68
ARTHUR E. HILL.
of a larger proportion of sulphate sulphur in these concentrated solutions than was found in the dilute solutions boiled in open air also favors the supposition that the reaction by which pentasulphide is produced a t the expense of thiosulphate is accompanied by the formation of ca.lcium sulphate. The snialler proportion of sulphate sulphur in some of the more Concentrated solutions seems to oppose this view, however. on the other hand, the analytical met hods for distinguishing quantitatively between these several forms of sulphur compounds are not yet thoroughly perfected and it may be that this apparent objection to the reaction would be removed if more exact methods of analysis were available. In the absence of better analytical methods, these results are presented as the best obtainable, and the conclusions suggested as a possible step toward a better knowledge of thc reactions between lime and sulphur under varying conditions. I,Al3ORATORY O F T H E ~~VSHINI’TON . k G X I C C L T U R A L k:XPRRIMENT 1’11
STATIOS,
L 1,nf A N, TV.4 S H .
[CONTRIBUTION FROM THE
HAVEXEYER CHEMICAL
LABORATORY, XEW I’ORK
UNIVER-
SITY.]
THE RELATIVE SOLUBILITY OF THE SILVER HALIDES AND SILVER SULPHOCYANATE. 15’1 A R T H C X
E. H I L L
Received October
22, 190;
The theory of electrolytic dissociation teaches that if saturated solutions of two very insoluble salts which have an ion in common could be mixed without increase of volume, precipitation of both salts would occur. For example, saturated solutions of AgCNS and AgCl would, upon mixing, precipitate both compounds in part, since by the addition of silver ion the solubility product of each salt would be exceeded. The quantities of chloride and sulphocyanate precipitated would be such as to leave the solution saturated in respect to both salts, and the equilibrium finally reached would be expressed for the respective compounds by the equations CAPx CCN~ = K, and CAg s C, = K,, where C stands for concentration in equivalents per unit volume, K, and K, are the products of the free ions (solubility products), and the subscripts denote the respective ions. By division,
The common term ,C,
may be cancelled out, the relation becoming
SILVER HALIDES AND SILVER SULPHOCYANATE.
69
The same equilibrium will be attained if t o a solution of AgC1, in presence of excess of the salt, a soluble sulphocyanate, such as KCNS, be added in sufficient quantity to exceed the solubility product of AgCNS. Precipitation of AgCXS will result, reducing the concentration of silver ions ; the solutioii will therefore become temporarily undersaturated with respect to AgC1, and the latter salt will dissolve, thereby increasing the concentration of chloride ions and also of silver ions, so that the cycle of reactions will again be set in operation. Simultaneous precipitation of AgCNS and solution of AgCl will continue until the mixture contains the
__
-
ions C1 and CNS in such quantities as are in equilibrium with the two silver salts, as expressed in equations I and 2. The solution thus obtained differs from the hypothetical solution discussed in the preceding paragraph in having present an additional cathion (potassium ion) and in having the chloride and sulphocyanate ions present in amounts measurable by ordinary analytical methods. The theory of solution equilibria of this character was first stated by Nernst.’ By the application of equation 2 to reactions such as the foregoing it becomes possible to determine the relative solubility of two difficultly soluble salts whenever ( I ) the concentration of the free ions (ec AgCNS -I. KCl. The solutions were analyzed for total chloride and 6% sulphocyanate by the method of Volhard and for sulpliocyanate by the colorimetric method, the chloride being dcterlilined by difference. For the Volhard determination 2 % j cc. samples were taken in experiments I U and I h , jo cc. samples in 2~ and z h , and I O O cc. saniples in 3a and 3b. For the colorinietric tests the standards were made in every way identical with the anal?-zed solutions by mixing solutions of potassiuni chloride and potassiun: sulphocyanate in such quantities that the total salt in the standard was equal to that in the sample tcsted, and by adding equal quantities of iron-ammoniuni aluiii to the two solutions. In the following table, column I gives the nuinher of the experiment, column 2 the approximate conccntration of tlir solublc salt used, colunin 3 the salts takrn for the reaction, and coluniiis 4 and j the final concentration of chloride and sulphocyanate as :malyticall!~ determined. Column 6 gives the ratio of the square roots of these conccntrations, which, according t o equation 4, is equal t o t h e ratio of the solubility of the two salts.
+
n,l
TABLE I
+ KCSS
Equilibrium XgCl
XgCSS
+ KC1 at
2 jo.
CCNS Approx. concen.
N/5 N/5 N/20 x/20
N/IOO S/roo
conc. CI.
Keacting salts.
+ + +
AgCKS KC1 AgCl KCNS A4gCNS KC1 .4gC1 KCNS AgCKS KCI AgCl KCNS
+ + +
COllC.
cxs.
0.197
o.ooro.+
0.193
0.00100
0,OjOI
0,000293
o.o@g o 0x0
o.000280
0.0726 0.0719 0.0764 0.0759
o.000060 o.0000~7
0.0774 0.0750
0.010
Mean ratio,
S-\gcl = 0.0748 hgc PI'S
s
TABLE 11. Equilibrium XgCNS -t KBr Approx. Colleen.
N/5 s/5 N/zo N / ~ o
I
Keacting salts.
AgCNS t KBr AgBr KCNS AgCXS KBr AgBr + KCNS
1. pr. Chern., 9 (N.I?.),
+
217
+
(1x74).
AgBr
+ KCNS at 25'.
Conc. nr.
Conc. CKS.
0.06.+j
0.1205
0.0665 0.0165 o 0176
0.1207
0.0312
v 0324
4:z 0,732 0,742 0.727 0.737
See also this journal, 29, 2 G 9 (190;)
SILVER HALIDES AND SILVER SULPHOCYANATE.
73
Table I1 shows the results obtained by comparing AgBr and AgCNS. The solutions were analyzed for bromides by the method of Rosanoff and Hil1,I the total bromide and sulphocyanate being determined by the method of Volhard. The columns have the same significance as in the previous table. ,4ttenipts were made to carry out experiments in solutions of hundredthnormal concentration, but fine suspensions of the silver salts resulted, which would not settle in reasonable time and could not be filtered out. The solutions were finally analyzed in the presence of the suspended solid, but under those conditions the results obtained by approaching the equilibrium from the two directions differed by 4 t o 5 per cent., the amount of bromide in solution being larger when silver bromide had been taken than when its original source was the soluble potassiuni bromide. The values for the ratio were 0.0709 and o.ojoo when silver sulphocyanate was the original solid phase, and 0.078 j and 0.0 763 when silver bromide was taken; the mean is almost identical with that of Table 11. Whether this peculiarity is referable simply to analytical errors due to the presence of the finely suspended silver salts or, as seems to be indicated by Table 11, t o some change in the nature of the solid phase (mixed crystals, solid-solution, or the like) must remain for the present unanswered. Analytical difficulties prevent the determination of the relative solubility of silver iodide and bromide from being as accurate as might be desired. The iodide in solution when equilibrium has been attained amounts to only 0.01 to 0.02 per cent. of the bromide present, expressed in equivalents, and no method is known by which such minute quantities of iodides can be exactly determined in the presence of such large amounts of bromides. Of the various methods proposed for this separation, t h a t of Fresenius,2 which depends upon the oxidation of the iodides by nitrous acid in sulphuric acid solution, extraction of the free iodine by means of an organic solvent and titration with Na,S,O,, was thought t o be best suited to the estimation of small amounts of iodine. Test analyses were made of solutions containing 0.1 to 0.5 mg. of iodide in presence of z to 3 grams of bromide dissolved in 150 cc. of water. The iodine freed by five drops of a solution of nitrous acid in concentrated sulphuric acid was collected by extraction of the aqueous solution with I O cc. portions of chloroform until no further coloration of the solvent could be detected; the iodine was then titrated with N/4oo Na,S,O,. The quantity found was always less than that taken, and by amounts varying between 1 2 and 40 per cent. If this maximum error be assumed to have occurred in the analysis of the mixtures recorded in Table 111, the error in the ratios calculated will be about 2 7 per cent., and the ratios This Journal, 29, 1461(1907). Quam Chem. Anal. (Braunschweig, 1898),p. 482.
74
WALTER E. MATHEWSOK.
given below may therefore differ from the true ratios by t h a t amount. In the experiments tabulated below, somewhat greater concentrations were used than in the previous experiments. in order t h a t the iodide present might become a measurable quxntit!;. TABLE 111.
Equilibrium AgBr Conceii. Appros.
Keactitig salts.
ra rh
N
2u
S/\i
2h
N/,j
AgI KBr AgBr KI AgI t KBr :lgBr + K I
No.
+ KI
+ KBr at 2jo. 2 4 . 4 X IO)-^ 1.3.9 X IO)-^ 4 . 8 3 x (Io:-.'
o.()jj
+
J
colic. I .
Colic. llr,
+
B
AgI
0.94j
0.185 0.187
2.2
x
CBr
0.016 0.012
0.016
(10)-
Mean ratio
C;
0.011
SAgI
SAgUr
=o.o14
I;rom the ratios recorded in 'fables I . I1 and I11 the relative solubility of the whole series may be calculated. The values thus obtained arc given in Table IV. Colunin I designates the salt, column 2 its relative solubility referred t o that of silver chloridi. :LS unity, and column 3 the absolute solubilitl- in gram-molecules per liter. taking Kohlrausch and Rose's' figure for the silver chloride as the standard. Column 4 gives the niasi~nuiii and minimum values obtained for the solubility of these salts by other methods; the figures are takc.11 from the table complied by .\begg and Cox.? 'I'ABLE IS'. Relative solubility.
Salt.
.
.4gCl
I 00000
A b s o l u t e zolubility.
1.6
l*:xtl-emevalues.
r: (IO)-?
I . 25-1.64
X
IO)-^
x (10l-h I .OS-I, 2 j x (101-6 0 . O jj O 0 8 . S x (Io)-; 6 . 6 -8.1 x (IO)-' o.ooo;.i 1 , 2 3 x (IO)-' 0.97-1.05 x (IO)+? '['he foregoing investigation is onc of L: series planned in collaboration with Professor AI. X. Rosanoff. A change of residence has made it advis;tblc to carry out the scparatc parts of t h e work independently. Credit is gladly given t o Professor Rosanoff for many of ihe ideas contained in this paper. AgCXS AgBr AgI
0 . 07480
I . ?
S 1. TV YO It K 1. S 1VI;I? S I T Y
October,
ty~:.
._
..
...
ON THE ANALYTICAL ESTIMATION OF GLIADIN. RY W A L T E R I-:n r A T r I m v s o x . Received O c t o l x r
2:. I!)(
:.
'1~Iiat tlic co.1111;o11methods for the determination of gliadin are unsatisfactor>-is generally recognized. The amount of nit rogeiious material
' Abegg .I
Lor. cit
and Cos, Z. physik. Cheni., 46,
II
(1903).
