Kinetics of the Oxidation of Aliphatic Aldehydes by Os(VII1) in Alkaline Solution

VINITA SHARMA, KAKULI CHOWDHURY, and KALYAN K. BANERJI

Department of Chemistry, University of Jodhpur, Jodhpur 342 001, India

Abstract

Kinetics of oxidation of six aliphatic aldehydes by Os(VII1) in alkaline solutions have been studied. The reaction is of first order with respect to each of the aldehyde and Os(VII1). The pseudo-first order rate constants decreased with an increase in the concentration of hydroxyl ions. The oxidation of deuterioacetaldehyde (MeCDO) exhibited a substantial primary kinetic isotope effect. Separate rate constants for the oxidation of hydrate and free aldehyde forms have been evaluated. The aldehyde hydrate is postulated as the active reductant. Ionic strength has no noticable effect on the rate. The rate-determining step is, therefore, postu- lated to be a bimolecular reaction between the aldehyde hydrate and [OSO,(OH),]~~. The value of the limiting rate constant exhibited an excellent correlation with Taft u* values; re- action constant being negative. A mechanism involving transfer of a hydride ion from the aldehyde hydrate to Os(VII1) has been proposed.

Introduction

The kinetics of oxidation of aliphatic aldehydes in acid solutions have been widely studied [l-31 but there are not many reports about the oxida- tion in alkaline solutions. The oxidations of formaldehyde by hexacyano- ferrate (111) 141 and Cu(I1) [5] in alkaline solution have been reported. Os(VII1)-catalyzed oxidation of benzaldehyde by hexacyanoferrate (111) has also been reported [6]. In view of this lack of information, we have studied the oxidation of a series of aliphatic aldehydes by Os(VII1) in alkaline solu- tion. The nature of the species of Os(VII1) present in alkaline solutions are well-defined [71. The oxidation of formaldehyde, in this system, has been studied IS].

Experimental

Materials

The aldehydes were commercial samples and were purified by the meth- ods described earlier [a ] . The isotopic purity of MeCDO (Sigma) was 98%. A stock solution of osmium tetroxide (Johnson Matthey) was prepared in 0.5 mol dm-3 sodium hydroxide by weighing and stored in a refrigerator. It was diluted as required. Sodium hydroxide solution was standardized against potassium hydrogen phthalate (BDH, AnalaR). Sodium chloride (Sarabhai M, GR) was used for adjusting the ionic strength. Doubly dis- tilled water was used through out.

International Journal of Chemical Kinetics, Vol. 22, 1039-1049 (1990) C 1990 John Wiley & Sons, Inc. CCC 0538-8066/90/101039-11$04.00

1040 SHARMA, CHOWDHURY, A N D BANERJI

Stoichiometry

Alkanals were oxidized by a known excess of Os(VII1). The residual Os(VII1) was determined spectrophotometrically at 420 nm using a caliber- ation curve. Several determinations, using different aldehydes, indicated that the reaction showed a 1 : 1 stoichiometry.

Product Analysis

The product analysis was carried out under kinetic conditions, i.e., with an excess of the alkanal over Os(VII1).

In a typical experiment, acetaldehyde (1.8 g, 0.04 mol), Os(VII1) (0.003 moll and sodium hydroxide (0.04 g, 0.01 moll were made up to 100 ml. The reaction mixture was allowed to stand for ca. 2 h to ensure completion of the reaction. Acetic acid formed in the course of the reaction would not be enough to neutralize completely the alkali added. The reac- tion mixture was acidified and extracted with ether (3 x 50 ml). The ether extract was dried (MgSO,) and treated with thionyl chloride (5 ml). The solvent was allowed to evaporate. Dry methanol (3 ml) was added and hy- drochloric acid formed was removed in a current of dry air. The residue was dissolved in dry ether (100 ml) and the ester content was determined colorimetrically as ferric hydroximate by the procedure of Hall and Schae- fer [91. The analysis showed that the yield of acetic acid was 0.16 g (89% based on the consumption of Os(VII1)).

Kinetic Measurements

Kinetic measurements were carried out under pseudo-first order condi- tions by keeping an excess (at least ~ 1 0 , generally 20 to 30 times) of the substrate over Os(VII1). The reactions were studied at a constant tempera- ture (k0.05 K). Reactions were followed on a Hi-Tech model SFL-44 stopped-flow spectrophotometer which was connected to an MCS-1 data processing system. The data were transferred to an Apple IIe PC for analy- sis and printing. The absorbance of Os(VII1) increases with an increase in the [OH-]. The A,,, of a n alkaline solution of Os(VII1) lies at about 320 nm. However, it was observed that the absorbance at 420 nm, did not vary noticably with [OH-]. In view of this and that A,,, is in UV region, the rates were followed by monitoring the decrease in the concentration of Os(VII1) at 420 nm. The application of Beer’s law at 420 nm was verified at different [OH-]. At least 5-7 runs were studied for each reaction and the pseudo-first order rate constants, k,, were reproducible within 23%. The experimental rate constant, k,,, was calculated by the relation: K,, = k,/[aldehydel.

Results

The rate data were obtained for all the aldehydes studied. Since the re- sults are similar, only the representative data are reproduced here.

KINETICS OF OXIDATION 1041

Product analysis and stoichiometry indicated the following overall reac- tion [eq. (1)l. (1) Os(VII1) + RCHO + 30H- + Os(V1) + RCOO- + 2H20

Rate Laws

The reaction is first order with respect to Os(VII1). The plots of log [Os(VIII)] against time were linear (r > 0.990). Further the pseudo-first order rate constants, k,, are independent of the initial concentration of Os(VII1). The reaction is of first order with respect to the aldehyde also (Table I). The dependence of l z , on alkali concentration was studied a t dif- ferent temperatures (Table 11). The rate decrease with an increase in the concentration of alkali. The inverse of k, is linearly related to [OH-] with a nonzero intercept (Fig. 1). The least squares values of the intercepts and slopes at different temperatures are recorded in Table 111.

Kinetic Isotope Effect

The oxidation of deuterioacetaldehyde (MeCDO) was studied to ascertain the importance of cleavage of C-H bond in the rate-determining step.

TABLE I. Dependence of the reaction rate on the concentration of Os(VII1) and alde- hyde.

103 [OS(vI I1 ) 3

no1 d r 3 mol dm-3 sec-1

0.75

1.50

3.00

4.50

6.00

7.50

0.75

0.75

0.75

0.75

0.75

0.050

0.050

0.050

0.050

0.050

0.050

0.005

0,010

0.020

0.030

0.040

204f2.1

207k1.7

20M2.5

200333.2

197f4.0

Z O Z Z .7

21. &kO .4

40.7tO. 6

82 .El. 1

124S. 3

16M3.0

[OH-] 0.8 mol dm-3, T 298 K.

1042 SHARMA, CHOWDHURY, AND BANERJI

TABLE 11. Dependence of the reaction rate on alkali concentration.

[OH-1 ki/sec-l

no1 dm-3 298 K 303 K 308 K 313 K

178f2.0 0.1 37. lfl.3 12m1.7 144fl. 3

0.2 81.5f1.4 99.31.4 125k1.8 15B1.5

0.3 69.6f l . 0 85.8fl. 1 106f2.1 13M1.9

0.4 61. @O .9 76.21.2 92. If1 .o lllf2.0

0.5 54.7f1.0 67.6t0.8 82.9t1.4 99. M2.0

0 .E 50.M0.7 €0. O f 1 .o 75.0-+1.3 90. If1 .5

0.8 41. Of0 .4 51.4f0.4 62.3f1.1 75 . 2 1 . 6

1.0 35.2fO. 6 43.9k0.7 52.e1.0 65.20.9

[MeCHOl 0.01 mol dm-3, [Os(VIII)I 7.5 x mol dm-3, I = 1.0 mol dm-3.

The results showed the presence of a large primary kinetic isotope effect (Table IV).

Effect of Ionic Strength

The dependence of k , on ionic strength was studied by varying the con- centration of sodium chloride. The rate is not affected by ionic strength (Table V). This indicates that the rate-determining step does not involve an ion-ion interaction.

Discussion

In an alkaline solution of osmium tetraoxide the following equilibria are present [51.

(2) H,OsO,(OH), + OH- 3 [HOsO,(OH),]- + H 2 0 K2

(3) [HOsO,(OH),]- + OH- I [OSO,(OH),]-~ + H 2 0

The value of K , has not been determined so far. However, it can be esti- mated from the first dissociation constant, K,, of osmic acid. The value [ 101 of K,, is 8 x The calculated value of K , is 4.4 x lo3. Hence it is ap- parent that under our reaction conditions almost whole of H,OsO,(OH), will be converted to the acid-anion, [HOsO,(OH),]- . The acid-anion reacts with hydroxyl ion to form the anion, [OSO,(OH),]-~. The value of K,, calcu- lated from the second dissociation constant [ll] of osmic acid (K2a = 3 x

is found to be 16.6 which is fair agreement with the literature value of 24 t 4. Hence we need to consider only the equilibrium (3 ) for the postu- lation of the mechanism and the choice of reactive oxidizing species is re-

KINETICS OF OXIDATION 1043

0.028

0-024

0.02 0

c Y . 0,016 c

0.01 2

0.008

0.004

0 , 0.2 0.4 0.6 0.8 1.0

[ O H - ] rnol drn-3

Figure 1. perature A 298 K; B 303 K; C 308 K; and D 313 K.

Plot of the inverse of observed rate against the concentration of alkali. Tem-

stricted to [HOsO,(OH),I- and [OSO,(OH),]-~ only. The decrease in 1 ne reaction rate with an increase in [OH-] led us to postulate [HOsO,(OH)J as the reactive oxidizing species.

The aldehydes are known to be extensively hydrated in aqueous solu- tions [ 131 and the dissociation constants for the equilibrium (4) have been reported.

Kd (4) RCH(OH), e RCHO + H,O

Two sets of rate data were calculated from the values of k,, and K,. Table VI records the experimental rate constants, k,,, and the values K,.

TABLE 111. [OH 1.

Values of intercepts and slopes of the linear plots between l /k , and

Temp. /K 298 303 308 313

lo3 Intercept 8.27f0.10 6.83t0.10 5.44f0.10 4.54f0.07

sec

102 Slope 2.01tO.02 1.6M0.02 1.34f0.02 1.OSfO.01

dm3 mol-1 sec

1044 SHARMA, CHOWDHURY, AND BANEMI

TABLE IV. Kinetic isotope effect in the oxidation of acetaldehyde.

[Aldehyde] Trpe kl/sec-’

rnol dm-3

0.005

0.010

0.030

0.010

0.020

0.030

MeCHO

MeCHO

MeCHO

M & K l

M e C W

M & K l

21.650.4

40.7f0.6

124+2.3

8.6SO . 2

17.1k0.3

26. @O .4

lpi = 4180 k 150 dm3 mol-1 sec-l

b 862 & 12 dm3 mol-1 sec-l

~ H / ~ I J 4.85 k0.24; k~ and k~ = ki/ CAdeh~del

[Os(VIII)I 6.5 X mol dm-3, [OH-] 0.8 rnol dm-3, T 298 K.

The values of kHy were calculated assuming that only the hydrated form participates in the reaction [eq. 51

(5) -d[Os(VIII)]/dt = k,,[RCH(OH)J [OS(VIII)]

Similarly the values of k, were calculated assuming that only the free aldehyde form participates in the reaction [eq. 61.

(6) - d [ Os(VIII)j/dt = k JRCHO] [ OS( VIII)]

The rates of oxidation of the aldehyde hydrate, kHy, showed an excellent correlation ( r = 0.9997) with Taft’s u* substituent constants [141 with a re- action constant p* of -1.58 * 0.01. On the other hand on such correlation exists between log k, and u* ( r = 0.0763). In particular formaldehyde was found to be much more reactive as compared to other aldehydes. No satis- factory log k, vs. u* correlation was obtained even after neglecting the rate data for the oxidation of formaldehyde ( r = 0.5883).

TABLE V. Dependence of the reaction rate on ionic strength.

I/mol dm-3 0 . 2 0.4 0 .6 0 .8 1.0

ki/sec-l 81.5-1.4 82 .31 .6 82.1k1.5 8 0 . 3 1 . 7 80.7f1.5

[OH-] 0.2 mol dm-3, [MeCHO] 0.01 mol dm-3, [Os(VIII)] 7.5 x lo-‘ rnol dm-3, T 298 K.

KINETICS OF OXIDATION 1045

TABLE VI. Rate constants for the oxidation of aliphatic aldehydes by Os(VII1).

Rate constants/dm3 mol-1 s-l Aldehyde K d

kax kA kHY

HCHO 5.5 x 10-4 1157= 231400 1157

He€HO 0.67 4070 10175 6783

EtCHO 1.4 4200

PrCHO 2.1 3274

724 1

4815

10000

10230

Pr'CHO 2.3 4230 6043 14100

ClCHzCHO 0.027 138 5308 142

ClKCHO 3.6 x 10-5 0.47 13055 0.47

"Data from ref. 8. [OH-I 0.8 rnol dm-3, T 298 K.

The existence of a good structure-reactivity correlation in the oxidation of aldehyde hydrate suggests that the aldehyde hydrates are involved in the oxidation process. Rocek and Ng 111 also observed a good correlation in the rate of oxidation of aldhyde hydrates and Taft c* values and concluded on that basis that the hydrate form participates in the chromic acid oxida- tion. Similar conclusions have drawn in several other oxidations of the aldehydes in acid solutions 121.

The presence of a substantial primary kinetic isotope effect (k,/k, =

4.85) confirms the cleavage of the aldehyde C-H bond in the rate- determining step.

The following mechanism is, therefore, proposed for the oxidation of aliphatic aldehydes by alkaline osmium (VIII). Reaction (7) was considered in view of the well-known equilibrium between an alcohol and aloxide ion in alkaline solution [151.

K2 (3) [HOs04(OH),l- + OH- [Os04(OH),1-2 + H,O

(7) RCH(OH), + OH- e RCH(0H)O- + H,O

(8) RCH(OH), + [OSO,(OH),]-~ e RCOO- + [OsO,(OH)]- + H,O

(71, and (8) is given by eq. (9).

K

k

The rate of disappearance of Os(VII1) in terms of reaction sequence (3),

(9) -d[Os(VIII)]/dt = kK,[HOsO4(OH)J [OH-] [RCH(OH),]

- - ~K,[OS(VIII)][RCH(OH),][OH-] (1 + K,[OH-])(l + K[OH-]

The values of equilibrium constant K are not known. We could, however, assume that in view of the known high value 1121 of K,, K,[OH-I 9 1 for

1046 SHARMA, CHOWDHURY, AND BANERJI

solution having [OH-] > 0.2 mol dm-3. Eq. (10) is reduced to eq. (11) which is the observed rate law of the reaction.

(12) 1 K[OH-]

k[RCH(OH),] + k[RCH(OH),] l / k , =

Equation (12) is in accord with the plots shown in Figure 1. The slope and intercept of these plots have the values equal to:

(13) Slope = K/k[RCH(OH),]

( 14) Intercept = l/k[RCH(OH),]

( 15) Slope/Intercept = K

The value of K at different temperatures were calculated from the slope and intercept of the plots and are listed in Table VII. The equilibrium con- stant, K, is found to be independent of the temperature. The values of k were determined using eqs. (13) and (14) (Table VIII). The concentration of RCH(OH), was calculated by using the value of Kd available in the litera- ture [ l l l . Similar calculations were made for all the aldehydes and the rate-limiting constant k were obtained (Table 1x1. The mean values of k have been used to evaluate the activation parameters (Table X).

The limiting rate constant k exhibited an excellent correlation with Taft’s u* substituent constants. The reaction constants, p* , have negative values (Table XI). The negative reaction constants point to an electron- deficient carbon center in the transition state.

The large kinetic isotope effect and the negative reaction constant sug- gest a transition state in which considerable positive charge is localized on carbon bearing the gem-diol. The transition state thus approaches a car- bocation in character. Therefore, the transfer of a hydride ion from the

TABLE VII. Values of equilibrium constant, K, calculated by equation (15).

K Aldehyde

298 K 303 K 308 K 313 K

NeCHO 2.43f0.03 2.37f0.04 2.46f0.04 2 . 4 2 0 . 0 5

E t a 0 2.38f0.05 2.41fO.M 2 . 4 2 0 . 0 6 2.30tO. 05

PrCHO 2.34f0.03 2.28f0.06 2.3-0.04 2.37f0.06

Pr -0 2.26?0.05 2.31f0.07 2.33f0.04 2.2W0.06

2.88f0.06 2.77f0.03 2.93f0.07 2.9MO. 07 ClcHZcHO

c 1 m o 3. lOfO .10 3. O M O . 08 3.35f0.11 3.1SO. 09

KINETICS OF OXIDATION 1047

TABLE VIII. oxidation of acetaldehyde hydrate.

Temp. I< k Mean k

Values of rate limiting constant, k, calculated by eqs. (13) and (14) for the

~

K (From eq. 13) (From eq. 143 dm-3 mol-1 s-1

29E: 20189f263 20153 133 20172230

303 26457+440 26482375 264€9+3€6

308 34642t543 34691k385 34667f425

313 4E283t923 458962475 4€08@€87

aldehyde hydrate to Os(VIII), in the rate-determining step is proposed. Reaction (8) may, hence, be visualized as reaction (15) and (16).

OH + I k (15) R- C -H + [OSO,(OH),]-~ RC(OH), + [HOSO,(OH),]-~ I

+ (16) RC(OH), + [HOSO,(OH),]-~ RCOO- + 2H20 + [OsO,(OH)]-

TABLE IX. peratures.

Mean values of k for the oxidation of the aldehyde hydrates a t different tem-

k/ dm-3 mol-1 s-1 A 1de hycie

298 K 303 K 300 K 313 K

HCHO- 3440f28 4530230 5&10?34 749m50

HeCHCi 20172f230 2646*336 34667f425 4608M687

EtCHO 28040+534 33200f519 4400@436 5 8 5 6 2 i 6 3 9

F rCHO 268002424 35425f367 460€@546 60257+524

Pr T H O 33000f483 42140f458 5370@651 72640+680

ClCHKWO 1410f21 2150f38 301@52 437Ok62

C 13CCHO 23 520 e 4 7 . 2 0 . 8 63.3f1.2 119k1.7

“based on the data from ref. 8.

1048 SHARMA, CHOWDHURY, AND BANERJI

TABLE X. Activation parameters for the oxidation of aldehydes by Os(VII1).

A1 I&1;.dc A !I* AS' AG* (at 298 K ) -~

i ~ 1 m i l l - ' J mol-1 K - 1 kJ mol-1

N 4 H C 41:. lt-i! .S -23 f 1.7 48.5f0.4

E Y H C ,?9.5+1.1 -23 * 3.7 47.3f0.9

r :.I'HC 311.3*11.2 23 +_ 0.7 47.8fO. 2

F r Ii'HCl J ' 5k1.3 -32 5 4 . 3 47.3fl. 0

5: .3? 1 0 0 ? 3 . 3 55.0+0 . 8 . l'..H;.CHC

i- 1s:r-o 7:1 Of42 44 f 13 € 5 . 1 9 . 2

,--

.. -

TABLE XI. Dependence of the reaction constant on temperature".

Tenp ., K 23e 303 3oe 313

P' 1. c m n . 04 -1.04f0.04 -1.02f0.05 -0.98f0.05

I' i l . 3373 0.3361 0.3352 0.3928

SD 0 .03 0.10 0.11 0.13

"No. of data points = 7

Acknowledgment

Thanks are due to UGC (India) and CSIR (India) for the financial sup- port to the project. Thanks are due to Dr. Raj Mehrotra for communicating the results of the oxidation of formaldehyde.

Bibliography

L11 J. Rocek and C. S. Ng, J . Org. Chem., 38, 5348 (1973). [21 A. L. J a i n and K. K. Banerji , J . Chem. Research, ( S ) , 60 (1983); (M), 678 (1983);

V. Sharma and K.K. Banerji, J . Chem. Research, (S), 340 (1985); (MI, 3351 (1985); K. K. Banerji and C. Goswami, Tetrahedron Lett., 5039 (1971).

[31 F. Freeman, D. K. Lin, and G.R. Moore, J . Org. Chem., 47, 56 (1982). 141 V.N. Singh, M. C. Gangwar, B.B. L. Saxena, and M. P. Singh, Can. J . Chem., 47, 1051

151 U. Shanker and M. P. Singh, Indian J . Chem., 6, 702 (1968). [61 P. S. Radhakrishnamurti and B. Sahu, Indian J . Chem., 17A, 93 (1979). [71 F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, Wiley Eastern, New

I81 R. N. Mehrotra, personal communication.

(1969).

Delhi, 1986, p. 1004.

KINETICS OF OXIDATION 1049

191 R.T. Hall and W.E. Schaefer in Organic Analysis, Vol. 11, J. Mitchell, Jr., Ed., Wiley- Interscience, New York, 1954, p. 55.

I101 0. M. Yast and R. J. White, J . A m . Chem. Soc., 50, 81 (1928). [11 I R. D. Sauerbrunn and E. B. Sandell, J . Am Chem. Soc., 75, 4170 (1953). [l21 D. Mohan and Y. K. Gupta, J . Chem. Soc., Dalton Trans., 1085 (1977). [131 R.P. Bell, Adu. Phys. Org. Chem., 4, 1 (1966). 1141 K. B. Wiberg, Phystcal Organic Chemistry, Wiley, New York 1963, p. 415. [I51 J. Murto in The Chemistry ofthe Hydroxyl Group, S . Patai, Ed., Wiley-Interscience, New

York, 1971, p. 1087.

Received April 28, 1989 Accepted April 24, 1990