- Email: [email protected]
Effect of hydration on the acidities of substituted acetic acids
Journal of Molecular Structure (Theo&m), 188 (1989) 45-54 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
EFFECT OF HYDRATION ON THE ACIDITIES OF SUBSTITUTED ACETIC ACIDS
CA1 JINFENG and RONALD D. TOPSOM Department of Chemistry, La Trobe University, Bundoora, 3083 (Australia) (Received 26 August 1988)
ABSTRACT
We report calculations at the ab-initio level on the hydration equilibria of substituted acetates XCH&O; + nH,O * XCH$O; with nH,O up to n = 3. Results are also given for the hydration energies of the corresponding acetic acids. These results allow the calculation of the acidities of the acetic acids in the presence of a limited number of water molecules. Such results are compared to experimental figures in the gas phase and in aqueous solution. Results for three water molecules per acetate ion and one per acetic acid molecule approach those found for aqueous solution.
INTRODUCTION
Most reactions are observed in solution, particularly in water, rather than in the gas phase. Thus, the effect of water molecules on the properties and reactions of organic molecules is of great importance. Much work has been done in recent years [l] on the interaction of specific molecules or ions with a discrete number of water molecules. Experimental studies, involving such techniques as pulsed electron beam mass spectrometry and ion cyclotron resonance, have included studies on the hydration of substituted ammonium ions with up to seven molecules of water [ 2,3], of pyridines and pyridinium ions [ 2,4-61, ROH$ ions [ 2,7,8], phenols [4] and anilines [9]. Theoretical studies include the hydration of substituted ammonium ions [3,10], pyridines and pyridinium ions [ 4,6], phenols, and phenolates [ 111 and various anions [ 121. Less work has been done on the effect of limited hydration on the equilibria of series of organic molecules. Results available include a study [6] of the proton affinity of substituted pyridines both in the gas phase [ eqn. (1) ] and with a single molecule of water [ eqn. (2) 1.
0166-1280/89/$03.50
0 1989 Elsevier Science Publishers B.V.
Experimental [4 ] and theoretical [6 ] (ab-initio STO-3G) results are available for both series and they are in reasonable agreement [ 1 ] (although the number of substituents was limited in equilibrium 2). Another example [ 3,131 is the proton transfer reaction of some substituted amines RNH,+ -zH, 0 + CH3 NH, .yH, 0 = CH3NH3+ .~cH~O+RNH~.~H~O
(3)
Here, the theoretical calculations reproduce the aqueous-phase figures for x = 3, y=2 where the water molecules lie along the NH axes. By contrast, the addition of several water molecules to the proton-transfer reactions of pyridinium ions [6] and phenols [4,11], goes only part of the way towards results found experimentally in aqueous solution. No explanation has been advanced for this difference. We have recently shown [14] that ab-initio STO-3G calculations give results for the relative acidities of substituted acetic acids [eqn. (4) ] in good agreement with experimental gas-phase results XCHz CO, H + CHBCO, = XCH, CO, + CHBCO2H
(4)
Results are also available for the acidities in aqueous solution and thus the series provides an excellent testing ground to investigate the effects of limited numbers of water molecules. CALCULATIONS
Calculations were performed at the ab-initio STO-3G basis using the GAUSSIAN-82 series of programs [ 151. The geometries of the acetic acids and acetates were fully optimized for the various possible conformers [ 141. The positions and structures of various interacting water molecules were then obtained
47
by geometry optimization (using the structures for the various acetate conformers ). Energies calculated for XCH&O; and its most stable mono-, diand tri-hy~ates are given in Table 1. In Table 2 we list the energies for molecules XCH,CO,H and their monohydrates. Previous work on organic equilibria has shown that calculations at the STO3G basis level give good agreement with experiment, for example, for the proton affinities of pyridine [ 161 and their monohydrates [ 1,4,6], cyanides [ 171, TABLE 1 Molecular orbital calculations ( -E) for XCH,C02 (H,O), at the STO-3G//STO-3G (the CX bonds are coplanar with the CO, plane, unless otherwise indicated)
basis level
E (hartree)
X
H Me Et Ph F Cl OMe CF3 CN
?t=O
n=l
n=2
n=3
224.04833 262.62979 301.21084 450.80853” 321.50457 678.67425 336.46345 555.03356 314.62462”
299.05711 337.63780 376.21846 525.81181” 396.51044 753.07535 411.46949 630.03663 389.62429”
374.05547 412.63547 451.21592 600.80788” 471.50655 828.06712 486.46571 705.03092 464.61648”
449.04742 481.6268? 526.20738’ 675.79866’ 546.495~ 903.05324” 561.45567” 780.01593” 539.60370”
eSubstituent at 90” to carboxylate plane. TABLE 2 Molecular orbital calculations ( -E) STO-3G basis level X
for XCHBCOpH and their monohydrates at the STO-3G//
E (hartree) XCH,CO,H
XCH&OeH(H,O)“
XCH,CO,H ( Hz0 )
224.81020 263.39050 301.97080 451.55678 322.25184 678.80271 337.21638 555.77179 315.35031
299.78960
299.78305 338.36256 376.94287 526.52843 397.22328 753.77410 412.18885 630.74196 390.31959
b
H Me Et Ph F Cl OMe CF, CN “-0-H..
*OHp.">C=O* *.HOH.
338.36975 376.94~ 526.53573 397.23261 753.78438 412.19566 630.75294 390.33242
substituted methylamines (the ab-initio STO-3G basis overall gives [ 13,181 as good agreement with experiment as split-valence basis calculation), and anilines [ 161 and for acidities of acetic acids [ 141, benzoic acids [ 19,207, 4substituted f 2.2.21bicycle-octyl carboxylic acids [ 203, and phenols 1211. RESULTS AND DISCUSSION
We first discuss the structure of the hydrated acetate ions and acetic acids. Structure of acetates In our previous paper, we calculated energies for optimized structures of the acetates having the CX bond coplanar with the COz group, as shown in structure 1.
O
*,
H
P
Hgg-c
xp$-“; 4
0
H
H
2
1
We have found, on a more intensive investigation, that this is indeed the most stable or equal most stable conformer for X = H, F, Cl, Me, Et, Me0 and CF3. For X = CN and Ph, conformer 2, with the XC bond at right angles to the COz plane, is more stable by 0.5 and 0.9 kcal mol-‘, respectively. St~~tures with the CH,X groups slightly rotated (ea. 30” ), compared to the CO, plane do lie in an energy well and are very close in stability to 1 for X=H, Me, Et, OMe and CF3. There is not such a well for X = F, Cl, CN and Ph. Str~eture of acetate hydrates We have investigated the interaction of one, two and three molecules of water with the acetates. For the monohydrates, the most stable structure found for the acetate component has the same conformation as found for the free acetate molecule. The preferred position of the water molecule in each case leads to the general structure 3. The process [eqn. (5) ] XCHzCO;
+H,O
* XCH,CO,
-Hz0
(5)
has energy here of - 26.9 kcal mol-l for X = H. The preference for structure 3 compared to structures such as 4 may reflect a weak attraction between the non-hy~ogen-bonded carboxylate oxygen and the far hydrogen on the water molecule.
49
b” ! i
O---H-O XCH,-C
/
‘H
-
XCH, -c-
'0 3
4
/O '0
The next most stable structure, 4, has, for X=H, a 1.6 kcal mol-’ lower calculated value of -LIE” for process (5). (Structure 5 lies, at best, in a very shallow energy well.) O---H
XCH,--C
’- ‘0 \ O---H / 5
The energy values for equilibrium (5) are given in Table 3 for other substituents. With the dihydrates, again the preferred structure [21] of the acetate part was as in the parent, that is, with the CX bond coplanar with the CO, group, except for X = CN and Ph. The most stable structure found for the dihydrates, in each case, was 6 (except that for X = CN and Ph, the XC bond was at right angles to the CO2 plane). TABLE 3 Calculated STO-3G XCH$O, (H,O),
X
H Me Et Ph F Cl OMe CF, CN
energies
(-AE”
)
for
the
equilibrium
XCH.$O;
+nH,O
*
- AE (kcal mol- ’ ) n=l
n=2
62-l
n=3
63-2
26.91 26.33 26.18 23.46 25.08 22.09 25.17 23.32 21.19
47.28 46.27 45.98 42.39 44.04 38.32 44.19 41.14 37.69
20.37 19.94 19.80 18.93 18.96 16.23 19.02 17.82 16.50
63.62 62.27 62.02 58.00 58.69 51.01 59.29 53.13 51.07
16.34 16.00 16.04 15.61 14.65 12.69 15.10 11.99 13.38
50
6
Thus, for X=H, the -LIE” value calculated kcal mol-l for 6 XCH,CO,
+2Hz0
for equilibrium
= XCH,CO,.BH,O
(6) was -47.3
(6)
Other results are given in Table 3. For the trihydrates, the structure 7 was the most stable for all of the substituents except X = H, F, where the corresponding structure but with the CX bond coplanar with the CO, group was preferred or equally stable.
d” I
H
I
H 6,
x D.&-c H’?‘
&__+4-J /
‘H
-
‘0
The calculated XCHzC!O;
value for equilibrium
+3H20
(7) for X = H was - 63.6 kcal mol-‘.
= XCH,CO,.3H,O
(7)
Table 3 gives the values of AE” for the equilibria (5), (6) and (7), for the most stable hydrates in each case. It is seen that the -dE” values for the formation of the monohydrates lie in the range 21-27 kcal mol-‘, for the dihydrates in the range 37-47 and for the trihydrates, in the range 51-64 kcal mol-l. Table 3 also lists the -LIE” values referring to equilibrium (8) CH3C0;
- (H,O),_,
+H,O
The values show that -dE”
= CH,C02
- (H,O),
(6)
values get smaller as n goes from 1 to 2 to 3, but
51
the fall off is gradual.Thus, for acetate itself for n= 1, -LIE” = 26.9, for n= 2, 20.4 and for n=3, 16.3 kcal mol-‘. A similar result was found ]2,3] for the successivehydrationof methylammoni~ ions. Structure of acetic acid hydrates
There are two stable hydratesfound as shown in structures8 and 9.
‘H 8
9
The structureswith the CX bond out of the plane of the carboxylic groupwere less stable, except for X= CN and Ph, where that had similar stability to the planar forms. The energiesfor equilibrium (9) are generallyconsiderablyless than XCHzCOzH+HzO = XCH2C02H*H20
(9)
for the monohydrationof the acetates.Thus, for X = H, the calculated -LB” for equilibrium (9) are 8.5 kcal mol-’ for structure8 and 4.4 kcal mol-l for structure9. Valuesfor other substitutedacetic acids are given in Table 4. TABLE 4
Calculated STO-3G energies ( -dE” ) for the ~~~ib~~ X
H Me Et Ph F Cl OMe CF, CN
XCH&O,H + Hz0 * XCH,CO,H (H,O
-AE” (kcal mol-‘) -OH. * .OHz
-C=O * - .HOH
8.47 8.38 8.34 8.19 9.33 9.90 8.40 9.57 10.17
4.36 3.87 3.87 3.61 3.4% 3.44 4.12 2.68 2.12
)
52
Acidities of substituted acetic acids
We can use the results above to examine the effects of hydration on the acidities of substituted acetic acids according to eqn. (10) XCH2COzH(OH,),
+CH3C0;
(H,O),
= XCH2CO; (HOO)y+CH&OaH(OH2),
(10)
Tbe energies for hy~ation at the carbonyl group of the acetic acids in process (9 ) were seen to be relatively small (Table 4 ) , and numerically smaller than for the energy for the dimerisation of water ( - 4.6 kcal mol-l at the STO3G basis). Thus, such structures should be of little importance in simulating the changes in going from the gas phase to aqueous solution. In Table 5, we list the experimental dG” values for both the gas-phase equilibria and for aqueous solution and also the calculated dE value [ 41 for the gas phase. Then, in Table 6, we compare the 6dG” values (gas -+ solution) with the calculated S&P values (gas --+process 10) and give the 6dEI” values for x = 0, y = 1, 2 and 3 and for 1c= 1, y = 3 where the water in the neutral acid is bonded to the OH group. Table 5 shows that the ab initio calculation gave results for the acidities close to the gas-phase results for X = F, OMe and CF3, whilst results for X= Cl and CN deviate by 3-5 kcal mol-I. The result for the chloroacetic acid is not surprising noting the inadequacy of minimal basis calculations for second-row elements such as chlorine. The average deviation between experimental and calculated results was only 1.4 kcal mol-1 (Cl omitted). Where, for the gas phase, the calculated results had been found to be significantly too high or too low, then a similar feature is found with the comparison TABLE 5 Acidities of substituted acetic acids (--AC” and -A_E” in kcal mol-‘)
H Me Et Ph F Cl OMe CF, CN
0.0 1.2 2.0 6.9 9.6 12.0 6.0 13.1 15.3c
“Ref. 13, STO-3G//STO-3G. ref. 13.
0.0 0.73 1.20 8.55 9.16 20.96 5.61 14.83 22.70
0.0 -0.08 - 0.09 0.61 2.96 2.58 1.62 2.31 3.12
0.0 - 0.44 - 0.27 3.21 3.37 6.93 1.37 3.24 8.45
bSee text, is for qn. (lo), x= 1, y=3. ‘Estimated from eqn. (2) in
53 TABLE 6 Calculated effect of hydration on the acidities of substituted acids [see text and eqn. (10) for explanation] X
z=O,y=l
x=o,y=2
x=o,y=3
x=l,y=3
6AG” k---q)
H Me Et Ph F Cl OMe CF, CN
0.0 0.5 0.7 3.5 1.8 4.8 1.7 3.6 5.7
0.0
0.0 1.3 1.6 5.6 4.9 12.6 4.3 10.5 12.6
0.0 1.2 1.5 5.3 5.8 14.0 4.2 11.6 14.3
0.0 1.3 2.1 6.3 6.6 9.4 4.4 10.8 12.2”
0.9
1.3 4.9 3.2 9.0 3.1 6.1 9.9
“-AG: estimated from eqn. (2) in ref. 13.
for the aqueous results. Therefore, in order to examine how far the various numbers of molecules of water explain aqueous results, we have constructed Table 6. This shows the change in acidity with amount of hydration and also lists the experimental change in going from gas to water. This clearly shows that three molecules of water on the carboxylate group and one on the carboxylic acid are sufficient to account for the change from gas to solution; the main effect is caused by hydration of the carboxylate group. The average deviation (Cl omitted) between the calculated and experimental change in going from the gas phase to aqueous solution is only 0.8 kcal mol-‘. We also calculated [ 231 values for y = 4 for a few substituents (H, F, CF3, CN). The fourth water, which is attached to a secondary solvation shell, brings the result for CN 2 kcal mol-’ closer to the experimental aqueous value but has almost no effect on the value for X=F and lowers that for CF3 by only 1 kcal mol-‘. This confirms that the inner solvation shell is responsible for the major changes here in going from the gas phase to aqueous solution. The results thus suggest that the change in going from gas to solution is accounted for by localized hydration near the charged site. A similar result has been reported for the proton affinities of a limited range of substituted methylamines. By contrast, for the acidities of molecules such as substituted pyridinium ions and phenols, hydration by up to three molecules of water near the charged site goes only part of the way (less than 50% ) to explaining the difference between results in the gas phase and in solution. One likely explanation is that the acetic acids and methylamines are relatively small molecules and most of the lines of force can pass through the molecular cavity. In the larger molecules, bulk solvent may become much more important in the transmission
54
of the field effect. Alternatively, the aromaticity of the larger molecules could be a factor. We are continuing the investigation into a larger range of methylamines as well as quinuclidines, bicycle [ 2.2.2 ] octyl carboxylic acids and octylamines, pyridines and phenols. ACKNOWLEDGEMENT
We are grateful to the Australian Research Grants Scheme for financial assistance. REFERENCES 1
I 8 9 10 11 12
13 14 15
16 17 18 19 20 21 22 23
For some general details see R.W. Taft, Prog. Phys. Org. Chem., 14 (1983) 247; R.W. Taft, in P. Ausloos (Ed.), Kinetics of Ion-Molecule Reactions, Plenum, New York, 1979. M. Meot-Ner, J. Am. Chem. Sot., 106 (1984) 1257,1265. Y.K. Lau, and P. Kebarle, Can. J. Chem., 59 (1981) 151. W.R. Davidson, J. Sunner and P. Kebarle, J. Am. Chem. Sot., 101 (1979) 1675. P. Kebarle, W.R. Davidson, J. Sunner and S. Meza-Hojer, Pure Appl. Chem., 51 (1979) 63. E.M. Arnett, B. Chawla, L. Bell, M. Taagerpera, W.J. Hehre and R.W. Taft, J. Am. Chem. sot., 99 (1977) 5729. K. Hiraoka, H. Takimoto and K. Morise, J. Am. Chem. Sot., 108 (1986) 5683. Y.K. Lau, S. Ikuta and P. Kebarle, J. Am. Chem. Sot., 104 (1982) 1462. Y.K. Lau, K. Nishizawa, A. Tse, R.S. Brown and P. Kebarle, J. Am. Chem. Sot., 103 (1981) 6291. D.A. MacDonaill and D.A. Morton-Blake, Theor. Chim. Acta, 65 (1984) 13. J. Bromilow, R.W. Taft and R.D. Topsom, unpublished results. W.L. Jorgensen and M. Ibrahim, J. Comput. Chem., 2 (1981) 7; S. Ikuta, Chem. Phys. Lett., 68 (1979) 179; S. Ikuta, J. Comput. Chem., 5 (1984) 374; J. Gao, D.S. Garner and W.L. Jorgensen, J. Am. Chem. Sot., 108 (1986) 4784. M. Taagepera, D. De Frees, W.J. Hehre and R.W. Taft, J. Am. Chem. Sot., 102 (1980) 424. C. Jinfeng, R.D. Topsom, A.D. Headley, I. Koppel, M. Mishima, R.W. Taft and S. Veji, J. Mol. Struct. (Theochem), 168 (1988) 141. J.S. Binkley, M. Frisch, K. Raghavachari, D. De Frees, H.B. Schlegel, R. Whiteside, R. Fluder, R. Seeger and J. Pople, Gaussian 82, Release A, Department of Chemistry, Carnegie Mellon University, Pittsburg, PA, 1981. M. Taagepera, K.D. Summerhays, W.J. Hehre, R.D. Topsom, A. Pross, L. Radom and R.W. Taft, J. Org. Chem., 46 (1981) 891. S. Marriott, R.D. Topsom, C.B. Lebrilla, I. Koppel, M. Mishima and R.W. Taft, J. Mol. Struct. (Theochem.), 137 (1986) 133. S. Marriott, R.W. Taft and R.D. Topsom, to be published. P.G. Mezey and W.F. Reynolds, Can. J. Chem., 55 (1977) 1567; S. Bohm and J. Kuthan, Int. J. Quantum Chem., 26 (1984) 21. C. Jingfeng and R.D. Topsom, J. Mol. Struct. (Theochem), in press. A. Pross, L. Radom and R.W. Taft, J. Org. Chem., 45 (1980) 818. A calculation showed that hydration of one of the carboxylate oxygens by the water dimer gave an optimized energy of - 374.04060 hartree. Energies (STO-SG//STO-3G) forXCH&Ol (H,O), were: H, 524.03480; F, 621.48112; CF,, 855.00130; CN, 614.587987 hartree.
EFFECT OF HYDRATION ON THE ACIDITIES OF SUBSTITUTED ACETIC ACIDS
CA1 JINFENG and RONALD D. TOPSOM Department of Chemistry, La Trobe University, Bundoora, 3083 (Australia) (Received 26 August 1988)
ABSTRACT
We report calculations at the ab-initio level on the hydration equilibria of substituted acetates XCH&O; + nH,O * XCH$O; with nH,O up to n = 3. Results are also given for the hydration energies of the corresponding acetic acids. These results allow the calculation of the acidities of the acetic acids in the presence of a limited number of water molecules. Such results are compared to experimental figures in the gas phase and in aqueous solution. Results for three water molecules per acetate ion and one per acetic acid molecule approach those found for aqueous solution.
INTRODUCTION
Most reactions are observed in solution, particularly in water, rather than in the gas phase. Thus, the effect of water molecules on the properties and reactions of organic molecules is of great importance. Much work has been done in recent years [l] on the interaction of specific molecules or ions with a discrete number of water molecules. Experimental studies, involving such techniques as pulsed electron beam mass spectrometry and ion cyclotron resonance, have included studies on the hydration of substituted ammonium ions with up to seven molecules of water [ 2,3], of pyridines and pyridinium ions [ 2,4-61, ROH$ ions [ 2,7,8], phenols [4] and anilines [9]. Theoretical studies include the hydration of substituted ammonium ions [3,10], pyridines and pyridinium ions [ 4,6], phenols, and phenolates [ 111 and various anions [ 121. Less work has been done on the effect of limited hydration on the equilibria of series of organic molecules. Results available include a study [6] of the proton affinity of substituted pyridines both in the gas phase [ eqn. (1) ] and with a single molecule of water [ eqn. (2) 1.
0166-1280/89/$03.50
0 1989 Elsevier Science Publishers B.V.
Experimental [4 ] and theoretical [6 ] (ab-initio STO-3G) results are available for both series and they are in reasonable agreement [ 1 ] (although the number of substituents was limited in equilibrium 2). Another example [ 3,131 is the proton transfer reaction of some substituted amines RNH,+ -zH, 0 + CH3 NH, .yH, 0 = CH3NH3+ .~cH~O+RNH~.~H~O
(3)
Here, the theoretical calculations reproduce the aqueous-phase figures for x = 3, y=2 where the water molecules lie along the NH axes. By contrast, the addition of several water molecules to the proton-transfer reactions of pyridinium ions [6] and phenols [4,11], goes only part of the way towards results found experimentally in aqueous solution. No explanation has been advanced for this difference. We have recently shown [14] that ab-initio STO-3G calculations give results for the relative acidities of substituted acetic acids [eqn. (4) ] in good agreement with experimental gas-phase results XCHz CO, H + CHBCO, = XCH, CO, + CHBCO2H
(4)
Results are also available for the acidities in aqueous solution and thus the series provides an excellent testing ground to investigate the effects of limited numbers of water molecules. CALCULATIONS
Calculations were performed at the ab-initio STO-3G basis using the GAUSSIAN-82 series of programs [ 151. The geometries of the acetic acids and acetates were fully optimized for the various possible conformers [ 141. The positions and structures of various interacting water molecules were then obtained
47
by geometry optimization (using the structures for the various acetate conformers ). Energies calculated for XCH&O; and its most stable mono-, diand tri-hy~ates are given in Table 1. In Table 2 we list the energies for molecules XCH,CO,H and their monohydrates. Previous work on organic equilibria has shown that calculations at the STO3G basis level give good agreement with experiment, for example, for the proton affinities of pyridine [ 161 and their monohydrates [ 1,4,6], cyanides [ 171, TABLE 1 Molecular orbital calculations ( -E) for XCH,C02 (H,O), at the STO-3G//STO-3G (the CX bonds are coplanar with the CO, plane, unless otherwise indicated)
basis level
E (hartree)
X
H Me Et Ph F Cl OMe CF3 CN
?t=O
n=l
n=2
n=3
224.04833 262.62979 301.21084 450.80853” 321.50457 678.67425 336.46345 555.03356 314.62462”
299.05711 337.63780 376.21846 525.81181” 396.51044 753.07535 411.46949 630.03663 389.62429”
374.05547 412.63547 451.21592 600.80788” 471.50655 828.06712 486.46571 705.03092 464.61648”
449.04742 481.6268? 526.20738’ 675.79866’ 546.495~ 903.05324” 561.45567” 780.01593” 539.60370”
eSubstituent at 90” to carboxylate plane. TABLE 2 Molecular orbital calculations ( -E) STO-3G basis level X
for XCHBCOpH and their monohydrates at the STO-3G//
E (hartree) XCH,CO,H
XCH&OeH(H,O)“
XCH,CO,H ( Hz0 )
224.81020 263.39050 301.97080 451.55678 322.25184 678.80271 337.21638 555.77179 315.35031
299.78960
299.78305 338.36256 376.94287 526.52843 397.22328 753.77410 412.18885 630.74196 390.31959
b
H Me Et Ph F Cl OMe CF, CN “-0-H..
*OHp.">C=O* *.HOH.
338.36975 376.94~ 526.53573 397.23261 753.78438 412.19566 630.75294 390.33242
substituted methylamines (the ab-initio STO-3G basis overall gives [ 13,181 as good agreement with experiment as split-valence basis calculation), and anilines [ 161 and for acidities of acetic acids [ 141, benzoic acids [ 19,207, 4substituted f 2.2.21bicycle-octyl carboxylic acids [ 203, and phenols 1211. RESULTS AND DISCUSSION
We first discuss the structure of the hydrated acetate ions and acetic acids. Structure of acetates In our previous paper, we calculated energies for optimized structures of the acetates having the CX bond coplanar with the COz group, as shown in structure 1.
O
*,
H
P
Hgg-c
xp$-“; 4
0
H
H
2
1
We have found, on a more intensive investigation, that this is indeed the most stable or equal most stable conformer for X = H, F, Cl, Me, Et, Me0 and CF3. For X = CN and Ph, conformer 2, with the XC bond at right angles to the COz plane, is more stable by 0.5 and 0.9 kcal mol-‘, respectively. St~~tures with the CH,X groups slightly rotated (ea. 30” ), compared to the CO, plane do lie in an energy well and are very close in stability to 1 for X=H, Me, Et, OMe and CF3. There is not such a well for X = F, Cl, CN and Ph. Str~eture of acetate hydrates We have investigated the interaction of one, two and three molecules of water with the acetates. For the monohydrates, the most stable structure found for the acetate component has the same conformation as found for the free acetate molecule. The preferred position of the water molecule in each case leads to the general structure 3. The process [eqn. (5) ] XCHzCO;
+H,O
* XCH,CO,
-Hz0
(5)
has energy here of - 26.9 kcal mol-l for X = H. The preference for structure 3 compared to structures such as 4 may reflect a weak attraction between the non-hy~ogen-bonded carboxylate oxygen and the far hydrogen on the water molecule.
49
b” ! i
O---H-O XCH,-C
/
‘H
-
XCH, -c-
'0 3
4
/O '0
The next most stable structure, 4, has, for X=H, a 1.6 kcal mol-’ lower calculated value of -LIE” for process (5). (Structure 5 lies, at best, in a very shallow energy well.) O---H
XCH,--C
’- ‘0 \ O---H / 5
The energy values for equilibrium (5) are given in Table 3 for other substituents. With the dihydrates, again the preferred structure [21] of the acetate part was as in the parent, that is, with the CX bond coplanar with the CO, group, except for X = CN and Ph. The most stable structure found for the dihydrates, in each case, was 6 (except that for X = CN and Ph, the XC bond was at right angles to the CO2 plane). TABLE 3 Calculated STO-3G XCH$O, (H,O),
X
H Me Et Ph F Cl OMe CF, CN
energies
(-AE”
)
for
the
equilibrium
XCH.$O;
+nH,O
*
- AE (kcal mol- ’ ) n=l
n=2
62-l
n=3
63-2
26.91 26.33 26.18 23.46 25.08 22.09 25.17 23.32 21.19
47.28 46.27 45.98 42.39 44.04 38.32 44.19 41.14 37.69
20.37 19.94 19.80 18.93 18.96 16.23 19.02 17.82 16.50
63.62 62.27 62.02 58.00 58.69 51.01 59.29 53.13 51.07
16.34 16.00 16.04 15.61 14.65 12.69 15.10 11.99 13.38
50
6
Thus, for X=H, the -LIE” value calculated kcal mol-l for 6 XCH,CO,
+2Hz0
for equilibrium
= XCH,CO,.BH,O
(6) was -47.3
(6)
Other results are given in Table 3. For the trihydrates, the structure 7 was the most stable for all of the substituents except X = H, F, where the corresponding structure but with the CX bond coplanar with the CO, group was preferred or equally stable.
d” I
H
I
H 6,
x D.&-c H’?‘
&__+4-J /
‘H
-
‘0
The calculated XCHzC!O;
value for equilibrium
+3H20
(7) for X = H was - 63.6 kcal mol-‘.
= XCH,CO,.3H,O
(7)
Table 3 gives the values of AE” for the equilibria (5), (6) and (7), for the most stable hydrates in each case. It is seen that the -dE” values for the formation of the monohydrates lie in the range 21-27 kcal mol-‘, for the dihydrates in the range 37-47 and for the trihydrates, in the range 51-64 kcal mol-l. Table 3 also lists the -LIE” values referring to equilibrium (8) CH3C0;
- (H,O),_,
+H,O
The values show that -dE”
= CH,C02
- (H,O),
(6)
values get smaller as n goes from 1 to 2 to 3, but
51
the fall off is gradual.Thus, for acetate itself for n= 1, -LIE” = 26.9, for n= 2, 20.4 and for n=3, 16.3 kcal mol-‘. A similar result was found ]2,3] for the successivehydrationof methylammoni~ ions. Structure of acetic acid hydrates
There are two stable hydratesfound as shown in structures8 and 9.
‘H 8
9
The structureswith the CX bond out of the plane of the carboxylic groupwere less stable, except for X= CN and Ph, where that had similar stability to the planar forms. The energiesfor equilibrium (9) are generallyconsiderablyless than XCHzCOzH+HzO = XCH2C02H*H20
(9)
for the monohydrationof the acetates.Thus, for X = H, the calculated -LB” for equilibrium (9) are 8.5 kcal mol-’ for structure8 and 4.4 kcal mol-l for structure9. Valuesfor other substitutedacetic acids are given in Table 4. TABLE 4
Calculated STO-3G energies ( -dE” ) for the ~~~ib~~ X
H Me Et Ph F Cl OMe CF, CN
XCH&O,H + Hz0 * XCH,CO,H (H,O
-AE” (kcal mol-‘) -OH. * .OHz
-C=O * - .HOH
8.47 8.38 8.34 8.19 9.33 9.90 8.40 9.57 10.17
4.36 3.87 3.87 3.61 3.4% 3.44 4.12 2.68 2.12
)
52
Acidities of substituted acetic acids
We can use the results above to examine the effects of hydration on the acidities of substituted acetic acids according to eqn. (10) XCH2COzH(OH,),
+CH3C0;
(H,O),
= XCH2CO; (HOO)y+CH&OaH(OH2),
(10)
Tbe energies for hy~ation at the carbonyl group of the acetic acids in process (9 ) were seen to be relatively small (Table 4 ) , and numerically smaller than for the energy for the dimerisation of water ( - 4.6 kcal mol-l at the STO3G basis). Thus, such structures should be of little importance in simulating the changes in going from the gas phase to aqueous solution. In Table 5, we list the experimental dG” values for both the gas-phase equilibria and for aqueous solution and also the calculated dE value [ 41 for the gas phase. Then, in Table 6, we compare the 6dG” values (gas -+ solution) with the calculated S&P values (gas --+process 10) and give the 6dEI” values for x = 0, y = 1, 2 and 3 and for 1c= 1, y = 3 where the water in the neutral acid is bonded to the OH group. Table 5 shows that the ab initio calculation gave results for the acidities close to the gas-phase results for X = F, OMe and CF3, whilst results for X= Cl and CN deviate by 3-5 kcal mol-I. The result for the chloroacetic acid is not surprising noting the inadequacy of minimal basis calculations for second-row elements such as chlorine. The average deviation between experimental and calculated results was only 1.4 kcal mol-1 (Cl omitted). Where, for the gas phase, the calculated results had been found to be significantly too high or too low, then a similar feature is found with the comparison TABLE 5 Acidities of substituted acetic acids (--AC” and -A_E” in kcal mol-‘)
H Me Et Ph F Cl OMe CF, CN
0.0 1.2 2.0 6.9 9.6 12.0 6.0 13.1 15.3c
“Ref. 13, STO-3G//STO-3G. ref. 13.
0.0 0.73 1.20 8.55 9.16 20.96 5.61 14.83 22.70
0.0 -0.08 - 0.09 0.61 2.96 2.58 1.62 2.31 3.12
0.0 - 0.44 - 0.27 3.21 3.37 6.93 1.37 3.24 8.45
bSee text, is for qn. (lo), x= 1, y=3. ‘Estimated from eqn. (2) in
53 TABLE 6 Calculated effect of hydration on the acidities of substituted acids [see text and eqn. (10) for explanation] X
z=O,y=l
x=o,y=2
x=o,y=3
x=l,y=3
6AG” k---q)
H Me Et Ph F Cl OMe CF, CN
0.0 0.5 0.7 3.5 1.8 4.8 1.7 3.6 5.7
0.0
0.0 1.3 1.6 5.6 4.9 12.6 4.3 10.5 12.6
0.0 1.2 1.5 5.3 5.8 14.0 4.2 11.6 14.3
0.0 1.3 2.1 6.3 6.6 9.4 4.4 10.8 12.2”
0.9
1.3 4.9 3.2 9.0 3.1 6.1 9.9
“-AG: estimated from eqn. (2) in ref. 13.
for the aqueous results. Therefore, in order to examine how far the various numbers of molecules of water explain aqueous results, we have constructed Table 6. This shows the change in acidity with amount of hydration and also lists the experimental change in going from gas to water. This clearly shows that three molecules of water on the carboxylate group and one on the carboxylic acid are sufficient to account for the change from gas to solution; the main effect is caused by hydration of the carboxylate group. The average deviation (Cl omitted) between the calculated and experimental change in going from the gas phase to aqueous solution is only 0.8 kcal mol-‘. We also calculated [ 231 values for y = 4 for a few substituents (H, F, CF3, CN). The fourth water, which is attached to a secondary solvation shell, brings the result for CN 2 kcal mol-’ closer to the experimental aqueous value but has almost no effect on the value for X=F and lowers that for CF3 by only 1 kcal mol-‘. This confirms that the inner solvation shell is responsible for the major changes here in going from the gas phase to aqueous solution. The results thus suggest that the change in going from gas to solution is accounted for by localized hydration near the charged site. A similar result has been reported for the proton affinities of a limited range of substituted methylamines. By contrast, for the acidities of molecules such as substituted pyridinium ions and phenols, hydration by up to three molecules of water near the charged site goes only part of the way (less than 50% ) to explaining the difference between results in the gas phase and in solution. One likely explanation is that the acetic acids and methylamines are relatively small molecules and most of the lines of force can pass through the molecular cavity. In the larger molecules, bulk solvent may become much more important in the transmission
54
of the field effect. Alternatively, the aromaticity of the larger molecules could be a factor. We are continuing the investigation into a larger range of methylamines as well as quinuclidines, bicycle [ 2.2.2 ] octyl carboxylic acids and octylamines, pyridines and phenols. ACKNOWLEDGEMENT
We are grateful to the Australian Research Grants Scheme for financial assistance. REFERENCES 1
I 8 9 10 11 12
13 14 15
16 17 18 19 20 21 22 23
For some general details see R.W. Taft, Prog. Phys. Org. Chem., 14 (1983) 247; R.W. Taft, in P. Ausloos (Ed.), Kinetics of Ion-Molecule Reactions, Plenum, New York, 1979. M. Meot-Ner, J. Am. Chem. Sot., 106 (1984) 1257,1265. Y.K. Lau, and P. Kebarle, Can. J. Chem., 59 (1981) 151. W.R. Davidson, J. Sunner and P. Kebarle, J. Am. Chem. Sot., 101 (1979) 1675. P. Kebarle, W.R. Davidson, J. Sunner and S. Meza-Hojer, Pure Appl. Chem., 51 (1979) 63. E.M. Arnett, B. Chawla, L. Bell, M. Taagerpera, W.J. Hehre and R.W. Taft, J. Am. Chem. sot., 99 (1977) 5729. K. Hiraoka, H. Takimoto and K. Morise, J. Am. Chem. Sot., 108 (1986) 5683. Y.K. Lau, S. Ikuta and P. Kebarle, J. Am. Chem. Sot., 104 (1982) 1462. Y.K. Lau, K. Nishizawa, A. Tse, R.S. Brown and P. Kebarle, J. Am. Chem. Sot., 103 (1981) 6291. D.A. MacDonaill and D.A. Morton-Blake, Theor. Chim. Acta, 65 (1984) 13. J. Bromilow, R.W. Taft and R.D. Topsom, unpublished results. W.L. Jorgensen and M. Ibrahim, J. Comput. Chem., 2 (1981) 7; S. Ikuta, Chem. Phys. Lett., 68 (1979) 179; S. Ikuta, J. Comput. Chem., 5 (1984) 374; J. Gao, D.S. Garner and W.L. Jorgensen, J. Am. Chem. Sot., 108 (1986) 4784. M. Taagepera, D. De Frees, W.J. Hehre and R.W. Taft, J. Am. Chem. Sot., 102 (1980) 424. C. Jinfeng, R.D. Topsom, A.D. Headley, I. Koppel, M. Mishima, R.W. Taft and S. Veji, J. Mol. Struct. (Theochem), 168 (1988) 141. J.S. Binkley, M. Frisch, K. Raghavachari, D. De Frees, H.B. Schlegel, R. Whiteside, R. Fluder, R. Seeger and J. Pople, Gaussian 82, Release A, Department of Chemistry, Carnegie Mellon University, Pittsburg, PA, 1981. M. Taagepera, K.D. Summerhays, W.J. Hehre, R.D. Topsom, A. Pross, L. Radom and R.W. Taft, J. Org. Chem., 46 (1981) 891. S. Marriott, R.D. Topsom, C.B. Lebrilla, I. Koppel, M. Mishima and R.W. Taft, J. Mol. Struct. (Theochem.), 137 (1986) 133. S. Marriott, R.W. Taft and R.D. Topsom, to be published. P.G. Mezey and W.F. Reynolds, Can. J. Chem., 55 (1977) 1567; S. Bohm and J. Kuthan, Int. J. Quantum Chem., 26 (1984) 21. C. Jingfeng and R.D. Topsom, J. Mol. Struct. (Theochem), in press. A. Pross, L. Radom and R.W. Taft, J. Org. Chem., 45 (1980) 818. A calculation showed that hydration of one of the carboxylate oxygens by the water dimer gave an optimized energy of - 374.04060 hartree. Energies (STO-SG//STO-3G) forXCH&Ol (H,O), were: H, 524.03480; F, 621.48112; CF,, 855.00130; CN, 614.587987 hartree.