Theories of Acids and Bases

In the Arrhenius model, an acid produces H⁺(aq) in water and a base produces OH^-(aq). HCl(aq) provides H⁺ and Cl^-; NaOH(aq) provides Na⁺ and OH^-. Neutralisation removes H⁺ and OH^- as water: H⁺(aq) + OH^-(aq) → H₂O(l).

The Bronsted-Lowry model follows a proton: an acid donates H⁺, and a base accepts it. A conjugate acid-base pair differs by exactly one proton. Losing that proton also reduces the signed charge by one; gaining it increases the signed charge by one.

Water accepts a proton from HCl but donates one to NH₃; a substance capable of either role is amphiprotic. In aqueous equations H⁺(aq) is convenient shorthand for the hydrated proton. Writing H₃O⁺ makes the transfer to water visible.

A strong acid is essentially fully dissociated in dilute aqueous solution. A weak acid establishes an equilibrium in which only a proportion is dissociated. Strong and weak bases differ similarly in the extent to which they generate OH^- in water. "Concentrated" and "dilute" instead describe amount per volume. A concentrated weak acid is still weak.

A weak acid can still react completely with enough strong base. As H⁺ is removed, more weak acid dissociates. Equal volumes of equal-concentration monoprotic strong and weak acids therefore require the same amount of NaOH for complete neutralisation, even though their initial pH and initial reaction rates differ. Comparing pH alone cannot establish acid strength unless concentration and the number of ionisable protons are also controlled.

For a monoprotic weak acid HA(aq) ⇌ H⁺(aq) + A^-(aq), K_a = [H⁺][A^-]/[HA] at equilibrium. At a fixed temperature, a larger K_a means a greater tendency to donate a proton. For a base B + H₂O ⇌ BH⁺ + OH^-, K_b = [BH⁺][OH^-]/[B]. Pure liquid water is omitted from these expressions.

pH = -log₁₀[H⁺(aq)], with the concentration expressed in mol dm^-3 for these calculations. One pH unit corresponds to a tenfold change in hydrogen-ion concentration. K_w = [H⁺][OH^-] is the ionic product of water. At 25 degrees C, K_w = 1.00 × 10^-14 mol² dm^-6.

At equivalence, the acid and base have reacted in the stoichiometric proportion. An indicator changes colour over a pH range; its endpoint should occur within the steep portion of the titration curve, where a very small added volume causes the whole indicator transition. Choose using the supplied curve or transition data, not merely a memorised colour.

For reliable observations, add titrant dropwise near the endpoint, swirl to mix, and use a consistent faint permanent colour. A darker final colour after extra titrant is not "more complete": it means the endpoint has been overshot.

An acidic buffer contains appreciable amounts of a weak acid and its conjugate base, for example CH₃CO₂H and dissolved sodium ethanoate. Added H⁺ is consumed by ethanoate; added OH^- is consumed by the weak acid. Because most added strong acid/base is removed, the pH changes much less than it would in water.

A basic buffer uses a weak base and its conjugate acid, such as NH₃/NH₄⁺. NH₃ consumes H⁺; NH₄⁺ consumes OH^-. Buffers help keep enzyme conditions suitable and prevent unwanted pH shifts in formulations. Their capacity is finite: a large acid/base addition can use up one component.

The required ocean pair is CO₃^2-/HCO₃^-. Carbonate can accept a proton to form hydrogencarbonate; hydrogencarbonate can donate a proton to react with added hydroxide. This is one part of the wider seawater carbonate system.

More atmospheric CO₂ promotes more dissolved CO₂. Its reaction with water can be represented overall as CO₂(aq) + H₂O(l) ⇌ H⁺(aq) + HCO₃^-(aq). The extra H⁺ is partly taken up by carbonate, increasing hydrogencarbonate and reducing carbonate availability. Buffering limits the change; it cannot cancel a sustained large input, so ocean pH falls.