ACID DISSOCIATION CONSTANT OF DIETHYLDITHIOCARBAMIC ACID
3625
Determination of the Acid Dissociation Constant for Diethyldithiocarbamic Acid. The Primary and Secondary Salt Effects in the Decomposition of Diethyldithiocarbamic Acid by Keijo I. Aspila, Serge J. Joris,’ and Chuni L. Chakrabarti* Department of Chemistry, Carleton University, Ottawa 1 , Ontario, Canada
A primary and a secondary salt effect has been observed at low and high pH values, respectively, for the decomposition of diethyldithiocarbamic acid (EtZDTCH). The acid dissociation constant for EtzDTCH has been investigated in solutions of various ionic strength ( p ) in order t o evaluate the influence of p on the acid dissociation constant K O . This apparent constant, K,, has been calculated from the ionic strength parameters and the apparent and limiting rate constants for the decomposition of the dithiocarbamic acid. A value of (4.2 =t0.4) X has been obtained for the true acid dissociation constant K. by extrapolation of K , to an ionic strength equal to zero.
Introduction A survey of the literature2*aon dithiocarbamic acids indicates that the acid dissociation constants (KO)reported for these compounds are not consistent. The following investigations were made on the influence of ionic strength ( p ) on the K Ovalue of diethyldithiocarbamic acid in order to determine if the differences in these reported values are due to variations in experimental conditions such as ionic strength. The structure of EtzDTCH is shown below. CZH5
\
N-C
2
Because of the unstable n a t ~ r e ~ ~ of pmany ~ - ~ dithiocarbamic acids, direct techniques7such as titrimetry and spectrophotometry are not very suitable for determining the K , values; so indirect methods are required. One such method is based on the study of the decomposition rates since there is a direct relationship between the rate of decomposition and the ratio of the protonated and nonprotonated forms of the dithiocarbamate in aqueous solutions. I n fact, extensive studies4 have shown that the decomposition occurs only when the dithiocarbamate is in the protonated form. A rate equation for the decomposition process has been established by these studies. I n addition to the rate constants and the proton concentration, this rate expression contains the acid dissociation constant K , for the dithiocarbamic acid. By studying the variation of all the parameters in this equation as a function of ionic strength ( p ) it should be possible to evaluate the acid dissociation constant K O (expressed in terms of concentrations). Furthermore, the results obtained
by this method can be extrapolated to zero ionic strength to obtain K , (dissociation constant in terms of activities). A comparison with the results obtained from Debye-Huckel theory should also be possible.
Experimental Section Reagents. The diethyldithiocarbamic acid used was obtained commercially as the trihydrated sodium salt. The percentage dithiocarbamate was determined by amperometric titration. The acid form of the dithiocarbamate was generated by addition of the sodium salt to the aqueous buffer solutions. The buffer solution corresponding to pH equal to 5.00 and of a specific ionic strength was prepared by the method of Bates.8 The solution corresponding to p H 1.0 was prepared to various ionic strengths by the addition of appropriate amounts of KC1. Reagent grade KCl was used for all solutions to obtain a desired ionic strength. Apparatus. A modified Bausch and Lomb Spectronic 505 recording spectrophotometer was used to study the kinetics of decomposition. Matched silica cells of 10-mm path length were used and maintained at a constant temperature of 25.0 f 0.1”. All pH measurements were made with a Fisher Accumet pH
* To whom all correspondence should be addressed. (1) Postdoctoral Fellow. (2) (a) I. M. Bhatt, K. P. Soni, and A. M. Trivedi, J . Indian Chem. Soc., 45, 354 (1968); (b) A. Hulanicki, Tulanta, 14 1371 (1967). (3) D.M. Miller and R. A. Latimer, Can. J . Chem., 40, 246 (1962). (4) S. J. Joris, K. I. Aspila, and C. L. Chakrabarti, Anal. Chem., 41, 1441 (1969). (5) D.J. Halls, Milorochim. Acta, 62 (1969). (6) H. Bode, Fresenius 2.Anal. Chem., 142,414 (1954). (7) A. Albert and E. P. Serjeant, “Ionization Constants of Acids and Bases,” Methuen and Co., Ltd., London, 1962. (8) N. G. Bates, “Electrometric pH Determinations,” Wiley, New York, N. Y.,1954. The Journal of Physical Chemistry, Vol. 74,No. 80, 1970
K. I. ASPILA,S. J. JORIS, AND C. 1,. CHAKRABARTI
3626 meter, Model 210. Standardization of this pH meter with several buffer solutions near pH 5.00 assured that the solutions of various ionic strength were accurate to within 0.05 pH unit. Procedure. For the kinetic studies, solutions of M EtzDTCH were prepared by adding a to measured amount of a solution of unbuffered dithiocarbamate salt to the various buffer solutions kept in the silica cells a t constant temperature. Mixing of these solutions of buffered dithiocarbamate was .done with a Teflon (Du Pont) stirrer or a pipet bubbler and was completed within an average time of 6 sec before absorbance measurements were recorded. The pH values were checked before the addition of the dithiocarbamate and at the conclusion of the experiment in order t o check for constancy ( f0.05 pH unit). The initial concentration for the decomposition M DTC acid. The studies ranged from to plot of the log concentration vs. time was found to be linear at pH 1.0 and 5.0 and at all ionic strengths. The rate constants were evaluated t o an accuracy of f 2 % . Absorbance measurements were made at the wavelength of maximum absorbance (about 280 nm). A small shift (5 to 10 nm) in the absorption maximum was found when the pH of the solution was changed from 5 to 1. This confirmed the findings of other worker^.^
*
This equation implies that a secondary salt effect10 should be observed when the apparent rate constant, k,,,, is measured at pH 5.00 as a function of ionic strength. This is apparent from eq 2 since the rate constant is proportional to the acid dissociation constant, which is a function of the ionic strength. The following equations describe how the value of K , can be evaluated from the above rate constants when all the parameters are known as a function of the ionic strength. The acid dissociation constant is related to the thermodynamic acid dissociation constant K , (expressed in terms of activities) by the relationship
K O=
[H+][DTC-] YDTCH = Ka [DTCH] YH +YDTC-
(3)
where y are molar activity coefficients and [I stands for molar concentrations. Rearrangement of eq 2 gives the dissociation constant K , in terms of the decomposition rate constants and the hydrogen ion concentration. This equation is
(4)
However, before eq 4 can be applied to evaluate K,, it is necessary to convert the pH meter readings of hydrogen ion activities (UH +) to hydrogen ion concentrations. Theory Strictly speaking, such a conversion would require the The stability of dithiocarbamic acids in aqueous use of single ion activity coefficients (ym+) for the solutions has been described in the l i t e r a t ~ r e . ~ ~ , ~hydrogen -~ ion. However, such values are difficult to The decomposition of these acids follows a pseudo firstobtain,” and in view of the errors associated with the order rate law that is shown by measurement of pH and rate constants, it seems appropriate to use, as a first approximation, the mean V = kl[H+l [DTCIt = kapp[DTC]t (1) activity coefficients (ymh) of HCl. Furthermore, it [H+l K , must be recognized that the variation of the ionic where V = rate of decomposition; kl = limiting rate stsength must not effect the junction potentials in the constant; K O = acid dissociation constant (based on measurement of the pH. The usel’ of saturated KCl in concentration) ; [DTC It = total dithiocarbamate conthe reference electrode normally gives a fairly constant centration; [H+] = hydrogen ion concentration; kapp = junction potential for solutions of ionic strength 0.01 to apparent rate constant. 1.0. This prevents the pH from deviating more than The overall rate p r ~ f i l efor ~ , the ~ decomposition of the 0.04 unit in the ionic strength range required for these dithiocarbamic acid as a function of pH reveals a kinetic studies. The conversion12 of hydrogen ion limiting rate constant at low pH (e.g., pH 3). Indeed, activities t o hydrogen ion concentrations is done by the at these pH values the above rate equation can be use of the equation reduced in such a manner that kl equals k,,, since the value of K , is almost negligible compared to the hydro~ c = h ymddom/c (5) gen ion concentration. Study of the variation of this where yo+ = mean molar activity coefficient; ~~k = limiting rate constant as a function of ionic strength at mean molal activity coefficient; m = molality of HCI; these low pH values then gives the magnitude of the do = density of pure water; and c = molar concentraprimary salt eff ect’0 on the limiting decomposition rate tion of HC1. Equation 5 relates the molal and the constant. On the other hand, at high pH values (e.g., pH 5.00), the apparent rate constant (which is the (9) K. I. Aspila, V. S. Sastri, and C. L. Chakrabarti, Talanta, 16, measured rate constant) becomes a linear function of 1099 (1969). the hydrogen ion concentration. When [H+]
