Electrochemistry of the Zinc-Silver Oxide System Part 1: Thermodynamic Studies Using Commercial Miniature Cells Mkhael J. Smith, University of Minho, Braga 4700, Portugal Colin A. Vincent' University of St. Andrews, St. Andrews, Fife KY16 9ST, Scotland The study of reversible electrochemical cells near equilibrium permits basic thermodynamic concepts such as maximum work and free enerw to be examined and understood. There have, however, he& practical difficulties in devising simple, safe lahoratory experiments that would permit large numbers of students to make measurements without the use of rather sophisticated devices such as hydrogen electrodes or hazardous materials such as mercury or corrosive solutions. Here we describe the use of readily available, sealed, miniature or "button" cells to obtain an accurate measurement of the enthalpy, entropy, and Gihhs freeenergy change of a chemical reaction. Of the many common commercial cell systems, most are unsuitable for fundamental studies because they have poorly defined or complex overall reactions. For example in the Leclanch6 (zinc-manganese dioxide) cell, the reaction scheme depends on both the rate and the "depth" (i.e., percentage) of discharge, and on the precise formulation of the electrolyte ( I , 2). A further important requirement for thermodynamic studies is that hoth electrodes act as truly reversible svstems. I t is also advantaeeous to avoid cells with liquid junctions that introduce some uncertainty into the measurement of emf (esneciallv in the case of variable tem, the majority of cell sysperature experiment.& ~ i n a l l i in tems, the emf is affected by the concentration (activity) of solution species, which greatly complicates hoth the emf measurement and interpretation of the experimental data (3). One of the most suitable cells for student experiments is the zinc-silver oxide alkali cell, which is written formally as
and for which anodic and cathodic discharge processes may he assumed to be
and AgzO(s)+ HzO(l)+ 2e
-
2Ag(s) + 20H-(aq)
is 9-10 mol dm+ NaOH or KOH giving solely ZnO(s) as anodic product (6). In general the cathodic process is straightforward in zinc-silver oxide cells. However, miniature ;ells from certain manufacturers have modified cathode compositions. For example, the sharp voltage cutoff on approaching complete discharge is unattractive for certain practical applications, so that a mixture of AgzO with MnOz may he used, resulting in some changes in the emf. (Cells are also manufactured which contain a mixture of Ago and Arlo to eive hieher ca~acitv. .. hut these are usuallv confieu;Ld so that t h i electrode potential is determined by tge latter (7).)Thus where comparison of thermodynamic data obtained from the cell with literature values is required, i t may he necessary to select cells from only particular manufacturers.
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Experimental Commercial button celh in the capacity range 35-190 mAh are readily available. Here we report on experiments carried out on D350 (Duracell UK) 100 mAh cells. A detailed description of the construction of these cells is given in I'an 2 (81. It rr important to make rehable electrical contact with [he terminals The deaian of two different rell holders. constructed from 10mm plastic sheets, is suggested in Figure 1. Both versions gave satisfactory results. The emf must be measured with a precision of *0.1 mV: this is mast readily achieved hy use of a digital voltmeter. (As each measurement is very rapid, such an instrument may be shared by many students.) There are a number of ways of thermostating the cell. A simple method is to immerse the cell holder in a test tube of paraffin oil. The test tube may then be heated in a beaker of water to, say, 40-50 ' C and allowed to cwl. A thermometer is retained within the test tube: provided that cooling is relatively slow, no serious temperature eouilibration madients arise. We have studied cells in which lone ~. . was carried out at a numher of tiaed temperatures using n thermostat, and also using heating and roolmg cycles in which the temperature wax varied a t rates 4 , f up to I O C pep minute. No significant difference was found in any of these experiments, although it was noted that it was simplerto achieve acontrolled temperature change during the cwling part of the cycle. The results of a typical experi~
~
respectively, so giving the overall spontaneous cell reaction,
Themost important feature of this reaction is the invariance of the composition of the electrolyte solution and the chemical potentials of the reactants and products as the discharge nroceeds. It should he nointed out. however. that in manv practical zinc-silver oxihe cells, the' anodic rdaction may be much more c o m ~ l e x14.5). l'hus in certain "wet" cells usinr a NaOH solution with concentration of about 6 mol dm-< soluble zincate species, especially Z n ( O H ) P , are initially formed, and the normal end-product is Zn(OH)n(s). For the miniature cells recommended here, however, the electrolyte
a
' Author to whom correspondenceshould be addressed.
Flgur. 1. Two posslbls types of cel holder eonrtrumo from 10-mm plastrc meet In (a) e sctr cal contact Is maos oy means of bra39 screw; in (b) phosphor bronze springs are used.
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Table 2. Glbba Frw Energy, Enthalpy, and Entropy Changes ZnO(s) Assoelated wHh the ReactlonaZn(s) Ag,O(s) 2Ag(s)
+
Experlmerdal
Flour0 2. Varlatlon of emf for a Duracell D350 mlnlature cell as a function of temperature. ment collected over a 30-min perid in which a cell cooled from 32 'C to 17 ' C are shown in F i r e 2. Results and Discuulon The emf, E, of a reversible electrochemical cell is related to the Gibbs free energy -~ change for the cell reaction, AG, by the expression
Calculaaxl
+
Hllls (10)
These calculated data are compared in Table 2 with the values obtained from the electrochemical cell: a very satisfactory agreement is obtained. There is also good agreement with the electrochemical values obtained from large commercial cells (Yardney International Corporation) by Hills (10). ~--,-
It is clear that the energy available from the zinc-silver oxide cell is primarily a result of the relatively large heat of formation for zinc oxide in comparison with the notably small value for silver oxide. This variation may he seen in turn to arise from a number of factors from the following Born-Haber cycles (values given in kJ):
where n is the number of electrons transferred and F is Faraday's constant. Further, since
determination of the temperature coefficient of the cell emf leads to the evaluation of A S and AH for the reaction. In the case of the zinc-silver oxide cell, reactants and products are all in their standard states, so that determination of Ee, the standard emf and (dEe/dT)p leads to A P , AHe, and ASe. Data derived from specific heat measurements published in the literature (9) for the cell reactants and products are given in Table 1.From these we have for the cell reaction:
= -317.5 kJ mol-'
Similarly,
Hence,
Table I. Standard Heats ofFormation and Standard Entropies OMalned from Calorlmetlc Data (9)
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Journal of Chemical Education
Zn(s) Zn(g) 12638 Zn2+(d
-348
!E
1
+ %Odd _,ZnO(s, -4036
02-(~)
Despite the high enthalpy of vaporization of silver, the heat of formation of 2Ag+(g) is considerably smaller than that of Zn2+(g).However, the lattice energy of ZnO(s) is over 1000 kJ mol-' more negative than that of AgpO(s), and this is therefore seen to he the critical term responsible for the relative stability of zinc oxide. Sherman (11) calculated lattice energy values of 2445 kJ mol-I and 4084 kJ mol-I for Ag*O(s) and ZnO(s), respectively, using an electrostatic model. The latter value agrees well with that derived from the Born-Haber cycle. Agreement for AgzO(s) is poor, reflecting the presence of significant covalent bonding in this com~ound. ~ i i v e oxide r has one of the highest standard entropies of all solid triatomic oxides, and since it is destroved during the cell reaction, ASe and hence (dE/dT), have significant, gegative values. Operation of this cell, even near to reversible conditions, thus results in the evolution of heat. However, as we will discuss in alater paper (8),themagnitude of the heat output associated with irreversible operation is usually much more significant. Conclusions We have examined the variation of emf with temperature for a numher of types of commercial miniature cell. One of the most suitable for deriving thermodynamic parameters is that based on the zinc-silver oxide couple, which gives results that are in good agreement with those derived from thermal measurements.
1. Cahoon.N.C.In TklfimryBoItery;Cahah,N.C.;Heise,C. W..Eds.; Wi1ey:New Ymk. 1976:Vol. 2. 2 Vincent, C. A,; Banino, F.: Laerari, M.: Semsati,B. MnlemBotteries;Amold: London. ISM. -~~~
3. ViocentC. A. J. Chem. Edue. 1970.47.386368. 4. Tye. F. I.InEkcfrockmicd Poww Sources; Bar& M., Ed.: Peregrious: Stove-, U.K.,1980.
5. FIei8ehu.A.: Lauder. J. J. ZincSilmr Ozidr Batteries; Wiky: NesvYork. 1911. 6. Sehumacher, E. A. In TheRimary Boltery; Heiae.0. W.; Cahwn, N. C.. Eda.:Wiley: New York, 1971:Vol. 2. 7. Schimizu. A,: Nerani, Y. In Handbook 01Sattrriea and Fuel Cella; Linden, D.. Ed.: McCraw-Hill: NeaVork, 1984. 8. 8mith.M. J.;Vineent.C.A. J. Chom.Edue.,inpress. 9. N.B.S. Cireuhr SOO.1952. LO. Hills. S . J ElecLmrhsm. Sm. 1961,108.81C4.11. 11. Shermsn J.Chem.Rau. 1932.11.95-170.
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