IGCSE Chemistry Key Definitions: The Ultimate Revision Sheet

Chemistry at IGCSE level demands that you know your definitions inside and out. Examiners frequently award marks for the precise use of scientific terminology, and a single missing keyword can cost you. This comprehensive definition sheet covers all the essential terms from the Cambridge IGCSE Chemistry (0620) syllabus, organised by topic to support your revision.

Bookmark this page or print it — it is your essential Chemistry reference for exam preparation.

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1. States of Matter

Solid — A state of matter with a fixed shape and fixed volume. Particles are closely packed in a regular arrangement and vibrate about fixed positions.

Liquid — A state of matter with a fixed volume but no fixed shape. Particles are close together but can move past each other.

Gas — A state of matter with no fixed shape and no fixed volume. Particles are far apart, move rapidly and randomly, and fill the container they are in.

Melting — The change of state from solid to liquid. Particles gain energy and overcome some of the forces holding them in fixed positions.

Boiling — The change of state from liquid to gas throughout the liquid at a fixed temperature (the boiling point).

Evaporation — The change of state from liquid to gas that occurs at the surface of a liquid at temperatures below the boiling point.

Condensation — The change of state from gas to liquid when a gas is cooled.

Freezing — The change of state from liquid to solid when a liquid is cooled to its freezing point.

Sublimation — The direct change of state from solid to gas without passing through the liquid state (e.g., solid carbon dioxide).

Diffusion — The net movement of particles from a region of higher concentration to a region of lower concentration due to the random motion of particles.

Brownian motion — The random, erratic movement of small particles suspended in a fluid, caused by collisions with the fast-moving particles of the fluid. Provides evidence for the kinetic particle model.

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2. Atomic Structure and the Periodic Table

Atom — The smallest particle of an element that can take part in a chemical reaction.

Element — A substance that consists of only one type of atom and cannot be broken down into simpler substances by chemical means.

Compound — A substance formed when two or more elements are chemically combined in a fixed ratio.

Mixture — Two or more substances that are not chemically combined and can be separated by physical methods.

Proton — A subatomic particle found in the nucleus with a relative mass of 1 and a relative charge of +1.

Neutron — A subatomic particle found in the nucleus with a relative mass of 1 and no charge.

Electron — A subatomic particle that orbits the nucleus in energy levels (shells) with negligible mass and a relative charge of −1.

Atomic number (Z) — The number of protons in the nucleus of an atom. It defines the element.

Mass number (A) — The total number of protons and neutrons in the nucleus of an atom.

Isotope — Atoms of the same element with the same number of protons but different numbers of neutrons, and therefore different mass numbers.

Relative atomic mass (Ar) — The weighted average mass of an atom of an element compared to 1/12 the mass of a carbon-12 atom.

Relative molecular mass (Mr) — The sum of the relative atomic masses of all the atoms in a molecule.

Relative formula mass — The sum of the relative atomic masses of all the atoms in the formula of a compound (used for ionic compounds).

Electron configuration — The arrangement of electrons in the energy levels (shells) of an atom (e.g., sodium is 2,8,1).

Period — A horizontal row in the periodic table. Elements in the same period have the same number of electron shells.

Group — A vertical column in the periodic table. Elements in the same group have the same number of electrons in their outer shell and similar chemical properties.

Alkali metals (Group I) — Highly reactive metals (e.g., lithium, sodium, potassium) that form alkaline solutions when they react with water and have one electron in their outer shell.

Halogens (Group VII) — Reactive non-metals (e.g., fluorine, chlorine, bromine, iodine) that have seven electrons in their outer shell and form salts with metals.

Noble gases (Group VIII/0) — Unreactive elements (e.g., helium, neon, argon) with a full outer electron shell, making them very stable.

Transition metals — A block of metallic elements in the centre of the periodic table (e.g., iron, copper, zinc) that often have coloured compounds, variable oxidation states, and catalytic properties.

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3. Chemical Bonding

Ionic bonding — The electrostatic attraction between oppositely charged ions, formed by the transfer of electrons from a metal atom to a non-metal atom.

Ion — An atom or group of atoms that has gained or lost one or more electrons, giving it a positive or negative charge.

Cation — A positively charged ion formed when an atom loses one or more electrons.

Anion — A negatively charged ion formed when an atom gains one or more electrons.

Covalent bonding — The sharing of one or more pairs of electrons between two non-metal atoms so that each atom achieves a stable electron configuration.

Single covalent bond — A bond formed by the sharing of one pair of electrons between two atoms.

Double covalent bond — A bond formed by the sharing of two pairs of electrons between two atoms.

Metallic bonding — The electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons.

Delocalised electrons — Electrons in a metallic structure that are free to move throughout the metal lattice, not associated with any particular atom.

Giant ionic lattice — A regular three-dimensional arrangement of alternating positive and negative ions held together by strong electrostatic forces in all directions (e.g., NaCl).

Simple molecular structure — A structure consisting of small molecules held together by weak intermolecular forces, with strong covalent bonds within each molecule (e.g., water, methane).

Giant covalent structure (macromolecule) — A structure where a very large number of atoms are bonded together by covalent bonds in a regular lattice (e.g., diamond, silicon dioxide).

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4. Stoichiometry and the Mole Concept

Mole — The amount of substance that contains 6.02 × 10²³ particles (Avogadro’s number). One mole of any element has a mass equal to its relative atomic mass in grams.

Avogadro’s number (constant) — 6.02 × 10²³, the number of particles in one mole of a substance.

Molar mass — The mass of one mole of a substance, measured in g/mol.

Empirical formula — The simplest whole-number ratio of atoms of each element present in a compound.

Molecular formula — The actual number of atoms of each element present in one molecule of a compound.

Balanced chemical equation — A symbolic representation of a chemical reaction showing the reactants and products with their correct formulae and balanced coefficients so that the number of atoms of each element is the same on both sides.

Concentration — The amount of solute dissolved in a given volume of solution. Usually measured in g/dm³ or mol/dm³.

Limiting reagent — The reactant that is completely used up in a chemical reaction and determines the maximum amount of product formed.

Excess reagent — The reactant that is not completely used up in a chemical reaction; some remains after the reaction is complete.

Percentage yield — The actual yield as a percentage of the theoretical yield. Percentage yield = (actual yield ÷ theoretical yield) × 100%.

Percentage purity — The mass of the pure substance as a percentage of the total mass of the sample.

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5. Chemical Reactions

Chemical reaction — A process in which one or more substances (reactants) are converted into one or more new substances (products) with different properties.

Exothermic reaction — A reaction that releases energy (usually heat) to the surroundings, causing an increase in temperature. ΔH is negative.

Endothermic reaction — A reaction that absorbs energy from the surroundings, causing a decrease in temperature. ΔH is positive.

Activation energy — The minimum amount of energy that reacting particles must have in order for a successful collision to occur and a reaction to take place.

Rate of reaction — The speed at which a chemical reaction takes place, measured by the change in concentration, mass, or volume of a reactant or product per unit time.

Catalyst — A substance that increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy, without being chemically changed itself.

Collision theory — The theory that reactions occur when particles collide with sufficient energy (≥ activation energy) and the correct orientation.

Reversible reaction — A reaction that can proceed in both the forward and backward directions under suitable conditions, represented by the symbol ⇌.

Equilibrium — The state in a reversible reaction when the rate of the forward reaction equals the rate of the backward reaction and the concentrations of reactants and products remain constant.

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6. Acids, Bases, and Salts

Acid — A substance that produces hydrogen ions (H⁺) when dissolved in water. Acids have a pH less than 7.

Base — A substance that reacts with an acid to form a salt and water. A base is a proton (H⁺) acceptor.

Alkali — A soluble base that produces hydroxide ions (OH⁻) when dissolved in water. Alkalis have a pH greater than 7.

pH scale — A numerical scale from 0 to 14 used to measure the acidity or alkalinity of a solution. pH 7 is neutral, below 7 is acidic, above 7 is alkaline.

Indicator — A substance that changes colour depending on whether it is in an acidic or alkaline solution (e.g., litmus, universal indicator, phenolphthalein).

Neutralisation — The reaction between an acid and a base to form a salt and water. H⁺(aq) + OH⁻(aq) → H₂O(l).

Salt — An ionic compound formed when the hydrogen of an acid is replaced by a metal ion or ammonium ion.

Strong acid — An acid that fully dissociates (ionises) into ions in aqueous solution (e.g., hydrochloric acid, sulfuric acid, nitric acid).

Weak acid — An acid that only partially dissociates into ions in aqueous solution (e.g., ethanoic acid, citric acid).

Titration — A technique used to determine the exact volume of one solution needed to react completely with a known volume of another solution.

Precipitation — A reaction in which two soluble substances react in solution to form an insoluble solid (precipitate).

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7. Oxidation and Reduction (Redox)

Oxidation — The loss of electrons, the gain of oxygen, or the increase in oxidation state of an atom or ion.

Reduction — The gain of electrons, the loss of oxygen, or the decrease in oxidation state of an atom or ion.

Redox reaction — A reaction in which both oxidation and reduction occur simultaneously.

Oxidising agent — A substance that oxidises another substance and is itself reduced (it gains electrons).

Reducing agent — A substance that reduces another substance and is itself oxidised (it loses electrons).

Oxidation state (number) — A number assigned to an atom that indicates the degree of oxidation, i.e., the number of electrons lost, gained, or shared unequally.

Rusting — The corrosion of iron in the presence of both oxygen and water, forming hydrated iron(III) oxide (rust).

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8. Electrolysis

Electrolysis — The decomposition of an ionic compound (electrolyte), when molten or in aqueous solution, by the passage of an electric current.

Electrolyte — A substance that conducts electricity when molten or dissolved in water because it contains free-moving ions.

Electrode — A solid electrical conductor through which current enters or leaves the electrolyte.

Anode — The positive electrode in electrolysis, where anions are attracted and oxidation occurs.

Cathode — The negative electrode in electrolysis, where cations are attracted and reduction occurs.

Electroplating — The process of depositing a thin layer of metal onto an object by electrolysis, using the object as the cathode.

Faraday’s laws of electrolysis — The mass of substance deposited or liberated at an electrode is proportional to the quantity of electric charge passed through the electrolyte.

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9. Metals and Reactivity

Reactivity series — A list of metals arranged in order of their reactivity, from most reactive (e.g., potassium) to least reactive (e.g., gold).

Displacement reaction — A reaction in which a more reactive element takes the place of a less reactive element in a compound.

Extraction of metals — The process of obtaining a metal from its ore. Reactive metals are extracted by electrolysis; less reactive metals are extracted by reduction with carbon.

Ore — A naturally occurring rock or mineral from which a metal can be economically extracted.

Alloy — A mixture of two or more metals, or a metal with a non-metal, that has improved properties compared to the pure metal (e.g., steel, brass, bronze).

Corrosion — The gradual destruction of a metal by chemical reaction with substances in its environment (e.g., oxygen, water, acids).

Galvanising — Coating iron or steel with a layer of zinc to protect it from rusting.

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10. Non-metals and Their Compounds

Hydrogen — A colourless, odourless gas (H₂) that burns with a squeaky pop when tested with a lit splint. It is the lightest element.

Oxygen — A colourless, odourless gas (O₂) that relights a glowing splint. It is essential for combustion and respiration.

Carbon dioxide — A colourless gas (CO₂) that turns limewater milky. It is produced by combustion and respiration and is a greenhouse gas.

Nitrogen — A colourless, odourless, unreactive gas (N₂) that makes up approximately 78% of the atmosphere.

Ammonia — A colourless, pungent gas (NH₃) that turns damp red litmus paper blue. It is alkaline and used in the manufacture of fertilisers.

Sulfur dioxide — A colourless, toxic gas (SO₂) with a pungent smell. It is a major cause of acid rain.

Haber process — The industrial process for manufacturing ammonia from nitrogen and hydrogen at high temperature (~450°C), high pressure (~200 atm), and using an iron catalyst. N₂ + 3H₂ ⇌ 2NH₃.

Contact process — The industrial process for manufacturing sulfuric acid, involving the catalytic oxidation of sulfur dioxide to sulfur trioxide using vanadium(V) oxide as a catalyst.

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11. Organic Chemistry

Organic chemistry — The study of carbon-containing compounds (excluding carbonates, carbon dioxide, and carbon monoxide).

Hydrocarbon — A compound containing only hydrogen and carbon atoms.

Homologous series — A family of organic compounds with the same general formula, similar chemical properties, and a gradual trend in physical properties. Each successive member differs by CH₂.

Alkane — A saturated hydrocarbon with only single covalent bonds between carbon atoms. General formula: CₙH₂ₙ₊₂ (e.g., methane CH₄, ethane C₂H₆).

Alkene — An unsaturated hydrocarbon that contains at least one carbon-carbon double bond. General formula: CₙH₂ₙ (e.g., ethene C₂H₄, propene C₃H₆).

Alcohol — An organic compound containing the hydroxyl functional group (−OH). General formula: CₙH₂ₙ₊₁OH (e.g., ethanol C₂H₅OH).

Carboxylic acid — An organic compound containing the carboxyl functional group (−COOH). General formula: CₙH₂ₙ₊₁COOH (e.g., ethanoic acid CH₃COOH).

Ester — An organic compound formed by the reaction of a carboxylic acid with an alcohol in the presence of an acid catalyst (esterification). Esters have characteristic fruity smells.

Polymer — A large molecule made by joining many small molecules (monomers) together in a repeating pattern.

Monomer — A small molecule that can join together with other monomers to form a polymer.

Addition polymerisation — A reaction in which many unsaturated monomer molecules (alkenes) join together to form a long-chain polymer with no other product.

Condensation polymerisation — A reaction in which monomers join together to form a polymer, with the elimination of a small molecule (usually water) for each bond formed.

Saturated — A molecule that contains only single covalent bonds between carbon atoms (no double or triple bonds).

Unsaturated — A molecule that contains one or more carbon-carbon double or triple bonds.

Isomer — Compounds with the same molecular formula but different structural arrangements of atoms.

Functional group — An atom or group of atoms responsible for the characteristic chemical properties of a homologous series (e.g., −OH in alcohols, C=C in alkenes).

Fractional distillation — The separation of a mixture of liquids (e.g., crude oil) based on differences in their boiling points by heating and condensing at different levels in a fractionating column.

Cracking — The breaking down of long-chain hydrocarbons into shorter, more useful molecules (shorter alkanes and alkenes) by heating with a catalyst (catalytic cracking) or with steam (thermal cracking).

Combustion — A chemical reaction in which a substance reacts rapidly with oxygen, releasing heat and light. Complete combustion produces CO₂ and H₂O; incomplete combustion produces CO and/or carbon (soot).

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12. Experimental Techniques

Filtration — A method used to separate an insoluble solid from a liquid by passing the mixture through filter paper.

Crystallisation — A method used to obtain a pure solid from a solution by evaporating the solvent until the solution becomes saturated and crystals form on cooling.

Distillation — A method used to separate a liquid from a solution or a mixture of liquids by heating to boiling point and condensing the vapour.

Chromatography — A technique used to separate and identify the components of a mixture based on their different solubilities in a solvent. Components travel different distances on the paper or plate.

Rf value — The ratio of the distance travelled by a substance to the distance travelled by the solvent front in chromatography. Rf = distance moved by substance ÷ distance moved by solvent.

Pure substance — A substance that consists of only one type of element or compound and has a sharp, fixed melting and boiling point.

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How to Use This Definition Sheet for Revision

1. Match each section to your syllabus and tick off topics as you revise them.
2. Cover the definitions and try to recall them from the term alone.
3. Write out key definitions — the act of writing helps memory retention.
4. Use past paper mark schemes to see exactly how examiners word definitions.
5. Revise with a partner — test each other and discuss any terms you find tricky.

Remember: in Chemistry exams, definitions often carry 1–2 marks each. Getting them right is one of the easiest ways to boost your overall grade.

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This definition sheet covers the key terms from the Cambridge IGCSE Chemistry (0620) syllabus. Always check your specific exam board’s syllabus for the most up-to-date requirements.

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