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 ( p H 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 t h e 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
salt and some alkxl pyridiniuni salts) has been investigated to establish the polarographic behavior and possible reduction niechanism of the 1)yridiniurii ion in pyridine solution. The electrocheniical reduction of Iiyritline (and h j - iiiililication 1iyricliniuiii ion) was apparently first ,studied syrt,ematically by Emniert ( 5 ); a thorough study is that by Orhiai and Kataoka ( 17 ) . The Iiolarographic reduction of 1)yridinium ion in aqueous solution was originally studied by Shikata and Tachi ( 1 8 ) , who observed two waves which were interpreted as due t o the reduction of pyridinium ion and of undissociated liyriciiniuni salt, respectively. Later ivork (12, 21) confiriiied the existence of these two waves, but postulated a catalytic hydrogen erolution as the source o f t h e firat vvave. 2 5
Hi-
"
"
l
~
u
- E ,valt s
Figure 1 . Electrocapillary curve for 0.1 M LiClOd solution in pyridine H+
Hz
2H"
(3)
and reduction of ~iyridiniumion as the cause of the second wave. The halfwa\-e potentials of both n-aves become more negative with increa5ing pH. The limiting current (wave not specified) is reported to be proportional to pyridiniurn concentration. llairanovskii and coworkers recently studied the electrocheniical behavior of pyridine ( 1 9 ) and the quaternarv pyridine (f4)with special empha on the effect of formation on the dropping mercury electrode of a surfaceactive polymeric product ( f 5 ) . Hydrogen evolution did not occur in solutions (presumably aqueous) I\-hich were more concentrated than 0.5.ll in pyridine; the current is then caused by the forma-
D.M.E. v s . mercury pool. Capillary: m = 1.68 mg./sec.; h = 4 0 cm. (uncorrected far back pressure)
tion of an organic product, rvhich polymerizes. Calculations of dinierization rate constants for some quaternary pyridine salts were reported ( 15 ) . Kuta and 1)ral)ek (11) 1ircviouslJ- observed that this parameter tlqienils on drop time. Recently, Zuman and \Tan-zonek (23) excellently summarized the polarographic reduction of the quaternary ammonium nitroi?;en-carhn system (solvents not q)ecified but presumably aqileous in most cases).
~
Table l. Background Electrolytes: Resistance and Decomposition Potentials of Saltz in Pyridine Cell resistConcn., ance, Conipd. M Kohni Edn volt LiCl 0 lb 50 -1.ilc 0.36 15 -1.5P 0.3 1. 0
LiClO,
KSCX
0 1 0.5 0.1 0.5
11
10 7 2 20 10
-1,5:3c
-1.4iC -1 -1 -1 -1
7iC
65c 65 53 -1 18 -1 57
N d 0 lb 4 NnB( CbHj)4 0.1 5 E d represents the decornposition potential i's. the rnercury anode-i.e., the potent,ial a t which the current reaches 1 *a. h Initial rnasiriiurii appeared (rf. test ). % , small prewnve was tr1)~erved (rf. text)).
212 *
ANALYTICAL CHEMISTRY
EXPERIMENTAL
Reagents. Although analytical reagent grade pyridine (,J. T . Baker or llallinckrotlt) m. satisfactory for use without, further 1) u r i fi ea t io n , it was dried with Lincie molecular sicves type 4.4 ( 3 ) . The particular pyridine sample used will-as subsequently discussed--determine the magnitutle of the prewaves encountered with the background solution and with the pyridinium Polution. LiCl (Baker Rr .4tlainson), LiC10, ( G . F. Smith, anhydrous), Sa1 (llallinckrodt), Sa13(C6Hb)4(lIallincC S tl3akrr) n-err dried , and q t o r d in a desiccator. Other reagent* were of analytical reagent grade and were used without further purification. Pl-ridiniuni nitrate was prel)ared by slowly adding nitric acid. with periodic agitation, t o Iiyridine. cooled in an ice bath: the solid Iiroduct, which precipitated, was filtered and then recrystallized several times from acetone.
-4rgori (99.S9yc pure), u e d for deoxygenation of solution-., \vas first dried arid then passed through dry pyridine to equilibrate it. Triple distilled niercwy wan u-ed. Apparatus. Polarogram-: were recorded rvith a Sargent l l o d e l XV I'olarogralih; a I m d s b S o r t h r u p used as a millivoltmeter t o m ~ a . u r t ~the potential of the mercur)- ~ioolreference electrode. LIarinr barometer tubing was used for the D.1l.E. c.apillary. 'l'he product, m? 3 f 1 6 , in 0.1.11 LiCIOI in pyridine a t olieri circuit \vas 1.749 ( t = 3.6 seconds; ut = 1.68 m g . second). jacketed H-cell ( I O ) was used with a mercury iiool a. reference electrode. The pottmtial of the latter was measured apainqt the normal silver electrode in 1)yidine (S.1~17) ( 3 ) . For iR drop c,orrection, the cell re-"t >is ance was measured with a conventional bridge. About midway through the study, the iR-drop corrector of .\nnino and Hagler (I) was constructed and was used throughout the remainder of the study. Coulonietry at controlled potential was carried out in a lleites cell ( 1 6 ) with the potential supplied by a HelvlettPackartl Model 712R power supply and controlled by a Philhrick PSA-3 operational amplifier and SK2-I3 booster aniplifier. -i Fiiher Elecdropode was used as the reference liotential source. Polarographic Procedure. Stock wlutioiis of reagents were prepared by dissolving weighed quantities and diluting with pyridine to k n o n n volume. Test solutions were generally prepared by pipetting appropriate quantities of the latter solutions into 10-ml. volumetric flasks arid diluting to volume with pyridine. The t e d solution was divided between the two compartments of the H-cell, and argon \vas bubbled through both solutions for 15 minutes through mediuiii porosity glass frits. lIercury to bc used as thc anode was then added to one compartment and the D.11.E. was placed in the other compartment. Polarograms were started at 0 volt with the potential at the 11.lI.E. being made more negative until the background discharge was reached. l'he current sensitivity was adjusted in each case to an approlxiate value. -in argon atmosphere was maintained over the solution throughout each run. Temperature of observation was 25" C. i- 0.2", escept where other temperatures are reported. Initially, the halfwave potentials were corrrcted for iR drop: later, this correction was made automatirally during the recording of each polarogram. Coulometric Procedure. Onehundred milliliters of 0.1X LiC10, in pyridine \vas added to the cathode coniliartment of the coulometric cell arid deaerated for 15 to 20 minutes; m e r c u r y was then added and stirred with a magnetic htirrer. The bridge comliartment was filled wtih background solution; a helis, 2 cm. long anti 4 mm. in diameter, of 20-gauge platinum wire served as the counter electrode and was placed in aqueous 0 . l X LiCIOl solution. The back-
I
0.3
I
0.6
I
0.9
I
I
I
1.2
1.8
1.4
- E 5 NAgE,volts
Figure 2. Effect of benzoic acid concentration on the prewave in pyridine solution A:
0.01mM; B: 0.03mM; C: 0.05mM; D: O.lmM; Background electrolyte: 0.1M LiCIod
ground solution was preelectrolyzed for 20 to 30 min., 1 or 2 ml. of a stock solution of elect,roactive species was added, and the system was deaerated for 10 minutes more. The electrolysis current was integrated with a conventional silver coulometer. RESULTS A N D DISCUSSION
Selection of Background Electrolyte, T h e possible utility of several compounds as background electrolyte in pyridine was first examined (Table I ) . An irreproducible initial maximum ( E , = -0.1 volt), varying between 0 and 1.2 pa., was occasionally observed in 0.1 t o 0.5N LiCl solutions, but not in 1.OM solutions. Lithium perchlorate a t 0.1.11 concentration was the most convenient background electrolyte for the objectives of the present study because of its large potential range and lotv resistance. The electrocapillary curve for this background electrolyte (Figure 1) shows a maximum at ca. -0.35 volt us. NAgE, which is close to the value found for O.ld1 LiCl in pyridine ( 3 ) . Prewaves. T h e magnitude of a prewave, which just precedes the dischmge for the background solution, varies with individual lots of pyridine, although its potential is apparently constant. The height has varied from
E:
0.2mM.
0.07 to 2.0 pa. When an acid is added to the background solution in ever increasing concentrations, a wave begins to appear of half-wave potential, E m , equal to approximately - 1.10 volts; simultaneously, the height of the background prewave decreases, with the sum of the two waves remaining constant.
When the wave at -1.10 volts has reached the height of the original background prewave, which in the case of the pyridine used in the present study corresponded to an acid concentration of 0.02mJf, further addition of acid does not affect the wave a t -1.10 volts, but causes the appearance of a wave of E,iz a t about -1.35 volts; the latt,er wave, which grows linearly with further acid addition, is designated in the present paper as the main Iiyridinium wave. Such a behavior pattern in which benzoic acid mas added to a sample of pyridine (Merck, lot No. 74751), for which the background prewave was 0.40 pa., is shown in Figure 2 . The direct addition of an alkylpyridinium salt or of an alkyl halide does not affect the background prewave or cause the appearance of the lirewave at -1.10 volts, but results in the direct appearance of the wave at - 1.35 volts. Polarography of Monobasic Acids. Pyridine solutions of each of several monobasic acids-acetic, benzoic, trifluoroacetic, and 2,4-dichlorophenolcontaining 0.1 31 lithium perchlorate as background electrolyte gave a well defined wave preceded by a small prewave (Table 11; Figures 3 and 4). X similar solution of pyridinium nitrate gave a n identical wave pattern. Since the values of and diffusion current constant, I , listed in Table I1 each represent the mean of a number of individual measurements (usually 10 or 12), the data discussed in this and s u b sequent sections are expressed as the mean plus-minus the standard deviation. The half-wave potentials are independent of concentration of electroactive species; the mean Eiip is -1.36 f 0.03 volts us. S h g E for the main wave, and -1.12 i 0.01 volts for the pre-
L 5 -4
-
3
-
2
-
I
-
0.0
I
0.2
0.6
0.4
-E
Vi!
0.8
LO
1.2
1.4
1.6
1.8
20
N A 9 E, V O L T S
Figure 3. Polarogram of a pyridine solution, which is l . O l m M in benzoic acid and 0.1M in lithium perchlorate VOL. 37,
NO. 2, FEBRUARY 1965
213
I
0.8pa
5
W
a a
3 0
I
E
F
D 4
,
-
1
0
0.4
0 .a
-E
I2
I .6
2 .o
vr NAgE,volta
Compound CFjCOOH C~HSCOOH CHsCOOH HpyrNOs
2,4-CI&eHaOH
rent variations in the prewave accompanying mercury height and temperature variation. However, the combined current of both waves shows a linear variation with the square root of drop time-Le., with the square root of the height or hydrostatic pressure of the mercury column (corrected for back pressure). Temperature coefficients for the combined current bteween 10' and 25' C. are between 1.0 and 1.8%/degree.
solution in pyridine did not show any polarographic waves in the range of 1 to 5mM phenol. The small magnitude of the prewave as compared to the main wave and the merging of the two waves, renders difficult the measurement of the small cur-
Polarographic Data for Several Acids and Pyridinium Salts in Pyridine Solution Containing 0.1 M LiC104
pKaa 0 231 4.20 4.75 5.22 7.85 9.89 2.97 13.00
Concn. No. of range, runs mM 12 0 072-4 33 12 0 139-13 9 12 0 314-15 7 0 575-11 5 15 10 0 528-2 11 No waves observed 12 0 98-6 86
Eiip,'
-1 -1 -1 -1
Prewave 11 f 0 17 10 f 0 03 12 i 0 04 13f005 /
C6H&OH -1 05 zt 0 02 Salicylic acid,d pK1 PKz 9 0 093-9 25 -1 11 f 0 08 ca. - 3 HzS04, pKi i .92 PKz 2.96 Phthalic acid, pK, 12 0 98-9 83 -1 09 i 0 03 5.51 PKz 0.071-1.42 8 EtpyrBr 0.109-2.18 12 n-BuBre a Values are those for aqueous solution, taken from authoritative references. * Potentials are us. N.4gE, corrected for iR drop. c I = il/Cm.?'3t1'6. Values are for combined currents of main wave and prewave. d Also exhibits a second wave: EllZ = -1.75 f 0.03 volt us. NAgE; I = 1.21 i 0.01. * n-BuBr would be expected to react with pyridine to form n-BupyrBr. f Prewave for this compound was not measured.
214
e
2.5
1.00mM HpyrNOa and 1mM n-PrN4Br; 6: 1.lOmM HpyrNOa and 0.1 mM n-PrN4Br; C: 1 .OOmM HpyrNOa; D : 1 .OOmM EtpyrBr. Background electrolyte: 0.1 M LiCI04
Background; 6: 1.12mM CFaCOOH; C: 1.OlmM CeHbCOOH; D: 1.25mM CH3COOH; E: 1.00mM HpyrNos; F: 1.03mM EtpyrBr; G: 1.1 3mM n-BuBr
II.
1.6
pyridinium ion
A:
Table
12
A:
F gure 4. Polarograms for pyridine solutions containing 0.1 M LiC104 as background electrolyte at 25°C.
wave. The mean diffusion current constant for the five monobasic acids is 2.04 f 0.17. Phenol (pKa = 9.9 in water) has been reported (13) to exhibit acid behavior in nonaqueous solvents such as dimethylformamide. However, its
I
I
08
04
C
ANALYTICAL CHEMISTRY
volt Main wave -1 30 f 0 02 -1 36 f 0 04 -1 36 f 0 07 -1 39 i 0 03 -1 36 i 0 04
2 04 f 0 10 2 23 & 0 25 2 16 & 0 10 2 01 i 0 07 1 78 i 0 13
-1 22 f 0 04
1 80 i 0 07
-1 35 i 0 03
3 26 f 0 55
-1 36 i 0 01
3 48 i 0 24
-1 34 f.0 03 -1 38 zk 0 04
2 15i.037 1 83 f 0 36
X square-root dependence of current on mercury height is usually characteristic of a diffusion-controlled wave, for which the temperature coefficient is usually between 1 and 2% per degree. The combined wave, therefore, appears to be diffusion controlled. Polarography of Dibasic Acids. & i nidentical polarographic pattern of a well-defined wave a n d a small prewave was observed with similar pyridine solutions of three dibasic acidsphthalic, salicylic, and sulfuricTable 11). T h e half-wave potentials are independent of acid concentration; the mean El,z for the main Tvave of the three dibasic acids studied is -1.31 volts us. K.igE and - 1.08 volts for the prewave; the means for sulfuric and phthalic acids are - 1.35 and -1.10 volts. I n addition, salicylic acid exhibits a second wave, whose Ell2 is also independent of concentration. The total limiting current of the prewave and the main wave (both waves in the case of salicylic acid) yaries linearly with concent,rat,ion. The mean diffusion current constants for the combined waves of the dibasic acids are 3.48 and 3.26 for phthalic and sulfuric acids, respectively, and 1.80 and 1.21 for waves I and I1 of salicylic acid. Polarography of Mixtures of Acids. As a check on the linearit'y and additive nature of the limiting current-concentration relationships, several mixtures of t,rifluoroacetic and benzoic acids (Table 111) were examined polarographically with 0.1M LiC104 as background. Polarograms were ident,ical with those obt'ained with single acids; the mean Eliz was -1.37 f 0.05 volts vs. NAgE for the main wave and -1.15 i 0.04 volts for the prewave. Total acid concentration calculated from the measured limiting currents (based on the mean I = 2.03 as described in the subsequent Summary of Polarographic Results) agreed wit,h the actual total concentration to within an average of 470. Polarography of S-Alkylpyridinium Salts. T w o N - alkylpyridinium salts-ethylpyridiniuni bromide and n-butylpyridiniuni bromide (generated in situ by the addition of n-butyl bromide to pyridine)-were also st'udied polarographically (Table 11). I'olarograms (Figure 4) were quite similar to those of the acids studied, except for the appearance of a large maximum on the rising portion of the wave and a shift of the anodic discharge to ca. 0.3 volt more negat,ive; the mean Eliz for t'he W-alkylpyridiniuni salts studied is -1.36 volts a's. NAgE. The mean limiting current constant for the alkyl~ ~ y r i d i n i usalts m is 1.99. The more negative anodic discharge found n i t h the S-alkyll)yridiniuli salts studied can be attributed to the bromide ion present-Le., oxidation of mercury
Table 111.
Polarographic Data
for Mixtures of Trifluoroacetic and Benzoic Acids in Pyridine Solution.
Mixture CFsCOOH, C6HSmH, mM mM
Taken, mi44
Total concentration Found,b mM Difference, %
0.00 1.12 1.22 0.10 0.50 1.62 1.01 2.13 1.01 1.12 1.01 1.01 0 Background electrolyte: 0 1M LiCIOa. Temperature: 25' C. Based on I = 2.03. 1.12 1.12 1.12 1.12 0.11 0.00
to mercurous bromide. This is readily seen by comparing polarograms of N-alkylpyridinium bromide solution with those of pyridinium nitrate to which bromide ion has been added (Figure 5 ) . The anodic discharge of a 1.OOmJZ solution of HpyrSOs containing 1.00m.ll n-Pr4NBr occurs a t precisely the same potential as that of a 1.00mJI EtpyrBr solution. However, the maximum observed in the latter case does not appear in pyridinium nitrate solutions containing bromide and must, consequently, be characteristic of the N-alkylpyridinium cation. Summary of Polarographic Results. T h e mean half-wave potential for the main wave of all the compounds studied (including five monobasic acids, t,hree dibasic acids, and two 5-alkylpyridinium salts) is - 1.34 i 0.05 volts us. NXgE. T h e mean value excluding for salicylic acid is -1.36 f 0.03 volts. T h e mean diffusion current constant for the five monobasic acids and two N-alkylpyridiniurn salts is 2.03 f 0.17; the values for the dibasic acids are somewhat lower than expected-Le., twice the above value. That for wave I of salicylic acid 5hows reasonably good agreement wit,h the value for the monobasic acids, which is to be expected considering the rather large magnitude of p K z (13.0) of salicylic acid. The mean E I l 2for the irewave of all compounds studied except the two S-alkylpyridinium salts and 2,4-dichlorophenol is - 1.10 i 0.09 volts z's. NAgE. The niagnitude of the prewave varies with the lot of pyridine used ( ~ . fpre.~ vious discussion). Coulometry. Coulometry on the crest of the main wave was performed on solutions of three representative compounds: pyridinium nitrate (3 runs), benzoic acid ( 2 runs), and sulfuric acid (3 runs). The faradaic n values found (range given), which were 0.94 i 0.03, 0.95 i 0.15, and 2.07 i 0.10 for the compounds as listed, clearly indicate a one-electron reduction. Macroscale Electrolysis. Macro-
scale electrolysis a t controlled electrode potential gave products which have not been isolated a n d identified, but which indicate t h a t the pyridine ring is attacked. For example, during evaporation (in vacuo) of the solvent from a 10mM n-butyl bromide solution, which had been electrolyted at -30' C., a blue color became evident as the solvent evaporated; the blue color deepened as the evaporation proceeded. Such a color is frequently characteristic of the presence of an unpaired electron as in a free radical. I t is hoped that further work will help in identifying the product or products produced.
REDUCTION MECHANISM
The behavior of pyridine as a Lewis base has been amply demonstrated (e.g., 7-9). The pyridine molecule acts as a base in donating an available electron pair from the nitrogen atom to form a coordinate bond with the potential hydrogen ion in an acid, thus promoting the dissociation of the acid and forming a pyridinium ion:
HA
+Q
0+ N
A-
(4)
H+
The extent of this reaction-Le., of the position of the equilibrium-is a function of the acidic strength of the acid involved (its tendency to accept an electron pair), and the basic strength of pyridine (the tendency of this solvent to donate an electron pair). The strength of t'he acid, at least in the particular solvent concerned, can be measured by the equilibrium constant for the reaction of Equation 4,but this approach is limited by the strength of pyridine as a base; in the case of a strong acid, the reaction of Equation 4 proceeds com1)letely to the right, as written. 'The solubility of some conipounds in pyridine, as in other solvents of low dielectric constant,, is understood VOL. 37, N O . 2, FEBRUARY 1965
215
in terms of acid-base phenomena (29)Le., coordination of the cation with molecules of solvent. The formation of ion pairs, which is favored in solvents of low dielectric constant, must also be considered. In the present experiments, an eseentially identical polarographic wave accompanied by a small prewave was observed for pyridine solutions of the following compounds: seven acids, whose pKa in aqueous solution varies from 7.9 to very strong, one pyridinium salt, whose parent acid is strong in aqueous solution, one :V-alkylpyridinThe iuni salt, and an alkyl halide. combined limiting current of the prewave and main wave increases linearly with concentration of added species. Concentration does not affect Ellz of either Fvave. The half-wave pot,entials observed for all of the compounds studied agree to within experimental error. Benzoic and hydroiodic acids have been reported (6) to give two waves ( E l l z= - 1.27 and -1.i volts ZJS. mercury pool) in dimethylformamide medium containing pyridine; t,he first wave was assumed to be caused by reduction of hydrogen ion and the second by reduction of undissociated molecules. Limiting current information was not reported, but the second wave was stated to increase with increasing acid concentration (the first wave presumably remained constant,). Qualitatively, this behavior agrees quite well with that found in the present study. However, it is unlikely that, the react,ions assumed to result in the two waves in diniethylformamide would occur in pyridine solution (cf. subsequent discussion). Half-wave potentials in the two solvents cannot be readily compared because of lack of knowledge of liquid junction potentials and common reference potentials. Potentiomet,ric titration curves in pyridine of a number of polybasic acids, among them citric, oxalic, and succinic, showed only a single inflection point which was equivalent to the total acidic hydrogen (20) ; maleic and salicylic acids, on the ot'her hand, exhibit two inflections. The authors postulate that differentiating titration of two acidic groups ran be successful only when the ratio p K l to p & is of the order of 4 or greater-e.g., the ratio pK,,'pK2 is of the order of 3 and 4 for oxalic and maleic acids, respectively. S o definite end point \vas obtained with phenol. .iI)parently, acids witha pK, (aqueous) greater than ea. 9 are too weak to form the I)yridiniuni ion in pyridine solution. This can account for the agreement of the diffusion current constant' for wave I of salicylic acid (pK2 = 13) with that of' the monobasic rather than the dibasic acid.. The rather negative ( - 1.75 volts) half-wave potential for wave I1 216
ANALYTICAL CHEMISTRY
t
1
I
,
0.6
0.9
1.2
0.0
1.5
I .a
2.1
-E G N A ~ E) V O L T S Figure 6. Polarogram of a pyridine solution, which is 0.03m M in benzoic acid and 0.1M in lithium perchlorate (Merck pyridine, lot No. 7475 1 )
of salicylic acid indicates that a totally different process--e.g., reduction of the carboxyl group-is taking place. In the present study, sulfuric acid, whose anion is nonreducible polarographically, also shows one wave and a small preivave at the same potentials as the organic acids studied; these waves are not found with (XH&S04. Since 97y0 sulfuric acid was used to prepare H2S04solutions in pyridine, the water content is 0.02 to 1.6mJI for 0.093 to 9.25mM solutions. However, no water Ivave was observed. Origin and Nature of Prewave. T h e magnitude of the prewave, found in polarograms of the pyridinium ion and related species, is dependent' on the quality of the solvent. I n the case of the pyridine mostly used in the present study, the magnitude of the prewave was very small; however, later studies using different lots of pyridine resulted in prewaves, whose magnitude was as much as t'hree times as great. This increase was also associated with the appearance of a small wave of about' equal height, just preceding the background discharge. Studies of dilute acid solutions indicate, as described, that, as acid concentration increases, the prewave magnitude increases (the sum of prewave and predischarge wave currents remaining constant) until it equals the original magnitude of the pre-discharge wave. Further increase in acid concentration results in the appearance and subsequent increase of the main pyridinium wave, with the prewave remaining constant and the pre-discharge wave being com-
pletely eliminated. This behavior suggests t'he presence of an impurity in the solvent, which is more easily reduced than the solution background (as evidenced by t,he pre-discharge wave), and which is a stronger base than pyridine, since it preferentially reacts with added acid to form a species which is more easily reduced than pyridinium ion. For analytical determination of acids in pyridine, it is advisable to use the sum of the limiting currents for both main wave and prewave. If the prewave is small, use of the main wave is often sufficient for appreciable concentrat,ions, especially if a standard series calibration is used. Work is in progress on the nature of the material causing the prewave as well as on its elimination from pyridine. Species Reduced. T h e only species present, which could give rise t o polarographic waves in solutions of Br@nsted acids in pyridine, are undissociated acid (HA), the corresponding acid anion pyridine, free hydrogen ion and pyridinium ion (see Equation 4). Due to the base behavior of pyridine, the existence of any but infinitesimal concentrations of free hydrogen ion is doubtful. Since neither HA nor pyridine is reducible at t,he D.M.E. within the potential range involved and the reduction of the anion does not seem to occur, the main wave observed must be caused by the reduction of pyridinium ion. Two reaction sequences giving rise to current flow on reduction may then occur: (1) formation of a free radical with subsequent dinierization to a t,etrahydrobipyridyl,
pyrH+
2 pyrHO
+e
-
+
pyrHo
Hpyr-pyrH
where Hpyr-pyrH such as
(5) (6)
may be a species
or (2) a cyclic process involving catalytic hydrogen evolution, pyrH+
+e
-
pyr
+ H”
2H” + Hz
(7) (8)
The latter process occurs in aqueous acid solutions containing small amounts of pyridine. I n pyridine solution, t,he available evidence indicates that the reduction of Iiyridinium ion follows the free radical scheme as opposed t o the cyclic process. The exact nature of the final product(s) formed has yet t o be determined. ANALYTICAL SIGNIFICANCE
The present investigation of the po1arogral)hic behavior of pyridinium ion in pyridine has indicated several types of analytical procedures based on the measurement of the pyridinium ion wave of equal to -1.36 volts in pyridine solution, which is 0.1JI in lithium perchlorate. ,Iny individual Brgnsted acid, whose pKa in aqueous solution is less than 9, can be polarographically determined. The analytical procedure is simplified by the fact that all monobasic acids meeting the PI
