Two Definitions of an Acid
The older definition is Arrhenius: an acid produces H⁺ ions in water, a base produces OH⁻ ions in water. Sodium hydroxide is an Arrhenius base because it releases OH⁻ directly. But ammonia, NH₃, contains no OH⁻ to release, yet behaves as a base: the Arrhenius picture has a gap.
The Brønsted-Lowry definition, the CED’s, is built on proton transfer: an acid is a proton (H⁺) donor, a base a proton acceptor. A proton can only be donated if something accepts it, so acid and base always appear together.
In Water, the Proton Travels as H₃O⁺
A bare H⁺ never exists alone in water: it attaches to a water molecule, forming hydronium, H₃O⁺. H⁺(aq) is shorthand for H₃O⁺(aq).
Hydrogen chloride dissolving: HCl + H₂O → H₃O⁺ + Cl⁻. HCl donates the proton (the acid); water accepts it (the base).
With ammonia, water’s role reverses: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, water donating the proton and acting as the acid. The OH⁻ was made by the transfer, not released from the ammonia.
No acid stronger than H₃O⁺ survives in water: it hands its proton to water instead, making H₃O⁺ the strongest acid present in aqueous solution. In NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺ the equilibrium therefore lies well to the left: only a small fraction of NH₄⁺ ions have donated at any moment.
Conjugate Pairs Differ by Exactly One Proton
When an acid donates its proton, what remains is its conjugate base. When a base accepts a proton, what forms is its conjugate acid. A conjugate pair is two species differing by exactly one H⁺, their charges differing by one too.
Conjugate base: remove one H⁺ (H₂SO₄ gives HSO₄⁻, not SO₄²⁻). Conjugate acid: add one H⁺ (NH₃ gives NH₄⁺). Every proton transfer contains two pairs: in HCl + H₂O → H₃O⁺ + Cl⁻, HCl/Cl⁻ and H₃O⁺/H₂O.
A weak acid together with its conjugate base forms a buffer: the acid neutralises added OH⁻, the conjugate base added H⁺, so the pH changes far less than in water alone.
Amphiprotic Species Play Both Roles
An amphiprotic species donates a proton in one reaction, accepts one in another; its partner decides which. Water accepts from HCl, donates to NH₃.
Hydrogencarbonate, HCO₃⁻, is amphiprotic: with water it acts as a base, HCO₃⁻ + H₂O ⇌ H₂CO₃ + OH⁻; with hydroxide as an acid, HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O. HSO₄⁻ and H₂PO₄⁻ behave the same way.
Strong Means Fully Ionised, Weak Means Barely
A strong acid donates its proton completely: no intact acid molecules remain. The strong acids are HCl, HBr, HI, HNO₃, HClO₄ and the first proton of H₂SO₄; single arrow. Strong bases ionise completely too: Group 1 hydroxides, such as NaOH and KOH, and Ba(OH)₂ are entirely cations and OH⁻ ions.
A weak acid ionises only slightly: CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻ is an equilibrium, and in 0.10 mol/L solution only about one molecule in a hundred has ionised at any moment. A weak base accepts protons only partially: most NH₃ in solution is still NH₃.
Strength is the fraction ionised; concentration is the amount dissolved. Independent ideas: a concentrated weak acid can hold more H₃O⁺ than a very dilute strong acid.
Neutralisation: A Proton Meets Hydroxide
An acid reacting with an OH⁻ base produces a salt and water: HCl + NaOH → NaCl + H₂O, the salt pairing the base’s cation with the acid’s anion.
For any strong acid with any strong base the net ionic equation is Topic 4.2’s: H⁺(aq) + OH⁻(aq) → H₂O(l); the spectators differ, the proton transfer never does. With hydronium: H₃O⁺(aq) + OH⁻(aq) → 2H₂O(l).
If the acid is weak it stays whole in the net ionic equation, because most of it is whole in the flask: CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l). If the base is ammonia, no OH⁻ appears: NH₃(aq) + H⁺(aq) → NH₄⁺(aq); with hydrochloric acid the only product is the salt NH₄Cl. Neutralisation is the proton transfer, not the appearance of water.
Strength changes how ionised an acid is, not its total proton capacity: a mole of CH₃COOH neutralises exactly as much NaOH as a mole of HCl.
Polyprotic Acids Donate One Proton at a Time
A polyprotic acid has more than one donatable proton: H₂SO₄ has two, H₃PO₄ three. They leave one at a time, in separate steps, never together.
Sulfuric acid’s first step is complete, H₂SO₄ + H₂O → H₃O⁺ + HSO₄⁻; the second is an equilibrium, HSO₄⁻ + H₂O ⇌ H₃O⁺ + SO₄²⁻. Hence the conjugate base of H₂SO₄ is HSO₄⁻: one proton removed, one step taken.
Carbonate runs the same staircase in reverse: CO₃²⁻ + H⁺ → HCO₃⁻, then HCO₃⁻ + H⁺ → H₂O + CO₂ (through H₂CO₃, which falls apart). Together: 2H⁺(aq) + CO₃²⁻(aq) → H₂O(l) + CO₂(g), the net ionic equation behind the Topic 4.7 result.
Worked Examples
Example 1: labelling a proton transfer. In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, NH₃ accepts a proton: it is the base, and NH₄⁺ is its conjugate acid. H₂O donates the proton: it is the acid, and OH⁻ its conjugate base. Two pairs, each differing by one proton.
Example 2: three neutralisations, three net ionic equations. Nitric acid with potassium hydroxide: K⁺ and NO₃⁻ spectate: H⁺(aq) + OH⁻(aq) → H₂O(l). Acetic acid with sodium hydroxide: Na⁺ spectates, the weak acid stays whole: CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l). Hydrochloric acid with ammonia: Cl⁻ spectates, the proton lands on ammonia: NH₃(aq) + H⁺(aq) → NH₄⁺(aq). One family, three net equations, because what is present in each flask differs.
Example 3: carbonate, one proton at a time. A solution contains 0.0200 mol of CO₃²⁻. Adding 0.0200 mol of H⁺ runs the first step only, CO₃²⁻ + H⁺ → HCO₃⁻: 0.0200 mol of HCO₃⁻, no gas. A second 0.0200 mol of H⁺ runs the second step, HCO₃⁻ + H⁺ → H₂O + CO₂, and the 0.0200 mol of CO₂ escapes: 0.0200 × 22.4 = 0.448 L at STP. Total acid: 0.0400 mol, two protons per carbonate.
Traps That Cost Marks in Topic 4.8
Removing both protons at once. The conjugate base of H₂SO₄ is HSO₄⁻. A conjugate pair differs by exactly one proton, so SO₄²⁻, two protons away, cannot be the answer.
Miscounting the gap in a pair. H₂CO₃ and CO₃²⁻ are not a conjugate pair: two protons apart. The real pairs are H₂CO₃/HCO₃⁻ and HCO₃⁻/CO₃²⁻. H₂O and OH⁻, by contrast, differ by one proton and are a genuine pair.
Splitting weak acids and bases in ionic equations. CH₃COOH and NH₃ are written whole. Only strong acids and strong bases appear as separated ions, exactly as in Topic 4.2.
Confusing strong with concentrated. Strong describes how completely a substance ionises; concentrated, how much is dissolved. Dilute HCl is still strong; concentrated CH₃COOH is still weak.
Expecting water, or OH⁻, in every neutralisation. NH₃ + HCl → NH₄Cl forms a salt, no water, and its net ionic equation has no OH⁻: the base took the proton itself.
Pinning a permanent label on a substance. Water is the base when HCl dissolves in it, the acid when NH₃ does; HCO₃⁻ is a base toward water, an acid toward OH⁻. Assign roles from the equation in front of you.
Forgetting the second water. H₃O⁺(aq) + OH⁻(aq) → 2H₂O(l): written with hydronium, neutralisation makes two waters, not one.
Key Relationships
- Brønsted-Lowry: an acid donates H⁺, a base accepts H⁺; every acid-base reaction is a proton transfer needing both partners. Arrhenius is the water-only version: acid gives H⁺, base gives OH⁻
- In water the proton travels as H₃O⁺ (H⁺(aq) is shorthand), the strongest acid that can exist in aqueous solution
- A conjugate pair differs by exactly one H⁺: base + H⁺ gives its conjugate acid, acid − H⁺ gives its conjugate base, the charge shifting by one each way
- Every proton transfer contains two conjugate pairs; a weak acid with its conjugate base forms a buffer resisting pH change
- Strong acids and bases ionise completely and are written as ions; weak ones reach equilibrium and stay mostly intact; strength (fraction ionised) is not concentration (amount dissolved)
- Neutralisation net ionic equations: H⁺ + OH⁻ → H₂O (strong with strong); CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O (weak acid); NH₃ + H⁺ → NH₄⁺ (ammonia)
- Polyprotic acids donate stepwise, H₂SO₄ → HSO₄⁻ → SO₄²⁻; carbonate accepts stepwise, CO₃²⁻ → HCO₃⁻ → H₂O + CO₂, so acid plus carbonate gives a salt, water and carbon dioxide
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