Bronsted-Lowry Acids and Bases

The Arrhenius definition limits acids to substances producing hydrogen ions in water and bases to those producing hydroxide ions. The Brønsted-Lowry definition is broader, defining acids as proton donors and bases as proton acceptors, allowing acids and bases to exist outside aqueous solutions and bases without hydroxide ions. This expanded framework better describes acid-base behavior across diverse chemical systems.

Conjugate acid-base pairs are species that differ by one proton. When an acid donates a proton, it becomes its conjugate base; when a base accepts a proton, it becomes its conjugate acid. For example, acetic acid donates a proton to water, forming acetate (conjugate base) and hydronium ion (conjugate acid). These pairs are central to understanding reversible acid-base reactions.

Acid strength and conjugate base strength are inversely related. A strong acid completely dissociates, leaving a very weak conjugate base unable to accept protons. Conversely, a weak acid partially dissociates, producing a relatively weak conjugate base that can accept protons in the reverse reaction. This inverse relationship determines whether acid-base reactions reach equilibrium or completion.

Ammonia acts as a base because it accepts a proton from water according to Brønsted-Lowry theory. When ammonia accepts a proton from water, it forms an ammonium ion while water becomes a hydroxide ion. This proton-accepting behavior defines ammonia as a base, demonstrating that bases need not contain hydroxide ions initially.

Strong acids and bases dissociate completely, producing extremely weak conjugate bases or acids unable to reverse the reaction, causing it to go to completion. Weak acids and bases partially dissociate, producing weak conjugate acids or bases capable of reversing the reaction, establishing equilibrium. The relative strengths of the weak species determine the equilibrium position.

Water acts as an acid by donating a proton to bases like ammonia, becoming a hydroxide ion. Water acts as a base by accepting a proton from acids like hydrochloric acid, becoming a hydronium ion. This amphoteric behavior demonstrates that substances can be classified as acids or bases depending on the other species present in the reaction.

A stronger acid more readily donates protons than a weaker acid. Acid strength is determined by its propensity to transfer protons to bases. Strong acids like hydrochloric acid donate protons completely in solution, while weak acids like acetic acid donate protons only partially, establishing equilibrium with their conjugate bases.