6774
J. Phys. Chem. 1993,97,6114-6115
Structure and Stability of Trinitramide J. A. Montgomery, Jr.,. and H. H. Michels United Technologies Research Center, East Harvord, Connecticut 06108 Received: April 28, 1993
Ab initiocalculations of the structureand thermochemistry of trinitramide, N(NO2)3, are reported. A vibrationally stable C3 structure having 1.545 A N-N bond lengths is found at the MP2/6-31G* level of theory. The heat of formation of trinitramide is estimated to be 59 kcal/mol. Thermal decomposition of trinitramide most likely occurs via N-N bond cleavage, which is estimated to require 26 kcal/mol. Introduction There has been growing interest recently in the development of chlorine-free oxidizers for use in advanced propellant formulations. Nitramines, which contain one or more covalentlybonded N-NO2 groups, constitute a promising class of chlorine-free energetic molecules. Examples of such compounds are the wellknown explosives RDX and HMXIJ and the recently synthesized hydrogen and ammonium dinitramide.3-s Although there are many possiblecompoundsof the form RR'NN02 and RN(N02)2, there is only one possible trinitramide, N(NO2)3. There has been speculation concerning the existence of trinitramide, but very few results have actually been reported to date. Miroshnichenko et aL6 have estimated the heat of formation of N(NO2)3 as 38 kcal/mol from group additivity arguments. An SCF calculation of the structure of trinitramide has been performed by Redfern and Politzer? following a suggestionby Olah.* To our knowledge, no experimentalobservations have been reported. In the present paper, we report ab initio calculations, including electron correlation, indicating the existence of a vibrationally stable trinitramide structure and from them derive estimates of some of its thermochemical properties.
theRHF/6-31G* level, the planar structureis 24 kcal/mol higher in energy than the pyramidal ground state. The calculated vibrational frequencies(see Table 11) for the pyramidal structure are all real, indicating this structure is a true minimum. Three of the RHF/6-3 1G*vibrational frequencies found for the planar trinitramide structure are imaginary, indicating that it is not a stable minimum or a transition state, and therefore no further calculations were performed on this structure. Due to strong electrostatic interactions (crowding), the N-N bond distance in the pyramidal trinitramide structure is considerably lengthened (by 0.11 A at MP2/6-31G1) over that previously reported5 for the dinitramide anion, N(N02)2-. As a consequence, the N-0 bond lengths are slightly shorter (0.03 A at MP2/6-31G*) in trinitramide. Other known oxides of nitrogenI2have even longer N-N bond lengths: N2O3 (1.86 A) and N204 (1.78 A). Thus, structural considerations do not preclude the existence of trinitramide as a stable species, and its stability will depend on the N-N bond strength. The heat of formation of trinitramide can be estimated from the heat of reaction for NH,
Theoretical Methods Ab initio calculations were performed on pyramidal (C3) and planar (D3h) structures of N(N02)3, using the Gaussian 909and Gaussian 9210 programs. Full geometry optimizations and harmonic vibrational frequency calculations were performed on both planar and pyramidal structures at the RHF/6-31G* level of theory. An additional MP2/6-3 lG* geometry optimization and harmonic vibrational frequency calculation was performed on the pyramidal structure. Thermochemicalcalculationsinclude vibrational zero-point effects, estimated from the RHF/6-3 lG* vibrational frequencies scaled by 0.8929.ll A summary of the calculated results is given in Tables I and 11.
-
+ 3HN03
or, alternatively, from
3NH2N02
N(NO,),
N(NO,),
+ 3H,O
+ 2NH3
While the latter process is isodesmic,13its use requires estimation of the heat the formation of nitramide. From G2 calculations,14 we find MP(NH2N02) to be 4.1 f 2.0 kcal/mol. Using experimental values15 for the heats of formation of nitric acid, water, and ammonia, we calculate the heat of formation of trinitramide from the ab initio heats of reaction. The results are summarized in Table 111. The cancellationof errors arising from the use of an isodesmic reaction is offset in this case by the uncertainty in the heat of formation of nitramide. Therefore, we Results and Discussion average the MP2/6-3 1G* heats of formation calculated from the two reactions to obtain our final estimate, MP(N,O,) = As in ammonia, nitrogen trifluoride, and trimethylamine, our calculations show that trinitramide is pyramidal, not planar. At 59.0 kcal/mol. Analogous calculationsSof the heat of formation TABLE I: Calculated Energies (hartrees) and Geometries (angstroms and degrees) of Trinitramide level C3 RHF/6-31GS C3 MP2/6-31GS D3h RHF/6-31GS a The
energy -666.445 82 (0) -668.265 37 (0) -666.405 19 (3)
N-N 1.4444 1.5453 1.4663
N-01 1.1713 1.2191 1.1698
N-02 1.1757 1.2209 1.1698
N-N-N 110.9 105.2 120.0
N-N-01 112.6 111.0 116.2
N-N-02 116.7 117.4 116.2
N-N-N-01 -70.8 -76.0 0.0
N-N-N-02 110.4 105.2 180.0
number of imaginary vibrational frequencies is indicated in parentheses.
TABLE 11: Calculated Vibrational Frequencies (cm-l) and Intensities (km/mol) of Trinitramide C3 RHF/6-31GS
A
E C3 MP2/6-31GS
A
E
127 (0.6), 241 (5.0), 444 (O.O), 651 (4.7), 937 (10.3), 997 (9.3), 1641 (31.5), 1934 (869.1) 53 (O.l), 249 (3.7), 508 (2.3), 801 (8.7), 908 (155.5), 1155 (79.3), 1537 (290.8), 1979 (550.1) 96 (O.l), 217 (2.1), 325 (O.l), 521 (0.2), 796 (12.1), 835 (6.0), 1350 (14.3), 1912 (90.5) 43 (O-l), 197 (4.2), 321 (58.0), 605 (59.7), 748 (172.1), 800 (70.9), 1263 (180.6), 1921 (58.9)
0022-3654/93/2091-6114$04.00/0
0 1993 American Chemical Society
The Journal of Physical Chemistry, Vol. 97, No. 26, 1993 6715
Letters TABLE IIk Calculated Heat of Formation (kcal/mol) of Trinitramide reaction RHF/6-3 1G* MP2/6-3 lG*
-
+
3NHzNO3 -N(NO2)3 2NH3 NH3 3HNO3 N(NO2)3 + 3H20
+
91.3 91.5
55.0 63.1
Acknowledgment. This work was supported in part by AF Phillips Laboratory under Contract F04611-90-C-0009. References and Notes
of nitric acid give an error of 1.7 kcal/mol. Our estimated error in the calculated heat of formation of trinitramide is f 5 kcal/ mol. As suggested earlier, the stability of trinitramide is determined by the strength of the long N-N bond. Using our calculated heat of formation of trinitramide, we find the heat of reaction for to be 78 kcal/mol and estimate the N-N bond strength as onethird of that, or 26 kcal/mol. Alternatively, if we estimate a value of 100 kcal/mol for the N-H bond strength in HN(N02)2, then from our previously reported5 heat of formation (28.4 kcal/ mol) of HN(N02)2 and the experimental heat of formation of NO2 (8.6 kcal/mol),15 we find the heat of formation of the N(N02)2 radical to be 76 kcal/mol. We may now directly calculate &(NOrN(N02)2) from and obtain a value of 26 kcal/mol, in agreement with the average bond strength. The enthalpy of dissociation into ions, via N(NO,),
ramide should be similar to or slightly greater than that of dinitrogen pentoxide.
-
N(NO2);
+ NO:
is estimated as 135 kcal/mol from the ionization potential of NO2 (9.75 eV)16andourpreviousestimateof theelectronaffinity of N(N02)2 (5.0 eV).5 From experimental thermochemical data,16 the &NO2 bond strength in N ~ O S is estimated to be 20-23 kcal/mol. Comparing this value with our estimated N-N bond strength, we conclude that the thermal stability of trinit-
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