(4)J. W. Cares, Am. lnd. Hyg. Assoc. J., 20, 463 (1968). (5) J. Leroux, Sfaub, 29, 157 (1969). (6) C. L. Luke, Anal. Chim. Acta, 37,267 (1967). (7)C. L. Luke, Anal. Chim. Acta. 41, 237 (1968). (6)P. Grenfelt, A. Akerstrom, and C. Brosset, Atmos. Environ., 5, 1 (1971). (9)C. Brosset and A. Akerstrom, Atmos. Environ., 8, 661 (1972). (10)J. Leroux, Occup. HeaRhRev., 21, 19 (1970). (11)H. R. Bowman, J. G. Conway, and F. Asaro, Environ. Sci. Technol., 8, 558 (1972). (12)C. L. Luke, T. Y. Kometani, Y. E. Kessler, and T. C. Loonls. Environ. Sci. Techno/., 8, 1105(1972). (13)L. Beitz and U. Jecht, Siemens Rev., 30, 3 (1972). (14)J. L. Johnson, A. C. Ottoiini, F. A. Forster, and R. B. Loranger, Microchim. Acta. 2, 145 (1974). (15)T. R. Dittrich and C. R. Cothern, J. Air Pollut. Contr. Assoc., 21 (1l),716 (1971). (16)J. V. Gilfrich, P. G. Burkhaiter, and L. S. Birks, Anal. Chem., 45, 2002 (1973). (17)R. H. Hammerke. R. H. Marsh, K. Rengan, R. D. Giauque, and J. M. Jakievic, Anal. Chem., 45, 1939 (1973). (18)J. R. Rhodes, A. H. Pradzynski, and R. D. Sieberg, /SA Trans., 11 (4), 337 (1972). (19)J. R. Rhodes, Environ. Sci. Technol., 8, 922(1972). (20)J. R. Rhodes, I€€€ Trans Nucl. Sci., NS-21, 608 (1974). (21)A. Rolla, P. Frigieri, A. Gireiia, and R. Trucco, Chim. lnd. (Milan), 55, 623 (1973). (22)J. Leroux and M. J. Mahmud, J. Air Polluf. Contr. Assoc., 20, 403 (1970).
(23)P. L. Sciaraffa and C. A. Ziegler, /sot. Radiat. Technol., 8 , 164 (1970& 1971). (24)J. Leroux and M. Mahmud, Anal. Chem., 38, 77 (1966). (25) P. Frigieri, R. Trucco, R. Anzani, and E. Caretta, Chim. lnd. (MNen),54, 12 (1972). (26) P. Frigieri and R. Trucco, X-Ray Spectrom., 3, 40 (1974). (27)F. P. Brady and T. A. Cahiii, "Development of X-Ray Fluorescence Anaiysis and Application"; Report UCD-CNL 166 (1973). (28)J. R. Rhodes and C. B. Hunter, X-Ray Spectrom., I,113 (1972). (29) H. L. Rook and E. A. Schweikert, Anal. Chem., 41, 958 (1969). (30)R. Tertian, Doctoral Thesis, Universite de,Paris Vi (1972). (31)D. De Soete, R . Gijbels, and J. Hoste, Neutron Activation Analysis", Wiiey-lnterscience, London, 1972. (32)R. Dams and R. Heindryckx, Afmos. Environ., 7, 313 (1973). (33)K. Rahn, "Sources of Trace Elements in Aerosols-An Approach to Clean Air", The University of Michigan Technical Report, 1971. (34) "Study of National Air Pollution by Combustion", instituut voor Nucieaire Wetenschappen, Rijksuniversiteit Gent & Service de Chimie Medicale, Toxicoiogie et Hygiene, Universite de Liege, Belgium, Progress Report
1972 (1973). (35)T. B. Johansson, R. E. Van Grieken, and J. W. Winchester, J. Geophys. Res.. in press.
(36)D. C. Camp, A. L. Van Lehn, J. R. Rhodes, and A. H. Pradzynski, X-Ray Spectrom., in press.
(37)Micro Matter Co., 197 34th St. East, Seattle, Wash.. 98102.
RECEIVEDfor review December 30, 1974. Accepted May 27, 1975.
Electrochemical Study of Lithium(1) Interactions with Radical Anions Derived from 9, IO-Anthraquinone and I-Hydroxy-9,lOAnthraquinone by Cathodic Reduction in N,NDimethylformamide Solutions Mihael TkalEec, Ivan Filipovic, and Ivan Piljac' Laboratory of lnorganic Chemistry, Faculty of Technology, University of Zagreb, PO6 179, Zagreb, Croatia, Yugoslavia
By application of direct current polarography and cyclic voltammetry to 1-hydroxy-9,lO-anthraquinone in DMF soiutions in the presence of Li+ ion, the complexity of the eiectrode reduction of HOAO to the radical anion was demonstrated and the factors were pointed out which must be considered in establishing the exact quantitative reiationship between the electrochemical magnitudes and the strength of the metal ion-ligand Interaction. Mathematical models were constructed by deriving a theoretical expression from the Nernst equation for different types of metal ion-ligand Interactions and from these models, a computer simulation of polarographic waves, as well as the caicuiation of "theoretical" half-wave shift was made.
Supporting electrolyte cations influence the electrochemical behavior of organic electroactive species in aprotic solvents. In polarographic experiments, this influence is manifested as a shift of half-wave potential. The shift observed, due to ion-pair or complex-species formation with products arising from the electrode reaction, depends on the properties of the cations involved and their concentration and it is directed toward the more positive potential. Association processes affecting the electrochemical properties of organic molecules have been investigated (1-9). The organic molecule-carrying electron-donor group can, however, itself enter into stable complex with the cation, this association being independent of any accompanying Author to whom correspondence should be addressed.
electrode reaction. This kind of complex formation will shift the half-wave reduction potential of the organic electroactive species to more negative values. The strength of the interaction between the cation and organic molecule can be calculated from the magnitude of the observed shift, as has been done by other authors (IO). Peover and Davies ( I ) derived a mathematical relationship comprising the half-wave potential shift, the metal ion concentration, and the association constant of the complex species formed, starting with the Nernst equation. For a strong interaction involving the product of the one-electron reaction, which leads to the formation of a complex species with the composition M,+R.-. (Symbols to be used throughout: M+, monovalent metal ion; R, organic electroactive species; Re-, radical anion; and R2-, dianion forms of the latter.) This relationship is given as
where ( E ~ / Zand ) ~( E ~ l zare ) ~ the half-wave potentials for processes occurring with and without complex formation, respectively; K , is the equilibrium constant for the M,+R-formation, p is the metal ion stoichiometric coefficient ( p > 1) and CM+is the total metal ion concentration (the symbol C will subsequently be used for total concentration, as opposed to free or equilibrium concentration which will be denoted by the species symbol in brackets); other symbols have the usual meanings. As stated above, Equation 1 holds true for strong interaction only, i.e., only under circumstances where the free product concentration a t the working electrode surface, [Ra-]", is much less than the correspond-
ANALYTICAL CHEMISTRY, VOL. 47, NO. 1 1 , SEPTEMBER 1975
* 1773
ing organic molecule concentration, [ R ] O , and total ion concentration is in large excess over total electroactive species concentration, so that CM+ may be set equal to [M+] as a first approximation. The unreacted organic electroactive species is implicitly supposed not to undergo complex formation. Using the above assumptions, Lasia and Kalinowski (6) deduced another relationship for the formation of M+(R.-)l type of complexes. They show that the half-wave potential shift caused by complex formation depends not only on metal ion concentration but also on the concentration of the organic electroactive species. In the present work, the influence of the complex formation of Li+ ion with the electrode reaction product of 1hydroxy-9,lO-anthraquinone(HOAQ), as well as with HOAQ itself, in the N,N-dimethylformamide (DMF) solutions, on the shift of half-wave potential was studied. In the previous work ( I I ) , it was shown that the products of electrode reduction of HOAQ, i.e., radical anion and dianion, form stable lithium complexes with metal-to-ligand ratios of 1:2 and 1:1,respectively. Logs of respective formation constants were determined: log p2 = 5.49 f 0.18 and = 2.94 f 0.14. In DMF, Li+ was found to associate log not only with the products of the electrode reaction, but also with the parent HOAQ; and this additional interaction is expected to influence the position of the half-wave potential of this quinone. The present study is related to the influence of complex formation involving only the first reduction step of HOAQ leading to the radical anion, HOAQ--. In evaluating the results, any association of the unchanged organic molecule and/or its radical anion with the cation from the supporting electrolyte was neglected. Also, the possible hydrogen bond formation (12) between the radical anion and the traces of water from the solvent was disregarded in view of thorough purification of the solvent. A reexamination of some earlier study ( 1 ) on the 9,lOanthraquinone (AQ) system was included at this time under our experimental conditions, in order to compare the behavior of the two quinones in the one-electron reduction to the respective radical anions in the presence of Li+. A mathematical model was constructed by deriving a theoretical expression from the Nernst equation. On this model, computer simulation of polarographic waves, as well as calculation of the “theoretical” half-wave potential shift, was based. The experimentally determined potential shift was compared to simulated curves and their fit or their discrepancies are analyzed. The program for dc polarographic on-line control and data acquisition, as well as for numerical and statistical treatment, is worked out.
EXPERIMENTAL Chemicals. N,N- Dimethylformamide (DMF) was freed of water by azeotropic distillation with excess benzene, followed by drying over anhydrous A1203. The dried solvent was subjected to a final purification by distilling a t reduced pressure (30 mm Hg) and kept under dry nitrogen. Conductivity measurement gave K < 2.0 X lo-’ cm-l. Tetraethylammonium perchlorate (TEAP) was prepared from tetraethylammonium hydroxide. The product was recrystallized twice from water and dried at 333 K in vacuo. 1-Hydroxy-9,lO-anthraquinone(HOAQ) was purified by vacuum sublimation. Its solutions in DMF were reasonably stable. 9,lO-Anthraquinone (AQ) was recrystallized twice from ethanol. DMF solutions of this quinone, however, were unstable and became fluorescent after short standing. Lithium perchlorate was obtained by neutralization of lithium hydroxide solution with concentrated perchloric acid. The salt was purified by recrystallizing twice from water and drying a t 333 K in vacuo. A p p a r a t u s a n d Procedure. The electrochemical apparatus was composed of an analog component based on operational amplifiers 1774
Table I. Effect of Lithium Ion on the Half-Wave Potential of 9,lO-Anthraquinone By sumultaneous
By logarithmic
refinement ( 1 4 )
analysis
rU ernst 12,
‘EI/?I
id
mmoldm-3
mV
LLA
mV
mV
0.0 3.9 18.0 33.0 97.0
866.5 866.1 864.1 863.2 856.0
3.97 3.68 3.70
866.6 864.6 864.1 862.6 855.1
59.6 55.8
CL~*,
3.71
3.52
-E1
slope,
60.0
58.6 58.9
containing a three-electrode cell and an on-line connected minicomputer ( P D P W E from Digital Equipment Corp.). The computer component was programmed to carry out the following functions: control of mercury dropping time (by means of an electromagnetic hammer, giving t = 2 sec); starting and stopping the Integrator; collecting and analyzing data; and actuating an X-Y Plotter. The integral device was utilized for dc polarography while the analog component alone served for cyclic voltammetry. Polarographic experiments were carried out using a mercury dropping electrode adjusted to deliver 2.103 mg sec-I Hg into 0.2 mol dm-3 of TEAP in DMF a t open circuit (head: 53 cm). In addition to recording of actual polarographic waves, graphical tracings of their log transforms were provided in these experiments. Voltammetric measurements and recordings were made with a hanging mercury drop electrode. In both series of experiments, all potentials were referred to a saturated calomel electrode immersed in 1 mol dm-3 aqueous KN03, connected to the main compartment by means of an electrolytic bridge. Ionic strength and temperature were the same in all experiments. Adjustmedt of the former t o I = 0.2 mol dm-3 was achieved by addition of the required amount of TEAP. Thermostating of working solutions was set a t 298.2 f 0.1 K during measurements.
RESULTS AND DISCUSSION Polarography of 9,lO-Anthraquinone. The polarographic one-electron reduction of AQ to AQ-- in a DMF solution, in the presence of varying Li+ concentrations, gave the results presented in Table I. The interaction of the radical anion and Li+ is very weak under the conditions of these experiments, as indicated by small changes in potential shift with variation of Li+ concentration. Actually, the association parameter calculated from the half-wave poten= 0.76. This value agrees well enough tial shift was log = 1.06 obtained under identical experimental with log conditions but using the spectroelectrochemical method with an optically transparent thin layer (OTTL) electrochemical cell (11). I t is, however, important to point out that AQ solutions in DMF were unstable. These solutions on standing after preparation started to emit a green fluorescence which is probably caused by photoreduction of quinone, as suggested by the disappearance of emission upon exposing the solution to atmospheric oxygen. Ethanolic solutions of AQ are reported to behave similarly (13). Interference of photochemical reduction, as well as slight variability of half-wave potential shift with change in Li+ concentration, precluded the more detailed investigation of this system. Polarography of 1-Hydroxy-9,lO-anthraquinone. The first polarographic wave corresponding to one-electron reduction of HOAQ to HOAQS- and its changes with Li+ concentration are shown in Figure 1 and a summary of numerical data is given in Table 11. These results show the apparent polarographic reversibility of the HOAQ reduction a t all Li+ concentrations examined. The Nernst slope (2.3 RTInF),as calculated from the log transforms of recorded polarograms, closely approaches the
ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
~
~~~
Table 11.Effect of Lithium Ion on the Half-Wave Potential of 1-Hydroxy-S, 10-Anthraquinone By simultaneous
By logarithmic
refinement ( 1 4 )
analysis Nernst
CLi+'
mmol dm-3
0.00 0.99 1.98 2.95 4.87 9.50 13.90 22.60 29.70 39.90 54.40
mV
LLA
-E1I23 mV
730.4 725.5 721.8 720.3 717.0 711.9 707.1 703.3 700.4 697.0 695.6
4.05 4.00 3.99 4.08 3.95 3.92 3.86 3.80 3.77 3.63 3.68
730.4 725.5 721.7 720.3 716.9 711.8 706.6 702.8 700.2 696.3 695.6
-E112'
id
slope,
59.6 59.6 59.6 59.8 59.8 59.5 58.2 58.4 58.7 57.2 58.4
mV
98165L321
Flgure 1. Experimental polarographic waves of 1.00 mmol dm-3 HOAQ with different concentration of Li+ and their slope obtained by logarithmic analysis Concentrations of Li+ from 1 to 9: 0.00, 0.99, 1.98, 4.87, 9.50, 13.90, 22.60, 29.70, and 39.90 mmol dm-3
theoretical value of 59 mV. Otherwise, however, the polarographic results were not consistent with the findings obtained by a different method ( 1 1 ) regarding half-wave potential shift and the stability of the complexes formed. This inconsistency calls for consideration of a t least two factors as possible causes, viz., sluggishness in reaching equilibrium, which would invalidate any expression derived from the Nernst equation; and the occurrence of additional interactions, not covered by the original assumptions. T o find a rule governing the change in shape of the polarographic wave and the half-wave potential shift caused by formation of the previously established ( 1 1 ) M+(R--)z complexes, we carried out a computer simulation of a composite process comprising HOAQ and HOAQe- association with Li+, varying the metal ion concentrations. Three alternative patterns of HOAQ and Li+ interactions were taken into account, viz., (A) interaction is negligible; (B) only type M+R complex is formed; and (C) only type M+(R)z complex is formed. Rapid equilibration of the species reacting a t the electrode surface was assumed, so the conditions for utilizing the Nernst equation were fulfilled. In all simulations, the value Pz = antilog 5.49 ( 1 1 ) was used as the formation constant of the complex M+(HOAQ.-)z. Simulation of Shapes a n d Half-Wave Potential Shifts of Polarographic Waves. (A) A one-electron reversible process producing a radical anion R
+ e-
3
R.-
(2)
is followed by a reversible interaction of the product with a metal ion 2R.-
+ M+
M+(R.-)z
(3)
From Equation 3, the formation constant for M+(R--)z is obtained as (4) Let only the free ,organic electroactive species, R, undergo the electrode reaction. The surface concentration of M+(R.-)z a t half-wave potential, is given by:
E/mV vs. EYn:O.O
mV
Figure 2. Simulated polarographic waves for different concentration of metal ion assuming tog P2 = 5.49 Assumed concentrations of metal ion from 1 to 10: 0.0,1.0, 2.0, 4.0, 8.0, 16.0, 32.0, 64.0, 128.0, and 256.0 mmol dm-3
where [XI" was substituted for [M+(R--)2]", and C O R . - is the total surface concentration of the radical anion. CM+is taken as a first approximation for free metal ion concentration [M+]. T o make use of Equation 5, antilog 5.49 was substituted for 6 2 , and the total depolarizer concentration was set as CR = 1mmol dm-3. Half-wave potential shifts for fixed total metal ion concentrations were calculated from: AE1/2
=
--RT In F
COR.-
- 2[X]" = -RT
PI "
F
CR - 4[X]" CR
(6)
Plotting AEllz vs. log CM+,one obtains the theoretical dependence of half-wave potential on metal ion concentration for interactions leading to the M+(R.-)z complex formation. Substitution of Equation 5 into Equation 6 gives, for the strong interaction
RT a i / z =2F In CR +
RT + RT 2 F In C M + + 2F In p2
(7)
a relationship also obtainable from the Lasia and Kalinowski (6) expression for 1 = 2. Polarographic waves associated with M+(R.-)z formation were simulated using Equation 6 by varying COR.- = 0.01 to 0.99 CR and [R]' = 0.99 to 0.01 CR, assuming Pz = antilog 5.49 and ( E ~ / Z = )0 ~volt (Figure ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
1775
The corresponding formation constant is given by
Half-wave potential shifts were obtained from 20r
0
ID
o
0
?
2
-2 5
-1 5
1
log ICH. /mole d m - ' l
Figure 3. Calculated half-wave potential shift for different concentration of metal ion assuming log 6," = 0.43 and log p2 = 5.49, together with experimentally obtained data
2 ) . Waves obtained in this manner became steeper with increasing metal ion concentration. The actual Nernst slopes of E&. vs. log [ i / ( i d - i ) ] lines were always below 59 m v , the theoretical value for a one-electron reduction. For strong interaction, the polarographic wave equation is given by Ed,,.
= E112
RT + RT -In 2 - -In k -t 2F 2F .
RT i d - i RT RT - In - In CM++ -In p2 (8) F 2F 2F where k = kR.- = kR = 706 nD1I2 m2J3t1/6,Le. k is the coefficient from the Ilkovic equation assuming the diffusion coefficients of species R and Re- are equal. At half-wave potential, i = i d / 2 , and i d = kCR. Substituting these relationships into Equation 8, one again obtains Equation 7. (B) We now postulate, in addition to our previous assumptions, that the metal ion reversibly associates with the unreacted organic electroactive species.
-+
M+
+ R + M+R
(9)
From Equation 9, the formation constant of complex M+R is
(Marking the formation constant with an asterisk helps to discriminate complexes involving reactants from those involving products of electrode reactions.) The interface concentration of Y can be calculated from an expression
Under the conditions assumed in (B), the half-wave potential shift is given by
RT C O R . - - 2[XI0 F CRO - [Y]" The theoretical dependence of the half-wave potential on the metal ion concentration derived from Equation 12 with regard to Equations 5 and 11 is shown in Figure 3, for PI* and p2 set equal to antilog 0.48 and 5.49, respectively, and CRO fixed to 1 mmol dm-3, together with experimentally determined points. The experimental points obviously fit quite well to the theoretical curve. When the same values for and 62 were used for simulation of polarographic waves, the Nernst slope of Ed,,. vs. log [ i / ( i d - i)] lines were again less than theoretical value of 59 mV for the one-electron reduction. (C) In this case, the metal ion is assumed to react with the unchanged depolarizer in a ratio of 1:2 instead of 1:l as taken in (B), Le., the following equilibrium is assumed s 1 / 2
=
-- In
M+ -I-2R * M'R2 1776
RT COR.- - 2[XI0 F COR - 2[zl0 As is the preceding alternatives, the simulated waves are steeper than theoretically required for one-electron reduction and steepness increases with increasing metal ion concentration. In summary, polarographic experiments conducted under the conditions described above gave half-wave potential shifts (respective to E112 obtained in the absence of Li+ ion) compatible with the association of Li+ with HOAQ, to form either type M+R (PI* = antilog 0.48) or M+R2 (p2* = antilog 3.48) complexes in addition to formation of the type M+(R.-)2 complex with the radical anion HOAQe- ( 0 2 = antilog 5.49). On the contrary, experimentally determined slopes of E d e vs. log [ i / ( i d - i ) ]lines were inconsistent with those calculated from simulated polarographic waves. At any Li+ concentration, the experimentally determined slope was from 57 to 59 mV, while that calculated from simulated polarographic waves was more less than 59 mV under any of the assumed conditions. The apparently reversible experimental waves are thus a t variance with the theoretically derived rules and that with respect to half-wave potential shift, as well as with respect to Nernst slopes. These inconsistencies may be attributed to sluggishness of the systems observed in establishing interface equilibria, primarily those involving radical ion complexes. The considerable stability of Li+(HOAQ.-)2 (log p2 = 5.49 ( 1 1 ) )seems to indicate that the formation of the complex is governed not only by electrostatic forces, but rather by a combination of these forces and chelate bonds formed between Li+ ion and two oxygens participating in the HOAQ.- structure. Electrochemical behavior of the HOAQ/Li+ system, however, suggests that not only do free organic depolarizer molecules undergo an electrode reaction, but also so do the complex species formed from organic molecules and metal ions present in the medium. This aspect was hitherto neglected and does not appear in the literature. I t may, however, become important when possession of electrodonor groups enables nucleophilic interaction of organic molecules with the metal cation, whereby the energy levels of organic molecules are altered. By virtue of its electron-acceptor action, the metal ion bound a t the coordinationforming side of the nucleus, therefore, facilitates the electroreduction of organic molecules. Cyclic Voltammetry of 1-Hydroxy-9,lO-anthraquinone in the Presence of Li+. The strength of the Li+ interaction with HOAQ leading to formation of complexes of considerable stability suggests substantial participation of such complexes in the electrochemical reduction imposed upon HOAQ. DC polarography, however, was rather poorly adapted as a method for investigating this problem. Polarographic reduction of HOAQ in the presence of Li+ appeared as a reversible process. A clear experimental support for participation of HOAQ complexes in electrochemical reduction of HOAQ had to be sought by a different method and was found by means of cyclic voltammetry. Cyclic voltammograms recorded with several concentrations of Li+, are presented in Figure 4. The shift of the cathodic peak in a positive direction a t constant metal ion concentration, decreases as the scanning rate increases as shown in Figure 5 together with the experimental polarosE1/2
0
(13)
ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
=
-- In
AE/mV 50
I
Pol.
40 30 20 10
-
1
l o g IC,+/ mole dm”1
Figure 5. Experimentally obtained dc polarographic half-wave potential shift and cyclic voltammetric cathodic peak potential shift for different concentration of Li+
Cyclic voltammograms of 1 mmol dm-3 HOAQ with different concentrations of Li’ at scan rate 440 mV sec-‘ Figure 4.
Concentrations of Li+ from 1 to 9: 0.0, 0.8, 2.2, 3.9, 7.6, 15.1, 35.2, 57.4, and 98.5 mmol dm-3
graphic data. This shift might be due to delayed equilibration of the species during the complex formation with the radical anion, or may be caused by irreversibility of the electrode process. Existence of the irreversible process is obvious from the shape of the anodic part of the voltammogram. The anodic part of the voltammogram contains two peaks, indicating the occurrence of two different electrode oxidation processes. The more negative peak must be associated with the oxidation of free HOAQm-, because its height decreases with increasing metal ion concentration. The other peak is related to oxidation of another molecular species. Since its height increases with increasing Li+ concentration, this peak obviously corresponds to oxidation of Li+(HOAQ.-)2. The shape of this peak indicates irreversibility. At very low scanning rates, the two anodic peaks merge, thereby showing that dissociation of Li+(HOAQ--)2 takes place a t a relatively slow rate. Reduction of the HOAQ complex must be shifted to a more negative potential by increasing scanning rate because of its irreversibility, but it is accompanied with the reversible reduction peak of free HOAQ. From the investigation presented in this paper, it can be seen that the electrode reduction of HOAQ to HOAQS- in the presence of Li+, in spite of its apparent dc polarographic reversibility, is actually a composite process, in which free HOAQ takes part, as well as the complex species formed between Li+ and HOAQ. For this reason, it must be stressed that the results obtained only by dc polarography in the study of similar problems must be considered with care. A general conclusion seems to follow from the above discussion. Any study of metal ion influence on the electrochemical properties of the aromatic electroactive species, undertaken with the purpose of establishing a quantitative relationship between the electrochemical magnitudes and the interaction of potentially complex forming species, requires consideration of several factors in turn. 1) The association of the metal ion with the electrode reaction product. This factor shifts the reduction potential in the positive direction. I t is important to analyze the kinetic aspect of this interaction, especially with respect to Lhe speed of reaching equilibrium.
(1) Scan rate, 88 mV sec-’; (2) scan rate, 440 mV sec-’
2) The association of the metal ion with the unreacted organic molecules by forming a covalent bond. This factor is especially important with aromatics carrying electrondonor groups or atoms and causes a shift of the reduction potential in the negative direction. 3) The direct participation in electrochemical reaction of the complex species formed by association of organic molecules with metal ion. This process may be of considerable importance, together with the free electroactive species reduction, especially when more than one electroactive group resides on the organic molecule. The complex species will be reduced a t a more positive potential than that of the free depolarizer. This factor is likely to be manifested as a change in dc polarographic reversibility. Surely, all these factors cannot be adequately studied by using only electrochemical techniques. The necessity of taking recourse to other physicochemical methods, panicularly to methods based on spectrophotometric and especially simultaneous spectroelectrochemical techniques, is quite obvious. In our own work (II), the optically transparent thin-layer cell proved advantageous in providing a satisfactory explanation of results observed, but not explained, by other techniques.
LITERATURE CITED (1) M. E. Peover and J. D. Davies, J. flectroanal. Chem., 6, 46 (1963). (2) T. Fujinaga. K. Izutsu, and T. Nomura, J. flectroanal. Chem.. 29, 203 (197 1). (3) L. Holleck and D. Becher. J. Electroanal. Chem., 4, 321 (1962). (4) M. K. Kalinowski, Chem. Phys. Lett., 7 , 55 (1970). (5) M. K. Kalinowski, Chem. Phys. Lett., 8, 378 (1971). (6)A. Lasia and M. K. Kalinowski. J. flectroanal. Chem., 36, 51 1 (1972). (7) T. M. Krygowski. M. Lipsztajn, and Z. Galus. J. flectroanal. Chem., 42, 261 (1972). (8) M. Lipsztajn, T. M. Krygowski, and 2. Galus, J. flectroanal. Chem., 49, 17 (1974). (9) A. Lasia, J. flectroanal. Chem., 42, 253 (1973). (10) E. Casassas and L. Eek, J. Chim. Phys., 64, 971 (1967). (11) I. Piljac, M. TkalEec, and B. Grabaric, Anal. Chem., 47, 1369 (1975). (12) M. Fujihira and S. Hayano. Bull. Chem. SOC.Jpn, 4, 644 (1972). (13) A. Carlson and D. M. Hercules, Anal. Chem., 45, 1974 (1973). (14) I. Piljac, B. Grabaric. and I. Filipovic. J. flectroanal. Chem., 42, 441 (1973).
RECEIVEDfor review January 14, 1975. Accepted May 22, 1975. The authors gratefully acknowledge the financial support provided by US. National Bureau of Standards Grant NBS(G)-154 Project 8709357 under which sponsorship this work was performed.
ANALYTICAL CHEMISTRY, VOL. 47, NO. 11, SEPTEMBER 1975
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