Drabkin, D. L., Waggoner, C. S., J . Biol. Chem. 89, 51 (1930). Eden, A., Green, H.. H.,. Biochem. J. 34, 1202 (1940).
Elvehjem, C. A,, Lindow, )I7. C., J .
Biol. Chem. 81,435 (1929). Forster, W. A., Analyst 78, 614 (1953).
Gshler, A. R., ANAL. CHEW 26. 577 1954).
'
)hartit, 13. T., Sommer, H. H., ISD.EYG.CHEV., ASAL. ED. 3, 24 (1931). Gilliq, J., IIoste, J., FernandezCaldiis, E , Anales edafol. y jisiol. zcgctal 9, 585 (1950). Gran. Q.. Anal. Chiin. Acta 14. 150 (1956,. Greenleaf, C. A,, Ibid , 2 8 , 269 (1945). Greenleaf, C. A., J . Assoc. O$c. Agr. Chemisfs 30, 144 (1947). Guest, It. J., AYAL. CHEU.2 5 , 1484 (1953). Hoste, J., Eeclrhout, J., Gillis, J., Anal. Chzni. Acta 9, 263 (1953). Hubbard, D bl., Spettel, E. C.,Zbid., 2 5 , 1245 (1953).
Hurd, L. C., Chambers, J. S., ISD. ENG. CHEll., ASAL. ED. 4, 236 (1932).
(31) Janardhanm, P. B., J . Sci. Ind. Research ( I n d i a ) 12B, 514 (1953); C.A . 48. 5718. (32) Kingsbury, R. M., Lake, C. L., T a p p i 35,527 (1952). (33) Lasausse, E., Frocrnin, L., J . Pharm. Chinz. ( 8 ) ,2 3 , i i (1936). (34) Loriente, E . , h'afure 159, 470 (1947). (35) hIcCu11v. C. R.. Pulv and Pawer Znd. 22, 4$ (1948,: (36) Maltby, J. G., aYature 160, 468 (1947). (37) hlorrison, S. I,,, Paige, H. L., 1x0. EIG. CIIEJI.,ASAL. ED. 18, 211 (1945). (38) Parker, W.E., Griffin, F. P., Can. J . Research 17B, 66 (1939). (39) Padais, R., Bull. soc. chim. France 1947, 1031. (40) Reed, J. F., Cummings, R. W., ISD. EXG.CHEII.. ASAL. ED. 13. 124 (1941). (41) Richardson, F. W,, Analyst 55, 323 (10501. ,- (42) Sandell, E. B., "Colorimetric Deter-
Sedivec, V., Vasak, V., Collection Czechoslou. Chem. Communs. 15, 260 (1950).
Sheets, O., Pearson, R. W., Gieger, ll.,IXD. ESG.CHEW,ASAL. ED. 7, 109 (1935). Shumilov, P. V., Kil'ter, il. Y., Zhur. Przklad. Khim. (1954).
27,
109
Smith, G. F., McCurdy, W. H., A N ~ CHEX L 24,371 (1952). Smith, G. F., Kilkins, D. H., Zbid , 2 5 , 510 (1953).
Supplee, G . C., Bellis, B., J . Dairy Sci. 5 , 455 (1922).
Sylvester, N . D., Lampitt, L. H
,
Analyst 60, 376 (1935).
Wetlesen, C. U., Gran, G., SLensk. Papperstidn. 5 5 , 212 (1952).
Williams, IT.,J . Dairy Research 3 , 93 (1931).
I
mination of Traces of Metals," 2nd ed., p . 300, Interscience, Xew York, 1950. (43) Schonheimer, R., Oshima, F., 2. physiol. Chent. 180, 249 (1929).
RECEIVEDfor review hIay 12, 1956. Accepted Sovember 7 , 1956. Division of Analytical Chemistry, 129th Meeting, ACS, Dallas, Tes., April 1956.
Reactions of Arsenic(ll1) and Arsenic(V) with Thioacetamide in Acid Solutions ELIOT A. BUTLER and ERNEST H. SWIFT Gates and Crellin laboratories o f Chemistry, Californiu Institute of Technology, Pasadena, Calif.
The rates of precipitation of arsenic(ll1) sulfide by thioacetamide have been measured in solutions having pH values from 3.78 to 1 and found to follow quantitatively the calculated rates of hydrolysis of the thioacetamide to hydrogen sulfide. Qualitative experiments with solutions of pH 4 to 6 have produced no evidence of a change of mechanism to a direct reaction, such as was observed in the case of lead. Studies of the reaction of arsenic(V) with thioacetamide have shown that initially arsenic(V) is reduced to arsenic(ll1) with formation of sulfur. The rate of reduction i s first order with respect to the concentration of arsenic(V), thioacetamide, and hydrogen ion. The hydrogen sulfide resulting from hydrolysis of the thioacetamide reacts initially with arsenic(V) to form thioarsenic acids, and then with arsenic(ll1) to produce arsenious sulfide.
T
results of studies of the acidcatalyzed hydrolysis of thioacetamide and of the precipitation of lead sulfide by thioacetamide in acid solutions were reported in a previous HE
paper (6). The precipitation of lead sulfide by thioacetamide proceeded through hydrolysis of t h e latter in solutions with hydrogen ion concentrations greater than about l O + M . However, in solutions having pH values from 5.1 t o 3.4, the precipitation proceeded at a rate appreciably greater than and independent of the rate of hydrolysis. This paper reports studies of the reactions of arsenic(II1) and (V) with thioacetamide in acid solutions. REACTIONS
OF
ARSENIC(I1I) ACETAMIDE
WITH
THIO-
Reagents. Solutions of thioacetamide, sodium thiosulfate, potassium iodate, a n d hydrochloric acid mere prepared as described by Swift and Butler (6).
Considerable variation in the quality of commercial thioacetamide has been noted, even in separate samples from a single source, and anomalous results in reaction rates and precipitation effects have been obtained from some specimens. Material used consisted of white crystals, which could be dissolved completely in water to give a 1VF (volume formal) solution (or left only a residue amounting t o not more than 1 mg.
from 100 ml. of such solution), and had a melting Doint range - of 111.0" to 113.2"C.A stock solution of arsenious acid. O.lOVF, was prepared by dissolving a weighed portion of arsenious oxide in hot water; the concentration was checked iodometrically. Sodium formate-formic acid buffer solutions were prepared from sodium hydroxide and 90% formic acid and the p H values mere checked Kith a Beckman hlodel G p H meter. Apparatus. T h e reaction apparatus was that used in t h e study of t h e hydrolysis of thioacetamide (6). Rate Measurements. T h e reaction solutions were prepared by mixing accurately measured volumes of t h e stock solutions of thioacetamide, arsenic(II1) , and hydrochloric acid or sodium formate-formic acid buffer a n d diluting t o 100 ml. The reaction solution was heated to 90" + 1" C. and maintained at t h a t temperature by a constant temperature bath. Samples of the reaction solution were forced by air pressure at timed intervals through the sintered-glass tube. Any excess of reaction solution which remained in the sintered-glass tube after the removal of a sample was forced back into the main body of the solution by a low pressure flow of nitrogen. XitroVOt. 29, NO. 3, MARCH 1957
419
Table I.
Comparison of Measured and Calculated Concentrations of Arsenic(1ll) at Timed Intervals
(Calculations are based upon the second-order hydrolysis constant, k = 0.21 liter mole-' minute-'. All solutions were O.1OVF in thioacetamide; T = 90" C.) As(II1) Time, Hours aH + Concn. 0 1 1.5 2 1 . 6 x 10-4 3leasd. 0.0099 0.0097 0,0095 .. 0.0096 Calcd. *.. 0.0098 9 . 3 x 10-4 hleasd. 0.0099 0.0089 0.0086 0.0084 Calcd. ... 0,0091 0.0087 0.0083 .
ctHr
a
9
x
8
x 10-2
IO-*
As(II1)
Concn.
I
Time, Minutes
Neasd. Calcd. Neasd. Calcd.
0
3
5
6
9
10
15
20
0,0099 ... 0,0099
...
0.0089 0.0093 ...
...
...
0.0083
0.0086
0.0081 0.0080
0,0076
0
...
...
0.0074 ...
...
...
...
...
0.0062
0.0065
...
I
.
...
.
0.0045 0.0032
Oa
...
Calculated to be complet,ely precipit.ated in 8.8 minutes.
gen was not bubbled through the reaction solution during the period of the run. The samples were cooled immediately in a n ice bath to quench the reaction; then about 0.01 gram of finely divided dry paper pulp 17-as added and the mixture was centrifuged. This procedure n-as necessary because some arsenic(II1) sulfide from the reaction tube passed through the sintered glass and couId be removed by centrifugation only if paper pulp was present. A 10.00-ml. portion of #le centrifugate was withdrawn by pipet and to this mas added 1 ml. of GVF hydrochloric acid. This acidified qolution JTas heated in boiling water for 5 minutes. The resulting Dreciuitate of arsenic(111) sulfide n-as'rem&ed by centrifugation and the arsenic was estimated by the following procedure ( 7 ) . The precipitate was dissolved in 1 ml. of 6 V F .odium hydroxide and bromine water was added until an excess was indicated by a slight yelloxv color. The solution was heated in boiling water for 2 minutes, after nliich 3 ml. of water were added, follon ed by 6VF hydrochloric acid until a permanent bromine color appeared. Sodium hydroxide solution was added dropwise until the bromine color was discharged. Two milliliters of 6 V F formic arid were added and the solution \vas heated in a water bath for 4 minutes. The solution was cooled, m e p t with carbon dioxide for 3 minutes, and washed into a 125-ml. flask which contained about 0.5 gram of potassium iodide in 0.5 ml. of n-ater. Hydrochloric acid, 12T'F, was swept with carbon dioxide and a volume equal to t h a t of the arsenic-potassium iodide solution was added. The resulting solution was cooled and titrated with standard sodium thiosulfate solution to the disappearance of the iodine color. Confirmatory determinations demonstrated t h a t this procedure gave results accurate t o within 1% with the quantities involved. I n the precipitation experimentsmade with solutions 6 V F in hydrochloric acid, 10-ml. portions of the reaction solution were prepared from 1 2 V F hydrochloric acid, the arsenious acid, and
10.0 10
0 9.5 Y 0
t
c
9.0
8
b E
r:8.5
-
.3
8.0
I
420
ANALYTICAL CHEMISTRY
0.5
1.0 1.5 TI ME (hours)
2.0
5
10
15
20
T I ME (min.)
Figure 1. Calculated and measured concentrations of arsenic(ll1) various hydrogen ion activities
vs. time a t
Solid lines are calculated values. Thioacetamide, O.lOVF, 90' C. Hydrogen ion activities 1. 1.6 x 10-4 2. 9.3 x 10-4 3. 9 x 10-3 4.
8
x
10+
the thioacetamide solutions. These solutions were placed in tightly stoppered tubes which Kere suspended in a constant temperature bath of the desired temperature. The tubes mere removed from the bath after timed periods, cooled in a n ice bath, and centrifuged. A large excess of thioacetamide was added to the centrifugate, which was then heated to 90" C. in order to test for completeness of precipitation of the arsenic.
Data and Discussion. Precipitation of Arsenic(ll1) from Dilute Acid Solutions. The results of a series of precipitation experiments made in solutions with p H values ranging from approximately 1 t o 4 are shown in Table I. There the analytically found arsenic(II1) concentrations a t various times are compared x i t h arsenic(II1)
concentrations calculated upon the assumption that the hydrolysis of thioacetamide is unaffected by the arsenic(111) in solution, and that the hydrogen sulfide from the hydrolysis is removed quantitatively by precipitation of arsenic(II1) as sulfide. The agreement between the analytical and calculated concentrations is within the limits of experimental error. I n Figure 1 are plotted the data of Table I ; there is no significant trend in measured concentrations. It appears that precipitation of arsenic(II1) sulfide by thjoacetamide at p H 1 to 4 proceeds through the hydrolysis of thioacetamide and there is no evidence of a direct reaction such as was found with lead(I1). I n preliminary experiments a stream of one to two bubbles of nitrogen a
second m s passed through the reaction solution during the reaction period. K h e n this was done, the arsenic was precipitated more slonly than was calculated from the rate of hydrolysis. I n precipitation of lead sulfide by hydrolysis oi thioacetamide, such s l o ~ v sneeping had a negligible effect upon the rate of formation of lead sulfide. This gives qualitative indication that the more acidic arsenic(II1) is less rapidly precipitated by the hydiogen sulfide than is lead(I1). I n solutions n i t h pH values greater than about 4, the solubility of arsenic(111) sulfide increases rapidly with pH (4). The acid concentration a t which the solubility commences to increase is a function of the total sulfide (H,S HS- S--) concentration, and varies from about l o F 4for solutions saturated with hydrogen sulfide to 5 x 10-6 for solutions containing no added sulfide. Quantitative rate measurements were not made in solutions having pH values above 4 because of the uncertainties resulting from the increasing solubility of the arsenic(II1) sulfide and the very loiv rate of hydrolysis of thioacetamide. Hon ever. qualitative comparisons of the rates of precipitation of arsenic by thioacetamide in buffered solutions shon-ed that rate of formation of precipitate continued to decrease as the p H n-as increased from 4 to 6. Thus, unlike lead(II), arsenic(II1) does not show a change of mechanism in this pH range in its reactions with thioacetamide to form the sulfide. Qualitative evpeiiments have shown that other basic cations suchas silver(I), cadmium(II), nnd copper(II), display e+ dence of a direct ieaction with thioacetamide in the same pH range as does lead. However, antiniony(II1) is siinilar in its behavior to arsenic(II1). Further work is being done in an effort to establish a general basis for this tliffeience in behavior.
+
+
Precipitation of Arsenic(ll1) Sulfide Frefrom 6VF Hydrochloric Acid.
quently arsenic is separated from other elements by precipitation with hydrogen sulfide from approximately 61'P hydrochloric acid solutions. Qualitative experiments n ere made t o determine if the precipitation by thioacetamide is hydrolysis-controlled under such conditions. The activity coefficient of 6VF hydrochloric acid is calculated from the data of Randall and Young (3) to be 4.3. K i t h this information and the previously determined effect of temperature it is possible to obtain by extrapolation the rate of hydrolysis of thioacetamide a t this acid concentration and a t various temperatures. Thus, if a solution is initially 0.03V-F in thioacetamide, 6T7P in hydrochloric acid, and 0.01VF in arsenic(III), and if the hydrogen sulfide reacts quanti-
tatively n-ith the arsenic as fast as it is formed, one can calculate by assuming an average thioacetamide concentration that a t 40" C. approximately 10 minutes n-ould be required for the production by hydrolysis of a quantity of hydrogen sulfide equivalent t o the arsenic. Experimentally, n.hen several identical solutions of the above concentrations nere maintained a t 40" C. for various periods of time. the precipitation n-as incomplete after 8 minutes but complete after 10 minutes. Similar experiments neie conducted a t 20" C. with solutions 0.1 T% in thioacetamide, 61'8' in hydrochloric acid, and O.01T'F in arsenic(II1). The calculated time was 16 to 17 minutes, and only a trace of arsenic remained in solution after 25 minutes. Thus, even with the approximations involved, the calculated time for complete precipitation was within 50% of the measured time. These experiments. besides indicating that the piecipitation of arsenic(II1) sulfide by thioacetamide is hydrolysiscontrolled even in 61% hydrochloric acid, show that the hydrolysis of thioacetamide can be caused to proceed a t room temperature a t a significant rate by the use of suitable acid concentrations. Statements in the literature ha\re implied that elevated temperatures are essential for the hydrolysis, and the effect of pH upon the hydrolysis rate has been overlooked. Effect of Chloride Ion. Arcand (1) has shown that in 0.1VF hydrochloric acid the ratio of arsenious dihydroxychloride to arsenious acid is 9 X This ratio increases to about 33 in 6 T T hydrochloric acid. Thus, in the quantitative rate measurements reported above, chloride-containing species of arsenic(II1) n ere not significant, but in the qualitative experiments in 6VF hydrochloric acid there was a change of predominant species to the dihydroxychloride. There was no apparent change in the rate of precipitation of the sulfide a t the higher acid concentration; this indicates that both of the above mentioned arsenic(II1) species react with hydrogen sulfide rapidly as compared with the rate with which hydrogen sulfide is produced by hydrolysis. Analytical Applications. These results make possible the calculation of the time required for quantitative precipitation of arsenic(II1) as the sulfide by thioacetamide under various conditions. For example, with 100 ml. of a solution 0.10VF in thioacetamide, and containing 300 mg. of arsenic(III), the approximate times required for complete precipitation a t various hydrogen ion activities and temperatures are given in Table 11. These calculated values show the importance of control of the temperature
Table 11. Calculated Time Required for Quantitative Precipitation of 300 Mg. of Arsenic(l1l) as Sulfide by Thioacetamide (Volume, 100 ml. ; thioacetamide, 0.10T7F Time, Minutes aH+ 90" C. 7 0 ° C . GO" C. 1
and acid concentration if arsenic(II1) is to be quantitatively precipitated. Such calculations are invalid if the solution is boiled or an inert gas passed through i t ; under such conditions the hydrogen sulfide can be expelled from the solution so rapidly that no or only partial precipitation is obtained. Thioacetamide has been extensively used as a substitute for hydrogen sulfide gas for the precipitation of the conventional hydrogen sulfide group elements from hot solutions approximately 0.3F in hydrochloric acid. The time and the thioacetamide required to obtain quantitative precipitation can be minimized by beginning precipitation in a small volume of solution and then diluting to the desired final volume and acid concentration. During the pretreatment the concentrations of thioacetamide and acid are correspondingly greater and advantage is taken of the fact that the rate of hydrolysis of the thioacetamide to give hydrogen sulfide is first order with lespect to the concentrations of both the thioacetamide and the acid. K h e n the solution is diluted to the final volume, thus decreasing the acid concentration, the dissolved hydrogen sulfide precipitates the more soluble sulfides. such as cadmium and lead. The solution should not be boiled during this pretreatment or hydrogen sulfide will be lost. The solution should be warmed, the container stoppered. and then heated in a bath of boiling water for the necessary time. REACTIONS
OF
ARSENIC(V1 ACETAMIDE
WITH
THIO-
Reagents. Solutions of arsenic(V) were prepared from reagent grade arsenic pentoxide and standardized by the following procedure:
A measured portion of the solution was swept for 3 minutes with a stream of carbon dioxide and then xvashed into a 125ml. flask which contained 0.5 gram of potassium iodide dissolved in 1 to 2 ml. of water. The solution was cooled as an equal volume of 12VF hydrochloric acid previously swept with carbon dioside was added. The resulting solution was titrated with standard sodium thiosulfate to the disappearance of the ioVOL. 29, NO. 3, MARCH 1957
* 421
dine color. The average deviation of four determinations by this procedure was 0.7 part per hundred. Experiments showed that the arsenic(V) solutions contained less than 0.170 of arsenic(111). Magnesium ammonium nitrate reagent, 0.5VF in magnesium nitrate, 3VF in ammonium nitrate, and 0.2VF in ammonium hydroxide, was prepared from reagent grade chemicals. Apparatus. The same apparatus was used as in the experiments with arsenic(II1). Procedure. The rate of reduction of arsenic(V) by thioacetamide was measured in the following way: The reaction solution was prepared from the standard solutions of arsenic(V), hydrochloric acid, thioacetamide, sodium chloride, and water. The sodium chloride was added to maintain the chloride concentration constant a t O.lVF, except in certain qualitative experiments made in more concentrated hydrochloric acid. The reaction solution was placed in a heating bath and, when the desired temperature was reached, samples of the solution were taken a t timed intervals. The samples were cooled rapidly in an ice bath and centrifuged to remove the suspended sulfur which passed through the sintered-glass plug in the sampling tube. Two 5-ml. portions of the sample were taken by pipet and analyzed separately for arsenic(V). To each portion were added 2 ml. of 3VF ammonium chloride, 1 ml. of 6 VF ammonium hydroxide, and 2 ml. of the magnesium ammonium nitrate reagent. The solution was stirred vigorously and then allowed to stand for a t least 2 hours. The tendency for magnesium ammonium arsenate to form supersaturated solutions made this latter treatment necessary. The precipitate was removed by centrifugation and washed twice with 2-ml. portions of the magnesium ammonium nitrate reagent. The centrifugate was removed in each case by a drawn glass capillary attached to an aspirator. The precipitate was dissolved in 5 ml. of 0.6VF hydrochloric acid, the resulting solution was swept for 3 minutes with carbon dioxide, and the arsenic(V) was determined iodometrically as described above. Confirmatory experiments demonstrated that a t least 99% of the arsenic(V) present was precipitated as magnesium ammonium arsenate by this procedure in the absence of thioacetamide; the possible effect of thioacetamide is discussed below. Data and Discussion. Qualitative experiments showed t h a t in solutions 0.3 to 1VF in hydrochloric acid, arsenic was precipitated appreciably faster by thioacetamide than by saturation with hydrogen sulfide under the same conditions. A white precipitate of sulfur was observed first; the quantity of white sulfur formed before appearance of any yellow precigi422
ANALYTICAL CHEMISTRY
Table 111.
Effect of Hydrochloric Acid Concentration upon Rate of Reduction of Arsenic(V) b y Thioacetamide (TAA)
(Calculated constants for the expression - d [As(V)I = k [TArl] [As(V)][H+].) Initial condt centrations. 0.100VF CHICSNH,, O.01OVF As(V). 0.1OVF total chloride. T = 90' C. Time. ~
0.0100 0.020 0.040
0.080 0.100
a
0.0037 0.0047
0,010 0.010
6 4.5 3
0.0040
0.010 0.010
0.0031 0.0036
0,010
50 38
20 10
38
33
34
. 4 ~ . 36 ~
Does not include li for 0.010F HC1.
1
I
I
I
I
1
I
I
IO
I 15
I
5
20
25
30
TIME
I
(mh)
Figure 2. Effect of acid concentration on rate of reduction of arsenic(V)
Thioacetamide, 0.100VF; total chloride O.1OVF; 90" C. Hydrochloric acid concentrations 1. 0.OlOVF 2. 0.020VF 3: 0:040VF 4. 0.080VF 5. 0.1OOVF
tate indicated that the reduction of arsenic(V) was proceeding to a considerable extent before precipitation of a sulfide. On the basis of these observations, quantitative measurements were made to determine the rate and order of the reaction involved in the reduction of arsenic(V) by thioacetamide. Effect of Hydrogen Ion Concentration.
Rate measurements were made on solutions initially 0.01OOVF in arsenic(V), 0.lOOVF in thioacetamide, O.1OVF in chloride ion, and 0.0100 to 0.1OOVF in hydrochloric acid. I n Figure 2
the arsenic(V) concentration is plotted logarithmically against the time for runs a t various hydrochloric acid concentrations. The linearity of the plots indicates that the reduction reaction is first order with respect t o the concentration of arsenic(V), and the rate constants shown in Table I11 indicate a first-order hydrogen ion dependence. The range of acid concentrations considered was extended by qualitative experiments with solutions from 0.1 to 4VF in hydrochloric acid. The time required for the first appearance of a yellow sulfide in O.1VF hydrochloric
IO
d 5
I
I
I
10
15
20
tion; the extrapolated and observed values for several acid concentrations are shown in Table IV. The agreement indicates a first-order dependence upon the hydrogen ion activity through this range. Qualitative experiments in the range p H 3.4 to 4.5 indicated that the dependence of the rate upon the hydrogen ion concentration was less than first order. Subsequent quantitative measurements yielded the rate constants shown in Table V when first-order and half-order hydrogen ion dependence were assumed. A distinct trend is apparent even with t h e half-order assumption. A change in rate is to be expected as the hydrogen ion concentration is decreased below about lO-ZJP, since there is a change in the predominant arsenic(V) species. The first acid constant of arsenic acid is given as 5.6 x 10-3 ( 5 ); therefore the calculated ratio of arsenic acid to dihydrogen arsenate decreases from about 2 at pH 2 to 0.07 a t p H 3.4. KOfurther experiments were made in this low range of hydrogen ion concentrations, as analytically such precipitations are of little importance. At pH 2 the third-order rate constant is already above its value a t higher acid concentrations (Table 11). This is in agreement with the above discussion of the species involved, as a t p H 2 about 33% of the arsenic('\;) exists as dihydrogen arsenate.
T I ME (min .) Figure 3. Effect of thioacetamide concentration on rate of reduction of arsenic(V)
Hydrochloric acid, 0.02VF; sodium chloride, 0.08VF; 90" C. Thioacetamide concentrations 1. 0.025VF 2. 0.050VF 3. O.IOOT/'F 4. 0.2OOVF
Table 1V. Effect of Acid Concentration upon Time Required for Initial Sulfide Precipitation at 90' C.
[Initial concentrations. 0.1VF thioacetamide, O.01VF arsenic(V)] Calcd. Observed HCI, Time, Time, hloles/Liter Sec. Sec. 0.5 36 45 1.0 18 20-30 2.0 9 10-12 4.0
3
acid a t 90" C. was approximately 180 seconds when the initial concentrations of thioacetamide and arsenic(V) were O.1VF and 0.01VF, respectively. If it is assumed that dependence upon acid concentration is first order, and that quantitative reduction precedes sulfide precipitation, estimated times for the first formation of the sulfide can be obtained by extrapola-
5
Table VI. Table V.
Effect of Thioacetamide Concentration.
I n another series of experiments the initial concentrations of arsenic(V), acid, and chloride ion were kept constant while various thioacetamide concentrations were used. The results of these experiments, plotted in Figure 3, again demonstrate the first-order dependence of the reaction upon arsenic(]') concentration. Calculations made from the plots, presented in Table 1'1, indicate a first-order dependence of the reduction reaction upon the concentration of thioacetamide. Mechanism of Reduction. The apparent first-order dependence of the rate of reduction upon the arsenic(v)
Calculated Rate Constants
For expressions - d [ WVj1
=
k,[H+]
dt [TAA] [As(V)] and - d[As(V)I = dt kb[H+],1'21TL\.4] [As(V)] in solutions having pH values from 3.1 to 4.5 PH x.6 kb 3.4 146 2 9 4.0 358 3.6 4.5 932 5 0 ~
Effect of Thioacetamide Concentration upon Rate of Reduction of Arsenic(V)
Calculated constants for expression
- d[As(V)I ____
=
L [TAA][As(V)][H+]. Initial con-
dt
centrations. 0.02VF HCl, 0.OlOVF As(V), 0.08VF NaCI, T [CHaCSYH2], Mole/Liter
Pvlole/Liter
0.025
0.050
0.010 0 010
0,200
0.0074
0.100
[As(V)]i,
0.010
[As(V))*.
hlole/Liter 0.0062 0.0053 0.0047 0,0056
= 90" C.
k, Time Interval, Liter2/RZole2 Min. &fin. 30 3-1 20 32
10 2
Av.
VOL. 2 9 ,
NO. 3, MARCH 1957
38 35 35
423
concentration indicates that the reduction of arsenic(V) by thioacetamide is independent of the rate of hydrolysis of thioacetamide. RIoreover, no yellow sulfide formed until the measured arsenic(V) concentration had dropped t o less than 10% of its original value. These observations led to evperiments designed to gire information regarding the mechanism of the reduction of arsenic(Y) and its subsequent precipitation. Rate of arsenic(II1) sulfide precipitation corresponds to the rate of hydrolysis of the thioacetamide. Xoreover, even rery slow sneeping of the solution with a n inert gas caused a decrease in the rate of precipitation of the sulfide. Thus, it should be possible to prepare an acid solution n-hich contains arsenic(II1) and thioacetamide, heat this to a specified reaction temperature, sv-eep it vigorously with an inert gas to remove hydrogen sulfide, and thus prevent precipitation of the sulfide. Such an experiment mas performed v i t h a solution O.01VF in hydrochloric acid, O.1OVF in thioacetamide, and O.OlVF in arsenic(II1) at 90' C. S o precipitate was obtained in 10 minutes when the solution was swept v-ith a rapid stream of nitrogen, but immediate precipitation occurred when the flon- of gas was stopped. On the other hand, vigorous sweeping of the reaction solution with nitrogen should haye no effect upon the rate of reduction of arsenic(V) by thioacetamide, if the reduction reaction depends not upon hydrolysis of the latter, but upon a direct reaction. Experiments showed that in solutions 0.02VF in hydrochloric acid, 0.lOVF in thioacetamide, and 0.011'F in ar$enic(V) the reduction of arsenic(V) proceeds during the first half of the reaction a t a rate n hich is independent of whether or not the solution is swept with nitrogen. Thus a direct reaction is indicated. Effect of Thioarsenic Acids. I n the rate determinations after about half of the arsenic(V) initially present had been reduced, the arsenic(V) concentration decreased somervhat more rapidly than would be estimated byextrapolation of the straight-line portion of the semilogarithmic plot. It is recognized that arsenic(T7) is reduced to some extent by hydrogen sulfide in acid solution (2) and that thioarsenic acids such as HIAs02Sare formed prior to the reduction; therefore the possibility t h a t hydrogen sulfide from hydrolysis of the thioacetamide reacts with arsenic(V) was considered. If thioarsenic acids are formed, a significant difference should be observed in the rate at which hydrogen sulfide can be swept from a n acid solution of thioacetamide containing arsenic(V) and t h a t from a solution free of arsenic. 424
e
ANALYTICAL CHEMISTRY
Table VII.
Solution Composition, V F Volume, ml. Temperature,
O
C.
Effect of Hydrogen Sulfide upon Arsenic Acid Solution A B c 0.010 h ( V ) 0 . 4 CdC12 0.10 T A P 0,010 HC1 6 NH40H 0,010 HC1 20.0 20.0 15 90 i 2 25 25
Time, Min. Experiment I H2S formed (calcd), 0.02 mmole Found bv analysis, - . mmole Experiment I1 HpSformed (calcd.), 0.12 mmole Found by analysis, mmole Thioacetamide.
5
As(V) 0.19
Sulfide 0.01
-4s(V) 0 19
Sulfide 0.08
30
Q
I n experiments \\.it11 solutions 0.02VF in hydrochloric acid and O.1OVF in thioacetamide, the rate a t which hydrogen sulfide could be removed from the solution was found to be decreased by approximately 90% if the solution was also 0.OlVF in arbenic(V). This observation, coupled with the apparent increase in the rate of reduction of arsenic(V) after the reaction TTas about half complete, indicates that hydrogen sulfide from the hydrolysis of thioacetamide reacts n i t h the arsenic(V) in solution to form complex thioarsenic acids. This reaction is sloiv relative to the direct reaction betn-een arsenic(V) and thioacetamide during the first half of the reduction, but becomes significant as the rate of the direct reaction decreases because of decreased arsenic(V) concentration and increased total hydrogen sulfide eoncentration. It appeared of interest to determine TT hether the thioarsenic acids, once formed, would remain as such and whether during precipitation of magnesium ammonium arsenate the thioacid was reconverted to oxygen acid and sulfide. Experiments were conducted in which a slow stream of air was drawn consecutively through three solutions (Table V I ) . I n each experiment the total quantity of hydrogen sulfide formed by hydrolysis was calculated, and the quantity collected in solution C was determined iodometrically (6). Thus the quantity of hydrogen sulfide retained by solution B v-as obtained by difference. Immediately after the reaction was stopped, portions of solution B were analyzed by tlie magnesium ammonium arsenate procedure described above. I n neither experiment was there a perceptible precipitate in solution B a t the conclusion of the reaction period, but, in both cases there mere precipitates of sulfur in remaining portions of solution B after approximately 1 hour. Subsequent boiling of the solutions did not cause the formation of additional precipitate. I n Experiment I all of the sulfide produced is accounted for, 0.01 mmole
having passed into solution C and 0.01 mmole retained in B. I n Experiment I1 0.03 mmole of the sulfide produced is not accounted for in the analyses. -4 possible cause of this discrepancy is the formation of thioarsenic acids containing more than one atom of sulfur. These and previous experiments show that the rate of formation of the thioarsenic acid from hydrogen sulfide is relatively fast compared to the subsequent decomposition of these acids. Pentapositive arsenic, present as a thio compound a t the start of a magnesium ammonium arsenate precipitation, may not be quantitatively precipitated. Effect
of
Chloride
Concentration.
I n all the quantitative rate experiments discussed, the chloride concentration was maintained a t O.lVF, in order that variations due to chloride complexes of arsenic(V) would not cause uncertainties in the results. As a check on the magnitude of the chloride effect, a rate determination was made in Ehloride-free perchloric acid solution, initially 0.10VF in thioacetamide, 0.010VF in arsenic(V), and 0.020VF in perchloric acid. The rate constant obtained from this experiment was 35 liter2 mole-2 min.-l, rrhich is in good agreement n-ith the values obtained from experiments made in chloride SOlutions. This indicates that a t the acid concentration involved, chloride does not have a significant effect upon the rate of reduction of arsenic(\-) by thioacetamide. Effect of Temperature. -4single rate measurement was made at 70" C. to determine whether the reduction of arsenic(V) by thioacetamide shows a normal temperature dependence. The third-order rate constant obtained a t 70" C. is 6.3 which indicates an energy of activation of about 20 kcal. per mole. This is near the values obtained earlier in this study for other thioacetamide reactions. Analytical Applications. I n comparing the precipitation of arsenic(\-) as sulfide by hydrogen sulfide and by thio-
acetamide, studies of the former method ( 5 ) have shown t h a t if 50 ml. of solution, 0.6VF in hydrochloric acid, which contains 500 mg. of arsenic(]-) is heated almost to boiling, saturated with hydrogen sulfide, diluted to 100 ml. and cooled, resaturated with hydrogen sulfide, and heated to 100’ C. in a pressure bottle, quantitative precipitation of the arsenic nil1 require a t least 30 minutes in the pressure bottle. By contrast it is calculated from the present study t h a t in 100 ml. of solution, 0.3T-F in hydrochloric acid and 0.5VF in thioacetamide, 500 mg. of arsenic(\-) should be 99.97, reduced a t 90’ C. in approximately 1 minute and that sufficient hydrogen sulfide should be produced to precipitate the arsenic completely in about 6 minutes. (This allows for the loss by volatilization of 10 to 20y0 of the hydrogen sulfide from the relatively concentrated thioacetamide solution.) I n experiments with a similar solution the arsenic iTas completely precipitated after 6 minutes under the above conditions. The time would bc reduced to about 3 minutes
if a pretreatment in 0.6VF hydrochloric acid were used. Thioacetamide has been used a t the California Institute of Technology in a recently developed system of elemental analysis as the precipitant for a sulfide group which included pentapositive arsenic. Much time is saved by the change from hydrogen sulfide, and the method has been consistently satisfactory in student use. This investigation has emphasized again that reactions of thioacetamide and inorganic ions do not consist only of hydrolysis to give hydrogen sulfide, which then reacts Fvith the inorganic ions. Studies of the reactions of thioacetamide n-ith other inorganic ions are being continued. ACKNOWLEDGMENT
This investigation has been supported in its latter stage by a grant from the National Science Foundation. The authors are indebted to Gerald Klaz for preliminary measurements on
the reduction of arsenic(V) and the precipitation of arsenic(II1) by thioacetamide. LITERATURE CITED
Arcand, G. AI., thesis, California Institute of Technology, 1955. McCay, L. W., A m . Chetn. J . 10, 459 (1888); J. A 7 1 ~Chem. SOC.24, 661 (1902). Randall, )I., Young, L. E., Zbid., 50, 989 (1928). Ringbom, A., “Solubilities of Sulfides. Preliminary Report to Commission on Physico-Chemical Data of -4nalytical Interest, Analytical Section, IUPAC,” July 1953. Swift, E. H., “System of Chemical Analysis,” Prentice-Hall, S e w York 1938: (6) Swift, E. H., Butler, E. A , , i 2 s . k ~ . CHEJI. 28,146 (1956). ( 7 ) Swift, E. H., Siemann, C., “System for Ultimate Analysis of Chemical Warfare Agents,” C.W.S. Field Lab Memo 1-2-4 (rev. August 1944). RECEIVEDfor review July 13, 1956. Acce ted Tovember 20, 1956. Contribution s o . 2169 from the Gates and Crellin Laboratories of Chemistry, California Institute of Technology, Pasadena 4, Calif.
Potentiometric Determination of Mercaptans in Presence of Elemental Sulfur J. H.
KARCHMER
Humble Oil and Refining
Co.,Bayfown, Tex.
In the potentiometric titration of mercaptans in gasoline using alcoholic silver nitrate, low results and poorly defined titration curves may b e obtained with samples having a mercaptanelemental sulfur ratio greater than 1 to 1, and when the sample remains in the alkaline alcoholic titration solvent for an appreciable time before the titration. Experimental results indicate that the poor results can be attributed to the presence of inorganic polysulfides whose formation is related to the amount of elemental sulfur and the mercaptan types. Because the postulated reaction for the formation of the inorganic polysulfide is promoted by the alkalinity of the solvent, a less alkaline solvent was employed and more accurate results were obtained. The reaction of elemental sulfur with n-butyl, fert-butyl, and phenyl mercaptans has been studied and a probable mechanism for the steps in this reaction has been postulated. Some evidence is presented for the existence of the (RSS)- ion, which is produced as an intermediate in the reaction,
T
HE
POTESTIONETRIC
METHOD
Of
Tamele and Ryland (11) for determining mercaptans (thiols) in gasoline is accurate, reliable, and unconiplicated in absence of elemental sulfur and hydrogen sulfide. Even in the presence of these materials, the determination can still be accurate for most samples, if the analyst is aware that certain precautions must be taken. This paper describes some of the complications encountered when elemental sulfur is present and recommends a slight modification, to minimize these difficulties. The original method consists of potentiometrically titrating the sample dissolved in alcohol, containing 0.1N sodium acetate to buffer the solution, with a n alcoholic solution of silver nitrate using a silver sulfide indicator electrode us. a mercury-sodium acetate half cell. More recently a glass electrode has been employed as the reference electrode (8)in place of the mercurysodium acetate half cell. External calomel cells have also been successfully used. These cells are electrically connected to the solution by an agar-
saturated potassium nitrate bridge t o avoid contaminating the solution with chloride ions from the calomel cell. I n the presence of hydrogen sulfide and/or elemental sulfur erroneous results may be obtained unless the titration curve is correctly interpreted. When hydrogen sulfide is present, the initial potential between the silver electrode and the solution is approximately -0.7 volt (see Figure 1). As the silver ion is added to the solution, silver sulfide is precipitated and after all of the sulfide ion has reacted, the potential breaks sharply to the voltage characteristic of the particular mercaptan present. This voltage is influenced largely by the solubility product of the silver mercaptide in the solvent. For n-butyl mercaptan it is about -0.35 volt. Upon continued addition of silver ion to the solution, silver mercaptide begins to precipitate and finally another sharp “break” in voltage is observed. which corresponds to the end point of the mercaptan titration. The presence of hydrogen sulfide presents no difficulty, as the volume of silver nitrate used to reach the first VOL. 29, NO. 3, MARCH 1957
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