Abstract
The kinetics of hydrolyses of N-substituted aliphatic amides was studied spectro-photometrically in aqueous solution at 30-95°. In concentrated hydrochloric acid solution up to 2.0M, the observed rate constants were found to increase to a constant value. Activation entropies for acidic hydrolyses of N-alkyl acetamides were largely negative, ranging -18--32 e. u. The results can be explained by a bimolecular reaction mechanism in which a nucleophilic attack by water molecules on a protonated amide molecule is the rate-determining step. In alkaline hydrolyses of N-alkyl amides, the observed firstorder rate constants k1 were found to follow a kinetic equation : k1=k2 [OH-]+k3 [OH-]2. The activation entropies were largely negative in the range from -29 to -38 e. u. and their activation enthalpies were approximately 4 to 7 kcal/mole smaller than in acidic hydrolysis. The results suggest that the rate-determining step is the nucleophilic attack on amide molecule by hydroxide ions, and that hydrolyses of N-substituted acetamides, acyl-substituted amides and esters have the same reaction mechanism. Using Taft's method, polar and steric substituent effects on the rates of hydrolyses of N-substituted acetamides are also studied.
