(17) D. A. Johnson, "Some Thermodynamic Aspects of Inorganic Chemistry", Cambridge University Press, 1968. (18) R . W. Wopschall and I. Shain, Anal. Chem., 39, 1514 (1967). (19) D. J. Barclay and F. C. Anson, J. Elecfroanal. Chem., 28, 71 (1970). (20) F. C. Anson. Anal. Chem., 38, 54 (1966). (21) A. Frumkin, Z. Phys., 35, 792 (1926). (22) A. Frumkin, and B. Damaskin, "Modern Aspects of Electrochemistry", Vol. 3, J. Bockris and B. Conway, Ed., Butterworths and Co., London, 1964. (23) F. C. Anson and R. S.Rodgers, J. Nectroanal. Chem., 47, 287 (1973). (24) S.N. Frank and F. C. Anson, J. Necfroanal. Chem., 5.4, 55 (1974).
(25) M. J. Weaver and F. C. Anson, J. Elecfroanal. Chem.. 60, 19 (1975). (26) R. D. Armstrong, W. P. Race, and H. R. Thirsk, J. Necfroanal. Chem., 23,351 (1969).
RECEIVEDfor review July 3, 1975. Accepted November 3, 1975. We acknowledge support of this research by the Materials Research Center, National Science Foundation U.N.C., under Grant GH-33623 and by National Science Foundation Grant GP-38633X.
Stepwise Titration of Some Anion Mixtures and Determination of Kspof Silver Precipitates with Silver Ion Selective Electrode Eric E. Chao and K. L. Cheng" Depadment of Chemistry, University of Missouri-Kansas City, Kansas City, Mo. 64 110
Anion mixtures, such as sulfide, arsenite, and arsenate can be stepwise titrated in alkaline media with silver solution using the silver ion selective electrode as an indicator electrode. Various mixtures of 15 anions in different combinations have been successfully determined by stepwise titration. A simple method for determining the equilibrium constants such as Ksp using an ion selective electrode has been developed. Equilibrium constants results for 15 silver compounds are reported. Most results agree well with the literature values and some are new data. Silver or arsenite can be determined gravimetrically or volumetrically based on the formation of AgSAs03 precipitate.
There are many methods for determining cations. However, fewer methods are available for determining anions. T h e commercially available silver ion selective electrode responds to both free silver ion and sulfide ion. T h e silver ion activity depends upon the sulfide or other ligand ion activity as it is controlled by the solubility product or equilibrium constants of silver sulfide or silver complexes. The solubility product of silver sulfide has been reported to be 1.48 X ( I ) . T h e Nernstian response of silver ion activity A4 has been reported by using a silver from 0.1 M to ion selective electrode ( 2 ) . Its use to determine the anions which form stable complexes or precipitates with silver ions may be advantageous. Further, it offers a simple method to determine their solubility product constant (Ksp)and stability constants ( K f ) . The use of a divalent ion selective electrode as an indicator electrode in the stepwise titration of triethylenetetraminehexaacetic acid has been reported (3). T h e present paper reports the application of the siiver ion selective electrode as an indicator electrode to titrate sulfide, arsenite, and arsenate. A stepwise method of titrating many anion mixtures has also been developed. We also report here the K,, and the Kf values of 15 silver compounds. Some of them are reported for the first time. Many equilibrium constants which appeared in the literature vary widely. It should be worthwhile to reevaluate them with currently available ion selective electrodes.
EXPERIMENTAL Apparatus. The Orion model 94-16 silver sulfide ion selective
electrode and the Orion model 90-02 double junction reference
Table I. Gravimetric Determination of Silver as Silver Arsenitea N o . of pH of experi- precipirnents tation
Silver taken,
Silver arsenite precipitate,
nig
nig
.
Sil\ e r b f o u n d . rng
Relative error, 9
-0.37 54.1 74.3 53.9 +0.19 108.1 149.4 11.0 108.3 +o. 1 9 162.3 224.3 3 11.0 162.6 4 11.0 216.4 299.7 217.2 +0.37 5 11.0 270.5 374.5 271.4 +0.33 324.6 447.6 6 11.0 324.4 -0.06 7 10.0 +0.46 108.1 149.8 108.6 8 8.0 108.1 150.4 109.0 +0.83 9 6.0 108.1 149.3 108.2 +0.09 a Amount of arsenite added: 1.0040 mmol. The precipitates were dried under vacuum for 17 hr. b Silver was calculated based o n the gravimetric factor 0.7248 assuming the precipitate was in the form of Ag,AsO,. 1 2
11.0
electrode were used for the silver ion activity measurements. All the pH and emf readings were taken from a Corning model 10 pH meter with expanded scale. Infrared spectra were taken from a Perkin-Elmer 621 grating infrared spectrophotometer and the CsI cell and Nujol oil were used. Reagents. The primary standard arsenic trioxide from the Thorn Smith Co. was used for the preparation of 0.1 M arsenite solution; a weighed amount of arsenic trioxide was first dissolved in a minimum amount of 0.1 M NaOH solution, then made up to volume. More diluted arsenite solutions were prepared by appropriate dilutions. The silver ion solutions were prepared from anhydrous silver nitrate. All chemicals used were reagent grade. The water used in the experiment was doubly deionized. Procedure. Determination of Arsenite. The silver ion selective electrode, the double junction reference electrode, and the combination pH electrode were immersed into a sample solution which contained arsenite ion and was adjusted to pH 11.0 with sodium hydroxide. The titration was initiated by adding the silver nitrate solution from a buret with constant stirring. The emf changes were recorded in each interval addition of 1 ml or 0.50 ml of the silver nitrate solution. Near the equivalence point, 0.25 ml or 0.10 ml of silver nitrate solution was added each time. The ratio of silver to arsenite was 3 to 1. The pH was constantly checked and adjusted to 11.0 with a dilute NaOH solution or a dilute "03 solution. Stepwise Titration of Sulfide, Arsenite, and Arsentate. The procedure for titrating a mixture of sulfide, arsenite, and arsenate as well as for titrating other mixtures was the same as the one described above. The resulting titration curves show three breaks for the titration of mixture of all three anions; the first break correANALYTICAL CHEMISTRY, VOL. 48, NO. 2, FEBRUARY 1976
267
E(m
PH
Figure 1. Species of arsenious acid as function of pH
Ye TITRATION
Figure 3. Titrations of arsenite ion with silver at different pH
Y 6
'I
8
IO
9
I1
12
13
PH PH
Figure 2. Response of silver ion selective electrode at various pH's M arsenite ion solution, in (a) doubly deionized water, ( b ) 1 X M silver nitrate solution and (c) 1 X
sponds to sulfide, the second one to arsenite, and the last one to arsenate. Nine glass sintered crucibles (grade F) were cleaned in an aqua regia solution and dried to constant weight. Exactly 10 ml of 0.1004 M arsenite solution were added to nine different beakers. At this point, each beaker was treated separately; for the No. 1-6 beakers, the pH was adjusted to 11.0; for the No. 7-9 beakers, the pH was adjusted to 10,8,and 6, respectively. Then 5,10, 15,20, 25, and 30 ml of 0.1002 M silver solution were added to the No. 1-6 beakers at pH 11.0, respectively; and 10 ml of 0.1002 M silver nitrate solution were added to the remaining beakers at pH 10, 8, and 6. The precipitates were filtered and washed with 40 ml of doubly deionized water of corresponding pH and dried in an oven at 110 O C to constant weight. This procedure was repeated and the results in Table I show the composition of the precipitate for Ag+: A s O ~ as ~ -3:l; the gravimetric procedure is thus established.
RESULTS AND DISCUSSION Few methods are available to determine the mixture of As(II1) and As(V) in the same solution. T h e redox titration method by using iodine and thiosulfate has long been established. However, a t least two separate titrations are required. A simple and stepwise method is needed to determine a mixture of arsenite and arsenate. Figure 1 shows that the major species present in an acidic medium is H3As.03. Because of the amphoteric property of arsenious acid and its low K , values (41, the solubility of the silver arsenite precipitate increases with the decrease in 268
ANALYTICAL CHEMISTRY, VOL. 48, NO. 2, FEBRUARY 1976
Figure 4. Percent recovery of arsenite titration with silver as a func25 % methanol in water. (x) aqueous solution tion of pH. (0)
pH. Thus, a quantitative titration of arsenite with silver is favored in a strongly alkaline medium. Effect of pH on the Silver Ion Selective Electrode. The silver ion selective electrode responds to the hydroxide ions in water a t high pH, because hydroxide ions either precipitate or complex the silver ions produced from the silver sulfide (Figure 2a). In the arsenite solution, the higher the pH, the more species are available for the formation of the complexes and precipitate (Figure 2 b ) , resulting in the more negative potential shift a t higher pH. For a silver ion solution, the higher the pH, the more silver hydroxide precipitate is formed and less free silver ion in the solution, also resulting in a more negative potential shift (Figure 2c). An optimum pH range must be selected for the precipitation since the marked potential differences among these three solutions occurred at high pH especially in the range of pH 9-12. Since the arsenite ion a t pH 11 exists mostly in the protonated form, after the formation of silver arsenite the hydrogen ion is released. Consequently, during the titration of arsenite solution with silver, a buffer solution is needed. A sodium hydroxide solution containing some carbonate may be used for this purpose. Effect of pH on the Titration of Arsenite with Silver. Forty-ml solutions containing 20.08 mol of arsenite M silver solution a t various were titrated with 1.002 X p H values. The titration curves and percentage of recovery as a function of pH are shown in Figures 3 and 4. At pH
Table 11. Titrations of Sulfide, Arsenite, and Arsenate with Silver at 20 “C, Using Silver Ion Selective Electrode as an Indicator Electrode, at pH 11.0 and initial volume: (a) 50.0 ml, ( b )40.0 ml Arsenite a d d e d , mmol
(a )
5.145 X 5.145 X 5.145 X 5.145 X 5.145 X (b)
lo-‘
... ...
... ...
2.058 X lo-’ 2.058 x 1 0 - 2 c 2.058 X 2.573 X 1 W 2 2.573 X 1 0 - 2 C 2.573 X 1 0 - 2 c c In the presence
Arsenate added, mmol
Sulfide a d d e d , mmol
... ... ...
...
...
... ...
2.434 X lo-’ 2.434 x 10-’c
...
3.651’; ‘lo-’ 2.434 x lo-’ 1.217 X 2.434 x lo-’ 2.434 X lo-’ 2.434 x of 25% methanol.
8.638’X ‘10-3 2.160 x l o T 2 1.728 X lo-’ 8.638 x 10-3
...
...
2.432 2.438
...
X X
lo-*
...
...
3.700 X lo-’ 2.471 X lo-’ 1.309 X 2.512 X 2.478 X 2.411 X lo-’
2.104‘; ‘10-2 2.037 x 2.184 x 2.592 X 2.612 x 2.432 X
below 6, no visible precipitate was ever observed. This is not surprising because arsenite a t this p H mainly exists as H ~ A s O much ~; fewer ions exist which are directly responsible for the precipitation of Ag3AsO3. At p H above 12, the recovery is again less than 100% because of the complexation of silver by hydroxide. From Figures 3 and 4, the optimum p H for the titration of arsenite with silver is 9-11; p H 11 is a better choice because most of the cation interferences can be reduced to a minimum a t this p H and it is rather easy to adjust the p H to 11.0 with a sodium hydroxide solution. Effect of Organic Solvent. Titrations were also carried out in the water-methanol media. I t was hoped t h a t the presence of organic solvent might increase the sensitivity of the method because of the decrease in the solubility of Ag3As03 and that the optimum p H range might become wider. However, the results show t h a t the completion of the precipitation was only slightly increased in the presence of 25% methanol a t p H 11 (Figure 4). Concentration Range. The titration limit for the arsenite concentration a t p H 11.0 was found to be approximately 1 X 10-5M. T h e titrant used, Le., silver ion solution, was 10-fold more concentrated than the corresponding arsenite solutions. p H was constantly adjusted to 11.0 by adding dilute NaOH solution. A distinctive break in pAg from a titration curve for M arsenite solution could not be observed (Table I1 and Figure 5 ) .
E(mv)
-
... 1.713 X lo-’ ...
1.743 x lo-’
...
v,
... +0.23 +1.12 +0.35 +0.46 -0.08 +O. 16 -0.86 +0.87
8.717’X ‘10-3 2.124 X 1.743 X 8.617 x 10-3
(31 (21
2w. (I1
JOO
100 % T I TRATl 0 N
Figure 5. Titration of arsenite of various concentrations with silver solution at pH 11.0.(1)1 X M arsenite ion solution with 1 X M silver ion solution, (2)1 X M arsenite ion solution with 1 X M silver ion solution, (3) 1 X M arsenite ion solution with 1 X M silver ion solution, and (4)1 X M arsenite ion solution with 1 X lo-* M silver ion solution
Determination of a Mixture of Sulfide, Arsenite, and Arsenate or Any Combinations. Distinctive breaks were observed in the titration curves of mixtures of sulfide, arsenite, and arsenate. The first break is located a t approximately -300 mV (vs. the double junction reference electrode used) corresponding to the formation of the silver sulfide equivalence point. T h e second break a t approximately +115 mV corresponding to the formation of the Ag3AsOs equivalence point and the third break a t about +265 mV corresponding to the formation of the AgaAsO4 equivalence point. T h e K,, of these precipitates are decreasing in the order o f K s p , ~ g g < s K s p , ~ g s ~
