Chow, T. J., Thompson, T. G., ANAL.CHEM. 27, 18 (1955). Claffy, E. W., Am. J . Scz. 247, 187 (1949). Davis, E. N., Van Kordstrand, R. -4., ANAL.CHEM. 26, 973 (1954). Diamond, J. J., Ibid., 27, 913 (1955). Fast. E.. Nielsen. J. R.. J. ODt. Soc. A&. 37, 614(1947).' Feldman, c., h . 4 ~ .C m x . 21, 1041 . (1949). Geological Survey Bull. 980, U. S. Government Printing Office, Washington 25, D.C., 1950-51. Herzorr. L. F.. Pinson. R. H..
Table
VII.
Comparison of Barium and Strontium Determinations
Method Spectrochemical ( I ) Spectrochemical (direct current arc) (SI) Isotope dilution
...
(31'i
...
(laj
X-ray fluorescence (1'7) Spectrochemical (high voltage Eolution )
...
Strontium,yo G-1 TV- 1 0.0280 0,0290 0,0287 0.0172
...
0.0262
...
0.0263
Barium, 70 G- 1 JT- 1 0.130 0.0270
... ...
...
0.108
0.0145
spark-
0 0178
o.oi%
o.oi??
0.0256
0.0180
0.0197
GeoFhim. et Cosmochim. Acta 8;
295 (1955). Hillebrand, W.F., Lundell, G. E. F., "Applied Inorganic Analysis," Wiley, Kew York, 1929. Hinsvark, 0. N., Wittwer, S. H., Sell, H. lf., A 4 S A L . CHEM. 2 5 , 320 (1953). Hitchcock, R. D., Starr, W. L., A p p l . Spectroscopy 8, KO. 1, 5 (1954). Hower. J.. Fancher. T. W.. Science 125.'498 (1957). ' Kallman, S , , ASAL. CHEx 20, 449 (1948). Kulp, J. L., Turekian, K., Boyd, D. W.,Bull. Geol. Sot. Am. 63, 701 (1952). Lewis, G. J., Jr., Goldberg, E. D., ASAL.CHEX 28, 1282 (1956). Lucchesi, C. .4.,Ibid., 29, 370 (1957).
Lucchesi, P. J., Lewin, S. Z., Vance, J. E., Ibid., 26, 521 (1964). Mitchell, R. L., Robertson, I. M., J . SOC. Chem. I n d . 5 5 , 2691' (1936). Moore, ' C. E., Contrib. Princeton Univ. Observatory No. 2 0 , 1-110 (1945). Odum, H. T., "Biogeochemistry of Strontium," Ph.D. thesis, Yale Universitv. " , 1950. Pagliassoti, J. P., Appl. Spectroscopy 9, S o . 4, 153 (1955). Shaw, 1). M.,Ibid., 9, 40. 1, 44 (1954). Smales, A. -4., Analyst 76, 348 (1951). Specfrographer's LVews Letter (Applied Research Laboratories) 2, S o . 9 (1948).
(30) Ta lor, 4.E., Paige, H. H., ANAL. 6HEM. 27, 282 (1955). (31) Turekian, K K., Gast, P. W., Kulp, J. L., Spectrochim. Acta 9, 40- (1957'i. \ - - -
(32) Vallee, Vallee',-B. B. L., Rlargoshes, M.,A s . 4 ~ . CHEM.28, 180 (1956). (33) Weaver, Weaver. J. R., Brattain, Brattain. R. R., R.. Ibid., 21, 1038 (1949). (34) Killard, H. H., Goodspeed, E. W., I N D . ENG.CHEM., ANAL. ED. 8, 414 (1936). (35) Kilska, S., Acta Chem. Scand. 5 , 1368 (1951). RECEIVEDfor review ,4u~ust 21. 1957. Accepted March 24, 195s. Pubhation S o . 139, Shell Development Co., Exploration and Production Research Division, Houston, Tex.
Hydrolysis Reactions of Thioacetamide in Aqueous Solutions ELIOT A. BUTLER,' DENNIS G. PETERS, and ERNEST H. SWIFT California Institute of Technology, Pasadena, Calif.
,Spectrophotometric measurements of the rate of disappearance of thioacetamide in acid solutions have confirmed previous chemical measurements; hydrolysis of the thio group is predominant and the rate of hydrolysis of the amide group is too slow to be measurable by the methods used. The second-order rate constant obtained spectrophotometrically for the sulfide hydrolysis is 0.21 liter mole-' min.-' at 90" C., the same as that calculated from measurements of the hydrogen sulfide evolved. In sodium hydroxide solutions amide group hydrolysis is predominant and the secondorder velocity constant is 9 f 1 liter C. The mole-' min.-' at 100.3' rate of thio group hydrolysis in sodium hydroxide solutions is approximately one fourth that of amide group hydrolysis. The thioacetate ion hydrolyzes to sulfide more slowly; the rate is first-order with respect to both thioacetate and hydroxyl ion, and the second-order rate is constant 0.01 9 f 0.0015 liter mole-1 min.-l at 90" C.
The energy of activation was calculated to be 19 kcal. per mole from 70" to
90" C.
A
STUDY of
the acid-catalyzed hydrolysis of thioacetamide (6) has shown that the rate of hydrolysis to acetamide and hydrogen sulfide is first-order with respect to the concentration of thioacetamide and of hydrogen ion and the second-order velocity constant has the value 0.21 =t0.023 liter mole-' m k - 1 at 90" C. The hydrolysis of thioacetamide in acid solutions t o give thioacetic acid and ammonium ion was negligibly slow, compared vvith the hydrolysis t o acetamide and hydrogen sulfide. However, as the conclusions of that study were based entirely on measurements of the rate of production of hydrogen sulfide and the change in the hydrogen ion concentration, there appeared justification for a n independent confirmatory method which would permit direct determination of the rate of disappearance of thioacetamide. Such measurements would afford more evidence as
t o the relative rates of the sulfide and amide hydrolyses. Rosenthal and Taylor (6) have followed the hydrolysis of thioacetamide spectrophotometrically by making use of the absorption peak of thioacetamide a t about 260 mp. This technique was used in the confirmatory study of the acid-catalyzed hydrolysis of thioacetamide reported belolv.
REAGENTS AND APPARATUS. Arapahoe thioacetamide, lot 1402, was used throughout this work for the preparation of stock solutions. This material had a melting point range 111.0-13.0" C. and dissolved to give clear, colorless s o h tions. Reagent grade chemicals were used for all solutions. A Beckman Model DU quartz spectrophotometer with hydrogen discharge lamp and 1-em. quartz cells was used. PROCEDURE. Reaction solutions containing thioacetamide and hydrochloric Present address, Department of Chemistry, Brigham Young University, Provo, Utah. VOL. 30, NO. 8; AUGUST 1958
1379
I
I
I
I
I
I
I
I
240
250
260
270
280
290
300
W A V E LENGTH ( m y )
Figure 1. Plot of per cent transmittance against wave length for aqueous thioacetamide solutions 1. 2.
1 O-sVF CHaCSNH2 1 O-4VF CHaCSNHn
acid were prepared as described by Swift and Butler (6). These solutions were maintained a t 90" C. for timed periods, then cooled rapidly in an ice bath. The absorbances were measured with the spectrophotometer.
3
1
ANALYTICAL CHEMISTRY
I5
Figure 2. Absorbance of thioacetamide solutions as a function of time Initial thioacetamide, 1 .OO X 10 -4VF Hydrochloric acid, 0.25VF. T, 90' C.
comDosition uroducts. It did not interfere with measurements in acid solutions, because the absorption curves for thioacetamide, as the latter hydrolyzed, maintained the characteristic shape for thioacetamide and showed no anomalies. Data and Discussion. ABSORPTION RATEO F HYDROLYSIS I N !LCID SOLUBY THIOACETAMIDE. A plot of per TIOKS. I n Figure 2 are shown data cent transmittance against wave from a typical hydrolysis rate determilength for 10-4 and lO-5VF (volume nation. The solution was initially 0.25 formal, formula weights per liter) thioVFin hydrochloric acid and 1.00 X IO+ acetamide in water is shown in Figure l. VF in thioacetamide. The second-orThe wave length of maximum absorpder velocity constant calculated from the tion is 261 to 263 mp. The calculated slope, with the use of 0.76 as the activity molar extinction coefficients are 11430 coefficient of hydrochloric acid ( d ) , is and 11400 liter per mole em. for 10-4 0.21 liter mole-' min.-' a t 90' C. and 1O-W' thioacetamide, respecWithin the limits of the experimental tively. Additional experiments indierrors, this is the same as the chemically cated that deviations from Beer's lairdetermined value reported previously were consistently less than 57*. This (6). The agreement b e h e e n the reaccuracy was satisfactory for the rate sults from the two methods gives strong determinations, as errors caused by support t o the mechanism previously finite periods of heating and cooling proposed (6)-that the evolved hydrointroduced errors of similar magnitude. gen sulfide results directly from hydrolyISTERFERING SUBSTANCES. A 10W3 sis of thioacetamide and that the rate of VF solution of reagent grade acetamide hydrolysis of thioacetamide to yield showed no appreciable absorption in the thioacetic acid and ammonium ion is too range 230 to 330 mp. slow t o be detected by these measureA solution approximately 10-4VF in ments. This slow rate of the amide both acetic acid and acetate ion did not hydrolysis of thioacetamide is of interabsorb significantly a t wave lengths est, as Crocker (2) found the hydrolysis above 250 m p , The absorption inof acetamide to be much more rapid, creased slightly n-ith decreasing wave thus showing the effect of substitution length below 250 mp. of sulfur for oxygen. Seither ammonium ion nor hydrochloric acid solutions absorbed in the HYDROLYSIS REACTIONS OF THIOACETAMIDE region of interest. IN ALKALINE SOLUTiONS An aqueous hydrogen sulfide solution absorbed below 245 mp but not a t longer Barber and Taylor (1) have stated wave lengths. that the hydroxide-catalyzed hydrolysis KO positive conclusions could be of thioacetamide proceeds more rapidly drawn from experiments with three than the acid-catalyzed hydrolysis; commercial samples of thioacetic acid. however, no confirmatory data were These samples were intensely colored found in the literature. and their aqueous solutions were colored This investigation of the hydroxideand absorbed appreciably throughout catalyzed hydrolysis of thioacetamide the ultraviolet region. This absorption was made t o extend into alkaline soluwas probably due to impurities and detions the studies of the reactions in1380
6 9 12 TIME, M I N U T E S
volved in the precipitation of metal sulfides by thioacetamide. Two sets of hydrolysis reactions could take place in aqueous sodium hydroxide solutions. I. A. CHqCSNHo OH- =
+
11. C. CH3CSNH2
+ ++
OH- = CH3COiYHt HSD. CHaCONHa OH- = CHaCOONH3
+
Preliminary measurements showed that the rate of production of sulfide was initially fast, but after a short time became relatively slom-. Subsequent experiments indicated that the rate of hydrolysis of thioacetamide to ammonia by Reaction I-.% was approximately four times faster than the rate of formation of sulfide by 11-C. The quantities of sulfide and ammonia produced indicated that essentially all the thioacetamide had hydrolyzed within a short time, and that the slow rate of formation of sulfide observed thereafter n-as due to hydrolysis of thioacetate ion, CH3CSO-, by Reaction I-B. Amide Hydrolysis of Thioacetamide. The rate of hydrolysis of thioacetamide to give thioacetate ion and ammonia must be known if data concerning the sulfide hydrolysis are to be interpreted correctly. REAGENTS.Solutions of hydrochloric acid and sodium hydroxide were prepared and standardized by conventional methods. An 0.80VF sodium perchlorate E O ~ U tion vas used for adjusting the ionic strength of solutions and was prepared as described by Srift and Butler (6). APPARATUS. The all-borosilicate glass apparatus consisted of a 200-ml. round-bottomed flask with a 30-cm. Liebig condenser in vertical reflux position with a connector to a 30-cm. Liebig condenser in condensing position.
t
ri
I?2
0.7
-
0.6
-
cn
0
r ~ )
-I 0
0
\
N 0.5 -
I 2 Ln
-
I 0.4
-
0
I
I
5 TIME, M I N U T E S
/
IO
Figure 3. Effect of initial thioacetamide concentration upon the rate of the amide hydrolysis 0.01 OVF NaOH.
100.3' C.
Initial thioacetamide concentrations:
0
0.001OVF 0 0.0020VF A 0.0040VF
An adapter led to a 25 X 200 mm. test tube. PROCEDURE FOR RATE &TEASUREhlENTS O F -4hlIDE HYDROLYSIS.A re-
action solution was prepared from the standard solutions of thioacetamide, sodium hydroxide, and sodium perchlorate. An excess of standard hydrochloric acid was used as a collecting solution for the ammonia. The reaction solution was rapidly heated to boiling, with both of the condensers in operation. The time was measured from the commencement of boiling. The solution was allowed to reflux for a period of time determined by the extent of the reaction desired, after which the reflux condenser was stopped and 10 ml. of distillate were collected. The distillation was carried out a t such a rate that from 3 to 4 minutes were required to distill 10 ml. The collecting solution was then titrated with sodium hydroxide solution to a methyl red end point. The ionic strength of the solutions \\-as maintained a t 0.02 with sodium perchlorate; however, there was no observable change when the ionic strength was lon-ered to 0.005. The ionic strength effect is apparently small compared with the experimental errors of the method. Data and Discussion. Attempts to measure the rate of formation of ammonia by a continuous sweeping technique such as vias used in the determination of the rate of formation of hydrogen sulfide in the acid-catalyzed hydrolysis (6) proved futile because of the very high solubility of ammonia in aqueous solutions. At 25" C. only half of the ammonia could be swept by a stream of nitrogen from a solution initially 0.1I'F in hydroxide and O.OITrF in ammonium chloride in 30 minutes. I n a similar experiment at pH 11 and 50" C. only 2 5 7 , of the ammonia ivas removed after 15 minutes of sweeping.
TIME, MINUTES
Figure 4. Effect of hydroxyl ion on rate of amide hydrolysis of thioacetamide 0.001 OVF thioacetamide, 100.3' C . 1 . 0.005VF NaOH 2. 0.OlOVF NaOH 3. 0.020VF NaOH
Even a t 90' C. there was a large time lag in the removal of ammonia. Thus, neither a continuous sweeping method nor a method involving the taking of samples which would be cooled and swept was feasible. Because of these difficulties, a distillation technique was used. This introduces uncertainties because of concentration changes in the reaction solution as water and ammonia are removed, and, it does not permit a series of determinations upon the same solution unless the reaction solution is kept a t constant volume by addition of water. However, in these measurements the formation of ammonia is not of primary analytical importance and is of interest only in clarifying the mechanism of the hydrolysis of thioacetamide; as a result, the errors introduced by the distillation method are not of serious magnitude. Confirmatory experiments demonstrated that with ammonia concentrations as high as O.OlVF and hydroxide concentrations of the order of O.lVF, 9970 of the ammonia was found in the distillate if 10 ml. of a 50-ml. portion were distilled and that the quantity of hydroxide carried over in spray in the apparatus used was not measurable if a t least 4 minutes were used for distillation of the 10-ml. sample. A plot of the data resulting from experiments made a t a constant hydroxide concentration (0.OlOVF) and a t various initial thioacetamide concentrations is shown in Figure 3. I n Figure 4 are plotted the data obtained from experiments made a t constant initial thioacetamide concentration (0.00101'F) and a t various hydroxide concentrations. The scattering of points is of the order to be expected from the above-mentioned experimental uncertainties. From
the data of Figures 3 and 4 is calculated the value 9 =k 1 liter mole-I min.-l a t 100.3' C. for the second-order hydroxide-catalyzed hydrolysis of thioacetamide t o give ammonia and thioacetate. This value can be compared with that found by Crocker and Lowe ( 3 ) of 0.28 liter mole-' min.-' a t 95.9' C. for the hydrolysis of acetamide in alkaline solutions. It again shows the effect of substitution of sulfur for oxygen. Sulfide Hydrolysis of Thioacetate. Because of the uncertainties involved in volumetric determination of the actual concentrations of thioacetamide and sodium hydroxide during the initial rapid amide hydrolysis, no chemical means were found for quantitative measurements of the rate of the sulfide hydrolysis of thioacetamide. Spectral measurements were abandoned because of the mutual absorption of thioacetate and thioacetamide and the lack of pure thioacetate or thioacetic acid solutions for calibration. Consequently, quantitative measurements m-ere restricted to the hydrolysis of thioacetate ion a t a time when the initial rapid hydrolysis reactions had ceased to be predominant. However, because of analytical considerations, an estimate of the rate of the sulfide hydrolysis of thioacetamide was made; this estimate was based on the rates of the amide hydrolysis of thioacetamide and the sulfide hydrolysis of thioacetate. REAGEKTS. A buffer solution was prepared by mixing two volumes of 6VF acetic acid with one volume of 6VF sodium hydroxide to give a solution 2VF in both acetic acid and sodium acetate. Hydrogen sulfide can be swept rapidly out of samples buffered with this solution by bubbling dry nitrogen gas VOL. 30, NO. 8, AUGUST 1 9 5 8 ,
1381
Table I.
Series 1 2 3 4 5 6 7 8
a
Second-Order Velocity Constant for Hydrolysis of Thioacetate to Sulfide in Aqueous Sodium Hydroxide Solutions at 90' C.
Initial Concentrations,a VF NaOH CHsCSNH2 0.080
0.020
0.100
O . ( )15
0,100 0,100
0.025
0.100 0.200 0.200 0.300
0.020
0.030 0.020
0.100 0.020
N ~ of.
Detns. 3 4 4 3 3 4 6 3
SecondOrder Velocity Constant, ,y, Liter/Mole Min. 0.018 0.022
0.016 0.016 0.017 0.020
0.019 0.020 Av. 0.019 =k 0.0015 Ionic strength was 0.30 in all experiments; all initial concentrations are for 25" C.
through the solutions, but the p H is sufficiently high to prevent significant hydrolysis t o produce additional hydrogen sulfide by acid catalysis. Collecting solutions for hydrogen sulfide mere prepared by dissolving cadmium chloride in 6VF ammonium hydroxide to give solutions 0.4VF in cadmium chloride. Standard potassium iodate solutions, 0.002 and O.OlVF, were prepared by weight from analytical reagent. -40.1V'F sodium thiosulfate solution was prepared and standardized against the 0.01 VF potassium iodate solution. APPARATUS.The apparatus is shown in Figure 5. The reaction tube consisted of a 38 x 200 mm. test tube fitted with a tlvo-hole rubber stopper. One hole was enlarged to allow a 10ml. pipet t o be inserted into the reaction tube for sampling, A thermometer was inserted through the second hole. The reaction tube was surrounded by a mater bath which maintained the temo perature of the reaction solution at 90 =I= 1 ° C . The apparatus for removing hydrogen sulfide from each sample consisted of a 22 X 175 mm. test tube, which contained the buffered sample. A sintered-glass gas bubbling tube, through which nitrogen was passed, reached to the bottom of the test tube. The hydrogen sulfide was swept through a drawn capillary tube into a 16 X 126 mm. test tube containing the collecting solution. This size of test tube allowed the resulting cadmium sulfide precipitate to be centrifuged in a subsequent step. Repeated tests showed that the collecting solutions quantitatively retained the hydrogen sulfide, and that the hydrogen sulfide was swept out of the sample in a maximum of about 10 minutes. PROCEDURE FOR RATE MEASUREMENTS. Solutions for specific experiments were prepared from distilled water and the standard stock solutions of sodium hydroxide and sodium perchlorate. The solution was preheated in the reaction tube to slightly above 90" C., depending on the volume of thioacetamide solution to be added next. The desired volume of thioacetamide solution was pipetted into the reaction solution, so that the initial concentration of thioacetamide a t 90" C. would be known. Afterward the temperature 1382
ANALYTICAL CHEMISTRY
of the reaction solution was maintained at 90" 1" C. At timed intervals up to 3 hours, samples of the solution were pipetted from the reaction tube and transferred into 22 X 175 mm. test tubes. These mere immediately placed in an ice bath to quench the reaction. T o each sample was added sufficient buffer solution to neutralize the sodium hydroxide and buffer the sample at p H 4 to 5. The test tube which contained the buffered sample was then attached to the sweeping apparatus. The hydrogen sulfide formed by acidification of the sulfide was swept out of the sample into the collecting solutions. The collecting solutions mere centrifuged and the clear centrifugate was drawn off. The cadmium sulfide precipitate was washed into a 200-ml. flask and treated as described by Swift and Butler (6).
Data and Discussion. Experiments were conducted with various initial thioacetamide and sodium hydroxide concentrations (Table I). The ionic strength of all reaction solutions was adjusted with sodium perchlorate t o 0.30. The reaction temperature of 90' C. caused hydrolysis of thioacetate
t o proceed at a readily measured rate. I n the study of the hydrolysis of thioacetate, samples of the reaction mixture were not taken until sufficient time had elapsed to ensure that the concentration of thioacetamide had been reduced to an insignificant value. From the data obtained the second-order velocity constant for the hydrolysis of thioacetate to sulfide and acetate was calculated. ClLCUL-4TION OF THIOACETATE CONCENTRATION. Determination of the concentration of thioacetate present in the reaction solution involved several assumptions. Sulfur could be present only as the three species: thioacetaniide, thioacetate, and sulfide. -4s sampling was not begun until the thioacetamide had undergone essentially complete hydrolysis, only thioacetate and sulfide remained in significant concentrations. Determination of sulfide in the reaction solution a t timed intervals gave a measure of the rate of hydrolysis of thioacetate. CILCULATIO~OF HYDROXYL ION CONCESTRATION.There were three effects to bP considered in calculating the hydroxyl ion Concentration to be used in the hydrolysis expression: the changes in the hydroxyl ion concentration resulting from the hydrolysis of thioacetamide, of acetamide, and of thioacetate. As essentially all the thioacetamide had hydrolyzed before the thioacetate concentration was determined, sodium hydroxide equivalent to the initial thioacetamide concentration had been consumed. For the hydrolysis of acetamide in aqueous sodium hydroxide solutions, Crocker and Lowe (3) give the value of 0.28 liter mole-' min.-' for the secondorder velocity constant a t 95.9' C. From their data the second-order constant was calculated to be 0.24 liter
SINTERED GLASS
HYDROLYSIS APPARATUS
SWEEPING APPARATUS
Figure 5. Apparatus used in studies of sulfide hydrolysis of thioacetate
mole-' min.-1 a t 90" C.; hence in solutions 0.05VF in sodium hydroxide approximately 51% of the acetamide formed by hydrolysis of thioacetamide would hydrolyze during the first 60minute period after addition of thioacetamide to the reaction solution. The hydrolysis of thioacetate is so slow that in all experiments the consequent change in hydroxyl ion concentration did not exceed 5y0 during the time of reaction, and in most cases was less than 3%. The calculated values of hydroxyl ion concentration used in the hydrolysis equation were obtained from the folloming empirical expression: (OH-) = (OH-)o - (CHZCSNH2)o [(CH3CSNHg)o - (CHaCSO- )] 1 = (OH-)a - 2(CH&SNHz)o (CHaCSO-)t
+
where (OH-)o and (CH3CSNH2)orepresent the initial concentrations of sodium hydroxide and thioacetamide, respectively, and (CH,CSO-), is the calculated concentration of thioacetate at the beginning of the hydrolysis measurements. This empirical expression takes into consideration the hydrolyses of thioacetamide and acetamide, but neglects changes in hydroxyl ion concentration resulting from thioacetate hydrolysis. ORDER O F REaCTION AND CALCULATION O F VELOCITY CONSTANT. For individual experiments the concentration of thioacetate a t a given time, t , was calculated by difference and designated as C f . Successive determinations of the thioacetate concentration a t time At, after the initial deintervals, t termination were made for values of At ranging from 10 t o 160 minutes. These values are designated as Ct + at. The first-order velocity constant, k , for the hydrolysis expression
+
\vas calculated by the integrated form of the equation.
The variation in the values was caused by the dependence of the rate of hydrolysis on the hydroxyl ion concentration. When values for k/(OH-) mere calculated, making use of the values of (OH-) calculated by the empirical expression, the rate constants were found to be essentially constant, which implied that the rate of hydrolysis is first-order with respect t o both thioacetate and hydroxyl ion concentrations. Therefore, the second-order
kinetic equation for hydrolysis of thioacetate can be expressed as
where k' = k/(OH-). The second-order velocity constant with standard deviation is 0.019 =t0.0015 liter mole-' min.-' a t 90" C. TEMPERATURE DEPENDENCE OF HYDROLYSIS OF THIOACETATE. A series of experiments x a s performed to determine the change in rate of the hydrolysis of thioacetate from 70" to 90" C. The reaction solutions n-ere initially 0.030VF in thioacetamide and 0.100VF in sodium hydroxide; the ionic strength of the reaction solutions was 0.30. In Table I1 are shown the values of the second-order velocity constant, k', obtained for the temperatures studied. The energy of activation, calculated from a plot of -log k' against 1/T, is 19 kcal. per mole.
Table II. Variation of Second-Order Hydrolysis Constant of Thioacetate with Temperature"
Temp., "C. 70 80
90
Second-Order Velocity Constant, k', Liter/Mole Min. 0.0041 0.0097 0.019
acetamide hydrolyzes directly t o give sulfide. Calculations based on the rate data reported above indicate that 100 ml. of a solution O.1OVF in sodium hydroxide and 0.05VF in thioacetamide would have to be heated t o 90" C. for approximately 15 minutes to obtain 1 mmole of sulfide. There is preliminary evidence, however, that certain sulfides may precipitate as a result of a direct reaction such as that reported (6) for lead sulfide in dilute acid solutions. Reactions of metal sulfides with thioacetamide to form soluble sulfide complex ions, such as those occurring in the conventional separation of the so-called copper and tin groups, do not appear to have been studied. The use of thioacetamide for sulfide precipitations in alkaline solutions having pH values such as might be obtained with ammonia-ammonium ion or carbonate-bicarbonate buffer systems would not be practical because of the slow rate of sulfide formation, unless a direct reaction between the metal cation and thioacetamide occurs or unless the buffer system itself has some additional catalytic effect upon the rate of thioacetamide hydrolysis. Studies of the rate of hydrolysis of thioacetamide in each of the buffer systems mentioned have shown that hydrolysis of thioacetamide to sulfide is faster than is predicted on the basis of the hydroxyl ion concentration. The results of these studies are t o be reported.
Thioacetamide, 0.03OVF; sodium hydroxide, 0.100VF; ionic strength, 0.30. All initial concentrations are for 25" C.
ACKNOWLEDGMENT
The authors are grateful to the Kational Science Foundation for financial support of these studies. They are indebted to David F. Bowersox for coSULFIDE HYDROLYSIS OF THIOACET- operation with the spectrophotometric measurements and to Phillip Helnian AMIDE. By assuming an activation and Edward Seidman for preliminary energy of 20 kcal. per mole, the secondexperiments. The thioacetamide used order rate constant for the amide hydrolwas supplied by Arapahoe Chemicals, ysis of thioacetamide was calculated to Inc. have the approximate value 5 liters mole-' min.-1 a t 90" C. From this result the second-order constant for the LITERATURE CITED sulfide hydrolysis of thioacetamide was estimated to be 1 to 2 liters mole-lmin.-l (1) Barber, H. H., Taylor, T. I., "Semia t 90" C. Extrapolation of the data from micro Qualitative Analysis," the thioacetate hydrolysis studies gave Harper, New York, 1953. ( 2 ) Crocker, J. C., J . Chem. SOC.91, 593 a value of similar magnitude. (1907j. For the over-all hydrolysis of thio(3) Crocker, J. C., Lovie, F. H., Ibid., acetamide to sulfide and ammonia, the 91,952 (1907). second-order velocity constant is esti(4) Randall, it., Young, L. E., J . Am. mated to be approximately 6 to 7 liters Chem. SOC.50,989 (1928). mole-' min.+ a t 90" C. (.5 ,) Rosenthal. D.. Tavlor. T. I.. Zbid.. ANALYTICALCOXSIDERATIOKS. In 79, 2684 (1957). separate solutions with equal hydrogen (6) Swift. E H.. Butler. E. -4..ANAL. ion and hydroxyl ion concentrations, the &EM. 28,146 (1956). hydroxide-catalyzed hydrolysis of thioacetamide to yield sulfide and acetaRECEIVED for review October 7, 1957. mide is eight to ten times faster than the Accepted March 31, 1958. Contribution acid hydrolysis, but in hydroxide solu2262, Gates and Crellin Laboratories of tions only about one fifth of the thioChemistry. "
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