pound on the electrode surface. Benzyl mercaptan is adsorbed strongly, as shown by the distorted current-time curves of individual mercury drops obtained in the presence of this substance. The adsorbed film could reduce the surface area of the electrode to a very small value and thus inhibit electroreduction of In(III), as has been observed for many other strongly adsorbed organic molecules (18). A number of carboxylic and hydroxylic ligands are inactive, including acetate, tartrate, and acetylacetonate. Oxalate shows slight activity, giving a current of 0.08 pa. Glycine, ethylenediamine, and a number of 0-amino alcohols also fail to catalyze reduction of In(II1). I n comparison, aromatic diamines are more active than the aliphatic analogs. Thus o-phenylene diamine shows moderate activity, although aniline, o-aminobenzoic acid, and o-aminophenol are inactive. Ligands bearing heterocyclic nitrogens capable of chelation were among the
most efficient catalysts studied and one of these, 2,2’-bipyridine exceeded the catalytic power of any other ligand by fivefold. LITERATURE CITED
(1) Britton, H. T. S., Robinson, R. A,, Trans. Faraday Soc. 28, 531 (1932). (2) Connick, R. E., Coppel, C. P., J . Am. Chem. SOC.81,6369 (1959). (3) Delahay P., “New Instrumental
Methods ’in Electrochemistry,” pp. 87-114. Interscience. New York. 1954. (4) De &faeyer, L., Kustin, K., Ann. Rev. Phys. Chem. 14, 12-20 (1963). (5) Hattox, E. &I.,De Vries, T., J . Am. Chem. Soc. 58, 2126 (1936). (6) Hepler, L. G., Hugus, Z. Z., Ibid., 74, 6115 (1952). ( 7 ) Kolthoff, I. &I., Lingane, J. J., “Polarography,” 2nd ed., Vol. 11, pp. -51 9-20. _ _ . Interscience. New York. 1952. (8) Lingane, J. J., J . Am. Chem.’Soc. 61, 2099 (1939). (9) RlairanovskiI, S. 0.) J . Electroanal. Chem. 6 , 77 (1963). (10) ;\lark. H. B. Jr., ANAL. CHEM.36, Ok0 (1964).
(11) Mark, H. B., Jr., J . Electroanal. Chem. 7,276 (1964). (12) Mark, H. B., Jr., Reilley, C. N.,
ANAL.CHEM.35,195 (1963). (13) Mark, H. B., Jr., Reilley, C. N., J . Ekctroanal. Chem. 4, 189 (1962). (14) Mark, H. B., Jr., Schwarts, H. G., Jr., Ibdd., 6, 443 (1963). (15) Meites, L., “Polarographic Techniques,” pp. 78-82, Interscience, New York, 1955. (16) Moorhead, E. D., MacNevin, W. M., ANAL.CHEM.34,269 (1962). (17) Nelson, I. U., Iwamoto, R. T., J. Electroanal. Chem. 6 , 234 (1963). (18) Reilley, C. N., Stumm, W., “Progress in Polarography, Vol. I, Chap. V, P. Zuman and I. M. Kolthoff, eds., Interscience, New York, 1962. (19) Schufle, J. A., Stubbs, M. E., Whitman, K. E., J . Am. Chem. SOC.73, 1013 (1951). RECEIVED for review September 16, 1964. Accepted November 30, 1964. Research supported in part by National Science Foundation through Grant GP 1996. A. James Engel was sponsored by National Science Foundation and participated in the 1964 Summer Program in Instrumental Analysis held at Rensselaer Polytechnic Institute.
Electrooxidation of Tetraphenylborate Ion at the Pyrolytic Graphite Electrode W. RICHARD TURNER and PHILIP J. ELVING The University o f Michigan, Ann Arbor, Mich. Tetraphenylborate ion is oxidized at the stationary pyrolytic graphite electrode in aqueous solution to produce two voltammetric waves. The first wave, which occurs as a welldefined peak (E,,z = 0.21 6 volt VS. saturated sodium chloride-calomel electrode) results from a 2-electron, pH-independent process which produces diphenylborinic acid and biphenyl:
B(CE“h--f B(CsH6)2+
B(C&)z+
+
(C6H&
+
2e-
+ H2O -+ B(CGH&OH + H+
The second wave, which is less well defined, represents a 2-electron process, which involves the oxidation of diphenylborinic acid:
+
B(C6HsLOH 3- 2H20 4 NOHIS (CsH& f 2H+ -k 2eThe latter process is linearly pH0.057 dependent: E,lz = 0.92 pH. The slope of -0.057 indicates that one hydrogen ion is produced for each electron liberated in the oxidation.
-
S
its discovery by Wittig (IO), sodium tetraphenylborate (NaTPB) has become a highly useful analytical reagent for the determination INCE
of potassium, certain heavy metal ions, and many basic nitrogen compounds; in fact, its ability to precipitate large cations seems to be very general. While the chemistry of the tetraphenylborate ion in aqueous solution is fairly well understood, a systematic investiga tion of its electrooxidation in water has until now not been reported. As a part of a study of organic oxidation reactions a t the wax-impregnated graphite electrode, Elving and Smith (3) observed two waves for the electrooxidation of the tetraphenylborate ion in aqueous media. This observation led to the development of a method for the direct titration of potassium using amperometric equivalence-point detection ( 7 ) . These results would seem to be somewhat contradictory to those of Tho reported being unable to Geske (4, observe the oxidation of T P B a t the rotating platinum anode in aqueous solution because of the interfering evolution of oxygen. However, Geske was successful in carrying out the oxidation in anhydrous acetonitrile, where a single wave was produced. I n one instance, there was an indication of a second wave; however, film formation prevented further examination of this wave. On the basis of his results, Geske proposed the following mechanism for
the electrooxidation of tetraphenylborate ion in acetonitrile:
Subsequently, Geske reported (5) that mass spectrographic analysis of the products of the electrolysis of a mixture of T P B and perdeutero-TPB in acetonitrile revealed the presence of a mixture of biphenyls containing either all hydrogen or all deuterium, thus indicating that both phenyl groups in each biphenyl had come from the same tetraphenylborate ion. I n the present investigation, the voltammetric and coulometric behavior of the tetraphenylborate ion was examined at the pyrolytic graphite electrode in both aqueous and nonaqueous media. EXPERIMENTAL
Sodium tetraphenylborate reagent (J. T. Baker; Fisher Scientific) was recrystallized from chloroform during the earlier phases of the present study; this was abandoned when the commercial material seemed to be of sufficient purity. Chemicals and Reagents.
VOL 37, NO. 2, FEBRUARY 1965
a
207
Table 1.
PH 2.1 3.1 4.6 5.9
Variation of Half-Peak Potential and Peak Current with pH for First Oxidation W a v e of Tetraphenylborate. E p / P , V.
0.218 0.214 0.221 0.214 0.201 0,206 0.233 0.210 0.206 0.224 0.222 0.218
i,, pa. 7.6 7.7 7.0 7.2 7.2 7.6 7.0 6.4 5.9 7.0 6.8 7.0
PH 6.9 7.7 10.0
11.7
E p / P , V.
ip, pa.
0.213 0.207 0.214 0.220 0.210 0.216 0.223 0.225 0.213 0.221 0.226
7.8 7.8 7.9 7.2 7.8 7.6 7.4 7.6 7.5 7.7
8.0
Av. 0.216 7.3 Std. dev. 0.007 0.3b a Solution. l.O& NaTPB; buffersused described in text. Electrode: 0.126sq. cm, pyrolytic graphite surface. Temperature: 25". Polarization rate: 1.66 mv./sec. b Value based on excluding three worst figures of 6.4 and 5.9 pa. at pH 4.6 and 8.0 Ma. at pH 6.9.
Diphenylborinic acid was prepared by hydrolyzing its ethanolamine complex (Aldrich) by heating with a 100% excess of 1M sodium hydroxide solution on a steam bath. The clear solution was cooled and neutralized to Congo Red paper with 164 hydrochloric acid. The resulting oil was extracted with chloroform and dried; the solvent was removed by evaporation with a stream of air, The residual yellow oil was placed in a side-arm flask, which was evacuated for about an hour. During this time, the oil solidified, producing a white, semicrystalline material, which was recrystallized from petroleum ether t o yield a crystalline material (m.p., 116' C.), Although this value is in agreement with the literature value (1) for the anhydride of diphenylborinic acid, the infrared spectrum of the . crystalline material is characteristic of the acid itself. Anhydrous lithium perchlorate (G. F. Smith) was dried a t 200' C. for a t least 2 hours. Acetonitrile (Eastman technical) and dimethylformamide (Eastman White Label) were dried over phosphorus pentoxide and fractionally distilled. Pyridine (Merck reagent) , dried over barium oxide, was also fractionally distilled. All other chemicals were of reagent or C.P. grade and were used with no further purification. Graphite Electrodes. All electrodes were made from pyrolytic graphite (General Electric). Electrodes for voltammetry were machined from blocks to form cylinders, 4 mm. in diameter and between 1.0 and 1.5 em. in length with the ab plane forming the base of the cylinder. One of these was cemented (either Epon Resin 815 or melted polyethylene was used to form a sticky coating around the cylinder) into the end of a 5-inch length of 6-mm. glass tube so that the base of the cylinder was flush with the end of the tube. Between runs, the electrode was resurfaced by polishing the end with No. 600 silicon carbide paper mounted
208
0
ANALYTICAL CHEMISTRY
on a rotating disk sander. Electrical contact was made by placing a small amount of mercury in the tube and inserting a piece of copper wire. For macro scale electrolysis and coulometry, a large-area pyrolytic graphite electrode was constructed from 14 graphite strips (8.5 X 1.5 cm.; 2 to 3 mm. thick), which were separated by thin graphite spacers a t one end; the electrode was held together with a battery clip attached to this end. Although the electrode occupies a relatively small volume, its total surface is over 300 sq. cm. Apparatus. Voltammograms were recorded with a Sargent Model XV polarograph. The water-jacketed Hcell was maintained a t 25.0' =t 0.1' C. The reference electrode was a saturated sodium chloride-calomel electrode, which was separated from the test solution by a sintered glass disk and a n agar plug saturated with sodium chloride. The potential of this electrode has been measured as 0.2458 volt at 20' C. (6); the potential of the saturated potassium chloride-calome1 electrode is 0.2477 volt a t 20' C. For studies in nonaqueous media, a 0.21W silver perchlorate-silver electrode was used, whose potential would naturally vary from solvent to solvent; no attempt was made to assign any potential to this electrode. I n nonaqueous media, iR drop compensation was applied using the circuit of Annino and Hagler ( 2 ) . Cyclic voltammetry was carried out using operational amplifier circuitry similar to that of Underkofler and Shain (9) and a Moseley Autograf X-Y recorder. The low frequency triangular wave used for cyclic voltammetry was generated with a Hewlett-Packard Model 202A function generator. Constant potential coulometry was carried out using a Fisher controlled potential electroanalyzer and later an operational amplifier potentiostat. The current during the electrolysis was integrated by feeding the potential across a precision 1-ohm resistor into a
Dymec Model 2210 voltage-to-frequency converter and counting the resulting frequency with a HewlettPackard Model 521AR counter. When the electrode was rotated, a Holtzer-Cabot 300-r.p.m. synchronous motor was used. The rate of rotation was varied by means of belt-driven pulleys. DISCUSSION OF RESULTS
Voltammetry of Tetraphenylborate. The voltammetric behavior of sodium tetraphenylborate was investigated over a wide range of pH, concentration, electrode rotation speed, and media (aqueous and nonaqueous). I n all cases, a t least one, and in most cases a t least two waves were observed. Table I summarizes the half-peak potentials (Ep,*) and peak currents (i,) for the first wave obtained a t the stationary pyrolytic graphite electrode for pH 2.1 t o 11.7 buffered solutions, which were I d in NaTPB and 0.5M in total ionic strength. Both EPiz and i, are essentially independent of pH, with average values of 0.216 volt and 7.2 pa., respectively. At all pH values, a second wave appeared a t about 0.7 volt. I n some cases, a third wave appeared a t about 0.5 volt as a prewave to the second wave; this third wave was always much less than 1 pa. Since these data were obtained a t a stationary electrode where the waves appear as peaks rather than plateaus, it is difficult to evaluate the voltammetric data obtained for waves appearing subsequent to the first wave. Therefore, it was deemed inadvisable to present the data obtained for the second wave as having any quantitative significance. I n general, however, the second wave shifted to less positive potential with increasing pH. I n 0.5M acetate buffer (pH 4.6), the peak current-concentration relationship was linear for the first wave in the range of 0.05 to 20 m;M tetraphenylborate. The second wave was larger than the first wave below 0.5 mM, but above 8 mM it could not be detected. Another anomalous effect was noted when the electrode was rotated. At speeds below 75 r.p.m. two waves were observed; above this rate of rotation, the second wave disappeared. Moreover, after reaching its full height, the first wave diminished with increasing potential. .4t higher speeds of rotation, the effect was more pronounced. Such behavior can be attributed to film formation, since a t higher speeds of rotation more material is transported and therefore oxidized a t the electrode. A plot of the peak current for the first wave against the square root of rotation speed was linear, indicating that convective mass transport was the controlling process, rather than a kinetic
b Figure 1. Voltammograms recorded potential electrolysis
during
%
controlled
I-
z
W LL
Diphenylboronic acid after passage in oxidation a t 1.1 volts (arrow) of 0 coulomb 9.65 coulombs 19.3 coulombs In oxidation at 6. TetraDhenvlborate after *Dassoge . 0.5 voli ( 1 o’nd 2) and at 1 . 1 volts ( 3 and 4)of 1 . 0 coulomb 2. 9.65 coulombs 3. 19.3 coulombs 4. 2 8 . 9 5 coulombs
A.
1. 2. 3.
3 0
z W a
process. The disappearance of the second wave can be attributed to the removal of the oxidation products by convection before they have an opportunity to react and be further oxidized. This would confirm Geske’s mechanism which has the diphenylboronium ion being formed and then reacting with trace amounts of water to form diphenylborinic acid. The formation of this acid directly in the electrode process according to the reaction,
+ H + + 2e-
(C6H5)2
(3)
is precluded by the fact that the first wave is pH-independent. Consequently, it seems probable that the reaction of diphenylboronium ion with water is sufficiently slow that this ion is swept away from the electrode surface before it has reacted, resulting in diphenylborinic acid’s not being available at the electrode for further oxidation. Voltammetry of Diphenylborinic Acid. T o see if the second wave observed in the oxidation of tetraphenylborate was due to the oxidation of diphenylborinic acid, the voltammetry of this compound was investigated over the p H range of 1 to 10 (Table 11). Comparison of the voltammograms of the two compounds (Figure 1) leaves little doubt that this is the case. The half-peak potential of diphenylborinic acid occurs a t a potential in the same general region (0.6 to 0.7 volt) as the second wave for the oxidation of tetraphenylborate. The half-peak potentials are shifted to less positive values with increasing pH. The plot of Eplz us. p H is linear:
E,,z
=
0.93 - 0.057 p H
B(C6Hb)zOH
+
II: 2 V
+
2Hz0 -,B(OH), (C&h -I- 2H+ f 2e- ( 5 )
When McIlvaine phosphate-citrate buffers were used, two waves were pro-
4
2
3
4
lI
l
0
0 2
0 4
06
duced in some instances. The peak currents for most McIlvaine-buffered solutions were considerably less than expected. I t is possible that the acid reacts with one of the buffer components. For example, although two waves were observed in p H 2.9, 4.0, and 4.8 XcIlvaine buffers, only single waves were obtained in p H 3.7 and 4.6 acetate buffers. The prewave frequently observed in the voltammetry of tetraphenylborate may be due to a similar reaction with the buffer, since McIlvaine buffers were used throughout the p H range of 3 to 8. Voltammetry in Nonaqueous Solvents. Voltammograms were obtained for 1 m J l solutions of sodium tetraphenylborate in anhydrous acetonitrile, dimethylformamide, and pyridine, using 0.5M lithium perchlorate as supporting electrolyte in each case. Two waves were recorded in acetonitrile; Epiz values are 0.32 and
PH 1 0 1 9 2 0 2.9
3.7 4.0 4.6
l
I
IO
08
POTENTIAL, V
Table II.
I
lo
2 -2
(4)
The variation with pH is indicative of a process where one hydrogen ion is produced per electron transferred in the oxidation :
:l&&++q
s 1
I-
l
1
1
12
14
SCE
1.04 volts. 4 well-formed wave (Ep,:! = 0.23 volt) was obtained in dimethylformamide; a second wave, which was obscured by the background discharge a t 0.9 volt, appears to have an Epiz of about 0.8 volt. The oxidation in pyridine was obscured by the presence of waves due to the solvent, which began to appear a t about 0.6 volt; an investigation is in progress to determine the nature of these waves and means of their elimination ( 8 ) . Coulometry and Macroscale Electrolysis. I n a n effort further to elucidate the mechanism of oxidation of tetraphenylborate a t the pyrolytic graphite electrode in aqueous solution, constant potential coulometry was carried out. Solutions, which contained 0.1 mmole of NaTPB in 0.5.11 acetate buffer (pH 4.6), were exhaustively electrolyzed at 0.35 and at 0.85 volt to determine n, the number of electrons transferred in the oxidation
Variation of Half-Peak Potential and Peak Current with pH for Oxidation of Diphenylborinic Acida
EPIZ,
EPi2,
V.
a,, Ma.
PH
V.
0 89 0 84 0 82 0.81 0 76b 0.71
5 1 4 8
4 8
0 69* 0 6Sb 0 58 0 59 0.57 0.56 0 54 0.54 0.56 0.54 0.40
0.70b
3 8 4.2 5.p 3.8 6.5c
0.68 0.67 0 65
5.2 4 0 5.9
5 5 6.0 6.4 6 9 7.0 7.1 7.9 9.9
I,,
pa.
1 4e 1 oc 3 5 1 9c 1 5c 2 2c 1 5c
2.8 2.2c 1.P 2 6
Solution. 1.0mM diphenylborinic acid: buffers used described in text. Electrode: Temperature: 25” C. Polarization rate: 1.66 mv. /sec. Data recorded are for first of these waves. b Two waves observed a t this pH. c McIlvaine buffer mixtures were used, which may account for some of these low i, values. a
0.126 sq. cm. pyrolytic graphite surface.
VOL. 37, NO. 2, FEBRUARY 1 9 6 5
209
a
j
1
5
7
I
9
II
13
WAVELENGTH , p
Figure 2. 0.
b.
Infrared spectrograms
Electrode fllm from oxidation of tetraphenylborate a t 0.5 volt Sample of diphenylborinic acid
I
1
7
3
'9
II
WAVELENGTH , p
Figure 3. 0.
b.
Infrared spectrograms
Residue from N a O H extroction of electrode film Sample of biphenyl
WAVELENGTH ,p,
Figure 4.
Infrared spectrograms
Electrode fllm heated a t 110' for 1 hour b. Sample of diphenylborlnlc acid treated in same w a y Sample of commercially available "phenylborlc ocid" C. o.
210
0
ANALYTICAL CHEMISTRY
13
of 1 molecule of electroactive species, for each of the two waves. It was frequently difficult to carry the electrolysis a t 0.85 volt to completion because of the formation of film on the electrode surface; it was necessary periodically to interrupt the electrolysis and to clean the electrode by dipping it in acetone. Even with this procedure, the electrolysis required 5 to 6 hours for completion and the results were invariably low. The n values in four controlled potential coulometric oxidation experiments a t 0.35 volt were 2.25, 2.17, 2.24, and 2.07; on then increasing the electrolysis potential to 0.85 volt, further n values of 2.14, 1.73, 1.89, and 1.97, respectively, were measured. The low values in the last three runs may possibly be due to film formation. The results of the coulometry indicate that each wave is due to a 2-electron oxidation, which is consistent with the proposed mechanisms. To confirm further that these are 2-electron processes, voltammograms were recorded during the electrolysis after nF/c coulombs were passed through the solution a t a constant potential. I n this case, c was 0.1 mmole, so that 9.65~1 coulombs were necessary to carry out the oxidation. By recording the voltammograms after multiples of 9.65 coulombs had passed through the solution, it was not necessary to electrolyze the solution exhaustively. One need only note the reduction in the wave height after 9.65 coulombs were passed-e.g., the wave will be diminished by half if the process is a 2electron oxidation. I n the case of the tetraphenylborate oxidation a t 0.5 volt, the first voltammetric wave was reduced to half of its original height after the passage of 9.65 coulombs and was completely obliterated after 19.3 coulombs; the voltammogram at the latter stage was identical to that of an equivalent amount of diphenylborinic acid (Figure 1, B). The potential was then increased to 1.1 volts and an additional 9.65 coulombs were passed, resulting in the reduction of the height of the second wave by half. Extensive film formation prevented further electrolysis; however, there is no doubt from these results that the second wave is due to a 2electron oxidation of diphenylborinic acid. I n a similar oxidation of diphenylborinic acid itself, the voltammogram recorded after the, passage of 19.3 coulombs (2 F / c ) indicated that virtually all of the diphenylborinic acid had been oxidized (Figure 1 , A ) . The film on the electrode surface was examined after the exhaustive electrolysis a t 0.5 volt of a solution of tetraphenylborate identical to the one described above. The film was dissolved in acetone and the acetone evaporated,
40
“phenylboric acid.” The spectra are 77compared in Figure 4.
0
0 2
04
06
08
10
12
I4
POTENTIAL .V v i S C E
Figure 5. a.
Cyclic voltammograms
Tetraphenylborate
b. Diphenylborinic acid Solution composition. 1.OmM NaTPB or diphenylborinic acid, 0.5M acetate buffer (pH 4.6) Scan rate. 26 mv./sec. Electrode orea. 0.1 26 sq. cm.
leaving a semisolid residue; the infrared spectrum of a portion of this residue, dissolved in carbon tetrachloride, was almost identical to one of diphenylborinic acid (Figure 2). When another portion of the residue was extracted with sodium hydroxide solution to remove the diphenylborinic acid, there remained a material whose infrared spectrum was identical to that of biphenyl (Figure 3). When a third portion of the residue was heated at 120” C. for an hour, the material was converted to the same compound as that formed when diphenylborinic acid is heated. This is probably phenylboronic anhydride, since its infrared spectrum is identical to that of commercially available (Aldrich Chemical)
Cyclic Voltammetry. A number of cyclic voltammograms were recorded a t several voltage scan rates to see if a reduction process could be detected on the reverse sweep. However, as can be seen in Figure 5 , no cathodic current was produced on the reverSe sweep during the voltammetry of sodium tetraphenylborate or diphenylborinic acid, indicating that these oxidations are highly irreversible. MECHANISM OF OXIDATION
It seems fairly certain that the oxidation of tetraphenylborate ion in aqueous solution a t a graphite anode occurs in three discrete steps. First, the tetraphenylborate ion is oxidized electrochemically by an irreversible, 2-electron, pH-independent process to produce the diphenylboronium ion and biphenyl. Then, in a chemical reaction. the diphenylboroniuni ion reacts with water to produce diphenylborinic acid and hydrogen ion. Finally, a t a higher potential, the diphenylborinic acid is electrooxidized by an irreversible, 2electron, pHdependent proces? to yield biphenyl, boric acid, and hydrogen ions. The formation of biphenyl is probably a concerted reaction, whereby a pair of electrons are transferred to the electrode and a pair of electrons simultaneously move in to form a bond between the two phenyl rings. This seems more
probable than free radical formation, in the light of Geske’s studies with perdeuterotetraphenylborate. ACKNOWLEDGMENT
The authors thank the U.S.Atomic Energy Commission and the Horace H. Rackham School of Graduate Studies of The University of Michigan, which helped support the work described. One author (WRT) thanks the Institute of Science and Technology of The t-niversity of Michigan for a postdoctorate fellowship. LITERATURE CITED
(1) Abel, E. W., Dandegaonker, S. H., Gerrard, W.,Lappert, 52. F., J . Chem. Sac. 1956. 4697. ( 2 ) Annirb,’ R.-A., Hagler, K. J., ANAL. C H E M . 35, 1555 (1963). ( 3 ) Elving, P. J., Smith, D. L., Ibzd., 32, 154‘5 (1960). (4) Geske, I). H., J . Phys. Chem. 63, 1962 i19593. (5) Zbid:. 6 6 . 1743 (1962). ( 6 ) Landhla;, A. ~ D . d., Page, J. E., S a t w e 151, 84 (1943). ( 7 ) Smith, I). L., Jamieson, I). It., Elving, P. J., ANAL.CHEM.32, 1253 (1‘360).
( 8 ) Turner, W. R., Greinke,, R.,, Elvine. P. J., work in progress. ( 9 ) Underkofler, W. L., Shain, I., ANAL. CHEM.35,1778 (1963). (10) Wittig, G., Keicher, G., Riickert, A , , Itaff, P., Ann. 563, 110 (1949).
I ,
RECEIVEDfor review October 5, 1964. Accepted November 27, 1964.
Polarographic Reduction of Pyridinium Ion in Pyridine Application to the Determination of Bronsted and Lewis Acids MICHAEL S. SPRITZER, JOSE M. COSTA, and PHILIP J. ELVING The University of Michigan, Ann Arbor, Mich.
b Pyridinium ion gives a l e diffusioncontrolled reduction wave at the D.M.E. in pyridine solution (lithium perchlorate as background electrolyte), which is linearly proportional to concentration. The reduction involves attack on the pyridine ring, as contrasted to the nature of pyridinium ion reduction in aqueous solution. The l e nature of the wave was established coulometrically, which indicates that the pyridinium ion can b e also determined by direct coulometry. A small prewave, whose magnitude depends on the particular pyridine sample used, is included with the main pyridinium ion wave for analytical purposes. Essentially the same halfwave potential and diffusion current constant are given by a pyridinium
salt, an alkylpyridinium salt, a Bronsted acid of aqueous pK, less than 9, or a Lewis acid such as an alkyl halide, which forms a quaternary salt with pyridine. Polyprotic acids give a multiple diffusion current constant. This behavior allows the determination of a large variety of inorganic and organic acids, as well as the determination of the total acidity of a sample.
U
OF NONAQUEOUS media for polarographic investigations, particularly of organic compounds, has, in recent years, increased considerably; the advantages of such media, both theoretical and practical, have become fairly well known. Only recently has serious attention been given to the use of SE
pyridine as a solvent for polarography. Pyridine is apparently the first aromatic-type solvent to be systematically considered for polarography. Investigations have included studies of the behavior of inorganic salts ( S ) , of benzophenone ( 4 ) , and of the electrochemical reductive fission of the pyridine ring in solutions containing a Lewis acid (aluminum chloride) ( 2 ) . The electrochemical reduction or oxidation of organic compounds frequently involves the addition or elimination of hydrogen ions, which, in pyridine medium, form pyridiniuni ions; consequently, it was necessary to become thoroughly familiar with the electrochemical behavior of the pyridiniuni ion itself. Accordingly, the behavior of several proton-donors, a pyridinium VOL. 37, N O . 2, FEBRUARY 1965
21 1
