Voltammetric Study of the Hydrogen lon/Hydrogen Couple in Acetonitrile/Water Mixtures John A. Lanning and James Q. Chambers’ Department of Chemistry, University of Tennessee, Knoxville, Tenn., 37976
A voltammetric study at polished platinum electrodes of perchloric acid in acetonitrile/water mixtures is reported. The half-wave potentials of the proton reduction wave (ferrocene scale) and the viscosity of acetonitrile/water mixtures indicate that acetonitrile structures water in the ,0.80 to 0.95 mol fraction of water region. The proton is most difficult to reduce (E1/2 most negative) at 0.85 mol fraction of water. Acetonitrile proved to be a poor solvent for the generation of dry protons. Chemical reactions occur in H+/H2/CH3CN mixtures which depend on the nature of the generation technique employed. Solvent effects on the cyclic voltammograms of p-hydrobenzoquinone in acetonitrile-water mixtures were used to aid in the identification of product reduction waves. The wave assigned to the reduction of protonated benzoquinone was independent of solvent composition, in marked contrast to the reduction of solvated protons.
Protons are commonly produced as a result of anodic oxidation and, consequently, cyclic voltammograms of organic compounds often exhibit waves due to the reduction of protons. The position and wave characteristics of these “proton waves” depend on solvent basicity, solvation effects, and the presence of basic solute species which tie up the protons and shift the proton waves to negative potentials. Strictly nonaqueous conditions for solution electrochemistry are not achieved in practice since trace amounts of water are essentially impossible to remove. Thus, the experimenter is usually faced with the use of a mixed solvent whether he likes it or not. “Dry” protons have probably not been observed in solvents which are commonly employed in nonaqueous electrochemistry. Acetonitrile is a popular solvent for anodic electrochemistry since i t has desirable properties for this purpose ( 1 ) . The work reported here concerns the characterization of proton reduction waves in acetonitrile/water mixtures. In recent years, considerable progress has been made in the determination of single ion activity coefficients and medium effects in mixed solvents. The medium effect of a solute can be defined as a measure of the difference between the standard free energy of a solute in water and in a given nonaqueous solvent. ,G,”
- ,,.Gi0 = RT
In my,
(1)
Equation 1 states the relation between the medium effect, m ~ i ,and the standard free energies of solute i in a nonaqueous solvent, s, and in water, w. The medium effect so defined is the activity coefficient of solute i a t infinite dilution in the given nonaqueous solvent referred to pure water as the standard state. Popovych has presented these definitions and surveyed work in the field in a recent reveiw ( 2 ) . Address correspondence t o this a u t h o r . (1) C. K. Mann, “Electroanalytical Chemistry,” Vol. 3, A. J. Bard, Ed., Marcel Dekker, New York. N.Y., 1969, pp61-70. (2) 0. Popovych, Crit. Rev. Anal. Chem., 1, 73 (1970).
1010
A N A L Y T I C A L C H E M I S T R Y , V O L . 45, N O . 7, J U N E 1973
In the case where solute i is an ionic species, extrathermodynamic assumptions are required to obtain an estimate of the medium effect. The reference electrolyte assumption in which medium effects are assumed to partition equally between large symmetrical ions (e.g. in tetraphenylarsonium tetraphenylborate) or the use of a reference redox couple whose E” is assumed independent of solvent ( e . g . for the ferricinium/ferrocene couple) are the most popular approximations. In his review, Popovych ( 2 ) recommends the former with the ferrocene assumption being a “second choice.” These methods and others have been critically compared in recent publications (3, 4). Values of the medium effect for the proton can be readily converted to an estimate of the standard potential of the H+/H2 couple us. the ferricinium/ferrocene couple, E O H + I H ~by , means of Equation 2
where ( E O H + / H ~is) ~the standard potential of the H+/H2 couple us. the ferrocene couple in water. This information along with the hydrogen ion concentration can be used to estimate half-wave potentials using the equation given by Shuman ( 5 )
E,,,(H+/H1)
=
RT
- s[ln(r/2)
- In
C*,+] (3)
where y 2 is the ratio of the diffusion coefficients of H+ and Hz, C*H+ is the bulk concentration of hydrogen ions, and ED‘H+,H~is the formal electrode potential of the H+/H2 couple. Thus knowledge of medium effects in mixed solvents readily permits estimates of half-wave potentials and vice uersa. Medium effects on the proton have been estimated in a variety of solvents ( 2 ) . Acetonitrile complexes the proton less strongly than water and log m y H values of 5.1 and 5.5 have been determined for the proton between acetonitrile and water (6, 7), where m~~ is the medium effect. Kolthoff and Chantooni (8, 9) have studied proton solvation in acetonitrile/water mixtures and have been able to estimate the successive formation constants of the hydrated proton in these media. Barraque, Vedel, and Tremillon (10, 1 1 ) have determined medium effects for the proton in acetonitrile water mixtures in the form of Strehlow’s Ro(H) function (12). At ~ u H= 0, the Ro(H) function is equal to -log m~~ if the ferrocene (3) R . Alexander, A. J. Parker, J. H. Sharp, and W. E. Waghorne, J. Amer. Chem. SOC., 94, 1148 (1972). (4) I. M. Kolthoff and M. K. Chantooni, Jr., J. Phys. Chem., 76, 2024 (1972). (5) M. S.Shurnan, Anal. Chem., 41, 142 (1969). (6) I. M . Kolthoff and F. G. Thomas, J. Phys. Chem., 69, 3049 (1965). (7) N. A. Izmaylov, Dokl. Akad. NaukSSSR, 149, 1364 (1963). 90, (8) I. M. Kolthoff and M. K. Chantooni. Jr., J. Amer. Chem. SOC., 3320 (1968). (9) M. K. Chantooni. Jr., and I. M. Kolthoff, J . Amer. Chem. SOC., 92, 2236 (1970). (10) C. Barraque, J. Vedel, and B. Tremillon, Bull. SOC. Chim. Fr., 3421 (1968). (11) J. Vedel,Ann. Chim. (Paris), 335 (1967). (12) H. Strehlow, “The Chemistry of Nonaqueous Solvents,” Vol. 1, J. J. Lagowski, Ed., Academic Press, New York, N.Y., 1966, pp 129171.
assumption is made. Vedel’s plot of RdH) us. the mole fraction of water was not simple and exhibited a minimum in the low mole fraction region (11). Proton waves in several solvents have been briefly examimed by Nelson and Adams (13) as a diagnostic tool for investigating organic electrode reactions. They found that the proton reduction wave was irreversible in acetonitrile but did not pursue the issue. We initiated this study in order to clarify the nature of the proton wave in acetonitrile/water mixtures.
EXPERIMENTAL Apparatus. All electrochemical instruments and electrolysis cells were of conventional design. In most of the experiments, uncompensated resistance was minimized by means of a positive feedback to the control amplifier of the potentiostat. The two working electrodes used in most of these studies were planar platinum electrodes sealed in glass. T h e electrodes were polished to a bright, flat surface using a n aluminum oxide polishing agent, sizes 22 to 0.05 pm (Buehler Micropolish). The large platinum electrode was a commercial Beckman button electrode, area 0.223 cm2, and the small electrode was sealed in soft glass, area 0.82 mm2. The electrode areas were measured mechanically with a magnifying comparator and electrochemically by chronoamperometry on a 4.0mM K4Fe(CN)e solution, 2.OM in KC1 (14). Reproducible electrochemical results were best obtained when the working electrode was polished prior to each cyclic voltammogram or current-time curve. Two types of reference electrodes were used in this investigation. For aqueous systems or mixed aqueous systems, an aqueous saturated calomel electrode, SCE, of conventional design was used. For nonaqueous acetonitrile systems a Ag lAgN03 reference electrode was used. The potential of this electrode was measured after each experiment and was +0.30 V us. SCE. Experiments under rigorously nonaqueous conditions were carried out in a nitrogen atmosphere (Vacuum Atmospheres Co., Model HE-43-2 Dri-Lab). This enclosure had provision for continuous removal of oxygen and water to less than the part per million level (Vacuum Atmosphere Corp., Model HE-493 Dri-Train). In some experiments, temperature control was achieved using a conventional water bath, but most of the results reported here were obtained a t room temperature, 24 f 2 “C. Reagents. T a p distilled water was redistilled from KMn04, passed through an ion exchange column of 20-50 mesh Amberlite MB-3 mixed ion exchange resin, and then redistilled again. The resulting water had a specific conductance of 1.2 X 10-6 cm-1 ohm-l. Acetonitrile was purified following a slightly modified procedure of O’Donnell, Ayres, and Mann (15). The final distillate was collected in a 1-1. glass storage vessel which was painted black and maintained in a dry nitrogen atmosphere. The water content in the acetonitrile was measured using a 5 ft X 0.25 in. column of “Porapak Q” at 155 “C in a Wilkens Instrument and Research Model A-90 gas chromatograph employing a thermal conductivity detector (1). A 10-pl sample of freshly distilled acetonitrile gave no measureable water peak which gives an estimated water concentration of less than 1 X M . It is our experience that, by using careful distillation procedures and by discarding generous forecuts, “better” acetonitrile can be prepared in this manner than by using the more convenient method of Osa and Kuwana (16). Under scrupulously anhydrous conditions, anodic background current densities of less than 40 p A cm-2 were obtained a t +1.90 V us. SCE a t a sweep rate of 50 mV sec-l. For comparison, Osa and Kuwana report a background current for spectro-grade acetonitrile passed through a n alumina column of 50 *A cm-2 at f1.75 V us. SCE a t sweep rate of 53 mV sec-1 (16). The supporting electrolytes, sodium perchlorate, tetraethylammonium perchlorate (TEAP), and tetra-n-butylammonium tetrafluoroborate (TBAF), were repeatedly recrystallized from water and dried in a vacuum oven a t 60 “C before use. The perchloric acid used in this investigation was J. T . Baker reagent 70-7270 perchloric acid. The concentration of the stock (13) R . F. Nelson and R . N. Adams, J. Elecfroanal. Chem., 16, 439 (1968) (14) M . von Stackelberg, M. Pilgrarn, and V . Toome, Z. Elektrochem., 5 7 , 342 (1953) (15) J. F. O’Donnell, J. T. Ayres, and C. K. Mann, Ana/. Chem., 37, 1161 (1965) ( 16) T. Osa and T. Kuwana, J. E/ectroana/. Chem., 22, 389 (1969).
380
300
200
140 0.2
0.4
0.6
0.8
Figure I.Variation
of half-wave potential of ferrocene in acetonitrile/water mixtures I
8
I
I
41r
w
1.6 . E
0 c 0
1.2
L I
I
0.2
I
Conc.
Figure 2. Variation of
0.6
0.4
0.8
NoC104,
apparent proton diffusion Walden product
with sodium perchlorate concentration
2
4 Conc. HC104,
6
mM
8
Figure 3. Variation of apparent proton diffusion coefficient with perchloric acid concentration
solution was 11.82M when titrated against standard sodium hydroxide. The stock solution was delivered from a 25-pl Hamilton 702-N microliter syringe with an error of ca. *O.l pl, The concentrations of several of the test solutions, including 0.1M supporting electrolyte, were measured independently by titration. The results indicated that the concentrations were accurate to ca. fl%. Potassium ferrocyanide was recrystallized from water (twice); 1,lO-phenanthroline iron(I1) perchlorate, G. F. Smith Chemical Co., No. 163, was dried in a vacuum oven prior to use; and practical grade dicyclopentadienyl iron, Eastman P 8292, was twice recrystallized from aqueous ethanol and dried in vacuum prior to use. Procedure. Half-wave potentials were referenced against the ferricinium/ferrocene couple in order to estimate medium effects on the proton in acetonitrile-water mixtures. This procedure is based on the extrathermodynamic approximation that solute-solvent interactions are identical for the oxidized and reduced forms of the ferrocene couple. The plot of the E112 of the ferrocene couple is shown in Figure 1; data to construct this curve are availA N A L Y T I C A L C H E M I S T R Y , V O L . 45, NO. 7 , J U N E 1973
*
1011
___________
_
~~~
_
_
_
_
_
~
~~
~
~
Table I. Typical Cyclic Voltammetric Parameters for the Reduction of Hydrogen Ions in Aqueous Solution (2H+ 2e- = H2) Experimental Theoretical Parameter resulta valuea.*
+
Peak current (ipc) Peak potential (EPc) Peak width ( E p 1 2 C - Epa) Peak separation (Epa - EpC) Peak current ratio (ipc/ipa)
372 p A -0.436 V 41.2 rnV 4 4 mV 0.91 3 1 . 3 mV
6EpC/6 log [H']
373 p A -0.434 V 40.2 mV 4 5 rnV 0.920 29.5 mV
All potential measurements vs. SCE. Based on a second-order process (20) using a 4.02 X M proton concentration, a sweep rate of 22.4 mV sec-', a proton diffusion coefficient of 8.6 X cm2 sec-', and a haif-wave potential of -0.414 V vs. SCE. Q
able (17). The data in Figure 1 agree well with the literature, although previous investigators have not reported a break a t 0.85 mol fraction of water. Kolthoff and Thomas give +0.145 V us SCE for the E1 2 of ferrocene in water, in good agreement with Figure 1, and a n overall A E l l z of 220 mV (6). We obtained a slightly larger AE1/2, 235 mV. The break a t 0.85 mol fraction of water was also present in El,z data for the phenanthroline complex in acetonitrile/water mixtures. Diffusion coefficients were measured by chronoamperometry using the Cottrell equation a s modified by Soos and Lingane (18).
In this equation K is a constant, empirically equal to 2.12, po is the radius of a n unshielded planar electrode, and the remaining terms have their usual significance. We have used this equation extensively and have found t h a t most accurate diffusion coefficients are obtained by extrapolation of the left-hand side of Equation 4 to zero time. We have established the reliability of this method by reference to known diffusion coefficients of ferrocyanide, silver ion, and hydrogen ion in aqueous solutions. This procedure permits the electrode surface to be easily accessed for polishing between runs since cumbersome electrode shields are not required. In general, values calculated directly from Equation 4 or from the slope of the it1 2 us. t l ' 2 curve were somewhat low. Errors in the diffusion coefficients are given as the standard deviation of the intercept of a least-squares line fitted to the current-time data between l and 10 sec. All viscosity measurements were made with a simple Oswalt viscometer with the temperature maintained by a constant temperature bath (*0.02 "C). The empirical equation
q = Apt
+ Ep/tn
(5)
was used to calculate the viscosity ( q ) , where p is the density of the fluid, t is the time required for solution flow through the capillary, and A , E, and n are empirical constants (19, 20).
RESULTS AND DISCUSSION Aqueous Solution. The voltammetric parameters for the reduction of hydrogen ions in 0.1 M TEAP, which are summarized in Table I, present no surprises. The data indicate a nonunity reaction order and a diffusion-controlled process a t slow sweep rates corresponding to the well known electrode reaction 2 p
+ 2e-
H2
(6)
Excellent agreement was obtained with t h e theoretical values calculated from the equations given by Shuman (5, 21) and by Saveant and Vianello (22). At sweep rates greater than ca. 0.1 V sec-1, deviations from reversibility ( 1 7 ) J. A. Lanning, P h . D . Thesis, University of Tennessee, Knoxville, Tennessee, 1972. (18) 2. G . Soosand P. J. Lingane, J. Phys. Chem., 68, 3821 (1964). ( 1 9 ) F. Daniels, J. W. Williams, P. Bender, R . A. Alberty, and C. D. Cornwell, "Experimental Physical Chemistry," McGraw Hili, New York. N . Y . . 1962, pp 147-155. (20) M . R. Cannon. R . E. Manning, and J. D. Bell, Ana/. Chem., 32, 355 (1 960). ( 2 1 ) M . S. Shuman. Ana/. Chem., 42, 521 (1970). (22) J . M. Saveant and E. Vianello, Eiectrochim. Acta, 1 2 , 1545 ( 1 9 6 7 ) .
1012
A N A L Y T I C A L C H E M I S T R Y , VOL. 45, N O . 7 , JUNE 1 9 7 3
appear in the form of increases in the peak widths (EP -
Ep/2)and the peak separation (Epa - E p C ) . The proton diffusion coefficient in water was determined by chronoamperometry using the Soos-Lingane correction for edge effects (18). Since the diffusion coefficient of the proton is extremely large in water, nonlinear diffusion to planar electrodes is significant a t moderately short times even with the largest electrode used in this study. At 10 sec, the second term in Equation 4 accounts for 19% of the it1'2 value with the Beckman electrode. The average value of D H + for nine experiments at 25.0 "C in 0.1M supporting electrolyte (NaC104) is 8.59 f 0.12 X cm2 sec-I. This value agrees with that of von Stackelberg and Pilgrim, 8.63 X 10-5 cm2 sec-1 ( 1 4 ) . In applying this D value, one must remember that it is dependent on both the concentration of supporting electrolyte and electroactive species. This is true for any D value determined electrochemically, but the dependence is especially marked here since the mobility of the proton is very large in water. The variations of the apparent diffusion coefficient (Dapp) with supporting electrolyte and hydrogen ion concentration are shown in Figures 2 and 3. In Figure 2 , the Walden product ( 0 7 ) is plotted in order to take into account the small changes in solution viscosity (7)as the NaC104 concentration varies from 0 to 1.OM (23). As the ratio [H+]/[Na+] increases, the contribution of proton migration to mass transport increases and, consequently, Dappincreases. A rigorous theoretical treatment of the chronoamperometric experiment that includes a migration term in the mass transport equations has not appeared in the open literature. Migration effects have been discussed, but either for the cases of rotating disk electrode voltammetry (24), chronopotentiometry (25), or chronoamperometry in the complete absence of supporting electrolyte (26). We find that the empirical equation (27)
(7) appears to give reasonable values for the diffusion coefficient corrected for migration effects, where t~ is the transference number of the proton. Using this equation, a value of 7.1 x 10-5 cm2 sec-l is obtained from the data in Figure 3 if the points a t 0.5mM are neglected. This agrees with the value which is obtained by extrapolating the straight line portion of the curve to zero concentration and is in reasonable agreement with values in other electrolytes (28). The anomolous Dapp values at 0.5mM HC104 (Figure 3) remain unexplained. They cannot be ascribed to a basic impurity in the system since the concentration was accurately determined by titration in the presence of supporting electrolyte. Furthermore, the chronoamperometric data on these solutions appear to be well behaved. Acetonitrile-Water Mixtures. Several dramatic changes in the proton reduction wave occur when acetonitrile is added to aqueous HClO4 solutions. The wave shape changes from that of a reversible second-order process to that of an apparently irreversible electrode reaction, the proton diffusion coefficient decreases, and the wave shifts in a positive direction by more than 0.5 V us. E . R. Nightingale, J. Phys. Chem., 63, 742 (1959) S. L . Gordon, J. S. Newman, and C. W . Tobias, Ber. Bunsenges. Phys. Chem. 70, 414 (1966) M . D . Morris and J. J. Lingane, J . E/ectroana/, Chem.. 6, 300 (1963) H . A . Laitinen, Trans. Eiectrochem. SOC.,82, 289 (1942). M . J. Pikal, Lilly Research Laboratory, P.O. Box 618, Indianapolis, Ind.. 46206, personal communication (1972). R . A . Robinson and R. H . Stokes, "Electrolyte Solutions." Butterworths. London, 1965. Table 1 1 . 7 , p 317.
650 I
’
‘O0I 350
a
I
0.2
1
0.4
0.6
0.8
1
I
I
0.2
XH20
0.4
X
I
I
0.6
0.8
I
H20
Figure 4. Shift of H f / H 2 couple relative to ferrocene in acetonit r ile / wa ter rni x t u res
Figure 5. Proton diffusion coefficients in acetonitrile/water mixtures
SCE. These effects, which can be attributed to changes in the solvation of the proton and in solvent structure, are discussed in the paragraphs which follow. Half-Wave Potentials. In order to take into account changes in the liquid-junction potential as the mole fraction of acetonitrile increased and to accurately estimate medium effects, the half-wave potential of the proton wave, E1/2(H+/H2), was referenced against the ferrocinium/ferrocene couple. The results are shown in Figure 4. There are several interesting features in these data. First, 0 is a the marked positive shift of E1,2(H+/H2) as XH,O reflection of weak solvation of the proton by acetonitrile compared to water. Variations of this type have been reported previously for proton polarographic waves (29, 30), but the magnitude of the shift in Figure 4 is greater than that previously reported. Coetzee and Kolthoff (30) report that there is no change in the proton reduction potential for water concentrations below 1.3mM. This seems unlikely in view of Figure 4. The results are also plotted as Ro, the solvation coefficient of Strehlow (12). Strictly speaking, the Ro function should be calculated using E” values and not El12 values which are dependent on concentration and the ratio of diffusion coefficients. However, the data presented in Figure 4 were obtained a t constant concentration of H t (4.0X 10-3 M ) . The variation of the apparent diffusion coefficient of H + in acetonitrile/water mixtures is shown in Figure 5. A maximum variation of 10 mV in E1,2(H+/H2) is calculated from Equation 3 using the data in Figure 5 and assuming that the diffusion coefficient of hydrogen is constant. This variation is small compared to the total change observed in Figure 4. A second complication involves the Ell2 measurement a t a low mole fraction of water. The E1/2 values were arrived a t by averaging the anodic and cathodic peak potentials. As the mole fraction of water goes to zero, the anodic peak current becomes less well defined and an estimate of E1,2 becomes more difficult. However, it is felt that this procedure gives a reliable estimate of E1,2(H+/ Hz) in all but the driest solutions. In dry acetonitrile, the peak potential is more positive and no reverse current is observed in the cyclic voltammograms. Thus, we have not been able to estimate a reliable E1,2(H+/H2) in dry acetonitrile and these points are not included in Figure 4. The problems associated with obtaining dry HC104/ CH&N solutions are discussed below. The qualitative features of Figure 4 are similar to the previous results of Vedel (II), with two exceptions. Our
E1,2
+
(29) A A Vlcek, Chem L/sty, 48, 1741 (1954) (30) J F Coetzee and I M Kolthoff, J Amer Chem SOC 79, 6110 (1957)
-
data span a considerably wider range than Vedel’s and the minimum in the Ro(H) curve of Vedel as X H ~ O 1 is not present in our data. On the other hand, the acidit y function data of Vedel are in better qualitative agreement with Figure 4. Our results are in better agreement with those of Kolthoff and Thomas (6) and Kolthoff and Chantooni (8, 9). Values for the proton medium effect (log scale) in acetonitrile of 5.1 and 5.3 can be calculated, respectively. from the potential data on the hydrogen electrode (6) and the formation constants of the hydrated proton in acetonitrile/water mixtures (8, 9). The latter study, which employed acid-base indicators to estimate ~ U H ,showed that the tetrahydrate species, H(H20)4+, is the principal proton species for water concentrations greater than cu. 0.3 M in acetonitrile. A value of 4.4 is calculated from the data in Figure 4 for the proton medium effect in acetonitrile. This low value probably results from the neglect of the significant amount of water present in the solutions which give the most positive E1,2(H+/H2) values in Figure 4. However, the data in Figure 4 exhibit a maximum in the region of 0.8 mol fraction of water. The maximum AE1,2(H+/H2) is 336 mV between acetonitrile and 0.75 mol 70of water solution. Chantooni and Kolthoff ( 4 ) have estimated an error of 2.3 log units in the ferrocene assumption in acetonitrile in acetonsolutions. They report a value of 8.1 for log itrile based on the reference electrolyte assumption (tetraphenylarsonium tetraphenylborate) which is significantly larger than the above values based on the ferrocene assumption. Using the “error” in the ferrocene assumption, we obtain a value of 6.7 for the proton medium effect in acetonitrile. It is disturbing that on the basis of 8.1 for log mYH, the aqueous SCE appears to be a better reference redox electrode than the ferricinium/ferrocene couple. Preferential solvation of (C6Hs)4As+ in water is known (31) and probably is a contribution to the large value of log m~~ obtained with the reference electrolyte assumption. Behavior in the 0.7 to 1.0 Mole Fraction of Water Region. Figure 4 shows that the proton reduction potential becomes less negative above ca. 0.7 mol fraction of water. We believe that this variation reflects a solvent structure effect on the activity of the proton in these media: Other possible perturbations on E1/2(H+/H2) in this relatively simple experiment are small compared with the variation seen in Figure 4. Variation of the ratio of diffusion coefficients ( D H + / D H will ~ ) alter E1 2(HA/H2) but the effect is small and in the wrong direction. Using Equation 3 for (31) J F Coetzeeand W R Sharpe. J Phys. Chem. 7 5 , 3141 (1971)
A N A L Y T I C A L C H E M I S T R Y , VOL. 45, NO. 7 , J U N E 1973
1013
I
8
Table II. Reduction Wave Parameters for Anhydrous Protons in Acetonitrile
10-
0.
Experimental conditions
u
-8-
About 2 m M HCIO4 in dry acetonitrile, 0.1 M T E A P , outside'drybox About 4 m M protons from electrolysis of 0.1 M NaCIO4 at f 2 . 5 V vs. SCE, outside
2 X
F 6 -
0.2
0.4 X
0.6
0.8
H2° Figure 6.
Viscosity of acetonitrile/water mixtures, 0.1 M in
El12, the
data in Figure 5 for the diffusion coefficient of
TEAP
D H + ,and a constant value of D H ~a, maximum A E l l z of less than +6 mV is obtained for E1/2(x = 0.7) - El& = 1.0). This would increase the break observed in Figure 4. A second possible reason for the increase in E1,2(H+/ H2) a t high water mole fractions is a change in the solvation of hydrogen gas. Data bearing on this question are in the literature. Ben-Naim and coworkers determined thermodynamic parameters for the transfer of inert gases from the gas phase to the solution phase of several mixed solvents (32-34). For example, for the transfer of argon to waterlethanol, waterlp-dioxane, and water/ethylene glycol solutions, abnormal changes in the free energy of transfer occur in the region from 0.7 to 1.0 mol fraction of water. Only in the case of ethylene glycol is the effect in the proper direction to rationalize the observed behavior in Figure 4. In this case the approximate 300 cal mol-1 change corresponds to a 27 mV potential shift in the positive direction. In the other mixtures the shift is less than 27 mV and in the negative direction. Furthermore, the solubilities of hydrogen in dimethyl sulfoxide/water mixtures have been determined and these data predict a shift of ca. 10 mV in the negative direction (35). Thus we are lead to conclude that the variation of Ellz(H+/Hz) in this region is due to a medium effect on the proton or the ferricinium/ferrocene couple. The former will make the greatest contribution. Thus, these results suggest that the basic solute, acetonitrile, increases the solvent basicity in the 0.8 to 0.9 mol fraction of water region. This increase in basicity is similar to that observed in alcohol/water mixtures by Brande and Stern (36) and by Popovych et al. (37). This phenomenon can be rationalized by postulating that the solute disrupts the water structure, which exposes more hydrogen-bonding sites, and therefore increases the solvent basicity (2, 3 6 ) . This behavior for acetonitrile/water mixtures is consistent with the suggestion by Armitage et al. (38) that intercomponent hydrogen bond formation predominates in these mixtures. There is also some evidence in our data for weak intensification of the water structure a t low concentrations of acetonitrile (mol fraction 0.95 to 1.00). Anomalous be(32) (33) (34) (35)
A. Ben-Naim and G. Moran, Trans. FaradaySoc., 61, 27 (1965). A.,Ben-Naim,J. Phys. Chern., 72, 2998 (1968). E. A. Symons, Can. J . Chem., 49, 3940 (1971). F. Franks and D. J. G. Ives, Quart. Rev., Chem. Soc., 20, 1 (1966). (36) E. A. Brandeand E. S. Stern.J. Chern. Soc., 1976 (1948). (37) 0. Popovych, A. Gibofsky, and D. H. Berne, Ana/. Chem., 44, 811 (1972). (38) D. Armitage, M. J. Blandemer. M . J. Foster, N. J. Hidden, K. W. Morcom. M. C. R. Symons, and M. J. Wooten, Trans. Faraday Soc., 64, 1193 (1968).
1014
ANALYTICAL C H E M I S T R Y , VOL. 45, NO. 7, J U N E 1973
drybox Addition of acetic anhydride, 6 m M HC104, 0.1 M TEAP, outside drybox About 2 mM protons from electrolysis of 0.1 M T E A P at f 2 . 5 V vs. SCE, inside drybox 2.5 mM protons from hydrogen oxidation at f 0 . 6 V vs. A g - A g N 0 3 , 0.1 M TBAF, inside drybox 2.5 mM protons from hydrosen oxidation at . +0.6 V vs. A g - A g N 0 3 , 0.1 M T E A P , inside drybox
Peak potentialQ
Peak widthQ
i-0.063
70
fO.OO1
60
-0.030
160
-0.142b
85
-0.330
78
-0.37ab 78 Peak potentials are volts vs. SCE, and peak widths are in mV. Cyclic voltammograms show evidenceof a small couple at +0.19 V vs. SCE.
havior in this region is shown by our data for the viscosity of acetonitrile/water mixtures (Figure 6) and there are hints of inflections in this region in the E1IZ(H+/Hz) and DH+ data (Figures 4 and 5 ) . However, this intensification of the water structure is much less than that which occurs in the 0.8 and 0.95 mol fraction of water region in many water/alcohol mixtures (39-41). Analytical Application. The use of Figure 4 for the quantitative determination of water in the 0.005 to 0.5 mol fraction of water region is suggested by the form of this curve in the low mole fraction region. The variation of E ~ Q ( H + / Hwith ~ ) respect to the SCE is even more spectacular a t low water concentration and is conveniently used in place of Figure 4. We have used this method, along with anodic background limits, as qualitative checks on the water content of ordinary acetonitrile. Interferences result from basic components of the solution to be analyzed. Analysis time should be shorter and the procedure simpler than for a gas chromatographic or volumetric water determination. Unfortunately, use of this analytical method a t very low mole fraction of water mixtures is precluded in acetonitrile solutions by the difficulty of producing protons in solution in the absence of water. Generation of Dry Protons in Acetonitrile. The use of concentrated perchloric acid as the source of protons in acetonitrile has a serious drawback, commercial concentrated perchloric acid is ca. 30% water. Three techniques were employed to produce dry protons (H+A.N) in anhydrous acetonitrile-one chemical and two electrochemical. The results are summarized in Table 11. Acetic anhydride was added to the acetonitrile-HC104 solution to convert the water to acetic acid which is undissociated in acetonitrile/water mixtures (30). Although this technique was not extensively researched, the addition of small amounts of acetic anhydride increased the background current and distorted the proton reduction wave. Furthermore, the water content, determined by gas chromatography, was not substantially reduced. This method can also be criticized in principle since the protons in these solutions would probably contain acetic acid in their solvation shells. Thus Hf4N would not be formed. (39) E. M. Arnett and D.R . McKeivey, J , Amer. Chem. Soc., 87, 1393 (1965). (40) R. Lurnryand S.R . Rajender, Biopolymers, 9 , 1125 (1970). (41) R. E. Robertson and S . E. Sagamori. Can. J. Chem., 50, 1353 (1972).
In any event, results using this technique were far from satisfactory, and proton reduction waves were too distorted to be of any use. A second literature method for the possible generation of H+*N involves the electrolysis of perchlorate solutions to form HC104 (42, 43). This method was used both under anhydrous and ordinary conditions. The results are given in Table II. The experiments indicate that protons can be generated by this method even though the overall reaction mechanism is not known. The experiment performed outside the drybox shows evidence of traces of water because there was a small anodic current associated with the main reduction wave. The somewhat negative peak potential of +0.001 V us. SCE also indicates a small amount of water. Under rigorously anhydrous conditions, a “proton wave” with a peak potential of -0.142 V us. SCE was produced by this procedure. This potential is too negative to be ascribed to the reduction of H + A N or H+(H20). A third attempt at the production of H+*N involved the oxidation of hydrogen gas a t platinum under scrupulously anhydrous conditions ( 4 4 ) .Hydrogen was passed through an electrolysis cell while a rotating platinum electrode, which was potentiostated a t +0.6 V us. Ag/Ag+, stirred the solution. After completion of the electrolysis, excess hydrogen was removed with high purity nitrogen. A cyclic voltammogram of the resulting solution is shown in Figure 7. A broad reduction wave is present with a peak potential of -0.38 V us. SCE. A small anodic current is evident at +0.23 V us. SCE which grows with repeated cycles. The peak potential of the main reduction wave in Figure 7 is again too negative to be assigned to H + A Nor H+ (HzO), reduction. Furthermore, addition of small amounts of water to the solution of Figure 7 caused no appreciable negative shift of the voltammogram ( c f . Figure 4). Clearly some component of the system, most likely acetonitrile, has reacted during the generation process. Barraque et al. (IO) state that acetonitrile can undergo reduction by hydrogen which is catalyzed by strong acid media. Since hydrogen was continuously passed through the acetonitrile solution during electrolysis, formation of an unknown basic species (B) is possible. Several attempts were made to identify B. The simplest reduction product of acetonitrile would be ethylamine. A solution of protonated ethylamine gave a broad, irreversible reduction wave with a peak potential of -0.835 V us. SCE. Alkyl amines have similar basicities which rule out other possibilities as di- and triethylamine (from TEAP) for base B. Acid-catalyzed hydrolysis of acetonitrile leads to the formation of acetamide, CH3CONH2. Protonated acetamide is reduced in an irreversible process (EP = -0.175 V us. SCE) in acetonitrile. This peak potential is in the region in which the reduction wave is observed from the NaC104 oxidation (Table 11). A third possibility for B is the cyclic trimer, 2,4,6-trimethyl-S-triazine, obtained from acetonitrile.
I
I
1
’
(45). (42) (43) (44) (45)
H. Schmidt and J. Noak, Angew. Chem., 69, 638 (1957). H. Schmidt and J. Noak, Z.Anorg. Ailg. Chem.. 296, 262 (1958). J. Vedel and B. Trernillon, J. Necfroanai. Chem., 1 , 241 (1959). D. Clark, M. Fleischmann. and D. Pletcher, J. f/ecfroana/. Chem., 36, 137 (1972).
I
7
0.04mA
.1 ........._.. ,..
3
*
I
1 0.0
#
.j
’
I
I
-0.2
I
I
1
-0.0
-0.6 Ag
-0.4 V o l t s vs.
Figure 7. Cyclic voltammogram after oxidation of hydrogen in anhydrous acetonitrile, 0.1M TEAP, 0.050 V sec-’ sweep rate, 0.223 cm2 electrode area Dashed line shows growth of couple after sixth repeated cycle
In an effort to completely eliminate all possible sources of water, experiments were carried out using tetrabutylammonium fluoroborate as the supporting electrolyte. Cottrell and Mann (46) have implicated the following reaction as a source of water in anhydrous acetonitrile.
+
2HC104 C120: H2O (9) The resulting reduction wave was again a broad, irreversible wave with a peak potential of -0.330 V vs. SCE. Coulometric analysis of wet (10mM) HC104 solutions gives slightly low n values. A value of n = 0.92 was obtained in repetitive experiments. The low value could be due to formation of protonated basic species during the electrolysis at -0.2 V us. SCE.
ZIP+ CH3CN
2e-
-
+ Hz
H~
(10)
B
(11)
--et
+
B H+ -F== BHf (12) Finally, in our hands, exhaustive electrolysis of acetonitrile/HClO4 solutions followed by attempts a t product isolation produced unidentified tars and Liebig titration of electrolysis solutions indicated the absence of cyanide. These experiments indicate that acetonitrile is chemically reactive under relatively mild electrochemical conditions. The above attempts a t the production of protons in anhydrous acetonitrile failed but lead instead to a variety of basic species. The differences in the reduction potentials in Table I1 point to the overall complexity of the chemistry of the H+/Hz couple in acetonitrile. The base B may vary with the generation technique employed and the availability of water. It is interesting to speculate on the nature of the small couple at +0.19 V us. SCE in the cyclic voltammograms, Table I1 and Figure 7. If the reduction of the BH+ species (Equation 11) produces hydrogen according to the half-reaction
B P This compound has been identified in acetonitrile electrolysis solutions under acidic conditions by Clark et al.
I
+ e-
-
+ %Hz
B-
(13)
this process could be the anhydrous H+/Hz couple. This wave has a Ellz of -0.19 V us. ferrocene (log m ~ = H 6.4) which is consistent with Figure 4. Identification of Proton Waves in Cyclic Voltammograms. One of our initial objectives in undertaking this study was to use Figure 4 to distinguish “proton waves” from reduction waves of other oxidation products in cyclic (46) P. T. Cottrell and C. K . Mann, J. €lectrochem. SOC., 116, 1499 (1969),
ANALYTICAL C H E M I S T R Y , VOL. 45, NO. 7, J U N E 1973
1015
Table 1 1 1 . Oxidation of 2.7 X M p-Benzohydroquinone in 0.1 M TEAP/CH3CN/HzO Mixtures, Platinum Electrode (0.223 cmz) Peak potentials.aV vs. Sweep rate, V sec-'
0.010 0.050 0.200 0.104 0.104 0.104 0.104
Mol fraction of water
5x 5x 5x 0.14X 0.86 X 7.3 x 12.1 x
10-4 10-4 10-4
lo-' IO-* 10-2 10-2
SCE
EPa
0.905 0.945 0.96 0.94 0.91 0.88 0.91
O.0EBiC
0.01b -0.01 -0.00 -0.01 -0.01 -0.04
a Measured vs. Ag/Ag+ or SCE: see Experimental Section. b A second cathodic peak was barely discernible at -0.20 V vs. SCE. ? A second cathodic peak was barely discernible at -0.16 V vs. SCE.
voltammograms in acetonitrile solution. Anodic oxidation of many organic compounds produces species which give reduction waves in the region of 0.0 V us. SCE. Peak potentials as positive as ca. +0.2 V and as negative as -0.3 V us. SCE have been ascribed to proton reduction in the literature. Waves with peak potentials more negative than a t least +0.06 V us. SCE (at a sweep rate of 0.042 V sec-I) should more properly be assigned t o reduction of protonated species and not H + A N .This value represents the most positive peak potential we have observed for a bona fide proton reduction wave in acetonitrile/water mixtures.
The oxidation of p-benzohydroquinone, which has been well studied in acetonitrile (47, 48), is a-good example of the use of these data for the analysis of cyclic voltammograms. Cyclic voltammograms of the hydroquinone (QH2) exhibit a two-electron oxidation wave followed by a reduction wave in the region of 0.0 V us. SCE. The latter wave has been ascribed to reduction of protonated quinone, QH+ (47, 48). The results of the oxidation of p-benzohydroquinone in acetonitrile/water mixtures are given in Table ID. Note the behavior of the peak potential (EpC)of the "QH+ wave" as a function of water concentration. The peak potential varies very little ,as the mole fraction of water is increased, in marked contrast with the behavior of the proton wave in these mixtures (Figure 4). This behavior and the somewhat negative peak potential of this wave in dry acetonitrile (+0.01 V as compared to +Os@ V us. SCE) further rule out an assignment of this wave to simple H + A Nreduction. Received for review September 18, 1972. Accepted January 29, 1973. This study was supported in part by a grant from the Petroleum Research Fund administered by the American Chemical Society. J.A.L. was an N.D.E.A. Fellow, 1970-1972, University of Tennessee. (47) 6. R. Eggins and J. Q . Chambers, J. Electrochem. SOC.,117, 186 (1970). (48) V. D. Parker, Chem. Commun.. 716 (1969).
Sodium Tungsten Bronze as a Potentiometric Indicating Electrode for Dissolved Oxygen in Aqueous Solution P. B. Hahn, M. A. Wechter, D. C. Johnson, and A. F. Voigt Ames Laboratory-USAEC
and Department of Chemistry lowa State University, Ames, lowa 50010
Sodium tungsten bronzes, nonstoichiometric compounds, NaxW03, with 0.5 < x < 0.9, were found to respond potentiometrically to dissolved oxygen in basic solutions. A Nernstian response, with a slope of approximately 120 mV per decade, was exhibited in the concentration range 0.2-8 ppm. Oxygen analyses were made in this range with precision and accuracy approaching 3 ~ 5 % .The large slope and other observations place serious doubt on customary oxygen redox reactions as possible mechanisms. An absorption mechanism is proposed which involves the displacement of adsorbed hydroxide ions by molecular oxygen.
The utility of sodium tungsten bronzes, highly conducting nonstoichiometric compounds of formula Na,W03, as potentiometric indicating electrodes has recently been reported by Wechter et al. ( I ) who showed that the potential between such electrodes and reference electrodes indicated the concentration of reducible metals and the course of acid-base and redox titrations. References to earlier research on the bronzes as indicating electrodes and in fuel cells can be found in that paper. In the present work, cubic sodium tungsten bronzes (0.5 (1)
M. A.
Wechter, H. R. Shanks, G. Carter, G. M. Ebert, R. Guglielmino, and A. F. Voigt, Anal. Chem., 44, 850 (1972).
1016
A N A L Y T I C A L CHEMISTRY, V O L .
45, NO. 7, JUNE 1973
< x < 0.9) were used as electrodes for the potentiometric determination of oxygen in basic aqueous solution. The bronze electrodes show a Nernstian response over an oxygen concentration range of 0.2 to 8 ppm and a useful range from approximately 0.1 to 40 ppm with an unusually large concentration dependence, 120 mV/decade. The method differs from commonly used electrochemical methods ( 2 ) (polarographic or galvanic) for dissolved oxy'gen in aqueous media in that a potential rather than a current indicates the oxygen concentration. Other potentiometric methods have been proposed ( 3 ) but the dependence on oxygen concentration was the predictable 15 mvldecade based on redox reactions, and the sensitivity to changes in concentration was correspondingly much less than in these electrodes. The mechanism for the response of the Na,W03 electrode to dissolved oxygen in basic solution is not a t all obvious. A number of observations, however, have suggested that adsorption and desorption of cations and anions, especially OH-, a t the electrode surface play a more important role in the potentiometric response than any appreciable reduction of molecular oxygen. (2) J. P. Hoare, "The Electrochemistry of Oxygen," Wiley-lnterscience, New York, N.Y., 1968. (3) I. M. Kolthoff and H. A. Laitinen, "pH and Electrotitrations," Wiley. New York, N.Y.. 1941, pp 96ff.
