Olmstead and Nicholson (12) have indicated the feasibility of evaluating k l by constructing working curves for the ratio of anodic to cathodic peak current U S . kf COT. Here C" is the bulk concentration of L2+and 7 is the time from E" to the switching potential, Ex. Since the appearance of the anodic wave occurred in a region extremely high capacitive current (Figure 4), it is not possible to utilize this diagnostic criterion to evaluate k Nevertheless, the appearance of the anodic peak at a value of log Ad = -1.58 substantiates the proposed mechanism and is consistent with the estimated rate constant. The scan rate data on Table I1 for DBA oxidation show no dependence of normalized peak current with scan rate. This is indicative of an electrochemical oxidation uncomplicated by chemical reactions. The experimental value of n, = 2 indicated a net two-electron oxidation. Figures 1 and 3 show no reversible cathodic wave in conjunction with the anodic DBA wave, indicating the two-electron oxidation of DBA to be a net irreversible process. Since true twoelectron oxidations in organic systems are unlikely, it is more probable that the oxidation passes through a one-electron oxidation intermediate. The one-electron oxidation product of DBA is L.'. By reference to Figure 4, it may be seen that L.+ itself will be oxidized to lucigenin at the anodic potential necessary to oxidize DBA to L . + . Thus, the oxidation of DBA appears to fit the scheme: (12) M. L. Olmstead and R. S . Nicholson, ANAL. CHEM.,41, 862 (1969).
DBA
- e-
--t
L e + - e- + L2+
or two, consecutive one-electron oxidations occurring at the same potential. In summary, the electrochemical behavior of lucigenin in DMSO appears to best fit the following scheme : L*+
2L.+ -% L2+ DBA
+ e- e L . +
+ DBA for the reduction of lucigenin
- e- e L . + -
e-
-+
L2+ for the oxidation of DBA
The value of k , has been estimated as 4.41 X 104M-l sec-l. Potential step methods by Booman (13) for following disproportionation reactions may yield a more refined estimate of k,. ACKNOWLEDGMENT
We thank David Hume and Leon Klatt for reading the manuscript and for their helpful suggestions.
RECEIVED for review February 1, 1972. Accepted March 30, 1972. This work was supported in part through funds provided by the U S . Atomic Energy Commission under Contract AT(30-1)905. K.D.L. was a Predoctoral Fellow, 1965-68. (13) G. L. Booman and D. T. Pence, ibid., 37,1366 (1965).
Ion-Selective Electrode Study of Copper(1) Cornplexes in Acetonitrile L. F. Heerman and G . A. Rechnitz Department of Chemistry, State University of New York, Buffalo, N . Y . 14214
A cuprous sulfide-membrane ion-selective electrode was used for the potentiometric measurement of copper(l) ion concentration in acetonitrile with tetraethylammonium perchlorate or sodium perchlorate as supporting electrolytes. The electrode showed an almost Nernstian behavior (slope 55-56 mV/decade) for copper(l) ion concentrations down to lo-5M in pure solutions and to at least 3 x l W 7 M in the presence of complexing ligands. The electrode was used for the measurement of the stepwise formation constants of copper(l) complexes with the halides in acetonitrile. Log p1 and log p2 values were as follows: CI-, 4.9 and 10.7; Br-, 3.8 and 7.8; and I-, 3.2 and 6.4. The electrode might be particularly useful for the study of interactions of copper(1) ions with organic ligands in acetonitrile. As an example, complex formation between copper(1) and thiourea has been investigated. The experimental data suggest Cu[S=C(NH2)&+ as the predominant complex species for which the formation constant has been evaluated as log p2 = 6.3. COPPER(I)IONS are stable in a number of organic solvents, e.g., acetonitrile ( I ) and nitromethane (2). The stability of copper(1) ions in such solvents has been explained by an (1) I. M. Kolthoff and J. F. Coetzee, J. Amer. Chem. SOC.,79, 1852 (1957). (2) I. V. Nelson, R . C. Larson, and R. T. Iwamoto, J . Inorg. Nucl. Chem.,22,279 (1961).
increase of the solvation energy of copper(1) ions and/or a decrease of the solvation energy of copper(I1) ions with respect to the solvation energies of these ions in water ( 2 ) . The formation constants of copper(1) complexes with chloride in acetonitrile have been measured by polarographic methods (3). The formation constants of copper(1) complexes with the halides and thiocyanate have been determined by potentiometric measurements using a copper amalgam indicator electrode (4). In this work, the formation constants of copper(1) complexes with the halides and with thiourea in acetonitrile have been measured using a novel cuprous sulfide-membrane ionselective electrode as the indicator electrode. Ion-selective electrodes have been used for a large number of complex formation studies in aqueous solutions (5, 6). Furthermore, ion-selective electrodes have been used for potentiometric titrations in a number of organic solvents, e.g., acetonitrile, and their application for the quantitative (3) S. E. Manahan and R. T. Iwamoto, Inorg. Chem., 4, 1409 (1965). (4) J. K. Senne and B. Kratochvil, ANAL.CHEM., 43,79(1971). ( 5 ) J. N. Butler, "Ion-Selective Electrodes," R. A. Durst, Ed., NBS Special Publication No. 314, U.S. Government Printing Office, Washington, D. C. 1969, p 143. (6) G. A. Rechnitz, Accounts Chem. Res., 3,69 (1970).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 9, AUGUST 1972
1655
LEADING WIRE
GLASS TUBE
ARALDITE CUPROUS SULFIDE MEMBRANE
Figure 1. Cross-section of the cuprous sulfidemembrane ion-selective electrode
-
!€I-5
CONCENTRATION,
M
Figure 2. Calibration curve for copper(1) ions in acetonitrile Total ionic strength = 0.1M Et,NCIO, (25 "C). Potentials are measured cs. SCE filled with saturated KC1-methanol electrolyte study of complex formation reactions in such nonaqueous media has been suggested (7). An ion-selective electrode method seems t o be particularly well suited for the measurement of copper(1) complex formation constants in acetonitrile since anhydrous solutions of copper(1) perchlorate in this solvent are easily obtained over a wide range of concentrations. EXPERIMENTAL Chemicals. Acetonitrile (Fisher Scientific Co., Certified Reagent) was used without further purification. Copper(1) perchlorate was prepared as the tetrakis (methylcyanide) complex by reaction of copper metal with copper(I1) perchlorate in acetonitrile (8). Analysis for copper gave 9 9 . 6 z C U ( C H ~ C N ) ~ C I OThe ~ . exact copper concentration of the solutions was determined by EDTA titration after evaporation of the acetonitrile and oxidation of the metal with dilute nitric acid. Tetraethylammonium perchlorate (Eastman Organic Chemicals) was recrystallized twice from deionized water and dried for 24 hr a t 110 "C. Silver nitrate tests for halides were negative. (7) G. A. Rechnitz and N. C. Kenny, Anal. Lett., 2,395 (1969). (8) B. J. Hathaway, D. G. Halah, and J. D. Postlethwaite, J . Chem. Soc., 1961,3215.
1656
0
Tetraethylammonium chloride, tetraethylammonium bromide, and tetraethylammonium iodide (Eastman Organic Chemicals) were purified by the methods recommended in the literature (9). Anhydrous sodium perchlorate (G. F. Smith Chemical Co.), sodium iodide (Fisher Scientific Co.) and thiourea (Baker Certified Reagent) were used as received. The Araldite No. 502 Resin was obtained from Chemical Coding and Engineering Co., Media, Pa. Preparation of the Electrode. A cross-section of the cuprous sulfide-membrane ion-selective electrode is shown in Figure 1. An electrode of almost identical design has been described by Hirata et al. (10) for the potentiometric measurement of copper(I1) ion concentration in aqueous solution, but was not used for copper(1) measurements or in nonaqueous solvents. The membrane (diameter: 13 mm) was prepared by heating finely powdered cuprous sulfide (Alfa Inorganics) (1.7-2.0 grams) a t 300 "C for 24 hr under a pressure of :ibotit 7 ton,' cm'. After fastening a copper leading wire directly I O th:. membrane, the pellet was imbedded in a glass tube with araldite (which proved t o be very resistant toward acetonitrile) and the electrode surface was polished with very fine abrasive cloth. The electrode was then conditioned overnight in a 10-'M copper(1) perchlorate solution in acetonitrile before the measurements were started. The replacement of the internal electrode and filling solution, normally used in glass electrodes and ion-selective electrodes, by a direct leading wire contact to the membrane proved to be useful for the preparation of several sulfide-based solid state ion-selective electrodes ( / U - / 4 ) . Leakage of the solution through the side-wall of the membrane and the electrode body, causing electrical contact between the solution and the leading wire and resulting in breakdohn of the electrode, is almost totally excluded and facilitates use of such electrodes in nonaqueous solvents. Procedure. All experiments were carried out in a donblewalled borosilicate glass vessel thermoitated at 25 'C. Electrodes and burets were inserted through a rubber stopper in the cell. No special precautions were taken to avoid the entry of oxygen in the cell since the solutions wrre stable toward air oxidation during the time needed for the measurements. Potentiometric measurements (+0.1 mV) were carried out with a Corning Model 12 research pH meter. In solutions containing tetraethylammonium perchlorate as the supporting electrolyte either a Beckman quartz fiber type SCE filled with saturated KCI-methanol electrolyte or a n Orion Model 90-02 double junction electrode with the supporting electrolyte solution in the outer chamber was used as a reference. In solutions containing sodium perchlorate as the supporting electrolyte, however, stable potentials could be obtained only with the double junction electrode as a reference. Calibration curves were recorded by adding increments of a 10-*M or a 2.5 x 10-2M copper(1) perchlorate solution to a known volume of solution containing the supporting electrolyte. Microburets used for calibration or titration were graduated to 0.02 ml (for recording calibration curves, a microsyringe was used for the lower concentrations). In ~-
(9) C. K. Maim "Electroanalytical Chemistrl," A. J. Bard. Ed., Vol. 3, Marcel Dekker, New York, N.Y. 1969, p 57, Appmdix 11. (10) H. Hirata, K . Higashiyama. and K . Date. Ami/. C//im.Acta.,
51,209 (1970). (11) H. Hirataand K. Higashiyama, ibid.,54,415(1971). (12) H. Hirata and K. Higashiyama, Birll. C h m . Sor. Jrrp.. 44, 2420 (1 971). (13) J. Czaban and G. A. Rechnitz, paper presented at the Third Northeast Regional ACS Meeting, Buffalo. N.Y., October 11-13, 1971. (14) G. A. Rechnitz, G. H. Fricke, and hl. S. Xlohan. AXAL. CHEM.,44,1098 (1972).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 9, AUGUST 1972
complex formation studies, the electrode was calibrated prior to each titration run. For complex formation studies with chloride, the total copper(1) concentration was held constant at -5 X 10-3M during the titration by adding the same amount of copper to the titrant. For complex formation studies with bromide or and iodide, the total copper(1) concentration was N ~ O - ~ M thiourea titrations were performed at different total copper(1) concentrations varying from 5 x 10-4M to 5 x 10-3M. The ligand concentration was varied over the range zero to 10+M for the halides (except chloride) and from zero to 2 X 10d2M for thiourea. All calculations were performed using a Hewlett-Packard 9100A calculator. A linear least square program was used for the determination of the slope and E" values.
\
RESULTS AND DISCUSSION A typical calibration curve is shown in Figure 2 . The electrode displayed an almost Nernstian response to copper(1) ion concentrations over the range 10-2M-10-6M with a slope of 55-56 mV/decade (theoretically 59.16 mV at 25 "C). According to Hirata et al. (10) who studied the behavior of cuprous sulfide-membrane ion-selective electrodes only in aqueous copper(I1) solutions, the slope of the calibration curve and the response time of the electrode depends on the stoichiometry of the sulfide phase. The slope decreased and the response time increased when x in Cu2-,S was varied from 0.21 to zero, the best results thus being obtained for C U ~ , ~ ~The S . cuprous sulfide used in this work had an almost perfect stoichiometric composition (analytical data supplied by the manufacturer) so that the small deviations of the slope from the theoretical value can be accounted for by the high metal to sulfur ratio of the membrane. No detailed study was made of the response time of the electrode. It was observed, however, that less than 60 seconds was required to obtain steady potentials for copper(1) concentrations >5 X lO-4M. Below this concentration level, the response became more sluggish and an interval of 5-10 minutes was allowed for equilibration. The E" values changed by several millivolts from one day to another but remained constant during the time needed for a titration run. The complex studies described below indicated that the electrode showed an almost Nernstian response, Le., slope of 55-56 mV, for copper(1) concentration down to at least 3 X lO-'M in the presence of complexing ligands. As an example, consider the titration of copper(1) ion with chloride shown in Figure 3. In this figure, the lower concentration level below which the electrode no longer showed a Nernstian response in pure solutions is also indicated. The linear part of the calibration curve was extrapolated to lower concentrations for the determination of the free copper(1) ion concentration used for the calculation of the complex formation constant. Only by this procedure was it possible to obtain a consistent set of log p2 values for the different points along the titration curve. Nernstian response down to very low concentrations in the presence of complexed metal ion has also been reported for several other sulfide based ion-selective electrodes (15, Zh). Overall complex formation constants, PI and p?, for the copper(1) complexes with the halides were calculated by combining the formation constant expression (15) Instruction Manual, Orion cupric ion electrode Model 94-29,
Orion Research, Inc., Cambridge, Mass., p 5. (16) R. A. Durst, "Ion-Selective Electrodes," R. A. Durst, Ed., NBS Special Publication No. 314, U.S. Government Printing Office, Washington, D.C., 1969, p 375.
c [ c ~ =]
IO-^ M
b
3 TOTAL LIGAND/TOTAL METAL MOLE R A T I O
Figure 3. Titration of copper(1) perchlorate with tetraethylammonium chloride in acetonitrile Total copper(1) ion concentration 5.25 X 10-3M held constant during titration. Ionic strength = 0.1M Et 4NClO4 (25 "C)
with the mass balance equation for total copper [Cut], and total ligand [L-I,, whereby the last term in Equation 4 takes into
+ [CUL] + [CuL*-I [L-] + [CUL] + ~[CULZ-] + K,[Cu+l[L-]
[Cu+l, [L-1,
=
=
[Cu+l
(3)
(4)
account the ion pairing between the cation of the supporting electrolyte and the ligand (K, = association constant). No corrections were made for teraethylammonium perchlorate which is completely dissociated in acetonitrile (17). For sodium perchlorate, a value of the association constant K, = 10 mol-' has been reported (18) but since sodium iodide is completely dissociated (19), no corrections must be made in this particular case. In the general case, however, correction may be necessary for both the titrant salt and the supporting electrolyte. No corrections were made for copper(1) perchlorate which is completely dissociated in acetonitrile according to conductance measurements by Kratochvil and Yeager (20). The results of Hathaway et al. (8), however, indicate a small (17) I. Y . Ahmed and C. D. Schmulbach, J. Phys. Chem., 71, 2358 ( 1967). (18) R. L. Kay, B. J. Hales, and G. P. Cunningham, ibid.,p 3925. (19) R. P. T. Tomkins, E. Andalaft, and G. J. Janz, Trans. Furuduy Soc., 65,1901 (1969). (20) B. Kratochvil and H. Yeager, J. Phys. Cliem.,73,1963 (1969).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 9, AUGUST 1972
1657
0
Table I. Formation Constants of Copper(1) Complexes in Acetonitrile (25 “C) K, for titrant Titrant salt (mol-1) Supporting electrolyte log P1 log P 2 0.1M EtrNClOa 4.8 10.6 EtJNCl 350 0.1M Et aNCIO4 5.9 IO. 8 0.1 M EtaNClOc 4.2,4.3 10.2,10.2 0.1M EtaNClOa 3.8 7.7 EtrNBr 101 0.1M EtaNClOd 3.4.3.5 7.3.7.2 5b.c 0.1M EtrNClOa 3.2 64 ETrNI 0.1M EtNcC104 3.1,3.1 5.8,5.9 0 . 1M NaCIOa 3.1 6.2 Od NaI 0 . 1M NaC104 3.0 5.6,5.6 Thiourea ... 0 . 1M EtaNClOi .. 6.3 Reference ( 2 2 ) . * Reference (23). c Reference (24). d Reference (19). 1 Present work.
possibility of weak ion-pair formation and Schneider and v.Zelewsky (21) estimated the thermodynamic value of the association constant for copper(1) perchlorate between 20 and 40 mol-’. N o data are available on the ion-pairing of such species as Et4N+, Na+, Cuiz-, . . . . The calculated constants, therefore, are formal constants in which no differentiation has been made with regard t o free and ion-paired copper and complex species. Using Equations 2-4 and assuming the concentration of CuL t o be negligibly small at ratios of total ligand to total copper greater than two (as is indicated by the presence of only one break in the titration curve), P2 can be calculated from the equation P2
=
[CU+l, - [Cu+l [CU+l
*
[
]
[L-I, - 2([CU+l, - [Cu+l) 1 K,[Cu+]
+
-2
(5)
which is readily derived from the equations at hand (in the actual calculations [Cu+] i Z[Et4NC10& was used as a simplification). Once the value of p 2 is determined, values of Pl were calculated from the data at total ligand t o total copper ratios of less than two. Combining expressions for P1, P2, and total copper(1) ion concentrations gives
Similarly, combining expressions for ligand concentration gives
P1, P2, K, and total
Reference p.w.f (3) (4)
p.w.1 (. 4,) p.w.1 (4)
p.w.f (4)
p.w.f
sults in a quadratic expression from which the free ligand concentration may be calculated PdCu+l[L-I2
+ (1 + Ka[Cu+l[L-l +
([CU+], - [L-], - [CU+]) = 0 (8) It is then possible t o calculate PI from either Equation 6 or Equation 7. The formation constants determined in this work are summarized in Table I. The values reported here are generally somewhat higher than those given by Senne and Kratochvil (4) but the relative stability of the different copper(1) halide complexes is the same in both studies. The cuprous sulfide-membrane ion selective-electrode might be particularly useful for the study of the interaction of copper(1) ions with organic ligands in acetonitrile. As a n example, the complex formation between copper(1) ions and thiourea was investigated. The experimental data could be treated in the same way as was done for the copper(1) halide complexes by assuming that CU[S=C(”~)~]~+ is the predominant complex species in solution. The data obtained for total ligand to total copper ratios greater than two fitted Equation 5 over the entire range of concentrations used in this study. For total ligand t o total metal ratios of less than two small deviations from Equation 5 possibly could indicate the formation of a weak 1 :1 complex. Only the value of the second complex formation constant has been determined, however, and is listed in Table I. It may be noted that Cu[S=C(”2)z]2+ is also the predominant complex species found in aqueous solutions under similar experimental conditions(25).
p 96. (23) A. C . Harkness and H. M. Daggett, Jr., Can. . I Chem., . 43,
RECEIVED for review February 1, 1972. Accepted March 29, 1972. L. H. is grateful to the National Fonds voor Wetenschappelyk Onderzoek, Belgium, for a travel stipend and to the Department of Chemistry, Katholieke Universiteit te Leuven, for granting leave of absence during his stay at SUNY a t Buffalo. Work supported by grants from the National Science Foundation and the Environmental Protection Agency.
1215(1965). (24) G . Kortiim, S. G. Gokhale, and H. Wilski, 2.Phys. Cltem. (Frankfurt am Main), 4,86 (1955).
(25) Gmelin’s “Handbuch der Anorganischen Chemie,” 8 Auflage, 60 Kupfer Teil B. Lieferung B, p 246, 1462.
Combining Equation 6 and 7 such that
PI is
eliminated re-
(21) W. Schneider and A. v.Zelewsky, Helo. Cliim. Acta., 46, 1848 (1963). (22) C. W. Davies, “Ion Association,” Butterworth, London, 1962,
1658
ANALYTICAL CHEMISTRY, VOL. 44, NO. 9, AUGUST 1972
