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A Peroxy Isomer of NOP (2)0.Kratky and G. Porod, Recl. Trav. Chim. Pays-Bas, 88, 1106 (1949). (3)P. E. Rouse, J. Chem. Phys., 21, 1272 (1953). (4)B. H. Zimm, J. Chem. Phys., 24,269 (1956). (5)A recent discussion of TSM and CM geometries applied to DNA may be found in ref 26. (6)H. Triebel, K. E. Reinert, and J. Strassburger, Blopolymers, 10, 2619 (1971)., (7)J. Garcia de la Torre and A. Horta, to be submitted for publication. (8)0. Kratky, Prog. Biophys. Mol. Biol., 13, 107 (1963). (9) V. Luzzati, A. Nicolaieff, and F. Masson, J. Mol. Biol., 3, 185 (1961). (10)V. Luzzati, D. Luzzati, and F. Masson, J. Mol. Biol., 5,365 (1962). (11)V. Luzzati, F. Masson, A. Mathis, and P. Saludjian, Biopolymers, 5, 491 (1968. (12)S.Bram and W. W. Beeman, J. Mol. Biol., 55, 31 1 (1971). (13)R. G. Kirste and R. C. Oberthur. Makromol. Chem., 127,301 (1969). (14)J. G. Kirkwood, Recl. Trav. Chim. Pays-Bas, 68, 649 (1949). (15)J. G. Kirkwood, J. Polym. Scl., 12,l(1954). (16)R. Zwanzig, J. Keifer, andG. H. Weiss, Proc. Nafl. Acad. Scl. U.S., 60,381 (1968). (17)H. Yamakawa andG. Tanaka, J. Chem. Phys., 57, 1537 (1972). (18)H. Yamakawa and J. Yamaki, J. Chem. Phys,, 57,1542(1972). (19)H. Yamakawa and J. Yamaki, J. Chem. Phys., 58,2049 (1973). (20)Y. Muss, J. Chambron, M. Daune, and H. Benoit, J. Mol. Biol., 27,579 ( 1967). (21)H. Yamakawa and M. Fujii, Macromolecules, 7,649 (1974). (22)H. Yamakawa and M. Fujii, Macromolecules, 7, 128 (1974). (23)C. W. Schmid, F. P. Rinehart, and J. E. Hearst, Biopolymers, I O , 883 (1971)., (24)J. Garcia de la Torre, J. J. Freire, and A. Horta, Biopo/ymers, 14, 1327 (1975). (25)J. B. Hays, M. E. Magar, and B. H. Zimm, Biopolymers, 8,531 (1969). (26)H. Yamakawa and M. Fujii, Macromolecules, 6,407 (1973). (27)S.Arnott, S.D. Dover, and A. J. Wonacott, Acta Crystallogr., Sect. 6,25, 2192 (1969). (28)S.Arnott, Prog. Biophys. Mol. Biol., 22, 179 (1971). (29)H. P. Hanson, F. Herman, J. D. Lea, and S.Skillman, Acta Crystallogr., 17, 1040 (1964).
(30)A. Guinier, Ann. Phys., 12, 161 (1930). (31)J. Garcia de la Torre, Ph.D. Thesis, Madrid, 1975. (32)P. Mittelbach and G. Porod, Acta Phys. Austriaca, 14,405 (1961). (33)W. Burchard and K. Kajiwara, Proc. R. SOC. London, Ser. A, 316, 185 (1970). (34)R. Koyama, J. Phys. SOC.Jpn., 36, 1409 (1974). (35)G. Porod, Acta Phys. Austriaca, 2,255 (1948). (36)0.Kratky and G. Porod, Acta Phys. Austriaca, 2, 133 (1948). (37)J. Rlseman and J. G. Kirkwood, J. Chem. Phys., 18,512 (1950). (38)J. Riseman and J. G. Kirkwood, “Rheology”, Academic Press, New York, N.Y., 1956,p 495. (39)J. M. Burgers, “Second Report on Viscosity and Plasticity of the Amsterdam Academy of Sciences”, Nordemann, New York, N.Y., 1938,Chapter 3. (40)S. Broersma, J. Chem. Phys., 32,1632 (1960). (41)V. A. Bloomfield, W. 0. Dalton, and K. E. Van Holde, Biopolymers, 5,135 11987),. (42)C. M. Tchen, J. Appl. Phys., 25,463 (1954). (43)P. Doty, B. McGiII, and S. A. Rice, Proc. Natl. Acad. Sci. U.S., 44,432 (1958). (44)K. Kawade and I. Watanabe, Biochim. Biophys. Acta, 19,513 (1956). (45)K. Is0 and I. Watanabe, Nippon KagFku Zashi, 78,1268 (1957). (46)L. Pivec, S.Zadrazil, J. Sponar, and Z. Sormova, Collect. Czech. Chem. Commun., 30,3928 (1965). (47)G. Cohen and H. Eisenberg, Biopolymers, 8, 45 (1969). (48)H. Eisenberg, Biopolymers, 8,545 (1969). (49)A. Prunell and G. Bernardi, J. Biol. Chem.. 248,3433 (1973). (50)M. T. Record, C. P. Woodbury, and R . B. Inman, Biopolymers, 14,393 (1975). (51)J. Eigner and P. Doty, J. Mol. Biol., 12,549 (1965). (52)J. E. Hearst and W. H. Stockmayer, J. Chem. Phys., 37, 1425 (1962). (53)K. C. Holmes and D. M. Blow, “The Use of X-Ray Diffraction in the Study of Proteins and Nucleic Acid Structures”, Wiley, New York, N.Y., 1965. (54)P. Sharp and V. A. Bloomfield, J. Chem. Phys., 48, 2149 (1968). (55)V. A. Bloomfield, Macromol. Rev., 3,225 (1968). (56)M. B. Tunis and J. E. Hearst, Biopolymers, 8, 1325 (1968). (57)H. B. Gray, V. A. Bloomfield, and J. E. Hearst, Biopolymers, 14, 393 (1975). ~
A Peroxy Isomer of Nitrogen Dioxide1 H. F. Schaefer, 111, University of California at Berkeley, Department o f Chemistry, Berkeley, California 94720
C.
F. Bender, and J. H. Richardson’
University of California, Lawrence Livermore Laboratory, Livermore, California 94550 (Received April 9, 1976) Publication costs assisted by the University of California
Ab initio SCF calculations on the relative stability of a peroxy isomer of NO2, compared with the normal CzU isomer, coupled with the experimental bond strength in the normal isomer suggest that a peroxy isomer is stable. Implications to photodetachment experiments are discussed. Recent photodetachment erperiments have demonstrated the existence of an anomalous nitrite ion.2 Photodetachment of an electron from NO2- was observed to occur at energies significantly less than the electron affinity of NO2. Three possible explanations were suggested for this unexpected observation: (1) The “normal” CzUisomer of NO2- (2 lAl) might have been vibrationally excited. Photodetachment at wavelengths greater than the adiabatic electron affinity would then be expected, similar to the appearance of hot bands in conventional absorption spectroscopy. This explanation was eliminated because there is no evidence to suggest that a sufficiently high vibrational temperature (5000 K) was present in either of the apparatus used to generate nitrite ions.
(2) An excited metastable triplet state of NOa-. The lowest reported3 triplet state of NO2- (3B1) is a t best marginally stable with respect to autodetachment to NO2 + e-.2a Furthermore, it is unlikely that such a metastable species could be produced copiously from a hot cathode plasma ion sourcezb or be contained in an ion cyclotron resonance cell for several seconds.2a (3) A peroxy isomer of NOz-. Such an ion would be structurely similar and isoelectronic to NOF.4 Furthermore, its photodetachment spectrum might be expected to be similar to a perturbed oxygen ion, which is consistent with the threshold and cross section near threshold. Finally, an anomalous, energetic form of NO2- was observed in cluster ion reactions of NO; similar reactions with NO2 resulted in the The Journal of Physical Chemistry, Vol. 80, No. 18, 1976
H. F. Schaefer, C. F. Bender, and J. H. Richardson
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TABLE I: Geometries and Energies for Civ and Peroxy Isomers of NO2- and NOza
Peroxy ‘Property
CZ”
r (N-0) r (0-0) 6 , deg E , hartrees E , eV
1.20 135 -203.9474 0.0
NOz-
I 2A‘
I 2A”
1.16 1.90 122 -203.8820 1.78
1.23 1.49 115 -203.8563 2.48
CZ” 1.264 117.0 -204.0336 0.0
Peroxy 1.245 1.493 118.5 -203.9161 3.20
Bond distances are in 8.Data for NO2- is taken from ref 6. formation of a peroxy nitrate ion: and it is suggested that the anomalous form of NO2- was also a peroxy isomer. To further substantiate the existence of a peroxy isomer of NQ2-, ab initio SCF calculations were performed6 (the results of these calculations are reproduced in Table I). A peroxy form of NO’- was found to have a well-defined minimum in the potential energy surface. Symmetry restrictions and energy limitations prohibit interconversion of the peroxy isomer to the normal Czuisomer via a ring intermediate. This prediction is also consistent with the experimental observation of two distinct, apparently noninterconverting ions of mle 46. Photodetachment of an electron generally does not result in subsequent dissociation; hence, photodetachment of the peroxy isomer of NO’- might be expected to result in formation of a peroxy isomer of NO’. It has been suggested2that in this case the photodetachment process results in dissociation, i.e.
Consequently, to.investigate this hypothesis, a theoretical study of the relative stability of the peroxy isomer of NO2 was undertaken. Further motivation was provided by the paramount importance of the presence and chemistry of nitrogen oxides in the atmosphere. While obviously not a major, longlived species, the presence of a peroxy isomer of NO2 in trace amounts or as an energetic intermediate would have major significance in atmospheric chemical reaction mechanisms and m ~ d e l i n g This . ~ paper presents the results of ab initio self-consistent-field calculations of a peroxy isomer of NO2. Atom-optimized primitive Gaussian basis sets8 of size (9s 5p) were centered on the N and 0 nuclei. This primitive basis set was contracted to (4sPp), providing a double {quality basis set.9 The peroxy isomer of NO2 has only a plane of symmetry (point group Cs), The following configurations were investigated: 1a’22af~3a’24a’25a’26a’27a’~8a’~1a’’29a’22a’f21Oa’ (I 2A’) 1a’22af~~a’2~a~25a”6a’27a‘28a’zla’’29a’z2a’’23a’’ (I11 2A’’)
1a’22a’z3a’24a’~~a’~6a~z7~fz8a’~la’’z9a’z2a’’lOa’lla’ (IV”) This first configuration, I ‘A’, corresponding to the most stable bound state, was observed to have a rather shallow minimum with respect to the 0-0 bond length (the N-0 bond length is very similar to that of NO (1.151&lo).The depth of this well, ca. 0.008 hartree or 1700 cm-l, is insufficient to prevent the molecule from dissociating at room temperature. However, this well depth can be estimated more reliably from semiempirical considerations (vida infra). The Journal of Physical Chemistry, Vol. 80, No. 18, 1976
The second configuration studied, I1 ‘A’’, corresponds to promotion of an electron from the ?r to u framework. This configuration appears very tightly bound, presumably because it cannot dissociate in the single configuration approximation to ground state NO 0. It is interesting to note that the geometry of this configuration is very similar to that of the peroxy ion (Table I).The last two configurations studied did not exhibit any minimum with respect to the 0-0 bond length; each had an angular minimum corresponding to a linear geometry. Simple SCF calculations do not reliably predict dissociation energies due to the restriction of having doubly occupied orbitals.ll The results for the Czu isomer are also included in Table I, and agree with similar calculations12done a t the experimental geometry. Other basis sets and contraction schemes, including the use of polarization functions (d orbitals), have also been used and compared.12J3 The double { basis set used here was found to yield reliable predictions for equilibrium parameters and simple electronic expectation values, although dissociation energies are poorly estimated. However, by combining the energy difference (1.78 eV) between the Czu isomer of NO2 and the peroxy isomer (I 2A’) with the accepted14 ON-0 bond dissociation energy for the Czu isomer (3.11 eV), the NO-0 dissociation energy for the peroxy isomer can be estimated. Our resulting estimate for the NO-0 dissociation energy is 1.33 eV. This estimate is likely to be accurate because of the electronic similarity between the two isomers13 and the success of SCF calculations in predicting geometries and relative differences near potential energy minima. Finally, it is interesting to note that even in the SCF approximation NO2 is predicted to have a positive electron affinity. The remaining question is whether photodetachment of the peroxy isomer of NOz- is accompanied by concomitant dissociation of the neutral fragment. This question cannot be irrefutably answered without knowing the complete energy surface. For comparison, the vertical detachment energy of the CZu isomer exceeds the adiabatic electron affinity by only 0.4 eV, and there is a substantial geometry change in that process also.2a Our results certainly suggest the possibility that N O 0 is a stable species and may play a role in atmospheric chemistry. Further calculations are being performed with a more extensive basis set (d orbitals) and the inclusion of significant configuration interaction (CI). These results will be presented in a subsequent article.
+
References a n d Notes (1) This work wag performed under the auspices of the United States Energy Research and Development Administration under Contract No. W-7405Eng-48. (2) (a)J. H. Richardson, L. M. Stephenson, and J. I. Brauman, Chem. Phys. Lett.,
2037
Strong Complex Formation by HOn Radical 25, 318 (1974); (b) E. Herbst, T. A. Patterson, and W. C. Lineberger, J. Chem. Phys., 61, 1300 (1974). (3) R. M. Hochstrasser and A. P. Marchetti, J. Chem. Phys., 50, 1727 (1969). (4)R. R. Sardzewski and W. E. Fox, Jr., J. Am. Chem. SOC., 96, 304 (1974). (5) (a) N. G. Adams, D. K. Bohme, D. B. Dunkin, F. C. Fehsenfeld, and E A . Ferguson, J. Chem. Phys., 52,3133 (1970); (b) E. E. Ferguson, D. B. Dunkin, and F. C. Fehsenfeld, ibid., 57, 1459 (1972). (6) P.K. Pearson, H. F. Schaefer, ill, J. H. Richardson, L. M. Stephenson, and J. I. Erauman, J. Am. Chem. SOC., 96, 6778 (1974). , (7) For an example of the importance of nitrogen oxides in the atmosphere,
see G. Brasseur and M. Nicoiet, Planet. Space Sci., 21, 939 (1973). (8) S.Huzinaga, J. Chem. Phys., 42, 1293 (1965). (9) T. H. Dunning, J. Chem. Phys., 53, 2823 (1970). (IO) G. Herzberg, “Spectra of Diatomic Molecules”, Van Nostrand-Reinhold, New York, N.Y., 1950. (11) H. F. Schaefer, “The Electronic Structure of Atoms and Molecules”, Addison-Wesley, Reading, Mass., 1972. (12) S. Rothenberg and H. F. Schaefer, Mol. Phys., 21, 317 (1971). (13) G. D. Gillispie, A. U. Khan, A. C. Wahl, R. P. Hosteny, and M. Krauss, J. Chem. Phys., 63,3425 (1975). (14) G. Herzberg, “Electronic Spectra of Polyatomic Molecules”, Van Nostrand-Reinhold, New York, N.Y., 1966.
Theoretical Calculation of Strong Complex Formation by the H02 Radical: H02*H20and H02*NHsia E. J. Hamilton, Jr.,+lb and C. A. Naleway Chemistry Division, Argonne National Laboratory, Argonne, Illinois 60439 (Received February 20, 1976) Publication costs assisted by Argonne National Laboratory
Exploratory ab initio calculations using a minimal basis set support the existence of the H02*H20 and HO2. NH3 complexes, as proposed earlier by experiment. For the minimum energy configurations having HO2 as the H donor within a linear H-bond structure, electronic stabilization energies of 9.1 and 12.0 kcal/mol are calculated for HOyH20 and HOz-NHs, respectively, compared with 5.3 kcal/mol for ( H 2 0 ) ~Values . of AH” and AS’ for the complex formations are estimated and found consistent with available experimental data. A plausible model is proposed to explain the reactivity of these complexes. These calculations indicate that in the troposphere a significant fraction of the HO2 is complexed with H2O.
Introduction Recent experiments in this laboratory2 have revealed that the observed rate of the atmospherically i m p ~ r t a n tself~,~ reaction of HO2 in the gas phase is increased by up to a factor of -2.5 (at -298 K) in the presence of a few Torr of H2O or NH3. Various considerations have led to the inference that this phenomenon is due to the formation of 1:l complexes HO2 + H2O e HO2 * H20
(1)
HO2 + NH3 e HO2 NH3 (2) which are more reactive than uncomplexed HO2 toward a second uncomplexed HO2 r a d i ~ a lBased . ~ on a kinetic model for this system, a preliminary equilibrium constant for (2) of K p 95 (based on a Po = 1atm standard state) at -298 K has been derived from the data.5 In connection with this proposed explanation of the data, ab initio calculations on hydrogen-bonded HOyH20 and HOyNH3 complexes have been carried out and equilibrium thermodynamic parameters for (1)and (2) estimated.
-
Ab Initio Calculations In the Hartree-Fock calculations reported here, a minimal Gaussian basis set [3s lp/ls] expansion was employed. The primitive basis consisted of a (10s 5p) expansion on the oxygen and nitrogen sites, while a (5s) expansion was used on the hydrogens. This primitive basis set is of essentially atomic double-l quality. The contraction scheme, composed from a
concatenation of Whitten’$ s-type orbitals with Huzinaga’s7 p-orbital set, has been outlined in detail earlier.x An effective Slater exponent of { = 1.2 for all hydrogen orbitals was found to be near optimum for each monomer. This basis set is equivalent to that employed in previous studies of hydrogen bonding in the molecular series HF, HzO, and N H S . This ~ basis set expansion yields a stabilization energy in good agreement with experiment (see next section) for the closedshell (H2O)z sy~tem,~JO although the approximate character of this basis expansion dictates that this agreement is partially due to a cancellation of errors. It has been the authors’ premise that agreement such as that found earlier could be extended to complexes between an uncharged radical and a neutral polar molecule wherein the open shell system is primarily removed from the region of H bonding. Essentially experimental monomer geometries were held rigid for all calculations. The geometry used for the water monomer has an 0-H bond length of 0.957 A and an H-0-H angle of 104.52’. The ammonia geometry has an N-H bond length of 1.012 A and an H-N-H angle of 106.7’. The HOz geometryll has an 0-0 bond length of 1.34 A, an 0-H bond length of 0.96 A, and an 0-0-H angle of 100.0’. Self-consistent field molecular orbital calculations were performed following Roothaan’s formalism for open-shell systems1* using the MOLE LCAO-MO program package. The SCF energies for the H20, “3, and HO2 monomers were calculated as -75.976 221 3, -56.142 436 9, and -150.101 432 0 hartrees, respectively. The Journal of Physical Chemistry, Vol. 80,No. 18, 1976
