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II E L S E V I E R Chemical Physics 192 (1995) 99-110

Chemical Physics

Is there a hydride transfer between N20H + and saturated hydrocarbons? *

Jordi Mestres a, Miquel Duran a, Juan Bertr~n a.., Emilio Ballesteros b, Marta Herreros b Jos6-Luis M. Abboud b,. 9

" Departament de Quimica, Universitat Autbnoma de Barcelona, E-08193 Bellaterra, Catalonia, Spain b Instituto de Qufmica F£sica Rocasolano, CSIC, E-28006 Madrid, Spain

Received 1 February 1994; in final form 4 October 1994

Abstract

Various reaction mechanisms for the reaction between N2OH + and CH 4 have been explored theoretically at the MP26-31G** level. Also computed were the structure and thermodynamic state functions for several neutral molecules, notably N20, CH4 and C2H 6 and ions (CH~, C2H~, C2H~-, N2OH ÷ ). These results were combined with a Fourier transform ion cyclotron resonance study of the reactions of N2OH ÷ with CH4 or C2H6. Both studies suggest that the title hydride transfer, while thermodynamically favorable, is not kinetically preferred under thermal conditions in the gas phase.

1. Introduction

Dinitrogen oxide, N20, is widely used in radiolytic studies, notably in the radiolysis of liquid hydrocarbons, RH [ l a - l i ] and silanes [ l j ] . There is strong evidence [2] suggesting that the key processes deter- mining its reactivity in these experiments are those summarized in Eqs. ( 1 ) - ( 3 ) , a variety of products originating in the ensuing reactions between N20- or O - and RH. Among them, dinitrogen and oxygen- containing compounds, such as water and alcohols are noteworthy [ l e - l f ,3 ] ,

N 2 0 + e - ~ N 2 0 - ,

N 2 0 + e - ~ N 2 + O - ,

N 2 0 + N 2 0 - ~ N 2 + N 2 0 + O - .

(1)

(2)

(3)

* This work is dedicated to Professor Y. Hatano. * Corresponding author.

0301-0104/95/$09.50 © 1995 Elsevier Science B.V. All rights reserved SSDIO301-OIO4(94)OO384-X

Reactions involving hydrocarbons and O(1D) formed in the decomposition of excited N20 are also known [41.

In this work, we investigate alternative reaction pathways involving N/OH + and saturated hydrocarbons, because reactions of this sort have been considered to be potential precursors of dinitrogen and oxygen-containing compounds in the y-radiolysis of hydrocarbon- N20 mixtures [ 5 ]. Quite generally, the overall process can be written as

N2OH + (g) + R-H(g)

R ÷ (g) + OH2(g) + N2(g) • (4)

This reaction is strongly exothermal for most hydrocarbons. Thus, for R = CH 3 and R = C2H 5 we find respectively A H % = - 28.4 and - 70.8 kcal. mol - l, using the enthalpies of formation [6] and proton affin- ities [ 7 ] of the appropriate species.

From a mechanistic point of view it can be thought that a hydride transfer from the hydrocarbon to N2OH ÷

100 J. Mestres et al. / Chemical Physics 192 (1995) 99-110

might take place, leading, in principle to N2OH2, as emphasized by the referees. The "solvation" of R ÷ by H20 to yield the protonated alcohols ROH~-, can also be envisaged on account of its exothermicity, - 66.2 kcal mol- 1 for R = CH3 i and - 35.3 kcal mol- ~ for R=C2H5 2. Two other possible processes deserve attention: one of them is the direct insertion of OH ÷ into the hydrocarbon molecule to yield ROH~-, and the other, the proton transfer from N2OH + to the hydrocarbon molecule:

N2OH÷ (g) +RH(g) ---,N20(g) +RH~-(g) . (5)

The possibility for reaction (4) to take place in the case R-H = C2H 6 has been envisaged by Mackay, Schiff and Bohme [ 10], who also carried out a systematic experimental study of the reaction between c 2 n 6 and several charged proton-donors. It is of interest that, when this manuscript was under revision, Cacace and Speranza reported [ 11 ] that protonated ozone, O3H + reacts with CH 4 to yield 02, CH~- and H20 in a reaction formally similar to reaction (4).

Here we report the results of a study of the systems N2OH + / c 2 n 6 and N2OH + / C H 4, our purpose being to assess the possibility for reactions (4) and (5) to take place in the gas-phase under thermal conditions. To this end, ab initio calculations [ 12] and Fourier Trans- form Ion Cyclotron Resonance mass spectrometry (FT ICR) [ 13], were used.

2. Computational techniques and experimental methods

The theoretical treatment has been carried out through ab initio calculations by means of the gradient techniques commonly used in studies of potential energy surfaces [ 14 ]. Stationary points were optimized without any symmetry constraints and then characterized by the correct number of negative eigenvalues of their cartesian second-derivative matrix. Frequencies were obtained from these Hessian matrices using the scaling factor discussed below. Geometry optimiza- tions were performed at the MP2 level in the full orbital space, using the 6-31G** [15] basis set, all calculations being carried out with the Gaussian-92 series of

1 Using the value of AHf ° (CH~-) given in Ref. [8]. 2 A/~f (C2H ~- ) was from Ref. 191.

programs [ 16]. With the same program we carried out several single-point configuration interaction calculations including simple and double excitations (CSID) to test the reliability of the MP2 results. We further determined the intrinsic reaction coordinate (IRC) from the two transition states found in the study, in order to ensure the connectivity with the corresponding intermediates. To construct the IRC, we have used the recently proposed reaction-path-following algorithm of Gonzfilez and Schlegel [ 17] which only requires the first derivatives at points along the path. This path begins at the transition state and the analytical Hessian is used to determine the transition vector that follows downwards. Finally, counterpoise corrections for basis set superposition error (BSSE) were determined by means of the Monstergauss program [ 18].

FT ICR experiments were conducted on a modified Bruker CMS-47 mass spectrometer, already used in other studies [ 19]. The foreline and high-vacuum sec- tions of the instrument were baked under vacuum for 48 h. All the reagents were Alphagaz products with nominal purities of at least 99.99%. Dinitrogen oxide and ethane were carefully degassed in a vacuum line and stored in vacuum-baked glass-bulbs. Methane was used as received. Ionization energies were kept as low as possible (nominal values in the 12.5-16 eV range). Pressures of the reagents were in the range 10-8-10-6 mbar. Ions were obtained by electron impact ionization of gaseous mixtures: (N20, CH4), (N20 , C2H6),

( N 2 0 , c n 4 , C2H 6). They were then selected using MS / MS techniques [ 13a,13d] and allowed to react. This was intended to simplify the interpretation of the experimental results.

3. Results and discussion

3.1. Theoretical treatment

We summarize in Table 1 the electronic energies at 0 K (EMp2), as well as the zero-point energies together with the thermal contributions at 298 K (ETHER) of the various species examined in this work. Also given are the corresponding standard entropies. These data are used to compute the changes in thermodynamic state functions for the reactions compiled in Table 2, wherein the corresponding experimental values have been col- lected.

J. Mestres et al. / Chemical Physics 192 (1995) 99-110 l 01

Table 1 Electronic energies (EMp2, in hartree), zero-point and thermal contributions at 298 K (ErnER, in hartree) and entropy (S, in cal/mol K) for the species involved in all processes considered

Molecule EMp 2 E, mE R S

I N T 1 A -224.809136 0.076422 86.107 INTIB -224.820298 0.075042 80.571 T S 1 2 -224.813543 0.069884 75.268 I N T 2 -224.813565 0.071663 79.162 TS13 -224.770165 0.071240 82.317 I N T 3 -224.974224 0.079628 78.792 NzOH + - 184.434724 0.024097 58.026 CH4 -40.369856 0.049466 44.416 N20 - 184.213684 0.013456 52.647 CH; -40.580284 0.056853 51.899 N2 - 109.261574 0.007326 45.874 H20 - 76.222449 0.024733 45.087 CH; -39.351195 0.035534 44.546 H20-CH~ -115.697355 0.069738 58.204 H2 - 1.157661 0.012861 31.084 C2H ~ -78.601183 0.066237 55.757 C2H¢ -79.784259 0.090703 61.693 C2H 6 - 79.553713 0.080985 54.158

The optimized geometries of the various stable species involved in all reactions considered, are presented in Fig. 1.

Oxygen protonation of N20 has been assumed throughout, as it is known from the high-level theoretical study [20] of experimental data [21 ] to be ener- getically preferred to nitrogen protonation. In

particular, Jones et al. [21c] have reported that the proton binding energies for the two potential basic sites differ by 6.5 _ 3.5 kcal m o l - ~. Del Bene and co-workers [ 20b] explici t ly crit icized the use of MP2 energies to reproduce the proton affinity of N20. Indeed, we have performed calculations at the M P 2 / 6 - 3 1 G * * / / SCF/6 -31G** level and found that nitrogen protonation is favored over oxygen protonation by 2.09 kcal

Table 2

mol-1 . At the C I S D / 6 - 3 1 G * * / / S C F / 6 - 3 1 G * * level,

the sign of the difference is correct, but now, the gap between the two protonated species is too large ( 12.81 kcal m o l - ~ or 9.88 kcal m o l - 1 if we introduce the size- consistency correction). This indicates the need for

carrying out the optimization of geometries at levels beyond SCF/6-31G**. In order to introduce electronic correlation effects, we have chosen the MP2 method as

the fastest possible approach on account of the size of the various systems examined. Thus, although Del

Bene 's considerations remain valid, we feel that the

MP2/6 -31G** level is sufficient for our present pur-

poses, on account of the satisfactory agreement between the calculated and experimental values of the

changes in thermodynamic state functions presented in Table 2. Furthermore, our calculations at this level of theory confirm the special stability of the bridged struc-

tures of C2H ~ [22,23] and C2H ~ [22,24]. From the

experimental values of the vibrational spectra of H20 and CH4 we obtain an average scaling factor for the

frequencies of 0.94. If we apply this factor to the com-

puted OH-stretching frequency for N2OH ÷, we find it

amounts to 3324 cm-~, in excellent agreement with Amano ' s [ 21 ] experimental value ( 3 331 c m - t ).

The three mechanistic possibilities envisaged above have been examined and a detailed discussion follows:

starting from N2OH + and CH4, two intermediates, labeled INT1A and INT1B (Fig. 2) are obtained. Ener-

getically, INT1A is found to be 7.00 kcal m o l - ~ less

stable than INT1B, thus reflecting the fact that, while

both are ion-induced dipole complexes, the latter is

further stabilized by hydrogen bond interactions

between the two reactants. Due to the electrostatic

nature of these two intermediates, we performed a counterpoise correction calculation at the SCF level in

order to estimate the effect of a basis set superposition

Comparison between calculated and experimental changes in state functions for some relevant reactions. All values in kcal m o l i

Reactions AHTaE AHExp AGTHE AGExp

(4a) N2OH + +CH4--*CH3 + +HzO+N 2 -23.0 (5a) N20+CH~- ~N2OH + +CI-I4 -4.6 (5b) N20 + C2H~- ---> N2OH ÷ + C2H6 6.6 (6) CH~" +C2H6~CH4+C2H~- - 11.2 (7) CH~- --'CHJ" +H2 39.5 (8) C2H~" ~C2H~- +H2 8.7

-28.4 a -6.51251, -7.3[26], -7.31271 -3.9

4.91261 7.3 - 12.21271 - 11.2

42.418.25] 11.8124a] 1.2

-6.2[25], -6.71271 4.4[281

-10.81271

4.4124a]

a Value determined (as indicated in the text) from experimental data for other reactions.


109.2

1.18 71.10

0 . 9 ~

6 B L51

109.7

0.73 ~ 9

t Q 1.30 09

Fig. 1. Optimized geometries of the species involved in this study.

error, and found to be too small (0.189 kcal mol - 1 for INT1A and 0.305 kcal mol - ~ for INT1B) to qualita- tively affect our results and conclusions.

3.1.1. Hydride and OH + transfers Starting with INT 1A we tried to find a transition state

for a hydride abstraction from methane by N2OH + but, when approaching methane to the NzOH + species, an

OH + abstraction was observed. As a result, the system evolved towards the transition state TS13 portrayed in Fig. 3. Notice that, for the sake of clarity, the arrows in this figure pertain to the evolution from the transition state towards INT1A. We emphasize here that at the computational level used in this work, the putative species N2OH2 does not appear as a minimum on the poten-

J. Mestres et al. / Chemical Physics 192 (1995) 99-110 103

", 2.9{) 3.32 ",

INT1A

For the purpose of analyzing the evolution of the electronic charge, we may consider the system as divided into three hypothetical fragments, namely, N2, OH and CHa. From values in Table 3 it can be seen that accumulation of charge on N2, OH and CH4 is respectively large, moderate and marginal, this being the case for both INT1A and INT1B.

At variance with this, in TS 13 there is a very clear delocalization of charge over the three fragments. This suggests that a substantial stabilization of TS 1A originates in charge delocalization which partially over- comes the energy loss associated to the loosening of the Nr-O and C-H2 bonds.

In order to get a further insight into the electronic nature of interactions between the fragments within TS13, we located the bond critical points (BCP) between atoms and characterized them by their values

// • " 1.85

1.42 . " " ' ' ' " 2~.(,"

TS13

INT1B Fig. 2. Structure of the intermediates INT1A and INTIB.

tial energy hypersurface, thus indicating that a direct hydride transfer from CH4 to N2OH ÷ will not take place. At variance with this, the structure of TS13 shows that the process actually involves a hydride transfer from methane together with a transfer of OH ÷ from N2OH +, as can be deduced from the normal mode associated to the single imaginary eigenvalue ( - 1065.1 cm-~) characterizing this transition state (Fig. 3). Thus, from a mechanistic point of view, we can dissect this process as follows: (i) the breaking of the N r O bond; (ii) the breaking of the C-H2 bond; and, (iii) the approach of both H2 and OHm-.

1¸5{)

I "11{} " 0 0 0 1{}.97

INT3

Fig. 3. Structure of the intermediate INT3 and the transition state TS13.


of the charge density and the Laplacian thereof. The results are summarized in Table 4. According to the Bader analysis [29], negative values of the Laplacian indicate a buildup of charge density and, therefore, a covalent character of the bond. On the other hand, positive values imply charge depletion, and, consequently, an electrostatic nature of the bond. From values in Table 4, the above mentioned aspects of the process are clearly observed: (i) the positive value of the Laplacian for the NI--O bond indicates an electrostatic interaction between both atoms; (ii) although the Laplacian for C-H2 still remains negative, its magnitude as well as that of the charge density are considerably smaller than those found in standard C - H bonds, this being a con- sequence of the weakness of the C-H2 bond, and (iii) the electrostatic interaction between the O-H1 fragment and H2 becomes clear from the positive value of the Laplacian at the BCP of O-H2 interactions. Finally, it is worth noting that no BCP between C and O has been found and that O-H2-C do not originate a ring critical point. Thus, although there is an OH 1 motion coupled with the transfer of H2 to this fragment, the process cannot be regarded as an insertion of OH into the C - HE bond.

From TS 13 we determined the intrinsic reaction path (IRP) at the SCF level in order to find which intermediates are connected to the transition state. The IRP leads to two intermediates, INTIB and INT3 (Figs. 2 and 3), the energy barriers being 31.5 and 128.1 kcal. mol-~, respectively, as indicated in Fig. 5. As shown in the diagram (Fig. 6), the corresponding Gibbs energy changes amount to 24.8 and 121.8 kcal m o l - t , respectively. It must be pointed out that, although the IRP falls into intermediate INT1B, TS13 has been for-

Table 4 Charge density p (in au- 3) and Laplacian of the charge density, V2p (in au-5) values at the bond critical point of the different bonds for the transition state TS 13

Bond a p Vp

C-H b 0.1886 - 0.5504 C-H2 0.1489 - 0.2504 O-H2 0.0520 0.1417 O-H1 0.2438 - 1.1750 O-N1 0.0896 0.3115 N1-N2 0.5294 - 1.4610

a See Fig. 2 for atom numbering. b Average values between C-H3, C-H4 and C-Hs.

mally connected to INT1A because of the correct ori- entation of the latter with respect to TS13. In fact, INT1A and INT1B can be interconverted by rapid rota- tion of the OH group within N/OH + . Hence, in principle, both intermediates can be regarded as being the origin of the process.

The INT3 species is essentially a protonated methanol solvated by a dinitrogen molecule (Fig. 3). From data in Table 3, it can be deduced that the latter is slightly polarized. Thus, if we compare the dinitrogen moiety in INT3 with the nitrogen atoms within a free dinitrogen molecule, we observe that a small amount of negative charge develops in N~ ( - 0.0177 au), and that the NI-N2 distance is reduced (from 1.1300 ,~ in the free molecule to 1.1287 ,~ in INT3) and, consequently, the Nl-N2 stretching frequency increases from 2179 cm -~ in free dinitrogen to 2192 cm -~ in INT3. The existence of a weak hydrogen bonding is revealed by comparison of the O - H distance in INT3 and in free protonated methanol. In the latter, they are equal to

Table 3 Mulliken atomic charges for the different stationary points located, using the generalized density corresponding to the second-order MP2 energy. For atom numbering see Figs. 1 and 2

Atom INT1A INT 1B TS 12 INT2 TS 13 INT3

N2 0.2212 0.1940 0.0915 0.0759 0.1842 0.0739 Nt 0.5961 0.5860 0.5599 0.5534 0.3508 - 0.0177 O - 0.3029 - 0.3433 - 0.4954 - 0.5045 - 0.3198 - 0.4813 H1 0.4756 0.4038 0.3396 0.3251 0.4308 0.4678 C - 0.5196 - 0.5688 - 0.5871 - 0.5856 - 0.5095 - 0.1441 H2 0.1343 0.1446 0.2208 0.2890 0.2440 0.4358 H3 0.1911 0.2099 0.2627 0.2752 0.2241 0.2108 FL, 0.1021 0.1440 0.2754 0.2828 0.2076 0.2258 H~ 0.1021 0.2300 0.2826 0.2887 0.1878 0.2290


0.98 ,~, while in the former, they respectively amount to 1.00/~ (O-Hi) and 0.97/~ (O-H2).

The calculated vibrational frequencies can be ana- lyzed as follows: we have computed the vibrational frequencies of free water at the same level and found the following (unscaled) frequencies: bending, 1684 cm- 1, symmetric stretching, 3895 cm- ~, and antisym- metric stretching, 4034 cm- 1. The corresponding frequencies in free protonated methanol are respectively, 1712, 3704, and 3810 cm- 1. As can be extracted from these results, the overall pattern is similar to that found for a water molecule under the effect of a uniform electric field: the bending has moved to higher frequencies, while stretchings move to lower frequencies, all line intensities becoming stronger [30]. This reflects the fact that the CH~ moiety of protonated methanol creates a strong electric field in and around the water molecule. It is worth pointing out that, even unscaled, the calculated frequencies are expected to be within 6% of the experimental values. This estimate is based on our results presented above as well as on similar results by Schleyer and coworkers [ 31 ].

3.1.2. Proton transfer In the study of the proton transfer process between

N2OH ÷ and CH4 to yield N20 and CH~', two minima were found (INT1B and IN'I2B), the transition state between those two minima (TS12B) being perfectly characterized by a single negative eigenvalue of the second derivative matrix, the associated eigenvector corresponding to the motion of the proton from CH~- to N20 to give INT1B. The structures of INT2 and TS12 are presented in Fig. 4.

As shown in the potential energy profile (Fig. 5), INT1B is stabilized by 9.86 kcal mol- ~ with respect to the reagents. As indicated, this stabilization originates both in ion-induced dipole interactions and in hydrogen bonding. The former is shown by inspection of the Mulliken atomic charges in the CI-I4 moiety (Table 3). Evidence for the latter is as follows: (i) The O-H~ distance increases from 0.99 A in free N2OH ÷ to 1.04 /~ in INT1B; (ii) 0.1595 au of electronic charge is transferred from CH4 to N2OH+; and (iii) the OH- stretching frequency in INT1B is lowered by 906 cm- 1 with respect to the value in free N2OH +.

In the potential energy surface, the intermediate complex INT2 is found 4.23 kcal mol- ~ above INT1B and is stabilized by 12.30 kcal mol-~ with respect to

~ " 1 .46

~ . ~ . ~ 1 . 1 6

TS12

1 1 3

• 1 .56

INT2 Fig. 4. Structure of the intermediate INT2 and the transition state TS12.

the final products. Analysis of its structure shows it to be essentially a CH~ cation solvated by N20. Stabili- zation mechanisms are similar to those in INT 1B. Thus, polarization of N20 by CH~- can be seen by comparing the charge distributions: the negative charge on the oxygen increases from -0.3932 au in free N20 to -0 .5045 au in IN'I2. This arises essentially from the relatively short distance to the positively charged H1. Hydrogen bonding can be inferred (i) from the distor- tion of the CH~- moiety. In free CH~-, the longest C- H distances are 1.18/k, the two loosely bound hydro- gens being separated by 0.96 ~,. In INT2 the C-H~ distance is stretched up to 1.25.4, and the H1-H2 distance up to 1.04 ik; (ii) the charge transfer from N20


TSI3

24.45 I

N2OH++CH4 /

TS12 9.~ ~

N20+CHs +

12.30

INT2 0.01

128.05

N~H20+CH3 ÷

/ I 77.63

Np'H20"CH3+ /

9.60

INT3

Fig. 5. Energy diagram for the various possible processes originating in N2OH + + CH4. Energies in kcai mol- t

to CH~- ( -0.1248 au); and (iii) the lowering of the

C-H1 stretching frequency in INT2. As regards the structure of TS 12, it is closely related

to that of INT2. As shown in Fig. 4, HI has been largely transferred to CI-I4. The O-H! distance is significantly different in IN'I2 from that in TS12 ( 1.56 and 1.46 .A respectively), but both structures have very similar energies (they differ only by 0.014 kcal mol- ~ in the potential energy profile). The extreme flatness of the potential energy surface is confirmed by the very low

value of the single imaginary frequency ( - 124.5 cm- i ), a most rewarding computational result.

3.1.3. Energetic aspects of the different processes Calculation of the second derivatives enables the

construction of the full Gibbs energy diagram for both reaction paths. The overall energy and Gibbs energy profiles for the various reaction pathways are shown in Figs. 5 and 6. The main conclusions drawn therefrom are as follows: (i) for the hydride-OH+coupled-trans -


TSI3

22.39 :

/ 8 ~ I ~ A TSI2

N2OH++CII4 ~2.42 2 . 5 ~ 8 \INT1B

N20+CH5 ÷

INT2 /3.88

0.06

121.80 N=+H20+CH3 +

/ I 62.32

INT3 N~"H20"CH3+ / f0.45

Fig. 6. Gibbs energy diagram for the various possible processes originating in N2OH + + CH 4. Energies in kcal mol - '.

fer process, the calculated activation barriers (from INT1B) are 31.5 kcal. mol-~ in energy and 28.6 kcal mol-1 in Gibbs energy. The process shows a large exothermicity. For instance the energy and Gibbs energy of the system decreasing respectively by 96.6 and 98.0 kcal mol- l in going from TS13 to INT3. It is noteworthy that, notwithstanding the high exothermicity of the process, the activation barrier is remarkably high; (ii) for the proton transfer, the activation barriers amount to 4.2 kcal mol -~ in energy and 2.58 kcal mol- ' in Gibbs energy, the depth of the well for INT2 being negligible; (iii) the most remarkable difference between both profiles lies in the fact that INT1A becomes unstable with respect to the reagents in the Gibbs energy profile.

From the theoretical treatment presented above, it follows that, although the hydride OH ÷-coupled transfer process is thermodynamically favored, proton transfer is favored kinetically. Larger hydrocarbons are

better able to stabilize positive and negative charges. This would likely lower both barriers, but proton transfer would likely remain strongly favored.

3.2. Experimental results

3.2.1. Mixtures N 2 0 - C H 4

Ratios of reagents varied between 1 : 3 and 1 : 5 with total formal pressures in the range (1.5-3.5) × 10 - 7

mbar. After reaction times of 0.5-1.0 s, all ions, except CH~ (m/z = 17) were ejected from the ICR cell by means of a radiofrequency "chirp". CH~ was allowed to react for variable periods of time (up to 10 s). The only reaction we observed was the fast (rate constant: 0.95 × 10 - 9 c m 3 molec- ' s 1 [25] ), essentially complete, proton transfer from CH~ to N20:

CH~ +N20--~N2OH + +CH 4 . (5a)


Given the small dynamic range of ICR (ca. 103), this result is in agreement with the AG o value for this reaction given in Table 2. It is important that Bohme, et al. [25] using flowing afterglow (thermal conditions) have reached the conclusion that no other reactions are observed in gaseous mixtures CH~- -N2OH ÷- CHn-N20. As regards the formation of protonated methanol, we have not been able to detect the formation of CH3OH~- (m/z = 33) or CH3 ÷ (m/z = 15).

The reaction

N20H + +CH4 -*N2 + CHaOH~- (10)

is exothermal by 94.6 kcal tool- 1 as indicated earlier and the reaction

N2OH ÷ + CH4 ~ N2 + H20 + CH~ (4a)

is exothermal by 28.4 kcal. mol-1 (the latter is also exergonic, because of the increase in entropy associated to this process). It is clear that while reaction (4a) and (10) are thermodynamically much more favorable than the inverse of reaction (5a), they cannot compete with it kinetically. This is in excellent agreement with the theoretical results presented above.

3.2.2. Mixtures of CH4--N20-C2H6 In a representative experiment, the reactants were in

a 5:20:1 ratio, the total nominal pressure being 2.6× 10 -6 mbar. Under these conditions, the most abundant ion 0.5 s after electron impact was N2OH ÷ (m/z=45). All other ions were ejected from the ICR cell after this time and NzOH + was allowed to react for variable periods of time. Under our working conditions we could observe the proton transfer from N2OH + to C2H 6. This transfer is not complete within a reaction time of 10 s, the ratio [C2H ~- ] / [N2OH + ] being then ca. 11. This behavior is consistent with equilibration and would indicate that a lower bound for AG ° for equilibrium (5b) is 3.5 kcal mol- 1. This gap is too wide for ICR to give a very reliable value and so, we did not attempt to reach full equilibration. It is rewarding, however, that flowing afterglow experiments [28], wherein full equilibration is reached lead to a AG o value of 4.4 kcal mol- 1. These results are also in fair agreement with the theoretical calculations presented in Table 2.

Table 2 also shows that the decomposition of C2H ~- (in its bridged, most stable structure) into C2H~- and H2 (Eq. (8)) is endothermal by 11.8 kcal

mol- 1 and endergonic by only some 4.3 kcal mol- 1. Work by Kebarle [24a], Bohme [28], Futrell [32] and Yeh [24b] and their co-workers shows that the dissociation reaction does indeed take place. In our experiments it was clear that the only significant process up to reaction times of 2-3 s (depending on the conditions) was the formation of C2H~- (identified by the high-resolution spectrum) through proton transfer from N2OH + tO f E n 6. At longer times, the CEH~- sig- nal (m/z = 29, identified by the high-resolution spectrum) became significant. After 10 s, its intensity was ca. 15% of that of C2H~". Flowing afterglow experiments indicate that ca. 5% of C2H ~- is formed under thermal conditions [ 24 ]. In tandem-ICR experiments, the formation of 11% of c 2 n ~- has been reported [ 27 ]. These results suggest: (i) the contribution from some non-thermalized NEOH + ions and; (ii) the activation of C2H4- by the energy released in the protonation process.

Clearly, neither our experiments nor those by Futrell [32] and Bohme [28] and co-workers can rule out the formation of at least a fraction of C2H~- through hydride transfer. In any case, this reaction would be slower than proton transfer for which the rate constant at 298 K amounts to 1.1 × 10 9 cm 3 molec- 1 s- 1 [ 10].

More clear-cut results were as follows: (i) The possibility of N2 H+ being formed through

the following reaction, also thermodynamically favorable has been considered [ 28 ] :

N2OH + + C2H6

~ N 2 H+ +C2HsOH/(C2H4 + H 2 0 ) • (11)

Our high-resolution analysis of the peak at m/z = 29 shows that this ion is not formed.

(ii) We have not been able to detect the presence of C2HsOH + (re~z=47). Now, it seems reasonable to expect that if this species were an intermediate, some termolecular process would be able to stabilize at least a fraction of it prior to decomposition.

To summarize: (1) The very satisfactory agreement between the

theoretical and experimental changes in state functions obtained in this work strongly supports the main conclusions of the mechanistic study.

(2) Theoretical calculations show that hydride transfer (4) is thermodynamically favorable but has a sub- stantially higher activation barrier than proton transfer.


It is thus expected not to take place in the gas phase near room temperature.

(3) Experimental FT ICR results indicate rather con- clusively that reaction (4) does not take place between N20H + and C H 4.

(4) Analogous experiments involving c 2 n 6 and N2OH ÷ show proton transfer from N2OH ÷ to C2H 6 tO be the main reaction channel under our working conditions. Formation of N2 H+ and~or C2HsOH~ are ruledout. The formation of C2H~" is most easily ration- alized as arising from the decomposition of C2H~-, but the formation of at least part of it by hydride transfer cannot be entirely ruled out on a purely experimental basis.

(5) The possibility exists, at least in principle, for reaction (4) to occur at very high energies and/or densities (the latter would likely stabilize the multi- polar transition state TS1A involved in the hydride transfer process).

Acknowledgement

This work was partially supported by grants PL87- 0357 and PB90-0228-C02-02 from the Spanish D.G.I.C.Y.T. Work by E.B. and M.H. was supported by a scholarship from C.S.I.C.J.M. acknowledges financial support from the Spanish"Ministerio de Edu- caci6n y Ciencia" under project PB89-0318. We are grateful to the referees for their very valuable com- ments, especially for their suggestions on the possibility of an OH + transfer.

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