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Solvent effect on the ion pair formation between 2,3,5,6-tetrachloro-1,4-benzoquinone anion radical and Mg2+ measured using a pulse electrolysis stopped flow method
j o ~ J m ¢ ~ OF
.......
11 Journal of ElectroanalyticalChemistry401 (1996) 243-246
ELSEVIER
Preliminary note
Solvent effect on the ion pair formation between 2,3,5,6-tetrachloro- 1,4-benzoquinone anion radical and Mg 2÷ measured using a pulse electrolysis stopped flow method Munetaka Oyama, Toyomasa Hoshino, Satoshi Okazaki * Division of Material Chemistry, Faculty of Engineering, Kyoto University, Sakyo-yu, Kyoto 606-01, Japan
Received 15 August 1995;in revisedform 11 September1995
Keywords: Ion pair formation;Quinone;Pulse electrolysisstopped flow
1. Introduction
2. Experimental
Many studies have been devoted to the ion pair formation processes between electrochemically reduced quinones and cationic species [1-5]. The shift of polarographic half-wave potentials provides useful information to determine the association constant and stoichiometry of the ion pair formed in the vicinity of the electrode. However, quite irreversible responses occasionally make it difficult to study the ion pair interactions by electrochemical methods. For example, the reduction of various quinones in the presence of Mg 2÷, which forms strong ion pairs, had been found difficult to analyze [3]. In the previous paper, we proposed a new analytical approach to reveal the dynamic interactions at the ion pair formation processes in homogeneous solution [6]. In spite of the complex electrochemical responses, the kinetic processes could be observed directly for the ion pair formation between 2,3-dichloro-5,6-dicyano-l,4-benzoquinone dianion and Na ÷ using a pulse electrolysis stopped flow method [7]. In the present paper, we applied this method to observe the solvent effect on the ion pair formation between 2,3,5,6-tetrachloro-l,4-benzoquinone (TCQ; chloranil) anion radical (TCQ-") and Mg 2+. While the electrochemical measurement indicated only that the interactions were very sensitive to the basicity of the solvents, the pulse electrolysis stopped flow measurement has permitted the various aspects of the ion pair formation processes to be observed in homogeneous solutions of acetonitrile (AN), N,N-dimethylformamide (DMF) and dimethylsulfoxide (DMSO).
Details of the pulse electrolysis stopped flow method have been described previously [6,7]. In the present work, the solution of T C Q - ' , which was quantitatively prepared with controlled potential electrolysis for 10 s, was mixed with the solution containing Mg 2÷, and then the changes in absorption spectra with time or the time decay curves of absorbance at fixed wavelength were observed in an optical flow cell. To deaerate the solutions, N 2 gas was bubbled into the reservoirs before measurement. The cyclic voltammograms (CVs) were measured using a PAR 174 analyzer. 2,3,5,6-Tetrachloro-l,4-benzoquinone (Nacarai tesque., GR grade) and Mg(CIO4) 2 (Aldrich), as a source of Mg 2÷, were used as received. DMF and DMSO were distilled under reduced pressure over Call 2, after being dried with 4A molecular sieves. The purification methods of AN and tetrabutylammonium perchlorate (TBAP) were described previously [8].
* Correspondingauthor. 0022-0728/96/$15.00 © 1996ElsevierScienceS.A. All rights reserved SSDI 0022-0728(95)04349-7
3. Results and discussion It has been reported that TCQ-" does not associate with Li ÷ in DMF because of the low charge density on the oxygen atoms caused by the chloro group [1]. In the present work, at first we observed the electrochemical response of TCQ in the presence of Mg 2+ in AN, DMF and DMSO. Fig. 1 shows the changes in the CVs of TCQ with the addition of Mg 2÷. Two reversible one-electron redox peaks (PJ, P2), which are attributed to the formation of TCQ-"
244
M. Oyama et al./ Journal of Electroanalytical Chemistry 40l (1996) 243-246
and TCQ dianion respectively, were observed in all the solvents examined when only tetrabutylammonium ion of a supporting electrolyte existed as a cationic species. In this measurement, a Pt 1(I-, 13) reference electrode in the same solvent [9] was used. While the solvent effect on the first and second reduction potentials of this sort of quinone has been reported in detail previously [10], in the present work a remarkable difference appeared as changes in the redox responses with the addition of Mg 2+. In DMF and DMSO the first redox peak (Pj) remained unchanged even in the presence of 100 mM Mg 2+, while positive shifts and changes in the shape for the second peak (P2) were observed. By contrast, the responses in AN were quite different from those in DMSO and DMF. With increasing Mg 2+ concentration, the reversible response of PI was gradually diminished and a new irreversible cathodic peak (P3) appeared at the potential region slightly positive to PI. The irreversible P3 became the only observable peak at higher Mg 2+ concentration (over 1.0 mM). It is, therefore, clear that the interaction between TCQ-" and Mg 2÷ in AN is much stronger than those in DMF and DMSO. This can be explained by the difference in solvating power for the cationic species [2,5]. Since the donor numbers of AN, DMF and DMSO are 14.1, 26.6 and 29.8 respectively [11], the basicity of DMSO and DMF is much higher than that of AN. Thus, the solvation of AN for Mg 2+ is assumed to be so weak as to let TCQ-" interact with Mg 2+. To investigate the details of the interactions in AN, the reaction TCQ-" and Mg 2+ was analyzed using the pulse electrolysis stopped flow method by mixing their AN solutions. Fig. 2 shows the changes in absorption spectra depending on the concentration of Mg 2+ after the reactions were completed. Although the kinetic processes were
AN P1 P2
[Mga']
DMF P1 P2
DMSO Pl P2
1.0m=
0.50
1/~--"" i
<
NW/
t.:
° 'J, 300
400
500
600
Wavelength / nm Fig. 2. Changes in absorption spectra when the solution of TCQ-" was mixed with the solution of Mg 2+. The solution of TCQ-" was produced by electrolyzing the AN solution of 0.50 mM TCQ containing 0.1 M TBAP with applied potential - 0 . 5 V for 10 s, then mixing it with the AN solution of Mg 2+. Each spectrum was measured 10 s from the mixing (see text). Concentration of Mg2+; (1) 0, (2) 0.05, (3) 0.10, (4) 0.15, (5) 0.20, (6) 0.25, (7) 0.30 mM.
observed immediately after mixing, a certain spectrum was observed in 10 s at any concentration of Mg 2÷. The spectrum measured without the addition of Mg 2÷ was similar to the reported one for free TCQ-', measured by dissolving the TCQ-" salts in acetone [12]. With increasing Mg 2+ concentration, the absorption spectra were gradually transformed with isosbestic points from that of free TCQ-" When the concentration of Mg 2÷ was increased over 0.30 mM, the observed absorption spectra were almost similar to that at 0.30 mM. Hence, the absorption spectrum having an absorption maximum at 290 nm, which is quite different from that of free T C Q - ' , was attributed to that of the ion pair formed in AN. In addition, from the plots of the absorbance at 290 and 449 nm versus the concentration of Mg 2+, the stoichiometry of the ion pair produced was clearly determined as Mg 2+ :TCQ-" equal to 1" 2. Thus it was confirmed that the contact ion pair, Mg "TCQ 2, was formed in AN. The kinetics of the contact ion pair formation in AN were also successfully analyzed with the present method. Fig. 3 shows the time decay curve of free TCQ-" measured after mixing the solutions of 0.50 mM TCQ-" and 0.50 mM Mg 2+. Although the low concentrations of both reactants made it difficult to analyze by the pseudo-reaction-order plot, the rate law was found to be -- d [ T C Q - ] / d t = k [ T C Q - ] 2 [Mg2+ ]
0
-1.0
0 -1.0 E / V ( 1 3 - , 1" )
0
-1.0
Fig. 1. Changes in cyclic voltammograms of TCQ in AN, DMF and DMSO with the addition of Mg 2÷. Working electrode, Pt disk electrode (diameter 1.0 ram). Reference electrode Pt I(I-, I~ ) electrode. Scan rate, 100 mV s - i. TCQ, 1.0 mM. Mg 2÷ was added by dissolving Mg(CIO4) 2. The total concentration of the electrolyte was adjusted to be 100 mM by adding TBAP.
from the simulation, as shown in Fig. 3. This means that the reaction proceeds via the following scheme: Mg 2+ + TCQ- • fast > Mg2+TCQ -" M g 2 + T C Q - + TCQ-"
rds > Mge+.TCQ2.
From the measurements in which the initial concentration of Mg 2+ was varied, the rate constant k was estimated to be 4.0 ( + 0 . 5 ) × 109 M - 2 s - j
M. Oyama et al. / Journal of Electroanalytical Chemistry 401 (1996) 243-246
However, the measurements of absorption spectra were also carried out in DMF and DMSO, The absorption spectra of T C Q - ", measured without the addition of Mg 2÷ in these solvents, were similar to that in AN without Mg 2÷. This result shows that T C Q - " exists in free state or as a solvent-separated ion pair in the three solvents in the absence of Mg 2÷. Furthermore, unlike the reaction of TCQ-" with Mg 2÷ in AN, the absorption spectra of free TCQ-" in DMF and DMSO did not change, even when it was mixed with 100 mM Mg 2÷, as implied from the CV results of Fig. 1. This result indicates that T C Q - " exists in free state in DMF and DMSO even when a large excess of Mg 2÷ is present. Therefore, the interactions of DMF and DMSO with Mg 2÷ are found to be stronger than that between T C Q - ' and Mg 2÷, though the latter is a coulombic interaction. Also, the competing interactions of DMF (or DMSO) and TCQ-" for Mg 2÷ could be observed in the kinetic measurements. Fig. 4 shows the time decay curves in the reaction of TCQ-" and Mg 2÷, measured with the addition of a small amount of DMF or DMSO in the AN solution of Mg z÷. As shown in this figure, the reaction of TCQ-" and Mg 2÷ in AN was decelerated by the strong-solvating species, DMF or DMSO. With increasing the added concentration, the deceleration was intensified. In this manner, the degree of the interactions was observed as the difference in the reaction rate forming the contact ion pair. The greater decelerating effect of DMSO over DMF was ascribed to the difference in their basicity, indicated from their donor numbers. In conclusion, the difference in ion pair formation due to the solvent used could be observed by using the pulse electrolysis stopped flow method. Although the electrochemical response in AN was difficult to analyze, the measurement in homogeneous solution clearly showed that the contact ion pair, Mg "TCQ 2, was formed in AN, and
0.1
'\ i'\\ '--'
0.05
0
\"
0.2
0"1 ~ x
0 0
A B C
' 100
200
Time / ms Fig. 4. Time decay curves of T C Q - ' measured after mixing the solutions of 0.50 mM TCQ-" and 0.50 mM Mg 2+ with the addition of a small amount of DMF or DMSO. The additive to the AN solution of Mg2+; (A) none, (B) 10 mM DMF, (C) 10 mM DMSO. Wavelength, 449 nm.
that the kinetics and mechanisms were successfully analyzed with this method. This approach would be useful in analyzing the kinetics of the complex electrochemical reactions, and conversely, information on the reactions in the homogeneous solution should be useful to interpret the complex heterogeneous electrochemical responses. By contrast, the solvent-separated ion pair was found to be formed in DMF and DMSO. In the present case, the changes in the basicity of the solvents caused the drastic change in the ion pair formation processes. The effect of DMF and DMSO was also observed through the kinetic measurement of the ion pair formation reaction. The competition of DMF (or DMSO) with TCQ-" for interacting toward Mg 2+ was observed with deceleration of the reaction rate forming Mg "TCQ2. With this approach the degree of the interactions would be comprehended as the difference of the reaction rates. The present method would also be useful in analyzing the specific solute-solvent or solute-solute interactions in homogeneous solution concerning the electrogenerated species.
Acknowledgement This work was supported in part by a Grant-in-Aid for Scientific Research from the Ministry of Education, Science and Culture, Japan.
'~
............. 2217 0
245
40
References
80
Time / ms Fig. 3. Time decay curves of TCQ-" measured after mixing the solutions of 0.50 mM TCQ-" and 0.50 mM Mg 2+. Wavelength, 449 nm. experimental result; . . . . . . , simulated from the rate law - d[TCQ- "]/dt = k[TCQ- "]2[Mg2+ ]; ( . . . . . ), the result from another possible rate law, - d [ T C Q - ' ] / d t = k[TCQ-'][Mg 2+ ], which would not fit the experimental result.
[1] M.E. Peover, in A.J. Bard (Ed.), Electrochemical Chemistry, Vol. 2, Marcel Dekker, New York, 1967. [2] T. Fujinaga, K. Izutsu and T. Nomura, J. Electroanal. Chem., 29 (1971) 203. [3] T. Nagaoka, S. Okazaki and T. Fujinaga, J. Electroanal. Chem., 133 (1982) 89. [4] T. Nagaoka and S. Okazaki, J. Electroanal. Chem., 158 (1983) 139. [5] C. Russel and W. Jaenicke, J. Electroanal. Chem., 199 (1986) 139.
246
M. Oyama et al./Journal of Electroanalytical Chemistry 401 (1996) 243-246
[6] M. Oyama, A. Takei and S. Okazaki, J. Chem. Soc., Chem. Commun., in press. [7] M. Oayama, K. Nozaki and S. Okazaki, Anal. Chem., 63 (1991) 1387. [8] M. Oyama, K. Nozaki and S. Okazaki, J. Electroanal. Chem., 304 (1991) 61.
[9] J.F. Coetzee and C.W. Gardner, Jr., Anal. Chem., 54 (1982) 2530. [10] K. Sasaki, T. Kashimura, M. Ohura, Y. Ohsaki and N. Ohta, J. Electrochem. Soc., 137 (1990) 2437. [11] V. Gutmann, Coordination Chemistry in Non Aqueous Solution, Springer, New York, 1969. [12] Y. Iida, Bull. Chem. Soc. Jpn., 43 (1970) 2772.
.......
11 Journal of ElectroanalyticalChemistry401 (1996) 243-246
ELSEVIER
Preliminary note
Solvent effect on the ion pair formation between 2,3,5,6-tetrachloro- 1,4-benzoquinone anion radical and Mg 2÷ measured using a pulse electrolysis stopped flow method Munetaka Oyama, Toyomasa Hoshino, Satoshi Okazaki * Division of Material Chemistry, Faculty of Engineering, Kyoto University, Sakyo-yu, Kyoto 606-01, Japan
Received 15 August 1995;in revisedform 11 September1995
Keywords: Ion pair formation;Quinone;Pulse electrolysisstopped flow
1. Introduction
2. Experimental
Many studies have been devoted to the ion pair formation processes between electrochemically reduced quinones and cationic species [1-5]. The shift of polarographic half-wave potentials provides useful information to determine the association constant and stoichiometry of the ion pair formed in the vicinity of the electrode. However, quite irreversible responses occasionally make it difficult to study the ion pair interactions by electrochemical methods. For example, the reduction of various quinones in the presence of Mg 2÷, which forms strong ion pairs, had been found difficult to analyze [3]. In the previous paper, we proposed a new analytical approach to reveal the dynamic interactions at the ion pair formation processes in homogeneous solution [6]. In spite of the complex electrochemical responses, the kinetic processes could be observed directly for the ion pair formation between 2,3-dichloro-5,6-dicyano-l,4-benzoquinone dianion and Na ÷ using a pulse electrolysis stopped flow method [7]. In the present paper, we applied this method to observe the solvent effect on the ion pair formation between 2,3,5,6-tetrachloro-l,4-benzoquinone (TCQ; chloranil) anion radical (TCQ-") and Mg 2+. While the electrochemical measurement indicated only that the interactions were very sensitive to the basicity of the solvents, the pulse electrolysis stopped flow measurement has permitted the various aspects of the ion pair formation processes to be observed in homogeneous solutions of acetonitrile (AN), N,N-dimethylformamide (DMF) and dimethylsulfoxide (DMSO).
Details of the pulse electrolysis stopped flow method have been described previously [6,7]. In the present work, the solution of T C Q - ' , which was quantitatively prepared with controlled potential electrolysis for 10 s, was mixed with the solution containing Mg 2÷, and then the changes in absorption spectra with time or the time decay curves of absorbance at fixed wavelength were observed in an optical flow cell. To deaerate the solutions, N 2 gas was bubbled into the reservoirs before measurement. The cyclic voltammograms (CVs) were measured using a PAR 174 analyzer. 2,3,5,6-Tetrachloro-l,4-benzoquinone (Nacarai tesque., GR grade) and Mg(CIO4) 2 (Aldrich), as a source of Mg 2÷, were used as received. DMF and DMSO were distilled under reduced pressure over Call 2, after being dried with 4A molecular sieves. The purification methods of AN and tetrabutylammonium perchlorate (TBAP) were described previously [8].
* Correspondingauthor. 0022-0728/96/$15.00 © 1996ElsevierScienceS.A. All rights reserved SSDI 0022-0728(95)04349-7
3. Results and discussion It has been reported that TCQ-" does not associate with Li ÷ in DMF because of the low charge density on the oxygen atoms caused by the chloro group [1]. In the present work, at first we observed the electrochemical response of TCQ in the presence of Mg 2+ in AN, DMF and DMSO. Fig. 1 shows the changes in the CVs of TCQ with the addition of Mg 2÷. Two reversible one-electron redox peaks (PJ, P2), which are attributed to the formation of TCQ-"
244
M. Oyama et al./ Journal of Electroanalytical Chemistry 40l (1996) 243-246
and TCQ dianion respectively, were observed in all the solvents examined when only tetrabutylammonium ion of a supporting electrolyte existed as a cationic species. In this measurement, a Pt 1(I-, 13) reference electrode in the same solvent [9] was used. While the solvent effect on the first and second reduction potentials of this sort of quinone has been reported in detail previously [10], in the present work a remarkable difference appeared as changes in the redox responses with the addition of Mg 2+. In DMF and DMSO the first redox peak (Pj) remained unchanged even in the presence of 100 mM Mg 2+, while positive shifts and changes in the shape for the second peak (P2) were observed. By contrast, the responses in AN were quite different from those in DMSO and DMF. With increasing Mg 2+ concentration, the reversible response of PI was gradually diminished and a new irreversible cathodic peak (P3) appeared at the potential region slightly positive to PI. The irreversible P3 became the only observable peak at higher Mg 2+ concentration (over 1.0 mM). It is, therefore, clear that the interaction between TCQ-" and Mg 2÷ in AN is much stronger than those in DMF and DMSO. This can be explained by the difference in solvating power for the cationic species [2,5]. Since the donor numbers of AN, DMF and DMSO are 14.1, 26.6 and 29.8 respectively [11], the basicity of DMSO and DMF is much higher than that of AN. Thus, the solvation of AN for Mg 2+ is assumed to be so weak as to let TCQ-" interact with Mg 2+. To investigate the details of the interactions in AN, the reaction TCQ-" and Mg 2+ was analyzed using the pulse electrolysis stopped flow method by mixing their AN solutions. Fig. 2 shows the changes in absorption spectra depending on the concentration of Mg 2+ after the reactions were completed. Although the kinetic processes were
AN P1 P2
[Mga']
DMF P1 P2
DMSO Pl P2
1.0m=
0.50
1/~--"" i
<
NW/
t.:
° 'J, 300
400
500
600
Wavelength / nm Fig. 2. Changes in absorption spectra when the solution of TCQ-" was mixed with the solution of Mg 2+. The solution of TCQ-" was produced by electrolyzing the AN solution of 0.50 mM TCQ containing 0.1 M TBAP with applied potential - 0 . 5 V for 10 s, then mixing it with the AN solution of Mg 2+. Each spectrum was measured 10 s from the mixing (see text). Concentration of Mg2+; (1) 0, (2) 0.05, (3) 0.10, (4) 0.15, (5) 0.20, (6) 0.25, (7) 0.30 mM.
observed immediately after mixing, a certain spectrum was observed in 10 s at any concentration of Mg 2÷. The spectrum measured without the addition of Mg 2÷ was similar to the reported one for free TCQ-', measured by dissolving the TCQ-" salts in acetone [12]. With increasing Mg 2+ concentration, the absorption spectra were gradually transformed with isosbestic points from that of free TCQ-" When the concentration of Mg 2÷ was increased over 0.30 mM, the observed absorption spectra were almost similar to that at 0.30 mM. Hence, the absorption spectrum having an absorption maximum at 290 nm, which is quite different from that of free T C Q - ' , was attributed to that of the ion pair formed in AN. In addition, from the plots of the absorbance at 290 and 449 nm versus the concentration of Mg 2+, the stoichiometry of the ion pair produced was clearly determined as Mg 2+ :TCQ-" equal to 1" 2. Thus it was confirmed that the contact ion pair, Mg "TCQ 2, was formed in AN. The kinetics of the contact ion pair formation in AN were also successfully analyzed with the present method. Fig. 3 shows the time decay curve of free TCQ-" measured after mixing the solutions of 0.50 mM TCQ-" and 0.50 mM Mg 2+. Although the low concentrations of both reactants made it difficult to analyze by the pseudo-reaction-order plot, the rate law was found to be -- d [ T C Q - ] / d t = k [ T C Q - ] 2 [Mg2+ ]
0
-1.0
0 -1.0 E / V ( 1 3 - , 1" )
0
-1.0
Fig. 1. Changes in cyclic voltammograms of TCQ in AN, DMF and DMSO with the addition of Mg 2÷. Working electrode, Pt disk electrode (diameter 1.0 ram). Reference electrode Pt I(I-, I~ ) electrode. Scan rate, 100 mV s - i. TCQ, 1.0 mM. Mg 2÷ was added by dissolving Mg(CIO4) 2. The total concentration of the electrolyte was adjusted to be 100 mM by adding TBAP.
from the simulation, as shown in Fig. 3. This means that the reaction proceeds via the following scheme: Mg 2+ + TCQ- • fast > Mg2+TCQ -" M g 2 + T C Q - + TCQ-"
rds > Mge+.TCQ2.
From the measurements in which the initial concentration of Mg 2+ was varied, the rate constant k was estimated to be 4.0 ( + 0 . 5 ) × 109 M - 2 s - j
M. Oyama et al. / Journal of Electroanalytical Chemistry 401 (1996) 243-246
However, the measurements of absorption spectra were also carried out in DMF and DMSO, The absorption spectra of T C Q - ", measured without the addition of Mg 2÷ in these solvents, were similar to that in AN without Mg 2÷. This result shows that T C Q - " exists in free state or as a solvent-separated ion pair in the three solvents in the absence of Mg 2÷. Furthermore, unlike the reaction of TCQ-" with Mg 2÷ in AN, the absorption spectra of free TCQ-" in DMF and DMSO did not change, even when it was mixed with 100 mM Mg 2÷, as implied from the CV results of Fig. 1. This result indicates that T C Q - " exists in free state in DMF and DMSO even when a large excess of Mg 2÷ is present. Therefore, the interactions of DMF and DMSO with Mg 2÷ are found to be stronger than that between T C Q - ' and Mg 2÷, though the latter is a coulombic interaction. Also, the competing interactions of DMF (or DMSO) and TCQ-" for Mg 2÷ could be observed in the kinetic measurements. Fig. 4 shows the time decay curves in the reaction of TCQ-" and Mg 2÷, measured with the addition of a small amount of DMF or DMSO in the AN solution of Mg z÷. As shown in this figure, the reaction of TCQ-" and Mg 2÷ in AN was decelerated by the strong-solvating species, DMF or DMSO. With increasing the added concentration, the deceleration was intensified. In this manner, the degree of the interactions was observed as the difference in the reaction rate forming the contact ion pair. The greater decelerating effect of DMSO over DMF was ascribed to the difference in their basicity, indicated from their donor numbers. In conclusion, the difference in ion pair formation due to the solvent used could be observed by using the pulse electrolysis stopped flow method. Although the electrochemical response in AN was difficult to analyze, the measurement in homogeneous solution clearly showed that the contact ion pair, Mg "TCQ 2, was formed in AN, and
0.1
'\ i'\\ '--'
0.05
0
\"
0.2
0"1 ~ x
0 0
A B C
' 100
200
Time / ms Fig. 4. Time decay curves of T C Q - ' measured after mixing the solutions of 0.50 mM TCQ-" and 0.50 mM Mg 2+ with the addition of a small amount of DMF or DMSO. The additive to the AN solution of Mg2+; (A) none, (B) 10 mM DMF, (C) 10 mM DMSO. Wavelength, 449 nm.
that the kinetics and mechanisms were successfully analyzed with this method. This approach would be useful in analyzing the kinetics of the complex electrochemical reactions, and conversely, information on the reactions in the homogeneous solution should be useful to interpret the complex heterogeneous electrochemical responses. By contrast, the solvent-separated ion pair was found to be formed in DMF and DMSO. In the present case, the changes in the basicity of the solvents caused the drastic change in the ion pair formation processes. The effect of DMF and DMSO was also observed through the kinetic measurement of the ion pair formation reaction. The competition of DMF (or DMSO) with TCQ-" for interacting toward Mg 2+ was observed with deceleration of the reaction rate forming Mg "TCQ2. With this approach the degree of the interactions would be comprehended as the difference of the reaction rates. The present method would also be useful in analyzing the specific solute-solvent or solute-solute interactions in homogeneous solution concerning the electrogenerated species.
Acknowledgement This work was supported in part by a Grant-in-Aid for Scientific Research from the Ministry of Education, Science and Culture, Japan.
'~
............. 2217 0
245
40
References
80
Time / ms Fig. 3. Time decay curves of TCQ-" measured after mixing the solutions of 0.50 mM TCQ-" and 0.50 mM Mg 2+. Wavelength, 449 nm. experimental result; . . . . . . , simulated from the rate law - d[TCQ- "]/dt = k[TCQ- "]2[Mg2+ ]; ( . . . . . ), the result from another possible rate law, - d [ T C Q - ' ] / d t = k[TCQ-'][Mg 2+ ], which would not fit the experimental result.
[1] M.E. Peover, in A.J. Bard (Ed.), Electrochemical Chemistry, Vol. 2, Marcel Dekker, New York, 1967. [2] T. Fujinaga, K. Izutsu and T. Nomura, J. Electroanal. Chem., 29 (1971) 203. [3] T. Nagaoka, S. Okazaki and T. Fujinaga, J. Electroanal. Chem., 133 (1982) 89. [4] T. Nagaoka and S. Okazaki, J. Electroanal. Chem., 158 (1983) 139. [5] C. Russel and W. Jaenicke, J. Electroanal. Chem., 199 (1986) 139.
246
M. Oyama et al./Journal of Electroanalytical Chemistry 401 (1996) 243-246
[6] M. Oyama, A. Takei and S. Okazaki, J. Chem. Soc., Chem. Commun., in press. [7] M. Oayama, K. Nozaki and S. Okazaki, Anal. Chem., 63 (1991) 1387. [8] M. Oyama, K. Nozaki and S. Okazaki, J. Electroanal. Chem., 304 (1991) 61.
[9] J.F. Coetzee and C.W. Gardner, Jr., Anal. Chem., 54 (1982) 2530. [10] K. Sasaki, T. Kashimura, M. Ohura, Y. Ohsaki and N. Ohta, J. Electrochem. Soc., 137 (1990) 2437. [11] V. Gutmann, Coordination Chemistry in Non Aqueous Solution, Springer, New York, 1969. [12] Y. Iida, Bull. Chem. Soc. Jpn., 43 (1970) 2772.