1044
Singh et ai.
Kinetics and Mechanism of the Ruthenium(ll1) Chloride Catalyzed Oxidation of Butan-2-01 and 2-Methyl-I-propanol by the Hexacyanoferrate(ll1) Ion in an Aqueous Alkaline Medium H. S. Singh," R. K. Singh, S. M. Singh, and A. K. Sisodla Chemical Laboratories. University of Allahabad, Allahabad, India (Received August 17, 1976; Revlsed Manuscript Received March 4, 1977)
Studies of the kinetics of ruthenium(II1)chloride catalyzed oxidation of butan-2-01and 2-methyl-1-propanol by aqueous alkaline potassium hexacyanoferrate(II1)were made at constant ionic strength. The results show that complex formation between the ruthenium(II1)species and the alcohol molecule takes place with the possible exchange of a hydroxide ion. The complex thus formed undergoes disproportionation yielding the carbonium ion derived from the alcohol molecule as an intermediateand the ruthenium(1II)hydride species. The carbonium ion rapidly goes to an aldehyde (or ketone with secondary alcohol) with one molecule of hydroxide ion. The Ru(II1) hydride thus produced is oxidized rapidly to Ru(II1) with the hexacyanoferrate(II1)ion. The oxidation products are confirmed and a probable reaction path is given.
Introduction There have been only very few investigations concerning the ruthenium tetroxide or ruthenium trichloride catalyzed oxidation of organic compounds with some cooxidants in acid or alkaline medium. Recently, it has been reported that the oxidation of a diols under alkaline condition gives extensive carbon-carbon bond cleavage, Le., adipic acid was obtained in 80-90% yields from the oxidation of cyclohexane-1,2-diolusing a catalytic amount of ruthenium trichloride along with sodium hypochlorite as a cooxidant in an aqueous medium.' However, the kinetics of the problem are still unknown. Thus, this study was undertaken to investigate the kinetic features of the ruthenium trichloride catalyzed oxidation of butan-2-01 and 2-methyl-1-propanol by the hexacyanoferrate(II1) ion in an aqueous alkaline medium. Accordingly, the details of the results are presented and a probable reaction mechanism has been proposed.
Results and Discussion The reduction of hexacyanoferrate(II1) by butan-2-01 and 2-methyl-1-propanolin the presence of ruthenium(II1) chloride as a homogeneous catalyst was investigated at varying initial concentration of hexacyanoferrate(1II). The details of the kinetic data are presented in Tables 1-111. The standard zero-order rate constants (k,)presented in these tables are the averages from a particular run. A typical zero-order plot for the rate of oxidation of, butan-2-01 and 2-methyl-1-propanol is shown in Figure 1. It is obvious from these plots that the reaction velocity follows zero-order kinetics even up to the end of the reaction, A close examination of Table I clearly indicates that the k, values are practically constant indicating zero-order kinetics with respect to the oxidant. The results presented in Table IIA and B contain the standard zero-order rate constants at varying concentrations of organic substrates, and from k,/ [alcohol] values it is quite obvious that the k, values obtained for molar concentration of organic substrates continue to decrease with its increasing concentration. Thus the results presented in Figure 2 indicate the direct proportionality of the reaction rate on the low alcohol concentration and it tends toward zero order at higher concentrations. The oxidation kinetics of these alcohols have been found to depend on the hydroxide ion concentration. A plot of k, values against hydroxide ion concentration showed that The Journal of Physical Chemistry, Vol. 81, No. 11, 1977
TABLE I: Effect of Hexacyanoferrate(II1) on the Rate of Oxidation
A [ 2-Methyl-1-propanol] = 2.0 X
M, [NaOH] = 2.0 X l o - ' M, [Ru(III)]= 1.20 X M [K,Fe(CN),]X l o 3 M = 1.0, 2.0, 3.0, 4.0, 5.0 h, X lo5 M min-l = 2.65, 2.81, 3.18, 3.28, 3.07 B [Butan-2-01]= 5.0 X l o - ' M, [NaOH]= 2.0 X l o - ' M, [Ru(III)] = 1.20 X M [K,Fe(CN),] X l o 3 M = 0.50, 1.0, 2.0, 3.0, 4.0, 5.0 h,X l o 5 M min-' = 1.53, 1.50, 1.58, 1.85, 1.91, 1.77 TABLE 11: Effect of Variation of the Alcohol Concentration on the Rate of Oxidation
A [K,Fe(CN),]= 2.0 X l o - , M, [NaOH]= 8.0 x 10" M, [Ru(III)]= 1.20 X M [2-Methyl-l-propanol]X 10' M = 1.0, 2.0, 3.0, 4.0, 5.0, 6.0, 8.0, 10.0 k , x l o 5 M min-' = 0.77, 1.33, 1.88, 2.30, 2.71, 3.03, 3.94, 4.64 {h,/[2-methyl-l-propanoll)X l o 3 min" = 0.77, 0.66, 0.62, 0.57, 0.54, 0.50, 0.49, 0.46 B [K,Fe(CN),]= 2.0 X l o - , M, [NaOH]= 2.0 x l o - ' M, [Ru(III)]= 4.80 X M [Butan-2-ol]X 10' M = 1.0, 2.0, 3.0, 5.0, 6.0, 7.0, 8.0, 10.0 h,X l o 5 M min-' = 2.47, 4.59, 6.60, 9.70, 11.85, 12.59, 13.91, 16.12 {hS/[butan-2-oll)X l o 3 rnin-l = 2.47, 2.29, 2.29, 1.94, 1.97, 1.79, 1.74, 1.61
the k, values decrease with its increasing concentration, i.e., itmggests an approximate inverse proportionality of the reaction rate on hydroxide ion concentration (Figure 3). Table IIIA and B presents the effect of the variation of ruthenium(II1) chloride concentration on the reaction rate. Again, the values of k, calculated for molar concentration of the catalyst are fairly constant for about a tenfold variation, so it is concluded that the order with respect to ruthenium(II1) chloride is unity. Now on the basis of these results the following rate expression is proposed: - d[ Fe(CN)J 3 - - h,[ruthenium(III)] [alcohol] (1) dt h, [OH-] + k,[alcohol]
1045
Ru Catalyzed Oxidation of Alcohols
TABLE 111: Effect of Variation of the Ru(1II) Concentration on the Rate of Oxidation A [K,Fe(CN),]= 2.0 X l o - , M,[2-methyl-l-propanol]= 2.0 X lo-' M,[NaOH]= 8.0 X lo-' M [Ru(III)]X l o 6 M = 0.48, 0.96, 1.44, 1.92, 2.40, 3.12, 3.60, 4.32, 4.80 k,X l o 5 M min-' = 0.53, 1.02, 1.42, 1.73. 2.62, 3.40, 3.64, 4.27, 4.92 {k,/[Ru(III)]} X l o - ' rnin-l = 1.10, 1.06, 0.98, 0.90, 1.09, 1.08, 1.01, 0.98. 1.02 B [K,Fe(CN), ] = 2.0 X l o - ' M,[butan-2-01]= 2.0 X 10" M,[NaOH]= 2.0 X lo-' M [Ru(III)] X l o 6 M = 0.48, 0.96, 1.20, 1.44, 1.92, 2.40, 3.60, 4.80 h, X lo5 M min-' = 0.86, 1.61, 2.00, 2.33, 3.01, 3.62, 5.21, 7.29 {k,/[Ru(III)l) X l o - ' min-l = 1.79, 1.67, 1.66, 1.61, 1.61, 1.50, 1.44, 1.51
20.0-
18.0 -
16.0 -
14.0 -
-15 12.0E -'2
-
lobo-
91
z 8.0 -
h
X
42 (A)-Isobufanol
6.0
Temp 30'
[ K 3 Fe(CN)6]=2
OxlC3M
[isobutanol J=I.OXI~' M [NaOH ] =8,0xi6'M
4.0 -
[RUC13] = 1 . 2 O ~ t 5 ~ M p = 0.5 M
.-
2.0 -
(8) -Butan-z-ol Temp.doo rks Fe(cN),]=2.ox 1z3hf
1.6
-
[Butan-2-ol] =3.0Xl~?M
E
0
[NaOH ] =4.0~16*M [RUC13] = 4.8 x iE6 M # =05M
1.2
1.0
2.0
3.0
2.0
4.0
6.0
4.0 8.0
5.0 (E) 10.0 ( A )
[OH-]X102M
Flgure 3.
Time in m i n .
Flgure 1.
12.01
VI
8.0 E
3.2 6.01
/
[Organic ~ubstrafe] M Figure 2.
where k,, kb, and k, are constants. Before presenting the probable oxidation scheme, it is interesting to know the probable species of ruthenium(II1) chloride in the alkaline medium. Electronic spectra studies have confirmed that the ruthenium(II1) chloride exists in the hydrated2 form as [RU(H@)~]~+. Metal ionsaaof the form [ R U ( H ~ F ) ~are ] ~ +also known to exist as [Ru(H,O),(OH)] in an alkaline medium. In the present study it is quite probable that the species [RU(H~O),OH]~+ might assume the general form [Ru(III)(OH)x]3-x.The value of x would always be less than six because there are no definite reports of any hexahydroxy species3bof ruthenium. The remainder of the coordination sphere will be filled up by water molecules. Thus, all ruthenium(II1) chloride existing in the above form in an aqueous alkaline medium will be the actual reacting species. It has frequently been observed that transition metal complexes are a good abstracting agent for the hydride ion.425They can adequately abstract a hydride ion even from a hydrogen molecule. In acidic medium the abstraction of a hydride ion from the a-carbon atom of the 2-propanol has also been observed.6 In a similar manner studies have been performed and it was concluded that hydride ion transfer takes place from the a-carbon atom of the alcohol or alkoxide ion to the metal atom.' This gives sufficient evidence that hydride ion transfer will take place during the course of the oxidation of the above-mentioned alcohols by the hexacyanoferrate(II1) ion in aqueous alkaline medium using ruthenium(II1) chloride as a homogeneous catalyst. On the basis of the above evidence, the oxidation of these alcohols might be described as shown in Scheme I. The proposed mechanism apparently shows that complex formation between the ruthenium(1I.I) species and the alcohol molecule takes place with the possible exchange of an hydroxide ion coordinated with the central metal ion. The Journal of Physical Chemistry, Vol. 81, No. 1 1 , 1977
1046
Singh et ai.
Scheme I FH3 [Ru(III)(OH)]Z+t H,C-CH-CH,OH alcohol FH3
(1)
I
k
a [Ru(III)~HOH,C-CH-CH,]~+ t OHcomplex C,
k-1
k
[complex]3t -?+
FH3 H,C-CH-C'CHOH t [Ru(III)(H)] 2 t
(11)
[RU(III)(H)]~'t 2 0 H - t 2Fe(CN),,-
fast
[Ru(III)(OH)12+t 2Fe(CN),4- t H,O
FH3 H3C-CH-C+HOH+ OH-
fast
7H3 H,C-CH-CHO t H,O 2-methylpropionaldehyde
+
(111) (IV)
The water molecule coordinated with it has been omitted for the sake of simplicity. The complex thus formed undergoes disproportionation yielding a carbonium ion via hydride ion abstraction from the a-carbon atom of the alcohol by the ruthenium(II1) species. The metal hydride then undergoes rapid oxidation with ferricyanide in the subsequent step, yielding the initial ruthenium(II1) species and the ferrocyanide ion. Although the oxidation of the metal hydride would take place in one-electron abstraction steps, for the sake of simplicity only one step in the stoichiometric equation has been given. The carbonium ion thus produced in the system will react rapidly with the hydroxide ion producing an aldehyde (or ketone with secondary alcohol) and a water molecule. If we assume a steady state condition for the concentration of complex C1we have
k [ Ru( III)(OH)] '+[alcohol] = k 2 + k - , [OH-] The total ruthenium(II1) concentration could be obtained from
+ [complex C1 3 ( 3 )
[Ru(III)]T = [Ru(III)(OH)]"
Substituting the value for the complex C1 in eq 3 and solving for [Ru(III)(OH)]~', we have [Ru(III)(OH)] '+
(4) Now the final rate law in terms of decreasing concentration of the hexacyanoferrate(II1) would be - d[Fe(CN),] 3 dt - 2klk2 [alcohol] [Ru(III)] T (5) k, + k-l[OH-] + k,[alcohol]
-
This rate law (eq 5) apparently accounts for pseudofirst-order kinetics with respect to ruthenium(II1) concentration. The retarding trend due to hydroxide ion is also quite obvious from eq 5. At high hydroxide ion and low alcohol concentration the inequality k 2 + k-l[OH-] >> kl[alcohol] would be evident and eq 5 reduces to -d[Fe(CN),] dt
3--
- 2 k l k 2 [ a l ~ ~ h[Ru(III)] ~l] T 122
+ k-I[OH-]
The Journal of Physical Chemistry, Vol. 81 No. 1 1 , 1977 I
(6)
Figure 4.
This rate law (eq 6) clearly explains the first-order kinetics with respect to alcohol at low concentrations. Further verification of rate expression 5 might be made on rearranging it as -1 =
Vi
1 212 1 [alcohol] [ Ru( 111)] T k-, [OH-] 2k,h2[alcohol] [Ru(III)] T
+
1 + 2k,[Ru(III)I T
(7)
where Vi = -d[Fe(CN),I3-/dt. According to eq 7, a plot of l / V i vs. [OH-] gives a straight line (Figure 4) and supports the velidity of eq 5. The slope gives values for Klk2as 5.1 and 32.3 for butan-2-01 and 2-methyl-lpropanol at 30 "C (where K1 = k 1 / k - J , respectively. Similarly, a plot of 1/Vi vs. l/[alcohol] gives a straight line (Figure 5) with a positive intercept at the 1/Vi axis. These straight lines again support the validity of rate law 5. From the intercept the value of k2 is 52.1 min-l for butan-2-01 and 2-methyl-1-propanol both at 30 O C . Now the value of Kl is 9.5 X lo-' and 59 X lo-' for butan-2-01 and 2methyl-1-propanol, respectively, at 30 O C . A close examination of K1 values clearly reveals that the rate of complex formation (kl)is slower than the reverse process, i.e., k-l, and for this reason the value of K1 is quite small. It is quite remarkable to note that the rate of complex formation of butan-2-01and 2-methyl-1-porpanol is in the following order: butan-2-01 < 2-methyl-lpropanol. The value of k2 obtained for these two alcohols clearly reveals that the rate of disproportion is of the same order. From the experimental results, it becomes apparent that for one molecule of alcohol two molecules of hexacyanoferrate(II1) are consumed. Thus, the actual amount of aldehyde or ketone produced in a particular run will be half of the initial hexacyanoferrate(II1)concentration and the stoichiometry of the reaction might be represented as ?H3 CH,-CH-CH,OH + 20H- t 2Fe(CN),,+ 2Fe(CN),4- t 2H,O
7% --t
CH,-CH-CHO
1047
Ru Catalyzed Oxidation of Alcohols
Scheme I1
YH3
Te rnp = 30’
k, [ R U ( O H ) , ] ~ -t~ H,C-CH,OH e [ c ~ r n p l e x ] ~+- OH~ k-,
(4) -1sobutanol
k OH, [ c ~ r n p l e x ]2 ~ - ~[(HO),-, R u ( I I I ) H ] ~ - “
.s 10.0
YH3
+ H,C-CH CHO YHz [ R U ( O H ) , - , ( H ) ] ~ -+ ~ 30H-
fast
Figure 5.
A similar equation can be written for butan-2-01, In the above-mentioned reaction mechanism the complex formed in step I disproportionates into the ruthenium(II1) hydride and the corresponding carbonium ion. The carbonium ion thus formed reacts with one molecule of hydroxide ion with the fast process resulting in the corresponding aldehyde or ketone and a water molecule. Another possibility which might be assumed is that the complex formed in step I would disproportionate in the subsequent step yielding the ruthenium(II1) hydride and the corresponding aldehyde or ketone. The steps then would be as shown in Scheme 11. Now, considering the general form of ruthenium(II1) chloride, the probable steps then would be as shown in Scheme 11. The coordinated hydroxide ions are given an unknown number x . Applying a steady state conditions, the rate law for this sequence of steps is
-d[ Fe(CN),] dt
3-
- 2klk2[Ru(III)]T[alcoholl k2
+h
1
[alcohol]
+ k-1 [OH-]
(8)
Rate laws 5 and 8 are the same although they have been derived from two different reaction schemes. Thus, the validity of Scheme I1 is again obvious.
Experimental Section Analar (BDH) grade samples of butan-2-01 and 2methyl-1-propanol were employed and were redistilled before use. The solution of ruthenium trichloride was prepared by dissolving the sample in very dilute hydro-
+ 2Fe(CN),’-
[Ru(OH),]~-+ ~ 2H,O
+ 2Fe(CN),4-
chloric acid with the final strength of HC1 and that of ruthenium trichloride kept at 16.38 X M and 4.8 X M, respectively. The sample of NaOH was of AR (BDH) grade but that of potassium hexacyanoferrate(II1) was of GR (S. Merck) grade. The initiation of the reaction is carried out by mixing the requisite quantity of alcohol solution maintained at a constant temperature into a solution of K3Fe(CN)G, NaOH, and RuC13 kept in a reaction bottle at the same temperature. The temperature of the reaction mixture was kept constant with an electrically operated thermostat with an accuracy of f O . l “C. The progress of reaction was followed by estimating the amount of hexacyanoferrate(I1)ion produced after definite time intervals with a standard solution of ceric(1V) sulfate using ferroin as a redox indicator.8 The method always gave reproducible results. This clearly confirms that ceric(IV) does not react with the alcohols under the present experimental conditions. In a separate set of experiments methyl ethyl ketone and 2-methylpropionaldehyde were confirmed as the main oxidation products via the catalytic route using chromatography.
Acknowledgment. One of the authors (R. K. Singh) wishes to express his thanks to U.G.C. (New Delhi) for financing the project.
References and Notes (1) S. Wolfe, S. K. Hasan, and J. R. Campbell, J. Chem. SOC.D, 1420 (1970). (2) R. E. Connick and D. A. Fine, J. Am. Chem. Soc., 82, 4187 (1960). (3) (a) F. A. Cotton and 0 . Wilkinson, “Advanced Inorganic Chemistry”, 2nd ed, Wiley, New York, N.Y., p 153; (b) W. P. Griffith, “The Chemistry of the Rarer Platinum Metals”, Interscience, New York, N.Y., 1967, p 27. (4) J. F. Harrod, S. Ciccone, and J. Halpern, Can. J. Chem., 39, 1372 (1961). (5) (a) G. Lee Donald and V. D. E. Mathus, Can. J. Chem., 50, 2000, 3129 (1972); (b) S. Yoel and R. Garryl, ibid., 52, 3825 (1974). (6) H. 6. Charman, J . Chem. SOC. 6 , 629 (1967). (7) J. Chatt and B. L. Shah, Chem. Ind., 290 (1961). (8) V. N. Singh, H. S. Singh, and 6. B. L. Saxena, J. Am. Chem. Soc., 91, 2643 (1969).
The Journal of Physical Chemistry, Vol. 81, No. 1 1 , 1977
