ROLF ALTSCHUL and HENRY ALTMAN Sarah Lawrence College, Bronxville, New York
component alloys or mixtures can be so analyzed provided the metal pair is far from isobaric. One aggregate titer, in this "indirect" analysis, yields the concentrations of both individual constituents.
T m s paper presents an extension of metal sulfide precipitations from qualitative separations to quantitative analyses. ~h~ proton exchange accompanying aci&metrically. the process (equation 1) is 2H20+ M++ + H2Se MS
+ 2HnO+
(M
=
metd ion)
Our results may be significant to the analytical chemist, since they allow the quantitative determination of various heavy metal ions by a speedy and uniform procedure. This approach is, however, probably more significant to students, who gain experience in general acidimetly as well as in the charting and complete interpretation of titration curves involving mixed strong and weak acids. We have provisionally incorporated it in the quantitative analytical chemistry course in this college. Equation 1 represents the basic principle; precipitation of metal ions with an excess of hydrogen sulfide sffords a mixture of strong and weak acid. The former can be determined selectively and is a measure of the original metal ion molarity. This analytical attack is somewhat similar in principle to recent elegant '~complexometric" methods based on protolyses attending specific complex formations.' The relative position and mobility of the equilibrium in equation 1decide the applicability of the acidimetric procedure. Both factors are sufficiently favorable for the majority of the common subgroup metals: some are precipitated quantitatively upon addition of hydrogen sulfide, displacing the equilibrium entirely to the right (e. g., silver ion or mercuric ion). For other metals, completion is gradually attained upon slow addition of alkali during the titration proper (e. g., zinc ion or nickelous ion). In order to obtain clear end points, the techniques is confined to those metals whose sulfides are insoluble a t pH four or below. This qualification has excluded, for instance, manganous ion. There is no interference from substances unreactive toward hydrogen sulfide. The presence of acid in an unknown solution requires a preliminary determination of a "blank" titer for subtraction from the h a 1 reading. A solution of several heavy metal salts must, of course, be subjected to separation prior to application of the procedure above. This can be accomplished by any of the means ordinarily required in conventional analyses. There exists, however, a special case for which a preliminary separation is unnecessary: two SCBWARZENBACH, G., Helv. Chim. Ada, 29,1338 (1946). BIER., AND G. SCHWARZENBACH, Chimia, Z,56 (1948). These paperscontainreferwce toearlierpublicationon thesame subject. 1
DERW,
EXPERIMENTAL PROCEDURE AND RESULTS
(1)
Solutions of silver, lead, cadmium, mercuric, cupric, zinc, and nickelous salts were prepared from Baker's Analyzed c. p. chemicals. The acidic solutions (zinc sulfate and mercuric chloride) were subjected to indicator titration in order to determine the acid molarity. An excess of aqueous hydrogen sulfide, saturated at ice temperature, was added to a sample of the unknown solution in a 100-ml. beaker. The aliquot contained about two millimols of metal. Three approaches were tested: Method A. A complete titration curve was determined experimentally with a Beckman pH Meter (Laboratory Model). The standard accessory glass electrode and calomel electrode did not appear to suffer from prolonged exposure to sulfide solutions. Measured portions of carbonate-free standard alkali were admitted from a buret to the stirred suspension, followed by a pH reading after each addition. These pH measurements were reproducible and steady for ions of the Copper Group and for silver ion, all of which are completely precipitated by hydrogen sulfide. Suspensions containing zinc, nickel, or cobalt ion showed a drift from a maximum to a minimum pH after each alkali addition; this equilibrated value, generally reached after three to four minutes, was recorded. Titrations were carried a t least into the bisulfide buffer range, pH 6.5, the complete operation taking from 15 to 20minutes. Three representative curves are plotted in Figure 1: No. 51, silver ion, typifies the case of the most insoluble sulfides; the equilibrium is completely displaced toward formation of products (equation 1). This particular titration was carried beyond the first equivalence point of HzS (pH 9). No. 75, nickel sulfide, is gradually precipitated during the titration at slightly lower acidity, establishing a "buffer" around pH 2.5 or 3. This heterogeneous process is known to be slow,2 fully in accord with the drifting pH readmgs described above. No. 52, manganous sulfide, is too soluble to yield to this procedure; it is precipitated within the range of the bisulfide buffer. No end point is observed a t the calculated titer (about 9.6 ml.). This method is recommended for student work.
* HAMMEW,L. P., "Solution of Eleotrolytes," McGraw-Hill Book Co., New York, 1936, Ch. 111.
278
MAY, 1950
First, it illustrates the manipulations involved in the construction of multipleacid titration curves. Second, location of the relevant end point is required, which, in turn, is the key to the analysis of the original unknown. This assignment thus combines fundamental features of acid-base equilibria with a specific application to analysis. Compared to conventional student analyses of simple binary electrolyte solutions, this procedure is presumably as accurate as most and more rapid than some. Method B. This was a simplification of the preceding techniques, dispensing with the charting of a complete curve. The previous procedure was followed, but no individual p H readings were taken. The meter was set initially a t approximately pH 3.7, a value considerably below that of the final reading. This served as a precaution in that the needle functions as a warning signal of the approaching end point. After resetting the meter to the proper pH (4.50), the titration generally required another four to eight drops of standard alkali before the needle crossed the center marking. The entire operation generally took from three to five minutes. We recommend this procedure for general analytical purposes. I n the presence of members of the Zinc Group the final stage of the titration should he carried out slowly t o allow for complete precipitation. Method A is superior for nickel solutions. TABLE - 1 Acidirnetric T i t r a t i o n o f S t e n d w d Metal Salt Solution Analysis
Components, mls/liter
Molarity found
51 85 59 60 76 55 57 78 32 75 101 40 14i 15i 91 89* 90* 18i' 42
0.1991 0.1992 0.3075 0.3064 0.3074 0.0990 0.0992 0.0987 0.1021 0.1025 0.1021 0.1483 0.1488 0.1485 0.1503 0.1500 0.1500 0.1501 0.2211
98
0.2198
34i
0.2207
35i
0.2206
45
0.0989
46
0.0988
77
0.0992
.,.. 52 * This run was carried out by Mr. Alan Hecht.
Method
10
20
13
Cc. of NaOH F i w m 1. Acidirndris Titration 0 w . s for Aqueous S Y S P ~ M ~ O of M Mat& Sulfidaa i n ths P m s ~ n c .of Ex-- Hydrogen SulSde
From top t o bottom: Mn. No. 52 ( r = 4 . 0 ) ; Ag. No. 51 (I No. 75 ((r = 0).
= 2.0);
Ni,
Method C. Simple indicator titrations were possible for suspensions of zinc sulfide and cadmium sulfide. Other metals form precipitates so highly colored as to preclude this approach. Congo red appeared most suitable. I n the presence of cadmium it gave a clear end point, from green to orange, easily discernible on comparison with a previously prepared standard. Table 1 compares the results of orthodox analyses (column 2) with those by methods A, B, and C (column 3). The mercuric chloride concentration in the second column is uncertain. No accurate data were obtained through the recommended m e t h ~ d . ~ TABLE 2 Acidirnet~icAnalvses o f Metal Pairs Analyses 69 73 71
Composition, % by wt. Silver, Copper, Mercury, Cobalt, Lead, Zinc,
90.00 9.98 76.1 23.9 85.2 14.8
Composition f m d (method A 1, % 89.94 10.06 76.45 23.55 85.88 14.12
Two-Component Mixtures. Three metal pairs were analyzed, without preliminary separation, by means of one acidimetric titration. This represents the special case referred to above. The constituents had widely differing equivalent weights in each instance. Accurately weighed quantities were dissolved in 1:l nitric acid, evaporated to a moist paste on a hotplate, transferred to a volumetric flask, and diluted to volume. Separated aliquots were: (a) analyzed for their residual acid content by indicator titration (bromphenol blue or methyl red), and (b), subjected s B 6 AND ~ OESPER, ~ ~ "Newer ~ ~ Methods of Volumetric Chemical Analysis," D. Van Nostrand Co., New Yark, 1938, Part IV.
280
to Method A . This titer and the weight of the original sample, on substitution into Equation 2 resolved the unknown into the concentrations of its individual components. 100 X ( m y - WY) (2) Weight per cent (XI = w(zy - Xy) where X = atomic weight of metal (X); Y = atomic weight of metal (Y); s = valence of metal (X); y = valence of metal (Y); w = weight of metal sample, in grams; m = mols NaOH, for titration of sample (corrected for molarity of acid in solution).
Table 2 summarizes our data. First, a silver-copper
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alloy was analyzed by orthodox means, involving quantitative separation. This was checked against acidimetric titration (No. 69). Second, semiquantitative data were obtained from coarsely mixed metal pairs prepared by weight (Nos. 71, 73). No orthodox analyses were carried out. The discrepancies between and the determined results are in the the direction of impurities in the metal specimens and should most reasonably be to this factor since some samples were only 99.0 to 99.5 per cent pure.
