Proton transfer reactions

A **proton (H⁺) donor**.

A **proton (H⁺) acceptor**.

A **hydrogen ion, H⁺** — a hydrogen atom that has lost its electron.

Two species that differ by **exactly one H⁺** (an acid and the base left after it donates).

**Remove** one H⁺ from the acid (e.g. HCl → Cl⁻).

**Add** one H⁺ to the base (e.g. NH_{3} → NH_{4}^{+}).

A species that can **both donate and accept** a proton (e.g. H_{2}O, HCO_{3}⁻).

**HSO_{4}⁻** (remove one H⁺ — not SO_{4}^{2-}, which is two H⁺ away).

**H_{3}O^{+}** (the oxonium / hydronium ion).

**H_{2}O** and **HCO_{3}⁻** — both can donate or accept a proton.

**HCl** — it donates the proton; water is the base.

An acid can only **donate** H⁺ if a base is there to **accept** it — every proton transfer has both.

A measure of acidity based on hydrogen-ion concentration: $\text{pH} = -\log_{10}[\text{H}^{+}]$.

$\text{pH} = -\log_{10}[\text{H}^{+}]$ — given in the data booklet.

$[\text{H}^{+}] = 10^{-\text{pH}}$ — the rearranged given equation.

The ionic product of water, $K_{w} = [\text{H}^{+}][\text{OH}^{-}] = 1.0\times10^{-14}$ at 25 °C.

pH < 7 acidic · pH = 7 neutral · pH > 7 basic (alkaline), at 25 °C.

[H_{+}] changes by a factor of **10** (pH is a log scale).

Strong = **fully** dissociated into ions; weak = only **partially** dissociated.

No — strength is the **degree of dissociation**; concentration is the amount dissolved.

Single arrow, e.g. $\text{HCl} \rightarrow \text{H}^{+} + \text{Cl}^{-}$ (full dissociation).

Equilibrium arrows, e.g. $\text{CH}_{3}\text{COOH} \rightleftharpoons \text{H}^{+} + \text{CH}_{3}\text{COO}^{-}$ (partial).

Strong acid has a **lower pH**, **higher conductivity** and a **faster** reaction (more H_{+} ions).

It is fully dissociated, so it gives a **higher [H_{+}]**, and a higher [H_{+}] means a lower pH.

The reaction of an **acid with a base** to give a **salt and water**; the H⁺ and OH⁻ cancel out.

The ionic compound formed when the **H⁺** of an acid is replaced by a **metal ion** (or NH_{4}⁺).

**salt + hydrogen** (e.g. Mg + 2HCl → MgCl_{2} + H_{2}).

**salt + water** (neutralisation; the base is a metal oxide or hydroxide).

**salt + water + carbon dioxide** (e.g. 2HCl + CaCO_{3} → CaCl_{2} + H_{2}O + CO_{2}).

A **chloride** (e.g. NaCl, MgCl_{2}).

A **sulfate** (e.g. Na_{2}SO_{4}, MgSO_{4}).

A **nitrate** (e.g. NaNO_{3}, Ca(NO_{3})_{2}).

**Hydrogen** gives a squeaky **'pop'** with a lit splint.

**Carbon dioxide** turns **limewater milky** (cloudy).

It is **diprotic** — it provides **two H⁺**, so it neutralises two 1+ bases: 2NaOH + H_{2}SO_{4} → Na_{2}SO_{4} + 2H_{2}O.

Balance the **ionic charges** (e.g. Mg²⁺ with Cl⁻ → MgCl_{2}), then balance the whole equation.

The **acid-dissociation constant** — the equilibrium constant for HA ⇌ H⁺ + A⁻: $K_a = \dfrac{[H^+][A^-]}{[HA]}$.

A **stronger** weak acid — the dissociation equilibrium lies further to the right (more H⁺).

The **base-dissociation constant** for B + H_{2}O ⇌ BH⁺ + OH⁻; a larger K_{b} = a stronger weak base.

$\text{p}K_a = -\log_{10} K_a$ — the negative logarithm of the acid-dissociation constant.

A **stronger** acid — the minus sign flips the scale, so the lowest pK_{a} is the strongest acid.

Line up their **pK_{a}** values; the **lowest pK_{a}** is the strongest (largest K_{a}).

For a conjugate pair, $K_a \times K_b = K_w$ (= 1.0 × 10⁻¹⁴ at 298 K).

For a conjugate pair, $\text{p}K_a + \text{p}K_b = 14$ at 298 K.

$K_b = \dfrac{K_w}{K_a}$ — divide K_{w} by the acid's K_{a}.

$[H^+] = \sqrt{K_a\,c}$, then $\text{pH} = -\log_{10}[H^+]$.

**1.** [H⁺] = [A⁻] (acid is the only source of H⁺). **2.** [HA] ≈ c (dissociation is negligible).

$[OH^-] = \sqrt{K_b\,c}$ → pOH = −log[OH⁻], then **pH = 14 − pOH** (298 K).

**No** — like K_{c}, you write the expressions yourself; only K_{w} = 1.0 × 10⁻¹⁴ is a booklet constant.

A solution that **resists pH change** when a small amount of acid or alkali is added.

A **weak acid + its conjugate base** (e.g. ethanoic acid + sodium ethanoate), both present in appreciable amounts.

A **weak base + its conjugate acid** (e.g. ammonia + ammonium chloride).

The **conjugate base** reacts with it: A⁻ + H⁺ → HA (so the ratio [A⁻]/[HA] barely changes).

The **weak acid** reacts with it: HA + OH⁻ → A⁻ + H_{2}O.

$\text{pH} = \text{p}K_a + \log_{10}\dfrac{[A^-]}{[HA]}$ — buffer pH from pK_{a} and the base/acid ratio.

When **[A⁻] = [HA]** (equal amounts), so the log term is zero — the **half-equivalence point**.

**Exactly 7** — the salt formed is neutral.

**Greater than 7** — the salt of a weak acid hydrolyses to a basic solution.

**Less than 7** — the salt of a weak base hydrolyses to an acidic solution.

A **weak acid**, HIn ⇌ H⁺ + In⁻, whose acid form and conjugate base are **different colours**.

Pick one whose **colour-change range (≈ pK_{a}(In)) lies inside the vertical jump** at the equivalence point.

**Phenolphthalein** (range 8.3–10.0) — it changes within the basic vertical jump.

**Methyl orange** (range 3.1–4.4) — it changes within the acidic vertical jump.