W. C. VOSBURGH AND SUMTER A. COGSWELL
2412
[CONTRXBVlXON FROM TEE
DEPARTMENT OF
c€iEMISTRY OF DUKE
Vol. 65
UNIVERSITY)
Complex Ions. VIII. Pyridine-Silver Ions BY W. C. VOSBURGH AND SUMTERA. COGSWELL It is of interest to compare the complex ions formed from silvex ion and pyridine with the ammine-silver ions.' Koch2 and Britton and William~~ have determined the dissociation constant of the ion Ag(C&,N)t+ on the assumption that this is the only pyridine-silver ion. Under their conditions a monopyridine ion would be unimportant. The constants are not in very good agreement, but great accuracy is not claimed. In this investigation the solubilities of silver iodate, silver sulfate and silver bromate in pyridine solutions have been measured. Preliminary calculations of the iodate results gave rapidly increasing dissociation constants for the most dilute solutions when it was assumed that only the dipyridine complex ion was formed. Assumption of the existence of the ion AgPy+ as well as AgPyzf, where Py stands for C6H6N, led to a satisfactory interpretation of all of the solubility data. Materials.-Silver sulfate and silver iodate were prepared by precipitation as previously described, except that the iodate was dried at 110'. Silver bromate was prepared by slow precipitation from silver nitrate and potassium bromate. Analysis gave 45.70% silver; calculated 45.75%. Pyridine of "practical" grade was refluxed over barium oxide for several hours and distilled through a Widmer column. A fraction having a boiling-point range of 0.2' was taken. Pyridine solutions were usually made by weighing the pure pyridine and water. In some cases the weighed pyridine was diluted to known volume, and the density was determined so that the molality could be calculated Other chemicals were of analytical reagent grade. Solubility Determinations.-The solubility of silver iodate in pyridine solutions of known molality was determined essentially as previously described for the solubility in m m ~ n i aexcept ,~ that the pyridine was not determined. The temperature was 25 0.05'. The results are shown in the first two columns of Table I. The silver ion molality was calculated from the iodate ion molality (assumed equal to the total iodate molality), the ionic strength, and the solubility product for silver iodate4 (p. 2672). The difference between the total silver and silver ion molalities was taken as the sum of the molalities of two complex ions, AgPy+ and AgPyz'. Because of the small solubility of silver iodate, the f
(1) (a) Bjerrum, "Metal Ammine Formation in Aqueous Solution," P. Haase and Son. Copenhagen, Denmark, 1941, p. 130; (b) Vosburgh and McClure, THTS JOURNAL, 65, 1060 (1943). (2) Koch, J . Chrm. Soc., 2083 (1930). (3) Britton and Williams, ibid., 798 (1935). (0 Derr, Stockdale s n d Vonburgh, TEISJOWXNAL, 63,2672 (1841).
TABLEI SOLUBILITY OF SILVER IODATEIN PYRIDINESOLUTIONS PY, total m
AgIOs ~n X 103
0.0497 0547 0371 0997 1013 1075 1493 1498 2018 2061 3026 409 409 516 516 .728 853
1.082 1 180 1 232 2 084 2 130 2.244 3 089 3.122 4 173 4 189 6 27 8 46 8 47 10 66 10 72 15 05 17 62
Ag+ ?E
X 105
3.064 2.819 2 704 1637 1.603
1.526 1.128 1 117 0.852 .849 584 445 445 ,361 360 267 232
AgPye+
Py,free
m
Kg X 1@
0 905
0.0477 .0525 .OM9 .0957 .0973 1032 .1433 ,1437 1936 ,1979 .2903 .3924 .3925 .494 ,495 ,698 818
7.70 7.73 7.72 7.84 7.74 7.85 7.94 7.82 7.99 8.28 8.08 8.28 8.28 8.42 8.37 8.75 8 90
m X 10'
1.004 1.057 1.911 1,960 2.072 2.916 2.950 4.00 4.01 6.10 8 28 8.29 10.48 10.54 14.86 17.43
second is more important than the first. Therefore the data of Table I were used for the calculation of Kz,the equilibrium constant for the reaction
+
AgPy,+ _r Ag+ 2 4 Kz = [ A ~ + ~ [ P Y I * / [ A ~ P Y ~ + J (1)
Activity coefficients should cancel approximately from Equation 1. To correct for the relatively small amount of the monopyridine ion, AgPy+, the instability eonstant, k1, of this ion, determined as described below, was used. k,
=
[Ag+][Py]/[AgPy+] = 1.0 X lo-*
(2)
Simultaneous equations were set up allowing the calculation of the molalities of free pyridine and the ion AgPy2+. These molalities are given in Table I, together with the corresponding values of KZ. The solubilities of two somewhat more soluble silver salts, the sulfate and the bromate, in pyridine were determined to obtain more favorable data for the calculation of kl. For silver sulfate the procedure was essentially as described previousiy.lb For silver bromate the procedure was like that for iodate except that ammonium molybdate was used as a catalyst in the titration.s The results for silver sulfate are given in Table 11. As the ion AgPy+ was present in larger molality than the ion AgPyz+, the molality of the latter was calculated by means of Equation 1with Kz = 7.5 X 10-5. Since kl was needed for the calculation of KS,an approximation method in(5) Kolthoff and Sandell, "Quantitative Inorganic Andy&." The MacmiUan CQ, New Y o t k , N Y , 1943,P. 624
PYRIDINESILVER IONS
Dec., 1943
volving also the data of Table IV was used for finding mutually consistent values for the two constants. The molalities of the ion AgPy+ and the uncombined pyridine could then be calculated, and finally kl. TABLE I1 SOLUBILITY Py, total m
OF SILVER SULFATE IN PYRIDINE SOLUTIONS AgtSO4 Ag+ AgPy+ P y , free m X 10: m X 10: m X 101 m X 10: &I X 10:
2413
TABLE IV SOLUBILITY OF SILVER BROMATE IN PYRIDINE SOLUTIONS Py, total m
0.00597
.00827 .01273 .01564
AgBrO: m X 10:
9.80 10.32 11.49 12.34
Ag+ AgPy* Py, free m X 10: m X 10: m X 10:
7.08 6.76 6.14 5.76
2.09 2.44 3.22 3.72
2.63 3.59 5.25 6.20
&I
X 101
8.9 9.9 10.0 9.6
centrated solutions. However, they show only erratic variations for total pyridine molalities up to 0.1 m,and the value adopted, KZ= 7.8 X lo-? is the average for the more dilute solutions. The constants of Britton and Williamsa increase with the pyridine concentration also. Their average Before attempting to measure the solubility of value, determined at 15O, when corrected to 25' silver bromate in pyridine solutions, its solubility by the temperature coefficient data of Kochz is in water and potassium nitrate solutions was de- 7.0 X 10". Koch's value for K2 seems to be termined. The solubility of silver bromate has about 11 X a t 25O, but is a little uncertain been measured by Hill,EReedy,' and Dalton, Wey- because the temperature for the accepted average mouth and Pomeroy,S the results being 8.29, 8.31 value is not given. and 8.06 millimoles per 1000 g. of water, respecIt is probable that the increase in K2 is the eftively, a t 25'. The first two agree within the pre- fect of some uncontrolled variable which varies cision of the second, and the result in this investi- appreciably only in the more concentrated solugation was only a little lower, 8.27 X m. tions. The dielectric constant of the medium The solubility results with this preparation were might be one such variable. Another possibility is more consistent among themselves than those ob- the activity coefficient of the pyridine, which was tained with a preparation of lower solubility in a taken as unity in all calculations. previous investigation.1b The effect of potassium The assumption of the existence of two complex nitrate on the solubility agreed with that found pyridine-silver ions gives a satisfactory interpreby Dalton, Weymouth and Pomeroy. The re- tation of the solubility data, whereas the assumpsults are given in Table 111, and can be calculated tion of the dipyridine ion alone was unsatisfacby means of the equations tory. This is to be expected by analogy with the ammines. K8.p. = [Ag+][BrOs-]p = 5.60 X 10-6 (3) It is interesting to compare the constants for the -1ogf 0.50561/F/(l 4) 0.104~ (4) pyridine complexes with those for the ammines. TABLE I11 The constant kz may be defined in terms of the aLUBILITY OF SILVER BROMATE IN POTASSIuM NITRATEconstants used above by the equation klkz = Kz. The ratio of k1 to ka for pyridine is 1.3 as comSOLUTIONS pared with 3 for ammonia. The monopyridine KNO:, m AgBrO;, m P %P. X 10' ion is more important relatively than the monam0.00827 0.00827 5.61 mine. The ratio of kr for ammonia to kt for py0.0477 .00949 .0572 5.59 ridine is 4 X while the ratio of the ioniza.loo2 .01027 .1105 5.60 tion constants of the two bases is about .401 .01304 .414 5.62 This comparison emphasizes the lack of relationThe solubility of silver bromate in pyridine ship between basic strength and stability of silver solutions is given in Table IV. The calculation complex ions which was pointed out by Britton of the silver ion molality was made by means of and Williams. Equations 3 and 4 and that of kl by the method Summary used for the silver sulfate data. The value of k~ The solubilities of silver iodate, silver sulfate used in the calculations for the silver iodate data (Table I) was the mean of the values in Tables I1 and silver bromate in dilute pyridine solutions have been determined. and IV, KI = 1.0 X The results have been interpreted in terms of Discussion.-The values for IG in Table I intwo complex ions, AgPy+ an$ AgPyz+, with discrease with the pyridine molality in the more consociation constants k~ and Ks equal to 1.0 X ( 0 ) Hill, THISJOURNAL, 39, 218 (1917). and 7.8 X 10-6, respectively. (7) Reedy, ibid., 43, 1440 (1821). 0.01475 .02075 .03607 .05492
31.82 33.29 37.68 42.79
52.3 51.9 51.0 50.3
9.5 11.0 16.1 20.5
+
1.65 2.35 3.53 4.8
9.1 11.1 11.2 11.8
+
(8) Dalton, Weymouth and Pomeroy, ibid., 46, BO (1924).
DURHAM, N. C.
RECEIVED
SEPTEMBER 16, 1943
