Coprecipitation of Thallium(!) with Silver Chloride Precipitation from Homogeneous Solution LOUIS GORDON, J. I. PETERSON', and B. P. BURTT Department of Chemistry, Syracuse Univers;ty,
Svracure 10, N. Y.
T h i s investigation was u n d e r t a k e n t o develop a m e t h o d for the precipitation of silver chloride from homogeneOUB solution and to s t u d y the copreeipitation of thallium(1) with this oarrier. Large crystals of silver chloride were produced. The thallium-silver ratio in the crystal was f o u n d to be dependent on the solution concentration of thallium. Under the experimental wnditions used the mole r a t i o of thallium to gilver in the precipitate is of the order of 10-7. The results indicate adherence t o neither the homogeneous nor heterogeneous distribution laws. However, a pseudohomogeneous distribution is obtained because so small a fraction of t h a l l i u m is wprecipitated f r o m a solution initially 10-M in thallium that the thallium-silver ratio in the crystal remains essentially constant t h r o u g h o u t the entire precipitation proeess. The gmss picture is thus one of an apparent homogeneous distribution of thall i u m within the crystals of silver chloride.
R
ECENT studies (2,,S, 16, 90, $ 1 ) in caprecipitation have utilized ' . the technique of precipitation from homogeneous solution ( I , $9). This paper reports a study of the coprecipitation of thallium(1) with silver chloride precipitated from homogeneous solution. Silver ions were slowly released from the silver ammonia aomplex in the presence of chloride. The complex is destroyed with hydrogen ions produced by the slaw hydrolysis of D-hydroryethyl acetate (18). Photomicrographs of silver chloride produced by this technique are shown in Figure 1. Thallium(1) chloride is known to coprecipitste with silver chloride in the ordinary analytical precipitation (8); adsorption studies have also been made with thallium(1) on silver halides (6, 7, 1 2 ) . Although thallium(1) chloride does not mix isamorphously with silver chloride ( l e ) ,i t is only slightly soluble; the thallous ion does not precipitate in ammoniacal medium. These and other factors led to the choice of thallium(1) as the ion to be used in this eoprecipitation study.
pylene chlorohydrins were found to produce chloride ion a t a desirable rate and t o result in complete precipitation of the silver. To obtaingoad results with these chlorohydrins (Matheson Co)., these reagents were first purified by distillation and then passed through a moist ion exchange column containing IRA-400 resin (Rohm and Haas) in the free base form. The initial water-containing eluant wa8 discarded, and the chloride ion-free reagents were used. These chlorohydrins hydrolyze very slowly a t room tempera, tnre; only a small amount of precipitate will be obtained in a week'stimefrom 100ml.ofrtsalutioncontainin~O.l~amofsilver and 10 ml. of the reagent. However, st 60" C. quantitative precipitation can be effected within a few hours. The precipitate is white in the absence of light and consists of fine crystalline partioles 2 to 8 microns in diameter. The method shows promise for the determination of silver, as was indicated by a limited gravimetric study. Another approach t o the problem resulted in the production of large crystals of silver chloride, some of which were of the order of 0.2 mm. in diameter. I n this method the silver cation was homogeneously released from it8 ammonia. complex, in contrast to the previous method whereby only chloride ion was released from a reagent. Ethylene chlorohydrin was first tested, because it produces upon hydrolysis the necessary hydrogen ion for the destruction of the silver ammonia.complex as re11 as chloride ion for precipitation. However, a more desirable method of precipitation was evolved by adding chloride ion to the solution containing the silver ammonia complex and effecting the litter's destruction through the hydrolysis of n water soluble ester such ~8 8-hydroxyethyl acetate. The use of the ester is advantageous in that i t results in simplercontmlof chloride ion concentration. I t ma8 thismethod which was adopted for the copreoipitation study. Determination of Thallium in Silver Chloride Precipitate. Preliminary work showed that the amount of coprecipitated thallium was smell. Themolar ratio of thallium to silver in the precipitate wa6 of the order of 5 X 10 -7. This necessitated a determination ' ' "' of thallium in the precipitate by the usc of radioacti
PRELIMINARY STUDIES ON PRECIPITATION OF SILVER CHLORIDE FROM HOMOGENEOUS SOLUTION
Several organic compounds containing chlorine were tested to determine whether they would release chloride ion into an aqueous solution by hydrolysis. These were ethylene chlorohydrin, 3-chloro-1,Z-propanediol (propylene chlorohydrin), chloroaeetic acid, dichloroacetic acid, trichloroacetic acid, chloroacetiamide, triglycol dichloride, methylene chloride, chloral hydrate, and ethyl ch1oroexrhonat.e. Of these, only ethylene and prar Present address. Sylvania Eleotric Products, Ino.. Towanda, Pa.
b.
Figure 1. Photomicrographs of silver chloride precipitated from honiogeneous solution Di-t illumination 6QoC. in 0.0118M chloride b. Direct illumination; 25' C. in J.WM chloride
0. C.
Vertical illumination, 25' C. in 0.0118M chloride
1770
1771
V O L U M E 2 7 , NO. 1 1 , N O V E M B E R 1 9 5 5 THALLIUM-204. Thallium-204 nitrate in better than 99% radiochemical purity, but not carrier-free, was obtained from the Oak Ridge National Laboratory for this work. A solution was treated with sulfur dioxide to ensure that all the thallium was in the unipositive state. ilfter removal of excess sulfur dioxide by boiling, the thallium solution was diluted and stored in a polyethylene bottle. Total thallium was determined by spectrophotometric comparison ( 4 ) with a gravimetrically standardized (9) thallium solution. RADIOACTIVE ASSAYO F STliSDARD SOLUTIOS. The assay Of the radioactive thallium solution was carried out by an adaptation of the procedure of Berg and coworkers ( 2 7 ) for the determination of thallium. The procedure was as follows: . An aliquot of the standard radiothallium solution was delivered with a calibrated micropipet and diluted with 1 ml. of water. Then 5.00 ml. of standard thallium(1) nitrate solution (1 ml. = 1 nig. of thallium) were added followed by 1.0 ml. of 10M sodium hydroxide solution. The solution was diluted to 10 ml. and heated to 60' C. Then 2.0 ml. of thionalide solution, 1% in acetone, freshly prepared, were added and the solution was stirred and allowed to cool to room temperature by standing for 1 hour. Sext, 1 to 2 drops of Anti-Creep solution, Schleicher and Schuell, were added and the solution was again stirred. A fine jet of water served to stir the solution whenever necessary. The precipitate was filtered through a Royal Berlin A-3 porcelain crucible, washed -xith water, then n-ith acetone, and finally air-dried. This method of precipitation leaves less than 0.1% of the original thallium in the filtrate as was determined by the spectrophotometric method ( 4 ) . The crucible n-ith its thionalide precipitate was then positioned under a Geiger-Muller tube and t,he activity determined. The precision of this method is shown in Table I. Because the subsequent determination 11-as to be that of thallium in silver chloride, the latter was added in some of the experiment,s given in Table I. The silver salt was dissolved with potassium cyanide and the thallium determined as in the preceding procedure. The oliserved specific activity of samples of the standard radiothallium solution measured by this method was found to be 187 X l o 2count's per minute per microgram of thallium.
Table I.
Precision of Radioactive Method for Thallium
Deviation from Over-all Counts per Minutea Average Average, % 11, 7253, 7266, 7202, 7238 7240 1.1 ?b 7256, 7360, 7345, 7365 7381 0.8 7470 2 1 3b 7519, 7439, 7480, 7441 7398 1.1 4b 7452, 7395, 7357, 7366. 7421 7371, 7364, 7290, 7240 7315 0 0 7153, 7226, 7137, 7180 7174 2.1 7C 7340, 7345, 7316, 7229 7307 0.3 8C 7274, 7282, 7377, 7188 7280 0.5 All counts were made over a 10-minute period. b These samples contained thallium but no silver chloride. C These samples contained 0.1 gram of silver chloride as well as thallium. NO.
i:
Table 11. Precipitation at 60" C. Time of Yield Preoipitaof SgC1, tion, Hours Grain 3.0 0.1114 4.0 0.1232 5.5 0.1224 23.0 0.1312 Weight ratio.
Tl/Ag in Precipitatea X 105 0,0955 0.140 0.150 0.201
Fraction of Ag Precipitated 0.838 0.928 0.922 0,988
Fraction of T1 Precipitated X 108 0.081 0.131 0.139 0,200
Determination of Thallium in Precipitate. The precipitate of silver chloride containing the radiothallium was dissolved by the addition of 0.5 gram of potassium cyanide and 1 ml. of water. Five milligrams of stable thallium were then added and the determination was made as before. The amount of radiothallium in the precipitates obtained from the coprecipitation experiments was then determined by reference to the standard value previously obtained. I n these determinations, and in the previous ones as well, cor-
rections were made for dead time and for radioactive decay of the thallium. PRECIPITATION PROCEDURE
Precipitations were carried out in 250-ml. polyethylene beakers in a water bath controlled to =!= 1 O C. Polyethylene vessels Tere used to minimize small adsorption losses of thallium on glass. The solutions were stirred by means of a glass rod passed through a stopper inserted in the beaker. The entire precipitation assembly was enclosed in a light-tight box. The solutions Contained, in a 100-ml. volume, 0.1 gram of silver as nitrate, ammonium chloride, ammonium hydroxide, p-hydroxyethyl acetate, and thallium sulfate. ,411 chemicals were of reagent grade and were used without purification except in the case of p-hydroxyethyl acetate. The latter (Distillat,ion Products Industries) was distilled and the fraction boiling between 180" and 190" C. was used. For the coprecipitation studies solutions were prepared as. described above; the reaction was allowed to proceed for various periods of time as noted in the subsequent tables. I n this manner different fractions of the total silver in solution were precipitated. The precipitate was filtered through a porcelain crucible, washed with sufficient water to remove mother liquor, dried, and weighed. A known weight of the solid phase was then analyzed for thallium. A previous experiment showed that washing the precipitate did not affect the thallium-silver ratio of the crystal. The fractions of silver precipitated were in general above 0.25 since it was observed especially a t 25' C. that when fractions less than this were precipitated a supersaturation effect caused the spontaneous formation of small neJv particles during filtrat,ion; this did not occur a t 5' and 15' C. The thallium-silver ratios were determined for the various fractions of total silver precipitated a t several temperatures and in solutions of different initial chloride ion concentration. RESULTS
Coprecipitation at 60' C. Precipitations were made in a 100ml. volume containing 0.0098 gram of silver as nitrate, 0,00118 mole of ammonium chloride, 99.3 y of radiothallium, 2.0 ml. of 14X ammonium hydroxide, and 6.0 ml. of 8-hydroxyethyl acctate. The data are s h o w in Table 11. The limited data indicate a mther rapid increase in thallium content of the crystal as a function of fraction of silver precipitated. This was not observed in the subsequent work a t lower temperatures. This observation is apparently intimately connected with the physical characteristics of the precipitate, m shown by Figure 1 where the spongy structure is in sharp contrast with the cubes obtained a t the lower temperatures. Coprecipitation at 25" C. a t Initial 0.01181M Chloride Ion Concentration. Precipitations were next performed a t 25" C. in order to obtain crystals of silver chloride with improved physical characteristics and to coprecipitate increased amounts of thallium(1) chloride because of the large decrease in solubility of the latter with temperature. In these experiments the solutions were identical ivith those used a t 60" C., except that either 0.0998 or 0.1003 gram of silver and 76.0 y of thallium were used. Figure 1 shoxvs a photomicrograph of the crystals. The data (Nos. 1 to 19) are given in Table 111. I n this series of experiments the composition of the precipitate, as expressed by t'he ratio of micrograms of thallium per gram of silver, changed very slightly with increase in fraction of carrier precipitated. The ratio decreased somewhat, but as might be expected it was very slight; because so little thallium is coprecipitated, its concentration remains essentially constant regardless of the fraction of carrier which has been precipitated. However, the concentration of silver can increase almost fivefold aa chloride is removed from solution by precipitation. Thus, a competitive adsorption effect between silver and thallium during crystal grorvth could explain the apparent slight decrease in thallium coprecipitation as the fraction of carrier precipitated becomes larger. Coprecipitation at 25' C. a t Initial 0.100M and 1.OOM Chloride Ion Concentrations. I n order to determine if a competitive adRorption effect caused t,he slight decrease in thallium coprecipitation with increase in fraction of carrier precipitated as not,ed in
ANALYTICAL CHEMISTRY
1772 the previous experiments, the concentration of initial chloride ion was increased. This resulted in a decrease of the silver ion concentration during precipitation acrording to the solubility product relationship. However, as the initial chloride ion roncentration increased, the change in chloride during precipitation-and thus the change in silver ion concentration-was not so great. Precipitations were made with solutions containing the following in 100 ml.: 0.0998 gram of silver, 0.0100 mole of ammonium chloride, 76.0 yof thallium, 4.0 ml. of 14M ammonium hydroxide, 12 ml. of 8-hydroxyethyl acetate; and 0.998 gram of silver, 0.100 mole of ammonium chloride, 76.0 y of thallium, 7.0 ml. of ammonium hydroxide, and 12.0 ml. of B-hydroxyethyl acetate. The crystals are shoan in Figure 1. At these increased chloride ion concentrations-i.e., a t 0.1OON
Table 111. Coprecipitation Results ~~
No 1 2 3 4 5 6 7
8 9 10 11 12 13 14 15 16 17 18 19
Fraction Fraction of Tl/Ag Yield of of A g T1 Pptd. of AgC1, in Ppt.' Pptd. x 103 x 105 Hr. Gram Initial Chloride Ion Concentration = 0.0118M; t = 25' C. Reaction Time,
5.1
6.0 6.3 6.9 9.7 10.0 13.0 16.2 20.0 22.0 24.0 25.0 30.2 43.5 47.3 61.3 67.0 75.3 96.3
0.05966 0.0569b 0.0561 0.0589 0.0797b 0.09OOb 0.0976 0.0972b 0.1103 0.1204b 0.1144b 0.1183 0.1255 0.1293b 0.1318 0.1338 0.13286 0.1337 0.1335 Av.
0.621 0,710 0.622 0.783 0.788 0.945 0.937 0,965 1.33 1.04 0.935 1.14 1.35 1.05 1.28 1.08 0.988 1.26 1.48
0.107 0.126 0.112 0.134 0.100 0.106 0,0968 0.101 0.121 0,0877 0.0838 0,0969 0.108 0,0823 0.0985 0.0815 0,0754 0.0950 0.112 0 101 + 0.012
20 21 22 23 24 25 26 27
Initial Chloride Ion Concentration = 0.100.M; t = 26' C . 2.8 0.144 0.0435c 0.327 0 570 4.3 0.113 0 810 0.0735C 0.553 1 27 0.140 6.0 0.694 0.0923C 8.5 0.818 1 34 0.126 0.1088C 11.1 1 55 0.1153C 0.868 0.138 19.3 0.959 0.12730 1 60 0.130 0.1311C 30.0 0.987 1 75 0.136 46.0 0 997 2 02 0.155 0.1325C A r . 0.1.35 + 0.009
28 29 30 31
Initial Chloride Ion Concentration = 1.00M; t = 25' C. 4.2 0.0346C 0.146 0.261 0.382 6.3 0.0646C 0.133 0.486 0.830 10.0 0.0877C 0.167 0.660 1.42 23.3 0.1126C 0.0995 0.847 1 11 A r . 0.136 k 0.021
32 33 34 35 36 37 38 39
[nitial Chloride Ion Concentration = 0.0118.1.1; 1 3.9 0.0099c 0.121 0.075 5.0 0.0329C 0.116 0.248 7.0 0.0525C 0.133 0.395 9.4 0.0681C 0.128 0.512 17.3 0.0942C 0.136 0.709 24.2 0.1085C 0.137 0.816 42.3 0.1265C 0.132 0,952 42.3 0.1261C 0.132 0.950 Av. 0.129 i: 0.006
=
15O C 0.109
0.378 0.692 0.862 1.26 1.46 1.64 1.65
and I.OOM-the respective silver ion concentrations are decreased by large factors. Table I11 (Sos. 20 to 31) shows that the thallium content of the precipitate is slightly greater than was obtained with the experiments with 0.0118M chloride solutions. However, the increase in coprecipitution of thallium is small compared to the decrease in silver ion concentration, and, if anything, one must conclude t,hat the amount of thallium coprecipitated is essentially independent of both silver and chloride ion concent,rat,ions within the limits studied. The data also shox an increasing trend of thallium coprecipitation with fraction of silver coprecipitated. Coprecipitation at 15' C. at Initial 0.0118M Chloride Ion Concentration. The solutions in this series of experiments contained the following in 100 ml. of solution: 0.0998 gram of silver, 0.001 18 mole of ammonium chloride, 2.0 ml. of 14M ammonium hydroxide, 10.0 ml. of 6-hydroxyethyl acetate, and 76.0 y of thallium. The rate of precipitation a t 15" C. is not appreciably different from that a t 25' C. Table I11 (Nos. 32 to 43), shows that the coprccipitation of thallium is essentially the same as at 25" C. Some of the precipitations a t 15' C. were carried out in 0.1OOM sodium nitrate, but again the data are essentially unchanged. Coprecipitation at 5" C. at Initial 0.00118MChloride Ion Concentration. These precipitations were carried out in 100.0-ml. volumes containing 0.1000 gram of silver, 0.00118 mole of ammonium chloride, 2.0 ml. of 14M ammonium hydroxide, 10.0 ml. of p-hydroxyethyl acetate, and 59.6 y of thallium. The data are shown in Table I11 (Nos. 44 to 51). The crystals are shown in Figure 1. The data indicate a somewhat greater thallium coprecipitation than in the experiments a t 15' and 25' C. The variability in the data makes it impossible t o determine whether there is an increase or decrease in thallium coprecipitation with fraction of silver precipitated. Coprecipitation at 25" C. with Varying Thallium Concentrations. These experiments were performed in a manner identical with the others carried out a t 25' C. in 0.0018M initial chloride solution, except that the thallium cont,ent was varied. The data are shown in Table IV. All of the precipitations were allowed to proceed for about the same length of time, so that the fraction of silver precipitated was in the range 0.77 to 0.79. The data show that the ratio of thallium to silver in the crystal is approximately proportional to the concentration of thallium in solution. Evidence of Internal Distribution of Thallium within Silver Chloride. Two precipitations were performed a t 15" C., as previously described for the work at that temperature. While the precipitates were on the filtering crucible they were partially dissolved with dilute ammonia, while suction was applied. -4fter 70% of the original weight had been dissolved away, the remaining crystals contained, respectively, 0.125 and 0.154 y of thallium per gram of silver as compared to 0.132 (cf. Table 111). This indicates internal distribution of thallium rather than adsorpt'ion on the outer surface of the crystals of silver chloride.
Initial Chloride Ion Concentration = 0.0118M, also 0.10051 in N a 4 0 r ; t = 150
40 41 42 43
5.5 11.4 21.8 45.3
0.0606C
0.0921C 0.1131: 0.1310 Av.
0.127 0.456 0.123 0.693 0.114 0.851 0.113 0,985 0,119 =t0.006
0,762 1.12 1.28 1.47
Initial Chloride Ion Concentration = 0.0118M; L = 5' C. 9.0 0.0265d 0.162 0.199 0.541 11.9 0.0384d 0,103 0,289 0.500 34.5 0.0815d 0.234 0.613 2.40 31.0 0.0839d 0.179 0.707 2.12 47.5 0.1108d 0.21') 0.833 3.07 0.1231d 0.181 0.927 2.81 70.0 86.0 0.1276d 0.960 2.6R 0.166 194.0 0.1339d 0.162 1.008 2.73 Av. 0.176 0 026 Weight ratio. b Initial silver = 0.0998 gram: all others above = 0.1003 gram of silver. C Initial silver = 0.0998 gram. d Initial silver = 0.1000 gram.
44 45 46 47 48 49
fy
DISCUSSION
c.
+
Because silver and thallium(1) chloride are not isomorphous, it would be expected that neither the homogeneous nor heterogeneous distribution laws would be followed. However, Hahn's classification ( 5 ) of coprecipitation considers anomalous mixed crystal formation as a possibility. That this is not the present case is shown in Table V, which summarizes the distribution coefficients (19) calculated for the present work. The values of the homogeneous distribution coefficient, D, xere calculated in the usual manner from (?=carrier r)
PPt.
carrier
aoln.
(1)
Likewise, the values of the het,erogeneous distribution coefficient, X, were calculated from th? Doerner-Hoskins equation.
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V O L U M E 27, NO. 11, N O V E M B E R 1 9 5 5 tracer during its gro*-th
Table IV. Coprecipitationat 25” C. with Varying Thallium Concentration
In view of the equilibria involved, the carrier ion was considered to be free silver ion rather than total silver (or the ammonia complex). The usual integrated form of the Doerner-Hoskins equation
would erroneously lead to a negative distribution coefficient, since the concentration of silver ion increases, because of removal of chloride, as the fraction of total silver precipitated increaPes. Thus, for the present system, Equation 2 should be:
a
Concentration of T1, Y h l .
Tl/Ag in Precipitateo
Fraction of T1 Precipitated
1.52 0.760 0.456 0 304 0 . 162 0,0780 0.0152 0,00152 0.000152
1.36 1 01 0 503 0 20 0 11 0 oc)2 0 021 0 007 0 003
0.695 1.05 0.86 0.40 0.54 0.95 1.1 3.8 15.0
x
106
Table V.
c.
ti0
25 25 2.5 ~.
15 15
Summary of Distribution Coefficients” [C171 Molarity
x
0.0118 0 0118 0.100 1.00 0 0118 0.0118b 0.0118
x
(Tl, - T1) = moles TI in solution (CI, - Ag) = moles Ag in solution T1 = moles T1 in precipitate g = moles iig in precipitate Clc = moles C1 initially added Ksp = solubility product of silver chloride ( I O ) x = distribution coefficient TIC = moles T1 initially added v = volume in liters l h e values of in Table V were calculated using Equation 5. The data of Table V show conformity to neither distribution law. This is especially evident from the data a t 25’ C., where an extensive range of chloride ion concentration was employed. Further calculations have shown that neither distribution law is obeyed even if “total silver” is considered, thus permitting the use of Equation 3. The present system might have confornied to one of the distribution laws-and thus be anoinalous mixed crystal formation according to Hahn’s classification-had it been possible to work with thallium concentrations of the order of 10-10Af. One of the unique aspects of the present system is that the concentration of the minor constituent, thallium, is much greater than that of the carrier ion, uncomplexed silver. This is in sharp contrast with other distribubion studies in which the carrier ion concentration is much greater than that of the tmcer. For the present system, coprecipitation can be considered t o be internal adsorption as specified by Hahn ( 5 ) ,where the thallium is adsorbed by the silver chloride as the latter is being formed during the precipitation process. As the crystal continue? to grow, each succeeding layer covers the adsorbed layer. This process is essentially that of occlusion, as Kolthoff ( I S , f 4 ) uses this term, except, that the aging process which, according to his views perfects the imperfect crystals first formed, probably assumes a much lesser role here in precipitation frnni homogeneous solution. The Langmuir equation (15)predicts that the adsorption of an ion \vi11 be proportional to its solution concentration if the fraction of surface covered hy t3hision is small. This apparently is the case here with thallium, n-here the observed mole ratio of thallium to silver in the precipitate is of the order of 5 x 10-7. Likewise, in t,he case of silver which should be strongly at,tracted to the silver chloride crystal, the Langmuir equation predicts that its adsorption would be nearly independent of its solution concentration. This essentially occurs in the present case ivhere the silver ion concentsration was varied by altering the chloride ion concentratmion. If the competitive adsorption of silver and thallium TVAI’P tlcpenti~nt,on the silver ion concen-
~~
-
x
109
3 . 3 t o 5.3 3 . 1 t o 10 0 . 6 4 =k 0 . 0 5 0.04.5 t o 0.073 1.3 =t 0 . 2 1.3 =t 0.05 0.4 t o 1.1
5 a C‘alrulated from data given in ~ i r e r i o u stables. b .I180 0.100.11 in sodiiiin nitrate.
where
10s
Weight ratio.
Temperatine,
Integration leads to:
x
10s
4 . 3 t o 10 2.8 i 0.3 0 . 3 2 =k 0 . 0 2 0 . 3 0 =k 0 . 0 4 0.96 t o 4 . 0
1.5 to 3 7 0 . 4 to 1.6
___
- -~~
D
-__
~. -
tration, a variation in the extent of thallium coprecipitation should have heen expected. However, thallium coprecipitation remained essentially constant with variation in silver ion roncentration. The data of Table Vindicate adherence to the equ:~tioii:
(E)
=
K,[Tl]I/n s K I I T 1 ]
crystal
For experiments under a given set of conditions, such as in Table 111, where hut a small fraction of the initial thallium is coprecipitated so that the concentIration of thallium reinains essentially constant throughout the entire precipitation process, this equation reduces to:
This equation actually describes a “homogeneous” distribution of the thallium, such as would normally characterize a system adheriiig to the homogeneous distribution law (19). However, in the present case, homogeneous distribution is attained because each layer of siiver chloride adsorbs essent,ially tmhesame small amount of thallium from the solution which contains an essent8i,zllyuiivarying thallium concentration. The gross picture is thus one of an apparent homogeneous distrihution of thallium ~vithinthe crystals of silver chloride. However, it is possible that a heterogeneous distrihution will be observed in the case byhere the snlution concentration of t,hallium decreases significant1)- diiring the precipitation process. ACKNOWLEDGMEYT
The authors wish t o thank the Atomic Energy Commission, the Research Corp., and the Society of Sigma Xi for their support of this investigation. They also wish to thank Stanley Gedansky for aid in the preparation of the photomicrographv. LITERATURE CITED
(1) Gordon, L.,ANAL.CHEM.,24, 459 (1952). (2) Gordon, L., Reimer, C . C.. and Burtt, B. P.,Ibid., 2 6 , 8 4 2 (1954). (3) Gordon, L., Teicher, H., and Burtt, B. P., Ibid., 26, 992 (1954). (4) Haddock, L. A., Analyst, 60,394 (1935). (5) H a h n , O., “Applied Radiochemistry,” pp. 644164. Cornell University Press, Ithaca, N. Y . , 1935. ( 6 ) H a h n , O . , 2. angew. Chsm., 4 3 , 871 (1930). (7) H a h n , O., and Imre, L., 2. ph,ysik. Cheni.. A144, 161 (1929).
1774
ANALYTICAL CHEMISTRY
(8) Hillebrand, W. F., Lundell, G. E. F., Bright, H. A., and Hoffman, J. I., “Applied Inorganic Analysis,” 2nd ed., p. 205, Wiley, New York, 1953. (9) Ibid., p. 478. (10) Hodgman, C. D., “Handbook of Chemistry and Physics,” 34th ed., p. 1559, Chemical Rubber Publishing Co., Cleveland, 1952-3. (11) lime, L.,2. physik. Chem., A146, 41 (1930). (12) “International Critical Tables,” vol. IV, p. 53, McGraw-Hill, New York, 1926. (13) Kolthoff, I. hl., J . Phys. Chem., 36, 860 (1932). (14) Kolthoff, I. M., and Sandell, E. B., Ihid., 37, 723 (1933). (15) Prutton, C. F., and Maron, S. H., “Fundamental Principles of Physical Chemistry,” pp. 6 5 2 4 , Rlacmillan, New York, 1944.
(16) Salutsky, M., Stites, J. F., and AIartin, A. R., ANAL.CHEV.. 25, 1677 (1953). (17) Sandell, E. B., “Colorimetric Determination of Traces of Metals,” 2nd ed., p. 562, Interscience, S e w York, 1950. ~. 25, 1519 (1953). (18) Stine, C. R., and Gordon, L., A 4 i v ~ 4CHEM., (19) Wahl, A. C., and Bonner, X . A,, “Radioactivity Applied to Chemistry,” pp. 105-7, Wiley, New York, 1951. (20) Weaver, B., A N i L . CHEX., 26, 477 (1954). (21) Ibid.,p. 479. (22) . , Willard, H. H., Ibid., 22, 1372 (1950). RECEIVED for review r e b r u a r y 26, 1955. .\crepted .Tuli 19. 1955. Presented before the Division of Analytical Chemistry a t tile 127th meeting of the AMERICAN CHENICAL SOCIETY, Cincinnati, Ohio, A\Iarcti IEA.
Titrimetric Assay of Trichloroacetate W. A. SCHNEIDER, J ~ . , a n d L. E. STREETER The Dow Chemical Co., Midland, Mich. A simple and accurate assay method based upon the decarboxylation of trichloroacetate is presented. A known amount of standard acid is added to the neutral sample, the solution is refluxed for at least an hour, and the residual acid is titrated. The acid consumed is a direct measure of the trichloroacetate content. Contaminants usually present in commercial grades of the acid or salt do not interfere and need not be determined. The proposed method requires much less time than the classical alkaline hydrolysis procedure.
T
H E accepted method for the assay of trichloroacetate is essentially a chloride determination corrected for the chloride contributed by the major contaminant, dichloroacetate. The total chloride is determined following a Parr peroxide bomb decomposition or an alkaline hydrolysis. The latter process also converts any dichloroacetate present to oxalate, which is then determined by the usual calcium oxalate precipitation and subsequent permanganate titration. This method, which mav be credited to Pool ( 4 ) , has been largely developed by industrial laboratories without publication. In this laboratory, it has been noted that different analysts could not always agree on trichloroacetate assay, the usual point of difference being the dichloroacetate determination. A study by Dalin and Haimsohn ( 1 ) discussed the errors of dichloroacetate determination as applied to the assay of monochloroacetate, but these authors found that low results were related to the quantities of reagents used. However, in the present work, high values of dichloroacetate were the major concern.
Table I.
Substance I Added None
Table 11. Effect of Contaminants ’$ of NaTCA Taken,“ NaTCA Found,
Mixture
G.
%
G. Recovery ... 2.053 2.051 99.9 ... 2,492 2.492 100.0 2.492 2 492 100.0 NaHCOJ 1.98 2.508 2,510 100.1 4.63 2.508 2.510 100.1 N~~HPOI 1.07 2.492 2.495 100.1 2.05 2,508 2.510 100.1 5.43 2 508 2.510 100.1 NaDCA 11.6 ‘ 2.053 2.051 99.9 31.9 2,053 2.049 99.8 50.2 2,492 2.473 99.2 NahICA 9.5 2.053 2.049 99.8 19.2 2.492 2.484 99.6 40.5 2.492 2.468 99.0 51.7 2.492 2,380 95.5 Weights of KaTCA taken were computed on the basis of assayed salt purity of 99.6%.
...
of trichloroacetate in aqueous solution. I n some of these, the reaction rate was followed by titration of the residual acid (8, 3). The present method is essentially a refinement of this principle and is defined by the following equation:
Although the procedure described was developed specifically for the assay of technical sodium trichloroacetate, the general method is not restricted to the salt, as it has been used with equal success for the acid and the ethyl ester.
Standardization of Decarboxylation Method
Standard Sample Trichloroacetic acid Ethyl trichloroacetate Sodium triohloroacetate Sodium trichloroacetate
% Compound Found 99.4,99.8,99.5,99.7 100 2 , 9 9 . 6 , 9 9 . 9 , 9 9 9,99 6 99.7,99.6,99.7,99.6 99.7,99.7,99.6,99.5
Samples of trichloroacetic acid, known to be free of dichloroacetate by infrared and freezing point data, invariably showed a dichloroacetate content of from 2 to 3%, which was found to be a function of the sodium hydroxide concentration used for hydrolysis. The conclusion reached was that the alkaline hydrolysis also converts some trichloroacetate to oxalate, thus causing spuriously high values for dichloroacetate. Clearly, a new assay method was needed for trichloroacetate. Several kinetic studies have been made of the decarboxylation
APPARATUS AND REAGENTS
Flasks, 250 ml., flat-bottomed, short-necked T24/40. Condenser, 50-cm., water-cooled, with glass joint $24/40. Methyl red indicator, 0.1% solution in 95% ethyl alcohol. Sulfuric acid, approximately IN solution. Sodium hydroxide, 1N standard solution. Dioxane, freshly distilled. PROCEDURE
Dissolve a 25-gram sample of sodium trichloroacetate in water and dilute to 100.0 ml. ( I t is best to take such a relatively large quantity t,o ensure a representative sample.) After dissolving, select an aliquot which will cause the final tit,ration volume to be about half of the blank titration volume. Pipet duplicate 10.00-ml. aliquots of sample solution into 250ml. reflux flasks, add 1 drop of methyl red, and neutralize by
