1536
J. Phys. Chem. 1996, 100, 1536-1544
Site of Ionization of Polyfunctional Bases and Acids. 1. Ab Initio Proton Affinities Alessandro Bagno* and Gianfranco Scorrano Centro CNR Meccanismi Reazioni Organiche, Dipartimento di Chimica Organica, UniVersita` di PadoVa, Via Marzolo 1, 35131 PadoVa, Italy ReceiVed: July 7, 1995X
The acid-base equilibria of a series of monofunctional and polyfunctional bases and acids have been investigated with ab initio theoretical methods; the structures and energies of the neutral form and of the forms ionized at all ionization sites have been determined. Predictions concerning the preferred site of ionization are made on the basis of the energetics. (a) For bases of the type NHm-XHn (X ) O, S) the preferred protonation site (by 20-30 kcal/mol) is the nitrogen, but the preference is less marked (∆E ) 7 kcal/mol) for X ) P. (b) Amides deriving from weak acids (carboxylic, thiocarboxylic, cyanic, nitrous, sulfinic, etc.) protonate preferentially at the acid residue rather than at the nitrogen, with ∆E’s between 10 and 30 kcal/ mol. (c) Amides deriving from strong acids (nitric, sulfonic) do not show a marked preference, ∆E for the two tautomeric protonated forms being ca. 2-3 kcal/mol. (d) Carboxylic acids and esters are protonated at the carbonyl oxygen, O-alkyl protonation being less favored by 20 kcal/mol. (e) Dimethyl sulfoxide is O-protonated. N-protonation of some non-carboxylic amides and O-alkyl protonation of carboxylic acids and esters leads to a substantial lengthening of the bond between N or O and the acid residue and to overall structural changes indicating dissociation. The gas-phase energetics are used for the estimation of relative solution basicities by means of the energies of solvation of simple models.
The site of ionization (protonation or deprotonation) of a given base or acid possessing more than one possible site may be difficult to determine experimentally, and if the various ions which can be formed have very different solvation energies, it may also depend strongly on the solvent.1 Some examples are given by (a) amides (both carboxylic1 and sulfenamides,2 sulfinamides,2 sulfonamides,3 phosphoramides,4 nitrosamides,5 etc.) which may undergo protonation at the nitrogen or at the acid residue, (b) nitrogen heterocycles possessing multiple sites which differ little in basicity,1 and (c) hydroxamic acids, which can deprotonate at nitrogen or oxygen.6 Experimental studies usually rely on indirect methods, like changes in NMR chemical shifts, UV or IR spectral changes, etc.1 All these methods are well-suited for quantitative studies but require additional assumptions in order to infer the structure of the species formed. NMR spectra of the ions under slowexchange conditions (e.g., in superacids)7 will often give a conclusive answer, but this method requires forcing conditions which may not be easy to attain and, more importantly, poorly solvating media, which may have an impact on the result if the above conditions concerning solvation hold. Similar considerations apply also to theoretical studies, which normally deal with isolated species; this approach is however profitable, in that it allows to calculate the relative proton affinity of any of the possible species, so that one can estimate whether solvation can alter the ordering in solution. We have recently undertaken an experimental work aimed at the determination of the site of ionization in polyfunctional acids and bases by NMR T1 measurements of the nuclei potentially involved.2,6,8 The most common atoms giving rise to basic and acidic species (nitrogen, oxygen, and sulfur) have an NMR-active isotope possessing a quadrupole moment (14N, 17O, 33S).9 The NMR relaxation of such nuclei is normally dominated by the coupling between the nuclear quadrupole moment and the electric field gradient at the nucleus (efg) arising from the surrounding electron distribution, which is expected X
Abstract published in AdVance ACS Abstracts, December 15, 1995.
0022-3654/96/20100-1536$12.00/0
in principle to be changed by the ionization process. Experimental NMR studies2,6,8,9 have shown that changes in line width indeed take place for the heteronucleus involved in the acidbase process. The best known case is the protonation of an sp3 nitrogen, for which protonation causes a decrease of the nuclear quadrupolar coupling constant χ and hence sharper 14N signals.9 For other cases (ionization at oxygen, sulfur or nitrogen in other chemical environments) the problem has been hardly investigated.10 Therefore, except for the case outlined above, no simple empirical criteria exist for predicting the change of χ upon ionization. The present study has two aims: (1) to establish general criteria for predicting the preferred site of ionization in a variety of polyfunctional compounds, in order to allow a comparison with experimental results and to make predictions for all those cases in which experimental results cannot be obtained or are not easy to interpret, and (2) to establish prediction rules for the change in efg induced by protonation or deprotonation of common quadrupolar nuclei, in order to provide guidelines for the interpretation of NMR spectral changes. The results of a comprehensive theoretical investigation of the proton affinities of a wide range of mono- and polyfunctional nitrogen, oxygen, phosphorus, and sulfur bases and acids are reported herein; efg (and related NMR line width) changes occurring upon protonation or deprotonation are the subject of part 2. Method The calculation of absolute proton affinities (PA) to an accuracy comparable to experimental gas-phase results is outside the scope of this paper; therefore, no attempt has been made to achieve “thermochemical accuracy”, e.g., by including zeropoint vibrational energies or using Gaussian-2 theory.11 In any case, the computational level must be adequate for an accurate estimation of relatiVe proton affinities, where some energy terms (which are necessary for the evaluation of absolute PA’s) often cancel out; absolute PA’s will however be presented for comparison. © 1996 American Chemical Society
Site of Ionization of Polyfunctional Bases and Acids. 1 The structures and proton affinities of some of the species presented herein, especially the simplest, have been calculated previously (see refs 11-17 for a general overview); however, in time a variety of methods have been employed, ranging from HF wave functions with minimal basis sets to correlated wave functions with large basis sets and Gaussian-2 theory. The present work aims rather at providing an extended but fully consistent set of data, evaluated under the same conditions, and hence a compatible set of energies. Thus, for energy calculations the correlation energy correction was estimated via secondorder Møller-Plesset perturbation theory (MP2) in the frozencore approximation.18 Maximum flexibility was allowed for in geometry optimizations, employing split-valence basis sets with all-atom polarization functions and all-atom diffuse functions for anions, and a number of conceivable conformations were examined. Thus, for species containing only first-row elements the 6-311G(d,p) basis set was used for all neutrals and cations and 6-311++G(d,p) for anions, and for sulfur- and phosphorus-containing species the 6-31G(d,p) and 6-31++G(d,p) basis sets were used correspondingly.18 For conciseness, we will denote this alternative as 6-31(1). Proton affinities were calculated from MP2 energies at the optimized geometries, using the same basis set used for optimization. Thus, the calculations can be denoted as MP2/6-31(1)G(d,p)//6-31(1)G(d,p) and MP2/6-31(1)++G(d,p)//6-31(1)++G(d,p), except neutrals giving rise to anions. The geometries of these were optimized with the 6-31(1)G(d,p) basis set, and the energy was calculated at that geometry by adding the diffuse functions (MP2/6-31(1)++G(d,p)//6-31(1)G(d,p)). Atomic charges were calculated by fits to electrostatic potential maps, using a grid of 1 point per atomic unit. All calculations were carried out with the program Spartan,19 running on an IBM RS/6000 workstation. Results A wide variety of inorganic and organic acids and bases have been selected, in order to provide an overview of the various environments in which acid-base processes take place. Because our ultimate goal is a comparison with experimental NMR data in solution, preference has been given, where possible, to species actually amenable to experimental study. In order to understand the effect of substitution at centers adjacent to the site of ionization on the efg, in some cases variously substituted derivatives (e.g., NH3, MeNH2, Me2NH, PhNH2) were compared. The energies of all species as well as absolute and relative proton affinities of the neutrals and anions are reported in Tables 1 and 2, together with available experimental PA’s. Table 1 collects the results for monofunctional species, i.e., species for which only one heteroatom can be reasonably conceived as basic or acidic site. Table 2 collects the same for all tautomeric ions deriving from polyfunctional species, together with their relative energies. Only the data for the most stable conformers of each species are reported; the complete data set is available as supporting information. In those cases where a compound is monofunctional for one process, but polyfunctional for the other (e.g., MeSO2NH2), the entry for the neutral species is repeated in both tables for readability. In the paper, we will report and comment only on the structures of species previously studied at a lower level of theory or for the first time and on those for which a disagreement exists with published results. In the structural drawings we report distances in angstroms and angles in degrees; the extent of nonplanarity, especially of nitrogen centers, is expressed as the angle formed by one of the bonds and the bisector of the angle formed by the other two. The full series of structure drawings,
J. Phys. Chem., Vol. 100, No. 5, 1996 1537 TABLE 1: Energies and Proton Affinities for Monofunctional Compounds PA species
E(HF)
a
E(MP2)a,b
H3O+ e H3O+ f H2O OHMeOH2+ MeOH MeO-
Water and Methanol -76.331 843 7 -76.547 610 1 -76.330 177 5 -76.544 797 2 -76.047 012 0 -76.263 276 2 -75.384 193 0 -75.602 134 4 -115.385 768 9 -115.740 974 4 -115.075 812 0 -115.435 498 8 -114.443 811 9 -114.819 807 3
absc
exptld
178.4 414.9
166.5 390.8
191.7 386.3
181.9 381.2
H3S+ e H3S+ f H2S SHMeSH2+ MeSH MeS-
Hydrogen Sulfide and Methanethiol -398.952 533 1 -399.091 398 3 -398.903 340 3 -399.042 689 5 -398.675 027 7 -398.810 028 3 176.6 -398.111 219 9 -398.241 571 4 356.7 -438.013 287 1 -438.292 844 2 -437.709 027 2 -437.987 639 8 191.5 -437.131 742 4 -437.410 134 2 362.4
170.2 351.3
NH4+ NH3 NH2MeNH3+ MeNH2 MeNHMe2NH2+ Me2NH Me2NPhNH3+ PhNH2 pyridineH+ pyridine NH3NH32+ NH2NH3+ NH2NH2
Amines and Hydrazine -56.558 768 3 -56.754 558 0 -56.210 396 9 -56.408 357 6 -55.541 551 2 -55.755 398 4 -95.608 587 9 -95.949 164 7 -95.242 514 2 -95.586 881 6 -94.572 904 8 -94.934 383 9 -134.655 969 5 -135.144 690 4 -134.278 282 1 -134.771 774 7 -133.617 061 2 -134.132 555 4 -286.156 939 7 -287.155 463 2 -285.798 030 7 -286.798 918 0 -247.125 842 0 -247.966 112 4 -246.749 672 7 -247.598 205 5 -111.662 352 7 -112.029 599 3 -111.563 528 4 -111.931 805 2 -111.211 683 7 -111.582 822 2 Hydrogen Cyanide -93.186 022 2 -93.477 089 7 -92.899 956 3 -93.196 010 6 -92.334 911 8 -92.632 040 8
HCNH+ HCN CNH2COH+ H2CO Me2COH+ Me2CO H2CSH+ H2CS
187.4 356.8
217.2 409.7
204.0 403.7
227.3 409.4
214.1 403.2
234.0 401.1
220.6 396.3
223.7
209.5
230.9
220.9
61.4 219.0
204.7
176.4 353.9
171.4 351.1
Formaldehyde, Acetone, and Thioformaldehyde -114.195 105 3 -114.518 704 5 -113.899 154 0 -114.233 282 0 179.1 -192.345 135 4 -192.966 829 3 -192.013 282 2 -192.646 172 6 201.2 -436.814 655 8 -437.071 572 0 -436.509 953 1 -436.770 592 5 188.9
171.7 192.1 185.028a
HCONH2 HCONH-
Formamide -168.982 285 1 -169.492 432 0 -168.389 992 4 -168.917 661 5
360.7
359.9
NH2CN NHCNNCN2-
Cyanamide -147.952 111 0 -148.421 112 4 -147.387 672 3 -147.859 894 5 -146.552 006 1 -147.050 682 1
352.2 507.8
349.8
MeSO2NH2 MeSO2NH-
Methanesulfonamide -642.402 184 5 -643.221 941 2 -641.836 069 2 -642.679 596 0
340.3
HCOOH HCOO-
Formic Acid -188.820 524 8 -189.348 973 7 -188.257 229 7 -188.800 293 5
344.3
Me2SO2H+ Me2SO2
Dimethyl Sulfone -626.730 422 7 -627.521 359 1 -626.397 334 2 -627.193 736 2
205.6
345.6
a
Energies in hartrees; see text for level of theory. b MP2 energies in the frozen-core approximation. c ∆E at MP2 level, in kcal/mol. d Experimental proton affinity in kcal/mol from ref 45a (cations) and ref 45b (anions), except where indicated. e Nonplanar (C3V) form. f Planar (D3h) form.
containing all geometrical parameters, is collected in the supporting information.
1538 J. Phys. Chem., Vol. 100, No. 5, 1996
Bagno and Scorrano
TABLE 2: Energies and Proton Affinities for Polyfunctional Compounds PA species
E(HF)a
E(MP2)b
absc
exptld
rel PAe
PA species
E(HF)a
Hydroxylamine and Thiohydroxylamine -131.745 394 4 (0.00) NHOH-130.372 691 5 -131.701 737 2 27.39 NH3SH+ -454.021 096 5 -131.416 229 3 206.5 NH2SH2+ -453.985 833 5 -130.796 368 3 389.0 (0.00) NH2SH -453.685 932 2
E(MP2)b
absc
exptld
rel PAe
NH3OH+ NH2OH2+ NH2OH NH2O-
-131.359 176 5 -131.316 157 3 -131.027 720 7 -130.391 610 0
NH3PH2+ NH2PH3+
-397.843 084 2 -398.136 636 0 -397.836 084 7 -398.125 398 5
HC(OH)NH2+ HCONH3+ HCONH2
Formamide and Thioformamide -169.323 341 4 -169.826 300 3 (0.00) HC(SH)NH2+ -169.290 036 2 -169.802 324 0 15.04 HCSNH3+ -168.982 285 1 -169.492 432 0 209.5 198.4 HCSNH2
-491.917 779 9 -492.340 315 8 -491.881 941 2 -492.317 406 0 -491.576 411 5 -492.002 030 7 212.3
HC(NH2)2+ HC(dNH)NH3+
-149.512 885 1 -149.997 793 4 -149.456 658 8 -149.948 789 2
Formamidine (0.00) HC(dNH)NH2 30.75
-149.123 606 2 -149.618 516 7 238.0
NH2CNH+ NH3CN+
-148.273 959 9 -148.738 361 0 -148.230 252 5 -148.702 708 7
Cyanamide (0.00) NH2CN 22.37
-147.952 111 0 -148.421 112 4 199.1
NH3NO+ NH2NOH+ NH2NHO+ NH2NO
-185.203 315 0 -185.206 040 0 -185.172 077 0 -184.881 978 1
Nitrosamides (0.00) Me2NNOH+ 5.93 Me2NHNO+ 15.82 Me2NNO
-263.316 789 4 -264.158 777 8 -263.288 053 4 -264.142 380 1 -262.961 376 5 -263.810 772 9 218.4
(0.00) 10.29
NH2NO2H+ NH3NO2+ NH2NO2
Nitramides -260.019 293 2 -260.773 536 1 182.138 (0.00) Me2NNO2H+ -259.989 692 5 -260.766 404 4 4.47 Me2NHNO2+ -259.715 980 1 -260.471 833 3 189.3 Me2NNO2
-338.119 365 0 -339.117 491 0 -338.103 296 3 -339.173 354 8 -337.790 921 1 -338.848 536 8 206.4
(0.00) 2.59
MeS(OH)NH2+ MeSONH3+
-567.908 842 5 -568.538 947 0 -567.874 597 9 -568.522 089 3
Methanesulfinamide (0.00) MeS(H)ONH2+ 10.58 MeSONH2
-567.838 911 1 -568.475 476 1 -567.544 481 4 -568.182 526 6 223.6
39.83
Methanesulfonamide (0.00) MeSO2NH2 2.79
-642.402 184 6 -643.221 941 2 201.9
-185.756 087 8 -185.746 639 9 -185.730 879 4 -185.428 849 3 205.3
-642.711 598 4 -643.543 774 5 MeSO2NH3+ MeSO2(H)NH2+ -642.724 106 5 -643.539 328 4 HP(NH2)3+ P(NH2)2(NH3)+
-507.990 559 6 -508.617 949 3 -507.978 435 2 -508.612 000 2
-582.897 820 7 -583.700 926 5 (NH2)3POH+ OP(NH2)2(NH3)+ -582.868 269 2 -583.683 355 8
Aminophosphine (0.00) NH2PH2 7.05
-130.779 245 1 -454.325 990 8 -454.290 531 1 -453.989 302 3 211.3
-397.493 768 2 -397.785 305 8 220.5
Phosphorous Acid Triamide (0.00) P(NH2)3 3.73
-507.599 245 5 -508.231 987 6 242.2
Phosphoric Acid Triamide (0.00) OP(NH2)3 11.02
-582.519 541 5 -583.330 174 0 232.6
Formic Acid and Methyl Formate -189.646 384 7 (0.00) HCO(OH)OMe+ -228.171 844 2 -228.837 315 9 -189.616 398 7 18.82 HC(O)OHMe+ -228.132 250 3 -228.805 923 4 -189.615 550 6 19.35 HCOOMe -227.851 295 0 -228.525 559 7 195.6 -189.348 973 7 186.6 178.8
HC(OH)2+ HC(O)OH2+ f HC(O)OH2+ g HCOOH
-189.125 744 9 -189.082 510 0 -189.085 209 3 -188.820 524 8
Me2SOH+ Me2(H)SO+
-551.917 778 4 -552.520 073 4 -551.850 301 3 -552.462 215 7
Dimethyl Sulfoxide (0.00) Me2SO 36.30
10.74 (0.00) 22.25
(0.00) 14.37
194.530
(0.00) 19.70 188.4
-551.547 180 1 -552.158 831 9 226.68 211.3
a,b See Table 1. c ∆E in kcal/mol between base and most stable protonated form. d See Table 1. e Relative calculated proton affinities (∆E for tautomeric ions relative to most stable one) in kcal/mol at MP2 level. f Nonplanar form. g Planar form.
1. Monofunctional Compounds. The results for simple hydrides (H2O,17,20 H2S,17,20c,21 NH320a,b,d,22), their methyl derivatives (MeOH,16,17,20a,c,d MeSH20c,21,23), amines (MeNH2,20d Me2NH,20d PhNH2,24 pyridine25), hydrazine,20a,d HCN,16,20a,d,26 H2CO,16,20a,c,d,27 and H2CS23a,27a,d,28 agreed with published data (sometimes obtained at higher levels), or their structures (PhNH2, pyridine) did not present peculiar features. These species were studied to provide the simplest references for efg calculations. Acetone.17,20c,27b,c,e,29 Protonated acetone presents a subtle problem, in that the nature of the structure than can most easily be conceived for Me2COH+ (which has Cs symmetry like the neutral form) varies according to the method used. Thus, Hartz et al.20c and Lee et al.29 characterized it as a minimum and a transition state, respectively. It has been proposed that the minimum structure (C1 symmetry) has twisted methyl groups,
and the added proton is slightly off-plane from the CCO skeleton.29 Our calculations have confirmed the latter result, although the energy difference with the Cs structure is so small (0.2 kcal/mol) that the characterization of the potential energy surface will probably require more sophisticated methods. Formic Acid and Carboxylic Amides22a (Deprotonation). The neutral acid has two conformers with the OH hydrogen s-cis or s-trans to the carbonyl oxygen. The s-trans form is more stable by 5.3 kcal/mol. Formamide has two nonequivalent acidic hydrogens; the two anionic forms differ by 3.5 kcal/mol in energy in favor of the isomer having trans hydrogens but are otherwise similar. Cyanamide (Deprotonation).22a,30 For the neutral species, both a planar and a pyramidal (nonplanar) arrangement of the NH2 group can be conceived; the pyramidal one is more stable, a feature which is shared by all non-carboxylic amides studied
Site of Ionization of Polyfunctional Bases and Acids. 1 SCHEME 1: Methanesulfonamide
herein (see below). In this case the difference between the two (1.64 kcal/mol) is quite small. NH2CN can undergo two ionizations to yield HNCN- and NCN2-; in the former, the NCN group is slightly nonlinear (∠NCN ) 176.7°) at both the HF and MP230 levels. Methanesulfonamide (Deprotonation).31 The neutral compound has been studied in some detail by Topiol et al. In that study it was found that the amino group may or may not be planar according to the method used, MNDO/3 and 3-21G(*) giving a planar nitrogen and MNDO and STO-3G a nonplanar one. We could demonstrate that with the more flexible 6-31G(d,p) basis set the NH2 group is indeed nonplanar. A notable feature is the fact that the amino hydrogens are eclipsed with the sulfonyl oxygens (Scheme 1). Dimethyl Sulfone (Protonation).23b,32 The conformation of the neutral species in such that one methyl group is staggered with respect to the SO2 group, while the other methyl hydrogens are eclipsed with the sulfonyl oxygens, like in the case of MeSO2NH2. This arrangement is lost in the protonated structure, which has the acidic proton eclipsed with the other SO bond. 2. Polyfunctional Compounds. 2.1. HmN-XHn Compounds (X ) O, S, P). Hydroxylamine20a,22 and Thiohydroxylamine.21,31a,33 Both species have been extensively investigated and can undergo protonation or deprotonation at N or the other heteroatom (O, S). Neutral NH2OH and NH2SH are in an anti conformation; in NH2SH, the HSN angle is quite small (98°) as in H2S. While NH2OH2+ has a straightforward anti structure, NH2SH2+ has a gauche geometry similar to that of the isoelectronic NH2NH2.21 NH2OH2+ and NH2SH2+ are predicted to lie higher in energy than the respective N-protonated forms by 27.4 and 22.2 kcal/mol, respectively. Hydroxylamine is also a bifunctional acid. NH2O- is predicted to be the most stable ion by 10.7 kcal/mol. Aminophosphine.21,22c,34 Neutral NH2PH2 has a gauche structure similar to that of hydrazine. Unlike the two previous cases, the two ions which can arise from protonation differ by only 7 kcal/mol, favoring the N-protonated form PH2NH3+. 2.2. Carboxylic and Non-Carboxylic Amides and Thioamides. Formamide22b,35 and Thioformamide (Protonation).28a For both bases, protonation at X ()O, S) can give rise to two conformers in which the added proton is cis or trans to the nitrogen; in both cases, the trans conformer is more stable,35a although the difference is rather small (3.43 and 0.85 kcal/mol for X ) O and S, respectively). N-protonation causes a marked
J. Phys. Chem., Vol. 100, No. 5, 1996 1539 (0.2 Å) lengthening of the C-N bond35c and a widening (by ca. 8°) of the H-C-X angle which, loosely speaking, suggests viewing HC(X)NH3+ as tending toward a complex between H-CdX+ and NH3; more striking examples will be shown below. Once again, the added proton in HC(SH)NH2+ forms a small angle (95.9°) with the CdS group. As known, Nprotonation is disfavored with respect to O- or S-protonation by a sizable amount of energy (14-15 kcal/mol). Formamidine.36 For neutral HC(NH)NH2, four minima were located on the potential energy surface, two pairs with the )NH hydrogen cis or trans to NH2, which in turn can be planar or not. We found the most stable one to be the trans nonplanar form, although the others differ by less than 3 kcal/mol. These results contrast with those obtained at the HF/3-21G level.36b Protonation can lead either to HC(NH2)2+ or to HC(NH)NH3+; the latter occurs again as two conformers with the )NH hydrogen cis or trans to the NH2 group. The trans conformer is more stable than the cis by 7.5 kcal/mol and is in turn less stable than HC(NH2)2+ by 30.7 kcal/mol. Thus, there is a strong preference for protonation of the imino nitrogen, as could be anticipated on the basis of simple resonance arguments. Cyanamide (Protonation).22b,30 All calculations agree in predicting a nonplanar NH2 group, although the planar form is only slightly higher in energy (by 1.6 kcal/mol). An interesting effect of electron correlation is the one reported by Cacace et al.30 at the MP2(full)/6-31G(d) level, i.e., that for NH2CN, HNCN-, and NH2CNH+ the NCN group is slightly nonlinear; for the latter, the CNH arrangement is also not linear, which is not borne out by HF calculations. The cyano-protonated tautomer (NH2CNH+) is found to be more stable than NH3CN+ by 22.4 kcal/mol. Nitrosamide and N,N-Dimethylnitrosamide.37 These two compounds show a different behavior in most respects. NH2NO has a slightly nonplanar NH2 and in principle possesses three basic sites, i.e., the amino nitrogen, the nitroso nitrogen, and the oxygen. The O-protonated form (NH2NOH+) can exist in cis and trans configurations; the trans is more stable than the cis by no less than 12.3 kcal/mol. This energy difference is the largest found here for conformers of this type; a similar difference, although smaller in magnitude (ca. 6 kcal/mol), was found in protonated primary amides35a and explained as an electrostatic repulsion between the added proton and the positively charged NH2 hydrogens. As was previously reported by Hegarty et al.,37 NH3NO+ exhibits a very long NN distance (2.28 Å), which strongly suggests a weakly bound structure like [NH3‚‚‚NO]+ and was checked to be a minimum by vibrational analysis. In fact, its geometry is quite close to that of isolated NH3 and NO+ molecules. The energies of the three tautomeric species are ranked as NH2N(H)O+ (+15.8) > NH2NOH+ (+5.9) > NH3NO+ (0.0). Thus, in this compound the amino group appears to be the most basic; at the HF level the conclusion is opposite (NH2NOH+ favored by 1.7 kcal).37 However, the peculiar structure of the preferred cation suggests that the parent compound may not be a suitable model of stable alkyl nitrosamides. For this reason, we also investigated Me2NNO and its protonated forms, except the least stable one (protonated at the nitroso N). This time, the basicity order is reversed in favor of the O-protonated form by 10.3 kcal/mol; the NN distance in Me2NHNO+ is again lengthened, albeit by a smaller amount, and the NO distance is similar to that of the neutral. Nitramide38 and N,N-Dimethylnitramide.38a It is interesting to compare our results with those obtained by Cacace et al.38b by MP2(full)/6-31G(d) optimization and Gaussian-1 theory. Both predict NH2NO2 to have a nonplanar amide nitrogen and a very small difference in stability of the two protonated forms
1540 J. Phys. Chem., Vol. 100, No. 5, 1996
Bagno and Scorrano
SCHEME 2: Methanesulfinamide
SCHEME 3: Phosphorous Acid Triamide
NH2NO2H+ and NH3NO2+: 5.2 (Gaussian-1) and 4.5 kcal/mol (this work), in favor of the first. A major discrepancy exists about the structure of NH3NO2+, in that our HF/6-311G(d,p) calculations yield a structure in which the NN bond is 1.523 Å (0.16 Å lengthening) and the ONO angle opens up by only 9°, relative to the neutral. On the contrary, MP2(full)/6-31G(d) optimization yields a weakly bound complex between NH3 and NO2+, as shown by the long NN distance (1.996 Å) and wide ONO angle (146.5°). This circumstance prompted us to also investigate the more stable N,N-dimethyl derivative, much in the same line seen above. The energy difference of the two protonated forms is even closer, the O-protonated form Me2NNO2H+ being again favored by only 2.6 kcal/mol. There is a major effect of electron correlation, in that HF energies predict the same ordering, but with a 10 kcal difference. Like in the case of nitrosamides, N-alkylation diminishes the tendency of the N-protonated form to dissociate; in fact, in Me2NHNO2+ the NN distance (1.482 Å) has increased by 0.13 Å and the ONO angle by 5°. Methanesulfinamide.31a Sulfinamides have three basic sites: N, O, and S. The unknown parent species HSONH2 has been studied by Topiol et al.,31a and the general structural features agree, notably the nonplanar amino group. The three protonated forms are ranked in energy as MeS(H)+(O)NH2 (+39.8) > MeS(O)NH3+ (+10.6) > MeS(OH)+NH2 (0.0); thus the Sprotonated form can be discarded. Once again, N-protonation causes a substantial lengthening of the S-N bond (Scheme 2). Methanesulfonamide.31,39 Its basicity presents features similar to those of nitramides; the energy of the O- and N-protonated forms is quite close, the latter being more stable by just 2.8 kcal/mol (cf. 11.1 with HF/3-21G(*)).31b The conformation of the OH bond in MeS(O)(OH)+NH2 is peculiar, the acidic proton being almost eclipsed with the other SO bond (the trans form is not a stationary point). The S-N distance increases to 1.86 Å in MeSO2NH3+, again suggesting a weaker bonding. Phosphorous Acid Triamide.34d,40 The amino groups in P(NH2)3 are also nonplanar and adopt a conformation in which one of the amino groups is approximately perpendicular to the
plane defined by the other two NH2 hydrogens. Protonation may take place at one of the nitrogens or at the phosphorus atom; the energies of the two species, P(NH2)2(NH3)+ and HP(NH2)3+, are quite similar (∆E ) 3.7 kcal/mol in favor of the latter). The structures of both ionic forms present some peculiarities: (a) in P(NH2)2(NH3)+ the H3N+-P bond is much lengthened (from 1.7 to 2.0 Å), as seen many times now; even considering the longer equilibrium distances for P-N bonds, this structure resembles a diaminophosphenium ion solvated by ammonia (H3N‚‚‚P(NH2)2+), like H3N‚‚‚NO+ seen before. (b) The structure of HP(NH2)3+ is a strongly distorted tetrahedron (Scheme 3). Phosphoric Acid Triamide.41 The neutral species adopts an arrangement of the three amino groups similar to that in P(NH2)3, again with slightly nonplanar nitrogens. O-protonation is more favored than N-protonation by 11.0 kcal/mol; this difference is much lower in magnitude than the 50-100 kcal/ mol calculated at the STO-3G level for H2P(O)NH241a and OP(NH2)3.41b The structure of OP(NH2)2(NH3)+ resembles that of HP(NH2)3+, with a distorted tetrahedral arrangement in which the oxygen atom points toward the NH3+ group and the OPNN moiety becomes almost planar. The H3N+-P bond (1.86 Å) is lengthened by 0.2 Å, and considerations similar to the above ones apply. On the contrary, in the structure of the more stable (NH2)3POH+ ion the overall features of the neutral are preserved (Scheme 4). 2.3. Carboxylic Acids and Esters. Formic Acid20c,42 and Methyl Formate43 (Protonation). CO-protonated formic acid, HC(OH)2+, has several conformers differing by 2-5 kcal/mol, and the most stable one has one OH trans and one cis to the formyl hydrogen; the C-O distances are equal. The formation of HC(OH)2+ is largely favored over that of HC(O)OH2+ (acylO-protonation) by 18.8 kcal/mol. For the latter, two unusual structures were obtained, both of which were found to be minima on the potential energy surface. One is planar with a C-OH2
Site of Ionization of Polyfunctional Bases and Acids. 1
J. Phys. Chem., Vol. 100, No. 5, 1996 1541
SCHEME 4: Phosphoric Acid Triamide
Figure 1. Correlation between calculated and experimental proton affinities for neutrals and anions. The least-squares fit yields: slope, 1.01; intercept, -11.24; correlation coefficient, 0.997.
most stable protonated form, S-protonation leading to a tetrahedral structure less stable by 36.3 kcal/mol (cf. 73 at the MINDO/3 level44c). This should be compared with semiempirical (PM3) results for diaryl sulfoxides, for which the difference is considerably lower (17 kcal/mol).44b Discussion
SCHEME 5: Formic Acid
bond length of 1.606 Å and a substantially increased HCO angle. Another structure, slightly less stable, features an almost linear (174°) H-C-O arrangement, with the OH2 moiety located at a long distance (rCO ) 2.331 Å) in an almost perpendicular fashion. Both structures suggest a weakly bound complex between HCO+ and water (Scheme 5). Similar considerations apply to methyl formate, for which carbonyl protonation is favored over acyl-O-protonation by 19.7 kcal/mol. The structure of HC(O)OHMe+ does not show the peculiar features seen above, although a lengthening of the C-OMe distance by 0.15 Å can again be noted. 2.4. Sulfoxides. Dimethyl Sulfoxide.23b,32,44 This base has two protonation sites, i.e., O and S. In the first case two conformers have been considered, with the acidic proton syn or anti to the methyl groups, the syn form being less stable than the anti by 3.3 kcal/mol.44a,c The anti form is also the
General Considerations. A point of general interest is the relative performance of HF and MP2 methods in this type of calculation. From a comparison of the ∆E values at the HF and MP2 levels, one can see that in most cases the stability ordering of the various ionic species is not affected by correlation, although the actual value is sometimes appreciably different. (For the ions derived from HCONH2, HCSNH2, HCOOH, Me2NNO, NH2NO2, Me2NNO2, MeSONH2, Me2SO, and PO(NH2)3 ∆E’s at the two levels differ by more than 5 kcal/mol.) The relative energies of conformers are very little affected. However, for ionic species which differ little in energy the inclusion of correlation may reverse the stability order; thus, while HF theory predicts NH2NOH+ to be more stable than NH3NO+ by 1.7 kcal/mol and MeSO2HNH2+ to be more stable than MeSO2NH3+ by 7.8 kcal/mol, at the MP2 level the stabilities are reversed and differ by 5.9 and 2.8 kcal/mol, respectively. These data indicate that, by and large, only MP2 energies should be employed for such comparisons, especially if HF energies differ by less than ca. 5 kcal/mol. Experimental proton affinities45 are plotted against the values calculated in this work for the most basic site in Figure 1. [Szulejko and McMahon have recently reevaluated some gasphase basicities, improving the quality of the absolute scale.17 However, these new data are available only for a few of the species studied in this work (eight of a total of 48); the changes introduced are small (less than 1 kcal/mol on average) and fall below the accuracy of this correlation; therefore, in order to ensure maximum homogeneity of treatment, we have relied on available reference data.] The least-squares fit of the data yields a slope of 1.01 and an intercept of -11.2 kcal/mol with r2 ) 0.997. Thus, the calculations at the level employed herein consistently overestimate proton affinities by 11.2 kcal/mol. However, the fact that an excellent correlation with almost unit slope extends over 250 kcal/mol is certainly remarkable and gives confidence on the overall quality of MP2 calculations. Furthermore, if PA’s relative to NH3 are employed for the correlation, the slope (1.01) is unchanged, and the intercept is reduced to +4.5 kcal/mol; i.e., calculated PA’s underestimate experimental values, however by a smaller amount. Therefore,
1542 J. Phys. Chem., Vol. 100, No. 5, 1996
Bagno and Scorrano
TABLE 3: Atomic Charges on Amide Nitrogen and Bond Lengths with the Acid Residue (R) in Amides and Their Changes following Protonation base
qN(B)
qN(BH+)
rR-N(B)
rR-N(BH+)
∆q
∆r
Me2NH NH2CN NH2NO2 HCSNH2 HCONH2 OP(NH2)3 MeSO2NH2 MeSONH2 Me2NNO P(NH2)3 NH2NO
-0.74 -0.81 -0.74 -0.54 -0.93 -1.01 (av) -1.02 -0.90 +0.36 -0.68 (av) -0.35
-0.02 -0.45 -0.53 +0.02 -0.42 -0.50 -0.71 -0.53 +0.36 -0.53 -0.83
1.448a 1.340 1.358 1.323 1.350 1.660 1.646 1.679 1.298 1.704 1.315
1.496a 1.397 1.523 1.491 1.526 1.862 1.861 1.927 1.586 2.000 2.277
+0.72 +0.36 +0.21 +0.55 +0.51 +0.51 +0.30 +0.36 0.00 +0.15 -0.48
0.048 0.057 0.165 0.168 0.176 0.202 0.215 0.248 0.288 0.296 0.962
a
Methyl-N distance.
even MP2 energies obtained with a moderately large basis set, but without vibrational energy correlations, can serve for the estimation of unavailable proton affinities.18 Structural Considerations. First of all, it should be remarked that in all amides except carboxylic ones the nitrogen atom is not planar. In carboxylic amides, planarity is commonly thought to arise from the favorable resonance interaction with the carbonyl group that can thus ensue; it is therefore straightforward to acknowledge that when such interaction is poorer (i.e., with an acid residue based on second-row elements like S or P) planarity is no longer a requirement. However, this behavior persists even in amides whose acid residue is based on nitrogen or carbon, like amidines, cyanamides, nitrosamides, nitramides, and even N-hydroxycarboxamides (hydroxamic acids).6 In the cases studied, the energy difference between planar and nonplanar forms is, however, rather small (1-3 kcal/ mol) so that further investigation, perhaps employing higher level methods, is warranted; one might also take advantage of the criteria employed by Wiberg et al.46 to analyze the origin of resonance in amides and esters. A general structural change caused by ionization is that the distance between the ionized atom and the rest of the molecule increases upon protonation and decreases upon deprotonation, which is well documented in the literature for many simple species. For N-protonated noncarboxylic amides, the length of the bond between N and the acid residue may reach values which suggest a weakly bound ion‚‚‚NH3 complex, especially in the cases of NH3NO+ and P(NH2)2(NH3)+. These structures may be even more stable than those protonated at the acid residue; however, the case of NH2NO shows that alkylation at nitrogen reverses the stability balance and leads to a shorter N-acid residue bond, so care should be exercised when dealing with N-alkylated amides, which are often better known experimentally. The above picture holds also for carboxylic acids and esters, but in this case O-alkyl protonation is much more disfavored. An analysis of the atomic charges on nitrogen (Table 3) provides further evidence for this picture; the reference term is the protonation of an amine, for which a large fraction of the overall charge is localized on the nitrogen, as one would intuitively guess. Conversely, for amides featuring a major lengthening of the N-acid residue bond the charge on nitrogen increases only slightly upon protonation (for NH3NO+ it actually decreases); accordingly, there is an inverse correlation between the change in atomic charge on nitrogen (∆q) and acid residuenitrogen bond lengths (∆r). Absolute and Relative Basicities. On the basis of the results obtained in this work, it is possible to give guidelines regarding the overall basic strength and the most basic or acidic site in
the species studied, which span a variety of functional groups. Of course, these criteria as such refer to gas-phase, isolated species; however, they provide a basis for their extension to the solution phase. We will consider a given ionization site as exclusive when the relative proton affinity exceeds ca. 5 kcal/ mol (see below). (1) Protonation of Tricoordinated Nitrogen in Amine-like Molecules. Absolute PA’s for nitrogen protonation decrease as the group bound to the nitrogen becomes more electronegative: CH3NH2 > PH2NH2 ≈ NH2NH2 > NH2SH > NH2OH. However, the heteroatoms (O, S, P) are not the preferred protonation site, although for the last one the balance is close. Thus, simple amine-like compounds behave as nitrogen bases. (2) Protonation of Amides. The intrinsic structural diversity of these compounds calls for some caution in their comparison. With the exception of cyanamide, absolute PA’s (with respect to the most stable protonated form) lie above 200 kcal/mol; MeSO2NH2 and Me2NNO2 exhibit the lowest values, while MeSONH2 and phosphorus-containing amides are the most basic, owing in part to their larger molecular size. The basicity of the amides studied herein can thus be ranked as (most basic atom in boldface) P(NH2)3 > HC(NH)NH2 > OP(NH2)3 > MeSONH2 > Me2NNO > HCSNH2 > HCONH2 > Me2NNO2 > MeSO2NH2 > NH2CN. Protonation at one of the atoms in the acid residue (generally oxygen) is favored over nitrogen by 10-30 kcal/mol in amides deriving from weak acids like (thio)carboxylic, cyanic, nitrous, sulfinic, etc. Conversely, for amides deriving from strong acids (nitric, sulfonic) relative PA’s lie between 3 and 4 kcal/mol, and the preference is not well defined; for these amides, the protonation site in solution may be strongly affected by solvation and may not be exclusive. An exception to this rule is NH2NO, which derives from a weak acid but should protonate at NH2; however, the peculiar (dissociated) structure of NH3NO+ makes Me2NNO a better example, which should match more closely the behavior of stable nitrosamides in solution. Indeed, Me2NNO conforms to the above rule (preferred O-protonation). The classification becomes blurred for phosphorus amides, which contain three amino groups. However, the overall picture is consistent with a diminished availability of the nitrogen lone pair due to interactions with the acyl group, except where its nature strongly withdraws electrons from its basic atom. The strong preference for protonation at the acid residue in formamidine, whose basic sites are nitrogens, is due to the formation of the symmetric species HC(NH2)2+. At the other extreme, in NH2CN CN-protonation remains distinctly favored over NH2protonation, despite the high electronegativity of the sp nitrogen, albeit at the cost of a very low absolute basicity. (3) Protonation of Carboxylic Acids and Esters. These are less basic than the corresponding carboxylic amides by 10-20 kcal/mol and present a similar but enhanced picture: O-alkyl protonation is even more disfavored and essentially leads to dissociation, consistent with the lower basicity of oxygen. (4) Dimethyl Sulfoxide. The protonation of Me2SO takes place exclusively at oxygen, consistent with the engagement of the lone pair on sulfur in bonding with the oxygen atom. The higher proton affinity of Me2SO relative to that of Me2CO (by 20 kcal/mol) is probably due to polarizability effects. The solvation of protonated dimethyl sulfoxide presents some anomalies which were tentatively ascribed either to an anomalous solvation of neutral Me2SO or to gas-phase S-protonation.47 In view of these results, it is apparent that O-protonation is the only process occurring in both gas and solution phases; therefore, the very exoergonic solvation of neutral Me2SO contributes to the solvent effect on the ionization equilibrium
Site of Ionization of Polyfunctional Bases and Acids. 1 to an extent comparable to that of the protonated species, unlike most other bases.47 (5) Basicity of Anions. The acidity ordering of simple monofunctional N- and O-acids is MeSO2NH2 > HCOOH > NH2CN > HCONH2 > MeOH > MeNH2. A noteworthy feature is the fact that methanesulfonamide is a stronger acid than formic acid. The gas-phase acidity ordering (∆PA ) 4 kcal/mol) is reversed in water, where MeSO2NH2 (pKa ) 10.848) is a much weaker acid than HCOOH (pKa ) 3.751) and ∆∆G amounts to -9.6 kcal/mol. This should largely reflect the stronger solvation of the HCOO- oxyanion relative to the MeSO2NH- nitranion, although other data are needed to prove this statement. The only multiple deprotonation process considered here is that of NH2OH. MeOH is a stronger acid than MeNH2 by 23 kcal/mol; the strong preference for O-deprotonation found for NH2OH (formation of NH2O- favored over NHOH- by 10.7 kcal/mol) is in line with this trend. The adjacent electronegative oxygen has, however, an acid-strengthening effect, NH2OH being a stronger acid than MeNH2 by 20 kcal/mol. Solvent Effects on Relative Basicities and Acidities. It is well recognized that when one compares theoretical (obtained for isolated molecules) and experimental (obtained in solution) data, solvation effects must be taken into account. Taking a proton transfer of the type B + H+ f BH+ as an example, the relationship between the thermodynamics of proton transfer in the gas phase and in water is given by a Born-Haber cycle; this is expressed in eq 1 for free energies but of course holds for any other thermodynamic function:
∆G(g) ) ∆Gaq(B) + ∆Gaq(H+) + ∆G(aq) - ∆Gaq(BH+) (1) In eq 1, ∆G(g) and ∆G(aq) denote the free energies of ionization in the gas phase and water, respectively, and ∆Gaq(i) is the free energy of hydration of species i. If two protonated forms (BH+ and BH+′) can derive from the same base B, the difference between gas and aqueous basicities will be given by eq 2:
(∆G(aq) - ∆G(aq)′) - (∆G(g) - ∆G(g)′) ) ∆Gaq(BH+) ∆Gaq(BH+)′ (2) i.e., the difference between gas and aqueous basicities is entirely due to the differential solvation of the ions formed, with no contribution from the neutral species. (Similar arguments hold for two anions originating from the same neutral acid.) Thus, if one of the ions is favored energetically but disfavored by solvation, the basicity order in solution may be reversed, depending on the relative magnitude of the appropriate terms in eq 2. However, the right-hand side must be known. This constitutes a major obstacle, because solvation energies of charged species are determined through a combination of various experimental data (eq 1) which are difficult to obtain; this is even more true when one of the species of interest (the least stable tautomer) is elusive. Some predictions can be made if suitable models are available, as follows. Absolute free energies of hydration of some simple onium ions are experimentally available.35a,49 When comparing these data, one should account for the fact that solvation energies also depend on “physical” factors like molecular size and hence compare only ions with the same overall size, e.g., sharing a similar carbon backbone and substitution pattern.49a We will employ these data (expressed in kcal/mol) to estimate the solvent effect on the acid-base tautomeric equilibria of NH2OH, NH2SH, HCONH2, and MeSONH2. (a) The ∆Gaq of NH3OH+ and NH2OH2+ can be modeled by those of MeNH3+ (-69) and
J. Phys. Chem., Vol. 100, No. 5, 1996 1543 MeOH2+ (-83), respectively. Thus, solvation favors Oprotonation by 14 kcal/mol which, however, does not overcome the intrinsic basicity difference of 27 kcal/mol in favor of N-protonation. (b) By the same token, we can employ Me2SH+ (∆Gaq ) -55) as a model for NH2SH2+ and Me2NH2+ (-62) for NH3SH+; in this case the least stable ion (NH2SH2+, by 22 kcal/mol) is also disfavored by solvation, so the difference in solution is increased to 36 kcal/mol. (c) HCONH3+ is predicted to be less stable than HC(OH)NH2+ by 15 kcal/mol in the gas phase. The experimental ∆Gaq for MeC(OH)NH2+ is -66,35a which should be compared to MeCONH3+ or to its model, EtNH3+ (-66). Thus, the similar solvation of the two tautomers will not alter the balance given by gas-phase basicities. (d) O-protonated MeSONH2 (MeS(OH)NH2+) can be modeled by Me2SOH+ and MeSONH3+ by EtNH3+; for these, ∆Gaq are -60 and -66, respectively. Again, the more exoergonic solvation of the latter (by 6 kcal/mol) is not enough to overcome the greater intrinsic basicity of the former (by 11 kcal/mol), although the difference is reduced to just 5 kcal/mol. The accuracy of currently available data, joined to that of the above rough model, leads to a value of ca. 5 kcal/mol as the limit below which gas-phase differences may be reversed by solvation. Unfortunately, comparisons spanning a broader scope cannot be made, owing to the lack of solvation energies for most other onium ions and especially of anions. Considerable advances have been made in the field of theoretical calculations which model the solvent as a continuum,50 but such calculations still rely on experimental values for their assessment. These circumstances render experimental methods still a key step to obtain relative basicities in solution. Conclusions We have theoretically investigated the ionization equilibria of a variety of mono- and polyfunctional acids and bases, aiming at the determination of their absolute acidity or basicity, and of the relative stability of the various possible ionic species. We have shown that the favored ionization site in the gas phase is often borne out by the relative energies of the isolated species; in such cases, it is straightforward to transfer these results to the solution phase. There are instances, however, in which relative energies are so small that solvent effects may reverse the stability order; such cases are quite useful for understanding solvation phenomena but require experimental investigation by appropriate methods. Supporting Information Available: Energies, structures, and electric field gradients of all species and conformers studied (48 pages). Ordering information is given on any current masthead page. References and Notes (1) Stewart, R. The Proton: Applications to Organic Chemistry; Academic Press: New York, 1985; Chapter 5. (2) Bagno, A.; Eustace, S. J.; Johansson, L.; Scorrano, G. J. Org. Chem. 1994, 59, 232. (3) Laughlin, R. G. J. Am. Chem. Soc. 1967, 89, 4268. (4) Bagno, A.; Scorrano, G. J. Am. Chem. Soc. 1988, 110, 4577. (5) Keefer, L. K.; Hrabie, J. A.; Hilton, B. D.; Wilbur, D. J. Am. Chem. Soc. 1988, 110, 7459. (6) Bagno, A.; Comuzzi, C.; Scorrano, G. J. Am. Chem. Soc. 1994, 116, 916 and references therein. (7) Olah, G. A.; Prakash, G. K. S.; Sommer, J. Superacids; WileyInterscience: New York, 1985. (8) Bagno, A.; Comuzzi, C.; Scorrano, G. J. Chem. Soc., Perkin Trans. 2 1993, 283. (9) Multinuclear NMR; Mason, J., Ed.; Plenum Press: London, 1987. (10) Ali, A. A. M.; Harris, R. K.; Belton, P. S. Magn. Reson. Chem. 1990, 28, 318. (11) Smith, B. J.; Radom, L. J. Am. Chem. Soc. 1993, 115, 4885.
1544 J. Phys. Chem., Vol. 100, No. 5, 1996 (12) Hopkinson, A. C.; Holbrook, N. K.; Yates, K.; Csizmadia, I. G. J. Chem. Phys. 1968, 49, 3596. (13) Kollman, P.; Rothenberg, S. J. Am. Chem. Soc. 1977, 99, 1333. (14) Del Bene, J. E. Chem. Phys. Lett. 1978, 55, 235. (15) Del Bene, J. E. Chem. Phys. Lett. 1983, 94, 213. (16) Del Bene, J. E. J. Comput. Chem. 1985, 6, 296. (17) Szulejko, J. E.; McMahon, T. B. J. Am. Chem. Soc. 1993, 115, 7839 and references therein. (18) Hehre, W. J.; Radom, L.; Schleyer, P. v. R.; Pople, J. A. Ab Initio Molecular Orbital Theory; Wiley-Interscience: New York, 1986. (19) Wavefunction Inc., Irvine, CA. (20) (a) Del Bene, J. E.; Frisch, M. J.; Raghavachari, K.; Pople, J. A. J. Phys. Chem. 1982, 86, 1529. (b) Del Bene, J. E. J. Phys. Chem. 1993, 97, 107. (c) Hartz, N.; Rasul, G.; Olah, G. A. J. Am. Chem. Soc. 1993, 115, 1277. (d) Ozment, J. L.; Schmiedekamp, A. M. Int. J. Quantum Chem. 1992, 43, 783. (21) Magnusson, E. J. Comput. Chem. 1984, 5, 612. (22) (a) Hopkinson, A. C.; Lien, M. H. Int. J. Quantum Chem. 1978, 13, 349. (b) Hopkinson, A. C.; Csizmadia, I. G. Theor. Chim. Acta 1974, 34, 93. (c) Magnusson, E. Tetrahedron 1985, 41, 5235. (23) (a) Curtiss, L. A.; Nobes, R. H.; Pople, J. A.; Radom, L. J. Chem. Phys. 1992, 97, 6766. (b) Wolfe, S.; Stolow, A.; LaJohn, L. A. Tetrahedron Lett. 1983, 24, 4071. (24) Catala´n, J.; Ya´n˜ez, M. J. Chem. Soc., Perkin Trans. 2 1979, 741. (25) (a) Catala´n, J.; de Paz, J. L. G.; Ya´n˜ez, M.; Claramunt, R. M.; Lo´pez, C.; Elguero, J.; Anvia, F.; Quian, J. H.; Taagepera, M.; Taft, R. W. J. Am. Chem. Soc. 1990, 112, 1303. (b) Ferna´ndez Sanz, J.; Anguiano, J.; Villarasa, J. J. Comput. Chem. 1988, 9, 784. (c) Voets, R.; Franc¸ ois, J. P.; Martin, J. M. L.; Mullens, J.; Yperman, J.; van Poucke, L. C. J. Comput. Chem. 1989, 10, 449. (26) Marriott, S.; Topsom, R. D.; Lebrilla, C. B.; Koppel, I.; Mishima, M.; Taft, R. W. J. Mol. Struct. (THEOCHEM) 1986, 30, 133. (27) (a) Bernardi, F.; Ciszmadia, I. G.; Schlegel, H. B.; Wolfe, S. Can. J. Chem. 1975, 53, 1144. (b) Del Bene, J. E.; Vaccaro, A. J. Am. Chem. Soc. 1976, 98, 7526. (c) Del Bene, J. E.; Radovick, S. J. Am. Chem. Soc. 1978, 100, 6936. (d) Speers, P.; Laidig, K. E. J. Chem. Soc., Perkin Trans. 2 1994, 799. (e) Wiberg, K. B.; Marquez, M.; Castejon, H. J. Org. Chem. 1994, 59, 6817. (28) (a) Abboud, J. L. M.; Mo´, O.; de Paz, J. L. G.; Yan˜ez, M.; Esseffar, M.; Bouab, W.; El-Mouhtadi, M.; Mokhlisse, R.; Ballesteros, E.; Herreros, M.; Homan, H.; Lopez-Mardomingo, C.; Notario, R. J. Am. Chem. Soc. 1993, 115, 12468. (b) Grein, F. Can. J. Chem. 1984, 62, 253. (c) Pope, S. A.; Hillier, I. H.; Guest, M. F. Chem. Phys. Lett. 1984, 104, 191. (29) Lee, E. P. F.; Dyke, J. M. J. Chem. Soc., Faraday Trans. 1992, 88, 2111. (30) Cacace, F.; De Petris, G.; Grandinetti, F.; Occhiucci, G. J. Phys. Chem. 1993, 97, 4239. (31) (a) Sabio, M.; Topiol, S. J. Mol. Struct. (THEOCHEM) 1990, 65, 335. (b) Topiol, S.; Sabio, M.; Erhardt, P. W. J. Chem. Soc., Perkin Trans. 2 1988, 437. (32) Wiberg, K. B.; Castejon, H. J. Am. Chem. Soc. 1994, 116, 10489. (33) (a) Bernardi, F.; Bottoni, A.; Mangini, A.; Tonachini, G. J. Mol. Struct. (THEOCHEM) 1981, 3, 163. (b) Gimarc, B. M.; Warren, D. S. Inorg. Chem. 1991, 30, 3276. (c) Magnusson, E. Aust. J. Chem. 1986, 39, 735.
Bagno and Scorrano (34) (a) Gonbeau, D.; Pfister-Guillouzo, G.; Mazie`res, M. R.; Sanchez, M. Can. J. Chem. 1985, 63, 3242. (b) Korkin, A. A.; Mebel’, A. M.; Tsvetkov, E. N. J. Gen. Chem. USSR (Engl. Transl.) 1988, 58, 900. (c) Korkin, A. A.; Tsvetkov, E. N. Bull. Soc. Chim. Fr. 1988, 335. (d) Panina, N. S.; Yakovlev, V. N.; Kukushin, Yu. N. J. Struct. Chem. USSR (Engl. Transl.) 1984, 25, 682. (e) Korkin, A. A.; Tsvetkov, E. N. Russ. J. Inorg. Chem. (Engl. Transl.) 1989, 34, 161. (35) (a) Bagno, A.; Lovato, G.; Scorrano, G. J. Chem. Soc., Perkin Trans. 2 1993, 1091. (b) Hopkinson, A. C.; Csizmadia, I. G. Can. J. Chem. 1973, 51, 1432. (c) Scheiner, S.; Wang, L. J. Am. Chem. Soc. 1993, 115, 1958. (36) (a) Del Bene, J. E. J. Comput. Chem. 1984, 5, 381. (b) Oszczapowicz, J.; Regelmann, C. U.; Ha¨felinger, G. J. Chem. Soc., Perkin Trans. 2 1990, 1551. (c) Zielinski, T. J.; Peterson, M. R.; Csizmadia, I. G. J. Comput. Chem. 1982, 3, 62. (d) Tortajada, J.; Leon, E.; Luna, A.; Mo´, O.; Ya´n˜ez, M. J. Phys. Chem. 1994, 98, 12919. (37) Nguyen, M. T.; Hegarty, A. F. J. Chem. Soc., Perkin Trans. 2 1987, 345. (38) (a) Politzer, P.; Sukumar, N.; Jayasuriya, K.; Ranganathan, S. J. Am. Chem. Soc. 1988, 110, 3425. (b) Attina`, M.; Cacace, F.; Ciliberto, E.; de Petris, G.; Grandinetti, F.; Pepi, F.; Ricci, A. J. Am. Chem. Soc. 1993, 115, 12398. (39) Belvisi, L.; Carugo, O.; Poli, G. J. Mol. Struct. (THEOCHEM) 1994, 318, 189 and references therein. (40) Magnusson, E. Aust. J. Chem. 1985, 38, 23. (41) (a) Modro, T. A.; Liauw, W. G.; Peterson, M. R.; Csizmadia, I. G. J. Chem. Soc., Perkin Trans. 2 1979, 1432. (b) Bollinger, J. C.; Houriet, R.; Kern, C. W.; Perret, D.; Weber, J.; Yvernault, T. J. Am. Chem. Soc. 1985, 107, 5352. (42) (a) Hillenbrand, E. A.; Scheiner, S. J. Am. Chem. Soc. 1986, 108, 7178. (b) Hopkinson, A. C.; Yates, K.; Csizmadia, I. G. J. Chem. Phys. 1970, 52, 1784. (43) Remko, M. J. Mol. Struct. (THEOCHEM) 1989, 60, 287. (44) (a) Akagi, K.; Kobayashi, K.; Yamabe, T. J. Chem. Soc., Perkin Trans. 2 1980, 1652. (b) Laali, K. K.; Houser, J. J. J. Phys. Org. Chem. 1992, 5, 244. (c) Olah, G. A.; Donovan, D. J.; Lin, H. C.; Mayr, H.; Andreozzi, P.; Klopman, G. J. Org. Chem. 1978, 43, 2268. (45) (a) Lias, S. G.; Liebman, J. F.; Levin, R. D. J. Phys. Chem. Ref. Data 1984, 13, 695 and updates. (b) Lias, S. G.; Bartmess, J. E.; Liebman, J. F.; Holmes, J. L.; Levin, R. D.; Mallard, W. G. J. Phys. Chem. Ref. Data 1988, 17 (Suppl. 1). (46) Wiberg, K. B.; Laidig, K. E. J. Am. Chem. Soc. 1987, 109, 5935. (47) Bagno, A.; Lucchini, V.; Scorrano, G. J. Phys. Chem. 1991, 95, 345. (48) Trepka, R. D.; Harrington, J. K.; Belisle, J. W. J. Org. Chem. 1974, 39, 1094. (49) (a) Taft, R. W.; Wolf, J. F.; Beauchamp, J. L.; Scorrano, G.; Arnett, E. M. J. Am. Chem. Soc. 1978, 100, 1240. (b) Taft, R. W. Prog. Phys. Org. Chem. 1983, 14, 247. (50) Tomasi, J.; Persico, M. Chem. ReV. 1994, 94, 2027.
JP9518419
