Precipitation of Lead Sulfide by Thioacetami TA Solutions Dennis G . Peters and Abdolreza Salajegheh Department of Chemistry, Indiana University, Bloornington, Ind.
KINETICS STUDIES have revealed that in acidic thioacetamide solutions two different mechanisms are involved in the precipitation of the sulfides of lead ( I ) , cadmium (Z), zinc (3), and nickel (4). First, an individual metal ion can combine with the hydrogen sulfide which is generated by hydrolysis of thioacetamide or, perhaps, by some specific interaction between thioacetamide and another substance in solution. Second, the metal ion may undergo B direct, pH-dependent reaction with thioacetamide itself. Although the two processes are competitive, experimental conditions can be established such that one of these mechanisms predominates. In an aqueous ammoniacal medium, zinc, nickel, and cadmium sulfides are precipitated according to the same general mechanisms encountered in acid solutions-that is, the metal-ammine complexes interact homogeneously with sulfide formed from the ammonia-thioacetamide reaction (5, 6) as well as with thioacetamide directly (7,8). On the other hand, since no stable lead(I1)-ammonia species exists, the formation of lead sulfide can only proceed heterogeneously through metathesis of lead hydroxide by either sulfide or thioacetamide. Unless a ligand which strongly complexes lead ion is present, the precipitation of lead sulfide by thioacetamide in an ammonia solution has limited analytical applicability. Flaschka (9, IO) has proposed that ethylenediaminetetraacetic acid (EDTA) be employed as an auxiliary complexing agent in ammoniacal thioacetamide solutions. One advantage in the use of EDTA is to prevent the precipitation of oxides or hydroxides of metals which are not complexed by ammonia. Furthermore, those metal ions which form especially stable EDTA complexes, but relatively soluble sulfides, might be precipitated very slowly or not at all by thioacetamide, so that the number of successful separations of metals could be increased. Although Flaschka observed the qualitative behavior of many metal ions in solutions containing both ammonia and EDTA, no detailed quantitative studies of these systems have been reported previously. In the present investigation we have sought to evaluate the suggestions of Flaschka through a study of the kinetics of precipitation of lead sulfide by thioacetamide in ammoniaammonium nitrate buffer solutions containing EDTA. One question of particular interest was whether a bulky, polydentate ligand such as EDTA, which forms a stable complex with lead ion, could totally prevent the direct reaction between thioacetamide and lead(I1). In such an event, the rate of pre-
(1) E. H. Swift and E. A. Butler, ANAL.CHEM., 28, 146 (1956). (2) D. F. Bowersox and E. H. Swift, ibid., 30, 1288 (1958).
(3) D. F. Bowersox, D. M. Smith, and E. H. Swift, Tulunta, 3, 282 (1960). (4) Zbid., 2, 142 (1959). (5) D. G. Peters and E. H. Swift, ibid., 1, 30 (1958). (6) D. G. Peters and A. Salajegheh, ANAL.CHEW, 38, 1824 (1966). (7) D. H. Klein and E. H. Swift, Tulanta, 12, 349 (1965). (8) D. H. Klein, D. 6. Peters, and E. H. Swift, ibid., p. 357. (9) H. Flaschka, Chemist Analyst, 44,l (1955). (10) H. Flaschka, 2.Anal. Chern., 137, 107 (1952).
cipitation of lead sulfide might be governed by and predictable from knowledge of the ammonia-thioacetamide reaction (5,6). EXPERIMENTAL
Reagents. Preparation of solutions of thioacetamide, ammonia, and sodium perchlorate was accomplished as described previously (6). Lead nitrate, ammonium nitrate, and potassium dichromate solutions were prepared by weight, and a sodium thiosulfate solution was standardized against the potassium dichromate solution. Disodium ethylenediaminetetraacetate dihydrate (EDTA) was dried at 80 "Cyand a weighed portion of the solid necessary for 1 liter of a solution of the desired concentration was dissolved in approximately 800 ml of distilled water. By addition of sodium hydroxide solution, the pW of the EDTA solution was adjusted to exactly 10 with the aid of a pH meter. Finally, the solution was transferred to a I-liter volumetric flask and diluted with water. Apparatus and Procedure. Apparatus was essentially the same as mentioned in a previous paper (6), except that the reaction vessel was equipped with a glass electrode and a saturated calomel reference electrode for measurement cf the pH of each reaction solution. A reaction mixture was prepared from distilled water and the stock solutions of lead(II), EDTA, ammonia, ammonium nitrate, and sodium perchlorate and adjusted, if necessary, to pH 10 with sodium hydroxide solution. Outlined elsewhere (6) are other details concerned with deaerating and thermostating the reaction mixture and with adding the thioacetamide solution to initiate the precipitation process. A two-step procedure was used to determine the initial concentration of ammonia in a reaction solution. First, immediately after the addition of thioacetamide, an aliquot of the reaction solution was titrated with standard hydrochloric acid to a methyl red end point. Second, a blank solution was prepared in which the concentrations of all substances present were the same as those in the reaction mixture, with the exception that both ammonia and ammonium nitrate were absent. A volume of this blank solution equal to that used in the first step was titrated with hydrochloric acid. The difference between the volumes used for the two titrations was employed to calculate the concentration of ammonia at the beginning of the reaction. This procedure for the determination of the ammonia concentration was verified with synthetic test solutions of known composition. To obtain measurements of the rate of precipitation of lead sulfide, samples of the reaction mixture were taken at timed intervals and were analyzed for lead(II) remaining unprecipitated. Each sample was filtered quickly through a fineporosity sintered-glass funnel into a suction flask placed in an ice bath. An aliquot of the clear filtrate was pipetted into a 250-ml flask, concentrated nitric acid was added, and the solution was boiled gently to destroy unreacted thioacetamide and to oxidize sulfide to sulfur. Next, concentrated sulfuric acid was added, the solution was boiled to remove excess nitric acid, and solid potassium persulfate was introduced in small portions to destroy EDTA as well as any remaining thioacetamide. Finally, the solution was evaporated to fumes of sulfuric acid, cooled, and diluted to a volume of 30 ml with distilled water. The resulting lead sulfate pre-
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Table I. Comparison of Initial Rate of Precipitation of Lead Sulfide with Initial Rate of Ammonia-Thioacetamide Reaction -4Pb(II)J/dt, -d[CHKSNHz]/dt: Initial concentrationsa mole/liter mole/liter . [Pb(II)I, M [EDTA], F WLI, F min X lo4 min X l o 4 Series 0.0250 0.58 1.84 1.85 1 0.0100 0.0200 0.58 1.80 1.85 2 0,0200 0.0250 0.57 1.73 1.79 3 0.0200 0 * 0250 0.58 1.86 1.85 4 0.0200 0.0250 0.74 2.95 3.01 5 0.0200 0.0250 0.97 5.21 5.18 6 0.0200 0.0350 0.58 1.78 1.85 7 0.0200 0.0450 0.45 1.16 1.11 8 0.0200 0.0500 0.45 1.18 1.11 9 0.0200 0.0500 0.58 1.83 1.85 10 0.0400 0.0500 0.63 2.16 2.18 11 0 0400 a In all experiments the initial concentration of thioacetamide was O.lOOF, the pH was 10, the temperature was 40 "C,and the concentration of ammonium nitrate was one fifth of the initial ammonia concentration; the concentration of free EDTA was the difference between the initial lead(I1) and EDTA concentrations. b Calculated from the rate expression -d[CW,CSNHz]/dt = ~'[CH~CSNHZ][NH~]~, where k' = 0.0055 liter2/mole2. min at 40 "C; see reference 6. 3
#
Table 11. Precipitation of Metal Sulfides by Thioacetamide in Ammoniacal EDTA Solutions0 Initial concentration Metal ionb of metal ion, M [EDTA], F Temperature, 'C Observations No precipitate after 2 hours 0.005 0.060 90 No precipitate after 2 hours 0.0015 0,001 90 No precipitate after 2 hours 90 0.001 0.010 No precipitate after 1 hour 0.001 0.0015 90 No precipitate after 1 hour 0.001 0.010 90 Rate of precipitation com0,005 90 0.006 parable to that of PbS Quantitative precipitation 25 0,005 0.006 within 5 minutes 25 0.005 0.010 Quantitative precipitation 25 0 .005 0.010 within 5 minutes Quantitative precipitation 25 0.001 0.002 within 5 minutes a In all experiments the initial concentrations of thioacetamide, ammonia, and ammonium nitrate were 0.100 0.60, and 0.12 F, respectively, and the pH was 10. * An aqueous stock solution of either the nitrate or chloride of each metal cation was prepared by weight and served as the source of the metal ion.
cipitate was collected, washed with cold dilute sulfuric acid, and dissolved in ammonium acetate solution, and the quantity of lead(I1) was determined according to the procedure published by Swift and Butler (1). In experiments with known amounts of lead(II), this method has an average error of only *0.2z;;. RESULTS AND DISCUSSION
Several preliminary considerations were necessary before any quantitative measurements of the rate of precipitation of lead sulfide could be attempted. Among these were a check on the possibility of a specific interaction between thioacetamide and EDTA, the establishment of an optimum reaction temperature, and a brief study of the effect of ionic strength on the rate of precipitation of lead sulfide. To determine if EDTA has any inhibitive or catalytic effect on either the rate of the hydroxide-catalyzed hydrolysis of thioacetamide to sulfide (11) or the rate of the ammonia-thioacetamide reaction, solutions containing various concentrations of thioacetamide and EDTA were prepared with and without ammonia and sodium hydroxide, and the disappearance of thioacetamide was followed spectrophotometrically at (11) E. A. Butler, D. 6.Peters, and E. H. Swift, ANAL.CHEM., 30, 1379 (1958). 2080
e
40,60, and 90°C as described in a previous publication (6). In
every solution examined, the rate of disappearance of thioacetamide in the presence of EDTA agreed quantitatively with the rate predicted when EDTA is absent. EDTA has no significant effect on the rate of generation of sulfide from thioacetamide. A set of experiments was performed at different temperatures and reactant concentrations to ascertain the optimum conditions for quantitative rate measurements. A reaction temperature of 40 "C provided conveniently rapid precipitation of lead sulfide. In addition, variations of ionic strength over the ten-fold range Erom 0.12 to 1.2 did not influence the rate of precipitation of lead sulfide, a result which is consistent with earlier observations concerning the arnmonia-thioacetamide reaction (6). Consequently, in view of the complex nature of the reaction solutions, the ionic strength was not rigorously controlled in subsequent work. Since previous studies of the ammonia-thioacetamide reaction revealed an abrupt increase in the third-order rate constant as the concentration of ammonia is decreased and since the hydroxidecatalyzed hydrolysis of thioacetamide becomes more important at low ammonia concentrations, all of the experiments reported in this paper were performed with initial ammonia concentrations greater than 0.45F,so that these two effects were minimized.
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Chloride is known to retard the precipitation of cadmium sulfide by thioacetamide in acidic media (12), because of the formation of less reactive chloro complexes of cadmium(I1). Considering the comparable stabilities of complexes which chloride forms with cadmium(I1) and lead(I1) as well as the possibility of mixed chloride-EDTA complexes of lead@), we substituted ammonium nitrate for ammonium chloride (used in earlier work) as a constituent of the reaction solutions to avoid any effect of chloride on the rate of precipitation of lead sulfide. The initial concentration ratio of ammonium nitrate to ammonia was 1 to 5 in all experiments to keep the pH of the reaction mixture essentially constant during the precipitation of lead sulfide. Rate of Precipitation of Lead Sulfide. Experiments were performed with various concentrations of ammonia, ammonium nitrate, EDTA, and lead(II), but with the same initial concentration o f thioacetamide. Table I compares the rate of precipitation of lead sulfide with the rate of formation of sulfide from the ammonia-thioacetamide reaction. The tabulated reaction rates are values corresponding to zero time. Initial rates of the ammonia-thioacetamide reaction were computed from the third-order expression observed in earlier studies (5,6),namely
- d[CH3CSNH23 = k'[CH3CSNH2][NHal2 dt
where k' is 0.0055 liter2/mole2.min at 40 "C. To obtain the rate of precipitation of lead sulfide, we subtracted the concentration of lead(I1) remaining unprecipitated in each sample of the reaction mixture from the initial lead(I1) concentration and divided this difference by the appropriate elapsed time from the beginning of the measurements. Since the rate of precipitation was essentially constant within experimental error throughout the duration of each kinetic run, the initial rate of lead sulfide precipitation was simply taken as the average of these results. For most of the kinetics studies, the formation of lead sulfide was followed for approximately 45 minutes until 35 to 40% of the lead(I1) had been precipitated. However, in experiments with initial ammonia concentrations of 0.74 and 0.97F, where the precipitation of lead sulfide was particularly rapid, the rate measurements were terminated after 30 minutes, in which time 45 and 78z,respectively, of the lead(I1) reacted. Several conclusions can be drawn from an analysis of the data shown in Table I. First, the rate of precipitation of lead sulfide is independent of the concentrations of free EDTA and the lead-EDTA complex and is controlled solely by the ammonia-thioacetamide reaction. Second, when the initial EDTA concentration is equal to or greater than the initial concentration of lead(II), the occurrence of a direct reaction between thioacetamide and lead(I1) cannot be detected, although one predicts from the findings of Swift and Butler ( I ) that the rate of the direct reaction in the absence of EDTA should be 6 X mole/liter smin at pH 10 and 40 "C for initiallead(I1) and thioacetamide concentrations of 0.02 and 0.10 M , respectively. In spite of the fact that the mechanism of the direct reaction between dipositive metal ions and thioacetamide has not been elucidated, it is evident that EDTA can protect lead(I1) from attack by thioacetamide (but not by sulfide ion produced from the ammonia-thioacetamide reaction). Third, the results of this investigation substantiate the proposal made in our previous study of the ammonia-thioacetamide reaction (6) (12) D. V. Owens, E. H. Swift, and D. M. Smith, Tulantu, 11,1521 (1964).
that the consumption of thioacetamide and the generation of sulfide both proceed at the same rate. In earlier work, the sensitivity of sulfide ion to air oxidation caused the apparent rate of formation of sulfide to be significantly less than the rate of disappearance of thioacetamide. However, in the presence of lead(II), which combines with sulfide as rapidly as it is produced, the two processes occur at identical rates. Some experiments were performed in which the initial concentration of lead(I1) was greater than that of EDTA. Lead (11) which is not complexed by EDTA precipitates as a dense, white solid whose composition was not identified in this study. It was assumed, provisionally, that the precipitate containing lead(I1) would react directly and immediately with a stoichiometric amount of thioacetamide to give lead sulfide, whereas the soluble lead-EDTA complex would combine with the sulfide ion formed from the reaction between ammonia and the excess thioacetamide. We observed that the rate of precipitation of complexed lead(I1) was always somewhat greater than the expected rate of the ammonia-thioacetamide reaction, because the white precipitate tended to coagulate at the bottom of the reaction vessel and to undergo relatively slow and incomplete metathesis to lead sulfide. consequently, since thioacetamide did not react rapidly with the white precipitate, additional thioacetamide was available for the ammoniathioacetamide reaction and the precipitation of complexed lead(I1) was correspondingly more rapid. Analytical Applications. From knowledge of the thirdorder rate constant for the ammonia-thioacetamide reaction (6),predictions can be made of the time required to precipitate lead sulfide quantitatively from a solution containing excess EDTA. For example, from 100 ml of a solution initially 0.5F in ammonia, 0.1F in ammonium nitrate, 0.lF in thioacetamide, and 0.01F in EDTA, one may achieve complete precipitation of 100 mg of lead(11) in approximately 36 minutes at 40 "C and in less than 4 minutes at 90 'C. Under otherwise identical solution conditions, the quantitative precipitation of 200 mg of lead(I1) would require twice as much time. Equally important from an analytical viewpoint is the possibility that EDTA can be employed as a masking agent to facilitate the sulfide separations of metal ions. Flaschka (9) found that mercury(II), bismuth(III), silver(I), lead(II), thallium(I), and tin(1V) are among the species whose sulfides are precipitated by thioacetamide from ammoniacal EDTA solutions, but did not state whether these metal ions combine with thioacetamide directly or with the sulfide formed from the ammonia-thioacetamide reaction. The same paper reported that precipitation of the sulfides of nickel(II), copper(II), cadmium(II), cobalt(II), tin(II), and thallium(II1) is either greatly retarded or totally prevented when EDTA is present in a solution containing thioacetamide and ammonia. However, each of the latter elements can be precipitated as the sulfide if a large excess of calcium ion is added to the solution to displace the metal ion from its EDTA complex. Unfortunately, for all of these experiments, details are lacking about the exact solution conditions as well as the relative rates of precipitation of the various sulfides. Semiquantitative measurements have been performed to determine the effect of EDTA on the rates of precipitation of several metal sulfides from ammoniacal thioacetamide solutions and to evaluate the potential usefulness of EDTA for the separation of metal ions. Results of these studies, which are summarized in Table 11, indicate that a metal ion can exhibit one of three modes of behavior: it may fail to react either with thioacetamide or with the sulfide ion formed from the ammonia-thioacetamide reaction ; it may combine only
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with the generated sulfide ion; or it may interact with thioacetamide directly as well as with sulfide ion. Nickel(II), zinc(II), and cadmium(I1) belong in the first category, being so effectively masked by excess EDTA that no sulfide precipitate is obtained during a 2-hour period at 90 "C. Interestingly, if a small amount of zinc(I1) is added to the solution after the ammonia-thioacetamide reaction has proceeded for some time, a zinc sulfide precipitate forms temporarily, but redissolves quickly as long as excess EDTA remains when the mixture is stirred. When the experiment is repeated with cadrnium(II), the resulting cadmium sulfide precipitate does not redissolve. Copper(II), along with lead@), displays the second kind of behavior, because the copper-EDTA complex does not undergo a direct reaction with thioacetamide and because the precipitation of copper sulfide appears to be controlled by the ammonia-thioacetamide reaction. We established these facts in a semiquantitative way by preparing a reaction mixture initially 0.005M in copper(II), 0.006F in EDTA, and 0.100F in thioacetamide at pZI 10 and by observing that at room temperature no precipitate of copper sulfide resulted from interaction of thioacetamide with the copperEDTA complex. However, when 0.60F ammonia was introduced, precipitation of copper sulfide commenced at a rate comparable to that of lead sulfide under similar conditions.
Inasmuch as the sulfides of mercury(II), silver(I), and bismuth (111) are precipitated rapidly at room temperature in the presence of EDTA, it is evident that a direct reaction occurs between these metal ions and thioacetamide and that these species are members of the third class of reactivity. A number of possible sulfide separations of metal ions can be proposed on the basis of the present study. Mixtures of metal ions, representing each of the three groups of characteristic behavior, should be separable through proper control of the concentrations of EDTA and ammonia and the reaction temperature. As a brief example, one might separate cadmium(II), lead(II), and mercury(I1) by adding an excess of EDTA at pH 10 to complex each of the metal ions and then excess thioacetamide to precipitate mercury sulfide at room temperature, by introducing ammonia and heating the sample at 90°C to form lead sulfide, and by adding excess calcium ion or by acidifying the solution to obtain cadmium sulfide. Future research must include investigations of the kinetics and mechanism of precipitation of other metal sulfides in the presence of EDTA, quantitative tests of the sulfide separations suggested by this work, and examinations of coprecipitation phenomena.
RECEIVED for review August 6, 1969. Accepted October 2, 1969.
K. E. Shuping, G. R. Phillips, and A. A. Moghissil Southeahtern Radiological Health Laboratory, P. 0. Box 61, Montgomery, Ala. 36101 KRYPTON-~Sfound in the atmosphere originates from atmospheric nuclear weapons tests and, more recently, from the operation of nuclear reactors including fuel reprocessing operations. Horrocks (1) first proposed the application of liquid scintillation to the determination of radioactive gases. Setszer et al. (2) and Sax et al. (3)used plastic scintillators for krypton counting. Cohen et al. measured the concentration of * K r dissolved in water using a dioxane-based scintillation solution (4). Curtis et al. (5) reported a liquid scintillation counting technique suitable for environmental 8GKr analysis. It consisted of introduction of krypton with a pressure of 25 mm Hg into a 25-ml evacuated plastic vial and subsequent addition of a toluene-based scintillation solution. This technique suffers from two disadvantages. The amount of krypton which may be introduced into the solution is limited to less than 1 mI (at STP) and the reproducibility of the technique is poor. Dis-
Present address, Southwestern Radiological Health Laboratory, P. 0.Box 15021, Las Vegas, Nev. 89114 ~~
(1) D. L. Horrocks and NI. H. Studier, ANAL.CHEM.,36, 2077 (1964). (2) J. L. Setser, T. C. Rozzell. and B. L. Smith. Radiochiin. Acta. 8, 18 (1967). (3) . . N. I. Sax, J. D. Dennv. and R. R. Reeves. ANAL.CHEM.. , 40., 1915 (1968). (4) J. B. Cohen, J. L. Setser, W. D. Kelley, and S. D. Shearer, Talanta, 15,233 (1968). (5) M. L. Curtis, S. L. Ness, and L. L. Bentz, ANAL.CHEM., 38, 636 (1966). _
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solved krypton escapes from the solution to the void space of the vial and thus changes the counting efficiency. This paper describes a reproducible liquid scintillation technique which is considerably more sensitive than the previously described methods. EXPERIMENTAL
Materials. A Beckman liquid scintillation counting system was used in this investigation. Scintillation grade solutions were prepared by dissolving 7 g of 2,5-diphenyloxazole (PPO) and 1.5 g of p-bis-(0-methylstyry1)-benzene (bis-MSB) per liter of toluene or per liter of p-xylene. The PPO and bis-MSB were received from Pilot Chemicals of Boston, Mass. All other reagents were analytical grade. Counting vials were made of borosilicate glass and were fused to male h e r fittings. The diameter and height of the vials were dictated by the specifications of the counting system. The volume was approximately 25 ml per vial. Stopcocks, couplings, and caps utilizing standard h e r taper fittings were obtained from the Hamilton Company, Whittier, Calif. Pure krypton samples were received from Cryogenic Rare Gas Laboratories, Newark, N. J.; Air Products and Chemicals, Inc., Emmaus, Pa.; and Union Carbide Gorp., kinde Division, East Brunswick, N. J. All samples had a known collection date. Procedure, Handling of krypton was carried out by use of a manifold connected to a closed-end U-lube manometer. The volume of the entire system, excluding the counting vials was approximately 10 ml. Krypton was introduced into the manifold and the attached counting vials to a pressure of 400--600 mm Hg. The pressure of the system was
ANALYTICAL CHEMISTRY, VOL. 41, NO. 14, DECEMBER 1969
