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LearnEdexcelYCH112018 specificationWCH14 · Unit 4: Rates, Equilibria and Further Organic ChemistryAcid-base equilibria

WCH14 · Module

Acid-base equilibria

2 subtopics≈ 95 min in total

Start the module: pH, Ka and buffers Open graphing calculator

pH, Ka and buffers
$\mathrm{pH}=-\log[\mathrm{H^+}]$ . Strong acids: $[\mathrm{H^+}]=$ acid concentration (× number of H⁺). Water: $K_w=[\mathrm{H^+}][\mathrm{OH^-}]=1.00\times10^{-14}$ at 298 K; for strong bases, $[\mathrm{H^+}]=K_w/[\mathrm{OH^-}]$ . Weak acids: $K_a=\frac{[\mathrm{H^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$ ; assume $[\mathrm{H^+}]=[\mathrm{A^-}]$ , so $[\mathrm{H^+}]=\sqrt{K_a[\mathrm{HA}]}$ ; $\mathrm{p}K_a=-\log K_a$ . A buffer (weak acid + its salt) resists pH change: $[\mathrm{H^+}]=K_a\frac{[\mathrm{HA}]}{[\mathrm{A^-}]}$ , or $\mathrm{pH}=\mathrm{p}K_a+\log\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}$ . Titration curves show a steep section at equivalence; choose an indicator whose colour change lies within it. At half-neutralisation, pH = pKa.
≈ 55 min
Brønsted–Lowry acids and bases
A Brønsted–Lowry acid is a proton (H⁺) donor; a base is a proton acceptor. Every acid–base reaction has two conjugate pairs: in $\mathrm{HCl+H_2O\rightarrow Cl^-+H_3O^+}$ , HCl/Cl⁻ and H₃O⁺/H₂O. Strong acids dissociate fully ( $\mathrm{HCl\rightarrow H^++Cl^-}$ ); weak acids partially ( $\mathrm{CH_3COOH\rightleftharpoons CH_3COO^-+H^+}$ ). $\mathrm{pH}=-\log_{10}[\mathrm{H^+}]$ ; for a strong monoprotic acid $[\mathrm{H^+}]=c$ . Water is amphoteric: it can act as an acid or a base.
≈ 40 min