J. N. PWW AND LEE
n. OUOH
[COXTREiUTION PEOM TEE PHYSICAL ClrEMlSTRY LA3OEATOP.Y OF THB STATB UNIVERSITY OP IOWA]
The Potential of the Silver-Silver Ferrocyanide Electrode at 25' J. N.PEARCE'
AND
LEE D. O U G H ~
Tb el&rQde potentials3 of met& in solutians tightly. ,4ir was displaced from the cells with of their ions as well as certain non-metallic elec- nitrogen. As quickly as possible, silver ferrocyatrodes bave been investigated. Among the latter nide mixed with silver crystals was added to the may be mentioned the calomel and the silver- silver crystals surrounding the silver-plated platisilver chloride electrodes, the potentials of which num wire and pressed in place. Solutions mere have been determined with a high degree of preci- forced into the cells under nitrogen pressure. sion. Electrodes of ari organometallic nature, Half-cells of equal molality, connected by salt however, have been studied less extensively. In bridges, were short circuited to hasten equilibrium this research we have undertaken to determine the between the electrodes Following this procedure potential of the silver-silver ferrocyanide elec- the salt bridges were removed and a modified Lamb and LWSOII~ flowing junction device intrade, M;at&als.-Electrolytic silver crystals were serled, connecting half-cells of unequal molality. used in the preparation of the silversilwr ferro- After msasuremats were made with this set-up cyanide electrades. the flowing junction apparatus was removed and Potassium ferrocyanide was twice precipitated the asodgam reservairs put in place. The amalfram 8 atered, saturated aqueous solution by the gwn r w v o i r s were w d e d out by alternately addition of absolute alcohol. The salt, dried evaowting and filling witb nitragen severd times spontaneously, was dissolved in oxygen-free water before filling them with potassium amalgam. and kept under nitrogen in the dark. The silver-dwr ferrocyanide electrodes wuld Silver ferrocyanide, prepared immediately be- be duplicated to 1 mv. Duplicate half-cells were fore Use in the cells, was precipitated in an atmos- found to differ by d l . 5 mv. Measurements of phere of nitrogen fram a diluted portion of the po- the cells with ion-transference would remain contassium ferrocyanide stack solution with tttl exass stant to within 0.5 mv. for a period fram two to of dilute silver nitrate solution. The precipitate four hours. As many measurements of the cells was washed by decantation until no turbidity was without ion-transference as were possible were detectable upon t h e addition of sodium chloride taken in the time (ten to fifteen min.) necessary solution. for the amalgam reservoirs to empty. The electroAll solutions were prepared on the weight molal motive force measurements of four c& are given basis with canduc water. All other chemi- in Table I where E and $, reprmnt the observed by approved methods. potentials of cells without ion-transference and cals used were Metastrrements were made of the electromotive with ion-transfeivnce, respectively. force of the concentration cells TABLE r Ag \ i$g$Fe(CN)a, g4%(cK), (0 1 m) \ (0.05m)&%(CN)a,
AgBe(CN)e I A% Ag
ELECTAOIH)TSVS Faaiercs,OF SILVSR-S~VBR PERROCYANIDE C~NCENTRATIQR CELLS AT 25'
(0.1 m) 1 KHg I (0.03 PZ) KJ?e(CN!o, Ag&(CN)G 1 Ag
1 AgSe(CN)a, K,Fe(CNfa
E, volt
E$, volt
0.0208
.OB1
0.0131 .0134
.02m .0194
,0137 ,0156
The electrode vessels used isi these experiments were of the same type as those described by Har-
ned.4 The electrodes were prepared by placing dried s,ihrer crystals upon a mat of glass wool held by the constriction in the electrode vessel. A silver plated platinum wire electrode was inserted and around this the silver crystals mere packed
.-
_*__
.0€39 AK.
.(.I207 AY.
Measurements of concentration cells afford a method for ohtahing t h q m e a transference numher af the c a k b €or the molalities used. The werage values for E and E, from Table I were substituted in the formula
-
~ ( 5 ) l,aruh
_
_
,tad I2dI'wn, i h l
ne = E L ' E 1 42, 121 (1'JLlI)
Jan., 1938
POTENTIAL OF THE SILVERSILVER FERROCYANIDE ELECTRODE AT 25b
where 12, is the mean transference number. The value obtained for nc is 0.673, Electromotive force measurements were made also of cells of the type Ag I AgaFe(CN)s, K4F4CN)dmlS I (m)KCl, HgZClz I H g
81
TABLE111 POTENTIALS OF TEE W X E L AND SILVER FERROCYANIDE ELECTRODES AT 25” Ezps E m calomel m, EZWAg, m, KCI calomel Harned’a y* K&(CN)r AgrFe(CN),
0.2 .1 .05
0.3180 .3337 .3498
0.3177 .3337 .3498
0.2167 .2196 .2241
0.05 .025 .0125
The electromotive force values obtained are the mean values of two to seven cells in which each half-cell contained two to five electrodes. The in- electrode in which the potassium ferrocyanide modividual electrodes of the calomel half-cells came lality is 0.05. Following a reasoning similar to to equilibrium quickly and were constant to within that of MacInnes,8the relation d0.03 mv. The individual electrodes of the ferEL Et(1 - 1/5n0) rocyanide half-cells reached equilibrium in from is obtained where EL, Et, and n, have their usual six to forty-eight hours and differed by about 0.5 significance. The value calculated for EL,using mv. The duplicate hdf-cells differed by about 0.672 as the cation transference number, is 0.0098 *3 mv. in 0.0125 m and by *0.3 mv. in 0.06 m v. and is subtracted from the observed potential solutions. of the cell. The cell potential corrected for liquid Liquid junction potentials were calculated by junction then is 0.0041 v. The potential, &a, for means of the Lewis and Sargent’ formula the 0.1 m electrode is 0.2126 v. EL = 0.05915 log Xi/X, In order to evaluate for the silver4lver where AI and A2 are the equivalent conductames ferrocyanide electrode, a method similar to that of ferrocyanide and chloride solutions’ in which the employed by Lewis and Randall9 was used. The potassium ion molalities are equal. In the range potential, EO, is given by the equation: of the concentrations used the liquid junction poE = E” - 0.01479 log U F ~ ( C N ) ~ - (1) tential adds to the true potential of the cell; there- Subtracting 0.01479 log W Z F ~ ( C N )from ~-both sides fore, its value must be subtracted from the obser- of (1) and rearranging, the relation ved potential of a cell in order that the single elec- 0.01479 log ( ~ / M ) F ~ ( c I N ) ~ = - - E” trode potential may be obtained. The experimen(E 0.01479 log ~ F ~ ( c N ) ~(2) - ) tal data are compiled in Table IT. is obtained. At infinite dilution log TABLEI1 is equal to zero and therefore EiBB = E 0.01479 ELECTROMOTIVE FOXCES OF GALOMEU-FERROCYANIDE log mFc(cN),--. If the right member, EO’, is plotted CELLSAT 25’ against any function of rptp,(cN),-- the intercept rn, ILFe(CN)e 0.05 0.025 0.0125 m = 0 is equal to E;9B. m,KCI .2 .1 .05 Plotting the EO’values against the square root of .1210 .1310 the molality groups the points so that a straight Eow., V. .lo96 .0069 .0053 EL,v. .0083 line more nearly fits the data. Using the method .1141 .1257 .lo13 E 2 9 8 , v. of averages,L0the equation To obtain the potentials of the ferrocyanide Eo’ 0.1942 f O.OlZ697fi (3) electrodes in the various solutions, the values of is obtained for the straight line. The constant, Epss are subtracted from the potentials of the cor0.1943 v., is the value for the standard electrode responding calomel electrodes. The values for the potential of the silver-silver ferrocyanide electrode latter were calculated using E!BI= 0.2676 v. for at 25’. the electrode Hg,Hg,Cls,Cl-(u = 1). These were By substituting for Bo’its value in (3), the rechecked using Harried's* mean activity c o d lation cients for the respective molalities of potassium E = 0.1943 0.012597G - 0.01479 log mFa(CNj8-chloride solutions. Table I11 contains the data (4) covering the foregoing. is obtained. Using this equation, single electrode The single electrode potential of the 0.1 m elecpotentids were calculated for several molalities of trode catl be evaluated from the data for the cells (8) MacInnes, THIS JOIJBNU, 87, 2301 (1915). with ion-transference and the potential for the (9) Lewis and Randall, “ T h e r m o d y n a h , ” McGraw-Hill Book
+
-
+
-
+
(6) Lewis and Sargent, THIS JOURNAL, 81, 362 (1909).
(7) “International Cdtlcal Tabh,” 19211, Val. VI, pp. 234 and 253.
Co , Inc., 336 Weat 424 St.,New Ywk, N. Y.,1923,p. 234. (10) Daniels, “Mathematical Preparation for Physial c h u m i s try,” McGtar-Hill Book Co., New York, N. Y., 1918. p. M b .
WILLIAM C. BRAY
82
Vol. 60
potassium ferrocyanide. From the calculated is obtained. Solving for s, the value 2.27 X electrode potentials the activities of the ferrocy- the molal solubility of silver ferrocyanide in water, anide ion were computed using equation ( I ) , The is obtained. activity coefficients of the ferrocyanide ion for the TABLEIV several molalities were determined by dividing the aFe(CN 6-m x Y-Fe(CN)&-- a&+ x lo‘p K a x 10‘‘ actil-ities by the corresponding molalities. 0.01 0.8143 0.81 2.089 1.552 In addition the activities of the silver ion can be .02 1.518 .76 1.786 1.546 calculated from the relation .04 2.742 .69 1.542 1.549 E2s8= 0.79iX + 0.1)5916 log .06 b.686 .61 1.429 1.537 where 0.7978 v. is the standard electrode potential .cW 4.585 .57 1.355 1.546 .1(1 5 35; .54 1.303 1.546 for the silver electr~de.~The activities of the ferrocyanide and silver ions being known, the activity 1.546Av. product constant €or silver ferrocyanide aAgd
h, =
(aig
summary
(UFc(Lh)b’-)
From a series of measurements at 2 5 O , the pot ential of t he electrode Ag ,AgaFe(CN)6, Fe (CN ) (u = 1)is found to be 0.1943 v. The activity product constant and the molal solubility in water of silver ferrocyanide, 1.546 X and 2.27 X respectively, have been calculated.
can be evaluated. In Table I V calculations are brought together under headings which are self explanatory. If s is the molal solubility of silver ferrocyanide, 4s and s are the molal solubilities of the silver and ferrocyanide ions, respectively. Taking the solubilities equal to the activities, the equation (4\)1
x
= 250,; = 1.546 >! 10-41
IOWA CITY, IOWA
[CONTRIBUTION FRO31 THE CHIEMICAL
LABORATORY OF TXiB
RECEIVED S’OVEMBER 2, 1937
UNIVERSITY OF CALIFORNIA]
The Interaction of Ozone and Hydrogen Peroxide in Aqueous Solution B Y WILLIAM
In 1917 Rothmund and Burgstaller’ concluded that hydrogen peroxide at low concentration is an efficient catalyst €or the decomposition of ozone 209 = 302
and that HzOz
4-0 8
= HzO
(1)
+ 202
(2)
is the net reaction when the peroxide is present in large excess. The latter result is well supported by their four quantitative experiments1 (p. 297), y the work of others. Proof of the catalytic action of hydrogen peroxide depends on the results of their three rate experiments in the presence of dilute sulfuric acid; they determined the concentrations of both ozone and peroxide, and calculated the ratios of decomposed ozone to decomposed peroxide. These ratios decreased rapidly during each experiment. In spite of evident inaccuracies in the experimental data, it will be shown in this paper that the results are in fair agreement with the relation
(1) Rothmund and Burgstaller, itiionatih., SB, 295-303 (1917).
c. BRAT
where V I and v2 are the rates of reactions (1) and (2), respectively. The two rates and the two corresponding concentrations are the values at a given instant during an experiment. At high and low values of the concentration ratio, the limiting stoichiometric results are reactions (1) and ( 2 ) , respectively. It is evident, therefore, that ozone a t low concentration is not a catalyst for the decomposition of hydrogen peroxide in dilute acid solution. Weiss2 recently has plotted the Rothmund and Burgstaller ratios of the “mean consumption values,” AOa/AH&, against the ratios of the “corresponding mean values of the concentrations.” Equation (3) was not revealed by this non-differential method; but the results did indicate #at relative rates and relative concentrations might be related, and led me to estimate actual instead of average rates. Table I contains the results of calculations made with the assistance of Mr. E. L. Derr. In expt. I -d(H200n)/dt could not be determined ( 2 ) Tosfph Weieiss, Trans. Faraday Soc , S1, 668-681 (1935).
