T H E OXIDATION POTENTIAL OF T H E SYSTEM POTASSIUM MOLYBDOCYANIDE-POTASSIUM MOLYBDICYANIDE, AND THE EFFECT OF NEUTRAL SALTS ON T H E POTENTIAL I. M. KOLTHOFF School of Chemistry, University oj Minnesota, Minneapolis, Minnesota AND
WM. J. TOMSICEK
Department of Chemistry, College of St. Thomas, St. Paul, Minnesota Received June $0, 1936
Collenberg (1) determined the oxidation-reduction potentials of molybdocynnide and molybdicyanide solutions. His measurements, however, were made in the presence of potassium chloride, the concentration of which was not known. In the preparation of his solution of potassium molybdicyanide an excess of potassium chloride was used to decompose precipitated silver molybdicyanide in order to form the soluble potassium molybdicyanide and insoluble silver chloride. His data do not allow an extrapolation of the potential to an ionic strength of zero. For this reason we have determined the oxidation-reduction potentials a t increasing dilutions in order to find the potential at an ionic strength of zero. In addition the effects of neutral salts on the potential were studied. MATERIALS USED
Potassium molybdocyanide This product was prepared according to the method of Olsson (5) and recrystallized three times from water by the addition of ethyl alcohol. The water content of the air-dried product was determined by heating in the electric oven at a temperature of 105°C. for five hours. The loss in weight from two samples gave 1.993and 2.006 moles of water, respectively. Two other samples were titrated at the same acidity as in the potentiometric method described below, with a permanganate solution standardized against Kahlbaum’s sodium oxalate, and the results agreed with the formula K&Mo(CN)*.2HzOwithin 0.5 per cent. Erio-grun was used as an indicator in the permanganate titration, since the golden-yellow color of the potassium molybdicyanide formed masks the permanganate endpoint. The volume of permanganate used to obtain the endpoint with erio-griin was found to be identical with that found in the potentiometric method. 247
248
I. &I. KOLTHOFF AND WM. J. TOMSICEK
In figure 1 the change in potential is given in the potentiometric titration of a mixture of 100 ml. of 0.007 molar molybdocyanide and 2 ml. of concentrated sulfuric acid with 0.1 normal permanganate, the saturated calomel electrode being used as the reference electrode. Aqueous solutions of potassium molybdocyanide are relatively stable, as shown by the fact that the titer was constant for a t least two days. In our work fresh solutions were prepared every second day.
flL 0lONKn~0+
FIG.1. Titration curve for KIM,(CN)~
Potassium molybdicyanide Solutions of potassium molybdocyanide, acidified with sulfuric acid, were oxidized with potassium permanganate until a pink color persisted. An excess of silver nitrate was then added to precipitate silver molybdicyanide. The precipitate was filtered off on a Buchner funnel and washed until no test for the silver ion was obtained in the washings. The moist silver molybdicyanide was suspended in water and shaken with somewhat less than the equivalent quantity of potassium chloride, leaving some silver molybdicyanide undecomposed. The filtered solution was used as a stock solution, its concentration being determined by electrometric titration with a standard solution of potassium ferrocyanide.
249
OXIDATION POTENTIAL
Fieser (2) used the same procedure for the determination, but does not state the conditions for the titration. We found good results in neutral or weakly alkaline medium, but no distinct jump was observed in acid medium. This is easily explained by the fact that the oxidation potential of the ferrocyanide-ferricyanide system increases much more with increasing hydrogen-ion concentration than that of the molybdocyanidemolybdicyanide system. Figure 2 shows the change of potential in the
AOID BOLUTIONB
Di5uselight
19 3
~
NEUTRAL 80LUTION8
Dark
Di5use light
l 3
23 4
I
:;
BAS10 SOLUTIONS
Diffuse light
29.3
1
Dark
3.3
250
I. M. KOLTHOFF AND WM. J. TOMBICEK
1 per cent in six days. Hydrogen ions have a stabilizing influence on solutions of potassium molybdicyanide. The effects of acetic acid and sodium carbonate on the decomposition of potassium molybdicyanide are shown in table 1. The light to which the solutions were exposed was diffuse sunlight in a north room of the laboratory. TABLE 2 Oxidation-reduction potentials of potassium molybdocyanide-potassium molybdicyanide solutions at 86°C. RATIO MOLYBDI
X,MO(CN)~
TOTAL #
x
0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0
03012 003012 001205 000602 02696 00539 002696 001080 00054 000270 000108 02112 00422 00211 00084 000422 000211
EQ.E.
E, .a. E.
-
0.2155 0.1800 0.1697 0.1638 0.2131 0.1870 0,1776 0.1675 0.1615 0.1577 0.1553 0.2045 0.1786 0.1688 0.1585 0.1826 0.1485
0.7919 0.7564 0.7461 0.7402 0.7895 0.7634 0.7540 0,7439 0.7379 0.7341 0.7317 0.7809 0.7550 0.7452 0.7349 0.7290 0.7249
0.7908 0.7553 0.7450 0.7391 0.7911 0.7650 0.7556 0.7455 0,7395 0.7357 0.7333 0.7903 0.7644 0.7546 0.7443 0.7384 0.7343
___
YOLYBDO Y
0.02888 0.002888 0.001155 0.000577 0.02872 0.00574 0.002872 0.001150 0.000574 0.000287 0.000115 0.03044 0.006088 0.003044 0.001218 0.000609 0.000304
1.0430 1,0430
1,0430 1,0430 0.9390 0.9390 0.9390 0.9390 0.9390 0.9390 0.9390 0.6940 0,6940 0.6940 0.6940 0.6940 0.6940
0.4695 0.04695 0.01878 0.00938 0.44896 0,08974 0,04489 0.01798 0.00898 0.00449 0.00180 0.43212 0.08622 0.04311 0.01725 0.00862 0.00431
0.6852 0.2167 0.1371 0.0968 0.6701 0.2996 0.2119 0.1341 0.0950 0.0670 0,0424 0.6574 0,2936 0.2076 0.1313 0.09285 0.06566
$0
f b refers to the potential of a n equimolecular solution of molybdocyanide and molybdicyanide calculated from the figures in columns 3 and 7. EO.=. is t h e E.M.F. as measured against the quinhydrone electrode in a solution being 0.01 N in hydrochloric acid and0.09 N in potassium chloride. E x . H Eis. the E.M.F. referred t o the normal hydrogen electrode as calculated from EQ,x,(see ref. 4).
EXPERIMENTAL PROCEDURE
To a solution of potassium molybdicyanide, prepared as described above, was added a weighed quantity of potassium molybdocyanide. The solution was then analyzed for molybdocyanide and molybdicyanide by titration with potassium permanganate and potassium ferrocyanide, respectively, and used as a stock solution for the dilution measurements. All flasks and apparatus used to contain the solutions were coated with black lacquer and kept in the dark. The measurements were made in a darkened room. For details regarding the experimental technique reference is
251
OXIDATION POTENTIAL
made to a previous paper (3). Table 2 gives the results of measurements of various dilutions of three stock solutions containing different ratios of molybdocyanide and molybdicyanide. By “Total F” is meant the sum of the ionic strengths of molybdocyanide and molybdicyanide. The values of e: found in table 2 were plotted against 4;on large cross section paper and extrapolated to determine the value of the normal potential at zero ionic strength. The extrapolated value was found to be 0.7260 volt, referred to the normal hydrogen electrode. The curve is shown in
0.79 076 0.77
0.76
on 5 0.74
0.73
0.71
0.70
o
ai
az
0.3
a4
0.5
0.6
0.7
ae
9 FIG.3. Change of normal potential
e: with increasing ionic strength p . a = experi menta1 data (table 2); b = calculated vaIues
figure 3. The straight line in the figure represents the values of lated on the basis of the limiting Debye-Htiokel equation:
e:
calcu-
where E a = 0.7260 volt, and f3 and f 4 are the activity coefficients of the molybdicyanide and molybdocyanide, respectively. Since e: refers to the value for equimolecular concentrations of molyb-
252
I. M. KOLTHOFF AND W. J. TOMSICEK
docyanide and molybdicyanide, then the calculated value of e:
= eo
6;
is
+ 0.0591 log j73
in which, according to the limiting Debye-Huckel expression, log
e
=
f4
3.5 &. TABLE 3 Log fa - i n dilute solutions
of molyhdo-molyhdicyanides in the presence of various
electrolytes
f4
KC1 0 0 0 0 0 0
01 025 05 1 25 5
0.01 0.025 0.10 0.5
0 1365 0 1833 0 2421 0 3295 0 5080 0 7132
0 4068 0 5695 0 7322
!I I1 j ~
LOG
i 4
HC1 0 0104 0 0522 0 10
NHiCl
0.1317 0,1797 0.3276 0.7123
0 0 0 0
0 1075
LiCl 0.01 0 025 0 1 0.25
0.1317 0.1797 0,3276 0.5073
0,1317 0.1797 0.3276 0.7123
1343 1850 3390 7448
0 3796 0 6305 0 7779
0 4136 0 5813 0 9237 1 4525
CaC12 0 01 0 025 0 1 0 5
0 3813 0 5085 0 8085 1 0593
/I
SrCh 0.01 0.025 0.1 0.5
0 1380 0 2466 0 3296
NaCl
0 8288
k.
~
1317 1797 3276 7123
0 6100 0 8050 1 1322 1 5340
BaC12
0 01 0 025
0 6186 0 8390 1 1814 1 6814
0 0 0 0
00 51
0 0 0 0
1317 1797 3276 7123
0 6610 0 8813 1 2203 1 7085
The effects of various salts and of hydrochloric acid on the potential of dilute solutions of molybdocyanide-molybdicyanide are shown in table 3. Instead of reporting the measured values of the E.M.F., the figures are given
253
OXIDATION POTENTIAL
for log J2 calculated from the equation: f4
cMoCnr-+ 0.059 log CMoCn8----.-
E =
fa f4
where E is the measured potential referred to the normal hydrogen electrode and eo is the potential a t infinite dilution (0.7260 volt). For the sake of brevity the composition of the very dilute molybdocyanidemolybdicyanide solution is omitted. Since these solutions were not
1.8
1.7 1.6
15 1.4 1.3 12 1.1 10
09
08 07
06 05 04
03
at
az
a3
a+
a5
06
07
0.8
os
sil
-
FIQ.4. Ratio of activity coefficients as a function of ionic strengths. a HCl; b = LiCl; c = NaCl; d = KCl and NHdCl; e = CaCll; f = SrCl,; g BaCls; h = calculated values. stable, fresh solutions had to be prepared for each set of measurements. 1/; represents the square root of the total ionic strength of the mixtures, whereas p a designates the ionic strength of the added salts. The values of log f-3 me found plotted against 1/;L in figure 4. The f4
straight line in figure 4 represents the theoretical values of log f-3 calculated f4
on the basis of the limiting Debye-Hiickel equation. The effect of various anions on the potential of a dilute solution of
254
I. M. KOLTHOFF AND W M . J. TOMSICEK
molybdocyanide-molybdicyanidewas investigated, using potassium chloride, potassium bromide, and potassium nitrate a t a concentration of 0.25 molar, respectively, in the mixtures. The values of e: in the presence of the three salts were calculated from the measured E.M.F. as described above, and gave 0.7965, 0.7957, and 0.7950 volts, respectively. These results are in accordance with the results obtained in solutions of ferrocyanideferricyanide, the various univalent anions having virtually the same effect on the potential. An attempt was made to determine the effect of sodium hydroxide on the potential of dilute solutions of molybdocyanide-molybdicyanide. Two solutions were measured, one containing 0.01 molar and the other 0.1 molar sodium hydroxide. The solutions were exceedingly unstable and no constant readings could be obtained. The potentials decreased very rapidly, but initial readings agreed with the values for eo obtained with hydrochloric acid. DISCUSSION O F T H E RESULTS
The oxidation-reduction potential of the molybdocyanide-molybdicyanide system, starting a t an ionic strength of zero, changes with increasing ionic strength markedly different from that of the ferrocyanideferricyanide system (3). In the latter case it was found that up to an ionic strength of 0.04, the normal potential was greater than that calculated on the basis of the limiting Debye-Huckel expression. With molybdocyanidemolybdicyanide, however, the experimental values were lower than the calculated ones a t all ionic strengths. This behavior is normal in dealing with a system, the potential-determining ions of which are of such a high valence type. On the other hand, it should be stated that even a t the greatest dilutions the slope of the experimental curve is decidedly less than that of the line calculated on the basis of the limiting Debye-Huckel expression. For this reason no attempt has been made to calculate average ionic sizes using the more extensive Debye-Huckel equation. As was to be expected, neutral salts were found to increase the oxidationreduction potential of dilute molybdocyanide-molybdicyanide solutions to a very large extent. With the univalent cations the effect decreases in the order K+ = NH4+ > Na+ > Li+ > H+. In all these cases the values of log f-3 calculated from the experimental data were less than those f4
derived from the simple Debye-Huckel expression. In the case of ferrocyanide-ferricyanide, however, the experimental figures a t lower ionic strengths were less than the ones calculated from the Debye-Huckel equation. The divalent cation salts, a t the same ionic strengths, have a much greater effect, the latter decreasing with decreasing ionic size: Ba++ >
OXIDATION POTENTIAL
255
Sr++ > Ca++. At lower ionic strengths the experimental figures for log
fsf 4 were found to be greater than those calculated on the basis of the simple Debye-Hiickel expression. Hydrogen ions have an effect similar to that of other univalent cations, the former being comparable to the lithium ion. From this behavior one may infer that both molybdocyanic and molybdicyanic acids behave as strong electrolytes. I n this respect molybdocyanic acid is entirely different from ferrocyanic acid, since it has been found that the fourth ionization constant of the latter is equal to 5.6 X 10-6 (4). This explains why the oxidation-reduction potential of the ferrocyanideferricyanide system increases much more with increasing hydrogen-ion concentration than that of the molybdocyanide-molybdicyanide system. SUMMARY
1. The normal potential of the molybdocyanide-molybdicyanide system was extrapolated to an ionic strength of zero and found to be 0.7260 volt a t 25°C. 2. The effect of various salts upon the oxidation-reduction potential a t various ionic strengths haszbeen determined. The univalent anions investigated have the same effect at corresponding ionic strengths. 3. At the same ionic strengths divalent cations exert a greater effect than univalent cations, the effect decreasing with decreasing size of the ion. 4. From the effect of hydrogen ions upon the potential, it is concluded that molybdocyanic acid is a strong electrolyte. (1) (2) (3) (4) (5)
REFERENCES COLLENBERG, 0.: Z. physik. Chem. 109, 353 (1924). FIESER,L. F.: J. Am. Chem. SOC. 62, 5204 (1930). I. M., AND TOMSICEK, WM. J.: J. Phys. Chem. 39, 945 (1935). KOLTHOFF, KOLTHOFF, I. M., AND TOMSICEK, WM. J.: J. Phys. Chem. 39, 955 (1935). OLSSON:Ber. 47, 917 (1914).
