Electrochemical Reduction of Alkali Metal Ions in Pyridine John Broadhead and Philip J . Elving
The University of Michigan, Ann Arbor, Mich. 48104 Electrochemical reduction of pyridine solutions of the alkali metal ions was investigated, using polarography, cyclic voltammetry, coulometry, and controlled potential electrolysis. Na(l), K(I), Rb(l), and Cs(l) are reversibly reduced at the DME and give a reversible cyclic voltammetric couple at the hanging mercury electrode. Reduction of Li(l), however, shows some degree of irreversibility at the DME, especially at higher concentrations. Cyclic voltammograms of Li(l) display a double reduction peak; the first peak height is directly proportional to Li(l) concentration and may involve reduction of a pyridine molecule in the Li(1) solvation sheath, rather than reduction of Li(1) to the metal amalgam; the second peak is proportional to the water and lithium contents and may involve reduction of pyridine in a (Pyr + HzO), Li(l) species. The diffusion current constants are considerably smaller than those in water and can be correlated with the ionic conductance of alkali metal salts in pyridine.
ANHYDROUS NONPROTON-RELEASING solvents are useful for investigating electrode processes, which occur at a relatively negative potential in water-Le., close to or beyond that of hydrogen evolution, e.g., electrochemical reduction of the alkali and alkaline earth metal cations. In some instances ( I ) , the solvent forms a complex or strongly solvated entity with the electroactive species and thereby modifies its behavior; in so doing, the solvent often dissolves substances which would otherwise be insoluble in low dielectric constant solvents. Nitrogen-containing compounds are especially active in this respect, since many have unshared electron pairs which readily contribute to the solvating or complexing process. Pyridine, in particular, is such a solvent; it is nonprotonreleasing and has an unshared pair of electrons localized on the nitrogen atom. Only a few of the electrochemical studies in pyridine (2) have been concerned with metal cations. The polarography and voltammetry of pyridine solutions of Cd(II), Co(II), and Zn(I1) (3); Tl(I), Zn(II), Pb(II), and Al(II1) (4); and Tl(I), Cd(II), Zn(II), and Mg(I1) (5) have been described. In the case of the alkali metal series, only the electrodeposition of lithium, sodium, and potassium at mercury, platinum, copper, and iron electrodes has been investigated (6, 7). The present study involves a polarographic and voltammetric investigation of pyridine solutions of the alkali metal ions at mercury electrodes, including correlation of the observed polarographic parameters with other available electrochemical data on these solutions. EXPERIMENTAL
Reagents. Lithium, sodium, and potassium perchlorates (G. Frederick Smith Chemical Co.) were recrystallized twice from water and dried at 130 "C. Rubidium perchlorate (1) R. Takahashi, Talanta, 12,1211 (1965). (2) R. F. Michielli, J. Broadhead, and P. J. Elving, personal communication, The University of Michigan, Ann Arbor, Mich. (3) F. Willeboordse, Ph.D. Thesis, University of Amsterdam, 1959. (4) A. Cisak and P. J. Elving, J . Electrochem. SOC.,110, 160 (1963). ( 5 ) R. F. Michielli, Ph.D. Thesis, University of Michigan, Ann Arbor, Mich., 1969. (6) S. Laszcynski and S . Gorski, 2.Elektrochem., 4, 290 (1897). (7) L. F. Audrieth and H. W. Nelson, Chem. Rev., 8,335 (1931). 1814
(Alfa Inorganics Inc.) was dried at 130 "C. Cesium nitrate (Fisher Scientific Co.) was recrystallized and dried at 130 "C. Tetraethylammonium perchlorate (TEAP) was recrystallized as necessary and vacuum-dried. Pyridine (J. T. Baker) was twice zone recrystallized (8),and then allowed to stand over Type 4A molecular sieves (Linde) ; the product contained 0.005% (2.5mM) of water (Karl Fischer titration). High purity argon (Matheson), used for deoxygenating the cell solution, was passed through Drierite, 4A molecular sieves, and, finally, pyridine maintained at the same temperature as the cell solution. Triple-distilled mercury was used for the dropping mercury electrode (DME) and the hanging mercury drop electrode (HMDE). Apparatus. Polarograms were recorded with a Leeds & Northrup Electro-Chemograph Type E, equipped with an operational amplifier-based IR compensator (9). A triangular voltage scan was generated by an operational amplifier control circuit (IO); resultant cyclic voltammograms were recorded by a Moseley Model 135 X-Y recorder. For controlled potential electrolysis, a potentiostat, similar to that in the cyclic unit, controlled the potential of the working mercury pool electrode and an electronic coulometer measured the total charge consumed. The DME capillary was a cracked-off Sargent S-29417 capillary; its constants (open circuit; h = 44.8 cm; 0.1M TEAP in pyridine) were t = 5.30 sec and m = 1.41 mg/sec-l. The HMDE suspension was a Sargent S-29314-30 platinum electrode plated with mercury; for cyclic voltammetry, one or two drops of mercury were collected from the DME. The three-compartment, water-jacketed cell (11) was thermostated at 25' =t0.2 "C, unless otherwise specified. A normal silver reference electrode in pyridine (NAgE) (4) was connected via pyridine-TEAP-methyl cellulose salt bridges as follows: j
Ag
~
AgNO, lM ~
1
0.1MTEAP i~ 0.1M methyl cellulose TEAP I ( 6 7 3 gel ;j 0.1M TEAP /! Cell methyl cellulose ii solution (62,') gel Ji
;;
The extra bridge allowed convenient disconnection of the reference electrode from the main cell. Reported potentials are us. NAgE. Procedures. Test solutions of electroactive species were made by dissolving weighed quantities of solid in 0.1M TEAP solution in pyridine; 10-ml samples were deoxygenated and then polarographed with argon passing over the solution surface. The mean diffusion current was determined by measuring the difference between the background current and the mid-points of the current oscillations on the polarographic plateau. The reference electrode was periodically checked against an aqueous saturated calomel electrode (SCE), directly and uia an extra 0.1M TEAP-methyl cellulose plug and frit assembly to simulate the polarographic cell arrangement. The potential of the NAgE US. SCE was 70 f 10 mV without (8) D. A. Hall and P. J. Elving, Anal. Chim.Acta, 39, 141 (1967). (9) R. Annino and K. L. Hagler, ANAL.CHEM.,35, 1555 (1963). (10) G. Dryhurst and P. J. Elving, ibid.,39, 606 (1967). (11) J. E. Hickey, M. S . Spritzer, and P. J. Elving, Anal. Chim. Acta, 35, 277 (1966).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
I
I
-16
-20 POTENTIAL,
-24
- 1.4
-I0 POTENTIAL,
V
Figure 1. D.C. polarogram of 4.2mM Li(1) at DME in O.1MTEAP in pyridine
-22 V
Figure 2. D.C. polarogram of 2.0mM Na(1) at DME in 0.1 M TEAP in pyridine
the plug and 30 =k 30 mV with the extra plug and frit. The ranges shown cover the extreme values recorded, which were always those measured before the reference half-cell was inserted in the polarographic cell. After use, the potential levelled off and approached the mean value. The reference half-cell should not be left in the arm of the polarographic cell as contamination of the second methyl cellulose plug with AgN03 can cause a considerable potential shift. The TEAP solution in the cell arm should also be changed frequently. POLAROGRAPHIC BEHAVIOR
Polarographic data are summarized in Table I ; typical polarograms are shown in Figures 1 and 2. Lithium. The Li(1) wave is not well-defined at concentrations above 2.5mM and exhibits a maximum of the first kind (Figure 1). The maximum disappears below 2.5mM, but the wave becomes more difficult to isolate from the increase in current because of background electrolyte electrolysis. The limiting current (il) is proportional to concentration, but is not strictly proportional to hl'z; the index of h is 0.44 for 4.0mM Li(1). The temperature dependences of il at 0.61, 2.5 and 4.2mM, over the range 25'-40 "C, are 1.0, 1.1, and 1.8% deg-I, respectively. At low concentration, the current appears to be diffusion controlled, but, above 2.5mM, some degree of complication is evident. Analytically, lithium is the only alkali metal ion which can be polarographically determined in the presence of the other alkali metal ions or which can be present when determining one of the others (Figure 3). The lithium wave is more easily measured in 0.1MBu4NC104solution, where EllZ= -2.11 V and the background decomposition is 80 mV more negative than in TEAP solution. Sodium. The Na(1) reduction wave is well-defined over the range of 0.7 to 4mM (Figure 2); it is proportional to concentration and h l / *with a temperature coefficient of 1.0% deg-1. Consequently, ir is diffusion controlled. The El,f of - 1.89 V is within 0.02 V of those of potassium, rubidium, and cesium. Potassium. K(I) also gives a well-defined reduction wave, whose i~ is proportional to concentration and the temperature coefficient is 1.2 deg-I.
I
1.6
2.0 2 1 POTENTIAL V Figure 3. D.C. polarogram of 0.15mM K(1) and 0.07m M Li(1) at DME in 0.1M tetrabutylammonium perchlorate in pyridine Rubidium. The Rb(1) reduction wave is well-defined with proportional to concentration and h0.69; the temperature coefficient is 1.1 deg-1. Cesium. The solubility of cesium salts in pyridine is so small that the limiting currents obtainable are only about 0.05 MA; therefore, measurement of the variation of ii with h and temperature is of little value. il
Table I. Polarographic Data for Alkali Metal Ions in Pyridine (0.1M TEAP)
Wavec Concn range, mM
slope, V Ib mV 0.61-7.70 2.17 =t0.01 1 . 3 A 0.1 62 0.69-4.10 1.89 f 0.01 1 . 6 f 0.1 56 0.04-0.15 1.91 f 0.01 1 . 5 f 0.15 55 0.04-0.20 Rb(1) 1.91 A 0.01 1 . 6 f 0.1 50 0.03 1.89 =k 0.01 1.4 0.2 63 CS(I) Data given generally represent the mean and average deviation of 5 to 20 measurements at 25 "C. Diffusion current constant, I = PA mM-' mg-*/3 sec1'2, for
Ion Li(1) Na(1) K(I)
-Ellla,
+
the mean diffusion current. Calculated from relation, slope
=
E114
- &la.
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
1815
H O ,
-60
-40
CONCENTRATION
0
-20
20
mM 40
60
I
I
I
I ooci INITIAL
POTENTIAL, V
Figure 4. Cyclic voltammogram at HMDE of 1.3mM Li(1) in 0.1MTEAP in pyridine
X
IO’
% H20
ADDED
Figure 5. Dependence of apparent “water” peak height on water content in 0.1MTEAP solution in pyridine
Li(1) concentration: A , 1.3 mM; B, 0.4 mM. Scan rate, 100 mV SeC-1
Scan rate, 10 mV sec-l
CYCLIC VOLTAMMETRIC BEHAVIOR
In order to investigate further their degree of reversibility, the electroreduction of the alkali metal ions was examined by cyclic voltammetry (Table 11). Lithium. Li(1) (Figure 4) displays a double cathodic peak, A and B, between -2.25 and -2.35 V with no corresponding anodic wave on the return sweep. Multiple cycling and clipping between -2.10 and -2.35 V also failed to produce a reoxidation wave. When the voltammogram is clipped between peaks A and B, the return anodic sweep does not show cathodic peak C. Because a couple is not observed even at 0.3 “C, where chemical decomposition reactions would be slowed down, the overall reduction appears to be irreversible; the -40 mV shift of lithium peak A (as opposed to that for the water peak B, discussed in the next paragraph) on a 10-fold increase in potential scan rate suggests some kinetic control of the electrode process. Metered addition of water shows peak B to be related to the water concentration, while peak A is directly proportional to the Li(1) concentration. The cell solution containing 0.1 % water was treated with 4A molecular sieves; after 4 hr, peak B completely disappeared. Peak B also appears to be more irreversible than A ; an increase of 10 in the scan rate shifts peak B more negative by ca. 85 mV. At a 100-mV sec-* scan rate, the two peaks are inseparable at a water concentra-
tion of 0.12% or greater-Le., the water peak becomes so large that peak A becomes a shoulder on peak B. Addition of water also shifts the background decomposition to more positive potential-e.g., by 200 mV at 1 % water. Figure 5 shows the effect of water on peak height. On the basis of proportionality of the second peak height to water concentration, extrapolation indicates that these solutions originally contained between 0.08 and 0.12% water, which probably originated mainly from atmospheric contamination during transfer of solution. Sodium, Potassium, Rubidium, and Cesium. These ions behave in essentially the same manner-e.g., in giving voltammetric couples with equal anodic and cathodic peak heights, whose peak potentials are separated by 65-70 mV (Figure 6). At slow scan rates-e.g., 10 or 20 mV sec-’-all of the alkali metals exhibit a group of irregular oxidation peaks on the return (anodic) sweep. When the sweep is reversed just after the metal ion reduction peak, these anodic peaks became very small or negligible. ELECTROLYSIS OF LITHIUM SOLUTIONS
Because cyclic voltammograms of Li(1) solutions show some abnormalities, controlled electrode potential electrolysis was used to determine the fate of the lithium after electroreduction. On electrolysis at -2.21 V of a 0.1M TEAP solution in
Table 11. Cyclic Voltammetric Potential Data for Alkali Metal Ions in Pyridine (0.1M TEAP) -Epbat scan rate of AEpcat scan rate of Ion V 2.17 1.89 1.91 1.91 1.89
Li(1) Na(1) K(I) Rb(1) CSU)
40 400 mV sec-1 mvsec-1 2.26 2.30 1.95 1.95 1.97 1.97 1.97 1.98 1.95 1.95
40 mVsec-I
400 rnVsec-’
70 65 68 68
70
. . .d
. . .d 65
68 70 I
Measured at the DME by d.c. polarography. Cathodic peak potential in V at the HMDE. Difference in potential in mV between the anodic and cathodic peaks. No anodic peak appears. Cathodic peak was displaced by -40 mV when the sweep rate changed from 40 to 400 mV sec-’. a
1816
a
-1.0
-2.0
-
Figure 6. Cyclic voltammogram at HMDE of 0.2mM Na(1) in 0.1MTEAP in pyridine Scan rate, 200 mV sec-l
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
pyridine containing 24 pmole LiC104, the current decreased exponentially to an asymptotic value 10-20x above the level determined for a background solution and continued at this level for at least 7 hours. For an end point taken as the time at which no further decrease in current occurred, the value of the faradaic n was 0.97. Samples of catholyte solution and mercury pool, removed without breaking the circuit, were analyzed for lithium by emission spectroscopy. Lithium was detected in the catholyte but not in the mercury. Analysis of the pool was carried out directly on the mercury and on an aqueous extract. Samples from the pool were also examined after an electrolysis in which 131 pmole LiC104 was reduced; again, lithium could not be detected. It thus appears that, after electrolysis, the majority of lithium remains or reappears in the catholyte solution and a negligible quantity occurs as an amalgam. With NaC104 as the electroactive species, a similar electrolysis generated a sodium amalgam at the mercury cathode. DISCUSSION
Table I11 summarizes the mean values of Ellzand diffusion current constant ( I ) for the reduction of the alkali metal ions in pyridine and in water, where data for the latter could be found in the literature. Values of I for water (as opposed to water-alcohol mixtures) were calculated from the data given by Zlotowski and Kolthoff (12). The spread of Ellzis not as great in pyridine as in water (12-14, but Eliz for Li(1) is still 0.26 V more negative than those for the other members of the group. More noticeable is the difference in Z values in pyridine, compared with those in water. The abnormally low I values for Li(I), Na(I), and K(1) in pyridine suggest that the ionic mobility of these ions is considerably less in pyridine than in water, possibly due, at least in part, to strong solvation of the ions by pyridine with the resulting ion-solvent complex being more bulky than that of water. It is significant that these ions also show abnormally low ionic conductances in pyridine (15). Diffusion coefficients calculated from the present I values and the data of Burgess and Kraus (15) show some similarity (Table IV). In general, the polarographic and voltammetric results indicate that, except for Li(I), the alkali metal ions are reversibly reduced by diffusion controlled processes in pyridine, forming reversible couples. The lack of concentration dependence of Ellz suggests that ion-pair formation is not a controlling factor. Because lithium was absent in the coulometric mercury pool electrode, it is possible that reduction of solvent in the primary solvation sheath occurred in preference to direct reduction of the metal ion; the unity value of n would indicate that only one solvent molecule per lithium ion was reduced. In the reduction of aluminum chloride solutions in pyridine (16), where a similar effect was observed, it was postulated that pyridine molecules solvating or complexing Al(II1) were opened by l e attacks on a bond adjacent to the nitrogen atom in the pyridine ring; this electroreduction was facilitated by the electron-withdrawing effect of Al(II1) on the nitrogen and (12) I. Zlotowski and I. M. Kolthoff, J. Amer. Chem. SOC.,64, . 1297 (1942). (13) I. Zlotowski and I. M. Kolthoff. IND. ENG. CHEM..ANAL. ‘ ED., 14, 473 (1942). (14) J. Heyrovsky and D. Ilkovic, Collection Czech. Chem. Commun., 7 , 198 (1935). (15) D. S . Burgess and C. A. Kraus, J . Amer. Chem. SOC.,70, 706 (1948). (16) A. Cisak and P. J. Elving, Electrochim. Acta, 10,935 (1965).
Table 111. Comparison of Polarographic Behavior of Alkali Metal Ions in Pyridine and Water
Pyridine medium Ion
Water medium
P -Eli2 1.29 2.35b 1.57 2.1@ 1.51 2.13b 1.56 2.12b 1.44 2 . Ogb constant, I = MA mM-1 mg-*’3
I 2.17 1 .93Ctd Na(U 1.89 2.2Jc K(I) 1.91 2 . 82c RbU) 1.91 CdI) 1.89 0. Diffusion current sec1’2, for the mean diffusion current. b Data taken from reference 18. Data calculated from information in references 12 to 14, apparently for normal polarographic current damping. ethanol-water solution to Data corrected from 50 volume water medium by a viscosity adjustment. Li(1)
--Ell2
Table IV. Comparison of Calculated Diffusion Coefflcients for Alkali Metal Ions in Pyridine Polarography Conductance Ion DO Da.b De Li(1) 0.45 0.41 0.61 NaU) 0.67 0.60 0.65 KO) 0.62 0.58 0.78 RWI) 0.66 0.59 ... CS(I) 0.56 0.51 ... Electro-Chemograph set at damping position 1 (equivalent to a normally damped galvanometer). b Calculated from the modified Ilkovic equation, 390 1/2t1/6 I = 607nD1/*(1 7).
+
c Calculated from the conductance measurements of Burgess and Kraus (15).
the resulting free radical anions were stabilized by the Al(II1). A similar process may well occur with Li(1) with formation of a strongly associated free radical anion-lithium(1) complex, analogous to the metal ketyls frequently postulated in nonaqueous electrochemistry. The effect of water on Lie) reduction in pyridine can be explained in terms of ion-water-solvent interaction. A volume concentration of 0.08 water would be approximately 44mM. Consequently, if the water itself were electroactive, peak B would be approximately 100 times the peak height because of reduction of the lithium-pyridine complex; experimentally, B is only one fifth of A (Figure 5B). If peak B is due to a differently solvated lithium species, the slightly more negative potential of this peak and its apparent greater irreversibility need to be explained. Thompson (17) has suggested that, at high concentrations of water in pyridine solutions of Li(I), the pyridine solvation sheath around Li(1) is replaced by a sheath of associated water-pyridine molecules to form
x
(17) W. K. Thompson, J. Chem. Soc., 1964,4028. (18) A. A. Vlcek, Chem. Listy, 48, 1485 (1954); Collection Czech. Chem. Commun.,20,413 (1955).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
1817
If such a species or an analogous one were present at low water concentrations, peak B could be due to reduction of a pyridine molecule attached to Li(1) via a water molecule. The reduction would be more difficult (more negative and more irreversible) than that of pyridine directly attached to lithium, because the bridging water molecule would weaken the electron-withdrawing effect of Li(1) and electron attack on the -N=C==- bond would, hence, be more difficult. The direct reduction of Li(1) may occur via the pyr-HzO
complex, but the large water concentration dependence suggests that this is not so. However, without direct coulometric evidence as to the nature of peak B, direct reduction of Li(1) cannot be completely dismissed. RECEIVED for review March 24, 1969. Accepted September 2, 1969. The work described was supported in part by the Petroleum Research Fund of the American Chemical Society and the National Science Foundation.
Study of the Chemical Reaction Preceding Reduction of Cadmium Nitrilotriacetic Acid Complexes Using Stationary Electrode Polarography Mark S. Shuman and Irving S h a h Department of Chemistry, University of Wisconsin, Madison, Wis. 53706 The theory of stationary electrode polarography has been considered for the electrochemical system k/
k,, an
ne
Ai=?O*R;
Two parallel homogeneous reactions precede this charge transfer
A
F
ki
R
Cd(NTA)-
kb
whereA and-0 are in chemical equilibrium in the solution. 0 undergoes a reversible charge transfer reaction at the electrode and A undergoes an irreversible charge transfer reaction. A numerical method was employed to solve the integral equations obtained from the boundary value problem. The electrolysis mechanism of cadmium and its complex with nitrilotriacetic acid (NTA) was selected to test the theoretical calculations and to demonstrate the use of cyclic stationary electrode polarography for this reaction scheme. Extensive correlations were made between the experiment and the theory. The rate constant for the direct dissociation of Cd(NTA)- was too small to be determined by stationary electrode polarography, but an upper limit of 5 sec-’ could be assigned. The value obtained for a parallel acid 0.5 X lO5M-l sec-l in assisted dissociation was 4.1 0.1M acetate buffer and 1.OM KN03.
*
THE POLAROGRAPHIC characteristics of cadmium in buffered solutions containing excess nitrilotriacetic acid has been studied by several investigators (1-5). In the pH range 2.5 to 6.0, two polarographic waves are observed. Koryta (1-3) investigated the polarographic behavior of CdNTA solutions in detail and has shown that the first wave is the reversible reduction of “free” cadmium; that is, cadmium in the aquated form or complexed with the buffer. The second wave was found to be the direct reduction of the CdNTA complex. Koryta ( 2 ) has proposed the following mechanism to explain the observed polarographic behavior. The first wave involves the charge transfer step Cd2+
+ 2e ~t Cd(ama1)
(1)
(1) J. Koryta, Sbornik Meziriarod Polarog. Sjezdu Praze, 1st Congr., 1951 Part I, p 798. (2) J. Koryta, Collect. Czech. Chem. Commun., 24, 3057 (1959). (3) Ibid., p 2903. (4) K. Morinaga and T. Nomura, Nippon Kugaku Zusshi, 79, 200 (1958). ( 5 ) N. Tanaka, K. Ebata, and T. Takahari, Bull. Chem. SOC.Jap., 35, 1836 (1962). 1818
k-
Cd2+ 1
+ NTAa-
(2)
kl
Cd(NTA)-
+ H+ k-i=? Cd2++ HNTAZ-
(3)
2
The direct dissociation of CdNTA through Reaction 2 is slow. The much faster acid-assisted path through Reaction 3 involves a protonated complex intermediate. This acidassisted path causes the observed currents of the first wave to be pH dependent. The second wave involves the charge transfer step Cd(NTA)-
+ H+ + 2e
k,, an
Cd(ama1)
+ HNTA2-
(4)
Because a protonated complex participates in the charge transfer, the potential at which the second wave appears is pH dependent. The dissociation rate constants kl and kz have been estimated from polarographic data by several workers ( I , 2 , 4 , 5 ) using the theoretical results of Koutecky (6). The values that have been obtained for k, are in the range from 1 sec-l to 3 sec-1, and the values of kz are in the range from 6.5 X lo4M-'set-lto1.5 X 1OGM-lsec-l. A typical stationary electrode polarogram of a CdNTA solution is presented in Figure 1. The solution contained 1.0 x lO-3M Cd(II), 4.0 X 10+M NTA, and 0.1M acetate buffer, pH 3.88. The scan rate was 0.062 V/sec, and the electrode area was 0.112 cm2. The peaks marked I and I1 correspond to the first and second waves of the classical polarographic experiment. Wave I11 corresponds to the oxidation of the cadmium amalgam formed during the reduction process. Because there is no anodic counterpart of wave 11, the reduction of the CdNTA complex must involve an irreversible electron transfer reaction. The changes of peak currents and potentials of the cathodic waves that occur with (6) J. Koutecky, Collect. Czech. Chem. Commun., 18, 597 (1953).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
