How acids behave

An acid in solution has a pH below 7 and turns litmus red. Cambridge wants you to know its three characteristic reactions:

• Acid + metal → salt + hydrogen
  e.g. Mg + 2HCl → MgCl₂ + H₂
• Acid + base (or alkali) → salt + water  (neutralisation)
  e.g. 2HCl + CuO → CuCl₂ + H₂O
• Acid + carbonate → salt + water + carbon dioxide
  e.g. 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂

Watch out: the salt formed takes its name from the acid: hydrochloric → chloride, sulfuric → sulfate, nitric → nitrate. The metal/base just supplies the positive ion.

Product of an acid reaction

How bases behave

A base is an oxide or hydroxide of a metal; it neutralises an acid. A soluble base is called an alkali (e.g. NaOH, KOH, aqueous ammonia) — these turn litmus blue and have pH above 7.

Characteristic reactions of bases/alkalis:

• Base/alkali + acid → salt + water (neutralisation).
• Alkali + ammonium salt → salt + ammonia + water.
  e.g. NaOH + NH₄Cl → NaCl + NH₃ + H₂O — this releases ammonia gas, used in the test for NH₄⁺.

Base vs alkali: all alkalis are bases, but only the soluble bases are alkalis. Copper(II) oxide is a base but not an alkali because it does not dissolve in water.

Alkali + ammonium salt

Indicators and pH

An indicator changes colour to show whether a solution is acidic or alkaline. You must know these three:

• Litmus: red in acid, blue in alkali.
• Thymolphthalein: colourless in acid, blue in alkali.
• Methyl orange: red in acid, yellow in alkali.

Universal indicator gives a whole range of colours and a number on the pH scale (0–14):

Watch out: pH 7 is neutral — pure water and many salt solutions sit here. Lower pH = more acidic; higher pH = more alkaline. Neutralisation moves the pH towards 7.

Reading an indicator

Acids and alkalis as H⁺ / OH⁻

Cambridge describes acids and alkalis in terms of ions (Core) and protons (Supplement):

• An acid is a proton (H⁺) donor Supplement: aqueous acids contain H⁺ ions.
• A base is a proton (H⁺) acceptor Supplement: aqueous alkalis contain OH⁻ ions.
• Neutralisation is simply: H⁺ + OH⁻ → H₂O.

Key idea: this is why every neutralisation makes water — the H⁺ from the acid combines with the OH⁻ from the alkali.

Why do only some acid reactions fizz?

Strong vs weak acids Supplement

"Strong/weak" describes how fully an acid ionises (dissociates) in water — not how concentrated it is:

• Strong acid: completely ionised in solution — all the molecules release H⁺. e.g. hydrochloric, nitric, sulfuric acid. (HCl → H⁺ + Cl⁻)
• Weak acid: only partially ionised — a reversible reaction, so only a few H⁺ are released. e.g. ethanoic, citric, carbonic acid. (CH₃COOH ⇌ H⁺ + CH₃COO⁻)

Watch out — strong ≠ concentrated. Strong/weak is about degree of dissociation; concentrated/dilute is about how much acid is dissolved. A dilute strong acid can have a lower pH than a concentrated weak acid.

Strong or concentrated? Supplement

Classifying oxides

Oxides are grouped by how they react:

• Basic oxides — metal oxides; react with acids to give salt + water. e.g. CuO, CaO.
• Acidic oxides — non-metal oxides; react with bases and often form acids in water. e.g. CO₂, SO₂.
• Amphoteric oxides Supplement — react with both acids and bases to give salt + water. e.g. Al₂O₃, ZnO.
• Neutral oxides (extra — not in the syllabus) — react with neither acids nor bases. e.g. H₂O, CO, NO.

Watch out: "amphoteric" means it reacts with both acids and bases — don't confuse it with "neutral", which reacts with neither.

Classify the oxide

Why add too much on purpose?

Which salts dissolve?

You must learn the Cambridge solubility rules — they tell you whether a salt can be made by crystallisation or must be made by precipitation:

• All sodium, potassium and ammonium salts are soluble.
• All nitrates are soluble.
• Most chlorides are soluble — except silver chloride and lead(II) chloride.
• Most sulfates are soluble — except barium, calcium and lead(II) sulfate.
• Most carbonates and hydroxides are insoluble — except those of sodium, potassium and ammonium (and hydroxides of calcium, slightly).

Why it matters: a soluble salt is made in solution then crystallised. An insoluble salt is made by precipitation — mixing two solutions so the salt drops out as a solid.

Making a soluble salt

To make a soluble salt, react a dilute acid with an insoluble reactant, then crystallise:

• Acid + excess metal (for reactive metals) — or acid + excess insoluble base/carbonate.
• Add the solid in excess so all the acid reacts, then filter off the leftover solid.
• Crystallise: evaporate the filtrate to the point of crystallisation, then let it cool so crystals form; filter and dry.
• Crystals are often hydrated (chemically combined with water); an anhydrous substance contains no water. Supplement The water of crystallisation is the water in hydrated crystals, e.g. CuSO₄•5H₂O, CoCl₂•6H₂O.

If the base is a soluble alkali (e.g. NaOH), you cannot add it in excess and filter, so you use the titration method: titrate acid against a measured volume of alkali (with an indicator) to find the exact volume needed, then repeat without indicator and crystallise.

Watch out: the titration method is the route for a soluble salt made from a soluble base (alkali) — you need it because excess-and-filter won't work when everything is dissolved.

Choose the prep route

Why add too much on purpose?

Insoluble salts by precipitation

An insoluble salt is made by precipitation: mix two soluble solutions that each contain one of the needed ions. The insoluble salt forms as a solid, which you filter, wash and dry.

Watch out: reach for precipitation whenever the target salt is insoluble (use the solubility rules). Soluble salts are made and crystallised instead.

Soluble or insoluble?

Match each salt on the left to its solubility (and therefore its prep route) on the right.

How can a lab name a mystery powder?

Tests for cations

Add aqueous sodium hydroxide (then a little, then excess) and aqueous ammonia to identify metal ions by the coloured hydroxide precipitate:

A flame test also identifies some cations by flame colour: Li⁺ crimson (red), Na⁺ yellow, K⁺ lilac, Ca²⁺ orange-red, Ba²⁺ light green, Cu²⁺ blue-green.

Identify the ion

Read each observation and tap the ion it reveals.

Tests for anions

• Carbonate (CO₃²⁻): add dilute acid → fizzes; bubble the gas through limewater → turns milky (CO₂).
• Chloride (Cl⁻): add dilute nitric acid, then silver nitrate → white precipitate (AgCl). (bromide → cream; iodide → yellow)
• Sulfate (SO₄²⁻): add dilute nitric acid, then aqueous barium nitrate → white precipitate (BaSO₄).
• Sulfite (SO₃²⁻): add acidified aqueous potassium manganate(VII) → it turns from purple to colourless.
• Nitrate (NO₃⁻): add aqueous sodium hydroxide and aluminium foil, warm → ammonia gas given off (turns damp red litmus blue).

Watch out: the white precipitate of chloride (with silver nitrate) and of sulfate (with barium) look alike — it's the added reagent that tells them apart.

Identify the anion

Identifying gases

• Ammonia (NH₃): turns damp red litmus paper blue.
• Carbon dioxide (CO₂): turns limewater milky/cloudy.
• Chlorine (Cl₂): bleaches damp litmus paper (turns it white).
• Hydrogen (H₂): lit splint gives a squeaky pop.
• Oxygen (O₂): relights a glowing splint.
• Sulfur dioxide (SO₂): turns acidified aqueous potassium manganate(VII) from purple to colourless.

Watch out: a lit splint with a pop = hydrogen; a glowing splint that relights = oxygen — don't mix these two up.

Match the gas test

Tap a gas, then tap the box for the result it gives.

How can a lab name a mystery powder?

The key facts to know

You've covered the key ideas in Cambridge IGCSE 0620 Topic 7 — Acids, bases and salts. Press Finish to see your score.