Table 1.
Precision and Accuracy of Automatic Titrator on Various Titrations
Results, Ml. Manual Std. dev. Automatic Titration“ 10.01 13.47 13,46 Fe++ with Ce(1V) 5.90 5.90 *o. 01 Mn++ with MnOa11.73 11.75 10.02 S2O3-- with Ce(1V) 11.72 11.74 10.01 C1- with Ag a Details given in text. * Averages and standard deviations of at least five replic,ate titrations. +
Table II. Effect of Flow Rate on Automatic Iron(l1)-Cerium(lV) Titrationa
Flow Rate. AIL p& Minute
End Std. Dev., Point, 1Il.b M1.b 1 5 90 1 0 01 2 5 92 i 0 01 1 0 01 4 5 92 6 5 93 i 0 02 12 5 90 1 0 07 a For manual titration: end point, 5.91 ml.; std. dev., 1 0 . 0 1 nil. Results are from at least five replicate titrations. averages and standard deviations of at least five replicate titrations.
manual titrations were performed by titrating to the starch iodine end point. The automatic titrations were performed at a flow rate of 4 ml. per minute and a cell current of 2.2 pa. Chloride with Silver. I n order t o test t h e automatic titrator on a precipitation titration, t h e titration of chloride ion with silver was investigated. I n this titration t h e potentials before t h e end point are determined b y t h e chloride oxidation for the anode and t h e reduction of dissolved oxygen for the cathode. As the chloride concentration becomes
Std. dev. fO.O1
f 0 . 03 10.05 1 0.03
low just before the end point, the anode shifts to more positive potentials and the anode reaction becomes the evolution of oxygen. Just after the end point, the cathode also shifts t o r a r d more positive potentials and the cathode reaction becomes the reduction of silver ions. These potential changes result in the peaked titration curve required for operation of the titrator. The titrations were performed by adding 25-ml. aliquots of 0.04M potassium chloride to 25 nil. of 1X nitric acid in the titration vessel and titrating with 0.111f silver nitrate. The manual titration was performed to the dichlorofluoroscein end point. The automatic titration was performed a t a flow rate of 3 nil. per minute and a cell current of 9 pa. Gelatin was added (0.05%) to prevent reduction of colloidal silver chloride. The flow rate of titrant in this titration is severely limited by the efficiency of the stirring. Effect of Flow Rate and Dilution of Solutions. Because the ferrouscerate titration involves a fast reaction and is relatively free from complications, i t was used t o test the effects of varying t h e rate of titrant flou- and diluting the solutions. Table I1 indicates t h a t for this ideal
titration, flow rates as high as 12 ml. per minute can be used and t h e results remain within the 1% error level. For slower reactions, of course, the maximum flow rate will be lower. \Then the concentrations of both the ferrous and cerate solutions Tvere reduced to 0.01JI and the titrations performed as before, the error increased considerably. Standard deviations increased from 0.02 ml. for 0.1S solutions to 0.09 nil. for 0.01.Y solutions. ACKNOWLEDGMENT
The authors wish to thank E. H. Schraut of the University of IYisconsin Electrical Engineering Department for his suggestions during the development of this titrator. The work was supported in part by the Research Committee of the University of \Irisconsin with funds provided by the Kisconsin Alumni Research Foundation. LITERATURE CITED
(1) Baumann, F., Shain, I., ANAL.CHEM. 29, 303 (1957).(( ( 2 ) Delahay, Paul, Ken- Instrumental Methods in Electrochemistry,” p. 382. Interscience. Sew York. 1954. (3) Hubei., C. O., Shain, I., ASAL. ’CHEV. 29, 1178 (1957). (1)Iiolthoff, I. M.,Tanaka, S . j Ibid., 26, 632 (1954). ( 5 ) Lingane, J. J., “Electroanalytical Chemistry,” p. 128, Interscience, S e x Tork, 1953. ( 6 ) lIalmstadt,, H. V., Fett, E. R., .$SAL. CHEM. 26, 1318 (1954). ( i )Reilley, C. S . ,Coolie, K. D., Furman, S . H., Ibid.,23, 1232 (1951). 181 Schmitt,. 0. H.. J . Sci. Instr. 15. 24 (1938). \
I
RECEIVEDfor review July 29, 1057. ;\ccepted February 6, 1958.
Precipitation of Cadmium Sulfide from Acid Solutions by Thioacetamide DAVID F. BOWERSOX and ERNEST H. SWIFT California Institute of Technology, Pasadena, Calif. ,The precipitation of cadmium as sulfide by thioacetamide from acid solutions has been studied. In solutions having p H values of 2 or less the rate of precipitation is controlled by the hydrolysis of the thioacetamide to give hydrogen sulfide and is first order with respect to the hydrogen ion and to the thioacetamide concentrations. In solutions having pH values from 6.3 to 3.3, a direct reaction occurs in which the rate of sulfide precipitation is first order with respect to both the cadmium ion and the thioacetamide concentrations, and is in-
1288 *
ANALYTICAL CHEMISTRY
Swift and Butler (4). They found that at p H values less than approximately 3 the rate of precipitation was hydrolysis controlled and was first order - d[Cd(II)]/dt = with respect to both the thioacetamide k[Cd(II)][CHBCSNH~]/[H+]”* and the hydrogen ion concentrations. is 8.1 X literliz min.-’ The precipitation of lead sulfide at a t 90” C. in 0.15M sodium formate higher p H values was governed b y a solution. The energy of activation for direct reaction which was dependent this reaction was calculated to b e 20.8 upon the lead ion and on the thioacet=k 1.2 kcal. per mole. amide concentrations to the first order, and on the hydrogen ion to the inverse half order. I n the precipitation of triHE precipitation of lead sulfide positive arsenic (0, no evidence was from solutions similar to those in found for other than the hydrolysisthis investigation were studied by
versely half order with respect to the hydrogen ion concentration. The velocity constant, k, in the expression
T
controlled reaction. Studies of the rates of precipitation of other metal sulfides are being made to obtain a better understanding of the principles governing these reactions and a rational basis for the analytical use of thioacetamide as a homogeneous phase precipitant for metal sulfides. Studies of the coprecipitation effects resulting from such use of thioacetamide have been initiated. EXPERIMENTAL
Reagents. Thioacetamide solutions, 1.00T'F (volume formal, formula weights per liter), nere prepared from Baker's Lot Yo. 9260 reagent grade thioacetamide. This material gave clear, colorless solutions which were never kept more than 2 weeks. Reagent grade chemicals n ere used throughout. Standard 0.05VF potassium iodate solution was prepared by weight. Sodium thiosulfate solutions, 0.05T'F 15-ere standardized against this solution. Cadmium nitrate solution, O.01VF was used. As the initial cadmium ion concentration \vas determined in each experiment, no standardization was made. Hydrochloric acid solution, W F , was prepared from concentrated acid which was free from oxidizing agents. Sodium formate-formic acid buffer solutions containing a constant sodium formate concentration were prepared from sodium hydroxide solutions and 90y0 formic acid. Solutions of various concentrations of formic acid were added to equal volumes of the sodium hydroxide solutions and then diluted to the same final volume. These buffers were used to investigate the rate of the direct reaction. Apparatus. T h e apparatus used a t 90" C. was essentially t h a t of Swift and Butler (4). Xitrogen was not bubbled through the reaction solutions. I n the experiments a t 25" C., the reaction solutions were p1acr.d in groundglass stoppered flasks. Procedure. T h e reaction solutions were prepared b y mixing measured volumes of stock solutions of thioacetamide, cadmium nitrate, and acid or buffer and diluting t o 100 ml. in t h e reaction vessel. T h e initial ionic strength \vas ronstant for each series of experiments. T h e reaction solution was then placed in a constant temperature bath. A t timed intervals approximately 12 nil. of the reaction solution nere removed from the vessel, immediately cooled in a n ice bath, and centrifuged to remove any cadmium sulfide. Duplicate 5.00-nil. portions of each sample were taken from the clear centrifugate and transferred to 15 X 125 mm. test tubes. An excess of ammonium hydroxide \\as added to each test tube and the solution was placed in a hot water bath to coagulate the cadmium sulfide precipitate. Then the mixture was cooled, the sample centrifuged, and the centrifugate carefully draivn off. The precipitate was washed with 2-ml. portions
of 3VF ammonium sulfate until the wash was sulfide free. The precipitate n a s wished into a 100-ml. conical flask; then water, potassium iodide, standard potassium iodate solution, and hydrochloric acid were added, and the solution was titrated with standard sodium thiosulfate. This method was checked by the analysis of cadmium nitrate solutions with similar quantities of cadmium. The average deviation was less than 1%. TIME,MIN.
DATA A N D DISCUSSION
Figure 1. Effect of cadmium ion concentration on rate of cadmium sulfide precipitation a t 90" C.
Qualitative experiments indicated that the rate of precipitation of cadmium sulfide from acid solutions by thioacetamide reached a minimum at approximately p H 3. This indicated that as n i t h lead(I1) (4),two types of reactions were involved in the precipitation. These types are designated as the hydrolysis-controlled and the direct reaction precipitations. Quantitative rate measurements were then made. The reactions were carried out a t 90"C. to obtain an easily measurable rate of precipitation. Hydrolysis-Controlled Precipitation. T h e rate of precipitation of cadmium sulfide from solutions of p H
0.1 OVF thioacetamide; Hydrogen ion concentration:
I. 2.
3. 4.
- dJCHaCSNH21 =
Comparison of measured cadniium(I1) concentration with that calculated from = k[H+] the expression - d[Cd(ll)l dt [CHISCSH,],where k = 0.21 liter mole-' min.-1 Initial CHSCSXH?, O.lOVF; temperature, 90" C. ~ i ~Cadmium(II), ~ , RIole/Liter Minutes Found Calculated Initial Perchloric ilcid, O.lOI'F 0 0 0 0062 . . 0.0045 0 0046 0 75 0 0024 1 75 0 0024 2 42 0 0011 0 0010 Initial Perchloric Acid, O.OlOVF
dt
k[CHSCSSH2] [ H + ] (1)
~
0.0071 0.0064 0.0058 0.0049
7.50 10.00
12.50
15.00
Table 11.
should also be the rate expression for the precipitation of cadmium(I1). The velocity constant, k , has the value 0.21 liter mole-' min.-' a t 90"C. (4). Table I compares the measured cadmium(I1) concentrations a t various times with those calculated from Expression 1; all hydrogen sulfide from the hydrolysis was assumed t o react rapidly Kith cadmium(I1). The agreement between the experimental and the calculated values is well within that of the experimental errors. As in the case of lead(I1) (4), no evidence was found of catalysis or inhibition of the reaction by cadmium ion or by the cadmium sulfide precipitate.
0.0088
5.00
0.0038
x
1 and 2 was independent of cadmium ion concentration a n d $vas first order with respect t o both hydrogen ion and thioacetamide concentrations. T h e precipitation is, therefore, controlled b y the hydrolysis of the thioacetamide; also there is rapid precipitation of the cadmium by the hydrogen sulfide thus formed. If this is true, the rate expression for the hydrolysis of thioacetamide, which can be expressed as
Table 1. Hydrolysis-Controlled Precipitation of Cadmium Sulfide
0.0
4.8 x i o - 4 ~ 1.2 10-4~ 2.8 X 10-6M 2.0 X 10-6M
0.0073 0.0065
0.0059 0.0049
0.0038
Effect of Hydrogen Ion Concentration on Rate of Cadmium Sulfide Precipitation
[Initial cadmium(II), 0.01OVF; thioacetamide, 0.1OVFl IH+1 dlCdi1I)l 4 . 8 X lo-,' 1 . 2 x lo-,' 2 . 8 X lo-& 2 . 0 x lo-"
5 . 0 x 10-'1
35 X 10-lo 14 x 1 0 - 1 0 6 . 1 X 10-lo 1.6 X , . .
18
X
9 . 1 x 10-7 4 . 3 x 10-7 1. I x 10-7
0.6 X Av. 1 . 3 i~ 0 . 2 X
.
1.6 X
8 . 1 X lo-' 1 . 3 X lo-' 8 . 4 X IO-' 1 . 2 x 10-7 8 . 1 x 10-6 1.1 x 10-7 7.8 x 1 0 - 6 ... 8.2 X 8 . 1 -f 0 . 2 X
VOL. 30, NO. 7,JULY 1958
1289
%
Experiments a t p H 1 and 5.7 rvere carried out a t 90" C. with solutions containing the initial sodium formateformic acid buffer constituents. The presence of the buffer constituents did not change the reaction mechanisms. The rates of the precipitations conformed to the hydrolysis-controlled mechanism a t pH 1 and the direct reaction mechanism at p H 5.7 with no change in the velocity constants.
5'
10-2
5
Table 111. Effect of Thioacetamide Concentration on Rate of Cadmium Sulfide Precipitation a t 90" C.
I
1
I
I
I
I
I
2
3
4
5
[Initial Cd(II), 0.010VF; hydrogen ion, 0.00012M]
1
i
PH
Figure 2. Rate of cadmium sulfide precipitation by thioacetamide as functions of pH
Analytical Considerations. Figure 2 shows t h e calculated rates of precipitation of cadmium sulfide b y t h e hydrolysis-controlled and b y t h e direct reaction in solutions which are O.1OVF in thioacetamide and 0.OlOVF in cadmium(I1) a t 90" C. and which have t h e range of p H values shown. The direct reaction is dependent upon t h e cadmium ion concentration; therefore, t h e direct reaction curve will be changed vertically by changes in the cadmium ion concentration. For solutions having the concentrations specified above, the following calculations can be made: Rd = Rh a t pH 2.95, where Rd is the rate of precipitation by the direct reaction, and Rh is the rate by the hydrolysis controlled reaction. Rd = Rh/800 at pH 1 Rd
O.1OVF CHsCSSH2; 0.OlVF Cd(N03)*; 0 050 0 10
o is 0
0 20 0 25 0 30
3 65 x 10-6 8 2 X 8 3 X 8 4X io 9 x 10-5 7 3x 16 5 X 10-6 8 4 X 20 2 x 10-5 8 0 X 23 9 X 10-5 8 0 X A\,. 8 07 zt 0 3 X
go9 10-4
10-4 10-4 10-4
Precipitation by Direct Reaction. Formic acid-sodiuln formate buffers tve1e used in experiments in the pH range from 6.3 t o 3.3. These Tvere designed to study any direct reaction similar to that found n ith lead(I1) and to show the effect of hydrogcn ion, cadmium ion, and thioacetamide concentrations on such a reaction. Effect of Concentration. CADIIIUM IO?;. As plots of the logarithm of the cadmium(I1) coiicciitration against the time for the x-nrious hydrogen ion concentrations were linear (Figure 1). the reaction is first older with respect t o cadmium ion concentr8tion. HYDHOGEN ION. The effect of hydrogen im concentration was studied tvitli solutions in which the total ionic strength and the thioacetamide and the cadmium ion concentrations were all initially constant. The hydrogen ion concentration varied from 5.0 x 10-4 to 5 X 10-7?'F. Rate constants calculated from the data (Table 11) show that in this range the rate of precipitation is dependent on the inverse half power of the hydrogen ion concentration. This is the same dependence as found for lead(I1) (4). As yet, a plausible mechanism n hich will explain thi. effect has not been formulated. THIOACETARIIDE. The effect of the thioacetamide concentration upon the rate of the reaction was investigated over a range of thioacetamide concentrations from 0.05 to 0.308F. Rate constants calculated from the data are given in Table 111. The rate of precipitation is dependent upon the 1290
ANALYTICAL CHEMISTRY
c.
0 Hydrolysis of thioacetamide 0 Direct reaction
thioacetamide concentration to the first power. CHLORIDEION. Chloride ion decreased the rate of the precipitation reaction: With the formal concentration of chloride five times that of cadmium, the rate found was only 50% of that calculated. ii quantitative investigation of this effect is in progress. Temperature Effect on Direct Reaction. The dependence of the rate on the hydrogen ion and cadmium ion concentrations a t 25' C. over the p H range from 5.7 t o 3.3 were studied. The results indicate first-order dependence on the cadmium ion concentration and inverse half-order dependence upon the hydrogen ion concentration (Table 11). The rate of precipitation was so slow that the change in the cadmium ion concentration could not be measured with an accuracy greater than approximately 10%. Consequently, the calculated constants vary more than those for the experiments a t 90" C. However, the data clearly indicate that the mechanism of the reaction is similar to that a t 90" C. in the same hydrogen ion concentration range. The calculated velocity constant is 1.3 =k 0 2 X liter1'2 mole-"* min.-l The energy of activation for the direct reaction mas calculated from data obtained a t 25", 70", and 90" C. A plot of log k us. 112' was a straight line, thus showing a temperature dependence conforming to the Arrhenius equation. The activation energy was 20.8 f 1.2 kcal. per mole. These data show the analytical potentialities inherent in controlling the rate of formation of a cadmium sulfide precipitate by temperature control.
= 1250 Rh
at pH 5
I n solutions of p H 2 and lower, precipitation of cadmium sulfide is primarily through hydrolysis. As the p H is raised, the direct reaction becomes increasingly important until under the above specified conditions it is predominant a t p H 4 and greater. These observations are similar to those in the precipitation of lead sulfide by thioacetamide (4). The two cations exhibit much the same behavior in their precipitation from acid solutions. If it is assumed that the precipitation of cadmium as the sulfide by thioacetamide below p H 2 is preceded by hydrolysis to hydrogen sulfide, as this investigation indicates, then the data obtained may be used to calculate the necessary conditions for controlled homogeneous phase precipitations. I n a solution O.1T'F in hydrochloric acid, O.lOVF in thioacetamide, and initially O.01VF in cadmium salt. the time required for the hydrolysis to furnish hydrogen sulfide equivalent t o the cadmium present u-ould be 5 minutes a t 90" C. At lower temperatures the time lvould be correspondingly longer. This type of precipitation does not fulfill the conditions used for sulfide precipitations made by means of gaseous hydrogen sulfide, as it is general practice to saturate the solution with the gas a t atmospheric pressure; this provides an aqueous hydrogen sulfide concentration of approximately 0.1V F . Complete hydrolysis of the thioacetamide under the above conditions would not saturate the solution; furthermore, at 90" C. the time required for half the thioacetamide to hydrolyze is calculated t o be 33 minutes. Care should be exercised in replacing gaseous hydrogen sulfide lvith thioacetamide without careful investigation, especially where only a slight excess of thioacetamide is added. Barhcr and Taylor (2) stated that
“the hydrolysis is more rapid in an alkaline solution than in a n acid solution of the same strength.” I n a solution of p H 6, in which the initial cadmium ion and the thioacetamide concentrations are 0,010 and O.lOVF, respectively, the time required for half of the cadmium to precipitate by the direct reaction is less than 9 minutes at 90” C. I n this time nnd under these conditions, the cadmium sulfide precipitated by the hydrolysis-controlled reaction would be less than 10-8 mole. This indicates that the more rapid precipitation observed a t higher p H values could be caused by the direct reaction mechanism rather than an increase in the rate of hydrolysis. An investigation of the hydrolysis of thioacetamide in alkaline solutions is in progress. The velocity constant (4) for the
precipitation of lead sulfide by thioacetamide in the p1-I range 5.1 to 3.5 was 1.15 X lop3 liter1’* mole-1’2 minute-1. This is 42% higher than the constant for the precipitation of cadmium sulfide in the same pH range. The concentration of cadmium(I1) in a saturated solution a t a given sulfide concentration should be 52% greater than that of lead(I1) in a similar solution (3). There may be some correlation between this increase in solubility and the decrease in the velocity constant. Such a relationship cannot be postulated until the velocity constants for the rate of precipitation of other basic cations arc determined. ACKNOWLEDGMENT
Preliminary experiments by L. B.
LIarantz were of value in planning this investigation. Comments on the manuscript by R. C. Greenough have heen helpful. LITERATURE CITED
( I ) Barber, H. H., Taylor, T. I., “Semimicro Qualitative Analysis,” Harper, Kew York, 1953. (2) Butler. E. 4.. Swift. E. H.. -4NAL. CHEX 29, 419 (1957). (3) Kolthoff, I. M., J . Phys. Chem. 35, \
,
2711 (1931).
(4) Swift, E. H., Butler, E. A , ,4s.4~. CHEM.28, 146 (1956).
RECEIVEDfor review -4iigust 5, 1957. llccepted January 30, 1958. Contribution No. 2181 from the Gates and Crellin 1,aboratories of Chemistry, California Institute of Technology Pasadena, Calif. Supported hy the Yational Science Foundation.
Spectrophotometric Assay for Reaction of N-Ethylmaleimide with Sulfhydryl Groups EUGENE ROBERTS and GEORGE ROUSER Department of Biochemistry, City o f Hope Medical Center, Duarte, Calif.
b Ultraviolet spectral changes accompanying the interaction with Nethylmaleimide in dilute aqueous solution a t p H 6.0 can b e utilized for the quantitative estimation of sulfhydryl compounds. The maximal changes in absorbance occur a t wave lengths a t which minimal interference would occur from many ultraviolet-absorbing substances of biological interest.
S
the first report of the quantitative reaction of S-dhylnialeimidc with thiol groups (6), this type of reaction has been applied to the determination of various sulfhydryl compounds on paper chromatograms ( I , 8 ) and in tissue sections ( 9 ) . Changes in the ultraviolet absorption spectrum of Sethylnialeiinide upon reaction with sulfhydryl compounds have been observed (4, 6 ) , and it has been suggested that thew changes could be used for quantitative purposes (6). The present report s h o m that a t pH 6.0 the loss in absorption of M-ethylmaleimide a t 300 mp, the absorption maximum, can be used for the quantitative estimation of sulfhydryl groups in dilute solution in the presence of excess S-ethylmaleimide. Thus, 1 pmole of cysteine per ml. yields a n absorbance change of 0.60. The change in absorbance is proportional to the ISCE
concentration of cysteine in the range studied. Interference by a number of compounds of biochemical interest, such as purines, pyrimidines, nucleosides, and nucleotides, is minimal a t the w i r e length and p H employed. If necessary, wave lengths up to 320 nip can be used without too great a loss in sensitivity. This procedure overcomes the objections to the spectrophotometric method for following the reaction of p-chloromercuribenzoate with sulfhydryl groups a t 250 mp (6) and yet makes possible kinetic measurements without removal of individual samples for analysis ( 3 ) froin a reaction mixture.
RESULTS
Spectral shifts accompanying rmction of cysteine with 4-ethylmaleiinide are shown in representative curves (Figure 1). The spectrum of ’Y-ethylmaleimide, c u r w 1, has a maximuin a t 300 mp and a miiiimuni at 248 nip, Cystrine alone liai io distinctive 3 p ~ 1.0 r
R
.
k\
0.9
O.6 0.7 Ly
EXPERIMENTAL
Reagents. L-Cysteine hydrochloride monohydrate and glutathione were purchased from Schwarz Laboratories, Inc., and crystallized bovine plasma albumin from Armour a n d Co. T h e N-ethylnialeimide was obtained from Delta Chemical Works, Inc. Buffers were prepared from analytical reagent grade salts. Apparatus. 911 determinations were made with a Beckman Model DK-2 recording spectrophotometer. Kinetic measurements were made with t h e aid of a n adapter which a l l o m continuous measurement a t a given wave length, t h e curve progressing a t t h e rate of 1 inch per minute.
1
0.1
230
250
270
3W
324
WAVE LENGTH ( n p l Figure 1.
Absorption spectra
0.1M phosphate buffer, p H 6.0 1. 2.
3. 4. 5. 6.
1.5 X 1 O - 3 M N-ethyimaleimide 0.1 5 X 1 0 - 3 M cysteine with 1.5 X 1 O-3M N-ethylmaleimide 0.75 X 10-3M cysteine with 1.5 X IO-jM N-ethylmaleimide 1.27 X 1 O-3M cysteine with 1.5 X 1 O-3M N-ethylmoleimide 1.50 X 1 O-3M cysteine with 1.5 X 10-M N-ethylmaleimide 1.5 X 10-3M cysteine VOL. 30, NO. 7 JULY 1958
1291
