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Ru(VI)catalysis in some oxidation reactions by alkaline hexacyanoferrate(III)
J. inorg, nucl. Chem. Vol. 43, No. II, pp. 2893.-2897, 1981 Printed in Great Britain
0022-1902/81/112893-05502.00/0 Pergamon Press Ltd
Ru(VI)CATALYSIS IN SOME OXIDATION REACTIONS BY ALKALINE HEXACYANOFERRATE(III) R. K. DWlVEDI,H. NARAYANand K. BEHARI* Chemical Laboratory,University of Allahabad, Allahabad, India
(Received 15 September 1980; received for publication 7 March 1981) Abstract--The oxidation of a primary and a secondary alcohol by hexacyanoferrate(IlI)under mild conditions
employingruthenium(VI)as homogeneouscatalysthas been studied. The results obtainedsuggestthat the oxidationof these alcohols proceeds via the formation of complex between ruthenium(VI) and the substrate which slowly decomposes, followedby a fast reaction between the reduced ruthenium species and hexacyanoferrate(III)ion. A product study shows the formation of the correspondingketone or aliphatic acid.
Catalysis by a number of metal ions has been observed in the oxidation of organic compounds with alkaline hexacyanoferrate(III)[l]. Sodium ruthenate has been employed to oxidise a number of unsaturated compounds [2]. We have used ruthenium(VI) as a homogeneous catalyst in the oxidation of pentane-3-ol and n-pentanol by hexacyanoferrate(III) in alkaline medium at constant ionic strength (maintained by sodium perchlorate). Both reactions are zero order in hexacyanoferrate(Ill), first order in ruthenium(Vl), the rate varies linearly with the substrate in the lower range of concentrations and a retarding effect of hydroxyl ion on the velocity of the reaction is found. RESULTS AND DISCUSSION
The details of the kinetic data for the ruthenium(VI) catalysed oxidation of n-pentanol and pentane-3-ol are presented in the Tables 1-4. The standard zero order rate constant (kJ presented in these tables is the average of
at least two runs. The value of k~ has been calculated as reported by Krishna et al.[3]. Figure 1 shows a typical zero order plot for the oxidation of both the compounds. It is obvious from these plots that the reaction velocity remains constant upto more than 70% of the reaction. Table 1 clearly indicates that the k~ values are almost constant for various initial hexacyanoferrate(III) concentrations. However, at higher concentrations, the values are somewhat higher, which might be due to secondary salt effect, and it seems valid to assume a zero order dependence of the reaction rate on the hexacyanoferrate(III) concentration. Table 2 contains the average standard zero order rate constants for different concentrations of the alcohols. It is evident that as the concentration of alcohol is increased, the values of the factor kJ[substrate] goes on decreasing showing that the reaction velocity is not linearly proportional to the substrate concentration but attains a constant value at higher concentrations of the
Table 1. Fem~. 30°C (A)
~CH-3
- -
r
{s) = 3.33
p e n t dr~3 -°1 3
x io-~
= 25.00
L ~,u,vi)]
f CH-]
x IO-3H
~-ientaro
2.47 x io-Grt jb
CRu~VI)~
= <,.2Obl
M
:I2.BO
x iO "M
= 6.82
x IO-7M
]h = 0 . 2 6 M [Fe(CN]
x 103
: 5.~O x iO-•:
rain -I
x IO3M
:3]
ks x
IO5M
rain -I
O. 40
1 .o1
O. 40
O, 52
0.50
1 o12
O.5~
O.61
0.66
1 .57
O .6o
0.84
I .00
1 .78
I .OO
1 .07
1.43
2.25
t.43
1.26
2.00
2.39
2.00
1.21
2893
2894
R. K. DWIVEDI et aL Table 2. ~emp. ~30°C
(a) [Fe(CN)6~
(s)
- 2.00 x iO-351
[Fe(CN)o33
FOH-3
: 3.33 x IO-2M
[OH-]'
: 5.OO x Io-2M
[Ru(VZ)]
= 2.47 x 10-e'N
E~,u,.VI)J
- 6 . 8 2 x 10-?M
wu = O.26M
= 2.OU x IO-3M
~U = O.2OM
Pentane-
ksXlO 5
xlO3M
M rain-I xlO 3 rain-I
xlO3M
51 rain-I
Pentanol~ xlO3min -
6.25
0.97
1.55
iO.OO
1.O3
1.O3
9.09
1 .26
1 .|8
12.50
1.31
1.05
12.50
1.58
1.26
16.66
1.68
i.O1
255OO
2.59
1.O4
20.00
1 .87
0.96
33.33
3.23
0.96
25.00
2.75
0.86
50.00
4.04
0.80
33.33
2.49
0.75
iOO.OO
7.02
0.70
50.00
3.25
0.65
3o1~
k=/[Pentane-
n-
~3o1~
[~o~o~]
~
~1°5
-
Table 3. Temp.=30°C
(A)
(B)
EFe(CN);31
[0. 3 Pentane-3oZ
= 2.00 x IO-3M
E,o~c.~ 3] : 200 ~ 10-3.
: 33~ x 10-2.
To.-]
: 2 5 . O O x IO-3M
[~ ~en~o~-12 50 x 10~.
]
o 500
~/ - O.261"I
M/ : O. 26M
[ ~-(~)3 xlO6M
x lO-2.
~]
k.lo 5
k/[~.(~)]
Ru( VI )J
ksXlO 5
ks/[au(vl
M min - I
xlo-lmln -I
xlO7M
M mln -I
xlO - imi n- i
O.49
O .54
1 .io
0.34
0.75
2.17
1.48
1.62
1.09
0.68
1.31
1.92
2.47
2.59
1.05
1.36
2.86
2. Io
3.45
3.50
1.O2
2.04
4.25
2.08
4.44
4.41
0.92
2.72
5.89
2.15
4.93
4 •94
1 .oO
3.41
7.2O
2.12
Table 4. r~e[l| Lg. : 3 0 ° C
(A)
E,(~,)~33
(8)
l,
= 2.00 x 10-3~.~
p~(c,)~ 3]
- 2.00 ~ ~0-3,
EPen~ane-3-oJ
: 2 5 . o o x io-351
En-Pentanol~
: 12.5oxlO-3M
[R~vz> 3
= 2.47 × 10-6.
fau(v~)3
/u = 0.2~ M ['O1-{-,3 x io~I
ksXlO5M'mi'n -I
= 682
× i0-7.
Ju = 0 . 2 6 M ~OH~x
iO"~I
ksXlO5M mln -~
20 .OO io.oo
1.71 1.97
20 .OO io.oo
O .70 O .58
5 ,OO
2.29
6.66
1 .28
3 ,33
2.59
5.00
1.31
2,50
2.65
3.33
2.24
2,00
2.72
2.50
2.52
i .25
2.71
Ru(VI) catalysis in some oxidationreactions
2895
A t5
B
5
T E
"Q
4
E
Jo
x
% Z ID ~
2
x
05 I
Z
20
40
'
60
Time,
I I
I 2
I 3
80
I 5
M
rain
Temp. = 30"C [Fe(CN)-: ] = 2 . 0 0 x l O - ~
Fig. l. Line-A, Line-B. Initial [Fe(CN)~ 3] x 103M = 2.00, 2.00. [OH] × IlFM = 5.00, 3.33. [Substrate] x I03M = 25.00, 12.50. [Ru(VI)] × 106M - 0.68, 2.47. u = 0.26. Temp. = 30'C.
compound. This behaviour has clearly been demonstrated in Fig. 2. Table 3 contains the average ks values at different concentrations of ruthenium(VI). A linear correlation is obtained (Fig. 3) between [Ru(VI)] vs k, which clearly indicates that the order is unity with respect to ruthenium(VI). Table 4 contains the average k, values at different alkali concentrations. These tables show that the reaction velocity increases with decrease in the concentration of hydroxide ion (Fig. 4). At higher concentrations of sodium hydroxide the uncatalysed oxidation, of the both alcohols, may also be significant, therefore, the experiments were performed at low concentration of sodium hydroxide and it has been checked that the oxidation by uncatalysed route by alkaline hexacyanoferrate(III) is negligible under the experimental conditions. Before writing the scheme for the oxidation, it will be more proper to discuss the probable species of ruthenium(VI) involved. Electronic spectral4] have confirmed that the lower oxidation state of ruthenium is
l 4
[ o . 7 -- +
M
M
[ P e n ' t a n e - B - o t ] = 2 5 x 0 0 x I0 -3 M #=0.26
Fig. 3. present as hydrated species, but the higher oxidation state of ruthenium need not be hydrated. Therefore, it is likely that ruthenium(VI) in present study can be represented as RuO22 not as a hydrated species as in the case [4] of Ru(III). Thus on the experimental evidences, the oxidation of these alcohols might be given by following steps: K
"(RuO+.OH) 3
RuO4- 2 + O H - . RuO22+AI.
K|
(i)
--9
' (Complex) -
(ii)
k
(Complex) - 2 - - ~ RuO2.xH20 + product + OH +x H20
(fii) fast
RuOvx H20 + 2Fe(CN);-3 + 4OH- ---~
28
T C
-g
6
TE .~
x
2.4
%
i
x .x°
I
J
4
J
8
[Penl'one-3-ot]x
i2
(CN)-:] = 2.00x
[oH-l=
I0-~ M
3 . 3 3 x 1 0 -2 M
[Ru(VI)]- 2.47 x 10-3 M #=0,26
Fig. 2. J I N C Vol. 43, No. I I - - U
0
16
I
i
J
L
5
IO
i5
20
I0 -~ M
Temp. = 3 0 ° C [Fe
20
Temp.= 30 oC [Fe(CN)J=2.00x,0-3 M [Pentone-:B-eL] = 2 5 . 0 0 x I0-' M
[Ru(VI)] =2.47xIC) 6 M # =0.26
Fig. 4.
R. K. DWIVEDIet al.
2896
Table 5. 'rPrap.
= 300(:
C xlO3M ~ _/ xlO~l
xlO6M
Pent dne-3-ol
2.00
3.33
2.47
n-Pentanol
2.00
5.00
0.62
RuO; 2 + 2Fe(CN)g4 + (x + 2)H20
(iv)
Considering steps (i)-(iv) a rate law can be derived as d[Fey] 2kK,[AI][RuO~2]T Rate = - d t = 1 +K,[A1]+K[OH-]"
(1)
"i~
(mln'll
9/./. 16.40
16.78
47.28
15.79
6
'ol,< ~" 3
Where AI stands for the alcohols; T for total and Fey for hexacyanoferrate(III). During the study of reaction between per-ruthenate and manganate ion[5] in alkaline aqueous medium, it was suggested that one or more of the anions could possibly have hydroxide ion associated with them. Hydroxide ion association can be represented by the equilibrium:
Where M is Ru or Mn and x is 1 or 2 depending upon the oxidation state. If the formation constant for above reaction is large, the effect of varying[OH-] would not be pronounced. But Table 4 clearly indicates a retarding effect of hydroxyl ion. Though the co-ordination of OHwith RuO; 2 is not common, during a kinetic study Syman et a/.[6] showed that the hydroxide ion coordinates with per-ruthenate ion to a very small extent and suggested that ruthenium represents an intermediate case between osmium and rhenium. Therefore, step one in above mechanism is justified. Therefore, the assumption that (1 + K,[A1]) ~>K[OH-] is not unjustified. Hence eqn (1) can be rewritten as: Rate = 2kK, [A1] [RuO42]T (2)
The eqn (2) can be rearranged as 1
Rate
_
1
1
2kK, [A1] [RuO;2]T + 2k[RuOZ2]T"
i
8
i t2
[Peril"one-3- oL]-' xlO-' M-' Temp= 30"C Fe(CN): = 2.00x 1()3M
[o.-] =3.33x,o-'M Ru(Vll =2.47xlO-6M /z = 0.26
MOY + OH- ~ MO4.OH(x+l)-.
1 +K, [A1]
I
4
(3)
The validity of the rate law (3) is further supported by following conclusions: (a) The linear correlation between the inverse of rate against[A1]-' is consistent with the rate law (3) and giving a positive intercept on the Y-axis (Fig. 5). From this intercept the value of k has been calculated and given in Table 5. The slope of this straight line is equal to 1/2kK,[RuOZ2]T and from this the value of K, is calculated and is given in Table 5. (b) The plots of ks vs [RuO~2]T were made and a straight line was obtained (Fig. 3) whose slope will be equal to 2kK,[A1]/(I+K,[AI]). The value of K, calculated from this slope is 15.09 for pentane-3-ol and
Fig. 5. 17.82 for n-pentanol. It is obvious that the value of K,, thus obtained, is quite close to the value obtained above (Table 5) which supports the proposed reaction mechanism. EXPERIMENTAL
Pentane-3-ol and n-pentanol were distilled before use. The solution of sodium ruthenate was prepared as reported elsewhere[2,7]. The concentration of sodium ruthenate was found using a Beckman spectrophotometermodel-26by measuring the absorbance at 460 nm, the extinction coefiicient[8] being 1220. The stock solutionof cerium sulphate was prepared by the standard method[91. The reaction was started by mixing substrate solution with the mixture of other reactants namely hexacyanoferrate(III),alkali, sodium perchlorate and sodium ruthenate and progress of reaction was followed by estimating the amount of hexacyanoferrate(III) produced after definite intervals of time titrating against with standard solution of cerium sulphate using ferroin as a redox indicator. The ionic strength of the medium was kept constant by addition of sodium perchlorate because perchlorate ion does not form the complex with ruthenium(VI) Acknowledgement--The financial support to H. Narayanis thankfully acknowledgedfrom S.S.I.R., New Delhi. REFERENCES
1. L. Losenthaler, Chem. Zentr. 56, 441 (1932). 2. D. G. Lee, H. Helliwell and V. Chang, J. Org. Chem. 41, 3644 (1976). 3. B. Krishna and H. S. Singh, Zeit fur phys. Chemie, 231, 399 (1966). 4. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn, p. 153. Wiley,New York (1966).
Ru(VI) catalysis in some oxidation reactions 5. E. V. Loma and C. H. Brubaker, Inorganic Chem. 5, 1637 (1964). 6. M. C. R. Syman and A. Corrington, J. Am. Chem. Soc. 82, 284 (1960). 7. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chem-
2897
istry, 2nd Edn, p. 1006. Interscience, New York (1966). 8. R. E. Connick and G. R. Hurley, J. Am. Chem. Soc. 3rd Edn. 74, 5012 (1952). 9. I. A. Vogel, A Text Book of Qualitative Inorganic Analysis, 3rd Edn. p. 317. Longmans, Green & Co. Ltd., London (1961).
0022-1902/81/112893-05502.00/0 Pergamon Press Ltd
Ru(VI)CATALYSIS IN SOME OXIDATION REACTIONS BY ALKALINE HEXACYANOFERRATE(III) R. K. DWlVEDI,H. NARAYANand K. BEHARI* Chemical Laboratory,University of Allahabad, Allahabad, India
(Received 15 September 1980; received for publication 7 March 1981) Abstract--The oxidation of a primary and a secondary alcohol by hexacyanoferrate(IlI)under mild conditions
employingruthenium(VI)as homogeneouscatalysthas been studied. The results obtainedsuggestthat the oxidationof these alcohols proceeds via the formation of complex between ruthenium(VI) and the substrate which slowly decomposes, followedby a fast reaction between the reduced ruthenium species and hexacyanoferrate(III)ion. A product study shows the formation of the correspondingketone or aliphatic acid.
Catalysis by a number of metal ions has been observed in the oxidation of organic compounds with alkaline hexacyanoferrate(III)[l]. Sodium ruthenate has been employed to oxidise a number of unsaturated compounds [2]. We have used ruthenium(VI) as a homogeneous catalyst in the oxidation of pentane-3-ol and n-pentanol by hexacyanoferrate(III) in alkaline medium at constant ionic strength (maintained by sodium perchlorate). Both reactions are zero order in hexacyanoferrate(Ill), first order in ruthenium(Vl), the rate varies linearly with the substrate in the lower range of concentrations and a retarding effect of hydroxyl ion on the velocity of the reaction is found. RESULTS AND DISCUSSION
The details of the kinetic data for the ruthenium(VI) catalysed oxidation of n-pentanol and pentane-3-ol are presented in the Tables 1-4. The standard zero order rate constant (kJ presented in these tables is the average of
at least two runs. The value of k~ has been calculated as reported by Krishna et al.[3]. Figure 1 shows a typical zero order plot for the oxidation of both the compounds. It is obvious from these plots that the reaction velocity remains constant upto more than 70% of the reaction. Table 1 clearly indicates that the k~ values are almost constant for various initial hexacyanoferrate(III) concentrations. However, at higher concentrations, the values are somewhat higher, which might be due to secondary salt effect, and it seems valid to assume a zero order dependence of the reaction rate on the hexacyanoferrate(III) concentration. Table 2 contains the average standard zero order rate constants for different concentrations of the alcohols. It is evident that as the concentration of alcohol is increased, the values of the factor kJ[substrate] goes on decreasing showing that the reaction velocity is not linearly proportional to the substrate concentration but attains a constant value at higher concentrations of the
Table 1. Fem~. 30°C (A)
~CH-3
- -
r
{s) = 3.33
p e n t dr~3 -°1 3
x io-~
= 25.00
L ~,u,vi)]
f CH-]
x IO-3H
~-ientaro
2.47 x io-Grt jb
CRu~VI)~
= <,.2Obl
M
:I2.BO
x iO "M
= 6.82
x IO-7M
]h = 0 . 2 6 M [Fe(CN]
x 103
: 5.~O x iO-•:
rain -I
x IO3M
:3]
ks x
IO5M
rain -I
O. 40
1 .o1
O. 40
O, 52
0.50
1 o12
O.5~
O.61
0.66
1 .57
O .6o
0.84
I .00
1 .78
I .OO
1 .07
1.43
2.25
t.43
1.26
2.00
2.39
2.00
1.21
2893
2894
R. K. DWIVEDI et aL Table 2. ~emp. ~30°C
(a) [Fe(CN)6~
(s)
- 2.00 x iO-351
[Fe(CN)o33
FOH-3
: 3.33 x IO-2M
[OH-]'
: 5.OO x Io-2M
[Ru(VZ)]
= 2.47 x 10-e'N
E~,u,.VI)J
- 6 . 8 2 x 10-?M
wu = O.26M
= 2.OU x IO-3M
~U = O.2OM
Pentane-
ksXlO 5
xlO3M
M rain-I xlO 3 rain-I
xlO3M
51 rain-I
Pentanol~ xlO3min -
6.25
0.97
1.55
iO.OO
1.O3
1.O3
9.09
1 .26
1 .|8
12.50
1.31
1.05
12.50
1.58
1.26
16.66
1.68
i.O1
255OO
2.59
1.O4
20.00
1 .87
0.96
33.33
3.23
0.96
25.00
2.75
0.86
50.00
4.04
0.80
33.33
2.49
0.75
iOO.OO
7.02
0.70
50.00
3.25
0.65
3o1~
k=/[Pentane-
n-
~3o1~
[~o~o~]
~
~1°5
-
Table 3. Temp.=30°C
(A)
(B)
EFe(CN);31
[0. 3 Pentane-3oZ
= 2.00 x IO-3M
E,o~c.~ 3] : 200 ~ 10-3.
: 33~ x 10-2.
To.-]
: 2 5 . O O x IO-3M
[~ ~en~o~-12 50 x 10~.
]
o 500
~/ - O.261"I
M/ : O. 26M
[ ~-(~)3 xlO6M
x lO-2.
~]
k.lo 5
k/[~.(~)]
Ru( VI )J
ksXlO 5
ks/[au(vl
M min - I
xlo-lmln -I
xlO7M
M mln -I
xlO - imi n- i
O.49
O .54
1 .io
0.34
0.75
2.17
1.48
1.62
1.09
0.68
1.31
1.92
2.47
2.59
1.05
1.36
2.86
2. Io
3.45
3.50
1.O2
2.04
4.25
2.08
4.44
4.41
0.92
2.72
5.89
2.15
4.93
4 •94
1 .oO
3.41
7.2O
2.12
Table 4. r~e[l| Lg. : 3 0 ° C
(A)
E,(~,)~33
(8)
l,
= 2.00 x 10-3~.~
p~(c,)~ 3]
- 2.00 ~ ~0-3,
EPen~ane-3-oJ
: 2 5 . o o x io-351
En-Pentanol~
: 12.5oxlO-3M
[R~vz> 3
= 2.47 × 10-6.
fau(v~)3
/u = 0.2~ M ['O1-{-,3 x io~I
ksXlO5M'mi'n -I
= 682
× i0-7.
Ju = 0 . 2 6 M ~OH~x
iO"~I
ksXlO5M mln -~
20 .OO io.oo
1.71 1.97
20 .OO io.oo
O .70 O .58
5 ,OO
2.29
6.66
1 .28
3 ,33
2.59
5.00
1.31
2,50
2.65
3.33
2.24
2,00
2.72
2.50
2.52
i .25
2.71
Ru(VI) catalysis in some oxidationreactions
2895
A t5
B
5
T E
"Q
4
E
Jo
x
% Z ID ~
2
x
05 I
Z
20
40
'
60
Time,
I I
I 2
I 3
80
I 5
M
rain
Temp. = 30"C [Fe(CN)-: ] = 2 . 0 0 x l O - ~
Fig. l. Line-A, Line-B. Initial [Fe(CN)~ 3] x 103M = 2.00, 2.00. [OH] × IlFM = 5.00, 3.33. [Substrate] x I03M = 25.00, 12.50. [Ru(VI)] × 106M - 0.68, 2.47. u = 0.26. Temp. = 30'C.
compound. This behaviour has clearly been demonstrated in Fig. 2. Table 3 contains the average ks values at different concentrations of ruthenium(VI). A linear correlation is obtained (Fig. 3) between [Ru(VI)] vs k, which clearly indicates that the order is unity with respect to ruthenium(VI). Table 4 contains the average k, values at different alkali concentrations. These tables show that the reaction velocity increases with decrease in the concentration of hydroxide ion (Fig. 4). At higher concentrations of sodium hydroxide the uncatalysed oxidation, of the both alcohols, may also be significant, therefore, the experiments were performed at low concentration of sodium hydroxide and it has been checked that the oxidation by uncatalysed route by alkaline hexacyanoferrate(III) is negligible under the experimental conditions. Before writing the scheme for the oxidation, it will be more proper to discuss the probable species of ruthenium(VI) involved. Electronic spectral4] have confirmed that the lower oxidation state of ruthenium is
l 4
[ o . 7 -- +
M
M
[ P e n ' t a n e - B - o t ] = 2 5 x 0 0 x I0 -3 M #=0.26
Fig. 3. present as hydrated species, but the higher oxidation state of ruthenium need not be hydrated. Therefore, it is likely that ruthenium(VI) in present study can be represented as RuO22 not as a hydrated species as in the case [4] of Ru(III). Thus on the experimental evidences, the oxidation of these alcohols might be given by following steps: K
"(RuO+.OH) 3
RuO4- 2 + O H - . RuO22+AI.
K|
(i)
--9
' (Complex) -
(ii)
k
(Complex) - 2 - - ~ RuO2.xH20 + product + OH +x H20
(fii) fast
RuOvx H20 + 2Fe(CN);-3 + 4OH- ---~
28
T C
-g
6
TE .~
x
2.4
%
i
x .x°
I
J
4
J
8
[Penl'one-3-ot]x
i2
(CN)-:] = 2.00x
[oH-l=
I0-~ M
3 . 3 3 x 1 0 -2 M
[Ru(VI)]- 2.47 x 10-3 M #=0,26
Fig. 2. J I N C Vol. 43, No. I I - - U
0
16
I
i
J
L
5
IO
i5
20
I0 -~ M
Temp. = 3 0 ° C [Fe
20
Temp.= 30 oC [Fe(CN)J=2.00x,0-3 M [Pentone-:B-eL] = 2 5 . 0 0 x I0-' M
[Ru(VI)] =2.47xIC) 6 M # =0.26
Fig. 4.
R. K. DWIVEDIet al.
2896
Table 5. 'rPrap.
= 300(:
C xlO3M ~ _/ xlO~l
xlO6M
Pent dne-3-ol
2.00
3.33
2.47
n-Pentanol
2.00
5.00
0.62
RuO; 2 + 2Fe(CN)g4 + (x + 2)H20
(iv)
Considering steps (i)-(iv) a rate law can be derived as d[Fey] 2kK,[AI][RuO~2]T Rate = - d t = 1 +K,[A1]+K[OH-]"
(1)
"i~
(mln'll
9/./. 16.40
16.78
47.28
15.79
6
'ol,< ~" 3
Where AI stands for the alcohols; T for total and Fey for hexacyanoferrate(III). During the study of reaction between per-ruthenate and manganate ion[5] in alkaline aqueous medium, it was suggested that one or more of the anions could possibly have hydroxide ion associated with them. Hydroxide ion association can be represented by the equilibrium:
Where M is Ru or Mn and x is 1 or 2 depending upon the oxidation state. If the formation constant for above reaction is large, the effect of varying[OH-] would not be pronounced. But Table 4 clearly indicates a retarding effect of hydroxyl ion. Though the co-ordination of OHwith RuO; 2 is not common, during a kinetic study Syman et a/.[6] showed that the hydroxide ion coordinates with per-ruthenate ion to a very small extent and suggested that ruthenium represents an intermediate case between osmium and rhenium. Therefore, step one in above mechanism is justified. Therefore, the assumption that (1 + K,[A1]) ~>K[OH-] is not unjustified. Hence eqn (1) can be rewritten as: Rate = 2kK, [A1] [RuO42]T (2)
The eqn (2) can be rearranged as 1
Rate
_
1
1
2kK, [A1] [RuO;2]T + 2k[RuOZ2]T"
i
8
i t2
[Peril"one-3- oL]-' xlO-' M-' Temp= 30"C Fe(CN): = 2.00x 1()3M
[o.-] =3.33x,o-'M Ru(Vll =2.47xlO-6M /z = 0.26
MOY + OH- ~ MO4.OH(x+l)-.
1 +K, [A1]
I
4
(3)
The validity of the rate law (3) is further supported by following conclusions: (a) The linear correlation between the inverse of rate against[A1]-' is consistent with the rate law (3) and giving a positive intercept on the Y-axis (Fig. 5). From this intercept the value of k has been calculated and given in Table 5. The slope of this straight line is equal to 1/2kK,[RuOZ2]T and from this the value of K, is calculated and is given in Table 5. (b) The plots of ks vs [RuO~2]T were made and a straight line was obtained (Fig. 3) whose slope will be equal to 2kK,[A1]/(I+K,[AI]). The value of K, calculated from this slope is 15.09 for pentane-3-ol and
Fig. 5. 17.82 for n-pentanol. It is obvious that the value of K,, thus obtained, is quite close to the value obtained above (Table 5) which supports the proposed reaction mechanism. EXPERIMENTAL
Pentane-3-ol and n-pentanol were distilled before use. The solution of sodium ruthenate was prepared as reported elsewhere[2,7]. The concentration of sodium ruthenate was found using a Beckman spectrophotometermodel-26by measuring the absorbance at 460 nm, the extinction coefiicient[8] being 1220. The stock solutionof cerium sulphate was prepared by the standard method[91. The reaction was started by mixing substrate solution with the mixture of other reactants namely hexacyanoferrate(III),alkali, sodium perchlorate and sodium ruthenate and progress of reaction was followed by estimating the amount of hexacyanoferrate(III) produced after definite intervals of time titrating against with standard solution of cerium sulphate using ferroin as a redox indicator. The ionic strength of the medium was kept constant by addition of sodium perchlorate because perchlorate ion does not form the complex with ruthenium(VI) Acknowledgement--The financial support to H. Narayanis thankfully acknowledgedfrom S.S.I.R., New Delhi. REFERENCES
1. L. Losenthaler, Chem. Zentr. 56, 441 (1932). 2. D. G. Lee, H. Helliwell and V. Chang, J. Org. Chem. 41, 3644 (1976). 3. B. Krishna and H. S. Singh, Zeit fur phys. Chemie, 231, 399 (1966). 4. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn, p. 153. Wiley,New York (1966).
Ru(VI) catalysis in some oxidation reactions 5. E. V. Loma and C. H. Brubaker, Inorganic Chem. 5, 1637 (1964). 6. M. C. R. Syman and A. Corrington, J. Am. Chem. Soc. 82, 284 (1960). 7. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chem-
2897
istry, 2nd Edn, p. 1006. Interscience, New York (1966). 8. R. E. Connick and G. R. Hurley, J. Am. Chem. Soc. 3rd Edn. 74, 5012 (1952). 9. I. A. Vogel, A Text Book of Qualitative Inorganic Analysis, 3rd Edn. p. 317. Longmans, Green & Co. Ltd., London (1961).