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Mechanisms of redox reactions of Tlaq3+ and TlOHaq2+ with quinols in acidic perchlorate media
1. inorg, nucl. Chem., 1976, Vol. 38, pp. 2005-2007. Pergamon Press, Printed in Great Britain
MECHANISMS OF REDOX REACTIONS OF T13q AND T1OH~aq WITH QUINOLS IN ACIDIC PERCHLORATE MEDIA EZIO PELIZZETTI and EDOARDO MENTASTI Istituto di ChimicaAnaliticadell'UniversitL Via P. Giuria, 5. 10125Torino, Italy (Received 18 November 1975)
Abstract--Further evidence of an electron-transfer controlled mechanism in the reactions of TI~ and TIOH~ with quinols has been achieved by investigating2,3,5-trimethyl-and 2,3-dicyano-benzene-l,4-diol,in order to have a larger range of AG°. For both ionic species a linear relationshipbetween AG" and AG° has been observed with slope 0.37 and 0.39 respectively. INTRODUCTION THERE have been recent developments in the chemistry of aquo-thallium(III). Studies of the reactions of TI(II), generated by flash-photolysis, by Laurence et al., also shed light on the mechanisms of reaction TI(III) [1]. In fact this ion can react by one- or two-electron transfer mechanisms. For the latter type of reaction it has been observed that there is a slight dependence of the free energy of activation on the free energy change and, for the same AG °, AG e is larger than for the one-electron path. In previous papers we have reported the kinetics and mechanism of oxidation of a series of quinols with TI(III) [2]. In the present note we investigate the kinetics of the oxidation of two quinol derivatives, 2,3,5-trimethylbenzene-l,4-diol (TMHzQ) and 2,3-dicyano-benzene-l,4diol (DCH2Q), with TI(III). The comparison of the specific rate constants for the different reaction paths allows us to formulate some interesting hypotheses about the reaction mechanism of Tl(III) when behaving as a two-electron acceptor. EXPERIMENTAL Reagents and procedure
TMH2Q and DCH2Q (K&K) solutions were freshly prepared. The other reagents and general experimentalprocedure have been described[2]. The ionic strength was brought to 2.0 M (NaCIO4). The kinetic parameters were computed with a weighted leastsquare method. The standard deviation computed for the rate parameters ranged from 2 to 5%. RESULTS The second order rate constants, ko, at different temperatures and acidities are collected in Table 1. The kinetic data can be accounted for by the following scheme, as previously reported[2], where Kh is the first TI~
+S
' (TI. . . . S)a]+
K2
hydrolysis constant of TI~g[3], K1 and K2 the formation equilibrium constants for the precursor complexes (if an outer-sphere mechanism is operating they represent JINC Vol. 38 No. II--F
outer-sphere association constants) and Q is the pquinone corresponding to the quinol S. It follows that ko(l+Kh[H+]-')=k'K~+k"KzKh[H+] -'.
(1)
By plotting the left-hand side of eqn (1) as a function of [H+]-~, straight lines were obtained and from the intercept and slope, the terms k'K1 and k"K2 were evaluated; the specific rate constants are listed in Table 2 together with the activation parameters. DISCUSSION For electron transfer reactions of the outer-sphere kind, the Marcus theory[4] predicts the relationship:
+ (AG°r)Z/8(AGII + AG$:)
(2)
where AG* = AG" - w., AG °`= AG°+ w. - w.; Wr and w. are the work terms required to bring the reactants or products together in the transition state and AGT, and AG 7: are the free activation energies for electron transfer for the self-exchange reactions. The expression can be rearranged to give the gradient of a plot AG e vs AG Oin o the form 12(1 + AGr/4AGo), where AG~ is the value of AG e for AG ° equal to zero. It must be taken into account that it is difficult in many cases to know or to evaluate the exchange rates and the work terms for the reactants; thus by comparing the reaction rates of different substrates with the same ion the dependence of AG ~ vs AG °, can be different from the prediction of eqn (2). Moreover, eqn (2) was derived for reactions of the outer-sphere class, but similar relationships might exist in other cases[5]. The reaction mechanism and the relationship between AG e and AG O for the reactions involving both TI(III) and Tiff) have been recently discussed by Laurence et al.[1]. By means of flash-photolysis, the kinetic behavior of TI(II) and the evaluation of the standard redox potentials of the couples TI(III)/TI(II) and Tl(II)frl(I) were investigated; thus it was possible to assign a single two electron-transfer or two successive one-electron steps to a series of reactions involving TI(III). For the reactions involving a single two electron transfer (exchange TI(III)--TI(I), TI(III)-Hg(0), TI(III)-U(IV), TI(III)-V(II))[6], only a slight variation in AG" was observed for large AG ° changes; thus the
2O05
2006
EZIO PELIZZETTIand EDOARDO MENTASTI Table 1 Second-order rate constants (lmol-' s-')" at different temperatures§ TI~2Q H0104] :
b
0.30
0.50
1.OO
2.00
6.O
1.36x105
1.11x105
7.5x104
4.7x104
18.0
2.60x105
2.21x105
1.33x105
8.4x104
t
(*o)
DOH2 Q c
18.0
0.33
0.50
1.00
2.00
2.63
2.07
1.28
0.86
8.0
6.2
40.0
11.5
a averaged on at least five separate kinetic runs
b [~}I2Q]= 5.0X10-4 ; ~1(11T~ = 6 " 20xlO-4M. c [Doa2Q3. 5 ox~o 4 ~ ~ l C ~ I ~ .
50
• 20 0 , 1 0 - 3 =
tho vaZ=e, of K h have bee~ oo=p=tcd, f r o =
ref.
[3b], t o be
6 . 5 X 1 0 -2 at 6.00 , 7.3x 10-2 at 18.O 0 and 8 . 8 • 1 0 - 2 =
at 40.000.
Table 2 Specific rate constants (lmol-I s ' ) and activation parametersfor the reactionsof TMH2Qand DCH2QwithTI(III) t(oO)
k'K 1
k"K 2
TMH2Q
6.0
2.42x104
8.1x105
18.0
5.2 xlO 4
1.35x106
AH* a
9.9 + 1.7
6.5 _+ 2.1
-3.2 + 5.5
-8.8 _+ 6.7
b AS*
DOH2Q
18.0
0.58
12.5
40.0
3.4
62
~H# a
14.0 ~ 1.3
12.6 Z 1.5
&S$ b
-11.5 ± 4.3
-10.2 ~ 5.0
a kcal mol -I
b cal mol-ldeg-I
intrinsic effect predominates and no definite assignment could be made as to the role of thermodynamic factors. We have previously investigated the reaction of a series of quinols with TI(III) and a two electron-transfer mechanism has been suggested[2]. It has already been pointed out that the reaction rate is increased by electron releasing groups and reduced by electron withdrawing ones. The present data concerning TMH2Q and DCH2Q, together with the evaluation of the standard redox potentials for these compounds (see Table 3) allow us to formulate some conclusions about the role of the change
in the overall free energy of reaction on the reaction rate of TI(III) when a two electron-transfer mechanism is operating. In fact this series of reactions involves the same metal ion and a group of substrates differing in free energy due to different substituents lying far from the reaction site, so that the intrinsic factors (AG 7, and w) can be assumed constant. The data obtained are plotted in Fig. 1 and show a linear dependence with a slope of 0.37 and 0.39 respectively for the two paths involving the species Tl~ and TIOI-I~. The values of the slopes are in satisfactory agreement with the prediction of eqn (2).
2007
TI(III) interactions with quinols Table 3. Specific rate constants and standard redox potentials for the reaction of quinols with I"1(II1) E aa
k~ b
k'9~ b
benzene-1,4-diol 2,3,5-trimethyl2-methyl-
0.699 0.516 0.644
4.1xi02 5.2xi04 8.0xI02
1.4xi03 1.35xi06 2.1xi04
2-chloro2-bromo-
0.712 0.712
50 1.2xI02
Io3xi03 2.4xi03
2-sulphonic-
0.787
18
1.2xi02
2,3-dicyano-
0.910
0.58
12.5
Ref. 2a this work 2a 2b 2b 2b this work
a standard redox potentials (volt) of Q/H2Q couples at [HCIO4~ = 1.0 M, /u = 2.0 ~ (NaCl04) , 25.0 °. b 1 mol-ls -I , 18.0 °.
mechanism were operating, the rate should be electrontransfer controlled rather than substitution determined. It can be noted also that in the present case, the -OH ligand leads to a marked catalysis in the redox reaction. Similar behaviour has recently been discussed by Davies['/]. Moreover it is possible that the intrinsic 2+.is more favourable than parameter for the species T lOH~q that for TEa 3+[6], lowering AG" for a given AG °.
~20 •
S@@1 4
% <~ 15
10 15
I 20
I 25 -Z~G °
I 30 kca[
I 35
mot .1
Fig. 1. Plot of AG ~ vs AG°, at 18.0° and 2.0 M ionic strength, for the reaction of TI~ and TIOH~: (open and closed circles respectively) with quinols. The reductants are shown as follows: (1) Benzene-l,4-diol; (2) 2,3,5-tfimethyl-; (3) 2-methyl-; (4) 2bromo-; (5) 2-chloro-; (6) 2-sulphonic-; (7) 2,3-dicyano-. Since ligand substitution for TI~; has a rate > 1081 mol -~ s-' [1], the presently observed rates are much smaller and the observation of the dependence of Fig. 1 does not by itself allow a distinction between an inner- or outer-sphere mechanism. Thus, if an inner-sphere
REFERENCES 1. B. Falcinella, P. D. Felgate and G. S. Laurence, J. Chem. Soc. (Dalton Trans.) 1367 (1974); 1 (1975). 2. (a) E. Pelizzetti, E. Mentasti, M. E. Carlotti and G. Giraudi, J. Chem. Soc. (Dalton Trans.) 794 (1975); (b) E. Pelizzetti, E. Mentasti and E. Pramauro, Gazzetta 105, 403 (1975). 3. (a) G. Biedermann, Rec. Tray. Chim. 75, 716 (1956); (b) T. E. Rogers and G. M. Waind, Trans. Faraday Soc. 57, 1360(1961). 4. R. A. Marcus, Ann. Rev. Phys. Chem. 15, 155 (1964). 5. R. A. Marcus, J. Phys. Chem.72, 891 (1968); J. E. Earley, Progr. Inorg. Chem. 13, 243 (1970). 6. P. J. Proll, Comprehensive Chemical Kinetics (Edited by C. H. Bamford and C. R. H. Tipper), Vol. 7, Chap. 2, Elsevier, Amsterdam (1972); A. M. Armstrong, J. Halpern and W. C. E. Higginson, J. Phys. Chem. 60, 1661 0956); A. M. Armstrong and J. Halpern, Can. J. Chem. 35, 1022 (1957); A. C. Harkness and J. Halpern, Z Am. Chem. Soc. 81, 3526 (1959); J. D. Wear, J. Chem. Soc. 5596 (1965); F. B. Baker, W. D. Brewer and T. W. Newton, Inorg. Chem. 5, 1294 (1966). 7. G. Davies, Coord. Chem. Rev. 14, 287 (1974).
MECHANISMS OF REDOX REACTIONS OF T13q AND T1OH~aq WITH QUINOLS IN ACIDIC PERCHLORATE MEDIA EZIO PELIZZETTI and EDOARDO MENTASTI Istituto di ChimicaAnaliticadell'UniversitL Via P. Giuria, 5. 10125Torino, Italy (Received 18 November 1975)
Abstract--Further evidence of an electron-transfer controlled mechanism in the reactions of TI~ and TIOH~ with quinols has been achieved by investigating2,3,5-trimethyl-and 2,3-dicyano-benzene-l,4-diol,in order to have a larger range of AG°. For both ionic species a linear relationshipbetween AG" and AG° has been observed with slope 0.37 and 0.39 respectively. INTRODUCTION THERE have been recent developments in the chemistry of aquo-thallium(III). Studies of the reactions of TI(II), generated by flash-photolysis, by Laurence et al., also shed light on the mechanisms of reaction TI(III) [1]. In fact this ion can react by one- or two-electron transfer mechanisms. For the latter type of reaction it has been observed that there is a slight dependence of the free energy of activation on the free energy change and, for the same AG °, AG e is larger than for the one-electron path. In previous papers we have reported the kinetics and mechanism of oxidation of a series of quinols with TI(III) [2]. In the present note we investigate the kinetics of the oxidation of two quinol derivatives, 2,3,5-trimethylbenzene-l,4-diol (TMHzQ) and 2,3-dicyano-benzene-l,4diol (DCH2Q), with TI(III). The comparison of the specific rate constants for the different reaction paths allows us to formulate some interesting hypotheses about the reaction mechanism of Tl(III) when behaving as a two-electron acceptor. EXPERIMENTAL Reagents and procedure
TMH2Q and DCH2Q (K&K) solutions were freshly prepared. The other reagents and general experimentalprocedure have been described[2]. The ionic strength was brought to 2.0 M (NaCIO4). The kinetic parameters were computed with a weighted leastsquare method. The standard deviation computed for the rate parameters ranged from 2 to 5%. RESULTS The second order rate constants, ko, at different temperatures and acidities are collected in Table 1. The kinetic data can be accounted for by the following scheme, as previously reported[2], where Kh is the first TI~
+S
' (TI. . . . S)a]+
K2
hydrolysis constant of TI~g[3], K1 and K2 the formation equilibrium constants for the precursor complexes (if an outer-sphere mechanism is operating they represent JINC Vol. 38 No. II--F
outer-sphere association constants) and Q is the pquinone corresponding to the quinol S. It follows that ko(l+Kh[H+]-')=k'K~+k"KzKh[H+] -'.
(1)
By plotting the left-hand side of eqn (1) as a function of [H+]-~, straight lines were obtained and from the intercept and slope, the terms k'K1 and k"K2 were evaluated; the specific rate constants are listed in Table 2 together with the activation parameters. DISCUSSION For electron transfer reactions of the outer-sphere kind, the Marcus theory[4] predicts the relationship:
+ (AG°r)Z/8(AGII + AG$:)
(2)
where AG* = AG" - w., AG °`= AG°+ w. - w.; Wr and w. are the work terms required to bring the reactants or products together in the transition state and AGT, and AG 7: are the free activation energies for electron transfer for the self-exchange reactions. The expression can be rearranged to give the gradient of a plot AG e vs AG Oin o the form 12(1 + AGr/4AGo), where AG~ is the value of AG e for AG ° equal to zero. It must be taken into account that it is difficult in many cases to know or to evaluate the exchange rates and the work terms for the reactants; thus by comparing the reaction rates of different substrates with the same ion the dependence of AG ~ vs AG °, can be different from the prediction of eqn (2). Moreover, eqn (2) was derived for reactions of the outer-sphere class, but similar relationships might exist in other cases[5]. The reaction mechanism and the relationship between AG e and AG O for the reactions involving both TI(III) and Tiff) have been recently discussed by Laurence et al.[1]. By means of flash-photolysis, the kinetic behavior of TI(II) and the evaluation of the standard redox potentials of the couples TI(III)/TI(II) and Tl(II)frl(I) were investigated; thus it was possible to assign a single two electron-transfer or two successive one-electron steps to a series of reactions involving TI(III). For the reactions involving a single two electron transfer (exchange TI(III)--TI(I), TI(III)-Hg(0), TI(III)-U(IV), TI(III)-V(II))[6], only a slight variation in AG" was observed for large AG ° changes; thus the
2O05
2006
EZIO PELIZZETTIand EDOARDO MENTASTI Table 1 Second-order rate constants (lmol-' s-')" at different temperatures§ TI~2Q H0104] :
b
0.30
0.50
1.OO
2.00
6.O
1.36x105
1.11x105
7.5x104
4.7x104
18.0
2.60x105
2.21x105
1.33x105
8.4x104
t
(*o)
DOH2 Q c
18.0
0.33
0.50
1.00
2.00
2.63
2.07
1.28
0.86
8.0
6.2
40.0
11.5
a averaged on at least five separate kinetic runs
b [~}I2Q]= 5.0X10-4 ; ~1(11T~ = 6 " 20xlO-4M. c [Doa2Q3. 5 ox~o 4 ~ ~ l C ~ I ~ .
50
• 20 0 , 1 0 - 3 =
tho vaZ=e, of K h have bee~ oo=p=tcd, f r o =
ref.
[3b], t o be
6 . 5 X 1 0 -2 at 6.00 , 7.3x 10-2 at 18.O 0 and 8 . 8 • 1 0 - 2 =
at 40.000.
Table 2 Specific rate constants (lmol-I s ' ) and activation parametersfor the reactionsof TMH2Qand DCH2QwithTI(III) t(oO)
k'K 1
k"K 2
TMH2Q
6.0
2.42x104
8.1x105
18.0
5.2 xlO 4
1.35x106
AH* a
9.9 + 1.7
6.5 _+ 2.1
-3.2 + 5.5
-8.8 _+ 6.7
b AS*
DOH2Q
18.0
0.58
12.5
40.0
3.4
62
~H# a
14.0 ~ 1.3
12.6 Z 1.5
&S$ b
-11.5 ± 4.3
-10.2 ~ 5.0
a kcal mol -I
b cal mol-ldeg-I
intrinsic effect predominates and no definite assignment could be made as to the role of thermodynamic factors. We have previously investigated the reaction of a series of quinols with TI(III) and a two electron-transfer mechanism has been suggested[2]. It has already been pointed out that the reaction rate is increased by electron releasing groups and reduced by electron withdrawing ones. The present data concerning TMH2Q and DCH2Q, together with the evaluation of the standard redox potentials for these compounds (see Table 3) allow us to formulate some conclusions about the role of the change
in the overall free energy of reaction on the reaction rate of TI(III) when a two electron-transfer mechanism is operating. In fact this series of reactions involves the same metal ion and a group of substrates differing in free energy due to different substituents lying far from the reaction site, so that the intrinsic factors (AG 7, and w) can be assumed constant. The data obtained are plotted in Fig. 1 and show a linear dependence with a slope of 0.37 and 0.39 respectively for the two paths involving the species Tl~ and TIOI-I~. The values of the slopes are in satisfactory agreement with the prediction of eqn (2).
2007
TI(III) interactions with quinols Table 3. Specific rate constants and standard redox potentials for the reaction of quinols with I"1(II1) E aa
k~ b
k'9~ b
benzene-1,4-diol 2,3,5-trimethyl2-methyl-
0.699 0.516 0.644
4.1xi02 5.2xi04 8.0xI02
1.4xi03 1.35xi06 2.1xi04
2-chloro2-bromo-
0.712 0.712
50 1.2xI02
Io3xi03 2.4xi03
2-sulphonic-
0.787
18
1.2xi02
2,3-dicyano-
0.910
0.58
12.5
Ref. 2a this work 2a 2b 2b 2b this work
a standard redox potentials (volt) of Q/H2Q couples at [HCIO4~ = 1.0 M, /u = 2.0 ~ (NaCl04) , 25.0 °. b 1 mol-ls -I , 18.0 °.
mechanism were operating, the rate should be electrontransfer controlled rather than substitution determined. It can be noted also that in the present case, the -OH ligand leads to a marked catalysis in the redox reaction. Similar behaviour has recently been discussed by Davies['/]. Moreover it is possible that the intrinsic 2+.is more favourable than parameter for the species T lOH~q that for TEa 3+[6], lowering AG" for a given AG °.
~20 •
S@@1 4
% <~ 15
10 15
I 20
I 25 -Z~G °
I 30 kca[
I 35
mot .1
Fig. 1. Plot of AG ~ vs AG°, at 18.0° and 2.0 M ionic strength, for the reaction of TI~ and TIOH~: (open and closed circles respectively) with quinols. The reductants are shown as follows: (1) Benzene-l,4-diol; (2) 2,3,5-tfimethyl-; (3) 2-methyl-; (4) 2bromo-; (5) 2-chloro-; (6) 2-sulphonic-; (7) 2,3-dicyano-. Since ligand substitution for TI~; has a rate > 1081 mol -~ s-' [1], the presently observed rates are much smaller and the observation of the dependence of Fig. 1 does not by itself allow a distinction between an inner- or outer-sphere mechanism. Thus, if an inner-sphere
REFERENCES 1. B. Falcinella, P. D. Felgate and G. S. Laurence, J. Chem. Soc. (Dalton Trans.) 1367 (1974); 1 (1975). 2. (a) E. Pelizzetti, E. Mentasti, M. E. Carlotti and G. Giraudi, J. Chem. Soc. (Dalton Trans.) 794 (1975); (b) E. Pelizzetti, E. Mentasti and E. Pramauro, Gazzetta 105, 403 (1975). 3. (a) G. Biedermann, Rec. Tray. Chim. 75, 716 (1956); (b) T. E. Rogers and G. M. Waind, Trans. Faraday Soc. 57, 1360(1961). 4. R. A. Marcus, Ann. Rev. Phys. Chem. 15, 155 (1964). 5. R. A. Marcus, J. Phys. Chem.72, 891 (1968); J. E. Earley, Progr. Inorg. Chem. 13, 243 (1970). 6. P. J. Proll, Comprehensive Chemical Kinetics (Edited by C. H. Bamford and C. R. H. Tipper), Vol. 7, Chap. 2, Elsevier, Amsterdam (1972); A. M. Armstrong, J. Halpern and W. C. E. Higginson, J. Phys. Chem. 60, 1661 0956); A. M. Armstrong and J. Halpern, Can. J. Chem. 35, 1022 (1957); A. C. Harkness and J. Halpern, Z Am. Chem. Soc. 81, 3526 (1959); J. D. Wear, J. Chem. Soc. 5596 (1965); F. B. Baker, W. D. Brewer and T. W. Newton, Inorg. Chem. 5, 1294 (1966). 7. G. Davies, Coord. Chem. Rev. 14, 287 (1974).