and β can be given as Eqs (4) - (7).
solvent polarity. The rates in different solvents were analysed in terms of Eq.(8), separately with A and B and with (A + B).
log k2 = − 4.11 + 1.65 (±0.20) π* + 0.19 (±0.17) β + 0.27 (±0.16) α … (4)
log k2 = 0.46 + (±0.04) A + 1.79 (±0.03) B − 4.32 … (9)
R2 = 0.8680; sd = 0.18; n = 18; ψ = 0.40 log k2 = − 4.17 + 1.74 (±0.20) π* + 0.10 (±0.17) β … (5)
R2 = 0.9962; sd = 0.03; n = 19; ψ = 0.07 log k2 = 0.20 (±0.29) A − 3.09 r = 0.0070; sd = 0.48; n = 19; ψ = 0.91 log k2 = 1.76 (±0.09) B − 4.17
R2 = 0.8413; sd = 0.19; n = 18; ψ = 0.42
… (10)
2
… (11)
2
log k2 = − 4.15 + 1.77 (±0.19) π*
r = 0.9614; sd = 0.09; n = 19; ψ = 0.20 … (6)
r2 = 0.8375; sd = 0.19; n = 18; ψ = 0.41 log k2 = − 3.15 + 0.41 (±0.38) β
… (12)
2
r = 0.7840; sd = 0.22; n = 19; ψ = 0.48 … (7)
2
r = 0.0676; sd = 0.46; n = 18; ψ = 0.99 Here, n is the number of data points and is the Exner's statistical parameter20. Kamlet's19 triparametric equation explains ca. 87% of the effect of solvent on the oxidation. However, by Exner's criterion20, the correlation is not even satisfactory (cf. Eq. 4). The major contribution is of solvent polarity. It alone accounted for ca. 84% of the data. Both β and α play relatively minor roles. The data on the solvent effect were analysed in terms of Swain's21 Eq. (8) of cation- and anionsolvating concept of the solvents also. log k2 = a A + b B + C
log k2 = 1.35 ± 0.17 (A + B) − 4.27
… (8)
Here, A represents the anion-solvating power of the solvent and B the cation-solvating power. C is the intercept term. (A + B) is postulated to represent the
The rates of oxidation of TLA in different solvents showed an excellent correlation in Swain's equation (cf. Eq. 9) with the cation-solvating power playing the major role. In fact, the cation-solvation alone account for ca. 96% of the data. The correlation with the anion-solvating power was very poor. The solvent polarity, represented by (A + B), also accounted for ca. 78% of the data. In view of the fact that solvent polarity is able to account for ca. 78% of the data, an attempt was made to correlate the rate with the relative permittivity of the solvent. However, a plot of log k2 against the inverse of the relative permittivity is not linear (r2 = 0.5300; sd = 0.33; ψ = 0.70). Mechanism
The lack of any effect of radical scavenger, such as acrylonitrile on the reaction rate and the failure to induce the polymerisation of acrylonitrile, point against the operation of a one-electron oxidation
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giving-rise to free radicals. The observed solvent effect supports a transition state, which is more polar than the reactant state. Therefore, a mechanism involving a direct hydride ion transfer from S – H group to MCC is suggested (Eqs 13 – 15). slow R − S − H + OCrOClOMH → + − R − S + [HOCrOClOMH]
+
fast
… (13) +
R−S + R − S − H → R − S − S − R + H … (14) fast H2O + CrOClOMH [HOCrOClOMH]− + H+ → (CrIV) … (15)
The above mechanism is supported by the observed activation parameters. Bimolecular reactions usually exhibit negative entropy of activation. As the activated complex is formed, the reactants lose their freedom to move independently. Further, as the charge separation takes place in the transition state, the charged ends get solvated by solvent molecules. This results in immobilization of large number solvent molecules. This also results in unfavourable entropy. The observed large negative entropy of activation thus supports the above mechanism. The formation of a sulphenium cation, in the ratedetermining step, is supported by the observed major role of cation-solvating power of the solvents. The proposed mechanism is supported by the relative reactivity also. The observed reactivity is TLA > TMA > TGA. TLA has an electronic donating methyl group, which makes the transfer of a hydride ion easy as compared to that in TGA. In TMA, the presence of a carboxylic acid reduces the electrondonating tendency of the methyl group, reflected in a lower rate. In the presence of hydrogen ions the reactive oxidizing species will be MCCH+. The above mechanism leads to the following rate law, which is identical to the experimental observations. Rate = k2 [Thioacid] [MCC] + k3 [Thioacid] [MCC] [H+] Although, there is no kinetic evidence for the formation of thioester, its formation in small amounts cannot be ruled out. It is of interest to compare here the reaction patterns of the oxidation of thiocaids by PFC22, QFC7, BPCC23 and MCC. PFC and QFC represented a Michaelis-Menten type of kinetics with respect to
thioacid, whereas the oxidation by BPCC and MCC exhibited a second order kinetics, first with respect to each reactant. This may be due to a very low value of the formation constant of the thioester. The solvent effect and hydrogen ion dependence are parallel in all the reactions. Acknowledgement Thanks are due to the UGC New Delhi, India for financial (Major Research project No. F. 31-135/2005 (SR) Dated 31.03.2005), and to Prof. K K Banerji for helpful discussions and suggestions. References 1 Corey E J & Suggs W J, Tetrahedron Lett, (1975) 2647. 2 Guziec F S & Luzio F A, Synthesis, (1980) 691. 3 Bhattacharjee M N, Choudhuri M K, Dasgupta H S, Roy N & Khathing D T, Synthesis, (1982) 588. 4 Balasubramanian K & Prathiba V, Indian J Chem, 25B (1986) 326. 5 Pandurangan A, Murugesan V & Palanichamy M, J Indian Chem Soc, 72 (1995) 479. 6 Shiekh H N, Sharma M, Hussain M & Kalsotra B L, Oxid Commun, 28 (2005) 887. 7 Khurana M, Sharma P K & Banerji K K, Indian J Chem, 37A (1998) 1011. 8 Kumbhat V, Sharma P K & Banerji K K, Indian J Chem, 39A (2000) 1169. 9 Chouhan K & Sharma P K, Indian J Chem, 43A (2004) 1434. 10 Pohani P, Anjana & Sharma P K, Indian J Chem, 45A, (2006) 2218. 11 Bishnoi G, Malani N, Sindal R S & Sharma P K, Oxid Commun, 30(3) (2007) 607; Bishnoi G, Sharma M, Sindal R S & Sharma P K, J Indian Chem Soc, 84 (2007) 892. 12 Leussing D L & Kolthoff I M, J Electrochem Soc, 100 (1953) 334. 13 Krammer H, J Assoc Agric Chem, 35 (1952) 385. 14 Perrin D D, Armarego W L & Perrin D R, Purification of Organic Compounds (Pergamon Press, Oxford), 1966. 15 Kapoor R C, Kachhwaha O P & Sinha B P, J Phys Chem, 73 (1969) 1627. 16 Kapoor R C, Kachhwaha O P & Sinha B P, Indian J Chem, 10 (1971) 499. 17 Brown H C, Rao G C & Kulkarni S U, J Org Chem, 44 (1979) 2809. 18 Bhattacharjee M N, Choudhuri M K & Purakayastha S, Tetrahedron, 43 (1987) 5389. 19 Kamlet M J, Abboud J L M, Abraham M H & Taft R W, J Org Chem, 48 (1983) 2877 and references cited therein. 20 Exner O, Collect Czech Chem Commun, 31 (1966) 3222. 21 Swain C G, Swain M S, Powel A L & Alunni S, J Am Chem Soc, 105 (1983) 502. 22 Agarwal S, Choudhary K & Banerji K K, Trans Met Chem, 16 (1991) 661. 23 Rathore S, Sharma P K & Banerji K K, J Chem Res, (S) (1994) 298.
