Anal. Chem. 1980, 52, 59-62
L
0
I
J
50
IO0
mol % S Figure 3. Comparison of free energies of transfer of copper(1)ion from water to acetonitrile and of copper(I1) ion from water to dimethyl-
formamide and dimethyl sulfoxide metal, it is stable in AN. The reason appears to be that the interaction of singly-charged d'O ions such as copper(1) and silver(1) [not multicharged d'' ions such as zinc and cadmium ( 3 3 ) ]with AN is strengthened by a large contribution from back-bonding to the 7r-system of AN. In Figure 3 preferential solvation of copper(1) ion by AN in AN-W mixtures is compared to that of copper(I1) ion by DMF and DMSO in DMF-W and DMSO-W mixtures.
LITERATURE CITED (1) J. J. Lagowski, Ed., "The Chemistry of Nonaqueous Solvents", Academic Press, New York, 6 volumes, 1965-1978. (2) J. F. Coetzee and C. D. Ritchie, Eds., "Solute-Solvent Interactions", Marcel Dekker, New York, 2 volumes, 1968 and 1976. (3) A. K. Covington and T. Dickinson, Eds., "Physical Chemistry of Organic Solvent Systems", Plenum Press, New York. 2 volumes, 1973. (4) E. M. Kosower, "An Introduction to Physical Organic Chemistry", Wiley, New York, 1968. (5) V. Gutmann, "The Donor-Acceptor Approach to Molecular Interactions", Plenum Press, New York, 1978.
59
(6) C. Treiner, P. Tzias, and M. Chemla. J . Chem. SOC.,Faraday Trans. 7, 72, 2007 (1976). (7) G.R. Hedwig and A. J. Parker, J . Am. Chem. SOC.,96, 6589 (1974). (8) J. F. Coetzee and E. J. Subak, Jr., Rev. Chim. MinErale, 15, 40 (1978). (9) 0.Popovych, in "Treatise on Analytical Chemistry", I.M. Kolthoff and P. J. Elving. Eds., Part I, Voi. 1, 2nd ed.,Wiley, New York, 1978, p 711. (IO) R. Alexander, A. J. Parker, J. H. Sharp, and W. E. Waghorne, J . Am. Chem. SOC..94, 1148 (1972). (11) A J. Parker, Electrochim. Acta, 21, 671 (1976). (12) A. J. Parker and W. E. Waghorne, Aust. J . Chem., 31, 1181 (1978). (13) T. Kakutani, Y. Morkhiro, M. Senda, R. Takahashi, and K. Matsumoto, Bull. Chem. SOC.Jpn., 51, 2847 (1978). (14) J. F. Coetzee, D. Frollini. C. G. Karakatsanis, E. J. Subak. Jr., and K. Umemoto, in "Characterization of Solutes in Nonaqueous Solvents", G. Mamantov. Ed., Plenum Press, 1978, p 1. (15) J. F. Coetzee and G. R. Padmanabhan, J . Phys. Chem., 66, 1708 (1962). (16) J. F. Coetzee and J. L. Hedrick, J . Phys. Chem., 67, 221 (1963). (17) G.Heijne, W. Van der Linden, and G. Den Boef, Anal Chim. Acta, 69, 287 (1977). (18) Yu-Keung Sze and D. E. Irish, Can. J . Chem., 53, 427 (1975). (19) J. Mitchell, Jr., D. M. Smith, E. C. Ashby, and W. M. D. Bryant, J . Am. Chem. SOC.,63, 2927 (1941). (20) J. F. Coetzee and D. Froilini, unpublished results, 1979. (21) G. Rechnitz and N. Kenny, Anal. Lett., 2 , 395 (1969). (22) A. Diamond, A. Fanelli, and S. Petrucci, Inorg. Chem., 12, 611 (1973). (23) R. P. Buck, in "Ion-Selective Electrodes in Analytical Chemistry", H. Frieser, Ed., Vol. 1, Plenum Press, 1978, p 1. (24) B. G.Cox, A. J. Parker, and W. E. Waghorne, J. Phys. Chem.. 78, 1731 (1974). (25) I.M. Kolthoff and F. G. Thomas, J . Phys. Chem., 68, 3049 (1965). (26) I. M. K o h f f and M. K. Chantooni, Jr., J. Phys. Chem., 76, 2024 (1972). (27) J. I.Kim, J . Phys. Chem., 62, 191 (1978). (28) M. Abraham and A. Nasehzadeh, Can. J . Chem,., 57, 2004 (1979). (29) C. G. Karakatsanis, Ph.D. thesis, University of Pittsburgh. 1978. (30) I. M. Kolthoff and M. K. Chantooni, Jr.. Anal. Chem., 50, 1440 (1978). (31) F. Basolo and R. G. Pearson, "Mechanisms of Inorganic Reactions", Wiley, New York, 1967, p 226. (32) J. F. Coetzee and K. Umemoto, Inorg. Chem., 15, 3109 (1976). (33) G.R. Hedwig, D. A. Owensby, and A. J. Parker, J . Am. Chem. SOC., 97, 3888 (1975). (34) J. F. Coetzee, Pure Appl. Chem., 49, 27 (1977).
RECEIVED for review July 23,1979. Accepted October 10,1979. We acknowledge support by the National Science Foundation under grant number CHE-7727699.
Solvent Dependence of Equilibrium Constants and Dissociation Rate Constants of the Mono-Complex of Nickel(I1) Ion with 4-Phenyl pyridi ne J. F. Coetzee' and C. G. Karakatsanis Department of Chemistry, University of Pittsburgh, Pittsburgh, Pennsylvania 15260
Equilibrium constants for formation of the mono-complex of nickel( 11) ion with 4-phenylpyridine in a range of protlc and aprotic solvents were determined spectrophotometrically. From these equilibrium constants and previously reported rate constants for formation of the complex, dissociation rate constants are calculated. Correlations between the equilibrium and rate constants and the donor strengths of the solvents are discussed.
The development of nonaqueous solution chemistry is still being impeded by a serious lack of necessary thermodynamic data. While a considerable reservoir of such information has been accumulated particularly during the past two decades for two important classes of chemical reactions, viz., those involving proton and electron transfer, much less has been 0003-2700/80/0352-0059$01.OO/O
done on the equally important class of coordinat.ion reactions. This is particularly true for the reactions of multicharged transition metal ions, many of which can be studied in nonaqueous solutions only with difficulty, owing to such experimental problems as lowering residual water concentrations sufficiently, doing without electrochemical sensors since these are not available for many such ions, etc. While Gutmann in particular has collected and systematized much useful information, especially of a qualitative and semiquantitative nature, on the coordination reactions of metal ions in nonaqueous solvents ( I ) , it would be beneficial for analytical and other practical reasons, as well as for the development of solution theory, to have more extensive quantitative information available. One example of the very few systematic quantitative studies carried out in this field is the detailed dissection by Ahrland (2) of the thermodynamics of formation of halide and pseudohalide complexes of zinc and cadmium 1979 American Chemical Society
60
ANALYTICAL CHEMISTRY, VOL. 52, NO. 1, JANUARY 1980
ions in dimethyl sulfoxide as solvent. We report here data illustrating the effect of a series of solvents of widely differing donor and other properties on the stability of nickel(I1) complexes. Nickel(I1) is a n example of those ions for which free energies of transfer cannot be obtained by the preferred methods described in the previous paper ( 3 ) . Similarly, evaluation of enthalpies of transfer of nickel(I1) ion is also complicated by experimental problems ( 4 ) . Hence, in order to learn more about relative donor strengths of different solvents toward this acceptor, it is necessary to resort to other, less direct, information, such as the equilibrium and rate constants reported here. Nickel(I1) is a n important ion for a variety of reasons, not the least of which is its exceptional significance in studies of solvent effects on the reaction kinetics of labile complexes ( 5 ) . T h e result is that, atypically, much more is known about kinetic than about thermodynamic aspects of the reactivity of this particular ion in nonaqueous solutions. However, its kinetic behavior is complex and supporting thermodynamic data are keenly needed to improve understanding of the kinetics. We also report here new kinetic data for the formation and dissociation of nickel(I1) complexes to supplement those summarized before ( 5 ) .
EXPERIMENTAL Reagents. Solvents were purified as described elsewhere (3, 5 ) . Nickel(I1) perchlorate hexahydrate was prepared by treating excess nickel(I1) carbonate (Baker analyzed reagent) with perchloric acid (Baker analyzed reagent). The resulting solution was filtered to remove excess nickel carbonate and the nickel perchlorate was precipitated by cooling, recrystallized from dilute perchloric acid (pH 3 to 4), and finally washed with small amounts of deionized-distilled water to remove traces of acid from the surface of the crystals. Nickel(I1) was introduced as Ni(C104)2-nH20( n 2), obtained by heating the recrystallized hexahydrate in a vacuum oven at 85-90 "C for more than a day. Its solutions were dried further with molecular sieves or by application of a vacuum, as described elsewhere for copper(I1) perchlorate (3). 4-Phenylpyridine (Aldrich) was used without further purification. Isoquinoline (Aldrich, 97 5%) was distilled alone in vacuo in a Perkin-Elmer Model 251 Auto Annular Still equipped with a Teflon spinning band at a reflux ratio of 200:1, with a pot temperature of 126 "C and a head temperature of 62 "C; the middle 60% was collected. Apparatus. Equilibrium constants were determined spectrophotometrically using a Bausch and Lomb Shimadzu Model UV-200 spectrophotometer. Temperature control was achieved by pumping water from a Forma Model 2095 bath through the cell holder. During measurements the temperature was maintained at 25.0 & 0.3 "C. Kinetic measurements were made with a Durrum Instrument Co. Model D-110 stopped-flow spectrophotometer equipped with a Kel-F flow system; details have been reviewed before ( 5 ) . Procedures. For the equilibrium constants measurements in each solvent, ultraviolet spectra of solutions containing only nickel(I1) perchlorate and only 4-phenylpyridine were recorded separately between 310 and 240 nm. Spectra of a series of solutions containing varying amounts of nickel(I1) as well as a constant amount of 4-phenylpyridine (usually 1.5 to 2.5 X M) were also recorded over the same wavelength range. The concentration of nickel(I1) in the latter solutions was typically varied through 5 to 7 values until formation of the complex was complete. Thus, depending on the solvent, nickel concentrations from a low value of 5 X lo4 M to a high value of 8 X M were used. Calculations were carried out at 272.5 nm in all solvents except DMSO for which they were done at 273.7 nm. Procedures used in the kinetic measurements have been described before ( 5 ) .
-
RESULTS AND DISCUSSION Rate constants, kl,f, and equilibrium constants, Kl,f, for formation of the mono-complex of nickel(I1) ion with 4phenylpyridine (L) in a variety of solvents (S)were determined
as described above. Previous studies (5) have shown that in all solvents tested such reactions occur by a dissociative interchange mechanism in which the first step is rapid formation of an outer-sphere complex (OSC), followed by a second step in which rate-limiting loss of a solvent molecule from the inner coordination shell of the metal ion occurs with concerted entry of the ligand to form an inner-sphere complex (ISC): Niss2+
M
+ L e NiSs2+L e NiLS$+ osc ISC
4- s
(1)
Under pseudo-first-order conditions with a large excess of metal ion (M) present, the simplified rate law is hobsd
=
kl,f[Ml
+ kl,d
(2)
where kobd is the experimentally observed pseudo-first-order rate constant (in s-l), kl,f is the overall second-order rate constant for ligant substitution leading to formation of the I s c (in L mol-' s-'), and k l , d is the first-order rate constant for dissociation of the ISC (in s-l). Hence, a plot of kobsd vs. [MI is linear with a slope equal to kzf and a n intercept equal to k1,d. It also follows that =
Kl,f
kl,f/kl,d
(3)
where Ki,f
=
~ I S C / ~ M ~ L
(4)
For nickel(I1) complexes in a variety of solvents, kl,f typically falls in the range of stopped-flow spectrophotometry or relaxation techniques, such as pressure-jump relaxation; many such values are already available ( 5 ) . Values for k l , d can be determined by inducing dissociation of the complex with such ions as mercury(I1) (6, 7 ) . Alternatively, for the special case of complexes of intermediate stability, kl,f and k1.d can be determined simultaneously from Equation 2, although circumspection is required because impurities (such as water) may contribute to the magnitude of the intercept obtained in such pseudo-first-order plots ( 5 ) . Evaluation of overall equilibrium constants, Kl,f, by kinetic measurements is therefore a viable alternative to thermodynamic methods; it should be especially useful for very stable complexes of metal ions for which electrochemical sensors are not available, but little has been done so far t o exploit this possibility. Turning now t o the conventional spectrophotometric determination of Kl,f,it can be shown (8) that for a solution containing (for example) nickel(I1) ion, 4-phenylpyridine and their mono-complex, the total absorbance is given by
AT = ( E ) ( B ) ( C X )
(5)
and that
where
In these equations D,X ,and DX represent nickel(I1) ion, 4-phenylpyridine, and complex, respectively; CD and C X are the initial concentrations of nickel(I1) ion and 4-phenylpyridine; [D] is the equilibrium concentration of nickel(I1) ion; ED,EX,and EDX are the molar absorptivities of nickel(I1) ion, 4-phenylpyridine, and complex, respectively; E is the so-called "apparent absorptivity"; B is the optical path
ANALYTICAL CHEMISTRY, VOL. 52, NO. 1, JANUARY 1980
Table I. Equilibrium Constants for Formation of Mono-Complexes of Nickel(I1) Ion with 4-Phenylpyridine and Isoquinoline in Various Solvents at 25 “ C log Kl,f isoquin-
4-phenylpyridinen
solvent dimethyl sulfoxide dimethylformamide water methanol ethanol 1-propanol
2-propanol 50 mol % methanol in 2-propanol acetonitrile propylene carbonate a
Table 11. Rate Constants for Formation and Dissociation of Mono-Complexes of Nickel(I1) Ion with
4-Phenylpyridine and Isoquinoline in Various Solvents a t 25 C 4-phenyl-
ridin e _p v_ -__ -~ log logb kl,P h , d
olineb
1.4,‘ -
1.4,
1.9,
solvent
1.8, 1.9 2.8 3.1 3.9
1.8, 1.7
dimethyl sulfoxide dimethylformamide water methanol ethanol 1-propanol 2-propanol 50 mol % methanol in 2-propanol acetonitrile propylene carbonate
-
-
2.4
-
4.Od
-
Determined spectrophotometrical1
61
6.,e
this work, except
3.2 3.3 3.3 2.1
3.9 4.2 5.2
1.8‘ -
1.4 0.2
isoquinoline
-
log
h,iZ
logb h,d
3.5 3.4 3.3
2.0 1.4 1.5
2.0
0.3
1.1 1.0
4.3
-
1.3
-
-
2.9
0.5
5.3 -
-
3.0 5.3
-
3.1
-0.9d
-
-
- 1.6d
value for water which is from ref. 9. {Determined spectrophotometrically ( 7) except where otherwise noted. Value calculated from Equation 2 using rate constants from ref. 6 is 1.5 for I = 0.21 M. Calculated from Equation 2 using rate constants from refs. 7 and 10. e Approximated by substituting in Equation 2 k l d value for isoquinoline from ref. 7 and h1.f value for 4-phenylpyridine from ref. 11.
a Measured directly; units: L mol-’ s-’. Values for isoquinoline in 1-PrOH and 2-PrOH are new; we have summarized sources of other values elsewhere ( 5 ). Calculated from directly measured values of h1.f arid K l , f by Equation 2, except where otherwise noted; units: s-’. ‘ Directly measured value a t I = 0.21 M is 1 . 7 (6). Directly measured value (7).
length in cm; and F is the composite value for the collection of activity coefficients, given by
where Yt is the “transfer activity coefficient” (12) which is related to the Gibbs free energy of transfer by Equation 10
F = -f D X
(8)
fDfX
Hence, a plot of [ E - E D ( C D / C X ) ]vs. [ E D ( C D / C X )+ E X - E ] / [ D ] should ] be linear with a slope equal to F/Kl,,. Good linearity was obtained for all solvents. For example, for a series of solutions in ethanol consisting of a constant total concentration of 1.83 x 10-5 M 4-phenylpyridine and total concentrations of nickel(I1) perchlorate varying from 1.13 X to 1.41 X lo-* M, the linear correlation coefficient was 1and j D jDx, F in Equation 8 will be close 0.998. Since fx t o unity. Equilibrium constants for formation of the monocomplex of nickel(I1) with 4-phenylpyridine determined in this manner are given in Table I; also included are corresponding equilibrium constants reported before (7)for the related ligand isoquinoline. T h e fact that the phenyl ring is fused to the pyridine ring in isoquinoline while in 4-phenylpyridine it is free to rotate has little influence on the equilibrium constant. We have been conservative in the number of significant figures claimed in Table I. While the reproducibility was generally good for solutions dried as described (for example, typical standard deviations in log Kl,f were fO.01 for ethanol and f0.03 for 2-propanol), the accuracy of the numbers is expected to be less good for the weak donor solvents in which even trace amounts of residual water and other donor impurities influence the reactivity of nickel(I1) ion ( 5 ) . We recommend that authors exercise restraint in claims made or implied for the reliability of data obtained in relatively inert solvents. However, it is to be noted t h a t the equilibrium constants in Table I cover a wide range of 5.5 powers of ten, so that the estimated uncertainties are unimportant. The only additional equilibrium constants known for nickel complexes in nonaqueous solvents are those for a few anionic ligands summarized before ( 5 ) . Variations in the equilibrium constant of a given complex in different solvents are determined by the free energies of transfer of the three solutes involved, viz., the free metal ion, the free ligand, and the complex, so t h a t Equation 4 can be written as
-
Kl,fo /Kl,f
-
=
Yt,ISC/?t,MYt,L
(9)
AGt,io = RT In Y ~ , ~
(10) and Klio and Kl,fare the equilibrium constants in a reference solvent and in another solvent, respectively. Free energies of transfer of uncharged ligands such as 4-phenylpyridine among nonaqueous solvents are generally relatively small; for transfers from water to nonaqueous solvents, AGto values are larger and negative owing to positive values of ASto for the transfer of solutes containing large hydrophobic moieties. For example, AG,O values for 4-phenylpyridine from water to the nonaqueous solvents listed in Table I are all near -4 kcal mol-’, and -?‘Asto of ca. +3 consisting of contributions from Mt0 and -7 kcal mol-’, respectively (13). T h e effect of such differences in solvation of the free ligand on the relative stabilities of the complex in different solvents will be counteracted by parallel differences in solvation of the ligand in the complex, particularly in aprotic solvents in which such solvation is largely entropic. It is therefore to be expected that the major contribution to observed differences in equilibrium constants comes from differences in solvation of the free metal ion and the metal ion in the complex. Hence, for the aprotic solvents, the order of increasing equilibrium constants, as listed in Table I, should in general also be the order of decreasing donor strengths of these solvents toward nickel(I1) ion. For the protic solvents, particularly water, the donor strengths should be somewhat lower than indicated by the equilibrium constants because enthalpic contributions by H...N hydrogen bonding to solvation of the free ligand will be absent from solvation of the ligand in the complex. We conclude that the overall order of equilibrium constants in Table I is consistent with an order of donor strengths towards nickel(I1) ion that is not substantially different from the order established for copper(I1) ion as acceptor (3). In Table I1 rate constants for formation and dissociation of the mono-complexes of nickel(11) ion with 4-phenylpyridine and isoquinoline are compared for different solvents. The total variations in formation and dissociation rate constants are ca. 3 powers of ten each and therefore contribute about equally to the total variation in equilibrium constant. The formation rate constants show little correlation with expected donicities (donor strengths) of the solvents. Such correlations are
62
Anal. Chem. 1980, 52, 62-66
generally not to be expected. Since a rate constant in a given solvent depends on the difference in free energy of transition and ground states, variations in solvent donicity may influence the rate constant in a complex manner. For formation of the complex, which proceeds by a dissociative interchange mechanism, an increase in donicity of S will decrease the free energy of the transition state, MS5,and also that of the ground state, MS6, but the difference between these two decreases in free energy need not correlate in a simple manner with the increase in donicity of S. Somewhat better correlations are expected for the dissociation process, because the lability of L in MLS5 should increase with increasing donicity of S in a more direct manner. In fact, correlations of log kl,d for the mono- and bis-complexes of nickel(I1) ion with thiocyanate ion ( 1 4 ) , and of m l , d * for the mono-complexes of nickel(I1) ion with thiocyanate ion (15) and isoquinoline (7), with Gutmann donor numbers have been reported. However, it was pointed out by Hoffmann ( 1 4 ) that such correlations of rate constants are expected to break down for solvents having the ability to donate hydrogen bonds and for leaving ligands t h a t are good hydrogen bond acceptors. The data in Table I1 show this to be the case. For 4-phenylpyridine and isoquinoline, the order of decreasing dissociation rate constants does follow the expected order of decreasing donor strengths of the aprotic solvents, but not of the protic solvents. In conclusion, rate constants for the formation of the mono-complexes of nickel(I1) ion with 4-phenylpyridine and isoquinoline show little correlation with the expected donor strengths of solvents. Dissociation rate constants show better
correlations, but only in aprotic solvents. Equilibrium constants show the best correlations for all solvents, including protic solvents except water. In order to make correlations of equilibrium constants of complexes with the donor strengths of solvents more quantitative, it will be necessary to obtain information on the solvation of the metal ion, as well as the ligand, in the complex. Quantitative information on these interactions is not available a t present and will be difficult to obtain.
LITERATURE CITED (1) V. Gutmann, "The Donor-Acceptor Approach to Molecular Interactions", Plenum Press, New York, 1978. (2) S. Ahrland, Pure Appl. Chem., 51, 2019 (1979). (3) J. F. Coetzee and W. K. Istone, Anal. Chem., preceding paper in this issue. (4) J. F. Coetzee and E. J. Subak, Jr., Rev. Chim. Mihrale, 15, 40 (1978). (5) J. F. Coetzee, D. Frollini, C. G. Karakatsanis. E. J. Subak, Jr., and K. Umemoto, in "Characterization of Solutes in Nonaqueous Solvents", G. Mamantov, Ed., Plenum Press, New York, 1978, p 1. P. Moore and D. M. W. Buck, J . Chem. SOC. Dalton Trans., 1602 (6) (1973). (7) P. K. Chattopadhyay and 9.Kratochvil, Inorg. Chem., 15, 3104 (1976). (8) R. W. Ramette, J . Chem. Educ., 44, 647 (1967). (9) J. F. Coetzee and T. Takahashi, unpublished results (1978). (10) P. K. Chattopadhyay and J. F. Coetzee, Anal. Chem., 46, 2014 (1974). (11) J. F. Coetzee and K. Umemoto, Inorg. Chem., 15, 3109 (1976). (12) B. TrBmillon and J. F. Coetzee, Pure Appl. Chem., 50, 587 (1978). (13) J. F. Coetzee and D. Frollini, unpublished results (1979). (14) H. Hoffmann, Pure Appl. Chem., 41, 327 (1975). (15) P. K. Chattopadhyay and B. Kratochvil, Can. J . Chem., 55, 1609 (1977).
RECEIVED for review July 23,1979. Accepted October 10,1979. We acknowledge support by the National Science Foundation under grant number CHE-7727699.
Reverse Pulse Polarography Janet Osteryoung * Department of Chemistry, State University of New York at Buffalo, Buffalo, New York 14214
Emilia Kirowa-Eisner Department of Chemistry, University of Tel-Aviv, Ramat-A viv, Tel-A viv, Israel
The technique of reverse pulse polarography is described and examples are given of applications to (i) characterization of reversibility, (ii) identification and quantitative determination of products of an electrode reaction, including studies of intermediates, and (iii) study of reactions poorly separated from solvent or electrode decomposition.
Pulse polarography is a voltammetric technique originally introduced by Barker as an outgrowth of his work on square wave polarography ( 1 ) . It has been developed and employed for the study of electrode processes and for analytical purposes most notably by R. A. Osteryoung (2,3). Exploitation of pulse polarography and the development and implementation of methods for analysis and for electrochemical studies has been stimulated and influenced by the production of inexpensive and reliable commercial instrumentation. Much of the research in pulse polarography and related techniques has involved instrumental refinement or elaboration of Barker's original ideas. Thus we have seen the development of al0003-2700/80/0352-0062$01.00/0
ternate drop and constant potential pulse polarography, square wave and staircase voltammetry, and differential pulse and other pulse voltammetric stripping techniques ( 4 ) . We have been more interested in the full exploitation of techniques which can be carried out simply with existing instrumentation through the discovery of applications in which these techniques display unexpected power for solving difficult chemical problems. In this paper we present a conceptually new way of applying normal pulse polarography (NPP) which employs standard commercially available instrumentation but which constitutes a new tool for electrochemical investigations. We call this application reverse pulse polarography (RPP). Oldham and Parry used R P P (which they called scan-reversal pulse polarography) to characterize electrode reversibility (5). I t has been employed recently for this purpose also by Saito and Himeno (6). Here we describe the application of R P P to characterization of electrode reactions, including the investigation of reversibility of electrode reactions, identification and quantitative determination of products of an electrode reaction, including intermediates and study of reactions poorly separated from electrode or solvent decomposition. Another 0 1979 American Chemical Society
