THE MERCUROUS ACETATE ELECTRODE W. D. LARSON Department of Chemistry, College of S t . T h o m a s , St. Paul, M i n n e s o t a AND
F. H. MACDOUGALL School of Chemistry, University of Minnesota, M i n n e a p o l i s , M i n n e s o t a Received September 8, 1936 ISTRODUCTION
The object of this study was to determine whether the electrode H g I Hgz(OAc)z, OAc- would function as a reversible acetate electrode. One test of such reversibility is the constancy of the standard E.M.F. of t h e electrode when i t is measured a t various acetate ion activities. The cell chosen for this study was
Pt
1
HOAc, (Satd. with quinhydrone)
HOAc , Hgz(OAc)z I H g (Satd. with (solid) Hgz(0Ac)z)
I
This cell was chosen because it is simple to make u p and measure. The quinhydrone electrode is known to be reliable in solutions whose hydrogenion concentrations are of the order of magnitude of those in acetic acid solutions. Previous studies of the mercurous acetate electrode were not made with the purpose of evaluating its standard E.M.F. nlernst (9) studied the cell H g [ Hgz(OAc)z, KOAC 1 KNO, 1 KOXc, Hgz(O-4c)z 1 H g
I1
in a study of concentration cells. Bugarszky (3) studied the cells: H g I Hgz(OAc)z, KOAc KOAc 1 KOAc, KBr , HgzBrz I H g (solid) 1 . O M 1.OM 1 . O M 0.01M (solid)
111
Hg I Hgz(O=lc)z, KOAc KOAc 1 KOAc, KI , Hg& I H g (solid) l . O M 1 . O M 1.OM 0.01M (solid)
IV
Hg 1 Hgz(OAc)z, KOAC 1 KOAC 1 KOAC, K O H , HgzO 1 Hg (solid) 1 . O M 1 . O M 1 . O M 0.OliM (solid)
V
493
494
W. D. LARSON AND F. H. MACDOCGALL
At 18.5OC., he found for cells 111, IT-, and T' electromotive forces of - 0.2474, - 0.4277, and -0.2626 volts, respectively. From these data, it is possible to calculate the standard E.1i.F. for the mercurous acetate electrode, E & z ( ~ ~ cif ) the z , standard E.BI.F.'S for the mercurous bromide, iodide, and oxide electrodes are known, as well as the activity coefficients of the salts involved. The standard E.M.F.'S of the mercurous bromide and iodide electrodes are = -0.1385 v., and @zgz12 = +0.0416 v. (6). The standard E.N.F. of the niercurous oxide electrode, Ekg20,equals -0.13 v. (1). T'aluesfor the activity coefficients of the salts must be estimated, since data for 0.01 inolar potassium bromide, potassium iodide, or potassium hydroxide in 1 molar potassium acetate solutions are not available. Randall, licBain, and White (10) give the value 0.699 for the molal activity coefficient for 1.0 molal pota.;'s-iuni acetate; this is an average value between 0" and 90°C. For potassium bromide and potassium iodide in 1 molal solutions, Harned and Douglas (4) give y = 0.613 and y = 0.649, respectively, at 25°C. At the same temperature, Scatchard (11) finds y = 0.743 for 1 molal potassium hydroxide. I n the calculations, we may assume that the activity coefficients of potassium bromide, potassium iodide, or potassium hydroxide in the solutions in which they are 0.01 molar and in the presence of 1.0 molar potassium acetate are the same as the coefficient for 1.0 molar potassium acetate; on the other hand, we may assume that the activity coefficients are the same as in 1.0 molar potassium bromide, potassium iodide, or potassium hydroxide. Yeither of these assumptions is very likely to be correct. Using the first alternative, we obtain from cells 111,11-,and l-.,E $ g n ( ~=~ c ) z -0.5016, -0.5018, and -0.508 volts, respectirely. By the second means, there result the values -0.5049, -0.5040, arid -0.507 volts from the same three cells. The data of Bugarszky therefore indicate a value of about -0.50 volt for the standard E.M.F. of the mercurous acetate electrode. MA4TERISLS AiiD APPARATUS
Grasselli C.P. glacial acetic acid was used. I t n as refluxed with chromic oxide (2 g. to 100 cc. of the acid) and clihtilled. The mercury used was a coniniercial product. I t was washed firbt by allowing it to fall from a capillary tip through a l-meter column of diluted nitric acid (1 part of acid to 4 parts of m-ater), and then in the same way with distilled water. I t n a c next dried by allowing it to run through a filter paper with a hole punched in the apes of the cone. T h e mercury was then distilled in a current of air under reduced preswre. 3Iercurous acetate was made by precipitation from approximately onetenth normal solutions of niercurou+ nitrate arid potaqsiuni acetate. Preparations were made in which the mercurou. nitrate contained excess nitric acid, and when it did not. They gave identical result< in the cells
495
THE MERCUROUS ACETATE ELECTRODE
The salt was washed by decantation, collected 011 a Buchner funnel, airdried, and stored over anhydrous calcium chloride. The quinhydrone was prepared by the method of Biilmann and Lund (2) and was recrystallized from water at 60°C. The water uqed in the preparation of the solutions waq of “conductiyity” grade. For the E.M.F. measurements, a Leeds and Sorthrup type li potentiometer, a Leeds and S o r t h r u p type R galvanometer, and an Eppley cell of the unsaturated type were used. The constant-temperature water bath was a t a temperature of 25.00 i0.03”C. The temperature was determined by comparison with a thermometer calibrated by the U. S. Bureau of Standards. The cell-vessels were of Pyrex glass. They were of the customary type. The cell-vessel uqed for the mercurous acetate electrodes had a stopcock in the side arm to prevent diffusion of mercurous ions into the solution containing quinhydrone. Junction between the two half-cells was made in a small beaker outside the thermostat. Bright platinum electrodes were used in the quinhydrone half-cell. TABLE 1 E x p e r i m e n t a l data C
E
C
E
The cells were measured in duplicate, and were constant in E.M.F. over a period of a t least six hours, when measurements were discontinued. Constancy was usually attained within a n hour after making up the cells. Preliminary experiments showed t h a t if the quinhydrone half-cell were renewed each day, the E.M.F. of the cell was constant over a period of several days. The E.M.F. of duplicate cells never differed by more than 0.15 niv. The acetic acid solutions were analyzed volumetrically with carbonatefree sodium hydroxide, using phenolphthalein as indicator. The solutions more concentrated than 0.15 molar were diluted in calibrated volumetric flasks, and aliquot portions of the diluted solutions were analyzed. E X P E R I M E S T A L DATA
Table 1 gires the experimentally observed data; C represents the moles of acetic acid per liter of solution, arid E is the electromotive force. The mercury pole wai: the positive one.
496
W. D. LARSON AND F. H . MACDOUGALL
CALCULATION O F T H E STANDARD E.M.F.
The values of E in table 1 can be converted to the value of the of the cell,
E.M.F.
I
(Pt) 1 H z , HOAc, HgZ(OAc)z(s) Hg
VI
P = 1 atm.
E=, by adding to them 0.6992 volt, the value of the standard E.M.F. of the quinhydrone electrode (5). The E.M.F. of the left half-cell of cell T'I is given by
where E.M.F.
is the activity of hydrogen ions in the acetic acid solution. The of the right half-cell is given by
a1
ER = E
o
+ RT log, a? F
where a4 stands for the actirity of acetate ions in the acetic acid solution saturated with mercurous acetate. If a2represents the activity of acetate ions in the acetic acid solution, and a3 that of hydrogen ions in the acid solution saturated with mercurous acetate, the liquid junction E.M.F. between the two solutions will be given by
Ei
=
RT al RT a4 n, - log, - + n, - loge F a3 F a2
(3)
where n, and n, represent the transference numbers of cation and anion, 1 a4 = -; and respectively. NOW,approximately ala2 = a3a4, so that aa3
Ex
=
EL - EB
a2
+ E, = - Eo - RT - log, a3a4 F
(4)
or, since = UH+
. aoLko-'v -
aHOAc
a3a4
C(1 - a )
if the activity coefficient of undissociated molecules is unity, we have
=
E
+ 0.6992
(5)
497
T H E MERCUROUS ACETATE ELECTRODE
I n table 2 are given results of calculations of Eo from this equation. The column headed EO' gives values of
)
= - (E
+ 0.6992 +
and Eo gives
Values of a were approximated from the relation a =
&
TABLE 2
E Qcalculated jrom equation 6
c 0.05002 0.07485 0.09956 0.1497 0.2094 0.3021 0.4138 0.4992 0.7509 0.9943 1.495 2.058
I
EH
E" ~
0.8643 0.8557 0,8487 0.8393 0,8307 0.8228 0.8146 0.8096 0.7993 0.7923 0.7818 0.7735
-0.5062 -0.5079 -0,5082 -0,5093 -0.5103 -0.5108 -0.5107 -0,5105 -0.5107 -0,5109 -0.5107 -0.5108
I
EQ
-0.5067 -0.5083 -0,5085 -0.5096 -0.5105 -0.5110 -0,5109 -0.5107 -0.5108 -0.5110 -0.5108 -0.5109
where Kl is the dissociation constant for acetic acid, and has the value 1.8 X (7). K i n equations 6, 7 , and 8 is the thermodynamic dissociation constant, and at 25°C. has the value 1.753 X (8). The difference between the values in the third and fourth columns of table 2 gives the value of the term
aRT F The E o values calculated in this way are uncertain in the more dilute solutions, since no allowance for the solubility of the mercurous acetate has been made. The effect of this solubility will be less at higher acid concentrations becauFe in this case the solubility is smaller, and because the contribution of acetate ions from the mercurous acetate is relatively less.
498
TV. D. LAKSON AND F. H. MACDOCGALL
Apparently at an acetic acid concentration of 0.3 molar and above, t h e effect of the solubility of inercurous acetate upon the degree of dissociation is less than the cxperimcntal error in the measurements. .Accordingly, t h e data in this range arc considered to girc the true value of EO. SVMMIAHT
1. The electroniotivc forw3 of the cell Pt
I
HOAc
,
HOAC
,
(Satd. with (Satd. with quinhydrone) Hgz(0Ac)z)
Hg,(OAc)z 1 Hg (solid)
have been measured at 25"C., at concentrations of acetic acid varying from 0.05 molar to 2.0 molar. 2. From these data the standard E.M.F. of the electrode Hg 1 Hgz(O.\c)z, OAc- has been found to be -0.5109 f 0.0002 volt a t 25OC. REFERENCES (1) Abhandl. deut. Bunseii Ges. S o . 5, p. 198 (1911). (2) BIILMANN, E., A N D LUND,H.: h n . chim. 16,321 (1921). (3) BUGARSZKY, S.: Z. anorg. Chem. 14, 145-63 (1897). (4) HARNED, H. S., ASD DOUGLAS, S. 11.:J. Am. Chem. Soc. 48, 3095 (1926). ( 5 ) HOVORKA, F., .4SD DEIRING,IT.:J. -k!tl. Chem. SOC.6 7 , 446 (1935). (6) International Critical Tables, F-01. V I , p. 332. RIcGraw-Hill Book C,o., Kew York (1929). ( 7 ) See MACDOUGALL, E'. H . : Physical Chemistry, p. 513. The Macmillan Co., New York (1936). (8) MACINNES, D., A N D SHEDLOYSKY, T . : J. Arn. Chem. s o c . 64, 1429 (1932). (9) NERNST,JT.: Z . physik. Chem. 4, 129-81 (1889). (10) RANDALL, LI,, ~ L C B A I S , J., .4ND I ~ H I TAE, :, J. Anl. CheIIl. S O C . 48, 2520 (1926). (11) SCATCHARD, 6 . :J . -kin, Chrni. Sur. 47, 648 (1925).
