~~
Table I. Separation of Tribromophenol from Its Mixtures with Other Analogous Compounds" Recovery Melting of Quantity point of tribromoof recovered phenol, adduct, tribromophenol, "C Adduct mg P-Benzoquinone 10 99 95 96 Naphthalene 10 99 95 96 Phenol 10 100 95 96 Pentabromophenol 10 100 95 96 Pentachlorophenol 10 95 96 99 a Tribromophenol(50 mg) was mixed.
z
---
Table 11. Melting Point of Tribromophenol Substance Melting point, "C Ordinary tribromophenol 90 91 Tribromophenol sublimated in tube furnace 95 96 Tribromophenol sublimated by ordinary method 93 94
-
N
-
N
To examine if it is possible to separate tribromophenol from other analogous compounds such as pentabromophenol, etc., 50 mg of each of these compounds was put in the sublimation tube, the pressure reduced to 10-1 mm Hg, and sublimed for 1 hour in the gradient temperature tube furnace. The resulting position of the sublimation zone of these compounds is shown in Figure 6. As the figure shows, the positions of the sublimation zone of these compounds are different from that of tribromophenol. It was, therefore, concluded that it might be possible to separate tribromophenol from these compounds. Accordingly mixtures of tribromophenol and other compounds were sublimed. By weighing the contents of sublimation zone corresponding to tribromophenol, the recovered quantity of tribromophenol was determined. The results are shown in Table I. Of course, sublimation zones
corresponding to the other compounds, which were added as impurities, appeared in the tubes. Table I shows that tribromophenol can be separated quantitatively from other compounds. Purification of Tribromophenol. Ordinary tribromophenol was purified by sublimation. The results are shown in Table 11. The tribromophenol which was purified in the gradient temperature tube furnace is higher in melting point and purer than when purified by the ordinary sublimation method. In the process of sublimating tribromophenol in the gradient temperature tube furnace, there was formed a secondary sublimation zone. The melting point at the secondary sublimation zone was 88 "C. Though we could not determine the quality and quantity of the substance forming the secondary sublimation zone at once, this method was recognized to be more convenient than the normal vacuum sublimation method for the purification of reagents and the concentration of impurities. RECEIVED for review October 15, 1971. Accepted August 179 1972.
Stability of Metal Dithiocarbamate Complexes R. R . Scharfe, V. S. Sastri, and C. L. Chakrabartil Department of Chemistry, Carleton University, Ottawa, Ontario K l S 5B6, Canada SEVERAL INVESTIGATIONS were carried out earlier in this laboratory to clarify important aspects of the chemistry of dithiocarbamic acids, such as the decomposition mechanism (1-3) and the basicity of dithiocarbamic acids (4). Complexation reactions of dithiocarbamic acids were reviewed by Hulanicki (5). A careful study of the literature (9,however, shows a lack of definitive information on the stability constants of metal dithiocarbamate complexes. Recently, stability constants of metal complexes with morpholine dithiocarbamic acid were determined by the solvent extraction method (6). The objectives of the present study were as follows: i) to 1
All correspondence to be addressed to this author.
(1) K. I. Aspila, V. S. Sastri, and C. L. Chakrabarti, Talanta, 16, 1099 (1969). ( 2 ) S. J. Joris, K. I. Aspila, and C. L. Chakrabarti, ANAL.CHEM.,
647 (1970). (3) K. I. Aspila, S. J. Joris, and C. L. Chakrabarti, ibid, 43, 1529 (1971). (4) S. J. Joris, K. I. Aspila, and C. L. Chakrabarti, ibid., 41, 1441 (1969). (5) A. Hulanicki, Talanta, 14, 1371 (1967). (6) V. S. Sastri, K. I. Aspila, and C L. Chakrabarti, Can.J . Chem., 47,2320 (1969). 42,
determine the stability constants of complexes of some divalent metal ions with a variety of N-substituted dithiocarbamic acids, and ii) to determine the effect of various substituents at the nitrogen atom in disubstituted dithiocarbamic acids on the magnitude of stability constants. Solvent Extraction. Solvent extraction of a metal ion into a chloroform solution of a dithiocarbamic acid (HDTC) may be represented as follows (6, 7).
and the equilibrium constant for the above extraction reaction is called the extraction constant, K , and is given by
K = IM(DTC>nIorgIH+l" [M"+I[(HDTc)or~I'
(2)
In the presence of a masking agent, X, which forms a nonextractable metal complex, MXy, where X stands for the anion of the masking agent,
(3) (7) J. Stary and K. Kratzer, Anal. Chim. Acta, 40, 93 (1968).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY 1973
413
*
Table 1. Stability Constants: Log 82 Determined by Solvent Extraction“ and PolarographicbMethods Znz+ Ni2+ Pbz+ Solvent PolarogSolvent PolarogSolvent extraction Polarography Ligand extraction why extraction raphy PyrrDTC 9.8-10.4 10.5 12.4-1 3 . 0 13.2 16.8 17.1 PipDTC 11.5 11.7 13.9 12.1 16.9 16.8 HexaDTC 12.9 13.1 14.3 14.7 17.7 ... EtzDTC 11.4 11.6 12.9 12.1 18.3 17.7 MezDTC 9.1 ... ... ... ... ... Solvent extraction was done at an ionic strength of 0.01 Polarography was done in 1M KCI. pK values taken from reference (11).
where q is the distribution ratio of the metal ion in the presence of the masking agent, and py is the stability constant of the metal complex, MXy. The stability constant, pnOrg,of the metal dithiocarbamate complex in the organic phase is given by the equation:
where PHDTC is the partition coefficient of HDTC, and is defined by
and KHDTC is the dissociation constant of HDTC, and is defined by
KHDTC=
[H+][DTC-1 [HDTC]
(7)
Polarography. When known amounts of a metal ion and two ligands, DTC and NTA, are present (NTA being present in large excess over that of the metal ion while DTC is present in twice the concentration of the metal ion), the following expressions hold good for the metal DTC and the metal NTA complexes.
MZf $- 2 DTC-
M(DTC)2
(8) (9)
M2+
+ NTAZ-
P1
=
M(NTA)
[M(NTA)I [M2+][NTA2-]
(10) (1 1)
After the equilibrium is attained, the concentration of free (uncomplexed) DTC can be determined polarographically. The following equations hold good at the equilibrium
Assuming that the concentration of free M2+is very small (a reasonable assumption because the complexes have high stability), p2 can be evaluated. Allowances must be made for the competition of H+ ion for NTA. The values of such allowances and 61 can be obtained from the literature (8). (8) A. Ringborn, “Complexation in Analytical Chemistry,” Interscience Publishers, New York, N.Y., 1963. 414
pKsc
3.29 3.20 3.28 3.38 3.22
EXPERIMENTAL
Apparatus. The absorbance was measured with a Bausch and Lomb Spectronic 505 recording spectrophotometer at a constant temperature of 22.0 “C, the cell compartment being thermostated with a Haake thermostat. The polarographic measurements were carried out with a Metrohm E261 recording polarograph. All pH measurements were made with a Fisher Accumet pH meter, Model 210. Reagents. The dithiocarbamic acids used in this study were prepared as the sodium salt by reacting CSZwith the appropriate amine in a basic medium (9, IO). All the other chemicals used were of analytical reagent grade quality. The water used was first distilled and then deionized. The purity of dithiocarbamic acids was determined by amperometric titration with a standard solution of Zn2+ at a pH of 6-7, and by measuring the diffusion current due to the reduction of ZnZf at the applied potential of -1.25 volts with respect to Ag/ AgCl reference electrode. The purity of dithiocarbamic acids was also determined by spectrophotometric titration with standard solution of Cu2+, and the results agreed with those obtained by the amperometric titration. The water content of the dithiocarbamic acids was determined by Karl Fischer titration (11). Procedure. SOLVENTEXTRACTION STUDIES. The general procedure is given in the literature (6, 12). Aqueous solutions of known quantities of DTC (4 X 10-6 to 4 X 10-4M), of the metal ion (10-5 to lO-dM), and of competing ligand (NTA or EDTA) (10-3M), were added to a separatory funnel, and the volume made up to 25 ml with deionized water. The pH was maintained constant with a borate or phosphate buffer. An equal volume (25 ml) of chloroform was added to the separatory funnel, and the solutions were gently shaken for 15 minutes. After 15 minutes (by which time the equilibrium had been attained), the phases were separated, and the spectrum of the organic phase was recorded. The actual concentration of free dithiocarbamic acid was obtained by measuring its distribution coefficient separately. The competition of H+ ion with the metal ion for the masking ligands NTA or EDTA was allowed for by calculating the effective concentration of the masking ligand (8). POLAROGRAPHIC STUDIES.An aqueous solution containing 10-4 to 10-5M of the metal ion was mixed with 4 X low4to 4 X 10-511/1’ solution of DTC and 0.1M solution of NTA, the final volume being made up to 50 ml with deionized water and 1 M in KC1; the resulting solution was left at 25 “C for an hour for the equilibrium to be attained. The equilibrium attained was independent of the order in which the reagents (9) D. J. Halls, A. Townshend, and P. Zuman. Anal. Chim. Acta, 40, 459 (1968). (10) H. L. Klopping and G. J. Van der Kerk, Rec. Truc. Chim., 70 917 (1951). (11) K. I. Aspila, M. Sc. thesis, Carleton University, Ottawa, Ontario, Canada, 1971. (12) J. Stary, “Solvent Extraction of Metal Chelates,” The Macrnillan Co., New York, N.Y., 1964.
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY 1973
were mixed. (An hour was found to be sufficient time for the equilibrium to be attained.) Then the concentration of free DTC ligand was determined by polarography at the applied potential of -0.6 to -0.7 volt relative to Ag/AgCl electrode. No interference due to free NTA was observed in the polarographic determination of the concentration of free DTC. f12 for the metal complex, M(DTC)?, was evaluated by employing this value of the free DTC. RESULTS AND DISCUSSION
The stability constants (expressed as log p2) of dithiocarbamate complexes of ZnZ4,Ni2+,and Pb2+ together with the acid dissociation constants are presented in Table I. There is a fair agreement between the values of stability constants obtained by the solvent extraction and the polarographic methods; the differences could be due to the fact that widely difrerent ionic strengths, which were respectively 0.01 and 1.0, were employed in the two methods. The polarographic method necessitated the use of 1M KC1 as the supporting electrolyte and, hence, a high ionic strength. It should be noted that the stability constants determined by the solvent extraction method at the low ionic strength of 0.01 should approach the thermodynamic stability constants. The values determined by the polarographic method at the high ionic strength of 1.0 are the concentration stability constants. In the case of Pb2+, there is a linear relationship between the stability constant values and the pK (the basicity) of three dithiocarbamic acids. In the case of Zn2+and Ni2+, there is no regular trend in the stability constants with the increasing basicity of dithiocarbamic acids. Thus, there is no general correlation of the stability constants with the basicity of the
sulfur in the dithiocarbamic acids in question. The order of stability constants follows the natural order of Irving and Williams (13). The order of the stability constants for Zn*+, Ni2+,and PbZ+is hexamethylene DTC > pipDTC > PyrrDTC (see Table I). The metal complexes of dithiocarbamic acids were employed by earlier workers to determine various metal ions, singly. The magnitude of the stability constants presented in this paper lends theoretical support to the existing analytical procedures for determining NiZ+ or Cu2+ in the undermentioned mixtures by a simple spectrophotometric titration: Ni2+ in the presence of Zn2+, or Cu2+in the presence of Zn2+ and Pb2+( 6 ) . It is clear from the stability constants that pyrrolidine DTC forms complexes of the same degree of stability as the widely used diethyl DTC. However, at low pH of 1.0, pyrrolidine DTC (rIl2 = 30 min for decomposition) is more stable than diethyl DTC (tlj2 = 7.5 sec for decomposition) (11). Thus, the substitution of pyrrolidine DTC in many existing procedures which call for diethyl DTC will add an additional element of flexibility to many analytical procedures.
RECEIVED for review June 13, 1972. Accepted September 25, 1972. The authors are grateful to the National Research Council of Canada for the financial support of this work. This work is part of an investigation performed by R. R. Scharfe in partial fulfillment of the requirements for an MSc. degree presented by Carleton University. (13) H. Irving and R. J. P. Williams, Nature, 162,746 (1948).
Ionization Equilibria in the Ground, Lowest Excited Singlet, and Lowest Triplet States of Benzamide W. Larry Paul and Stephen G. Schulman College of Pharmacy, University of Florida, Gainesoille, Fla. 32601
DURINGTHE COURSE of fluorimetric and photochemical investigations of some arylcarboxamides of pharmaceutical significance, it became obvious that very little was known about the electronic spectroscopy and sites of ionization of this important class of organic compounds. Consequently, the present study of the pH dependence of the electronic spectra of benzamide was undertaken. EXPERIMENTAL
Absorption spectra were taken on a Beckman DB-GT spectrophotometer. Fluorescence spectra were taken on a Perkin-Elmer MPF-2A fluorescence spectrophotometer whose monochromators were calibrated against the xenon line emission spectrum and whose output was corrected for instrumental response with a rhodamine B quantum counter. Reagents. Sulfuric acid and chloroform were purchased as the reagent grade materials from Mallinckrodt Chemical Works, St. Louis, Mo., and used without further purification. Benzamide and benzoic acid were obtained from Matheson, Coleman and Bell, Inc., East Rutherford, N.J., and each was recrystallized several times from ethanol before use. Apparatus.
RESULTS AND DISCUSSION
Katritzky et a f . (I) have studied the protonation of benzamide in sulfuric acid media and reported a pK, of -2.1. Yates and coworkers (2, 3 ) have studied the protonation of the same compound in perchloric acid and determined a pK, of -2.0. The data points from which the pK, values of benzamide in both sulfuric acid and in perchloric acid were determined were found to yield straight lines with slopes of less than unity when fitted to the Henderson-Hasselbach equation ( I , 2). It has been suggested that this is indicative of protonation at the oxygen atom of benzamide ( 4 ) ; the Ho scales in H2S04and HCIOl failing to refer to the standard states of the protonated and unprotonated oxygen atoms of benzamide. However, this does not appear to be consistent (1) A. R. Katritzky. A. J. Waring, - and K . Yates, Tefruhedrarz, 19,
465 (1963). (2) K. Yates and J. B. Stevens, Can. J . C h e m . , 43, 529 (1965). (3) K. Yates and H. Wai, ibid., D 2131. (4) A. Albert, in “Physical Methods in Heterocyclic Chemistry,” Vol. 111, A. R. Katritzky, Ed., Academic Press, New York, N.Y., 1971, p 5.
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