3.1 Proton transfer reactions

Pair species that differ by exactly one H⁺ to identify conjugate acid–base pairs. Charge alone does not decide the role: in NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺, NH₄⁺ is the proton donor and water is the acceptor.

NH₄⁺/NH₃ and H₂CO₃/HCO₃⁻ are conjugate pairs because each pair differs by one H⁺. Removing H⁺ lowers charge by one; adding H⁺ raises it by one. Do not pair species merely because they occur on opposite sides of an equation—trace the specific proton transfer.

For HCO₃⁻, donation gives CO₃²⁻ whereas acceptance gives H₂CO₃. Showing both reactions is the evidence for amphiprotic behaviour; one acid–base equation alone is insufficient.

For [H⁺] = 2.0 × 10⁻³ mol dm⁻³, pH = 2.70; the leading 2 makes the answer non-integer. A colour indicator estimates a range, whereas a calibrated probe supports a numerical measurement.

At 25 °C, Kw = 1.0 × 10⁻¹⁴, so a neutral solution has [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ mol dm⁻³. Neutrality always means equal ion concentrations; neutral pH is not necessarily 7 when temperature changes.

At a fixed temperature, Kw is constant, so [OH-] = Kw/[H+]: a higher [H+] means a lower [OH-]. For example, at pH 9.3 and 25 C, [OH-] = 2.0 x 10^-5 mol dm^-3. Classify a solution from the ion comparison; do not assume neutral pH is 7 at every temperature.

Represent a strong acid with essentially complete ionization and a weak acid with an equilibrium containing substantial undissociated acid. Strength is an equilibrium property, while concentration is an initial amount per volume; pH depends on both, so strength alone cannot rank arbitrary solutions.

Balance proton capacity as well as atoms: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O, while an acid–carbonate reaction also releases CO₂. To identify parents of Na₂SO₄, trace SO₄²⁻ to the acid and Na⁺ to the base rather than treating the salt name as a reaction equation.

For a strong acid titrated with strong base at 25 °C, calculate the initial pH from excess acid, locate equivalence from stoichiometric moles, and place the steep section around pH 7. Equivalence is a mole condition; it is not the same as equal solution volumes unless concentrations and stoichiometry make it so.

A reliable route is [OH⁻] → pOH → pH → [H⁺], or the reverse, with each logarithm shown. Use pH + pOH = pKw; replacing pKw by 14 is valid only at 25 °C.

Worked pOH example: for $0.025\,\mathrm{mol\,dm^{-3}}$ $\ce{KOH(aq)}$, complete dissociation gives $[\ce{OH^-}]=0.025\,\mathrm{mol\,dm^{-3}}$. Therefore $\mathrm{pOH}=-\log_{10}(0.025)=1.60$. The low pOH is consistent with a basic solution; at 298 K, $\mathrm{pH}=14.00-1.60=12.40$.

Large Ka and small pKa both indicate the stronger acid; large Kb and small pKb indicate the stronger base. Strength describes extent of ionization, whereas concentration describes amount per volume—dilute and weak are not synonyms.

Match the constants to one conjugate pair: Ka(HA) × Kb(A⁻) = Kw. A stronger acid therefore has a weaker conjugate base at the same temperature; do not multiply constants belonging to unrelated species.

Worked conjugate-constant example at 298 K: methylamine has $\mathrm{p}K_b=3.34$, so for its conjugate acid $\ce{CH3NH3+}$, $\mathrm{p}K_a=\mathrm{p}K_w-\mathrm{p}K_b=14.00-3.34=10.66$. Equivalently, $K_aK_b=K_w$. This relationship applies only to a conjugate acid–base pair at the same temperature.

Spectator ions from strong parents do not control pH. NH₄Cl is acidic because NH₄⁺ donates a proton to water, while a carbonate salt is basic because CO₃²⁻ accepts one; write the hydrolysing ion, not the intact salt, in the equilibrium.

A weak acid–strong base curve starts at a higher pH than an equally concentrated strong acid, contains a buffer region, has pH = pKa at half-equivalence, and has an alkaline equivalence point from conjugate-base hydrolysis. Weak–weak titrations often lack a sufficiently steep jump for a simple indicator endpoint.

Added acid shifts HInd ⇌ H⁺ + Ind⁻ toward the HInd colour; added base favours Ind⁻. The visible transition spans a range around pKa because both colours must change in relative abundance. Universal indicator combines several such equilibria and is not one substance with every colour.

Overlay the indicator transition range on the titration curve and require the whole visible change to fall inside the steep region. A weak acid–strong base equivalence is alkaline and favours an alkaline-range indicator; a strong acid–weak base equivalence is acidic. Endpoint proximity, not a memorized indicator name alone, is the criterion.

In an ethanoic acid/ethanoate buffer, CH₃COO⁻ consumes added H⁺ and CH₃COOH consumes added OH⁻, so the conjugate ratio changes only slightly. A buffer resists small additions but has finite capacity; once one component is nearly exhausted, the pH can change sharply.

When a small amount of acid is added, A⁻ removes it to form HA; added base is removed by HA to form A⁻. Dilution leaves their ratio nearly unchanged but reduces buffer capacity, so resistance to a large addition is not unchanged.

Worked buffer example: an ethanoate buffer contains $0.100\,\mathrm{mol\,dm^{-3}}$ $\ce{CH3COOH}$ and $0.200\,\mathrm{mol\,dm^{-3}}$ $\ce{CH3COO^-}$, with $\mathrm{p}K_a=4.76$. Substitution gives $\mathrm{pH}=\mathrm{p}K_a+\log_{10}([\ce{A^-}]/[\ce{HA}])=4.76+\log_{10}(0.200/0.100)=5.06$. The pH is above $\mathrm{p}K_a$ because the conjugate base is more concentrated than the acid.