J. Phys. Chem. 1995, 99, 15822-15829
15822
Gas-Phase Positive- and Negative-Ion-Molecule Reactions in NF3 Kenzo Hiraoka,* Masayuki Nasu, and Susumu Fujimaki Faculty of Engineering, Yamanashi University, Takeda-4, Kofu 400, Japan
Shinichi Yamabe" Department of Chemistry, Nara University of Education, Takabatake-cho, Nara 630, Japan Received: March 8, 1995; In Final Form: July 3, 1995@
The positive- and negative-ion-molecule reactions in N F 3 were studied using a pulsed electron-beam mass spectrometer. It was found that the nature of bonding in the clusters F-(NF3)n is mainly electrostatic with the bond dissociation energies less than 8 kcallmol. The shell formation in F-(NF3),, with n = 2 and 4 is suggested. It was predicted that the NF2+(NF3)1 ion has two isomeric structures, i.e., F2N+-NF3 and F2NfFF2N. The former was found to be more stable than the latter. These predictions were confirmed by ab initio calculations. The lower limit of proton affinity (PA) for N F 3 was determined to be 132.4 kcallmol by measuring the temperature dependence of the rate constants for the endothermic proton-transfer reaction H'(N2)2 N F 3 NF3H' 2N2. The G2+ theory gives PA = 139.6 kcallmol.
-.
+
+
1. Introduction The demand for nitrogen trifluoride (NF3) as an etchant in plasma technology is rapidly growing for fabrication of semiconductors. N F 3 is a familiar reagent gas in the gas-phase ionmolecule study for producing F- ions by a dissociative electron capture reaction. Contrary to the wide applications of N F 3 to plasma chemistry and to gas-phase negative-ion-molecule reactions, the ion-molecule reactions in N F 3 have not been examined so extensively. The proton affinity (PA) of N F 3 ( 2132.4 kcdmol, present work) is much lower than that of NH3 (204.0 kcallmol). This is due to the fact that both N and F atoms are highly electronegative (electron-withdrawing) and the positive charge in H+(NF3) is much less dispersed than that in N b + . Fisher and McMahon' suggested that the more favorable protonation site of N F 3 is the F atom rather than the N atom. In order to distinguish between the nitrogen-protonated H+NF3 and the fluorine-protonated F2Nf-FH, they examined the reactions of Hf(NF3) with n-bases, C02 and HCl. They observed reactions 1 and 2, in which HF is displaced by C02 and HCl. H+(NF3) f CO, H'(NF3)
-
+ HC1-
+ HF NF2+HC1+ HF NF2'C02
Abstract published in Aduance ACS Abstracrs, October 1, 1995.
-
(1)
(2)
This observation suggests that H'(NF3) is basically a nitrenium ion (NF2+) weakly bound to a HF molecule and that reactions 1 and 2 are nucleophilic displacement reactions (SN~).The relative stabilities of H+NF3 and F2Nf-FH were examined using ab initio calculations, which confirmed slightly greater stability of the fluorine-protonated form. Additional structural evidence for NF*+-FH was obtained from a CID-MIKES experiment for H+(NF3) ion. The CID-MIKES spectrum revealed a dominant loss of HF from H+(NF3). This strongly suggests a fluorine-protonated N F 3 , i.e., F2N+-FH. Schwarz and his co-workers examined the energy surface of [N, F3, H]+ ions by means of ab initio MO studies and massspectrometric techniques.2 The global minimum corresponds to the fluorine-protonated isomer F2N+-FH, in agreement with @
the results of Fisher and McMahon.' The ion-dipole complex NF2+-FH was found to be 6.4 kcdmol more stable than the nitrogen-protonated form H'NF3 by the G1 theory. They found that the barrier for the intramolecular isomerization reaction NF2+-FH HfNF3 is significantly large (52.6 kcallmol), thus preventing facile isomerization. The isomer F2N+-FH has a low-energy dissociation channel to produce N F 2 + and HF, with a heat of reaction of +(NF3),,(reaction 4).
IOOO/T(K)
Figure 5. Van't Hoff plots for the clustering reaction NF~+(NF~),-I
+ NF3
d io4 z z a
NF3+(NF3), (reaction 5 ) .
6 lo3
[L
2
IO'
v)
52
IO'
0
0
0.5
I
1.5
2
2.5
'0 354 4 K
TIME "
0
0.5
I 1.5 TIME (ins)
2
2.5
Figure 4. Temporal profiles of NF2+ and N F ~ + ( N F ~ )Upper I. figure: ion source temperature = 326.4 K; pressure of Nz = 2.93 Torr; pressure of NF3 = 36.9 mTorr; integration time for NF2+ = 40 s; integration time for N F ~ + ( N F ~=) I60 s; 2 keV ionizing electron pulse = 20 ps. Lower figure: ion source temperature = 354.4 K; pressure of N2 = 2.97 Torr; pressure of NF3 = 34.8 mTorr; integration time for NF2+ = 40 s; integration time for N F ~ + ( N F ~=) I80 s; 2 keV ionizing electron pulse = 20 pus. Repetition rate of electron pulse = 30 s-l.
creased below 340 K. This effect is evident from the ion intensity changes in Figure 4. At 354 K, the equilibrium between N F 2 + and NF2+(NF3)1 is seen to be established right after the electron pulse. In contrast, it takes about 1.5 ms to establish the equilibrium at 326 K. Generally, ion-molecule clustering reactions proceed faster at lower temperatures because the lifetime of an intermediate ion-molecule complex becomes longer and the probability for the collisional stabilization by the third body increases at lower temperature. The observed positive temperature dependence of the rate of the reaction N F 2 + NF3 NF2+(NF3)1 is against the general trend. Fisher and McMahon' and Schwarz et aL2 predicted that the F-site-protonated NF3H+ is basically an ion-dipole complex NF,+-FH and that the bond dissociation energy between N F 2 + and FH is 14.6 kcaVmoL2 This rather high value arises from the large dipole moment (1.9 D) of HF. The experimentally determined bond dissociation energy (17.7 kcdmol) of NF2+(NF3)1 is greater than that of NFz+-FH. In view of the small dipole moment of NF3 (0.2 D), the observed bond dissociation energy seems to be too large if the complex NF2+(NF3)1 is an ion-dipole complex, F2N+-F6-F2N6+. Here, we may propose that the complex has a structure of F*N+N F 3 with a covalent character. This proposal will be confirmed theoretically in the next section. The van't Hoff plots for reaction 5 are displayed in Figure 5. The enthalpy and entropy changes obtained from Figure 5 are summarized in Table 1. The large gap in the van? Hoff plots between n = 1 and 2 is found in Figure 5, as in the case for
+
-
(ms)
Figure 6. Ions observed after the electron pulse in 2.58 Torr of N2, 0.68 Torr of H2, and 2.06 x Torr of NF3 at 372.1 K. Duration of 2 keV electron pulse is 20 ps.
-
reaction 4. This corresponds to the large falloff of binding energies, 7.8 3.4 kcal/mol in Table 1. 3.3. Evidence for Two Isomers of FtN+(NF3). We will describe here experimental results which indicate that F2Nf(NF3) can exist in two isomeric forms. These results also have some bearing on the proton affinity of NF3. Figure 6 represents temporal profiles of ions observed in 2.58 Torr of N2, 0.68 TOITof Hz, and 2.06 x Torr of N F 3 at 372 K after the electron pulse. The temporal ion profiles in Figure 6 indicate the following reaction sequence.
+ NF3 - Hf(NF3) + N, HfN2 + N, - Hf(N2)2 H+(N2)2+ NF3 - Hf(NF3) + 2N2 Hf(NF3) - NF2++ HF F2Nf-FH + NF3 - F2N+-FNF2 + HF F,N+-FNF, - NF,+ + NF, H+N2
(6) (7)
(8) (9) (10) (1 1)
The initial rapid growth of H+(NF3) is mainly due to protontransfer reaction 6. The proton affinity of NZ (118.2 kcal/mol) is much lower than that of NF3, and reaction 6 is highly exothermic. In Figure 6, H+N2 is seen to be quickly converted to H+(N2)2by reaction 7. The slow decay of H+(N2)2 must be due to reaction 8. The rapid decay of H+(NF3) corresponds to the growth of N F 2 + (Le., reaction 9) and NF2+(NF3). The initial growth of NF2+(NF3) is followed by the subsequent decomposition (reaction 11) to produce N F 2 + .
IodMolecule Reactions in N F 3
J. Phys. Chem., Vol. 99, No. 43, 1995 15825
- 5000,
-0
IO
20
30
40
PRESSURE O F NF3 ( m T o r r ) Figure 7. NF3 pressure dependence on rate constants for the unimoNF3. Ion source lecular decomposition reaction NFZ+(w3) NF2' temperature = 392 K in reaction 11.
+
\
1
2
3
4
5
6
7
IOOO/T(K)
a 3
i,,,,*',,l
'O2.0
2.5
30
IOOO/T(K)
Figure 8. Arrhenius plot for the rate constants of the unimolecular decomposition reaction NF2+(NF3)1 zNF2+
+ NF3 (reaction 11).
In the previous section, the thermochemical stability of NF2+(NF3)l has been determined. The NF2+(NF3)1 ion was proposed to have the structure F2N+*"NF3. From the van't Hoff plot for reaction 4 with n = 1, the ion intensity ratio [NF2+(NF3)1]/ [NF2+]can be calculated to be 7 x under the experimental conditions in Figure 6. As shown in Figure 6, the relative intensity of NF2+(NF3)1 compared to that of N F 2 + is much greater than the estimated value. This clearly indicates that the species NF2+(NF3)1 observed in Figure 6 is not the product from the clustering reaction of NF2+ with N F 3 (i.e., reaction 4 with n = 1). Then, what is the precursor for the formation of this ion? In Figure 6, the only possible precursor considered for this ion is H+(NF3). In the previous section, it was described that protonated NF3 has the structure of F2N+ interacting with FH weakly. The strong appearance of F2Nf(NF3)1 in Figure 6 may be explicable if one assumes that the FH molecule in F2N+...FH is substituted by the N F 3 molecule to form F2F+*..FF2N, Le., reaction 10. In reaction 10, the structure of the product ion NF2+(NF3), is assumed to be the normal iondipole complex F2N+-."FF2N6+. The occurrence of such a substitution reaction seems to be reasonable, since Fisher and McMahon' observed similar substitution reactions 1 and 2. Figure 7 represents the N F 3 pressure dependence on the rate constants for reaction 11 at 392 K, which are evaluated by plotting the logarithm intensity of NF2+(NF3)l vs time. The rate constants are independent of the N F 3 pressure, indicating that reaction 11 is unimolecular. It was found that the decay rate became faster with an increase of temperature. Figure 8 displays the Arrhenius-type plot of the rate constants for reaction 11. From the slope of the straight line, the activation energy for the unimolecular decomposition reaction 11 can be estimated to be 8.5 kcal/mol. If the ion N F 2 + ( N F 3 ) l observed in Figure 6 has the structure of an ion-dipole interaction, F2N+--P-F2N6+, the reaction N F 2 + F3N -c N F 2 + ( N F 3 ) 1 should not have any energy barrier, and thus the activation energy observed for reaction 11 may correspond to the bond dissociation energy (D) of F2Nf***FF2N.That is, D[F2N+.-FF2N] = 8.5 kcallmol. This
+
Figure 9. Arrhenius plot for the rate constants of the proton-transfer H+(NF3) 2N2 (reaction 8). reaction H+(N2)2 NF3
+
+
value is much smaller than D[F2N+.-*NF3] = 17.7 kcal/mol in Table 1. The former value seems reasonable considering that the dipole moment of N F 3 is only 0.2 D. The slow establishment of equilibrium for reaction 4 (n = 1) at lower temperatures (Figure 4) may be interpreted as follows. At higher temperatures, the thermochemically less stable isomer F2N+-FNF2 cannot exist md the equilibria between N F 2 + and F2N+-NF3 are established quickly (e.g., Figure 4,354 K). With a decrease of temperature, the less stable isomer F2N+-FNF2 starts to be formed. Because the isomer F2N+-FNF2 is basically an ion-dipole complex, this may be the dominant initial product (see the theoretical section). This kinetically favored primary product is gradually converted to the thermochemically more stable isomer F2N+-NF3 (Figure 4, 326 K). This results in the apparent slow establishment of the equilibrium for reaction 4 with n = 1 in the lower temperature region. The proton affinity (PA) of N F 3 recommended by Lias et al. is 144 kcal/mol." McMahon and Kebarle determined PA(NF3) to be 136.9 kcallmol by high-pressure mass spectrometry.I2 Schwarz calculated the N- and F-site proton affinities for NF3 to be 131.8 and 138.2 kcal/mol, respectively, by means of highlevel ab initio MO calculation.2 The theoretical value 138.2 kcal/mo12 is in excellent agreement with the experimental value 136.9 kcallmol. ' , I 2 In Figure 6, the decrease of the H+(N2)2 ion is only explicable by proton-transfer reaction 8. It was found that the decay rate of H+(N2)2 increases with an increase of temperature. This suggests that reaction 8 is endothermic. Figure 9 shows the Arrhenius-type plot of the rate constants for reaction 8. From the slope of the straight line, the activation energy of reaction 8 can be estimated to be 1.8 kcaYmo1. From the values of PA(N2) = 118.2 kcaYmolIOand the bond dissociation energy of Hf(N2)2 (16.0 kcal/mol)I3, the lower limit of PA(NF3) can be estimated to be PA(NF3) 1 118.2 16.0 - 1.8 = 132.4 kcaY mol. This value is in line with that obtained by McMahon and Kebarle.I2
+
4. Theoretical Results and Discussion Figure 10 shows results of N F 3 . N F 3 is of C3, symmetry and has a quite cationic nitrogen atom (f0.80 in Figure lO).I4 Seemingly, the nitrogen atom plays a role exclusively as an electrophile. However, the frontier orbital, HOMO, demonstrates that N F 3 can also be a nucleophile.
15826 J. Phys. Chem., Vol. 99, No. 43, I995
Hiraoka et al.
NF3
n=l lt4.80)) F O :
(charge-acceptor MO)
0-
n
d (charge-donor
n=2
(-027)
MO)
Figure 10. Geometry of N F 3 (left) optimized with RHF/6-31G+G and RHF/6-3 1G in parentheses and its frontier orbitals (right). Numbers in double parentheses attached to the geometry are RHF/6-31G electronic net charges (positive, cationic).
NF;
n=3
I
W
Strengthened due to loss of an electronin the N-Fr-type antibonding.
(-0.05)
NF2'
LUMO (-0.39
NF3 H' lro.93)
(4.441
lii\
M = 139.6 kcal/mol
- a$/"
Figure 11. Geometries of the three cation species and their orbital characteristics. For NF3H+, the proton affinity (PA) was evaluated here with the GAUSSIAN-2+ theory?
Figure 11 shows three cationic species derived from NF3. is found to have the shorter N-F bond and the better planarity than NF3. The geometric difference between NF3+ and N F 3 is explicable by the electron loss from HOMO of NF3. The loss leads to the weakening of the n-type antibonding nature, i.e., the N-F shortening and the better planarity in W3+. In NF3+, the positive charge is localized on the nitrogen atom (+1.15). Next, N F 2 + has a vacant orbital perpendicular to the molecular plane as LUMO. In NF3H+, this LUMO of NF2+ is used to accept the electronic charge of a lone-pair orbital of HF. In this respect, the F-protonated NF3H+ is regarded as a weak Mulliken charge-transfer complex. In this work, the most accurate value of the proton affinity (PA) of NF3 is obtained by the G2+ theory to be 139.6 kcaymol. The G2+ method9 is a refined version of the G1 theory used in ref 2. The present value 139.6 kcaVmol is similar to that of 138.2 kcaVmol of Schwarz et a1.2 Figure 12 shows the geometry of F-(NF3),, (n = 1, 2, and 3). Although binding energies are of the electrostatic magnitude
Figure 12. Geometries of F-(NF3)", n = 1, 2, and 3, optimized with the RHF/6-31+G (without parentheses, for n = 1 and 2) and RHF/63 1 (+)G (in parentheses) methods.
in Table 1, the position of F- toward NF3 is controlled by the frontier-orbital interaction in n = 1. This interaction mode is for the back-side attack of sN2 reactions. A N-F bond is elongated (1.442 A) for the F- displacement.
F-( NF+
NF3+
The second NF3 is attached to F-(NF3)1 to form a symmetric geometry practically in the C2 point group. In n = 2, also the sN2-type coordination is obtained. Likewise, the n = 3 geometry is of C3 symmetry. In accordance with the result of the small absolute value and the n-dependence of binding energies in Table 1, symmetric structures of F-(NF3),, have been obtained. The n = 2 3 gap of AHO,,-l,,, is reflected in the change of the F-***Ndistances (2.596 2.619 8).There is some steric crowding between NF3 molecules in n = 3. The n = 4 geometry, though not calculated due to a too large size, is predicted to be of a tetrahedral symmetry. In Table 1, theoretical binding energies of F-(NF3),, in square brackets and parentheses are in good agreement with observed ones.
-
-
IodMolecule Reactions in
J. Phys. Chem., Vol. 99, No. 43, 1995 15827
NF3
NFz* + NF3
NFP + NF3
F F
\!
+ N--------F--N...
-12.04
* ':.
Figure 14. Potential energy diagrams of NF2+(NF3)1 by MP4(SDTQ)/ 6-3 1 G*//RHF/6-3 1G calculations. Energies are in kcallmol, and those in parentheses are experimental ones obtained here. The geometries of a, TS, and b are shown in Figure 13.
process, in line with Hammond's p~stulate.'~Next, NF2+(NF3)2 (n = 2) geometries are examined in Figure 13. For the stable F*N-NF3+ species, two coordinations of the second N F 3 are obtained, a and b. The electrostatically bound isomer (a) is computed to be more stable than isomer b. However, the binding energy AE = -2.66 kcaVmol (exptl, -2.8 kcaVmol in Table 1) is quite small in the former cluster. This energy indicates that the positive charge is delocalized well in F2NNF3+.
. ---.
The enthalpy change AH'"o,, = -17.7 kcal/mol in Table 1 has been derived from the van't Hoff plot of Figure 3. The activation energy (8.5 kcal/mol) has been obtained from the Arrhenius plot in Figure 8. This energy corresponds to the bond energy in model a. Those energies are evaluated by the MP41 /6-31G*//MP2/6-31G method and are shown in Figure 14. The binding energies of isomers a and b are reproduced beautifully by theoretical calculations. The activation-energy barrier along the a b reaction is computed to be 3.36 kcal/mol. According to calculations, the dissociation of b to two fragments, N F 2 + and N F 3 , does not give rise to an activation-energy barrier on the right side of Figure 14. The direct production of b is thought to be unfavorable due to a quite small cross section.
-
AE a -266
LUMO
/
The anti-bonding nature narrows the reaction channel.
mF3
/
small cross section for (b).
Figure 13. Two geometric isomers (a and b) of NF2+("F3)1. TS stands for a transition state of another isomerization. For the two n = 1 isomers a and b, the -AE values are MP4/6-31G* binding energies. For the two n = 2 isomers (a and b), the -AE values are RHF/6-31G binding energies.
Figure 13 shows optimized geometries of two isomers (a and b) of NF2+(NF3), (n = 1). The experimentally deduced coexistence of the two isomers a and b is confirmed theoretically. Isomer a is mainly composed of the electrostatic attraction. However, the charge-transfer works also to some extent, because of the out-of- lane coordination of N F 3 onto N F 2 + and an elongated (1.46 ) F-N bond of NF3. Isomer b is dominantly composed of the charge transfer for formation of a N-N covalent bond. But the bond with 1.553 A is incomplete in view of the standard N-N length, -1.45 A. An isomerization transition state (TS) has been obtained, and its geometry is close to that of isomer a. Isomer a is less stable than isomer b, and the TS is located at an early stage of the a b exothermic
'exchange repulsion narrows the reaction channel.
vs.
R
-
[ 00 -cm
large cross section for (a).
F
the probability of three N-F bonds
1
1
Hiraoka et al.
15828 J. Phys. Chem., VoL. 99, No. 43, 1995
The handicap in b corresponds to the slow establishment of equilibria for reaction 4, n = 1. A similar contrast is observed in the gas-phase FC2HsF reaction.I6 While the S N ~ activation energy is much smaller than the E2 one, E2 occurs excl~sively~~ due to the difference of cross sections.
n=l
+
(-0.06)
Q
(4.22)
un4
Figure 15 exhibits two geometric isomers (a and b) of NF3+(NF3)i (n = 1). In a, the intermolecular F***Ndistance (2.360 A) and the F-N bond distance (1.430 A) are larger and smaller than those for NF~+(NF~)I in Figure 13, respectively. This comparison demonstrates that NF3+ is a poorer electrophile than N F 2 + . NF3+ has a singly occupied frontier orbital (SOMO), while NF2+ has a vacant orbital (LUMOin Figure 11). Due to the half vacancy in NF3+, the N***Nbond formation in isomer b of Figure 15 is quite incomplete with the bond energy 6.22 kcaymol. In isomer b of Figure 15, it is noteworthy that an eclipsed F3N***NF3+form is shown, where a corresponding staggered form is less stable. The stability of the eclipsed form is ascribed to the secondary charge transfer from the HOMO of N F 3 to the SOMO of NF3+. SOMO
AE = -6.22
n=2 (-0231
(-0.0s)
(0.06)
AE -5.25
w (-023)
HOMO
ec I i psed configuration
The computed stability of a is almost the same as that of b. The much smaller energy difference between a and b in NF3+(NF3)l than that in NF2+(NF3)1 comes from the difference between HOMO SOMO and HOMO LUMO chargetransfer interactions. In n = 2 of Figure 15, three geometric isomers are exhibited. Isomer a is a double electrostatic coordination, isomer b is an electrostatic and charge-transfer coordination, and isomer c is a double charge-transfer one. Isomers a and b are found to have similar stability. Whether the n = 1 geometry is of the electrostatic attraction or of the charge-transfer bonding, the second N F 3 is coordinated electrostatically with similar stability.
-
$974,
-
5. Concluding Remarks In this work, clustering reactions of N F 3 by several ions have been studied. Competitions of (a) electrostatic and (b) chargetransfer bonding modes have been demonstrated. In particular, it is a striking result that isomer b of NF2+(NF3), is much more stable than isomer a in spite of the large cation-cation repulsion in the former isomer. In conclusion, NF3 and its fragment ions cannot be bound so strongly as is expected from the large difference of electronegativities between nitrogen (3.0) and fluorine (4.0) atoms.
(-0.06) (-0.08)
(-022) (-022)*
AE = -0.22
Figure 15. Two geometric isomers (a and b) of NF$(NF3)1 (n = 1). -AE is an MP4(SDTQ)/6-3lG* binding energy in kcaVmol taken from Table 1. Three geometric isomers (a, b, and c of NF3+(NF& (n = 2). -AE is a ROHF/6-31G binding energy.
Acknowledgment. The authors thank the Information Processing Center of Nara University of Education for the allotment of CPU time on the CONVEX C-220 computer. The present work is supported in part by a Grant-in-Aid for Scientific Research on Priority Area ‘‘Theory of Chemical Reactions’’ from the Ministry of Education, Science and Culture, Japanese Government, and the Morino Foundation for Molecular Science. We thank Miss Miwako Kanno for her partial assistance in theoretical calculations. References and Notes (1) Fisher, J. J.; McMahon, T. B. J. Am. Chem. SOC.1988, 110,7599.
(2) Grandinetti, F.; Hrusak, J.; Schroeder, D.; Karrass, S.; Schwarz,
H.J. Am. Chem. SOC. 1992, 114,2804.
Ion/Molecule Reactions in NF3 (3) Kebarle, P. In Techniquesfor the S m d y of lon-Molecule Reactions; Farrar, J. M., Saunders, W. H., Jr., Ed.; Wiley: New York, 1988. (4) Kebarle, P. J . Am. SOC. Mass Spectrom. 1992, 3, 1. (5) Hiraoka, K.; Takimoto, H.; Yamabe, S. J . Phys. Chem. 1986, 90, 5910. (6) Hiraoka, K. J . Chem. Phys. 1987, 87, 4048. (7) Frisch, M. J.; Trucks, G. W.; Head-Gordon, M.; Gill, P. M.; Wong, M. W.; Foresman, J. B.; Schlegel, H. B.; Robb, M. A.; Replogle, E. S.; Gomperts, R.; Andres, J. L.; Raghavachari, K.; Binkley, J. S.; Gonzalez, C.; Martin, R. L.; Fox, D. J.; DeFrees, D. J.; Baker, J.; Stewart, J. J. P.; Pople, J. A. Gaussian 92, Revision C; Gaussian Inc.: Pittsburgh, PA, 1992. (8) For anion systems, inclusion of diffuse orbitals (+) is indispensable. See: Clark, T.; Chandrasekhar, J.; Spitznagel, G. W.; Schleyer, P. v. R. J . Comput. Chem. 1983, 4 , 294. (9) (a) Curtiss, L. A.; Raghavachari, K.; Trucks, G. W.; Pople, J. A. J . Chem. Phys. 1991, 94, 7221. (b) Curtiss, L. A,; Jones, C.; Raghavachari, K.; Trucks, G. W.; Pople, J. A. J. Chem. Phys. 1990, 93, 2537. (c) Pople, J. A.; Head-Gordon, M.; Fox, D. J.; Raghavachari, K.; Curtiss, L. A. J . Chem. Phys. 1989, 90, 5622.
J. Phys. Chem., Vol. 99, No. 43, 1995 15829 (10) Keesee, R. G.; Castleman, A. W., Jr. J . Phys. Chem. Ref. Data 1986, 15, 1011.
(1 1) Lias, S. G.; Liebman, J. F.; Levin, R. D. J . Phys. Chem. Rev. Data 1984, 13, 695. (12) McMahon, T. B.; Kebarle, P. J . Am. Chem. SOC. 1985, 107,2612. (13) Hiraoka, K.; Saluja, P. P. S.; Kebarle, P. Can. J . Chem. 1979, 57, 2159. (14) Apparently, this large charge separation seems to be inconsistent with the observed small dipole moment, 0.2 D. However, the calculated moment is also small, 0.41 D. The small moment of N F 3 is ascribed to the short distance between cation and anion centers along the principal axis. (15) Hammond, G. S. J . Am. Chem. SOC. 1955, 77, 334. (16) Minato, T.; Yamabe, S. J . Am. Chem. SOC.1985, 107, 4621. (17) Ridge, D. P.; Beauchamp, J. L. J . Am. Chem. SOC. 1974, 96,637. JP9506705
