Table I. Recovery of Lead Added to Samples Sample, Pb Total Pb Pb 5.00 g added, mg found, mg recovered, mg Cr Ni NBS No. 55e NBS No. 344
0.200
0.225 0.024
0.201
O.Oo0 0.200
0.204
0.194
O.Oo0
0.010
0.200
0.207
0.191
O.Oo0
0.016 0.197 0.008
0.189
0.200
O.Oo0
Table 11. Determination of Lead by Atomic Absorption US. Polarography (8) Lead. m m (ExtractionSample (Ion exchange-AA) polarography) 1 2 3 4
tion nearly to dryness, redissolving any soluble salts in 40 ml of 6N HCl, diluting to 50 ml, and then filtering through a medium porosity paper. The paper and residue were washed several times with 1N HC1. The filtrate was evaporated to a volume of 15 to 20 ml. The volume of filtrate was measured with a cylindrical graduate and then diluted with exactly 5 volumes of water to give a solution of 1N HCl. The solution was passed through the ion exchange column. The column was rinsed with 75 ml of 1N HC1 and then the lead was eluted with 150 ml of 12N HC1. The first 75 ml of eluate was returned to the column to remove small quantities of iron eluted from the column by the HC1. The eluate was evaporated to near dryness and 10 drops of HNO, plus 5 ml of water were added to redissolve any residue. The solution was diluted to volume in a 10- or 25ml volumetric flask. The concentration of lead in the test solution was determined by a Perkin-Elmer Model 303 Atomic Absorption Spectrophotometer using an air-acetylene flame and a wavelength of 217 nm. A Westinghouse No. 22927 lead hollow cathode tube provided the resonance line. Calibration. Aqueous calibration standards were prepared from Baker Analyzed lead nitrate. A stock solution containing 0.500 mg of lead per ml was diluted to give standards containing 0.0025, 0.0050, 0.0100, 0.0200, and 0.0300 mg of lead per ml. An appropriate volume of H N 0 3was added to match the concentration of the test solution.
5
4 12 30 65 136
6 7 8 9
49 28 22
7
4 10 28 65 126 6 47 28 20
(NBS) samples was determined. NBS sample No. 55e is ingot iron while NBS sample No. 344 is a chromium-nickel stainless alloy containing 2.40 Z molybdenum. Two replicates of each sample were taken through the procedure. To one replicate, 0.200 mg of lead was added initially while the second replicate served as a blank. The results of the analyses are shown in Table I. The loss of lead in the acid insoluble portion of seven steel samples was determined to be less than 1 ppm for all samples. A comparison between the ion exchange-atomic absorption procedures and a liquid-liquid extraction-polarographic procedure (8) for nine steel samples is presented in Table 11. Good agreement is evident for the lead analyses obtained by these two procedures. These methods are different, both in the preliminary separation and in the measurement of lead. Therefore, except for a cancelling of different errors, the values obtained are satisfactory. The standard deviations of the ion exchange-atomic absorption technique is indicated by the values of lt0.9 ppm at a level of 12 ppm and lt2.0 ppm at a level of 65 ppm for 10 replicate determinations on each of two samples.
RESULTS
The recovery of lead added to 5.00-gram samples of nickel, chromium, and two National Bureau of Standards
RECEIVED for review May 27, 1971. Accepted September 14,1971.
Determination of SolubiI ity Product Constants of Some 1:l Silver Salts of Thiols by Direct Competition with Silver Iodide Z. C. H. Tan, R. C. Vought, and R. B. Pontius Eastman Kodak Company, Rochester, N. Y . 14650 MOSTSOLUBILITYPRODUCT constants of sparingly soluble salts have been determined by electrometric methods--e.g., potentiometry, conductometry, etc.-either by precipitation titration or by measurement of the concentrations of the dissolved species in a saturated solution of each salt. In these types of measurements, adsorption of the precipitants or dissolved species on the electrodes is generally negligible. However, whenever the precipitants adsorb strongly on the electrode, measurement of their solubility product constants electrometrically is usually unreliable. This is the case for
some heterocyclic compounds (1-4). For example, Faerman (3) reported a value of 9.3 X 10-'6 (at 20 "C in neutral solution) for the silver salt of 1-phenyl-5-mercaptotetrazole (1) R. J. Newmiller and R. B. Pontius, Photogr. Sci. Eng., 5, 282 (196 1). (2) E. Gunther, Photogr. Korresp., 102, 108 (1966). (3) G. P. Faerman and A. P. Pletner, Usp. Nauch. Fotogr., 5, 114 (1957) (4) F Evva, 2. Wiss. Photogr., Photophys. Photochem., 60 (9-12), 145 (1967).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 2, FEBRUARY 1972
411
Table I. Solubility Data for Competition Reaction between Iodide and PMT Ion for Ag+ Initial iodide Initial PMTInitial AgN03 -PMT-1 concn, M concn, M concn, M Equil [I-], M Equil [PMT-I, M [I-]
Sample No. 1
x
x
10-4 0.57 x 10-4 = 0.05 with KN03) 2 x 10-4 2 x 10-4 0.54 x 10-4 3 x 10-4 AgN03, p = 0.05 with KNOI) (mixture of iodide and PMT2 x 10-4 2 x 10-4 2 x 10-4 1.19 x 10-4 2 x 10-4 2 x 10-4 2 x 10-4 1.14 x 10-4 (No. 3 and No. 4; liquid mixture of iodide and PMTAgNOa only)
2
10-4
2
10-4
3
x
4.4
x
10-5
0.77
0.91
4.5
x
10-5
0.83
0.94
9.88 x 10-5 9.76 x
0.83 0.86
0.80 0.84
(solid AgI and AgPMT in solution, p
2 3 4
+
+
Av 0.82
[I-l~efti n aolution 1
5
x (p
6 7 8
1 2 2
x x x
10-4 = 0.05 with 10-4 10-4 10-4
2
x
10-4
1
x
10-4
5.8
1 1 1
x x x
10-4 10-4 10-4
5 . 8 x 10-5 1.56 x 10-4 1.56 x 10-4
KN03)
x x 1 x
2 1
10-4
10-4 10-4
(PMT), while Evva (4) obtained values at 30 “C that ranged from to depending on the P M T concentration, which was varied from 1.7 X to 8 X 10-4M,and the pH of the solution. These previously reported values are larger than the solubility product constant of AgI (8.28 X lo-’’ at 25 “C) (5). On the other hand, we observed that in solution, PMT- can displace the iodide ion from AgI and that at high PMT- concentration (>1 X lO-*M), the addition of a small amount of silver ion to a mixture containing more iodide than PMTwill always give a mixture of solid AgPMT and AgI, with AgPMT present in the larger quantity. These observations led us to believe that the solubility product of AgPMT is either smaller than or at least comparable to the solubility product of AgI. We ascribe the discrepancies in the values reported for AgPMT by other workers to the adsorption of PMT ion to the silver surfaces (1,2). The present work describes the accurate measurement of the solubility product constant of AgPMT by direct competition with AgI. The solubility product constants of the 1 :1 salt of 2-benzimidazolethiol, 2-mercaptobenzothiazole, and 2-benzoxazolethiol were also determined by the competition method.
THEORY If two anions form slightly soluble salts with the same cation or if two cations form slightly soluble compounds with the same anion, the less soluble compound will precipitate first on the addition of a precipitating agent to a solution containing both. Consider the case in which we add silver nitrate to a mixture of iodide and PMT-. The less soluble silver salt will precipitate first, and when the second silver salt begins to separate, the supernatant liquid is in equilibrium with the two solid phases, AgI and AgPMT. This means that the following relations are satisfied : [Ag+][I-]7~g+71-= 8.28
x
Ksp,AgI
=
10-1’ at 25 “C and zero ( 5 ) ionic strength (1)
(5) “The Theory of the Photographic Process” 3rd ed., C. E. Kenneth Mees, and T. H. James, Ed., Macmillan, New York, N.Y., 1966, p 6. 412
[I-]consumed [PMT-lc,n,u,,~
x
10-5
[PMT-l~eftin solution 1.43 x 10-4 1.47 x 10-4 4.87 x 10-5 4.48 x 10-5
0.74 0.79 0.86 0.80 Av 0.83
where K,, denotes solubility product constant of the subscripted salt, brackets denote equilibrium concentration in solution, and 7’s denote activity coefficient of the species. At equilibrium (3) or when both AgI and AgPMT are present as solid phases :
If the activity coefficients of PMT- and I- in dilute solution are essentially the same and I- and PMT- are determined in the supernatant solution, the solubility product constant of AgPMT can be obtained from the above equation. The K s p , ~value g ~ calculated ~ ~ ~ from Equation 4 relative to the at zero ionic strength is the solubility product value of Ksp,AgI constant of AgPMT at zero ionic strength.
PROCEDURE AND RESULTS Recrystallized reagent grade AgN03, KI, and K N 0 3 were used as starting materials. The potassium salt of PMT (KPMT) used was prepared as follows: A 3-liter, 3-necked, Morton flask, equipped with air stirrer and drying tube, was charged with 1500 ml of absolute alcohol and 95 grams of 85% KOH. To the resulting solution was added 261 grams of PMT (Eastman P5220 l-Phenyl-2-tetrazoline-5-thione recrystallized from benzene) and stirring was continued for 2 hr. Subsequent standing at room temperature yielded a white precipitate which was collected on a filter; washed with absolute alcohol, triturated twice with boiling benzene, and dried under vacuum overnight. Final yield was 73 % melting at 247-249 “ C . Elemental analysis gave: C, 38.5%; H, 2 . 4 z ; N, 26.1z; S, 15.0%; K, 17.8%; compared to calculated values for C7HsN4SK: C, 38.9%; H, 2.3%; N, 25.9%; S, 14.8%; K, 18.1%. Table I summarizes the procedure and results of the experiments for AgPMT. In sample No. 1 , 2 X IOW5mole of aqueous KI and 2 x 10-5 mole of aqueous KPMT, each containing 1.5 X 10-5 mole of AgN03, were mixed and diluted to 100 ml (By this dilution the concentrations shown in sample 1 of Table I resulted). The ionic strength ( p ) was kept constant at 0.05 by adding KN03. In sample No. 2, KI and KPMT were mixed first followed by addition of KNOI and
ANALYTICAL CHEMISTRY, VOL. 44, NO. 2 , FEBRUARY 1972
Table 11. Measured Solubility Product Constants at 25 "C Salt KBP Ag-phenylmercaptotetrazolate Ag-benzimidazolethiolate Ag-mercaptobenzothiazolate
Ag-benzoxazolethiolate
(6.8 i 0.3) (3.7 i0.6) ( 2 . 3 i 0.6) (1.0 i. 0.4)
x 10-17 x 10-17 x 10-17 x 10-18
AgN03 and dilution to 100 ml. In samples No. 3 and No. 4, the liquid mixture of KI and KPMT was equilibrated with AgN03 without the addition of KN03. Each sample was agitated at a constant temperature of 25 "C for several days until equilibrium was reached. The supernatant solutions were separated from the solids by filtration through 0.45-p millipore filters and analyzed for PMT- and iodide concentrations. The PMT- concentration was analyzed by reading the UV absorption at 267 nm. The iodide concentration of each sample was analyzed by two methods : colorimetrically according to the method of Chapman (6), whereby the iodide was made to react with palladous sulfate in acid solution and the absorbance of the reaction product (Pd12) was read at 390 nm and corrected for the PMT- absorption at this peak; and by measurement of the catalytic action of iodide on the reaction between cerium (IV) and arsenic (111) in acid solution (7). The results from both methods agree fairly well. As shown in Table I, the ratio of the equilibrium concentration of PMT- and iodide approaches a constant value with an average of 0.82; the KBpof AgPMT was, therefore, calculated by Equation 4 to be (6.8 f 0.3) X lo-". To verify the closeness in values of the solubility product constants of AgI and AgPMT, various mixtures of KI and KPMT with excess of one anion were equilibrated with AgN03 (Samples No. 58). Since the amount of AgN03 added is not sufficient to establish the equilibrium relationship represented by Equation 4, the concentration ratios of I- and PMT- left in the solution were not used; instead, the ratios of the anion concentrations that went into precipitates, as [I-]conaumed/ [PMT-Icmmmed (where [I-]consumed = [I-]initial - [I-liert in aolutian) were calculated as given in the last column of Table I. These ratios are very close to an average value of 0.83. The ratio is almost constant regardless of whether the solution is initially more concentrated in PMT- (samples No. 5 and No. 6, with
No. 5 kept at constant ionic strength of 0.05>,or initially more concentrated in I- (samples No. 7 and No. 8>, or the concentrations of PMT- and I- are the same initially (samples No. 1-4, with samples No. 1 and No. 2 kept at constant ionic strength of 0.05). This confirms the closeness of the two solubility product constants because one would expect large ratio if the soludeviation of the [I-]oonsumed/[PMT]oansumed bility product constants were about an order of magnitude apart, since the initial AgN03 concentration used is low. The larger deviation of the [I-]oons~med/[PMT-]o~n~u~edratio in the first two samples is probably due to the fact that not all [I-]consumed and [PM~]consumed are completely in the solid phase. Thus, the K,, of AgPMT is slightly smaller than the Kspof AgI at 25 "C. This explains the observations mentioned in the Introduction. The solubility product constants of the 1 :1 silver salts of other thiols-namely, 2-benzimidazolethiol, 2-mercaptobenzothiazole, and 2-benzoxazolethiol were also determined by this method. The compounds used were in the form of their potassium salts prepared by a method similar to that for KPMT. Table I1 summarizes the measured solubility product constants at 25 "C and zero ionic strength for the four compounds studied. The uncertainty of each constant increases as the solubility product constant becomes smaller compared with the K,, of AgI. The value for 2-benzoxazolethiol should probably be taken only as an indication of order of magnitude since about 96-99Z of the anion goes into the solid phase, making the accuracy of the measured equilibrium concentration of 2-benzoxazolethiolate by spectrophotometry very poor. CONCLUSION The solubility product constants of Ag-phenylmercaptotetrazolate and Ag-mercaptobenzothiazolate as determined by us are much lower and more accurate than values previously obtained from electrochemical measurements ( 3 4 . No published values are available for Ag-benzimidazolethiolate and Ag-benzoxazolethiolate. Generally, when other techniques are not reliable, the competition method can be the easiest and most efficient way of determining the solubility product constant of any sparingly soluble salt. For best results, the solubility product of the competing salt should be similar to that of the salt of interest unless a highly sensitive technique for the analysis of the anion of interest is available.
(6) F. W. Chapman, Jr., and R. M. Sherwood, ANAL.CHEM., 29,
ACKNOWLEDGMENT We are grateful to Dr. P. Neddermeyer for valuable discussions.
(7) "Treatise on Analytical Chemistry," I. M. Kolthoff and P. J. Elving, Ed., Part 11, Volume 7, Wiley-Interscience, New York, N.Y., 1961, p 382.
RECEIVED for review July 6, 1971. Accepted September 15, 1971.
172 (1957).
ANALYTICAL CHEMISTRY, VOL. 44, NO. 2, FEBRUARY 1972
413
