Scales of hydrogen-bond acidity and basicity are now known both for solvents (a, and ti,) and for solutes (a2 and f32). Incorporation of the a, and tr parameters ...
MICHAEL H. ABRAHAM PRISCILLA L. GRELLIER The Chemistry Department University of Surrey Guildford, GU2 5 XH, UK
MORTIMER J. KAMLET RUTH M. DOHERTY Naval Surface Weapons Center White Oak Laboratory, Silver Springs Maryland 20910, USA
ROBERT W.TAFT The Chemistry Department University of California Irvine, California 92717, USA
JOSE-LOUIS M. ABBOUD Consejo Superior de Investigaciones Cienti fi cas Instituto de Quimica Fisíca Rocasolano Serrano 119 28008, Madrid, SPAIN
THE USE OF SCALES OF HYDROGEN-BOND ACIDITY AND BASICITY IN ORGANIC CHEMISTRY
Scales of hydrogen-bond acidity and basicity are now known both for solvents (a, and ti,) and for solutes (a2 and f3 2 ). Incorporation of the a, and tr parameters into a general equation for solvent effects has enabled the influence of solvents on a large number of physicochemical processes to be rationalised. Similarly, the solute a 2 and r2 parameters can be incorporated into a general equation for solute effects that has enabled a host of physico-chemical and bio-chemical processes to be explained and rationalised.
Rev. Port. Quím., 31, 85 (1989)
The acidity, or basicity, of compounds has long been regarded as a fundamental chemical property. Before the advent of the dissociation theory of Arrhenius, in 1887 [1] rather indirect measures of relative acid strength were used, such as catalysis of the inversion of sugar or of the hydrolysis of methyl acetate, but from the time of Arrhenius [2] to the present day [3] the dissociation of acids in aqueous solution has been taken as a quantitative measure of relative acid strength. Similarly, the relative strength of bases is measured [3] by either their pK b value, or the pK a value of the conjugate acid. The only generally accepted alternative to the proton-transfer theory has been that of G N Lewis [4] based on electron-pair acceptance (acids) or electron-pair donation (bases), but it is now clear that there are a large number of chemical structures and processes for which the pronton-transfer approach is not directly relevant, and for which the electron-pair approach does not lead to any quantitative methodology. For example, the three-dimensional structures of water and alcohols are based on hydrogen-bonded networks that involve no actual proton transfers, and the secondary structures of proteins are held in place by hydrogen bonds of the type N-H- - -O=C. A large number of processes, such as the solubility of nonelectrolytes in water or organic solvents, or the distribution of nonelectrolytes between two solvent phases, depend not at all on proton transfers but, as we shall see, on the ability of solutes or solvents to take part in hydrogen-bonding. To investigate such processes in any quantitative way, it is necessary to relate acidity or basicity specifically to hydrogen-bonding ability. We therefore define hydrogen-bond acidity (a.) as the ability of a molecule to act as a hydrogen-bond donor (HBD) [5] and define hydrogen-bond basicity (0) as the ability of a molecule to act as a hydrogen-bond acceptor (HBA) [6]. Because these parameters were first obtained by spectrophotometric methods, they are know as solvatochromic parameters. Some example of (HBD) acids and (HBA) bases are given in Table 1. 85
M. H. ABRAHAM, P. L. GRELLIER, M. J. KAMLET, R. M. DOHERTY, R. W. TAFC and J.-L. M. ABBOUD
(1 ) and the solute acid (a 2 ). In addition, there
Table 1 (HBD) acids ^— - -H-OMe
(HBA) bases Me 2 C = O - -
f— - -H-CH 2 Et 2 0 -
