A STUDY OF THE ZIKC ELECTRODE BY F R E D E R I C K H. GETXIISN
In our previous studies of the copper' and cadmium2electrodes it has been shown that single crystals of each of those metals function as constant and reproducible electrodes when immersed in solutions of their respective salts. The investigations of Anderson3 and Straumanis' on the electrochemical behavior of single crystals of zinc have revealed a similar constancy and reproducibility in the electrode potential of that metal. Straumanis also pointed out that when a single crystal of zinc is split so as to expose definite cleavage planes to the solution in which it is immersed, the resulting difference of potential is identical with that which is developed when a polycrystalline electrode of pure zinc is immersed in the same solution. Having previously used single crystals of copper and cadmium to determine the normal electrode potentials of those metals, it has seemed of interest to determine the electrode potential of zinc in a similar manner. Relatively few determinations of the normal electrode potential of zinc have been made. Among the first to measure the electromotive force of zinc cells was Jahn5 who studied the galvanic system, Zn(amalgamated), ZnClz,(M) AgCI, Ag, in which the concentration of zinc chloride ranged from 0. j j 6 to 2 . 2 2 ?(.I.Cells of the type, Zn, ZnClZ(M), AgCl, Ag, AgCl, ZnCl@'), Zn, were investigated by GoodwinI6 the concentrations of the zinc salt ranging from 0.001to 0.2 XI. The first accurate determination of the normal electrode potential of zinc was made by Horsch' who made use of cells of the type, Zn, ZnCl@), AgC1, B g , the concentration of zinc chloride ranging from 0.0003 to 0.01RI. Basing his calculations upon the best available conductivity data for solutions of zinc chloride he found the normal electrode potential of zinc to be 0.758 i 0 . 0 0 2 volt at z j"C. Quite recently the electromotive force of a similar cell has been measured by Scatchard and Teffts at 2 j o with concentrations of zinc chloride ranging from 0.003 to I . j M. From their measurements they find for the cell Getrnan: J. Phgs. Chern., 34, Iqj.+ ( 1 9 3 0 ) .
* Getman: J. Phys. Chern., 35, j88 (1931).
3Anderson: J. Am. Chern. SOC.,52, 1000 (1930). Stmumanis: Z. phgsik. Chem., 47, 161 (1930). Jahn: \Vied. Ann., 2 8 , 21, 491 (18861. " Goodwin: 2. physik. Chern., 13, 577 (1894). Horsch: J. Am. Chem. SOC., 41, 1787 (1919). Scatchard and Tefft: J. Am. Chem. Soc., 52, 2272 (1930).
2 7 50
FREDERICK H. GETMAS
ZnHg(z phase), ZnClp(RZ),hgC1, Ag,
E, = 0.9834 volt. On combining this value with Cohen’s value of 0.0006 volt for the difference of potential betxeen pure zinc and the tn-0-phase amalgam1 and Scatchard’s value of - 0 . 2 2 2 4 volt for the potential of the silver-silver chloride electrode,? they obtained for the normal electrode potential of zinc, Zn, Zn++, E, = 0.;616 volt. In the present investigation the cells employed were set up according to the scheme Zn, ZnC12(M),Hg2C12,Hg, with concentrations of zinc chloride ranging from 0 . 0 0 2 to 1.0X. Relatively few measurements of the electromotive force of this particular galvanic combination appear to have been made.
Experimental Muten’als. The single crystals of zinc were prepared for the writer by Mr. J. H. Dillon of the physics department of the L-niversity of 11-isconsin according to the method which he has developed for the preparation of single crystals of metals of relatively low fusibility.3 The zinc used in making these crystals was of a high degree of purity. The polycrystalline electrodes were made either from very pure zinc supplied by Merck, or from a sample of so-called “spectroscopic” zinc kindly furnished the writer by kZr, H. RZ. Cyr of the research staff of the S e w Jersey Zinc Co. and guaranteed to be 99.99970 pure. The granulated zinc from which the zinc chloride was prepared was obtained from Merck and was of the grade supplied for forensic analysis. All of the other materials employed in making up the cells were prepared as described in our previous papers treating of the copper and cadmium electrode^.^ Preparation of Electrolyte. The mother solution of zinc chloride was prepared by dissolving Kahlbaum’s granulated zinc in a solution of pure hydrochloric acid. The latter was prepared by slowly dropping pure concentrated sulphuric acid into a solution of pure hydrochloric acid and absorbing the resulting gaseous hydrochloric acid in conductivity water until an acid of approximately 0.5 molal was obtained. An excess of granulated zinc was then added to the acid and the mixture warmed gently on the water-bath until the metal ceased to dissolve. The solution was then filtered to remove the remaining particles of zinc and diluted to a concentration approximating I 31. After transferring to a glass-stoppered flask, the solution was allowed to stand for some days in contact with a small piece of zinc to insure against excess Cohen: Z. physik. Chern., 34, 612 (1900). J. Am. Chem. SOC., 47, 2098 (1925). Dillon: Rev. Sci. Inst., 1, 36 (1930). LOC.cit.
* Scatchard:
A STUDY OF THE ZINC ELECTRODE
275'
acidity. The concentration of the solution was finally established by gravimetric determination of its chlorine content as silver chloride and volumetric determination of its zinc content by means of standard potassium ferrocyanide The concentration thus determined was found to be 1.0375molal. Apparatus. The apparatus employed was the same as that used in our previous studies of electrode potentials and the same experimental procedure was followed in setting up the cells and measuring their electromotive force. Electromotive Force Measurements The cells were immersed in an electrically heated and controlled thermostat bath, set at z j o i o.ozo, and sufficient time was allowed for the establishment of thermal equilibrium before any measurements were made. Readings were taken at frequent intervals oVer a period of six hours, it having been previously observed that no appreciable deterioration of the cells is likely t o occur during that interval of time. The tendency to erratic fluctuations in electromotive force which was noted in our study of the cadmium electrode was also observed with polycrystalline electrodes of zinc, but to a less degree. Single crystal electrodes of zinc were, in general, found to resemble single crystal electrodes of copper and cadmium in tending to give rise to a greater difference of potent ial than the corresponding polycrystalline electrodes. The difference, however, was of the same order of magnitude as the experimental error and therefore, no attempt has been made to distinguish between the potentials of the two types of electrode. Cells in which the concentration of the electrolyte was less than 0.00 j hI were found to deteriorate so rapidly that measurements a t lower concentrations could not be made with any degree of satisfaction. TABLE 1 E.M.F. of the Cell, Zn, ZnCl*(lI), Hg2C12,Hg(z j")
m (rnols ZnCl?/Iooo g. H20) 0.00200
E' 1.24497
E
-
0.00j 0 0
I. 2244
I . 2244
0.01000
I ,2 0 3 5 0
0,01999 o.03000 0 . 0 j003 0 . I0019
I . 18000
I . 2035 I . 1800
I . 16765
I . 1677
1.14713 I , 13077
I . 1285
0 . Ij047
I . 11717
-
0.20000
0.25148 0.35339
I . 1008j I . 08960 -
0.50000
0.50786 0.76770 I . 00000
1.0375
I
.oj816
I . 06489 -
1.oj679
1.1joo
I . IOjO -
I . 1078
1,0573 -
FREDERICK H. GETMAN
27 jZ
The experimental data are summarized in Table I, where m denotes the concentration in mols per 1000 prams of solvent and where E' is the electromotive force in volts. As has been pointed out, the differences in the electromotive force of cells with single crystal and polycrystalline electrodes were too uncertain to aarrant separate tabulation. The values aqsigned to E' In the table represent the mean of forty or more measurements of the electromotive force of a series of cells among which two contained single crystal electrodes.
The values of E' in foregoing table together with the experimental data of Jahn, Horsch, and Scatchard and Tefft obtained with cells of the type Zn, ZnC12(fiX),AgCl, Ag(2 j") are shown graphically in Fig. I . From a similar plot drawn on a large scale the smoothed values of the electromotive force, E, given in the last column were obtained.
Calculation of Results The electromotive force of the cell Zn, ZnClZ(M), HgzC'12, Hg, is represented by the equation
E
=
E.; - R T / n F ln(4m373),
A STUDY O F THE ZINC ELECTRODE
2i53
where E denotes the measured electromotive force of the cell, EL the normal electrode potential of the cell, m the concentration of the electrolyte in mols per 1000grams of solvent and y the activity coefficient. The symbols R,T, n and F have their usual significance. On simplifying and transforming to common logarithms ( I ) becomps
E
=
E: - 0.08873 log (1.588 my).
According to Randall’ equation l o g y - ~.E:
(2)
(2)
may be written in the form
+
= -
0.08873
0.2207
+ logm
1
,
(3)
If the right-hand side of equation (3) is plotted against the square-root of the ionic strength, pj, and the resulting curve is extrapolated to xero-concentration, the value of EA can then be readily calculated by means of equation (3). In order to employ this method, however, it is necessary to know the values of E corresponding to the region of smaller concentrations of the electrolyte where, as has been pointed out, satisfactory results cannot be obtained with the galvanic combination employed. In fact, the only available data for very dilute solutions of zinc chloride are those recorded by Horsch2 and even those are not of a high order of accuracy. Horsch points out that because of the lack of reproducibility of the electromotive force of cells in which the electrolyte is very dilute he was unable to extend his measurements below 0.0003 molal, and in concentrations below 0.001 molal the deviations of the individual observations from the mean was as much as 0.002 volt. Applying Randall’s method for the determination of EA to the measurements of Horsch, the data given in Table I1 are obtained.
m o 00034j8 0 0003995 o 000649 0 000772
o 0012j3 0 001453
TABLE I1 Data derived from Measurements of Horsch Pi E E o oS@j + o 2007 o 03241 0 03463 0 04413 0 04771 o 06131, o 06603
I
2jIO
11
I
11
I
2660 -7497 2440
I
2272
11
I
2219
11
1
7
log m
0660
I1 0 7 0 2
11
0969 1083 1286 13jo
On plotting the figures given in the fourth column against the square-root of the ionic strength, the curve is found to intersect the axis of zero concentration a t 11.039, On substituting this value in equation (3) we find EA = 0.9800. The potential of the silver-silver chloride electrode, as given in the International Critical Tables, is Ag, AgCl, C1-, E, = - 0 . 2 2 2 1 volt, Randall: Trans. Faraday Yoc., 23, jog 2
L O C . Clt.
(1927).
2754
FREDERICK H. G E T N A S
and therefore the normal electrode potential of zinc, as derived from Horsch's measurements, is Zn, Zn++, E, = o j i 9 \olt As has already been stated, Horsch computed the electrode potential to be E, = o 7;8 volt, making use of conductivity data If r+e accept E, = o i s 8 volt, as the normal electrode potential, ab derived from Horech's data by t x o different methods, the actility coefficients of zinc chloride may be calculated by means of equation ( z ) , and from these in turn the value of EA for the cell, Zn, ZnC'ln(bI),Hg?C12,Hg, can be computed as shown in Table I11
TABLE 111 Evaluation of E: for the Cell, Zn,ZnC'I? (AI), Hg2Cl2,Hg (Activity Data derived from Horsch) -
I.
0.8jj 0.79i
-
0.723 0.680 0.619 0.533 0,486
I .02j;
0.01
2605 23 j 6 I . 2108 I . 2244 I . 203 j
0.880
I.
0 . 0 0 0j 0.001
0.002 0.005
0.02
I . I800
0.05
I , 1500
0 .I
r.128j
(25')
I
,0281
I
,0286
1.0281 I . 0289
RIean
0.7j8
0.761 0.761 0 . ;61 0.761 0.760
The values of E, tabulated in the last column of the table arederivedfrom the corresponding values of EL in the preceding column by subtracting the potential of the mercury-mercurous chloride electrode. Hg, Hg2C12, C1-, E, = 0.2676 volt, as given in the International Critical Tables.. In their paper treating of electromotive force measurements with zinc chloride, Scatchard and Tefft' have computed the values of the activity coefficients from 0.0001to I . j molal. On substituting their values of y in in equation ( z ) , together with the corresponding values of E given in Table I, another set of values of EL and E, can be computed. In this manner the data recorded in Table IT.' have been obtained. The mean value of E, calculated from the activity data of Scatchard and Tefft is slightly greater than that derived from data based upon the electromotive force measurements of Horsch. Iloc. cit., p. 2281.
A STUDY O F T H E ZINC ELECTRODE
2755
TABLE IF' Evaluation of E: for the Cell, Zn, Zn C12(RI),Hg2C12,Hg ( 2 j") (Activity Data derived from Scatchard and Tefft) E m 7 E, E: 1,0279 0.767 0,760 I . 2244 0.OOj I . 0297 0.708 0.762 1.2035 0.01 0.642 I . 0300 I . 1800 0.762 0.02 I ,029; o.jj6 0.762 I . I joo 0.oj 0 .j 0 2 I . 0301 0.762 I . 128j 0.1 1.0319 I . 1070 0.2 0.448 0 764 I . 03 14 I . 0780 0.376 0,764 0.j I . 0309 0.32j 0,763 I .o 1,9573 Mean
0.762
In a discussion of the available methods for the extrapolation of electromotive force data to infinite dilution, Randall and Toung' have emphasized the experimental difficulties involved in securing reproducible measurements in extremely dilute solutions and have suggested that the solubility of the glass of which the cells are made is probably sufficient to vitiate all electromotive force data pertaining to extremely dilute solutions. To overcome this difficulty, Hitchcock* has recently proposed a method for the extrapolation of electromotive force data to unit ionic activity based upon a partially expanded form of the familiar Debye-Huckel equation. According to this method when EA - 0.1j jdF is plotted against where c denotes the molar concentration, the resulting curve should approach a straight line asymptotically in the region of small concentrations. Hence, if a straight line is thus obtained over the range of concentrations where experimental data are trustworthy, a linear extrapolation to zero concentration may be made with considerable confidence. Employing the interpolation formula c/m = 0.99707 - 0.0134m - 0.0129m2 0.0288 m3, derived by Scatchard and Tefft from the density tables given in the International Critical Tables, to compute the values of c corresponding to the values of m given in Table IF', and plotting the values of EA - 0.1jj a g a i n s t d g the value of EL a t infinite dilution was determined as shown in Fig. 2 . The value of EL, as read from the plot, is 1.300 and, therefore, the value of E, is E, = 1.0300- 0.2676 = 0.7624 volt. This value will be seen to agree substantially with the mean value of E, given in Table IV. Taking the average of the latter with the two preceding calculated values for E, we obtain as the normal electrode potential of zinc. Zn, Znii, E, = 0.7613 volt -
dm,
+
Randall and Young: J. Am. Chem. SOC., 50, 989 (1928). Hitchcock: J. Am. Chern. SOC.,5 0 , 2076 (1928).
2756
FREDERICK H. GETYAS
FlG. 2
Summary of Results (I)
The electromotive force of the cell Zn, ZnC12(M),Hg2C12,Hg
(25')
has been measured with concentrations of zinc chloride ranging from 0 . 0 0 2 to 1.0 11. (2) Electrodes of pure zinc in the form of both single crystals and polycrystalline aggregates have been employed. (3) The single crystal electrodes were found to resemble single crystals of copper and cadmium in exhibiting a tendency t o develop a difference of potential toward the electrolyte slightly greater than that developed by the polycrystalline electrodes. The differences, however, being of the same order of magni'ude as the experimental error no attempt has been made to differentiate between them. ( A ) The normal electrode potential of zinc has been computed from the electromotive measurements of the cell by three different methods as follows: (a) By substituting the values of y derived from the measurements of Horsch in the equation
A STUDY OF THE ZINC ELECTRODE
E
=
E'
- o.oSS73 log (1.588
2757
my),
the resulting value of E, being 0.760 volt; (b) By substituting the values of y derived by Scatchard and Tefft in the same equation, the resulting value of E, b e i s 0.762 volt; (c) By plotting E:- 0 . I j j d c a g a i n s t v'm,asproposed by Hltchcock, and noting the value of E, determined by the intersection of the rcsulting curve with the axis of zero concentration. This was found to be 0.762 volt. (;I The average of these three results has been taken as representins the probable value of the normal electrode potential of zinc as derived from electromotive force measurements of the cell employed, viz., Zn, Znu, E, = 0.7613 volt. Htllside Laboratory, Stamford, Conn.
