In the Laboratory
Zinc Electrodes and the Thermodynamics of a Galvanic Cell Donald A. Probst and Giles Henderson Department of Chemistry, Eastern Illinois University, Charleston, IL 61920
A popular physical chemistry experiment in which students measure the temperature dependence of a galvanic cell potential has been described by McSwiney (1). ∆G, the change in free energy for nF coulombs discharged through a potential E, is dependent on both ∆H, the change in enthalpy, and ∆S, the change in entropy, in accordance with ∆G = {nFE = ∆H { T∆S
(1)
where T is the absolute temperature. Thus ∆G for a galvanic cell reaction is obtained at a given temperature from the measured potential, while ∆S is given by the slope of a plot of E vs. T. Finally, ∆H is calculated from ∆G and ∆S and can be converted to ∆H° with appropriate corrections for enthalpies of dilution (2). Conversion to standard state permits students to compare their results with ∆H° obtained from data for standard enthalpy of formation. The galvanic cell described previously (1) is based on the reduction of iron(III) by zinc: Zn(s) + 2Fe(CN)63{(aq) → Zn2+ (aq) + 2Fe(CN)64{(aq) (2) This cell reaction exhibits a large potential (>1 V) and a large entropy change. Moreover, it employs simple and inexpensive electrodes: Zn|Zn2+(1.0 M)||Fe(CN)63{(0.1 M), Fe(CN)64{(0.1 M)|Pt (or C)
These electrodes are more convenient than the mercury or lead amalgams cited by other authors (3). However, our students find that when monitored for several minutes, the cell potential ramps steadily downward (Fig. 1). This process is found to be ongoing whether or not the cell is connected to the measuring circuitry. Moreover, the initial voltage is not restored upon replacement of the electrolyte or salt bridge solutions. The initial potential is restored after drying and polishing the zinc electrode or by rinsing the zinc electrode in dilute acid. If the zinc elec-
trode is replaced with a less active metal such as Pb, the cell potential is stable and well behaved. These observations suggest that the downward ramping in cell potential results from the electrical resistance of a surface coating formed by some reaction(s) of the zinc electrode with the surrounding solution. Aqueous corrosion of zinc is well known and has been extensively reported (4). Hoxeng and Prutton (5) have also reported potentials of zinc in open beakers of various aqueous solutions that decrease with exponential time constants of approximately 1 hour. Finally, we note that a fine, cloudy precipitate often forms during preparation of Zn2+ electrolyte solutions. Filtrates are found to be weakly acidic and the cloudy solutions clear when a few drops of acid are added. We wish to understand this behavior and identify appropriate modifications of McSwiney’s experiment that will permit stable EMF measurements. Experimental Procedure Our students prepare electrodes and the electrolyte solutions in accord with reference 1. We have also used a “lead” pencil carbon electrode as an alternative for the more expensive platinum electrode (6, 7). A carbon electrode also eliminates the use of mercury, which is commonly used as a low-resistant junction to a copper potentiometer lead. Anode and cathode half-cells can be constructed from test tubes equipped with Vycor glass tips1 to isolate the half-reactions from the surrounding salt bridge electrolyte. Students use ice and thermostated temperature baths with a Fluke model 8860A digital multimeter to measure the cell potential at 5 different temperatures ranging from 0 to 60 °C. Calculations We use parameters listed in Table 1 to calculate equilibrium conditions depicted as zero free energy surfaces in the (pH, [Zn2+], T) domain for two important zinc reactions: Zn(s) + 2H2O(l) → H2(g) + Zn2+ (aq) + 2OH{(aq)
(3)
Zn2+(aq) + 2OH{(aq) → Zn(OH)2(s)
(4)
and
The surface coordinates are obtained by incrementally changing the temperature and zinc ion concentration over a desired grid and at each point, calculating the pH for which the corresponding reaction is at equilibrium. The equilibrium conditions of reaction 3 are depicted as the upper surface in Figure 2. In this case it is convenient to obtain the equilibrium conditions from standard potentials: Figure 1. When the zinc electrolyte is at ambient pH, the galvanic cell potential is unstable and decreases with time.
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Journal of Chemical Education • Vol. 73 No. 10 October 1996
∆G = {nFE° { RT ln Q = 0.0
(5)
In the Laboratory
where Q is the usual product/reactant activity ratio. The lower surface of Figure 2 identifies the pH, [Zn2+], and temperature at which the solubility product equilibrium is satisfied. Conditions described by a point above this surface correspond to a supersaturated solution in which reaction 4 is spontaneous in the forward direction. Reaction 4 is spontaneous in the reverse direction for conditions described by a point below this surface. Thus a neutral 1.0 M Zn2+ solution at room temperature (depicted as the upper, circled point) can precipitate Zn(OH)2(s) until the pH is reduced to 5.94 (lower solid point), consistent with the cloudy precipitate and acidic filtrate described above.
Open Circuit Chemistry The zinc electrode can be oxidized by reacting with dissolved oxygen (4): 2Zn(s) + O 2 + 2H2O → 2Zn2+(aq) + 4OH {
(7)
{ The increase in [Zn2+][OH ]2 may then exceed K
·
Zn + 2H2O → H2 + Zn2+ + 2OH{ {upper surface} Zn2+ + 2OH{ → Zn(OH)2(s) {lower surface} pH = 4 (buffered electrolyte) unbuffered electrolyte
Figure 2. ∆G = 0 surfaces depict the [Zn2+], T, and pH at which two important zinc reactions are at equilibrium. The direction of reaction spontaneity is determined by whether the reaction conditions are described by a point above or below the respective ∆G = 0 surface (see text). The circled point characterizes typical conditions in which a zinc electrolyte may be initially prepared. These calculations predict a subsequent spontaneous precipitation of Zn(OH)2 (s), resulting in a saturated condition depicted as the lower surface. The zinc anode is protected from surface poisoning at pH = 4.0, depicted by the solid line.
Table 1. Parameters Used with Equations 5 and 6 To Calculate the Zero Free Energy Surfaces in Figure 2 (9) ∆H (reaction 4)a ∆H (ionization of
28.4 kJ H2O)a
57.7 kJ
Ksp298
Zn(OH)2a
7.0 x 10-17
Kw297
(ionization of H2O)
1.0 x 10-14
E° (reaction 3) a
-0.0653 V
Calculated from ∆ H° formation values ( 9).
Knowing E°, R, T, and the temperature dependance of the ionization constant for water (8), the [OH{] (and pH) for which Q = K and ∆G = 0.0 can be easily calculated. The equilibrium conditions for reaction 4 are depicted in the lower surface of Figure 2. The equilibrium constant for this reaction (the inverse solubility product constant) is well known (7) and varies with temperature: d ln K = ∆H° dT RT 2
Results and Discussion
(6)
sp and can result in the precipitation of Zn(OH)2(s). It is known that Zn(OH)2 can form a very adherent film or coating on a zinc surface (9), which can effectively protect zinc from further corrosion. Zinc can also be oxidized by reacting with water (eq 3). Reaction 3 is spontaneous in the forward direction for all points below the upper surface of Figure 2 and is spontaneous in the reverse direction for points above this surface. The forward reaction produces an increase in [Zn2+][OH{]2 and as in the case of reaction 7 can also result in the formation of a Zn(OH)2(s) coating on the surface of the zinc electrode. The formation of a strongly adhering surface film protects the zinc from further oxidation, impeding its function as an anode.
Closed Circuit Chemistry If the zinc electrode is coupled through an external circuit to the Fe(III)/Fe(II) half-cell, zinc is oxidized in accordance with eq 2. If the [Zn2+] is increased from the equilibrium electrolyte conditions described by the lower solid point in Figure 2, the formation of a Zn(OH)2(s) coating on the surface of the zinc electrode is predicted. All three of the above processes can contribute to poisoning the anode surface if the cell is operated at ambient pH. Schikorr has reported (10) that if solid zinc hydroxide is in contact with atmospheric carbon dioxide, further reaction can produce a very stable basic zinc carbonate, ZnCO3?3Zn(OH)2. A similar process might be expected in a solution containing dissolved CO2. Figure 2 suggests that if the Zn2+ electrolyte is maintained at a pH below the lower surface boundary, a nonconducting electrode surface coating of zinc hydroxide or basic zinc carbonate cannot form. We find that the downward ramping of the cell potential is indeed eliminated by simply buffering the Zn2+ electrolyte at pH = 4.0 (solid line in Fig. 2). Since Fe(CN)63{/ Fe(CN)64{ solutions are unstable under acidic conditions, the cathode electrolyte was left unbuffered. The small junction potential that arises from differences in half-cell/salt bridge pH is neglected. Under these conditions, students obtain stable measurements throughout the desired temperature range. With a carbon electrode we observe cell potentials that are within ± 0.002 volts of those measured with a platinum electrode. The use of a buffered anode electrolyte permits a reliable study of the temperature dependence of the cell potential. Representative data are depicted in Figure 3.
Vol. 73 No. 10 October 1996 • Journal of Chemical Education
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In the Laboratory
Acknowledgments The authors wish to acknowledge C. Robert Leidner for acquainting us with Vycor and Jonathan Blitz for helpful suggestions. Note 1. Porous Vycor rod, tubing, and tips are available from Bioanalytical Systems, Inc., 2701 Kent Avenue, West Lafayette, IN 47906.
Literature Cited 1. McSwiney, H. D. J. Chem. Educ. 1982, 59, 165. 2. Hamer, W. J., Ed. The Structure of Electrolytic Solutions; Wiley: New York, 1959; pp 136–142. 3. See for example: Daniels, F.; Williams, J. W.; Bender, P.; Alberty, R.; Cornwell, C. D. Experimental Physical Chemistry, 6th ed.; McGraw–Hill: New York, 1962; pp 208–211; or Shoemaker, D. P.; Garland, C. W.; Steinfeld, J. I.; Nibler, J. W. Experiments in Physical Chemistry, 4th ed.; McGraw–Hill: New York, 1981; pp 240–244. 4. A comprehensive bibliography of zinc corrosion appears in: Slunder, C. J.; Boyd, W. K. Zinc: Its Corrosion Resistance, 2nd ed.; International Lead Zinc Research Organization: New York, 1986. 5. Hoxeng, R. B.; Prutton, C. F. Corrosion 1949, 5, 330–338. 6. Selig, W. S. J. Chem. Educ. 1984, 61, 80–81. 7. Arena, J. V.; Mekies, G. J. Chem Educ. 1993, 70, 946–947. 8. Lide, D. R., Ed. CRC Handbook of Chemistry and Physics, 71st ed.; CRC: Boca Raton, 1990–1991. 9. Porter, F. Zinc Handbook; Marcel Dekker: New York, 1991; pp 97–127. 10. Schikorr, G. Werkstoffe Korrosion 1964, 15, 537–543.
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Figure 3. The temperature dependence of a galvanic cell potential is used to fully characterize the thermodynamics of the cell reaction. The potential of a zinc anode if found to be well behaved when buffered at pH = 4.0 as suggested by Figure 2.
Journal of Chemical Education • Vol. 73 No. 10 October 1996
