A. The Methylamines
The basic strength of methylamine is appreciably greater than that of ammonia. This is sometimes attributed to the +I inductive effect of the methyl group, leading to an increased electron availability at the nitrogen atom. It would perhaps be better to regard this as a mutual polarisation of the methyl and amino groups. At the same time it is of importance to note, as will be discussed later, that the difference in ΔG0 and hence in pKa arises more from the difference between the entropy changes associated with the dissociation of the ions than from the difference between the enthalpy changes.
The substitution of a second hydrogen atom of the ammonia molecule by a methyl group to give dimethylamine leads to only a small further increase in base strength, whilst the substitution of the third hydrogen atom leads to a considerable decrease, trimethylamine being only slightly stronger as a base than ammonia. This peculiar order of base strengths was attributed by Brown and his coworkers to what they call 'B-strain'. This they define as the strain introduced into a molecule as a result of changes in the normal bond angles of an atom brought about by the steric requirements of bulky groups attached to that atom. Thus they suggested that in trimethylamine the three methyl groups are crowded around the small nitrogen atom, and that their steric requirements are met by a spreading of the CNC bond angles to a value greater than the tetrahedral angle.
On this view the addition of a proton to the lone pair, which would tend to reduce the bond angle to the tetrahedral value, is resisted by the molecule. Very similar ideas have been expressed by Fyfe. These views, however, seem to be inconsistent with the observations of Lide and Mannl", who have deduced from the microwave absorption spectrum that the CNC bond angle in trimethylamine is 108.7 ± 1 degree.
The apparent anomaly is clarified to a considerable extent, however, by reference to the precision measurements of Everett and Wynne-Jones on ammonia, methylamine, dimethylamine, and trimethylamine, which covered a wide range of temperatures and of ionic strengths. These permitted the evaluation of ΔG0p, ΔS0, ΔG0, and ΔH0 for the dissociation of each of the cations. The final results for 25 degC and zero ionic strength are included in Table 1. [Not included, but can be redrawn when if required - when time allows; p_d]
For the dissociation of the ammonium ion ΔS0 is small and ΔC0p is zero. Unlike methane, which is repelled from water due to the large negative ΔS value involved, the ammonium ion must be very strongly attracted by water. As a result there is only a very small decrease in entropy on dissociation corresponding with the slight changes in the orderliness of arrangement attending the replacement of the NH+ ion by an H3O+ ion and a water molecule by an ammonia molecule. The equilibrium is determined principally, therefore, by ΔH0, which in turn depends mainly on the relative base strengths of ammonia and water.
As the hydrogen atoms of the ammonia molecule are successively replaced by methyl groups the value of ΔS0 changes systematically to more negative values. Trotman-Dickenson explained this as arising from the progressive decrease in the number of hydrogen atoms on the amine cation available for hydrogen bonding to solvent molecules, thus decreasing the constraint produced in the solvent through this cause. At the same time ΔC0p acquires a progressively larger positive value, this being associated with a large decrease in ΔH0.
For the change from ammonia to methylamine the entropy effect predominates and leads to a large increase in ΔG0 and hence in pKa, but for the subsequent changes from methylamine to dimethylamine and from dimethylamine to trimethylamine almost similar increases occur in -ΔS0, but with each step ΔC0p shows increasing increments, with associated decreases in ΔH0. The result is that a maximum value of ΔG0 is reached at dimethylamine.
In addition to the effect produced by the reduction of the number of hydrogen bonds to the cation, and the further ordered structure produced near the aminium group, the successive additions of methyl groups would be expected to yield increasing areas of surface over which hydrophobic hydration can occur in the neutral amine molecule, and to reduce slightly the hydrogen bonding of its amino group. It is difficult to assess the area over which hydrophobic hydration can occur but it is not unreasonable to suppose that this may be very much greater over the almost complete hemisphere formed by three methyl groups than over the smaller surfaces presented by two or one methyl groups. Being close to the centre of positive charge this hydration atmosphere will be almost wholly dispersed in the cation, so the heat capacity of the system in the dissociated state will be higher than that for the cation-solvent system, since it contains a term representing the heat required to disperse this solvent layer.
This has a corresponding effect on the enthalpy of the amine-solvent system, since a considerable amount of heat would be required to bring the solvent atmosphere around the methyl groups isothermally to the more random state which exists around these groups in the cation. These effects and theories regarding hydrophobic hydration are discussed fully by Ives and Marsden.
It is interesting to note that the effects become modified in aqueous methanol solution, where for ammonia ΔS0 becomes positive, possibly through the ammonium ion exerting a greater sorting effect in its hydrogen bonding than does the hydroxonium ion. Although ΔC0p; has the rather unexpected value of - 12 cal/deg/mole for methylamine, it is very low for the other amines, indicating that no phenomenon analogous to hydrophobic hydration occurs to an appreciable extent in this solvent mixture. The changes in ΔS0, therefore, may reflect the decrease in the number of hydrogen atoms in the cation available for hydrogen-bond formation. The ΔG0 values for dimethylamine and triethylamine must be lower at zero ionic strength than at I = 0.10, so the net result is that in this medium methylamine seems to be the strongest of the bases.
