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Kinetics of oxidation of iodide by vanadium(V)—oxalic acid catalysis
J ,m~rt, nu,! C i w m 1978, Vol 40, pp 29",-297. Pergamon Press.
Printed in Great Britain
KINETICS OF OXIDATION OF IODIDE BY VANADIUM(V)-OXALIC ACID CATALYSIS P. V. SUBBARAO. P. S. N. MURTY and R. V. :q. MURTY Department of Chemistr}, Andhra University, Waltair-530003. India
(Received 5 April 19771 Abstract--The kinetics of the oxidation of iodide by vanadium(V) catabsed by oxalic acid was studied in perchloric acid medium at constant ionic strength. The reaction was found to be first order with respect to both the vanadium(V) and iodide. A plot of l/k~ versus I/[OX]*" is a straight line. Hydrogen ions were found to accelerate the reaction. Increase of ionic strength resulted in a decrease in reaction rate suggesting that the rate determining step involves a cation and anion. The kinetic pattern is in favour of the formation of a 1:2 complex between vanadium(V) and oxalic acid which is supposed to be a more powerful oxidant. Activation parameters of the reaction and the stability constant of the complex were calculated from the kinetic data.
INTRODUCTION
Although the efficacy of oxalic acid to catalyse vanadium(V) oxidations is k n o w n [ l - 3 ] , the kinetics and m e c h a n i s m of oxalic acid catalysis have not been studied so far. We report a kinetic study of the oxalic acidcatalysed oxidation of iodide ion by vanadium(V). The analytical application of this reaction in the iodometric estimation of vanadium(V) has been reported by Rao et al. The kinetics of the uncatalysed reaction have already been investigated [4-9].
EXPERIMENTAL
Materials Vanadiumll/) solution. An approx. 0.1 M solution was prepared from ammonium vanadate and sodium carbonate. This was standardised with iron(II) sulphate which was previously standardised with potassium dichromate. Iodide solution. An approx. 0.l M solution was prepared in water and standardised with standard silver(l) nitrate. Oxalic acid solution. 0.1 M solution was prepared by dissolving E. Merck pro-analyst grade oxalic acid in water. Sodium perchlorute solution. 5 M sodium percfilorate was prepared by mixing equal volumes of 5 M sodium hydroxide and 5 M perchloric acid. All the other materials used were of analytical reagent grade. Kinetic procedure. The course of the reaction was followed by titrating the iodine formed from time to time against standard sodium thiosulphate solution using 1% starch as indicator. The ionic strength was maintained constant by the addition of sodium perchlorate. Most of the runs were performed under isolation conditions taking a large excess of iodide in comparison with vanadium(Vl. A small but constant amount of EDTA (2.0x 10 ~M/ was present in all runs to eliminate the possibility of catalysis by traces of heavy metal ions. 1"he reaction rate was found to be uneffected by the addition of this small amount of EDTA. Stoichiomet O, of the reaction. The stoichiometry of the reaction was determined by adding a known excess of vanadium(V} to potassium iodide in the presence of 0.0l M oxalic acid and at pH 1.0. After the reaction is complete, the unreacted vanadium(VI was determined by titration against standard iron(ll). We found one mole of vanadium(V) reacts with one mole of iodide ion. This result extablishes that no induced oxidation of iodide ion by atmospheric oxygen takes place and further, oxalic acid is not consumed through oxidation by vanadium(V). Thus oxalic acid acts merely as a catalyst. Kinetic investigation has been made under these experimental conditions to avoid the complications due to aerial oxidation. Separate 295
experiments have shown that oxalic acid is not oxidised to a detectable extent under these conditions. Kinetk" orders. A plot of log(f.,-/',) vs time /where )", is the titre value at infinite time and T, is the titre value at time t/i~ a straight line tinder the conditions 11 ,>lVVl. 4thus isolating vanadium(V)) showing first order kinetics with respect to vanadium(V). The pseudo first order rate constant k~ was calculated from the slopes of these plots. The pseudo first order rate constants kt were determined at different concentrations of iodide km keeping oxalic acid concentration and pH constant. The data are presented in TaMe 1. It is evident from the data that k, is directly proportional to the concentration of iodide ion showing thai the reaction is first order with respect to iodide ion. kt was also determined at different oxalic acid concentrations keeping the concentration of iodide ion and the pH constant. The quantitative relation between the rate and oxalic acid concerntration is shown by the ,,traight line obtained with a positwe intercept on the I/k~ axis when Ilk) is plotted against I/[OX]-~: [OX] represents the oxalic acid concentration. lnfluence of hydrogen ion concentration: The rate data obtained by varying the concentration of hydrogen ion are presented in Table I. The data shov, that the reaction is catalysed bx h~drt~2:en ions. Influence of ionic streneth. The values of k~ determined at different ionic strengths are presented in Table 2. The data show that rate decreases with increase in ionic strength. No correlation of rate with ionic strength has been made because the Bronsted equation is not applicable. However it may be inferred from the observed decrease of rate with increase in ionic strength that the rate determining step is a cation-anion reaction. Similar behaviour was also observed in the uncatalysed oxidation of iodide ion by vanadium(V)[7]. Activation parameters. Activation parameters were determined by measuring the second order rate constants at four different temperatures. The values are presented in "Fable 3 A plot of log k vs I/T was a straight line. The activation energ?~ E and entropy of activation AS ~ are found to be 6.0-+ 0.5 kcal mole ~ and 53-+ I e.u. respectively. For the uncatalysed reaction an activation energy of l l.7kcalmole ~ has been reported[6l: thu,. the catalyst oxalic acid reduces the activation energy by 5 kcal.
DISCUSSION
Rossotti and Rossotti[10] have shown that vanadium(V) ;at pH ~<1 mostly exists in the form of VO2 + which with increasing hydrogen ion concentration is believed to be transformed into V(OH)32~. Waters and J o n e s [ I l l reported evidence for the formation of u 1:2
296
P. V. SUBBARAO
etal.
Table 1. Effect of iodide and effect of acid on rate constants. [Vv] = 2.0 × 10-3 M, [OX] = 1.0× 10 2M, Ix =0.412, temp.= 25±0.1°C SI.No.
[I-] ×
1. 2. 3. 4. 5. 6. 7. 8.
10 2 M
[H +] M K I x 103 sec i
2.6 2.2 1.8 1.4 1.0 1.0 1.0 1.0
0.1 0.1 0.1 0.1 0.1 0.2 0.3 0.4
12.4 10.3 8.7 6.8 5.0 7.4 8.6 9.8
K 1X
kKIK2[VV][OX]2[H+][I -]
dt
1 + K1[H +] + KtK2[H+][OX] 2
or
kl =
kKIKz[OXI2[H+][I -] 1 + K1[H +] + K1K2[H+][OX] 2
and
1
1
1
1/kl = kKIK2[OX]2[H+][I _] ~ kK2[OX]2[I_] ~ k[I-]
Table 2. Effect of ionic strength on rate constants. [Vv] = 2.0 × 10-3 M, [OX] = 1.0 × 10-2 M, [I ] = 1.0 × 10 -2 M, [H +] = 0.1 M, temp. = 25.0 ± 0.1°C Ionic strength (ix)
d[V v]
thus plots of l/k, vs I/[OX] 2 and Ilk, vs I/[H +] must be straight lines with positive intercepts. This has actually been observed (Figs. 1 and 2) showing the validity of the
103 sec-t 2.£
0.412 0.612 0.812 1.012 1.212
5.0 4.6 4.3 4.1 3.8
1.6
? _o Table 3. Activation parameters and effect of temperature on the rate constants. [Vv] = 5.0 × 10-3 M, [I-] = 5.0 x 10 3 M, [OX] = 1.0 × 10-2 M, [H +] = 0.1 M, ix = 0.11 Temperature (°C)
k × 104 (1.mol-j sec -~)
E (kcals)
AS ~' (e.u. at 26°C)
26 30 35 40
8.7 10.0 11.4 14.0
6.0±0.5
-53±1
1.2
x
- I ~ - o. e
0-4
I
Eox:~ x ,o -~
complex of oxalic acid with vanadium(V) at acidities below 4 M. K 1
M-z
Fig. 1. Plot of reciprocal rate constant vs reciprocal of [OX]:, [H+]=0.1M [VV]=2.0×10 3M [I-]=l.0×10 2M ix=0.412, temp = 25 _+0.1°C.
V02 + + H + + H20 7 - " V(OH)32+ 2.4 OC-O V(OH)3 2 + + 2 O X .
,,2 •
OH
I \
+
/
2.0
/v\ OC-O
O.CO.COOH 1.6
complex t~
The authors believe that such a complex is formed and acts as an active oxidising species. This complex is a stronger oxidant than uncomplexed vanadium(V) because the carboxyl groups being highly electrophilic render vanadium(V) more strongly oxidising; the ~r-electron character of oxalate ligand provides easy way for electron transfer. Further, oxalic acid forms stronger complex with vanadium(IV) increasing the red-ox potential of the v V / v TM system.
I-2 x
-I..~ 0-8
O-~
slow
complex + I
) v'Vox2 + I k
M -I
[.+1
fast
I+I This leads to the rate law
) 12.
Fig. 2. Plot of reciprocal rate constant vs reciprocal .of [H+], [VV]=2.0×10 3M [OX]= 1.0×10 2M [I ]=1.0×10 ZM ix= 0.412, temp. = 25 ± 0.1°C.
Kinetics of oxidation of iodide by vanadium(V)-oxalic acid catalysis rate law and the mechanism. F r o m the slopes and intercepts of these plots the values of K~, K2, K~K2 and k were calculated to be K, = 1.74 M ' K~=2.33×104M
2
K,K~ = 4.05 × 104 M ~ k-1.92M
'sect
Acknowledgements--Two of us (P.S.N. and R.V.S.) are grateful ~o UG.C. and ('.S.IR India for the award of Fellowships. REFERENCES I G. G. Rao, C. R. Viswandham and J. V. S. R. Anjaneyulu, Pr~c. Natl. lnst. Sci. (India) II, 333 (1945).
2t)7
2. G. G. Rao and C. R. Viswanadham, Curr. Sci. 12, 306 (1943). 3. G. G. Rao and M. N. Sastry, Curr. Sci. 21, 189 (1952). 4. W. C. Bray and J. B. Ramsay, J. Am. Chem. Soc. 55. 2279 (1933). 5. D. R. Rosseinsky, J. lnorg. Nucl. Chem. 33, 3976 (19711. 6. G. St. Nikolm and D. Mihailova. J. lnor~,. Nucl. ('hem. 31. 2499 (1969). 7. F. Secco. S. Celsi and C. Grati, J. Chem. Soc. I Daltotl 7"raJl~. 1675 (1972). 8. F. Secco, S. Celsi and M. Ventumini, J. Chem. Soc. 7!)3 (1974). 9. R. Jrutille and B. Bodriquez, Riogs. Alonso 7, 139 (1943!. 10. F. J. C, Rossolti and H. Rossotti, Acta. Chem. &'and 10, 9::;7 (1956). 11. J. R. Jones and W. A. Waters. J. ('hem. Soc. 4757 (1%1i.
Printed in Great Britain
KINETICS OF OXIDATION OF IODIDE BY VANADIUM(V)-OXALIC ACID CATALYSIS P. V. SUBBARAO. P. S. N. MURTY and R. V. :q. MURTY Department of Chemistr}, Andhra University, Waltair-530003. India
(Received 5 April 19771 Abstract--The kinetics of the oxidation of iodide by vanadium(V) catabsed by oxalic acid was studied in perchloric acid medium at constant ionic strength. The reaction was found to be first order with respect to both the vanadium(V) and iodide. A plot of l/k~ versus I/[OX]*" is a straight line. Hydrogen ions were found to accelerate the reaction. Increase of ionic strength resulted in a decrease in reaction rate suggesting that the rate determining step involves a cation and anion. The kinetic pattern is in favour of the formation of a 1:2 complex between vanadium(V) and oxalic acid which is supposed to be a more powerful oxidant. Activation parameters of the reaction and the stability constant of the complex were calculated from the kinetic data.
INTRODUCTION
Although the efficacy of oxalic acid to catalyse vanadium(V) oxidations is k n o w n [ l - 3 ] , the kinetics and m e c h a n i s m of oxalic acid catalysis have not been studied so far. We report a kinetic study of the oxalic acidcatalysed oxidation of iodide ion by vanadium(V). The analytical application of this reaction in the iodometric estimation of vanadium(V) has been reported by Rao et al. The kinetics of the uncatalysed reaction have already been investigated [4-9].
EXPERIMENTAL
Materials Vanadiumll/) solution. An approx. 0.1 M solution was prepared from ammonium vanadate and sodium carbonate. This was standardised with iron(II) sulphate which was previously standardised with potassium dichromate. Iodide solution. An approx. 0.l M solution was prepared in water and standardised with standard silver(l) nitrate. Oxalic acid solution. 0.1 M solution was prepared by dissolving E. Merck pro-analyst grade oxalic acid in water. Sodium perchlorute solution. 5 M sodium percfilorate was prepared by mixing equal volumes of 5 M sodium hydroxide and 5 M perchloric acid. All the other materials used were of analytical reagent grade. Kinetic procedure. The course of the reaction was followed by titrating the iodine formed from time to time against standard sodium thiosulphate solution using 1% starch as indicator. The ionic strength was maintained constant by the addition of sodium perchlorate. Most of the runs were performed under isolation conditions taking a large excess of iodide in comparison with vanadium(Vl. A small but constant amount of EDTA (2.0x 10 ~M/ was present in all runs to eliminate the possibility of catalysis by traces of heavy metal ions. 1"he reaction rate was found to be uneffected by the addition of this small amount of EDTA. Stoichiomet O, of the reaction. The stoichiometry of the reaction was determined by adding a known excess of vanadium(V} to potassium iodide in the presence of 0.0l M oxalic acid and at pH 1.0. After the reaction is complete, the unreacted vanadium(VI was determined by titration against standard iron(ll). We found one mole of vanadium(V) reacts with one mole of iodide ion. This result extablishes that no induced oxidation of iodide ion by atmospheric oxygen takes place and further, oxalic acid is not consumed through oxidation by vanadium(V). Thus oxalic acid acts merely as a catalyst. Kinetic investigation has been made under these experimental conditions to avoid the complications due to aerial oxidation. Separate 295
experiments have shown that oxalic acid is not oxidised to a detectable extent under these conditions. Kinetk" orders. A plot of log(f.,-/',) vs time /where )", is the titre value at infinite time and T, is the titre value at time t/i~ a straight line tinder the conditions 11 ,>lVVl. 4thus isolating vanadium(V)) showing first order kinetics with respect to vanadium(V). The pseudo first order rate constant k~ was calculated from the slopes of these plots. The pseudo first order rate constants kt were determined at different concentrations of iodide km keeping oxalic acid concentration and pH constant. The data are presented in TaMe 1. It is evident from the data that k, is directly proportional to the concentration of iodide ion showing thai the reaction is first order with respect to iodide ion. kt was also determined at different oxalic acid concentrations keeping the concentration of iodide ion and the pH constant. The quantitative relation between the rate and oxalic acid concerntration is shown by the ,,traight line obtained with a positwe intercept on the I/k~ axis when Ilk) is plotted against I/[OX]-~: [OX] represents the oxalic acid concentration. lnfluence of hydrogen ion concentration: The rate data obtained by varying the concentration of hydrogen ion are presented in Table I. The data shov, that the reaction is catalysed bx h~drt~2:en ions. Influence of ionic streneth. The values of k~ determined at different ionic strengths are presented in Table 2. The data show that rate decreases with increase in ionic strength. No correlation of rate with ionic strength has been made because the Bronsted equation is not applicable. However it may be inferred from the observed decrease of rate with increase in ionic strength that the rate determining step is a cation-anion reaction. Similar behaviour was also observed in the uncatalysed oxidation of iodide ion by vanadium(V)[7]. Activation parameters. Activation parameters were determined by measuring the second order rate constants at four different temperatures. The values are presented in "Fable 3 A plot of log k vs I/T was a straight line. The activation energ?~ E and entropy of activation AS ~ are found to be 6.0-+ 0.5 kcal mole ~ and 53-+ I e.u. respectively. For the uncatalysed reaction an activation energy of l l.7kcalmole ~ has been reported[6l: thu,. the catalyst oxalic acid reduces the activation energy by 5 kcal.
DISCUSSION
Rossotti and Rossotti[10] have shown that vanadium(V) ;at pH ~<1 mostly exists in the form of VO2 + which with increasing hydrogen ion concentration is believed to be transformed into V(OH)32~. Waters and J o n e s [ I l l reported evidence for the formation of u 1:2
296
P. V. SUBBARAO
etal.
Table 1. Effect of iodide and effect of acid on rate constants. [Vv] = 2.0 × 10-3 M, [OX] = 1.0× 10 2M, Ix =0.412, temp.= 25±0.1°C SI.No.
[I-] ×
1. 2. 3. 4. 5. 6. 7. 8.
10 2 M
[H +] M K I x 103 sec i
2.6 2.2 1.8 1.4 1.0 1.0 1.0 1.0
0.1 0.1 0.1 0.1 0.1 0.2 0.3 0.4
12.4 10.3 8.7 6.8 5.0 7.4 8.6 9.8
K 1X
kKIK2[VV][OX]2[H+][I -]
dt
1 + K1[H +] + KtK2[H+][OX] 2
or
kl =
kKIKz[OXI2[H+][I -] 1 + K1[H +] + K1K2[H+][OX] 2
and
1
1
1
1/kl = kKIK2[OX]2[H+][I _] ~ kK2[OX]2[I_] ~ k[I-]
Table 2. Effect of ionic strength on rate constants. [Vv] = 2.0 × 10-3 M, [OX] = 1.0 × 10-2 M, [I ] = 1.0 × 10 -2 M, [H +] = 0.1 M, temp. = 25.0 ± 0.1°C Ionic strength (ix)
d[V v]
thus plots of l/k, vs I/[OX] 2 and Ilk, vs I/[H +] must be straight lines with positive intercepts. This has actually been observed (Figs. 1 and 2) showing the validity of the
103 sec-t 2.£
0.412 0.612 0.812 1.012 1.212
5.0 4.6 4.3 4.1 3.8
1.6
? _o Table 3. Activation parameters and effect of temperature on the rate constants. [Vv] = 5.0 × 10-3 M, [I-] = 5.0 x 10 3 M, [OX] = 1.0 × 10-2 M, [H +] = 0.1 M, ix = 0.11 Temperature (°C)
k × 104 (1.mol-j sec -~)
E (kcals)
AS ~' (e.u. at 26°C)
26 30 35 40
8.7 10.0 11.4 14.0
6.0±0.5
-53±1
1.2
x
- I ~ - o. e
0-4
I
Eox:~ x ,o -~
complex of oxalic acid with vanadium(V) at acidities below 4 M. K 1
M-z
Fig. 1. Plot of reciprocal rate constant vs reciprocal of [OX]:, [H+]=0.1M [VV]=2.0×10 3M [I-]=l.0×10 2M ix=0.412, temp = 25 _+0.1°C.
V02 + + H + + H20 7 - " V(OH)32+ 2.4 OC-O V(OH)3 2 + + 2 O X .
,,2 •
OH
I \
+
/
2.0
/v\ OC-O
O.CO.COOH 1.6
complex t~
The authors believe that such a complex is formed and acts as an active oxidising species. This complex is a stronger oxidant than uncomplexed vanadium(V) because the carboxyl groups being highly electrophilic render vanadium(V) more strongly oxidising; the ~r-electron character of oxalate ligand provides easy way for electron transfer. Further, oxalic acid forms stronger complex with vanadium(IV) increasing the red-ox potential of the v V / v TM system.
I-2 x
-I..~ 0-8
O-~
slow
complex + I
) v'Vox2 + I k
M -I
[.+1
fast
I+I This leads to the rate law
) 12.
Fig. 2. Plot of reciprocal rate constant vs reciprocal .of [H+], [VV]=2.0×10 3M [OX]= 1.0×10 2M [I ]=1.0×10 ZM ix= 0.412, temp. = 25 ± 0.1°C.
Kinetics of oxidation of iodide by vanadium(V)-oxalic acid catalysis rate law and the mechanism. F r o m the slopes and intercepts of these plots the values of K~, K2, K~K2 and k were calculated to be K, = 1.74 M ' K~=2.33×104M
2
K,K~ = 4.05 × 104 M ~ k-1.92M
'sect
Acknowledgements--Two of us (P.S.N. and R.V.S.) are grateful ~o UG.C. and ('.S.IR India for the award of Fellowships. REFERENCES I G. G. Rao, C. R. Viswandham and J. V. S. R. Anjaneyulu, Pr~c. Natl. lnst. Sci. (India) II, 333 (1945).
2t)7
2. G. G. Rao and C. R. Viswanadham, Curr. Sci. 12, 306 (1943). 3. G. G. Rao and M. N. Sastry, Curr. Sci. 21, 189 (1952). 4. W. C. Bray and J. B. Ramsay, J. Am. Chem. Soc. 55. 2279 (1933). 5. D. R. Rosseinsky, J. lnorg. Nucl. Chem. 33, 3976 (19711. 6. G. St. Nikolm and D. Mihailova. J. lnor~,. Nucl. ('hem. 31. 2499 (1969). 7. F. Secco. S. Celsi and C. Grati, J. Chem. Soc. I Daltotl 7"raJl~. 1675 (1972). 8. F. Secco, S. Celsi and M. Ventumini, J. Chem. Soc. 7!)3 (1974). 9. R. Jrutille and B. Bodriquez, Riogs. Alonso 7, 139 (1943!. 10. F. J. C, Rossolti and H. Rossotti, Acta. Chem. &'and 10, 9::;7 (1956). 11. J. R. Jones and W. A. Waters. J. ('hem. Soc. 4757 (1%1i.