Nitrobenzene Free !on as a Hydrogen Bond Acceptor g vapors. It seems therefore that the photosensitizer must be some other factor in the glass surface. A similar phenomenon was discovered in the photolytic decomposition of hydrogen molecules on a silica gel surface.21 It is remarkable that the number of radicals observed in these experiments was always found to be smaller by approximately four orders of magnitude than the substituent groups available. This cannot be explained kinetically by equilibrium between production and decay, because maximum intensity is reached in minutes whereas half-times of decay are measured in days, sometimes even weeks. Thus one cannot rehate the low saturation concentration to a steady state between these two processes and we conclude that the production ~ r thec,e ” stable radicals must require some additional, rather rare conditions on the surface. One possibility is that the pholtolytic process needs the presence of &me trace impurities which exist only on a few sites. This
1027
impurity could be a metal oxide similar to the findings of Fujita, et al.,22,23 who observed photosensitization by trace amounts of Vz05 and Moo3 in the formation of adsorbed methyl radicals. It would be interesting to try to create different kinds of radicals by irradiation of other chemically bound species. Our attempts to observe radicals by irradiation of aminated glass were unsuccessful. Acknowledgments. We wish to thank Professor B. L. Silver for his helpful advice. We would also like to thank Professor D. Wolf from the Weizmann Institute of Science, Rehovoth, Israel, for a gift of 1 7 0 enriched HzO and CH30H. (21) G. 8. Pariiskii, I. D. Micheykin, and V. 8. Kazanskii, Dokl. Akad. Nauk SSSR, 1,180 (1 968). (22) Y. Fujita, K. Hatano, M. Yanagita, T. Katsu, M. Sato, and T. Kwan, Bull. Chem. SOC.Jap., 44, 2884 (1971). (23) T. Katsu, M. Yanagita, and Y. Fujita, J. Phys. Chem., 75, 4064
(1971).
Equilibrium Studies by Electron Spin Resonance. I I I . The Nitrobenzene Fre Acceptor erald R. Stevenson” and Hector Hidalgo Department of Chemistry, University of Puerto Rico, Rio Piedras, Puerto Rico 00931 (Received August 23, 1972) Publication cosis assisted by The University of Puerto Rico
Hydrogen bonding proton donors have been added to the PhNOz anion radical in hexamethylphosphoramide (HMPA). An equilibrium has been observed between hydrogen bonded HMPA plus the free PhNOz anion radical and free HMPA plus hydrogen bonded PhN02. The equilibrium constant for this reaction has been algebraically related to the nitrogen hyperfine coupling constant of the PhNOz anion radical. A temperature-dependent study of this equilibrium has resulted in the first relative hydrogen bonding energies by electron spin resonance. At lower temperatures (from 0 to -10”) line width alternat,ion is observed, due to rapid forming and breaking of the hydrogen bonds to the PhNOz anion.
The variation of mr coupling constants upon solvent and counterion interactions in organic anion radicals has been a topic of much recent interest.l Aromatic anion radicals containing polar groups, such as the nitrobenzene (PhNQ2) anion radical, exhibit extreme sensitivity to these effectsS2-‘ In the casF-of the PhNOz anion radical, the magnitude of the isotopic nitrogen coupling constant, A,, always increases with external forces such as solvent and counterion interactions.5 To date it has been difficult to obt,ain quantitative solvent effects upon the esr parameters and th$ thermodynamics of these interactions, since there always existed the possibility of interference with the ion pairing interactions and the interactions of the solvent with itself. Recently, the PhN& anion radical has been recorded essentially free from solvent and counterion interactions by lithium reduction in hexamethylphosphoramide MPA).e This “free” PhNOz anion radical (a)was ob-
served simultaneously with its ion pair (@), allowing the enthalpy of reaction 1 to be determined from temperature dependent esr ~ t u d i e s . ~ , ~
P
a Utilizing the lowest nitrogen splitting for the PhNO2 anion radical an empirical equation was developed which (1) M. T. Jones in “Radical Ions,” E. T. Kaiser and L. Kevan, Ed., Interscience, New York, N. Y., 1968,Chapter 6. (2) J . M. Gross, J. D. Barnes, and G. N. Pillans, J. Chem. Soc., 109
(1969). (3) J. M. Gross and M. C. R. Symons, Trans. Faraday Soc., 63, 2117 (1967).
(4) W. M. Gulick. Jr., and D. H. Geske, J. Amer. Chem. SOG.,87.4049 (1965). (5) C. Y. Ling and J. Gendall, J. Chem. Phys., 47,3475 (1965). The Journal of Physical Chemistry, Vol. 77, No. 8,1973
Gerald R. Stevenson and Hector Hidalgo
1028
relates the total A,, to the contributions made to it by the counterion and solvent A, = 8.48
+ P,A,(solvent) + P,A,(M+)
N
(2)
A,(solvent) and An(M+) represent the maximum contributions to A, due to the solvent and cation interactions, respectively. Pb and P, are the fractions of these interactions found in a particular system. Several workers have observed line width alternation phenomena, resulting from the modulation of the nitrogen hyperfine splitting, for the PhNO2 anion radical.*-Io In all cams, however, the line width alternation has been attributed to the rapid equilibrium between two different ion pairs.638.9 In this report, we wish to communicate the effect of hydrogen bonding upon the nitrogen coupling constant of the PhN02 free anion radical, the effect of hydrogen bonding upon the equilibrium between the free ion and ion pair (eq I), and the tbermodynamics of this hydrogen bonding.
Results and Disa:ussion Addition of srnall amounts of alcohol or water to the system PhNID2-HMPA-Li affords increases in the nitrogen coupling constant for the formally free ion ( a ) due to hydrogen bonding of the acidic proton of the alcohol with the nitro group of the PhN02 anion radical. However, no change in the nitrogen splitting of the ion pair (B) is observed (Figure E I. Hydrogen bonding does not occur with the ion pair as evidenced by the lack of change of the nitrogen splitting upon addition of the alcohol or water. This is presumably due to the fact that the nitro group of the ion pair is too strongly complexed with the lithium cation to act as an acceptor in hydrogen bonding. From Figure 1, it is clear that A, is a linear function of the alcohol or water concentrations for very low concentrations of alcohol or water Thus Pd,(alcohol) is equal to a constant times the alcohol concentration. The constant is a function of the alcohol used and the concentration of the anion radical Since the observed nitrogen splitting smoothly increases upon addition of proton donor, this value must be a time average between the free ion (a) and the hydrogen bonded ion (a'). As more proton donor is added to the solution the concentration of the time averaged species increases a t the expense of the ion pair (Figure 2). From this, it is clear that hydrogen bonding stabilizes the formaly free ion relative to the ion pair. This is an expected result if the enthalpy of hydrogen bonding to a is finite. At high concentrations of proton donor (greater than about 1 M ) no esr signal can be obeerved for the ion paired species. If our conclusilon that the experimental A, i s due to a time averaged between a: and a' is correct, we would expect a slower modulation of the nitrogen splitting as the temperature is lowered resulting in line width alternation. The g values of the two radicals should be the same or very close resultitng in broadening of the m = -1 and m= + I tines in the region of line width alternation. The m = 0 lines should reimain sharp. In agreement with this prediction, dramatic line width alternation is observed when these systems containing proton donor are studied under low temperature conditions (Figure 3). The line width.alternation cannot be due to rapid exchange between 01 and /3, since both can be observed simultaneously at room temperature. The line width alternation is due to the rapid exchange between the free ion and the hydrogen The Jcurnal of Physical C,%emistry,Vol. 77, No. 8, 1973
-
10.85
9.2
-
9.0
-
8.8
-
8.6
-
8.4
-
AN
[M~OH] Figure 1. Plot of A n for the free ion (lower) and the ion pair (upper) vs. the methanol concentration ( M ) .The coupling constants are in Gauss.
3
2
&0.4
0.2
[MeOH] Figure 2. Plot of the concentration of the free ion (a') divided by the ion pair concentration (B) vs. the methanol concentration (M).
bonded ion. For dilute solutions of alcohol or water the line width alternation increases rapidly with increasing proton donor concentration. Figure 3 shows that the esr spectrum a t -10" for the PhNO2 system containing 0.38 M methanol exhibits strong line width alternation while the same system not containing methanol does not show any line width alternation effect. Essentially only the 18 esr lines resulting from the proton splitting of the m = 0 component of the nitrogen coupling can be seen a t -10". At still lower temperatures, the simultaneous observation of a: and a' would be expected. This cannot be accomplished due to the fact that the solvent freezes a t about -10". A simple plot of esr line width us. 1/RT allows us to estimate the magnitude of the energy of activation for reaction 3. The energies of activation estimated in this manner are shown in Table I. For the case of water, R G. R. Stevenson, L. Echegoyen, and L. R. Lizardi, J . Phys. Chem.,
76,1439 (1972). G. R. Stevenson, L. Echegoyen, and L. R. Lizardi. J. Phys. Chem., 76,2058 (1972). G. R. Stevenson, L. Echegoyen, and L. R. Lizardi, J. Phys. Chem.,
submitted for publication. F. J. Smentowski and G. R. Stevenson, J. Amer. Chem. SOC., 90, 4661 (1968). J. M. Gross and J. D. Barnes, J. Phys. Chem., 74, 2936 (1970).
1029
Nitrobenzene Free Ion as a Hydrogen Bond Acceptor
TABLE I:Thermodynamic Parameters Con~ro~li~g Equilibrium 3 and A,‘ for the Proton Donors E,,
AS”, kcal/ Proton donor
An’,
Methanol fert-Butyl alcohol Water
G
14 13 14.5
K e g at 25” AN”, kcal/mol
3.5f 0.2 1.0 f 0.4 0.9 f O . l
1.2 & 0.2 1 1.6 f 0 . 2
eu
mol
1.6
1.8 5.0 2.8
5.5
sents the concentration of the hydrogen bonded ion, A,’ represents the esr splitting for the hydrogen bonded species, and A , is the observed nitrogen splitting. The thermodynamic equilibrium constant is given by Keg = (a’)(HMPA)/(a)(HMPA’), where (HMPA’) represents the concentration of the HMPA h ~ ~ r o gbonded e ~ i to the proton donor. Combining this equation for with eq 4, we obtain the expression -
ICeq= ( A ,
Figure 3. Esr spectra for the PhN02-HMPA-Li system as a function of temperature. The spectra on the left are the low-field halves for the system without added methanol. The spectra on the right are for the system contaiining 0.38 M methanol: (A) -lo”, only the free ion on tho left and the time-averaged species on the right can be observed; (B) 25”,the ion pair is observed simultaneously with the free ion and with the time-averaged species: (C) TO”, only the ion pair (can be observed. represents the other proton hydrogen bonded to a molecule of HMPA.
a
H-OR ! !
+
HMPA
(3)
a! HMPA itself i s a good acceptor for hydrogen bonds.ll For very dilute solutions of alcohol there is very little selfassociation of the alcohol, and the hydrogen bonding equilibrium can be described by eq 3. The time averaged esr nitrogen splitting constant, A,, i s given by &{(a)
+
(a’)) = 8.48(a)
+ (a‘)An’
(4)
( a ) represents the concentration of the free ion, (a’)repre-
-
8.48)(HMPA)/(An‘ - A,)(MMPA’)
(5)
A,’ is estimated from the extrapolation of a plot of A, us. the concentration of the proton donor to infinite concentration. The slope of this plot falls off rapidly at higher concentrations of alcohol (Figure 4 ) . For the case of methanol, this yields a value of 14 G for AI,’. The values used for A,’ are given in Table I for the alcohols and for water. For all cases the observed splitting, A,, is much closer to 8.48 than to A,’. This means that the~concentrat~on of the hydrogen bonded HMPA is much greater than that for .the hydrogen bonded ion. This i s an expected result, since HMPA forms relatively strong hydrogen bonds,ll and nitrobenzene neutral molecule is a poor proton aeceptor.12 We expect the anion radical of PhNO2 to be a better proton acceptor than the neutral molecule. Since (HMPA’) is much larger than (a’), we can use the eoncentrat,ion of proton donor for the (HMPA’) term in eq 5 . Using the technique described above, the equilibrium constants were determined for the PhN02-HMPA-Li system with added methanol, tert-butyl alcohol, and water. The equilibrium constants determined in this manner gave quantitative agreement for alcohol concentrations between 0.05 and 0.5 M . The equilibrium constants are given in Table I together with the standard deviation determined from six measurements. The enthalpy of the reaction consists of four predominant terms: the enthalpy of hydrogen bonding to the HMPA, the enthalpy of hydrogen bonding to the free PhNOz anion radical, the heat of solution of the products, and the heat of solution of the reactants. ~ x p e r ~ ~ e n t a l ~ y the AW can be determined from a plot of In {(A, 8.48)/(An‘ - A,)} us. l / R T . This plot should yidd a straight line if the concentration of the HMPA’ does not vary appreciably. This condition is met since at all t,emperatures (HMPA’) is much greater than (a’). All, of the systems studied do yield linear plots (Figure 5 ) . ‘The enthalpies for the various systems are given id Tab1.e 1. The enthalpy for all three systems studied are positive and between 0 and 2 kcal/mol. This indicates that HMPA i s a much better hydrogen bond acceptor than the anion radical of PhNQ2. Since PhNO2 is known to be a very weak hydrogen bond acceptor, this small that the anion radical of PhN02, free from counterion in( I t ) T. Olsen, Acta Chem. Scand., 24,3081 (1970). (12) W. F. Baitinger, P. v. R . Schleyer, T. S. S . R . Murty, and i. Robinson, Tetrahedron, 1635 (1964). The Journal of Physical Chemistry, Vol. 77, No. 8 , 7973
Gerald R. Stevenson and Hector Hidalgo
1030
-3.1
-3.2
21 2
i 4
-3.3
0
50
100
Percent Methanol
0
-3.4
Variation OF the nitrogen coupling constant with added methanol for the timcs-averaged species. The anion radical was generated in HMPA by electrolytic reduction using 0.1 M tetran-butylammonlum perchlorate as a counterion. Figure 4.
~
0.0016
0.0017
0.0018
1/R!
Modified vant' Hoff plot for the system PhNOa-HMPALi containing 0.38 M methanol.
Figure 5.
teractions, is a much stronger hydrogen bond acceptor. The enthalpy for the system containing tert-butyl alcohol as a proton donor is too small to yield measurable coupling constant differences at different temperatures. For this system we ci3n only say that AN" is greater than zero and iess than one. The entropies have been calculated from the Eyring equation for the three systems and are given in Table I. The signs of A s g i are positive, indicating that there is a reorganization of the solvent in going from the free ion to the hydrogen bonded ion. The solvent is more ordered around the free ion than it is around the hydrogen bonded ion. This is as expected, since there is more dispersion of charge in the hydrogen bonded ion than there is in the free ion. Solutions of PhN& in HMPA containing 0.1-1 M methanol also give line width alternation when reduced electrolytically (tetra-n-butylammonium perchlorate), but the effect is not as, strong as for the lithium reductions. There is some ion pairing in these systems with the counkerion as evidenced by the fact that the A , is a function of the counterion This weak ion pairing re-
The Journal of Physical Chemistry, Vol. 77, No. 8, 1973
duces the ability of the NO2 group to act as a hydrogen bond acceptor.
Experimental Section Nitrobenzene (Eastman Organic Chemicals) was vacuum distilled prior to use. Hexamethylphosphoramide was vacuum distilled from calcium hydride and stored over molecular sieves 4A. The HMPA was distilled from the solvated electron under high vacuum mm pressure) directly into the reaction vessel. Weighed portions of proton donor were added directly to the anion radical solutions through a break seal. The entire system above the reaction vessel was then warmed to ensure that all of the proton donor had entered the reaction vessel. The esr spectra were recorded on an X-band Varian E-3 esr spectrometer. The temperature was controlled with Varian V-4557 variable-temperature controller. Acknowledgments. We are grateful to Research Corporation for the support of this work.
