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Kinetics of the oxidation of Hg(I) by Fe(III)-2,2′-bipyridyl complex
J. inorg. ,ucl. Chem. Vol. 42, pp. 1181-1184 Pergamon Press Ltd., 1980. Printed in Great Britain
KINETICS OF THE OXIDATION OF Hg(I) BY Fe(III)-2,2'BIPYRIDYL COMPLEX P. V. SUBBA RAO,* M. SURYANARAYANA,K. V. SUBBAIAH and P. S. N. MURTY Department of Chemistry, Andhra University, Waltair-530 003, India
(Received for publication 22 October 1979) Abstract--The kinetics of the oxidation of Hg(I) by the Fe(III)-2,2'-bipyridylcomplex has been investigated spectropbotometrically. The reaction obeys first order kinetics with respect to Hg(I) and Fe(III). Plots of l/kt vs l/[2,2'-Bipyridyl]z and l/kt vs [H+]2 are straight lines with positive intercepts on the llkt-axes. A reaction mechanism is proposed in which the rate-determining step is believed to involve concomitant breaking of the (Hg-Hg)2+ bond and electron transfer. Hg(II) solution: 0.05M solution of Hg(II) was prepared by dissolving mercuric nitrate in 0.5 M nitric acid. All solutions were prepared from deionised water and analytical reagent grade chemicals. The source of H+ ion was nitric acid. Kinetic procedure. The course of the reaction was followed by measuring the absorbance of the product ferrodiin at 510nm using a Hilger UV Speck spectrophotometer. At this wavelength all the other materials concerned have negligibleabsorption. To avoid any effect of light the reaction vessel was coated outside with black paint. The ionic strength was maintained constant with sodium nitrate. The kinetic runs were performed under the conditions [Hg(I)] and [Bipy]§>>[Fe(III)], thus isolating the latter. Kinetic orders. A plot of log (D®-/9,) vs time (where D~ and D, are absorbances at infinite time and at time t respectively) is a straight line under the above conditions, showing first order EXPERIMENTAL behaviour with respect to Fe(III). The pseudo first order rate Materials. Hg(l) solution, 0.01 M: Hg(1) nitrate solution was constant, kl, was calculated from the slope. prepared in 0.025 M nitric acid and standardized titrimetrically Effect of Hg(I) concentration. The pseudo first order rate with standard ferric alum in the presence of thiocyanate. constants, k~, were determined at different concentrations of Fe(HI) perchlorate: Iron wire (Riedel, Analar) was digested Hg(I), keeping the concentrations of other materials constant. with 70% perchloric acid (G. R. E. Merck) and evaporated to The data were presented in Table 1. The results show that k I is crystallization. The material was twice reccrystallized from 60% directly proportional to the concentration of Hg(I), i.e. the perchloric acid and stored under vacuum. An approximately reaction is first order with respect to Hg(I). 0.1 M solution was prepared in perchloric acid. The acid content Effect of 2,2'-bipyndyl concentration, k~ was also determined and Fe(llI) in the solution were determined by the method varying the concentrations of 2,2'-bipyridyl and at three temdescribed by Milburn et al.[3]. peratures, keeping [Fe(III)], [Hg22+] and [H+] constant. The plots 2,2'-bipyridyl solution: 0.05M solution was prepared by dis- of 1/kl vs l/[Bipy]2 were straight lines with positive slopes and solving in 50% (v/v) methanol. intercepts on the 1/kl axis (Fig. 1). Effect of acid concentration. The values of k~ were determined at different hydrogen ion concentrations keeping the ionic tFerridiin denotes Fe(III)-tris(bipyridyl)complex. strength constant with sodium nitrate. Hydrogen ions were found ;tFerrodiindenotes Fe(II)-trisCoipyridyl)complex. to inhibit the reaction. A plot of 1/kl vs [H+]2 is a straight line §Bipy is the abbreviation used for 2,2'-bipyridyl. with positive intercept on the 1/kl axis (FIB. 2). INTRODUCTION
Blau[l] was one of the first to detect the complex between Fe(lll) and 2,2'-bipyridyl by direct reaction. The complex is different from ferridiin,t which cannot be obtained by direct mixing, but only by the oxidation of ferrodiin~[2]. A study of the oxidizing action of this complex has not been made. As a part of the investigation of the reactions of this complex, we have studied the kinetic study of the oxidation of Hg(1) by Fe(IH) in the presence of 2,2'-bipyridyl. The results show that the active oxidizing species is the 1:2 complex of Fe(lll) and 2,2'-bipyridyL Fe(lll) in the absence of 2,2'bipyridyl does not oxidize Hg(1) under the experimental conditions employed.
Table 1. Effect of Hg(1)concentration on the rate of oxidation: [Fe(III)] = 1.0 x 10-4 M; [H +] = 1.0 x 10-2 M;/~ = 0.1; [Bipy] = 2.0 x 10-2 M; MeOH = 20% (v/v); Temp = 35 _+0.1°C
[Hg[+] ~ io3, s
k I ~ 1o3, .i~-l
kl/ (ag~+~
0.8 1.0 1.3 1.6 2.0
5.O 6.3 8.1 i0.0 12.5
6.25 6.30 6.27 6.25 6.25
2.3
14.5
6.~0
1181
P. V. S. RAO et al.
1182
Table 2. Effect of Hg(II) concentration on rate: [Fe(IIl)] = i.0 x I0-~ M; [Hg22+]= 2.0 x 10-3 M; MeOH = 25% (v/v); [H÷] = 2.0 x 10-2 M; Temp = 35 + 0.1°C;/~ = 0.1 [Hg 2+] x 10 5, M
[Blpy] • 10 2 , M
k I x i0 ~, mln "I
0.0
2.00
4.15
0.5 1.0
2.00 2.00
5.80 3.10
1.5 2.0
2.00 2.00
2.75 2.25
0.5
2.15
4.12
1.0
2.90
4.20
1.5 2.0
2.45
4.15 4.10
2.60
Effect of Hg(II) concentration. Kinetic runs were carried out at different initial concentrations of Hg(II). The data presented in Table 2 show that the rate decreases with increasing concentration of Hg(II). Since Hg(II) forms a stable 1:3 complex with bipyridyl (log t3 = 19.54)[4], the authors believe that the inhibition is due to the decrease in the concentration of free bipyridyl ligand because of complex formation with Hg(II). This view is supported by the fact that when the decrease in the concentration of bipyridyl is compensated by adding an additional amount of bipyfidyl (three times to that of added Hg(II)) the rate of the reaction is practically unaffected by the added Hg(II) (Table 2). Influence of ionic strength. Kinetic runs were carried out at different ionic strengths by varying the concentration of sodium nitrate in the range 0.1-0.9. The rate has been found to increase with increase in ionic strength indicating that the rate-determining step involves ions of similar charge. The data were presented in Table 3. DISCUSSION
Nature of species of Fe(III) The results of Milburn and Vosburgh[3] show that Fe(III) at concentrations less than 1.0 x IO-3M does not
3.6 3.2
Table 3. Influence of ionic strength on rate: [Fe(III)]= 1.0 x 10-4 M; [H+] = 1.0 x l0-2 M; [Hg22+]= 1.0 x 10-3 M; [Bipy] = '2.0 x 10-2 M; MeOH = 20%(v/v); Temp = 35 + 0.1°C Ionic
strength,
~
k I x 10 3 , m l n -I
o.1 o.5 0.5 0.7 0.9
630 7.Ol 7.83 8.57 9.25
dimerize. A similar observation has also been made by Siddall and Vosburg[5]. Ciavatta and Grimaldi[6] have reported the negative logarithm of the hydrolysis constant of Fe(III) to be 3.1-+0.2. In the present investigation, the Fe(III) concentration employed was the order of 1 0 - ' M and hydrogen ion concentration was in the range 0.01-0.02 M. Hence it is reasonable to assume that there is no appreciable dimerization under these conditions, and the concentration of hydroxylated Fe(III)
2.4
2.8 2.0
2,4 1.6
2.0 .=_ :E
1.2
'9 1.2
.w.. 0 . 8
0.8 0.4 O.4
II I0-~' [81PY] 2, M -2
Fig. 1. Plots of l/k1 vs 1/[Bipy]2. [Fe(I[])]= 1.0xI0-4M; MeOH = 25% (v/v). [Hg22+l-- 1.0 x 10-3 M; [H÷] = 1.0 x 111-2M; (9-49, 31~C; A--A, 35°C; [B-..~, 41PC.
I 2 104x[H+] 2, M 2
I 3
I 4
Fig 2 Plot of l/k, vs [H+] = [Fe(I[])] = 10 x 10-4 M; MeOH = 20%(v/v). [Hg22.] = 20 x 10-3 M; [Bipy] = 2 0 x 10-= M ; / ~ = 002;
Temperature = 35-+O.1°C.
Kinetics of the oxidation of Hg(I) by Fe(lll)-2,2'-bipyridyl complex
has been neglected in view of the low hydrolysis constant. The hydrolysis is also suppressed, since the ligand (2,2'-bipyridyl) was present in large excess. It has been reported [7] that 2,2'-bipyridyl exists in singly protonated form (pKA = 4.35). Since the kinetic investigation has been carried out in the range of pH 2-3, it is reasonable to suppose that bipyfidyl is in the form of Bipy H +. Mecury(II) is known to form metal chelates with 1,10phenanthroline and 2,2'-bipyridyl. Andegg[4] has reported that the stability constants of Hg(II)-Bipyridyl complex are: log K~ =9.64, log K2 =7.10 and log K3 = 2.8 (log f13 ~ 19.54). Thus there is autoinhibition due to the removal of 2,2'-bipyridyl through complex formation with Hg(II) which is the product of the reaction, (Table 2.) Hence, in all the kinetic runs, 2,2'-bipyridyl was taken in large excess (more than 175 times) in comparison to Fe(III) and the decrease in the concentration of the ligand removed by complex formation with Hg(II) and Fe(II) is therefore negligible, thereby preventing the complications due to autoinhibition. To explain the kinetic results, the authors propose that Fe(III) when mixed with bipyridyt forms 1:2 complex with bipyridyl ligand and the complex is a better oxidizing species than uncomplexed Fe(III). The greater oxidizing ability of Fe(III) in the presence of bipyridyl is presumable due to its stabilizing the lower valence state Fe(II) through back-bonding with a consequent increase in the redox potential of Fe(III)/Fe(II) system. But the acceleration by 2,2'-bipyridyl may also be due to the possibility that electron transfer is facilitated by 7r-electron system present in the ligand. To account for the detailed dependence of rate on the concentrations of Hg22+, Bipyridyl and Fe(III), the authors propose the following scheme.
O
~+~ %
'7 02 ,..-t o
Et
O
,-t
I
r.)
CO r-t
.g~D O
B~
~O
tO tc'x
O
t~
,.-4
,4
E) r~
4 .=_ o
lcx
r.4
4
1183
r.) v
.rig ,--t <~
t~
o~
~o
Fe3aq+ 2 Bipy H+. K [Fe(Bipy)2], + + 2 H + (Complex)
(I)
1,4 4~ C
"E
Complex + (Hg-Hg) 2+
k
slow
Ferrodiin + Hg ~÷ + Hg 2+
(H~u O~
(2) .,4
Complex + Hg I+ rex i-4
0"~
o .~® v
-o gt v o
O
fast ~ Ferrodiin + Hg2÷.
(3)
'~¢xl
The rate-determining step in the proposed mechanism is the breaking of the (Hg-Hg) 2+ bond, concomitant with the transfer of an electron; the Hg 1+ formed reacts with complex in a fast step to form Hg2÷. In the case of oxidation of Hg(I) by Tl(III)[8] and Mn(III)[9], the formation of elemental mercury through the step H
g22 + ~ H g TT O -. H g 2+
(4)
followed by rate-determining oxidation of Hg ° has been proposed. But such a mechanism is not plausible in the present reaction in view of lack of any effect of Hg(II) on the rate, apart from the inhibitory effect due to the removal of bipyridyl by way of complex formation. The mechanism proposed now is similar to the one proposed by McCurdy et aL[10], in the case of oxidation of Hg(I) by Ce(IV), who also did not observe any inhibitory effect of Hg(II) on the rate.
1184
P. V. S. RAO et al.
Rate-law, The mechanism leads to the rate-law,
d[Fe(III)] _ kK[Fe(III)][Bipy]2[Hg~2+] dt K[Bipy]2 + [H+]2 under the conditions [Bipy] and [Hg22+]>>[Fe(III)] the pseudo first order rate constant, kl, is given by kK[Bipy]2[Hg22+] k, = K[Bipy]2 + [H+]2 taking reciprocals one gets I
E =~
I
[H+] 2
+ kK[Bipy]2[Hg22+]"
According to the rate-law, the plots of 1/k~ vs 1/[Bipy]2 and l/k1 vs [H+]2 should be straight lines with positive intercepts on the l/kraxes, which were actually observed. The values of k and K have been calculated from the values of the slopes and intercepts of the plots of l/kj vs 1/[Bipy]2 at three different temperatures. Using the value of k, i.e. the rate constant for the rate-limiting step at three different temperatures, the activation energy, E, of rate-limiting has been computed from the plot of log k vs I/T. Using this value of E, the entropy of activation
(AS") has also been computed. From the values of the stability constant (K) of the iron(III)-bipyridyl complex determined at three temperatures, enthalpy, the free energy and entropy of formation of this complex have been calculated. All these results are presented in Table 4. Acknowledgements--Two of the authors (K. V. S. and P. S. N. M) thank the Council of Scientific and Industrial Research, India and the University Grants Commission for the award of Research Fellowships.
REFERENCES 1. F. Blau, Monatsh. Chem. 19, 647 (1898). 2. A. Gaines, Jr., L. P. Hammett and G. H. Walden, J. Am. Chem. Soc. 58, 1668(1936). 3. R. M. Milburn and W. C. Vosburgh, J. Am. Chem, Soc. 77, 1352 (1955). 4. G. Anderegg, Helv. Chim. Acta 46, 2397 0963). 5. T. H. Siddall and W. C. Vosburgh, J. Am. Chem. Soc. 73, 4270 (1951). 6. L. Ciavatta and M. Grimaldi, J. Inorg. Nucl. Chem. 37, 163 (1950). 7. J. H. Baxendale and P. George, Trans. Faraday Soc. 46, 55 (1950). 8. A. M. Armstrong and J. Halpern, Can. J. Chem. 35, 1020 (1957). 9. D. R. Russeinsky, J. Chem. Soc. 1181 0963). 10. W. H. McCurdy, Jr. and G. G. Guilbault, J. Phys. Chem. 64, 1825 (1960).
KINETICS OF THE OXIDATION OF Hg(I) BY Fe(III)-2,2'BIPYRIDYL COMPLEX P. V. SUBBA RAO,* M. SURYANARAYANA,K. V. SUBBAIAH and P. S. N. MURTY Department of Chemistry, Andhra University, Waltair-530 003, India
(Received for publication 22 October 1979) Abstract--The kinetics of the oxidation of Hg(I) by the Fe(III)-2,2'-bipyridylcomplex has been investigated spectropbotometrically. The reaction obeys first order kinetics with respect to Hg(I) and Fe(III). Plots of l/kt vs l/[2,2'-Bipyridyl]z and l/kt vs [H+]2 are straight lines with positive intercepts on the llkt-axes. A reaction mechanism is proposed in which the rate-determining step is believed to involve concomitant breaking of the (Hg-Hg)2+ bond and electron transfer. Hg(II) solution: 0.05M solution of Hg(II) was prepared by dissolving mercuric nitrate in 0.5 M nitric acid. All solutions were prepared from deionised water and analytical reagent grade chemicals. The source of H+ ion was nitric acid. Kinetic procedure. The course of the reaction was followed by measuring the absorbance of the product ferrodiin at 510nm using a Hilger UV Speck spectrophotometer. At this wavelength all the other materials concerned have negligibleabsorption. To avoid any effect of light the reaction vessel was coated outside with black paint. The ionic strength was maintained constant with sodium nitrate. The kinetic runs were performed under the conditions [Hg(I)] and [Bipy]§>>[Fe(III)], thus isolating the latter. Kinetic orders. A plot of log (D®-/9,) vs time (where D~ and D, are absorbances at infinite time and at time t respectively) is a straight line under the above conditions, showing first order EXPERIMENTAL behaviour with respect to Fe(III). The pseudo first order rate Materials. Hg(l) solution, 0.01 M: Hg(1) nitrate solution was constant, kl, was calculated from the slope. prepared in 0.025 M nitric acid and standardized titrimetrically Effect of Hg(I) concentration. The pseudo first order rate with standard ferric alum in the presence of thiocyanate. constants, k~, were determined at different concentrations of Fe(HI) perchlorate: Iron wire (Riedel, Analar) was digested Hg(I), keeping the concentrations of other materials constant. with 70% perchloric acid (G. R. E. Merck) and evaporated to The data were presented in Table 1. The results show that k I is crystallization. The material was twice reccrystallized from 60% directly proportional to the concentration of Hg(I), i.e. the perchloric acid and stored under vacuum. An approximately reaction is first order with respect to Hg(I). 0.1 M solution was prepared in perchloric acid. The acid content Effect of 2,2'-bipyndyl concentration, k~ was also determined and Fe(llI) in the solution were determined by the method varying the concentrations of 2,2'-bipyridyl and at three temdescribed by Milburn et al.[3]. peratures, keeping [Fe(III)], [Hg22+] and [H+] constant. The plots 2,2'-bipyridyl solution: 0.05M solution was prepared by dis- of 1/kl vs l/[Bipy]2 were straight lines with positive slopes and solving in 50% (v/v) methanol. intercepts on the 1/kl axis (Fig. 1). Effect of acid concentration. The values of k~ were determined at different hydrogen ion concentrations keeping the ionic tFerridiin denotes Fe(III)-tris(bipyridyl)complex. strength constant with sodium nitrate. Hydrogen ions were found ;tFerrodiindenotes Fe(II)-trisCoipyridyl)complex. to inhibit the reaction. A plot of 1/kl vs [H+]2 is a straight line §Bipy is the abbreviation used for 2,2'-bipyridyl. with positive intercept on the 1/kl axis (FIB. 2). INTRODUCTION
Blau[l] was one of the first to detect the complex between Fe(lll) and 2,2'-bipyridyl by direct reaction. The complex is different from ferridiin,t which cannot be obtained by direct mixing, but only by the oxidation of ferrodiin~[2]. A study of the oxidizing action of this complex has not been made. As a part of the investigation of the reactions of this complex, we have studied the kinetic study of the oxidation of Hg(1) by Fe(IH) in the presence of 2,2'-bipyridyl. The results show that the active oxidizing species is the 1:2 complex of Fe(lll) and 2,2'-bipyridyL Fe(lll) in the absence of 2,2'bipyridyl does not oxidize Hg(1) under the experimental conditions employed.
Table 1. Effect of Hg(1)concentration on the rate of oxidation: [Fe(III)] = 1.0 x 10-4 M; [H +] = 1.0 x 10-2 M;/~ = 0.1; [Bipy] = 2.0 x 10-2 M; MeOH = 20% (v/v); Temp = 35 _+0.1°C
[Hg[+] ~ io3, s
k I ~ 1o3, .i~-l
kl/ (ag~+~
0.8 1.0 1.3 1.6 2.0
5.O 6.3 8.1 i0.0 12.5
6.25 6.30 6.27 6.25 6.25
2.3
14.5
6.~0
1181
P. V. S. RAO et al.
1182
Table 2. Effect of Hg(II) concentration on rate: [Fe(IIl)] = i.0 x I0-~ M; [Hg22+]= 2.0 x 10-3 M; MeOH = 25% (v/v); [H÷] = 2.0 x 10-2 M; Temp = 35 + 0.1°C;/~ = 0.1 [Hg 2+] x 10 5, M
[Blpy] • 10 2 , M
k I x i0 ~, mln "I
0.0
2.00
4.15
0.5 1.0
2.00 2.00
5.80 3.10
1.5 2.0
2.00 2.00
2.75 2.25
0.5
2.15
4.12
1.0
2.90
4.20
1.5 2.0
2.45
4.15 4.10
2.60
Effect of Hg(II) concentration. Kinetic runs were carried out at different initial concentrations of Hg(II). The data presented in Table 2 show that the rate decreases with increasing concentration of Hg(II). Since Hg(II) forms a stable 1:3 complex with bipyridyl (log t3 = 19.54)[4], the authors believe that the inhibition is due to the decrease in the concentration of free bipyridyl ligand because of complex formation with Hg(II). This view is supported by the fact that when the decrease in the concentration of bipyridyl is compensated by adding an additional amount of bipyfidyl (three times to that of added Hg(II)) the rate of the reaction is practically unaffected by the added Hg(II) (Table 2). Influence of ionic strength. Kinetic runs were carried out at different ionic strengths by varying the concentration of sodium nitrate in the range 0.1-0.9. The rate has been found to increase with increase in ionic strength indicating that the rate-determining step involves ions of similar charge. The data were presented in Table 3. DISCUSSION
Nature of species of Fe(III) The results of Milburn and Vosburgh[3] show that Fe(III) at concentrations less than 1.0 x IO-3M does not
3.6 3.2
Table 3. Influence of ionic strength on rate: [Fe(III)]= 1.0 x 10-4 M; [H+] = 1.0 x l0-2 M; [Hg22+]= 1.0 x 10-3 M; [Bipy] = '2.0 x 10-2 M; MeOH = 20%(v/v); Temp = 35 + 0.1°C Ionic
strength,
~
k I x 10 3 , m l n -I
o.1 o.5 0.5 0.7 0.9
630 7.Ol 7.83 8.57 9.25
dimerize. A similar observation has also been made by Siddall and Vosburg[5]. Ciavatta and Grimaldi[6] have reported the negative logarithm of the hydrolysis constant of Fe(III) to be 3.1-+0.2. In the present investigation, the Fe(III) concentration employed was the order of 1 0 - ' M and hydrogen ion concentration was in the range 0.01-0.02 M. Hence it is reasonable to assume that there is no appreciable dimerization under these conditions, and the concentration of hydroxylated Fe(III)
2.4
2.8 2.0
2,4 1.6
2.0 .=_ :E
1.2
'9 1.2
.w.. 0 . 8
0.8 0.4 O.4
II I0-~' [81PY] 2, M -2
Fig. 1. Plots of l/k1 vs 1/[Bipy]2. [Fe(I[])]= 1.0xI0-4M; MeOH = 25% (v/v). [Hg22+l-- 1.0 x 10-3 M; [H÷] = 1.0 x 111-2M; (9-49, 31~C; A--A, 35°C; [B-..~, 41PC.
I 2 104x[H+] 2, M 2
I 3
I 4
Fig 2 Plot of l/k, vs [H+] = [Fe(I[])] = 10 x 10-4 M; MeOH = 20%(v/v). [Hg22.] = 20 x 10-3 M; [Bipy] = 2 0 x 10-= M ; / ~ = 002;
Temperature = 35-+O.1°C.
Kinetics of the oxidation of Hg(I) by Fe(lll)-2,2'-bipyridyl complex
has been neglected in view of the low hydrolysis constant. The hydrolysis is also suppressed, since the ligand (2,2'-bipyridyl) was present in large excess. It has been reported [7] that 2,2'-bipyridyl exists in singly protonated form (pKA = 4.35). Since the kinetic investigation has been carried out in the range of pH 2-3, it is reasonable to suppose that bipyfidyl is in the form of Bipy H +. Mecury(II) is known to form metal chelates with 1,10phenanthroline and 2,2'-bipyridyl. Andegg[4] has reported that the stability constants of Hg(II)-Bipyridyl complex are: log K~ =9.64, log K2 =7.10 and log K3 = 2.8 (log f13 ~ 19.54). Thus there is autoinhibition due to the removal of 2,2'-bipyridyl through complex formation with Hg(II) which is the product of the reaction, (Table 2.) Hence, in all the kinetic runs, 2,2'-bipyridyl was taken in large excess (more than 175 times) in comparison to Fe(III) and the decrease in the concentration of the ligand removed by complex formation with Hg(II) and Fe(II) is therefore negligible, thereby preventing the complications due to autoinhibition. To explain the kinetic results, the authors propose that Fe(III) when mixed with bipyridyt forms 1:2 complex with bipyridyl ligand and the complex is a better oxidizing species than uncomplexed Fe(III). The greater oxidizing ability of Fe(III) in the presence of bipyridyl is presumable due to its stabilizing the lower valence state Fe(II) through back-bonding with a consequent increase in the redox potential of Fe(III)/Fe(II) system. But the acceleration by 2,2'-bipyridyl may also be due to the possibility that electron transfer is facilitated by 7r-electron system present in the ligand. To account for the detailed dependence of rate on the concentrations of Hg22+, Bipyridyl and Fe(III), the authors propose the following scheme.
O
~+~ %
'7 02 ,..-t o
Et
O
,-t
I
r.)
CO r-t
.g~D O
B~
~O
tO tc'x
O
t~
,.-4
,4
E) r~
4 .=_ o
lcx
r.4
4
1183
r.) v
.rig ,--t <~
t~
o~
~o
Fe3aq+ 2 Bipy H+. K [Fe(Bipy)2], + + 2 H + (Complex)
(I)
1,4 4~ C
"E
Complex + (Hg-Hg) 2+
k
slow
Ferrodiin + Hg ~÷ + Hg 2+
(H~u O~
(2) .,4
Complex + Hg I+ rex i-4
0"~
o .~® v
-o gt v o
O
fast ~ Ferrodiin + Hg2÷.
(3)
'~¢xl
The rate-determining step in the proposed mechanism is the breaking of the (Hg-Hg) 2+ bond, concomitant with the transfer of an electron; the Hg 1+ formed reacts with complex in a fast step to form Hg2÷. In the case of oxidation of Hg(I) by Tl(III)[8] and Mn(III)[9], the formation of elemental mercury through the step H
g22 + ~ H g TT O -. H g 2+
(4)
followed by rate-determining oxidation of Hg ° has been proposed. But such a mechanism is not plausible in the present reaction in view of lack of any effect of Hg(II) on the rate, apart from the inhibitory effect due to the removal of bipyridyl by way of complex formation. The mechanism proposed now is similar to the one proposed by McCurdy et aL[10], in the case of oxidation of Hg(I) by Ce(IV), who also did not observe any inhibitory effect of Hg(II) on the rate.
1184
P. V. S. RAO et al.
Rate-law, The mechanism leads to the rate-law,
d[Fe(III)] _ kK[Fe(III)][Bipy]2[Hg~2+] dt K[Bipy]2 + [H+]2 under the conditions [Bipy] and [Hg22+]>>[Fe(III)] the pseudo first order rate constant, kl, is given by kK[Bipy]2[Hg22+] k, = K[Bipy]2 + [H+]2 taking reciprocals one gets I
E =~
I
[H+] 2
+ kK[Bipy]2[Hg22+]"
According to the rate-law, the plots of 1/k~ vs 1/[Bipy]2 and l/k1 vs [H+]2 should be straight lines with positive intercepts on the l/kraxes, which were actually observed. The values of k and K have been calculated from the values of the slopes and intercepts of the plots of l/kj vs 1/[Bipy]2 at three different temperatures. Using the value of k, i.e. the rate constant for the rate-limiting step at three different temperatures, the activation energy, E, of rate-limiting has been computed from the plot of log k vs I/T. Using this value of E, the entropy of activation
(AS") has also been computed. From the values of the stability constant (K) of the iron(III)-bipyridyl complex determined at three temperatures, enthalpy, the free energy and entropy of formation of this complex have been calculated. All these results are presented in Table 4. Acknowledgements--Two of the authors (K. V. S. and P. S. N. M) thank the Council of Scientific and Industrial Research, India and the University Grants Commission for the award of Research Fellowships.
REFERENCES 1. F. Blau, Monatsh. Chem. 19, 647 (1898). 2. A. Gaines, Jr., L. P. Hammett and G. H. Walden, J. Am. Chem. Soc. 58, 1668(1936). 3. R. M. Milburn and W. C. Vosburgh, J. Am. Chem, Soc. 77, 1352 (1955). 4. G. Anderegg, Helv. Chim. Acta 46, 2397 0963). 5. T. H. Siddall and W. C. Vosburgh, J. Am. Chem. Soc. 73, 4270 (1951). 6. L. Ciavatta and M. Grimaldi, J. Inorg. Nucl. Chem. 37, 163 (1950). 7. J. H. Baxendale and P. George, Trans. Faraday Soc. 46, 55 (1950). 8. A. M. Armstrong and J. Halpern, Can. J. Chem. 35, 1020 (1957). 9. D. R. Russeinsky, J. Chem. Soc. 1181 0963). 10. W. H. McCurdy, Jr. and G. G. Guilbault, J. Phys. Chem. 64, 1825 (1960).