IB Chemistry R3.1 Proton Transfer Reactions Key Definitions

The equilibrium constant for the dissociation of a weak acid in water, $K_a = \mathrm{\frac{[H^+][A^-]}{[HA]}}$. A larger value of $K_a$ means the acid dissociates more fully, so it is a stronger acid.

Acid–base indicators are weak acids that exist in equilibrium between their protonated form (HInd) and deprotonated form (Ind⁻).

An amphiprotic species is a chemical species that can either donate a proton (acting as an acid) or accept a proton (acting as a base), depending on the reaction conditions.

A Brønsted–Lowry acid is a proton (H⁺) donor, which gives up a hydrogen ion in a reaction.

A Brønsted–Lowry base is a proton (H⁺) acceptor, which gains a hydrogen ion in a reaction.

The Brønsted–Lowry theory, introduced in 1923, defines acids and bases based on their ability to transfer protons (H⁺ ions).

Buffers are solutions that resist changes in pH when small amounts of acid ($H⁺$ ions) or base ($OH⁻$ ions) are added.

A conjugate acid–base pair consists of two species that differ by exactly one proton ($H^+$).

The equilibrium constant $K$ is a ratio that provides a snapshot of the relative concentrations of products and reactants at equilibrium for a reversible chemical reaction.

The equivalence point in a titration is the precise moment when the amount of base added is chemically equivalent to the amount of acid initially present.

When a salt dissolves in water, its ions may undergo hydrolysis, reacting with water to form either H$^+$ (acidic) or OH$^−$ (basic) ions.

The equilibrium constant for the self-ionisation of water. It equals the product of the hydrogen ion and hydroxide ion concentrations, Kw = [H⁺][OH⁻], and has a value of 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.

Neutralization is a chemical reaction where an acid reacts with a base to produce a salt and water.

The pH scale is a logarithmic measure of the concentration of hydrogen ions ($ [H^+] $) in an aqueous solution.

The negative base-ten logarithm of the hydroxide ion concentration, pOH = −log₁₀[OH⁻]. A lower pOH means a higher [OH⁻] and a more basic solution.

Strong acids and bases ionize completely in water. All their molecules dissociate into ions, leaving no undissociated molecules in the solution.

Weak acids and bases only partially ionize in water. At equilibrium, the solution contains both ions and undissociated molecules.