1595
Communication:; to the Editor use solutions vvith varying concentrations of the organic compound but a constant concentration of a single electrolyte. Because the thermodynamic theory of electrocapillarityz requires that the relative surface excess of the organic compound be calculated from measurements made a t constant electrolyte activity, the practice of using constant electrolyte concentration is equivalent to making the implicit assumlption that the activity of the electrolyte is unaffected by the presence of a neutral organic compound. During a n investigation of the electrosorption of 2-butanol on mercury from aqueous sodium sulfate solution^,^ we tested this assumption by measuring the activity of the sodium sulfate in the three-component mixtures a t 25" by the emf method. We conclude, on the basis of this test, that for this system at least the assumption is completely untenable. A galvanic cell without liquid junction consisting of a Corning NAS 11-18 sodium ion electrode4 and a twophase lead amalgam-lead sulfate electrodes was used. In the absence of organic compound a plot of the emf of this cell us. the natural logarithm of the mean ionic activity6 over a concentration range 0.05-1.4 m had a least-squares slope which agreed with the theoretical value of 1.5RTIF to within better than one part per thousand. To prove that the cell also behaved correctly in the presence of 2butanol we measured the emf in the presence and absence of the organic compound with solutions saturated with solid sodium sulfate. The emf in the two cases agreed to within 10 p V . Emf measurements were then made on a series of solutions having different concentrations of NaZS04 and of 2butanol in order to determine what concentration of Na2S04 would be required, for a given concentration of 2-butanol, to yield the same emf (ie.,salt activity) as 0.1 M NazS04 in pure water. The results are shown in Figure 1. Had the assumption that the organic compound does not seriously affect the electrolyte activity been true, a horizontal line with ordinate 0.1 M would have been obtained. From Figure 1 it may be seen that this assumption is definitely false. For example, when the concentration of 2-butanol is 1.0 M the concentration of NazS04 required to give the same emf as the solution of 0.1 M NaZS04 in pure water is only 0.0498 M. Moreover, a solution which was 1.0 M in 2-butanol and 0.1 M in NaZS04 was found to have a salt activity nearly five times larger than that of 0.1 M NaZS04 in pure water. These results show that for this. system very serious errors would result if the electrosorption data were collected from a series of solutions with constant salt concentration. Conclusions about the nature of the electrosorption isotherm reached from the analysis of such data on the assumption that the salt activity was constant would be doubtful if not meaningless. For such systems the first step must be to obtain data such as that of Figure 1 to provide the recipe for preparation of a series of solutions of varying organic concentration a t constant salt activity. The observed effect of 2-butanol on the electrolyte activity may be related to the structure-making properties of this alcoh01,~and, on this assumption, the linear relation shown in Figure 1 suggests a simple correlation. If the structured water in the form of cages around the alcohol molecules is, on the average, unavailable to the ions of the salt, then the addition of the alcohol to the electrolyte solution will effectively raise the salt concentration in the remaining unstructured water. On this assumption we calculated the number of water molecules associated with
0.0
.J
.6
.9
1.2
1.5
WOLARITY OF 2-BUTANOL
Figure 1. Plot of the molar concentration of sodium sulfate which yields the same electrolyte activity as 0.1 M Na2S04 in pure water vs. the corresponding molar concentration of 2-butano1 in the solution at 25".
one alcohol molecule and obtained an average number of 24.2. This number could be consistent with a cage of 24 water molecules around the alcohol molecule in the form of a tetrakaidecahedrons which has 12 pentagonal and 2 hexagonal faces. Molecular models indicate such a cage would be able to contain a molecule of 2-butanol.
Acknowledgment. This work was supported by the U.'S. Air Force Office of Scientific Research under Grant No. AF-AFOSR-70-1887. We thank Professor T. N. Solie for helpful discussions. (1) For recent reviews ci. B. E. Damaskin, 0. A. Petrii, and V. V . Batrakov, "Adsorption of Organic Compounds on Electrodes," Pienurn Press, New York, N.Y., 1971; R. Payne, J. Eiectroanal. Chem., 41, 277 (1973). (2) For a review cf. D. M. Mohilner in "Electroanalytical Chemistry," Voi. 1, A. J. Bard, Ed., Marcel Dekker, New York, N.Y., 1966, pp 241-409. (3) H. Nakadornari, D. M. Mohilner, and P. R. Mohiiner, to be submitted for publication. (4) G. Eisenrnan, Advan. Anal. Chem. Instrum., 4, 213 (1965). (5) H. S. Harned and J. C. Hecker, 2. Amer. Chem. SOC.,56, 650 (1934). (6) H. S. Harned and B. B. Owen, "The Physical Chemistry of Electrolytic Solutions," 3rd ed, Reinhold, New York. N.Y., 1958, p 553. (7) G. Nernethy and H. Sheraga, J. Chem. Phys., 36, 3401 (1962); G. Jap., 35,646 (1962). Wada and S. Umeda. Bull. Chem. SOC. (8) J. L. Kavanau. "Water and Solute interactions," Holden-Day, San Francisco, Calif., 1964, p 17.
Department of Chemistry Colorado State University Fort Collins, Colorado 80521
David M. Mohilner* Hisamitsu Nakadomari
Received March 79, 7973
Substituent Effects on Excited-State Acidities of Some Substituted 8-Hydroxyquinolinium Cations'
Sir: Considerable interest has recently appeared concerning the acidity of organic molecules in their excited states. It has been shown that in most cases acidities of these molecules were significantly different in the ground (1) Research was carried out as a part of a study on the phosphorimetric analysis of drugs in blood and urine, supported by U. S. Public Health Service Grant No. GM-11373-08.
The Journal of Physical Chemistry, Voi. 77, No. 72, 1973
Communications to the Editor
1596
state and the fiist excited singlet state,2,3 which is of importance in photochemistry. Prototropic equilibria involved in the lowest singlet excited state of quinolinols have especially received great attention.4-11 However, significant study of substituent effects upon excited-state protolytic equilibrium constants of these heterocyclic compounds has not been reported until now. It is only in the homocylic aromatic series that some quantitative studies based on the Hammett equation have been described.12-14 For example, Wehry and Rogers13 demonstrated by means of the Hammett and Taft equations that, in a series of substituted phenols, conjugative effects on acidity were more important than inductive effects, in excited states compared to the ground state. Wehryl4 reported also that the extent of electron-withdrawing conjugation by sulfone, sulfonium, and sulfoxide groups greatly enhanced acidities of phenols in lowest triplet, and first excited singlet states, relative to the ground state. The purpose of the present study was to examine the influence of halogeno-, sulfo- and thiocyano-substituents upon the first singlet excited-state dissociation constants of a series of substituted derivatives of %-hydroxyquinoline (8-HQ). We were particularly concerned by the possibility of correlation of excited-state acidities of these heterocyclic compounds with ground-state substituent constants. First excited-state protolytic equilibrium constants (pKa*) were determined oia fluorescence titration for the equilibria between the excited singlet states of the cations and the zwitterions of 8-hydroxyquinoline substituted derivatives shown in Scheme I. We have already described in B: Prototropic Equilibrium between First Excited Singlet State of Cation and Zwitterion of Substituted 8-HQ (X and Y are 5 and 7 Substituent). X X ~~~~~~
H+
detail the methods of fluorescence titration used for the determination of pKa*.ll Data are given in Table I for 22 mono- and disubstituted 8-HQ derivatives. The pMa* values are between -5.9 for '7-iodo-$-hydroxyquinoline and -9.6 for the 5,7-disulfo-S-hydroxyquinoline, which represents a large range of nearly 4 decades. For such a large acidity range, substituent effects can be considered as significant, in spite of the fact that substituents include mainly halogens. All of the groups in this study exhibit electron-withdrawing effects in the ground state. As expected, sulfo and thiocyano derivatives are relatively strong acids; this is partly due to electron-withdrawing conjugation of sulfur d r orbitals probably enhanced in singlet excited states, as it has been observed in the case of sulfur-containing phenols.14 In the case of' the halogeno derivatives, it is of interest to notice that excited state acidities of 5- and 7-halogeno-%-hydroxyquinoline are decreasing with increasing electron-withdrawing power (or electrone,gativity) of the halogen atom; for example, 7-fluoro-$-hydroxyquinoline is a stronger acid than 7-iodo-8hydroxyquinoline in the first excited singlet state. Similar sequences are observed for the dihalogeno derivatives. This acidity order i s the reversal of that which would be The Jourrial of Physical Chemistry, Vol. 77, No. 12, 1973
TABLE I: PKa" Values for 8-Hydroxyquinoline Sybstituted Derivatives in the Singlet Excited Statea Substituent 5-Br
PK~* -8.60 7.30
-
5- I 5-SO3H 5-SCN 7-F 7-CI 7-Br 7-1 7-SO3H 5-F-7-CI 5-CI-7-CI
'
-9.30 -8.95 -9.40
-9.20 -8.65 -5.90 -9.15 -8.80 -8.75
Substituent 5-Br-7-CI 5-1-7-CI 5-F-7-Br 5-CI-7-Br 5-Br-7-Br 5-I-7-Br 5-F-7-1 5-Br-7-1 5-1-7-1 5-1-7-503H 5-SO3H-7-SO3H
PK~* -8.40 -7.10 -8.60 -8.55 -8.20 -6.95
-7.95 -7.50 -6.40
-6.15 -9.60
a Values recorded at room temperature in sulfuric acid solvent." experimental values given to nearest 0.05 unit. Error 50.1 5.
All
expected on the basis of the values of ground-state substituent constants, and as a result, the Hammett equation can not be applied with any success to the substituted 8-HQ. These anomalous sequences of pKa* for the halogeno-8hydroxyquinoline could not result from the enhancement in the excited state of proximity steric or polar interactions between ortho halogeno substituents (in I position) and the phenolate oxygen atom. Indeed, the sequence of pKa* values for the 5-para substituents I > Br > C1 > F is similar to the one observed for the 7-0-halogeno series and appears to be independent of the nature of the 7-halogeno substituent. An alternative explanation would be a balance which would occur differently in the excited state, between conjugative and inductive contributions to the global electron-withdrawing effect of the halogen atoms. However, that the anomalous acidity order observed in this study is also the reverse of the one noticed by Wehry and Rogers13 for 3-halogenated phenols in the first excited singlet state suggests to us that there also may be some influence of the heterocyclic part of the molecule on the acidity sequence of 8-HQ derivatives in the excited state. In order to evaluate the respective contribution of inductive and resonance effects of the halogen substituents upon the excited state acidities of derivatives of 8-HQ,16 we have used the Taft equation16 (a) C. A. Parker, "Photoluminescence of Solutions," Elsevier, Amsterdam, 1968, pp 328-341. (b) R. S. Becker, "Theory and Interpretation of Fluorescence and Phosphorescence," Wiiey, New York, N. Y . , 1969, pp 239-244. E. L. Wehry and L. B. Rogers, "Fluorescence and Phosphorescence Analysis," D. M. Hercules, Ed., Wiley, New York, N. Y., 1966, Chapter 3, p 125. C. J. Haylock, S. F. Mason, and 8. E. Smith, J. Chem. SOC., 4897 (1963). R. E. Bal!ard and J. W. Edwards, J. Chem. SOC., 4868 (1964). S. Schulman and Q. Fernando, J. Phys. Chem., 71,2668 (1967). S. Schulrnan and Q.Fernando, Tetrahedron, 24,1777 (1968). S. F. Mason, J. Philp, and B. E. Smith, J. Chem. SOC.A, 3051 (1968). S. G. Schulman and H. Gershon, J. Phys. Chem., 72,3693 (1968). M. Goldman and E. L. Wehry, Anal. Chem., 42, 1178 (1970). M. P. Bratzel, J. J. Aaron, J. D. Winefordner, S. G. Schulman, and H. Gershon. Anal. Chem., 44,1240 (1972). H. H. Jaffe and H. L. Jones, J. Org. Chem., 30, 964 (1965). E. L. Wehry and L. E. Rogers, J. Amer. Chem. SOC., 87, 4234 (1965). E. L. Wehry, J. Amer. Chem. SOC.,$9, 41 (1967). Because of the lack of available values of aI.and uR , SOsH and SCN substituents were not included in the correlations.' C. D. Ritchie and N, F. Sager, "Program in Physical Organic Chemistry," Vol. 2, 5. G. Cohen, A. Streitwieser, and R. W. Taft, Ed., Wiley, New York, N. Y., 1964, p 323.
1597
Communications to the Editor
0-no
/
5 - 1 7-1
6
Figure 1. Correlations between
ductive substituent constants. log K = P I U I
excited-state dissociation constants of the monohalogeno and dihalogeno 8-hydroxyquinolines and in-
+
~ R u R
+
log KO
(1)
which permits a semiquantitative evaluation of inductive (I) and resonance (R) effects. Values of UI and UR proposed by T a f P were used. Therefore, the electrical effect of the halogeno substituents were assumed to be approximately equal in ortho and para positions, as suggested by several authors.17-20 Mono- and disubstituted derivatives are separated according to two correlations and these correlations are described below. The least-squares line for the five monosubstituted 8-HQ is log R,* = 21.101 4.90, + 0.05 (2)
+
The reaction constants are PI = 21.1 and PR = 4.9. The multiple correlation coefficient r = 0.89. For the 11 disubstituted 8-HQ's, the least-squares line is log K,"
= 9.42~14- 0.82aR
-
0.05
(3) with PI = 9.4, PR = 0.8, and r = 0.90. That satisfactory Taft correlations are found confirms that only electrical effects are involved and that there are no important steric or polar proximity interactions of 7-halogeno substituents with the ortho phenolic group. The considerably larger PI values compared to PR in eq 2 and 3 indicate that inductive effects are relatively more important than the conjugative effects in the first excited singlet states of the halogeno 8-hydroxyquinolines. The sensitivity of the prototropic equilibria in the excited singlet state to the inductive effects is also demonstrated by the satisfactory correlations found with only UI constants (see Figure 1).For the mono- and dihalogeno derivatives, the least-squares lines are, respectively
+
log K,* = 1 6 . 6 ~ ~ 1.05
(4)
log K," 12.4ZuI - 3.25 (5) with r = 0.95 for eq 4 and r = 0.96 for eq 5. By comparison with previous studies which indicated that enhanced conjugative interactions occurred in the excited s t a t e , 1 2 ~especially ~~ in the case of phenols, the present results may appear surprising. However, we must point out that the only substituents included in our correlations are for the halogeno species which are well known
for their ability to simultaneously use p r orbitals to donate a electrons to the ring and d r orbitals to accept electrons from the a orbitals of the aromatic ring.z1 As this last factor is included in the total inductive electron-withdrawing power, represented by UI constants,Z2 it is probable that the increase of the inductive effect observed in the correlations is partly due to enhanced interactions of the aromatic a electrons with the halogen dr orbitals in the first singlet excited state of the 8-HQ derivatives. Another part of the enhancement of the inductive effect would result from the influence of the heterocyclic part of the molecule on the acidity in the excited state. Most of the halogen electron-donating conjugative interactions involved in the zwitterion-cation system would be mobilized in the direction of the nitrogen of the heterocyclic ring in accordance with the enhanced basicity of the nitrogen atom observed in the lowest excited singlet state. In the mesomeric forms A and B (Scheme 11), the inductive effects of the halogeno substituents would be greatly enhanced and would consequently strengthen the acidities of the excited cations. Occurrence of the forms A and B in the first singlet excited state is supported by Huckel molecular orbital calculations,23 which predict a migration of electronic charges from the homocyclic to the heterocyclic ring of the 8-hydroxyquinoline upon excitation to the first excited singlet state. Different behavior of the monohalogeno and dihalogeno derivatives is indicated by the different slope and ordinate values found for the two series of compounds (Figure 1 and eq 4 and 5). Electrical interactions of the halogen substituents with each other in the dihalogenated derivatives may occur in the singlet excited state and partly cause the separation of the data in the two correlations. This interpretation is confirmed by the existence of a very satisfactory multiple regression of the general type de(17) M. T. Tribble and J. G . Traynham, J. Amer. Chem. SOC., 91, 379 (1969). and references cited therein. (18) C. L. Liotta, Chem. Commun., 338 (1968). (19) This view has recently been disputed by Charton,zo but the results of this controversy would not affect the validity of the correlations obtained in the present work. 91,6649 (1969). (20) M. Charton, J. Amer. Chem. SOC., (21) L. N. Fergusson, "The Modern Structural Theory of Organic Chernistry," Prentice-Hall, Englewood Cliffs, N. J., 1963, pp 393-401. (22) R. W. Taft, Jr., J. Chem. Phys., 26,931 (1956). (23) R. E. Burton and W. J. Davis, J. Chem. SOC.,1766 (1964). The Journal of Physical Chemistry, Vol. 77, No. 12, 1973
1598
Communications to the Editor
Scheme 11: Mesomeric Forms of the 5-Halogeno-8hydroxyquinoline (A) and of the 7-Halogeno-8hydroxyquinoline (B).. :X:,
@
:x:
scribed by Leffler and Grunwald,24 with an interaction term proportional to the product of the two UI values, for the dihalogeno derivatives
log K,*
10.6Zlq
- 1 . 7 Z u ~ u-~ 215
(6) with axand UP inductive constants of the halogens X and Y, PI = 10.6, interaction constant qI = -1.7, and multiple correlation coefficient r = 0.998. The significantly different PI value compared with the PI value of 16.6 for the monohalogeno derivatives (eq 4) indicates a greater sensitivity of these last compounds to the inductive effect. =
(24) J. E. Leffler and E. Grunwald. “Rates and Equilibria of Organic Reactions,” Wiley, New York, N. Y., 1963, pp 192-194. (25) Present address, Carleton University, Ottawa, Ontario, Canada. (26) On leave from the Laboratoire de Chimie Organlque Physique, Paris, France.
Department of Chemistry University of Florida Gainesville, Florida 32601
M. P. BratzeIz5 J. J. Aaronz6 J. 0 . Winefordner*
College of Pharmacy University of Florida Gainesville, Florida 32607
S. G. Schulman
Boyce Thompson lnstitute for Plant Research, Inc., Yonkers, New York 10701
H. Gershon
Received November 3, 1972
Reply to the Comments of Desnoyers on the Paper, “lOniC Solvation Numbers from Compressibilities and Ionic Vibration Potentials Measurements” Publication costs assisted by The Flinders University
Sir: We wish to discuss the points in Desnoyers’ communicationl in order. The Journalof Physical Chemistry, Vol. 77, No. 12, 1973
(1) “It is better to use the apparent molar compressibility and apparent molal volumes to determine solvation numbers from compressibility data. In this way, one eliminates interionic effects.” This could be done; we think the advantage would be nugatory. Our work is aimed at models of the solvation shell at infinite dissolution. The changes with codcentration are better interpreted in terms of the interpenetration of solvation shells. (The correlations with this model are shown in our paper,2 Figure 10 and Table VIII.) A similar model has been used by Ramnathan and Friedman.3 ‘(2) “The partial molar expansivity can be related to hydration numbers and, in this case, certain assumptions give rise to a negative value for the hydration number.” We do think this relates to our paper, where we have used the relative compressibilities of solvent and solute to calculate solvation numbers. The assumption (we correct for the partial degree of its applicability in our calculation2) is that the water molecules are completely incompressible a t sufficiently high field. The assumption of Dr. Desnoyers’ equation relating expansivity to hydration numbers is that both the hydration number and the volume of the water molecules in the hydration sheath are independent of temperature. This assumption seems to have a different status from that which we have used. We suggest the contradiction in sign arises from the lack of applicability of the assumptions. Corresponding remarks may be applied to his comment on heat capacity data. (3) In respect to the heat capacity data for sodium chloride solutions in H20 and DzO, we cannot agree that they are relevant to the interpretations of the compressibility data. They tell us something about how the heat capacities of solutions and salts vary in these two solvents, and it seems likely that the effects of structural changes outside the solvation sheath may be different, for heat capacities than for compressibilities. Now let us move to the general thought behind the comments of Dr. Desnoyers, and provide a general answer. In the work of Bockris and Saluja,4 the ionic vibration potentials have been combined with compressibility measurements to give individual solvation numbers, because the first measurement gives the sum and the latter measurement the difference of solvation numbers. This is the essential new thing about the experimental approach of Bockris and Saluja and the essential new thing about the theoretical approach is to distinguish quantitatively for the difference between the coordination number of the ion in aqueous solution and the number of water molecules temporarily attached to it, while it moves in the soluti~n.~-? Dr. Desnoyers’ essential point is to suggest that the fact that we have assumed (for univalent ions) that the change of the solution compressibility can be given predominantly by the change in the layer which the solvation and coordination waters inhabit, and then one further struc(1) J. E. Desnoyers, J. Phys. Chern., 77, 567 (1973). (2) J. O’M. Eockris and P. P. S. Saluja, J. Phys. Chern., 76, 2140 (1972). (3) P. S. Ramnathan and H. L. Friedman, J. Chem. Phys., 54, 1086 (1971). (4) J. O’M. Bockris and P. P. S. Saluja, J. Phys. Chern., 76, 2298 (1972). (5) J. O’M. Bockris, Quart. Rev. Chem. SOC.,3, 173 (1949). (6) J. O’M. Eockris and A. K. N. Reddy, “Modern Electrochemistry.” Plenum Press, New York, N. Y.. 1970. (7) 0. Ya Samoilov, ”Structure of Electrolyte Solutions and Hydration of Ions,” Consultants Bureau, New York, N. Y., 1965; Discuss. FaradaySoc., 24, 141 (1957).
