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pK's of the Naphthoic Acids
p K s of the Naphthoic Acids in the First Excited Singlet State 0. Chalvet, H. H. Jaffe,"' and Ernest0 de la Serna Centre de Mechanique Ondulatoire Appliquee, CNRS, 750 19, Paris, France (Received January 17, 1975)
The acid-base equilibria of a- and P-naphthoic acids, with their respective conjugate acids, 1,l-dihydroxya- and 0-naphthylcarbonium ions, and conjugate bases, a- and P-naphthoate anions, in the excited states are reexamined theoretically. The ultraviolet absorption spectra of the six compounds are discussed in an Appendix.
Wehry and Rogers,:! in a paper on deuterium isotope effects of excited state pK's, have given the pK's of the first singlet state of a- (1) and 0-naphthoic acids (2). Vander Donckt and P ~ r t e rbecause ,~ of a dissimilarity of the values with those of the anthroic acids: have reexamined 1 and 2, and concluded that the values given by the earlier authors involved, at least in part, the two-step acid-base equilibria between the protonated forms of 1 (3) and 2 (4)and the corresponding anions 5 and 6. We wish to show in the present note that the conclusion of Vander Donckt and Porter are theoretically reasonable, and that the discrepancy might have been detected by purely theoretical methods. Accordingly, we have performed CNDO/S5 calculations of the absorption spectra of 1-6 and applied the Forster cycle6 to the data obtained. The results of this work are given in Table I. Since experimental determinations of the pK's of the lowest triplets are also available? we have equally made the calculations for the triplet states. From the data of Table I one can readily calculate the various ApK according to
TABLE I: Absorption Spectra of Molecules 1-6 (AllEnergies in eV) Obsd Molecule 3 1
5 4 2
6
Singlet calcd
Absorptn
3.18 3.96 3.97 3.08 3.89 3.93
3.86 3 -96 3.70 3.77
Fluorscn 2.70 3.14 3.51 2.67 3.26 3.48
dl
4,OO
ApK = pK* - PKG = 17AE at 298 K, where AE is in electron volts. The values obtained for the equilibria 1 & 5 and 2 F! 6, 0.2 and 0.7,respectively, are smaller than the experimental values of Vander Donckt and P ~ r t e rThis . ~ discrepancy arises largely from the fact that these authors have used mean values of absorption and fluorescence spectra, and that the Stokes shift decreases sharply upon ionization of the free acids. The pK's of protonation of the free acids, i.e., for the equilibria 3 e 1 and 4 a 2 are 13.3 and 13.8, respectively. The values of ApK found by Wehry and Rogers2 agree with neither of the two equilibria calculated, nor with the overall, two-proton equilibra 3 & 5 and 4 F! 6; however, their values are near the average of the values for the protonation of the free acids, and the overall process. This, then, is in perfect accord with the postulate of Vander Donckt and Porter3 that the measurements involved the absorption of the free acids and the emission of their conjugate acid. Our results are shown graphically in Figure 1, in which the transition energies for the transition to the first excited singlet are shown. The slopes of the line segment drawn represent relative values of ApK. Figure 1 strikingly confirms the conclusion of Werner and Hercules' that, under the influence of light, Le., in the first excited state, the increase in basicity of a carbonyl oxygen atom is substantially greater than the increase in acidity of an OH group. Since we have had to perform rather extensive calcula-
3'001r c,,H,c(oH);
C,,H,
COOH
C,, ti, COO -
Figure 1. Variation of the transition energy as a function of protonation in the acid-base equilibria of the naphthoic acids.
tions of the spectra of the molecules and ions 1-6 here discussed, we are reporting the results obtained in an Appendix which is available as supplementary material. Supplementary Material Available. An Appendix will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper only or microfiche (105 X 148 mm, 24X reduction, negatives) containing all of the supplementary material for the papers in this issue may be obtained from the Business Office, Books and Journals Division, American Chemical Society, 1155 16th St., N.W., Washington, D.C. 20036. Remit check or money order for The Journal of Physical Chemistry, Vol. 79, No. 23, 1975
Helen 8. Brooks and F. Sicilio
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(6) T. Forster, Z. Nektrochem., 54, 42 (1950). (7) T. Werner and D. Hercules, J. fhys. Chem., 73, 2005 (1969): 74, 1030 (19701. ( 8 ) K. Nishimoto and N. Mataga. Z. Phys. Chem. (Frankfurt am Main), 12, 335 (1957); 13, 140 (1957). (9) J. R. Platt, J. Chem. Phys., 18, 1168 (1950); J. Opt. SOC.Am., 43, 252 (1953). (IO) Note the similarity of this behavior to that of doublet and triplet states of many other molecules, discussed elsewhere: H. M. Chang, H. H. Jaffe, and C. A. Masmanidis, J. fhys. Chem., 79, 1109, 1118 (1975). (1 1) E. Clar, “Aromatische Kohlenwasserstoffe”, 2nd ed, Springer, Berlin, 1952. (12) Y. Hirschberg and R. Norman, Can. J. Res., 278, 437 (1949); R. M. Hochstrasser, Can. J. Chem., 39, 1776, 1853 (1961); V. N. Lisltsyn, Didenko, and Dashevskii, Zh. Org. Khim., 4, 1086 (1968). (13) J. T. D’Agostino and H. H. Jaffb, J. Am. Chem. SOC.,91, 3384 (1969).
$4.50 for photocopy or $2.50 for microfiche, referring to code number JPC-75-2543. References and Notes (1) Author to whom correspondence should be addressed at the Department of Chemistry, University of Cincinnati, Cincinnati, Ohio 45221. (2) E. L. Wehry and L. B. Rogers, J. Am. Chem. SOC.,88, 351 (1966). (3) E. Vander Donckt and G. Porter, Discuss. Faraday Soc.. 64, 3215 (1968). (4) E. Vander Donckt and G. Porter, Trans. Faraday SOC., 64, 3218 (1968). (5) J. Del Bene and H. H.Jaffe, J. Chem. fhys., 48, 1807, 4050 (1968); 49, 1221 (1968); 50, 426 (1969); R. L. Ellis, G. Kuehnlenz, and H. H. Jaffe, Theor. Chim. Acta, 26, 131 (1972).
Electron Spin Resonance Study of the TiF2+-H202 Reaction System Helen B. Brooks and F. Sicllio* Department of Chemistry, Texas A&M University, College Station, Texas 77843 (Received June 25, 1975) Publication costs assisted by The Robert A. Welch Foundation
The Ti(III)-HzO2-substrate reaction system has been used extensively to generate organic free radicals in aqueous media. The Ti(II1) decay in this system is not normally observed by ESR. We have studied the reaction of TiF2+ with H202 in which decay of the reactant TiF2+ can be followed by ESR. Inclusion of an organic substrate to the TiF2+-H202 system leads to the formation of organic radicals as in the case of the Ti(III)-H202 system. The TiF2+-H202 system has the distinct advantage that both a reactant and the organic radical can be followed simultaneously, leading to an obvious delineation of mechanism,
Introduction The reaction of H202 with Ti(II1) in the presence of an organic substrate has been used to generate organic radicals in aqueous media.1,2 Kinetic studies have been made with a continuous flow system monitored by ESR.2,3However, the reactants cannot be monitored by ESR. The kinetics of Ti(II1) decay have been studied by a stopped-flow spectrophotometric t e ~ h n i q u e . However, ~ these studies must be carried out with excess Ti(III)* and the ESR studies must be carried out in excess H ~ 0 2 . ~ In this work fluoride was added to the Ti(II1) reactant solution to form the complex species TiF2+ which can be monitored by ESR.5 This system allows a reactant and also radical intermediates to be observed simultaneously so that the mechanism may be delineated more critically. Experimental Section Reagents. Titanium(II1) solutions were prepared from W. H. Curtin & Co. technical grade 20% Tic13 and were standardized by titrations with KMn04 (KyGz04 was the primary standard). H202 solutions were prepared from Baker Analyzed Reagent grade 30% solutions and standardized with KMn04. Baker reagent grade NaF and Fisher Certified ACS NaF were both used for preparing solutions. Methanol, sulfuric acid, and hydrochloric acid were reagent grade. All solutions were deaerated by bubbling nitrogen for 20 min. Equilibrium Studies. Two mole ratio studies were done The Journal of Physical Chemist!y, Vol. 79, No. 23, 1975
on a Cary 14 spectrophotometer. Concentrated Ti(II1) and F- solutions were mixed and diluted to volume with deaerated water. Two wavelengths, 4350 and 4860 8,were selected, and additional data for mole ratio and continuous variation experiments were collected with a Beckman DU. The ESR data were developed using a Varian 4502-15 spectrometer and a Varian V-4556 flat cell. Each set of data was collected continuously to minimize errors. The flat cell was removed between solutions, but the instrument was retuned by adjusting the flat cell position rather than the instrument settings. Known runs gave agreement within f5% for this technique. Two sets of data were collected a t different pW, using both ESR and a Beckman DU on the same solutions. The pH was monitored with a Corning Model 101 digitaI electrometer and HCl was added to maintain constant pH. Kinetic ESR Procedures. The flow system and ESR apparatus has been described previously.6 Solutions were prepared from deaerated water and stored in 10-1. glass reservoirs. Under nitrogen pressures of 2 atm, the separate streams were forced through a standard Varian 4547 mixer and a Varian 4548 quartz flat cell via polyethylene tubing. The flow rate was regulated by a needle valve on the exit side of the flat cell. An average flow rate was measured by timing the collection of 100 ml of product solution. The system was set up so that the flow rates from each reservoir were identical. The “dead times”, t, for the kinetic data were calculated by the relationship t = V / r where V is the dead volume and r is the flow rate. Spacers were inserted
