Chemical Physics Letters 398 (2004) 560–563 www.elsevier.com/locate/cplett
The enormous apparent gas-phase acidity of cubylamine Jose´-Luis M. Abboud a,*, Ibon Alkorta b, Peeter Burk c,*,1, Juan Z. Da´valos a, Esther Quintanilla a, Ernest W. Della d, Ilmar A. Koppel c,*,1, Ivar Koppel c,1 Instituto de Quı´mica Fı´sica ÔRocasolanoÕ, CSIC, C/Serrano, 119, E-28006 Madrid, Spain b Instituto de Quı´mica Me´dica, CSIC, C/Juan de la Cierva, 3, E-28006 Madrid, Spain c Institute of Chemical Physics, University of Tartu, 2 Jakobi St., 51014 Tartu, Estonia Department of Chemistry, Flinders University, Bedford Park, South Australia 5042, Australia a
d
Received 21 September 2004; in final form 28 September 2004 Available online 18 October 2004
Abstract The high acidity of cubylamine (1NH2) seems to originate in the release of strain energy attending the breaking of some C–C bonds in 1NH. This process is greatly facilitated by the strong stereoelectronic interactions in 1NH. The anionic species thus formed are less strained, and their corresponding conjugate acids seem unable to Ôborrow strengthÕ from the residual strain, at least within the time-scale of the FT ICR experiments. Ó 2004 Elsevier B.V. All rights reserved.
1. Introduction The protonation of cubane (1) in the gas phase is an irreversible process leading to the opening of the cubane framework [1] and the same holds for the protonation of 1NH2 [2]. The protonation of cubylamine, 1NH2, in the gas phase takes place on the hydrocarbon moiety and leads to protonated amines, 2aH+, 2bH+ and 2cH+ [2] (see Fig. 1). Again, this is a highly exothermal and exergonic process, essentially driven by the release of strain attending the opening of the cubic moiety. In aqueous solution, however, protonation takes place on the nitrogen, leading to 1NHþ 3 (aq) [2]. An important feature of 1NHþ (aq) is the small value of its aqueous pKa, over 3 two units lower than that of aliphatic and alicyclic primary amines and even lower than that of ammonium ion. This originates in the significant stabilizing interac*
Corresponding author. Fax: +34 91 5855184. E-mail addresses:
[email protected] (J.-L.M.
[email protected] (I.A. Koppel). 1 Fax: +372 7 375 264.
Abboud),
0009-2614/$ - see front matter Ó 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.cplett.2004.09.127
tion between the lone pair of the amino group and the antiperiplanar, antibonding C–C orbital of the cubic framework of the neutral amine, 1NH2 [3,4]. From the standpoint of molecular geometry, the interaction is reflected by the C–C and C–N bond lengths around the a carbon atom (see Fig. 2a) [2]. These facts have prompted us to explore the gasphase deprotonation of 1NH2 both experimentally (using Fourier Transform Ion Cyclotron Resonance Spectroscopy, FT ICR) and by means of DFT calculations at the B3LYP/6-311+G(d,p) level.
2. Experimental The FT ICR experiments were performed on a modified Bruker CMS-47 FT ICR mass spectrometer used in previous studies [2]. The sample of 1NH2 is of the same batch used in protonation experiments [2]. Geometries of the neutral acids and the related anions were fully optimized at the density functional B3LYP/6-311+G** level using the GAUSSIAN 98 program [5]. Harmonic frequency analysis was used to
J.-L.M. Abboud et al. / Chemical Physics Letters 398 (2004) 560–563
561
NH3+
H+(aq)
NH2
1NH3+ H
1NH2
H
NH2+
CH2
NH2+ H H
+
+
H+(g)
NH2+
2bH+
2aH+
2cH+
Fig. 1. Protonation of cubylamine in the gas phase and water.
1.571 Å 1.571 Å
1.592 Å
1.658 Å
1.427 Å
1.360 Å
1.568 Å
1.569 Å 1.569 Å
1.576 Å
1.570 Å
(a)
(b)
Fig. 2. Selected geometry parameters for neutral: (a) deprotonated; (b) cubylamine (B3LYP/6-311+G(d,p)) data).
verify that either minima or first order saddle points (corresponding to transition states) on the potential energy surface were located. Unscaled B3LYP/6-311+G** frequencies were used to calculate the enthalpies and free energies, taking into account zero-point energies, finite temperature (0–298 K) correction, the pressure volume work term and entropy term as appropriate. The raw results of these calculations are given in Table 1. The gas phase acidity of an acid AH was calculated as the standard Gibbs energy change for deprotonation reaction (1): AHðgÞ ¼ A ðgÞ þ Hþ ðgÞ
ð1Þ
3. Results and discussion We have found experimentally that in the gas phase, iso-amoxide anion (i-C5H11O), generated from isoamyl nitrite by electron capture, deprotonates 1NH2 to yield the anion C8H8N (3). This is remarkable in itself, because the gas phase acidity [6] of iso-amyl alcohol amounts to 366.7 ± 2.1 kcal/mol, while the gas phase acidity of di-iso-propylamine (the most acidic aliphatic
Table 1 Raw computational results obtained in this work Species
Total energy
Enthalpy
Gibbs free energy
1NH2 1NH 3a1H 3a 1 3a2H 3a 2 3b 1 3b1H 3b2H 3b 2 1-AdNH2 1-AdNH 1H (CH3)3CH (CH3)3CNH2 TSð1NH ! 3a 1Þ TSð1NH ! 3b 1Þ TSð3a 1 ! 3a2 Þ
364.9080 364.2836 364.9621 364.3397 365.0081 364.4104 364.3649 364.9619 364.9690 364.3494 446.1952 445.5550 309.5327 158.5064 213.8754 364.2701 364.2819 364.3345
364.7516 364.1444 364.8047 364.1984 364.8511 364.2688 364.2239 364.8051 364.8128 364.2090 445.9270 445.3041 309.3944 158.3698 213.7198 364.1356 364.1437 364.1949
364.7885 364.1816 364.8419 364.2358 364.8930 364.3118 364.2632 364.8449 364.8521 364.2485 445.9686 445.3456 309.4251 158.4033 213.7566 364.1723 364.1798 364.2321
All values in hartree, enthalpies and free energies at 298 K.
secondary amine studied to date) is 382.8 kcal/mol. Furthermore, 1NH2 is seen to clearly protonate the anions of benzyl alcohol and 2,2,2-trifluoroethanol (the gas
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phase acidities of the corresponding neutral acids being, respectively, 363.4 and 354.1 kcal/mol [6], 1NH2 ðgÞ þ A ðgÞ ¼ 3 ðgÞ þ HAðgÞ
ð2Þ
Ion-selection experiments involving widely different concentrations of the neutral species reveal that methyl-, ethyl-, and tert-butyl mercaptide anions as well as phenoxide anion (PhO) (gas phase acidities of the corresponding neutral acids being 350.6, 348.9, 346.2 and 342.9 kcal/mol, respectively [6]) are protonated by 1NH2. In turn, the corresponding mercaptans and even phenol are able to protonate 3(g). Further notice that we were not able to find any base weaker than PhO and still able to deprotonate 1NH2. The apparent acidity of 1NH2 is thus some 40 kcal/ mol higher than that of the most acidic aliphatic secondary amine so far examined (di-iso-propylamine, DGacid ¼ 382:8 kcal/mol [6]). This suggests large structural changes following deprotonation. Our computational results (Table 1) fully confirm this contention. Deprotonation of any of the C–H bonds of 1NH2 can be ruled out. These reactions are endergonic by some 390 kcal/mol. This is less than for the deprotonation of the parent hydrocarbon (396.5 kcal/mol) [7,8] but well above the experimental value. NH-deprotonation of the amino group to yield 1NH is easier, as it is calculated to be endergonic by
H H
NH-
H H
H+
NH2
374.6 kcal/mol. This value, while still well above experiment, lies between those for di-iso-propylamine and aniline (359.1 kcal/mol) and is quite consistent with the concept of strong stereoelectronic interactions between the nitrogen lone pairs on 1NH(g) and the hydrocarbon moiety. Fig. 2b clearly shows the influence of these interactions on the C–C and C–N bond lengths in 1NH. Furthermore, the 1-adamantylamine, a bulkier, more polarizable amine [2] is computed (this work, see Table 1) to be in the gas phase 10 kcal/mol less acidic than 1NH2. NBO calculations [9,10] similar to those performed on 1NH2 indicate that the stabilizing interactions between the lone pairs on the NH group and the corresponding antiperiplanar antibonding C–C orbitals in the hydrocarbon framework is twice as large as in the case of neutral amine [2]. An estimate of the stabilizing interaction energy between the amino group and the cubane moiety in neutral 1NH2 can be obtained through the computation of the standard enthalpy and free energy changes for the isodesmic reaction (3), respectively, DrH(3) and DrG(3). At the B3LYP/6-311+G(d,p) level we obtain that this reaction is exothermic, respectively, by 4.5 and 6.3 kcal/mol, 1HðgÞ þ iso C4 H9 NH2 ðgÞ ! 1NH2 ðgÞ þ iso C4 H10 ðgÞ
H H
H+ H
H
H
∆ Go
∆ Go
= -358.7
3b1H
3b1-
-35.4 (-33.6)
-51.2 (-49.9)
H
H
= -363.2
N-
H H
H+
NH
ð3Þ
H
∆ G acid = 372.5
3b2-
3b2H -39.9 (-38.4)
-42.0 (-40.5) H+ CH2 CN
CN
CH2 C-
∆G
NH2
H
3a2-
= 1.1
∆G
= 5.6
-65.6 (-62.4)
= 2.3 N-
∆G
C H H
= -358.4
3a2H
-81.7 (-78.1)
NH-
H+
∆ Go
NH H+ CH2
CH2
∆ Go = -374.1
∆ G acid = 374.6
1NH2
1NH-
0.0 (0.0)
0.0 (0.0)
3a1-
-34.0 (-33.9)
3a1H -33.5 (-33.3)
Fig. 3. Energetics and mechanistic pathways for the deprotonation of 1NH2 and the isomerization of 1NH in the gas phase. Numbers given on figure are free energies (enthalpies are given in parentheses), computed at the B3LYP/6-311+G(d,p) level and given in kcal/mol. Values in bold-face are for neutral species relative to 1NH2 and in italics for anionic species relative to 1NH.
J.-L.M. Abboud et al. / Chemical Physics Letters 398 (2004) 560–563 1NH can evolve towards species 3a 1 and 3b1 (Fig. 3) following two nearly barrierless paths (actual barriers are 5.6 and 1.1 kcal/mol in free energy scale). In turn, à 3a 1 can practically without a barrier (DG = 2.3 kcal/ mol) rearrange into structure 3a2 . These structures ð3a 1 ; 3a2 ; and 3b1 Þ are, respectively, 34.0, 81.7 and 51.2 kcal/mol more stable than 1NH and involve the breaking of one or two C–C bonds. The fact that these bonds break so readily is fully consistent with the stereoelectronic interactions mentioned above as well as with the release of internal strain. However, it should be kept in mind that the effect of the release of strain can not be observed fully due to the relatively high basicity of the anions formed by the opening of the cube. The apparent gas phase acidity of 1NH2 to yield any of these species would thus amount to 340.6, 292.9, and 323.4 kcal/ mol, respectively. The first value is the closest to the experimentally observed one. Further evolution of obtained anions towards the thermodynamically most stable cyclooctatetraene derivative (COT-NH) involves activation barriers well above 20 kcal/mol and is thus unlikely under the experimental conditions. Species 3a 1 ; 3a2 ; 3b1 ; and 3b2 are in turn, the anionic forms derived from structures 3a1H, 3a2H, 3b1H and 3b2H. These species have acidities in the 358.4– 374.1 kcal/mol range (Fig. 3). Therefore, 3a 1 ; 3a2 ; 3b1 and 3b2 are able to deprotonate phenol and the thiols indicated above. Our experimental results can be easily rationalized on the basis of these results: 1NH2(g) is deprotonated to yield 1NH(g), which in turn evolves leading to 3b 1 and 3a1 , the energy released upon the formation of the collision complex allowing the system to go over the small activation barriers involved. According to Fig. 3, the most favored anion both thermodynamically and kinetically, formed upon opening of 1NH in a single step reaction, is 3b 1 . It is interesting that the combination of our computational results with the experimental gas phase acidities for the various acids used in this study allow the determination of DrG(4) and DrG(5), as defined through Eqs. (4) and (5):
563
1NH2 þ PhO ! 3b 1 þ PhOH 3b 1 þ PhOH ! 3b1 H þ PhO
Dr Gð4Þ Dr Gð5Þ
ð4Þ ð5Þ
DrG (4) = 18.9 and DrG (5) = 20.9 kcal/mol. It follows that reaction (5) is more favorable than reaction (4) by 2.0 kcal/mol. This is not enough to prevent the slow re-protonation of PhO. It is significant that these are the only processes fully consistent with the experiments.
Acknowledgements This work was supported by Grants BQU2003-5487 (Spanish DGES) and 5226 (Estonian Science Foundation) and by a Joint Project CSIC-Est. Acad. Sci. ÔThe Reactivity of the Strained Cage MoleculesÕ.
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