1680
k' = C(
s>,,-'
exp(
J. Phys. Chem. 1984,88, 1680-1681
12 E ) kBT
(1 3b)
where
AE
= AV,, = V,,- Vi
(14)
It should be noted that eq 13a and 13b are similar to the Arrhenius empirical relation for k and k'except that a factor of ' I 2appears in the exponents for k and k', if -hE/kBT is set in the exponent for K, furthermore, the preexponential factors for k and k'are the same. Although the dynamics is fully accounted for in the treatment, k, and k & are determined from equilibrium properties. Once k, and k', are known, the detailed dynamic behavior of a,(t) can be obtained by solving eq 1. The-expressions for the Laplace transform of a,(t) with respect to t6 and a,(t) as an
explicit function o f t in the special case of k l = k2 ... = k,l and were obtained previously. It should be noted k', = k i = ... = k',' that the Smoluchowski equation with an arbitrary force fieldflx) can be put into a set of rate equations arising from chemical reaction in I. In obtaining k, k', and K , we did not need elaborate calculations2q3and the assumption of a transition state.24 The previous studies were concerned with the calculation of rate constants without directly connecting them to a specific reaction mechanism which immediately enables us to write the rate equations. Hence the calculated rate constants were too general to relate them explicitly to phenomenological rate constants appearing in the rate equations. The present treatment is free from this difficulty. (8) A. Morita, Chem. Phys. Lett., 91, 197 (1982).
Hydrogen Bonding In Aqueous Solutions of D-Ribose and 2-Deoxy-D-ribose J. H. Stern* and P. M. Hubler Department of Chemistry, California State University, Long Beach, California 90840 (Received: January 27, 1984)
Enthalpies of transfer of D-ribose and 2-deoxy-~-ribosefrom pure water to 3 m aqueous solutions of ethanol and urea were determined via enthalpies of solution at 25 O C . The enthalpies of transfer to aqueous ethanol were positive, while those to aqueous urea were negative, indicating their opposite effects on the solvent properties of water. Significant differences in enthalpies were observed for the two monosaccharide solutes, which differ in only one OH group and are important constituents of the nucleic acids. The results are explained by hydrogen-bonding differences between the solutes and the solvent media.
Introduction D-Ribose and 2-deoxy-~-riboseare constituents of the nucleic acids R N A and DNA, respectively, and for this reason these monosaccharides are of greatest importance biochemically. They are very similar, differing by only one OH group. Their interactions with the solvent medium are very important, since they may have a direct bearing on the stabilities of the native configurations of R N A and DNA in aqueous solutions. A number of paters deal with the physical chemistry of aqueous monosaccharide~.'-~ There appears to be no consensus on their mechanism of hydration, and it has been suggested, for example, that hydrogen bonding between glucose and water is stronger and more extensive than that between water molecule^,^ with ribose being less hydrated than g l ~ c o s e .In ~ any case, there is a lack of experimental information on such complicated systems. This contribution describes the enthalpies of transfer, AH, of ribose and deoxyribose from pure water to 3 m aqueous solutions of ethanol and urea at 25 O C . Urea and ethanol have drastically different effects on the solvent properties of and their interaction with a variety of aqueous electrolytes, smaller nonelectrolytes, and macromoleculesg have received considerable attention. Enthalpies of transfer of electrolyte and nonelectrolyte solutes from pure water to mixed aqueous solvents have proven (1) H. Hoiland and H. Holvik, J . Solution Chem. 7, 587 (1978). (2) F. Franks, J. R. Ravenhill, and D. S. Reid, J . Solution Chem., 1, 1 (1972). (3) A. Suggett in "Water, A Comprehensive Treatise", Vol. 4, F. Franks Ed., Pleunum Press, New York, 1975, Chapter 6. (4) J. H. Stern and L. R. Beeninga, J . Solution Chem., 5 , 617 (1976). (5) J. B. Taylor and J. S . Rowlinson, Trans. Faraday SOC.,51, 1183 (1955). (6) J. H. Stern and J. D. Kulluk, J . Phys. Chem., 73, 2795 (1969). (7) F. Franks and D. J. G. Ives, Q. Reu. Chem. SOC.,20, 1 (1966). (8) F. Franks and R. G. Bates in "Hydrogen-Bonded Solvent Systems", A. K. Covington and P. Jones, Eds., Taylor and Francis, London, 1968. (9) See D. Eagland, ref 3, Chapter 5.
TABLE I: Enthalpies of Solution and Transfer at 25 "C --____
fl, solute Dribose
2-deoxy-D-ribose
a
system
runs
H2O 3 m ethanol 3murea H,O 3 methanol 3mured
6 5 6 7 4 6
N o a n d AH, cal/mol
cal/ mol
*
3184 l g a ( A H " ) 4 2 3 8 ~20 1054 2153 A 9 -43 1 2 9 0 0 ? 10 (AH") 3765 i 1 9 86 5 2629 A 5 -271
All uncertainties dre standard deviations of the means.
to be valuable in a variety of solute-solvent interaction systems.4,6,10-1 2 Experimental Section The calorimeter and calorimetric methods have been described e1~ewhere.l~D-Ribose (Sigma) and 2-deoxy-~-ribose(Pfanstiehl) were dried in vacuo at 65 O C prior to use. Ethanol and urea were AR grade and the water was distilled and deionized. Enthalpies of transfer were determined from the difference of enthalpies of solution in the mixed solvent ( A H ) and in pure water (AHo): All = AH - AHo. The overall range of final solute concentrations was 0.0039-0.01 1 mol per 1000 of water or mixed solvent. Results and Discussion Enthalpies of solution and transfer are shown in Table I. Ethanol in water-rich solutions has a structuring effect on water' (10) A. C. Rouw and G. Somsen, J . Chem. Thermodyn., 13,67 (1981). (11) J. H. Stern and S . L. Hansen, J . Phys. Chem., 85, 3713 (1981). (12) J. H. Stern, S. J. Stoner, and G. L. Doyle, J . Solution Chem., 10,263 (1981). (13) J. H. Stern, B. L. Goeders, G. L. Withers, and S.L. Wujs, J . Chem. Eng. Data, 24, 314 (1979).
0022-365418412088-1680$01.50/0 0 . 1984 American Chemical Societv
J . Phys. Chem. 1984, 88, 1681-1690 and also competes for hydration and hydrogen-bonding sites with the transferring solute monosaccharide. Thus a net number of solute-water hydrogen bonds must be broken during transfer, with resulting endothermic values of AH. There is a significant difference of -190 cal/mol in AH between ribose and deoxyribose (1054 and 865 cal/mol). Since ribose has one more hydrogen bonding OH site than deoxyribose, more hydrogen bonds will be broken in the former upon transfer, in satisfactory agreement with the observed differences in enthalpies of transfer. For transfer to aqueous urea both values of A H are negative, and its presence in water thus appears to enhance hydrogenbonding opportunities of the solutes. In this case, going from ribose to deoxyribose increases AI? by 160 cal/mol since less hydrogen bonds will be formed by the latter. It may be noted that the absolute magnitudes of the changes in AZ7 for both solvent systems are nearly the same, ca. 0.2 kcal/mol. Thus hydrogen bonding satisfactorily explains the enthalpies of transfer and their dif-
1681
ferences in these very complicated solutions. It may be noted that interaction coefficients between aqueous urea and eight linear and one cyclic (myoinositol) polyols were recently reported,I4 based on enthalpies of dilution. The coefficients are negative and become more so with increasing molecular weight, in qualitative agreement with the present study on cyclic ribose and deoxyribose. Calculated pairwise urea-hydroxyl group contributions are also negative. Quantitative comparisons are not possible since significant differences were observed in the behavior of linear and cyclic polyols, and both studies were carried out in different concentration ranges.
Acknowledgment. The authors are grateful for financial support by the California Heart Association and the California State University at Long Beach. (14) G. Barone, V. Elia, and E. Rizzo, J . Solution Chem., 11,687 (1982).
FEATURE ARTICLE Study of Empty Electronic States by Inverse Ultraviolet Photoemissiont V. Dose Physikalisches Institut der Universitat Wurzburg, Wurzburg, West Germany (Received: October 6, 1983)
The basic concepts and current state of the rapidly developing technique of ultraviolet inverse photoemission are discussed. This new technique allows one to determine unoccupied electronic states in solids and at surfaces. The information obtained for empty electronic states is exactly analogous to that provided by ordinary photoemission for occupied states. The special merit of the method compared to other empty-state spectroscopies is its angle-resolution capability. Examples from recent work are presented and include density of state measurements, chemisorption, and band mapping. Future developments and goals are suggested.
I. Introduction Photoemission and inverse photoemission are spectroscopies based on effects which are discovered in the 19th century. In fact, the first observation of electron emission by electromagnetic radiation was reported by Hallwachs in 1888.’ About 10 years later X-rays were discovered by Rontgen at the physics department of the University of Wiirzburg.2 This 10-year delay appears as a rigid shift in the further development of photoemission and inverse photoemission. In 1905 Einstein presented his famous interpretation of the relation between the incident light frequency and the maximum photoelectron kinetic energye3 The first experimental confirmation of the inverse, that is, the minimum X-ray wavelength emitted by a sample under bombardment of monoenergetic electrons, was reported in 1915.4 Both of these experiments were milestones in the development and understanding of quantum theory. Application of both effects, photoemission and bremsstrahlung production, or in more general terms electronic transitions in solids by emission or absorption of light, had to wait for roughly 60 years. It appears now that such a delay was inevitable since a systematic development of ultrahigh-vacuum (UHV) techniques was a necessary prerequisite. During the 1960s the photoelectric effect was developed into a powerful and widely used photoemission spectroscopy (PES) for the study of occupied electronic states in solids and at surface^.^-^ The rapid growth Dedicated to Prof. K. Ulmer on the occasion of his 65th birthday.
of this field was last but not least due to the modest experimental investment necessary. A PES experiment needs an electron spectrometer and a monochromatic light source which is readily available in the form of a noble gas resonance lamp. In the course of the development of PES the necessity to vary the quantum energy was recognized quite early. The importance attributed presently to PES can hardly be better characterized than by the large investment involved in the construction of synthrotrons and storage rings as dedicated variable energy light sources for photoelectron spectroscopy. A new impetus to inverse photoemission spectroscopy occurred in the late 1970s when an essential technical step was reported! Equipment for an inverse photoemission (IPE) experiment is also modest. The source of excitation is an electron gun delivering monochromatic electrons. Radiation emitted by the sample under electron bombardment must be recorded by an energy-selectivequantum detection device. If the detection energy is fixed, we obtain the inverse of the PES resonance lamp exW. Hallwachs, Ann. Phys., 33, 301 (1888). W. C. Rontgen, Si&. Ber. Med. Phys. Ges. Wiirzburg, 137 (1895). A. Einstein, Ann. Phys., 17, 132 (1905). W. Duane and F. L. Hunt, Phys. Reu., 6, 166 (1915). M. Cardona and L. Ley, Eds., ”Photoemission in Solids, I, 11”. Springer-Verlag, West Berlin, 1978. (6) F. J. Himpsel, Adu. Phys., 32, 1 (1983). (7) N. V. Smith and F. J. Himpsel in “Handbook on Synchrotron Radiation”, E. E. Koch, Ed., North-Holland Publishing Co., 1983, pp 905-54. (8) V. Dose, Appl. Phys., 14, 117 (1977). (1) (2) (3) (4) (5)
0022-3654/84/2088-168 1%01.50/0 0 1984 American Chemical Society
