Every key formula you need for Cambridge IGCSE Chemistry 0620 is gathered here on one page, organised by topic so you can find exactly what a question is asking for in seconds. The mole calculations that students find hardest are broken down into simple step-by-step methods you can follow in the exam. Keep this sheet open while you revise, then test yourself with our free IGCSE Chemistry practice tests.
| Formula / Equation | What it means / How it is used |
|---|
| Relative atomic mass, Ar | The average mass of one atom of an element compared with 1/12 of the mass of a carbon-12 atom; read it straight from the Periodic Table. |
| Relative molecular mass, Mr | Add together the Ar values of every atom in the formula, e.g. Mr of CO2 = 12 + (16 × 2) = 44. |
| Amount of substance (mol) = mass (g) / Mr | The central mole equation: divide the mass you are given by the Mr to find the number of moles. |
| Mass (g) = moles × Mr | Rearrange the mole equation when you know the moles and need the mass in grams. |
| Number of particles = moles × 6.02 × 1023 | One mole of any substance contains 6.02 × 1023 particles (the Avogadro constant) — atoms, molecules or ions. |
| Concentration (mol/dm3) = moles / volume (dm3) | For solutions: divide the moles of solute by the volume of solution, remembering to convert cm3 to dm3 (divide by 1000). |
| Concentration (g/dm3) = mass (g) / volume (dm3) | Use when the solute is measured by mass rather than in moles. |
| Volume of gas (dm3) = moles × 24 | At room temperature and pressure, one mole of any gas occupies 24 dm3 (24,000 cm3). |
| Percentage yield = (actual yield / theoretical yield) × 100 | Compares the mass of product you actually obtained with the maximum mass the calculation predicted. |
| Percentage purity = (mass of pure substance / total mass of sample) × 100 | Shows how much of a sample is the substance you want, rather than impurities. |
| Percentage composition = (Ar × number of atoms of the element / Mr of compound) × 100 | Finds the percentage by mass of one element inside a compound. |
| Empirical formula method: (1) write each element’s mass (or %), (2) divide each by its Ar to get moles, (3) divide every answer by the smallest value, (4) round to the simplest whole-number ratio and write the formula | The four-step routine that works for every empirical formula question, whether you are given masses or percentages. |
| Reacting-mass method: (1) convert the known mass to moles using mass / Mr, (2) use the mole ratio from the balanced equation, (3) convert back to mass using moles × Mr | The standard three steps for finding the mass of a product formed, or the mass of a reactant needed. |
| Titration method: moles = concentration × volume (dm3); apply the ratio from the balanced equation; then calculate the unknown concentration, volume or mass | Find the moles of the solution you know, use the equation ratio to get the moles of the other solution, then scale up to the answer. |
| Formula / Equation | What it means / How it is used |
|---|
| Acid + metal → salt + hydrogen | Only metals above hydrogen in the reactivity series react; the salt formed depends on the acid used. |
| Acid + base (alkali) → salt + water | Neutralisation: the base cancels out the acid, leaving a salt solution. |
| Acid + carbonate → salt + water + carbon dioxide | The fizzing (effervescence) is carbon dioxide, which turns limewater milky. |
| Acid + ammonia → ammonium salt | No water is formed here; the product is a fertiliser such as ammonium sulfate or ammonium nitrate. |
| Hydrocarbon + oxygen → carbon dioxide + water | Complete combustion of any hydrocarbon fuel in a plentiful supply of air. |
| Metal + oxygen → metal oxide | Metals tarnish or burn in air to form their oxides, e.g. 2Mg + O2 → 2MgO. |
| Reactive metal + cold water → metal hydroxide + hydrogen | Group 1 metals and calcium react with cold water to give an alkaline hydroxide solution. |
| Metal + steam → metal oxide + hydrogen | Less reactive metals such as magnesium, zinc and iron need steam, and give the oxide instead of the hydroxide. |
| Formula / Equation | What it means / How it is used |
|---|
| Molten compounds: positive metal ions move to the cathode (−) and are discharged as metal; negative non-metal ions move to the anode (+) and are discharged as the non-metal | With a molten ionic compound only the compound’s own two ions are present, so the products are predictable. |
| Aqueous solutions: at the cathode, hydrogen forms unless the metal is below hydrogen in the reactivity series; at the anode, a concentrated halide gives the halogen, otherwise oxygen forms | Water provides H+ and OH− ions too, so the reactivity series decides which ion wins at each electrode. |
| Pb2+ + 2e− → Pb | Example reduction half-equation: lead(II) ions gain electrons at the cathode and deposit as lead metal. |
| 2H+ + 2e− → H2 | Hydrogen ions gain electrons at the cathode to form hydrogen gas. |
| 2Cl− → Cl2 + 2e− | Oxidation half-equation: chloride ions lose electrons at the anode to form chlorine gas. |
| 4OH− → O2 + 2H2O + 4e− | Hydroxide ions are discharged at the anode in dilute aqueous solutions, releasing oxygen. |
| Charge passed = current × time | The larger the charge passed through the electrolyte, the more product is formed at each electrode. |
| Formula / Equation | What it means / How it is used |
|---|
| Energy change = energy to break bonds − energy released making bonds | Add up all the bond energies broken in the reactants and subtract all the bond energies formed in the products. |
| Exothermic reaction: energy is released to the surroundings | The temperature of the surroundings rises; combustion and neutralisation are common examples. |
| Endothermic reaction: energy is absorbed from the surroundings | The temperature of the surroundings falls; thermal decomposition is a common example. |
| Energy transferred (J) = mass of water (g) × 4.2 × temperature change (°C) | Simple calorimetry: the heat given out by a fuel or reaction is found from the temperature rise of a known mass of water. |
| Formula / Equation | What it means / How it is used |
|---|
| pH scale: 1–6 acidic, 7 neutral, 8–14 alkaline; universal indicator goes red/orange in acid, green at neutral, blue/purple in alkali | Lower pH means a stronger acid (more H+ ions); higher pH means a stronger alkali (more OH− ions). |
| H+ + OH− → H2O | The ionic equation for every neutralisation between an acid and an alkali. |
| Solubility rules: all nitrates are soluble; most sulfates are soluble except barium sulfate and calcium sulfate; chlorides are soluble except silver chloride and lead(II) chloride; carbonates are insoluble except Group 1 and ammonium carbonates | Use these rules to predict whether a salt can be made by precipitation and whether a precipitate will form when solutions mix. |
| Formula / Equation | What it means / How it is used |
|---|
| Higher temperature → faster rate | Particles move faster and collide more often with enough energy to react. |
| Higher concentration (or pressure for gases) → faster rate | More particles are packed into the same volume, so collisions happen more frequently. |
| Larger surface area of a solid → faster rate | Powdering a solid exposes more particles to the other reactant, increasing the collision frequency. |
| A catalyst speeds up a reaction without being used up | It provides an alternative pathway with lower activation energy and can be recovered unchanged afterwards. |
| Mean rate = change in mass or volume / time taken | Calculate the average rate from experimental results, e.g. volume of gas collected divided by the time it took. |
| Formula / Equation | What it means / How it is used |
|---|
| Alkanes: CnH2n+2 | General formula for the saturated hydrocarbon family, e.g. propane is C3H8. |
| Alkenes: CnH2n | General formula for the unsaturated hydrocarbons containing one C=C bond, e.g. ethene is C2H4. |
| Alcohols: CnH2n+1OH | General formula for alcohols, e.g. ethanol is C2H5OH. |
| Carboxylic acids: CnH2n+1COOH | General formula for carboxylic acids, e.g. ethanoic acid is CH3COOH. |
| Functional groups: −OH (alcohol), −COOH (carboxylic acid), C=C (alkene) | The functional group gives a family of compounds its characteristic chemical reactions. |
| Addition: alkenes add atoms across the C=C bond, which breaks to a single bond; Substitution: alkanes swap one atom for another, needing ultraviolet light | A quick way to tell the two reaction types apart: addition needs a double bond, substitution swaps atoms one-for-one. |
| Carboxylic acid + alcohol → ester + water | Esterification, carried out by warming with a concentrated sulfuric acid catalyst; esters have sweet, fruity smells. |
| Formula / Equation | What it means / How it is used |
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| Oxidation = loss of electrons, or gain of oxygen | Remember OIL: Oxidation Is Loss of electrons. |
| Reduction = gain of electrons, or loss of oxygen | Remember RIG: Reduction Is Gain of electrons; oxidation and reduction always happen together. |
| Reactivity series: K, Na, Ca, Mg, Al, Zn, Fe, Pb, Cu, Ag (most reactive → least reactive) | Potassium down to silver: the higher a metal sits, the more vigorously it reacts and the harder its oxide is to reduce. |
| Displacement rule: a more reactive metal displaces a less reactive metal from its compound | Use the series to predict reactions, e.g. zinc displaces copper from copper(II) sulfate solution, but copper cannot displace zinc. |
This formula sheet is an original revision aid created by QuaidianPoint. It is not affiliated with, endorsed by, or connected to Cambridge Assessment International Education.
