Topic 12.1.1 The Bronsted-Lowry theory

• Acid-base equilibria involve the transfer of protons between substances. Therefore, substances can be classified as acids or bases depending on their interactions with protons.
• Bronsted-Lowry Theory defines an acid as a proton (H+ ion) donor, and a base as a proton acceptor.
• A Bronsted-Lowry acid is a proton donor. For example, ammonium ions.
• A Bronsted-Lowry base is a proton acceptor. For example, hydroxide ions.
• In order to accept a proton, a base has to have a lone pair of electrons so it can form a dative covalent bond with the proton.
• A conjugate acid is the species formed when a base accepts a proton.
• A conjugate base is the species formed when an acid donates a proton.
• Conjugate acids and conjugate bases form Bronsted-Lowry Conjugate Acid-Base pairs.
• A monoprotic/monobasic acid (e.g. hydrochloric acid) can only donate one proton.
• A diprotic/dibasic acid (e.g. sulfuric acid) can donate two protons.
• Diprotic/diacidic bases (e.g. carbonate ions) can accept two protons.
• An amphoteric substance can act as either an acid or a base in a reaction.
• Acid strength doesn't relate to the concentration of a solution. Instead it relates to the ability of an acid to dissociate in aqueous solution.
• A strong acid is defined as one that completely dissociates in aqueous solution, where the solution has a pH 0-1.
• A weak acid is defined as one that only slightly dissociates in aqueous solution, when in solution of pH 3-7.
• The definitions for strong and weak bases are identical to the definitions for strong and weak acids.
• Strong bases have a pH from 12-14, whilst weak bases have a pH from 7-11.
• Generally, the dissociation of a strong acid is represented using a single direction arrow in the equation.
• The dissociation of a weak acid is represented using the reversible (double) arrow symbol in the equation.