Conjugate acid–base pairs differ by one proton.

The hydrogencarbonate ion, $HCO_3^-$, is amphiprotic in aqueous solution.

A sample of rainwater has a pH of 5.60 at $298\ \text{K}$.

At $298\ \text{K}$, the hydroxide ion concentration in an alkaline cleaning solution is $3.20 \times 10^{-4}\ \text{mol dm}^{-3}$.

The acid dissociation data for two weak acids at $298\ \text{K}$ are:

$HX$: $pK_a = 3.18$

$HY$: $K_a = 1.6 \times 10^{-5}\ \text{mol dm}^{-3}$

The diagram shows observations when dry hydrogen chloride and ammonia gases are allowed to mix, and an aqueous proton transfer reaction involving ethanoic acid.

Equal volumes of $0.100\ \text{mol dm}^{-3}$ hydrochloric acid and $0.100\ \text{mol dm}^{-3}$ ethanoic acid are tested separately under identical conditions.

A $25.00\ \text{cm}^3$ sample of a monoprotic strong acid is titrated with $0.100\ \text{mol dm}^{-3}$ sodium hydroxide. The equivalence point occurs after $18.60\ \text{cm}^3$ of sodium hydroxide has been added.

Aqueous solutions of ammonium nitrate, $NH_4NO_3$, and sodium carbonate, $Na_2CO_3$, are prepared at the same concentration.

A monoprotic acid is titrated with a strong base. The pH curve starts above pH 2, has a buffer region, and has an equivalence point above pH 7. The half-equivalence point occurs at pH 4.74.

A student investigated the behaviour of aqueous hydrogencarbonate ions by adding them to acidic and alkaline solutions. The measured pH changes are summarized.

The pH of four household solutions was estimated using universal indicator and then measured using a calibrated pH probe.

At $298\ \text{K}$, the ion product constant of water is $K_w = 1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$. Some ion concentrations for dilute aqueous solutions are shown.

At $298\ \text{K}$, several alkaline cleaning solutions were analysed. Some pOH, pH and hydroxide ion concentration values are shown.

The acid dissociation data for three weak acids at $298\ \text{K}$ are shown.

Ethanoic acid has $K_a = 1.80 \times 10^{-5}\ \text{mol dm}^{-3}$ at $298\ \text{K}$. A solution of sodium ethanoate has concentration $0.100\ \text{mol dm}^{-3}$. Use $K_w = 1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$.

A buffer solution is prepared by mixing ethanoic acid and sodium ethanoate so that the final concentrations are $0.200\ \text{mol dm}^{-3}$ $CH_3COOH$ and $0.300\ \text{mol dm}^{-3}$ $CH_3COO^-$. For ethanoic acid, $pK_a = 4.76$ at $298\ \text{K}$.

Equal volumes of $0.100\ \text{mol dm}^{-3}$ hydrochloric acid and ethanoic acid were separately reacted with excess magnesium ribbon. The conductivity and pH of each acid were also measured before the magnesium was added.

A $25.00\ \text{cm}^3$ sample of hydrochloric acid was titrated with $0.100\ \text{mol dm}^{-3}$ sodium hydroxide. The pH was recorded after each addition of sodium hydroxide.

The table gives dissociation constants for weak acids and their conjugate bases at $298\ \text{K}$. Use $K_w = 1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$.

Equal concentrations of four salt solutions were prepared at $298\ \text{K}$. Their pH values were measured after calibration of the pH probe.

The hydrogencarbonate ion can take part in several proton transfer reactions.

Reaction 1: $HNO_3(aq) + H_2O(l) \to H_3O^+(aq) + NO_3^-(aq)$

Reaction 2: $HCO_3^-(aq) + H_2O(l) \rightleftharpoons CO_3^{2-}(aq) + H_3O^+(aq)$

Reaction 3: $HCO_3^-(aq) + H_2O(l) \rightleftharpoons H_2CO_3(aq) + OH^-(aq)$

At $298\ \text{K}$, an aqueous sample from a lake affected by acid drainage has a measured pH of $3.82$. The value of $K_w$ at this temperature is $1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$.

Neutralization reactions are used to prepare salts and to treat excess acidity. Magnesium oxide, sodium carbonate and ammonia are common bases.

A student monitors the acidity of a coloured fruit drink during storage. The hydrogen ion concentration in one sample is $4.0 \times 10^{-5}\ \text{mol dm}^{-3}$.

An acid-base indicator can be represented as the weak acid $HInd$. The undissociated form, $HInd$, is yellow and the deprotonated form, $Ind^-$, is red. The indicator has $pK_a = 5.1$.

A weak monoprotic acid was titrated with sodium hydroxide. For comparison, other possible acid-base titration curves and several indicator ranges are shown.

Several ethanoic acid/ethanoate buffer mixtures were prepared at $298\ \text{K}$. The table shows their composition and the pH measured before and after dilution or addition of acid. For ethanoic acid, $pK_a = 4.76$.

Equal volumes of $0.100\ \text{mol dm}^{-3}$ hydrochloric acid and $0.100\ \text{mol dm}^{-3}$ ethanoic acid are separately reacted with excess calcium carbonate at the same temperature. The carbon dioxide produced is collected over time.

A $25.00\ \text{cm}^3$ sample of hydrochloric acid is titrated with $0.1000\ \text{mol dm}^{-3}$ sodium hydroxide at $298\ \text{K}$. The equivalence volume is $21.40\ \text{cm}^3$.

Methylamine, $CH_3NH_2$, is a weak base. At $298\ \text{K}$, $K_b$ for methylamine is $4.4 \times 10^{-4}\ \text{mol dm}^{-3}$. A solution has an initial methylamine concentration of $0.150\ \text{mol dm}^{-3}$.

At $298\ \text{K}$, propanoic acid has $pK_a = 4.87$. Ethanoic acid has $pK_a = 4.76$.

The pH of a salt solution depends on whether the ions in the salt hydrolyse in water. Four salts are sodium methanoate, ammonium nitrate, sodium hydrogencarbonate and potassium chloride.

Two $25.00\ \text{cm}^3$ samples of monoprotic acids are titrated with sodium hydroxide of the same concentration. The pH curves are shown.

An acidic buffer is prepared at $298\ \text{K}$ by mixing $50.0\ \text{cm}^3$ of $0.200\ \text{mol dm}^{-3}$ ethanoic acid with $30.0\ \text{cm}^3$ of $0.200\ \text{mol dm}^{-3}$ sodium hydroxide. The $pK_a$ of ethanoic acid is $4.76$.