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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Table I. Electronic Absorption, Fluorescence, and Phosphorescence Spectra of Benzamide, Benzoic Acid, and Their Conjugate Cations. Spectral Maxima Are Reported in cm-I X Benzamide Benzoic acid Neutral Neutral molemoleCation cule Cation cule v1L,(inH200rH2S04) 4.10 4.40 3.83 4.33 EIL, clLb
(in HpOor H2S04)
E 1Lb
(in H 2 0 or H2S04, 298 “K) If (in H 2 0 or HISO,, 77 “K) cf (in CHCI,, 298 “K c, (in H20 or H?SOI, 77 “ K )
13,950 3.64 1,675
9,650 3.72 800
19,200 3.35 2,300
12,000 3.65 1,130
2.61
3.14
if
c~
2.80
-a
3.08 2.98
3.12 3.10
2.85 3.08
3.33 3.25
2.50
2.53
2.44
2.60
No fluorescence was observed from benzamide in water.
with the pK, values of other carbonyl derivatives of benzene (e.g., benzoic acid, benzaldehyde, acetophenone) which lie between -7 and -9 (3) and are known to be protonated at the carbonyl oxygen atom. Moreover, if it is considered that the effect of replacing a hydrogen atom of water by a benzoyl group is to lower the pK, by about 11 units, it might be anticipated that benzoyl substitution of the ammonium ion would have approximately the same effect. Consequently, the benzoyl ammonium ion which is identical with benzamide protonated at the nitrogen atom, is predicted to have pK, 2, in excellent agreement with the experimentally determined pK, of benzamide. On the other hand, the carbonyl oxygen atom of benzoic acid is slightly more difficult to protonate than those of either benzaldehyde or acetophenone (3),and it does not seem likely that the differences in inductive or resonance properties between -OH and -NH2 would be sufficient to make benzamide seven orders of magnitude stronger as a base than benzoic acid, with respect to protonation at the carbonyl oxygen atom. Consequently, it is proposed that the ground state prototropic equilibrium of benzamide occurring in moderately concentrated Hi304 involves protonation and dissociation at the -NH2 group according to Scheme 1.
--
0
Scheme 1
416
The electronic spectra of benzamide are summarized in Table I. Those of benzoic acid are also given for comparison. The dissociation constant (pK,*) of the benzoic acidium cation in the lowest excited singlet state has been previously determined by Weller and Urban (5) to be 0 and corresponds to dissociation from the carbonyl oxygen atom. Fluorimetric titration of the benzamidium cation fluorescence yields a pK,* of -0.3. That the dissociation of the benzamidium cation occurs from the nitrogen atom in the lowest excited singlet state (i.e., that proton migration to the carbonyl oxygen atom in the excited state does not occur) is strongly suggested by the approximate values of pK,* of -0.4 and -1.3 calculated from the small protonation shift of the benzamide ‘Lb absorption band and that of the fluorescence of benzamide in frozen aqueous and sulfuric acid solutions, respectively, in conjunction with the Forster cycle (6) and the ground state pK,. These values are in reasonably good agreement with the value of pK,* of -0.3 determined by fluorimetric titration and are calculated from the spectral properties of a cation which is constrained to remain protonated at the nitrogen atom. In benzoic acid where protonation occurs at the carbonyl oxygen atom in ground and lowest excited singlet states, the protonation shifts of ’Lb band and fluorescence spectrum in frozen aqueous media are much larger, yielding approximate pK,* values of -2.0 and f2.0, respectively, for the benzoic acidium cation. The triplet state dissociation constants (pKaT) of the benzamidium and benzoic acidium cations were calculated from the Forster cycle employing the respective ground state pK, values and the protonation shifts of the structureless phosphorescence spectra (7). These were found to be -1.5 and -4.9, respectively. As is typical of triplet state dissociation constants, they are intermediate between the ground and lowest excited singlet state dissociation constants. That the difference between the ground and lowest triplet state dissociation constants of benzoic acid is greater than that of benzamide, supports the conclusion that the site protonated in the triplet state of benzamide is the nitrogen atom because protonation at a carbonyl oxygen atom which is directly involved in the conjugated system should have a greater effect on the energy of the triplet state than protonation at an amide nitrogen which is only weakly coupled to the aromatic system. RECEIVED for review May 17, 1972. Accepted September 14, 1972. (5) A. Weller and W. Urban, Angew. Cliem., 66,336 (1954). (6) T. Forster, Z . Elektruchem., 54,42 (1950). (7) G. Jackson and G. Porter, Pruc. Roy. Suc., Ser. A , 260, 13 ( 1961).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY.1973
