Acid–Base Equilibria and pH

A Brønsted–Lowry acid donates a proton and a base accepts one. A conjugate pair differs by one proton. A strong acid essentially dissociates completely in dilute aqueous solution; a weak acid establishes an equilibrium with mostly undissociated acid. Concentration describes how much acid was dissolved per unit volume, whereas strength describes dissociation. A dilute strong acid may have a higher pH than a concentrated weak acid. Never infer strength from the label concentrated alone.

Use $\mathrm{pH}=-\log_{10}[H^+]$. For a dilute monoprotic strong acid, hydrogen-ion concentration is approximately the stated acid concentration. For a weak acid HA, $K_a=[H^+][A^-]/[HA]$. In an initially pure weak-acid solution, [H+] and [A⁻] are approximately equal, and if dissociation is small [HA] is approximately its initial concentration. Then $[H^+]\approx\sqrt{K_ac}$. Check that the calculated dissociation is indeed small before accepting this approximation.

Water’s ionic product is $K_w=[H^+][OH^-]$. At 298 K its usual value is $1.00\times10^{-14}$ mol² dm⁻⁶. For a strong alkali, find [OH⁻] from its concentration and stoichiometry, then calculate [H+] using Kw. Kw changes with temperature, so neutral does not always mean pH 7: neutrality means equal hydrogen- and hydroxide-ion concentrations. A pH meter needs calibration with suitable buffers and rinsing between readings; its display resolution alone does not establish total experimental accuracy.

Attempt these problems on paper before opening the worked reasoning. Explain your steps as well as giving a final answer.