Apply the Bronsted-Lowry theory, identify conjugate acid-base pairs, distinguish strong from weak acids and bases, and calculate pH using Kw

What this dot point is asking

The Bronsted-Lowry theory defines acid-base behaviour in terms of proton transfer. An acid is a proton ($\text{H}^+$) donor; a base is a proton acceptor. This is broader than the older idea that acids produce $\text{H}^+$ and bases produce $\text{OH}^-$, because it works for any proton transfer, including reactions in which ammonia acts as a base without any hydroxide being involved.

Conjugate acid-base pairs

When an acid donates a proton it becomes its conjugate base. When a base accepts a proton it becomes its conjugate acid. The two species in a pair differ by exactly one $\text{H}^+$. Consider:

Here ethanoic acid donates a proton to water. The conjugate pairs are $\text{CH}_3\text{COOH}/\text{CH}_3\text{COO}^-$ and $\text{H}_2\text{O}/\text{H}_3\text{O}^+$. Water acts as a base in this reaction. A species like water that can act as either an acid or a base, depending on what it reacts with, is amphiprotic (for example $\text{HCO}_3^-$ and $\text{H}_2\text{O}$).

Strong versus weak

The strength of an acid or base describes how completely it ionises in water, not how concentrated it is.

Strong acids (such as $\text{HCl}$, $\text{HNO}_3$, $\text{H}_2\text{SO}_4$) ionise essentially completely. Their reaction with water is written with a single arrow: $\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$.

Weak acids (such as $\text{CH}_3\text{COOH}$) only partially ionise, so they sit at equilibrium with the un-ionised molecule and are written with an equilibrium arrow. At any moment most molecules are not ionised.

The same distinction applies to bases. $\text{NaOH}$ is a strong base (fully dissociated), while aqueous ammonia is a weak base because the equilibrium $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$ lies well to the left. Note the difference between strength (degree of ionisation) and concentration (mol per litre). A concentrated weak acid can have a higher pH than a dilute strong acid.

Self-ionisation of water and Kw

Water ionises very slightly: $2\text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{OH}^-$. The ionic product is

In pure water $[\text{H}^+] = [\text{OH}^-] = 1.0 \times 10^{-7}\ \text{mol L}^{-1}$, giving pH 7 (neutral). Because $K_w$ comes from an endothermic ionisation, it increases with temperature, so neutral pH is below 7 at higher temperatures, though the solution is still neutral.

pH calculations

The pH scale measures hydrogen ion concentration:

For a basic solution, find $[\text{OH}^-]$, use $K_w$ to get $[\text{H}^+]$, then take the pH.

When you answer pH questions in the WACE examination, always check whether the acid or base is strong (full ionisation, so the concentration equals the ion concentration) or weak (an equilibrium that needs $K_a$, beyond a simple calculation), and quote pH to the correct number of decimal places.