The full lecture — every idea is shown on the right as an everyday object you already know: a lemon, an antacid tablet, a pool-test kit, a car's shock absorber. Scroll down; the panel keeps pace and the narration reads along with you.
You meet acids and bases every day. Acids taste sour, turn litmus red, and release H⁺ ions (really H₃O⁺) in water — the citric acid in a lemon, the acetic acid in vinegar, the carbonic and phosphoric acid in a cola. Bases feel soapy, turn litmus blue, and release OH⁻ ions — the lye in soap, the ammonia in a cleaner, the hypochlorite in bleach.
| Acids | Bases |
|---|---|
| sour; turn litmus red | bitter, soapy; turn litmus blue |
| release H⁺ (H₃O⁺) in water | release OH⁻ in water |
| react with metals → H₂ gas | react with acids → salt + water |
So an acid and a base are really opposites: one floods the water with H⁺, the other with OH⁻, and the whole chapter is about how chemists measure and balance the two.
Three theories define an acid, from narrow to broad. Arrhenius: an acid gives H⁺ and a base gives OH⁻ in water. Brønsted–Lowry: an acid is a proton (H⁺) donor and a base is a proton acceptor. Lewis: an acid accepts an electron pair and a base donates one — broad enough to cover BF₃ (no hydrogen at all).
| Theory | Acid | Base |
|---|---|---|
| Arrhenius | gives H⁺ | gives OH⁻ |
| Brønsted | H⁺ donor | H⁺ acceptor |
| Lewis | e⁻-pair acceptor | e⁻-pair donor |
When an acid hands its proton over, what is left behind is its conjugate base; the base that caught the proton becomes a conjugate acid. They differ by just one H⁺.
Strength is the degree of ionisation — how much of the acid splits into ions in water, not how concentrated it is. A strong acid (HCl, H₂SO₄, HNO₃) ionises almost completely; a weak acid (CH₃COOH, H₂CO₃) only partly. The same is true for bases: NaOH and KOH are strong, NH₃ is weak.
Even pure water ionises a tiny bit: H₂O ⇌ H⁺ + OH⁻. The product of the two is fixed — the ionic product of water.
Because those numbers are awkward, we use a log scale. pH = −log[H⁺], and pOH = −log[OH⁻], with pH + pOH = 14. Each step on the scale is a ten-fold change in acidity — exactly the colour chart on a swimming-pool or soil test kit.
For a weak acid we need a number for how weak. Its ionisation reaches equilibrium, and the equilibrium constant is the acid dissociation constant Ka.
A smaller Ka (a larger pKa) means less ionisation — a weaker acid. The same logic gives Kb for a weak base, and for any conjugate pair Ka × Kb = Kw. The simple litmus test only says "acid or base"; Ka is the precise version chemists and pharmacists rely on.
When an acid meets a base they cancel out: acid + base → salt + water. The net ionic reaction is simply H⁺ + OH⁻ → H₂O (ΔH ≈ −57 kJ/mol). A salt is what is left when the H⁺ of the acid is replaced by a metal (or NH₄⁺) ion — this is exactly how an antacid tablet calms stomach acid.
But the salt is not always neutral. Its ions can react with water (hydrolysis), and the weaker parent decides the pH.
| Salt from | Example | Solution |
|---|---|---|
| strong acid + strong base | NaCl | neutral (pH 7) |
| strong acid + weak base | NH₄Cl | acidic (pH < 7) |
| weak acid + strong base | CH₃COONa | basic (pH > 7) |
A buffer resists a change in pH when small amounts of acid or base are added — it is the chemical version of a car's shock absorber, soaking up the bump. It is made from a weak acid plus its salt (acidic buffer) or a weak base plus its salt (basic buffer).
Add a splash of acid to plain water and the pH crashes; add it to a buffer and the pH barely moves, because the salt mops up the extra H⁺.