Lewis Acids “An acid is an electron pair acceptor, A base is an electron pair donor”

Comparison of definitions All Bronsted Lowry acids are also Lewis acids. Bronsted-Lowry – Proton donor Eg HCl → H+ + Cl- Lewis – electron pair acceptor H+ has no electrons, so can accept a pair to form a dative bond. H+ + (:OH)- → H←:OH

Cu2+ + 6H2O → [Cu(H2O)6]2+ Cu2+←:OH2 But other Lewis acid/bases would not qualify under the Bronsted-Lowry definition. Eg; When a water molecule acts as a ligand in a complex ion, forming a dative bond, it acts as a Lewis base. By accepting an electron pair the transition metal is acting as a Lewis acid. Cu2+ + 6H2O → [Cu(H2O)6]2+ Cu2+←:OH2 Acid Base

Acidity of complex ions [Fe(H2O)6]2+ solutions are not noticeably acidic. But [Fe(H2O)6]3+ is more acidic than ethanoic acid. (pKa = 2.2) This is because Fe3+ has a higher charge density than Fe2+ and is more polarising. Fe3+ more strongly attracts the electrons on oxygen, weakening the OH bonds, releasing protons. [Fe(H2O)6]3+ ⇌ [Fe(H2O)5OH]2+ + H+

This type of reaction is often referred to as hydrolysis; [Fe(H2O)6]3+ + H2O ⇌ [Fe(H2O)5OH]2+ + H3O+ The complex is thus acting as a Bronsted Lowry acid. Generally the M3+ aqua ions are more acidic than the M2+. (The same is also true for Al3+.)