Skip to content

Chemistry Glossary

Short, exam-ready definitions for the key terms used across this site's Chemistry pages, with a link back to where each one is taught properly. Covers all of Year 11 Chemistry; Year 12 terms will be added as those pages are built.

Hover for a reminder

The trickiest of these terms are also set up as hover tooltips on every Chemistry content page — hover over (or tap) a dotted-underlined word to see its definition without leaving the page.

Command verbs (Identify, Explain, Evaluate…) aren't here — they work the same way in every subject, so they live in the Command Verb Glossary on Resources.

Jump to: A · B · C · D · E · F · G · H · I · K · L · M · N · O · P · R · S · V

A

  • Absolute zero — 0 K (−273.15 °C) — the lowest possible temperature, where particles would have no kinetic energy. → Properties of Gases
  • Activation energy (\(E_a\)) — the minimum energy colliding particles need for a reaction to happen. The \"hump\" on an energy profile diagram. → Rates of Reactions, Energy Changes
  • Activity series — Metals ranked from most to least reactive (K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > H > Cu > Ag > Au). A metal displaces the ions of any metal below it. → Classifying Chemical Reactions
  • Allotropes — Different structural forms of the same element, e.g. diamond, graphite, graphene and fullerenes are all carbon. → Chemical Bonding
  • Alpha radiation (\(\alpha\)) — a helium nucleus (2 protons + 2 neutrons) emitted from an unstable nucleus. Charge +2, highly ionising, stopped by paper. → Atomic Structure
  • Anion — A negative ion — an atom (or group of atoms) that has gained electrons, e.g. \(\ce{Cl-}\). → Atomic Structure
  • Atomic number (\(Z\)) — the number of protons in an atom. It decides which element the atom is. → Atomic Structure
  • Atomic radius — The size of an atom. Decreases across a period, increases down a group. → Periodic Trends
  • Avogadro constant (\(N_A\)) — \(6.022 \times 10^{23}\), the number of particles in one mole. → Mole Concept

B

  • Background radiation — The low level of radiation around us all the time, from rocks, soil, radon gas, food and space. → Atomic Structure
  • Beta radiation (\(\beta\)) — a fast electron emitted when a neutron in the nucleus turns into a proton. Charge −1, stopped by a few mm of aluminium. → Atomic Structure
  • Block — The s, p, d or f section of the periodic table, named after the type of sublevel the last electron goes into. → Atomic Structure, Periodic Trends

C

  • Calorimetry — Measuring the heat released or absorbed by a reaction from the temperature change of a known mass of water. → Energy Changes
  • Catalyst — A substance that speeds up a reaction without being used up, by providing an alternative pathway with a lower activation energy. Doesn't change \(\Delta H\). → Rates of Reactions, Energy Changes
  • Cation — A positive ion — an atom that has lost electrons, e.g. \(\ce{Na+}\). → Atomic Structure
  • Closed system — A sealed system: energy can get in or out, but matter can't. The measured mass stays constant. → Conservation of Mass
  • Collision theory — For particles to react they must collide, with at least the activation energy, and with the correct orientation. → Rates of Reactions
  • Combustion — A reaction with oxygen that releases energy. Complete combustion of a hydrocarbon gives \(\ce{CO2}\) and \(\ce{H2O}\). → Classifying Chemical Reactions
  • Compound — A pure substance made of two or more elements chemically bonded in a fixed ratio. → Separation of Mixtures
  • Concentration — How much solute is dissolved in a given amount of solution. Measured in mol L⁻¹, g L⁻¹, %, ppm and more. → Measures of Concentration
  • Core charge — Protons minus inner (non-valence) electrons — a measure of how strongly the nucleus pulls on the outer electrons. → Periodic Trends
  • Covalent bond — A shared pair of electrons between two non-metal atoms, attracted to both nuclei. → Chemical Bonding
  • Covalent network — A giant structure where every atom is covalently bonded to its neighbours (e.g. diamond, \(\ce{SiO2}\)). Very high melting point. → Chemical Bonding

D

  • Decomposition — One substance breaking down into two or more simpler ones, e.g. \(\ce{CaCO3 -> CaO + CO2}\). → Classifying Chemical Reactions
  • Delocalised electrons — Electrons not attached to any one atom, free to move through a structure — the \"sea\" in metallic bonding. → Chemical Bonding
  • Dilution — Adding solvent to lower a solution's concentration. The moles of solute stay the same: \(C_1V_1 = C_2V_2\). → Measures of Concentration
  • Dipole–dipole forces — Attractions between the \(\delta^+\) end of one polar molecule and the \(\delta^-\) end of another. → Chemical Bonding
  • Discrete molecule — A small, separate covalent molecule (e.g. \(\ce{H2O}\), \(\ce{CO2}\)). Strong bonds inside, weak forces between. Low melting point. → Chemical Bonding
  • Dispersion forces — Weak attractions between all molecules, caused by temporary uneven spreads of electrons. Stronger for bigger molecules. → Chemical Bonding
  • Distillation — Separating a liquid from a mixture by boiling it and condensing the vapour. Fractional distillation separates liquids with different boiling points. → Separation of Mixtures

E

  • Electron configuration — How an atom's electrons are arranged in sublevels, e.g. sodium is 1s² 2s² 2p⁶ 3s¹. → Atomic Structure
  • Electron dot diagram (Lewis diagram) — a diagram showing only the valence electrons, as shared pairs and lone pairs. → Chemical Bonding
  • Electronegativity — How strongly an atom attracts the electrons in a chemical bond. Increases across a period, decreases down a group. Fluorine is highest. → Periodic Trends, Chemical Bonding
  • Element — A pure substance made of only one kind of atom. → Separation of Mixtures
  • Emission spectrum — The set of coloured lines given off when an element's excited electrons fall back to lower energy levels. Unique to each element. → Atomic Structure
  • Empirical formula — The simplest whole-number ratio of atoms in a compound, e.g. \(\ce{CH2O}\) for glucose. → Mole Concept
  • Endothermic — A reaction that absorbs heat from the surroundings, so they get colder. \(\Delta H\) is positive. → Energy Changes
  • Energy level — The shells electrons occupy around the nucleus (\(n = 1, 2, 3 \dots\)). Further out = higher energy. → Atomic Structure
  • Energy profile diagram — A graph of chemical potential energy against reaction progress, showing reactants, products, activation energy and \(\Delta H\). → Energy Changes
  • Enthalpy change (\(\Delta H\)) — the heat energy released or absorbed by a reaction, usually in kJ mol⁻¹. \(\Delta H = -\dfrac{q}{n}\). → Energy Changes, Energy Changes
  • Excess reagent — The reactant that isn't fully used up — some is left over after the reaction. → Mole Concept
  • Exothermic — A reaction that releases heat to the surroundings, so they get hotter. \(\Delta H\) is negative. → Energy Changes

F

  • Filtrate / residue — In filtration, the filtrate is the liquid that passes through the filter paper; the residue is the solid left behind. → Separation of Mixtures
  • First ionisation energy — The energy needed to remove one electron from each atom in a mole of gaseous atoms. Increases across a period, decreases down a group. → Periodic Trends
  • Flame test — Heating a metal compound in a flame to identify the metal from the colour it gives off. → Atomic Structure

G

  • Gamma radiation (\(\gamma\)) — high-energy electromagnetic radiation from an unstable nucleus. No charge or mass; reduced by thick lead or concrete. → Atomic Structure
  • Group — A column of the periodic table. Elements in a group have the same number of valence electrons, so they react similarly. → Periodic Trends

H

  • Half-equation — An equation showing just the oxidation or just the reduction, with electrons included, e.g. \(\ce{Zn -> Zn^{2+} + 2e-}\). → Classifying Chemical Reactions
  • Half-life — The time for half of the radioactive atoms in a sample to decay. → Atomic Structure
  • Heterogeneous mixture — A mixture that isn't the same all the way through — you can often see separate parts (e.g. sand and water). → Separation of Mixtures
  • Homogeneous mixture — A mixture that's the same all the way through (e.g. salt water, air). → Separation of Mixtures
  • Hydrogen bonding — The strongest intermolecular force: between an H atom bonded to N, O or F and a lone pair on an N, O or F of another molecule. → Chemical Bonding

I

  • Ideal gas — A model gas whose particles take up no space, don't attract each other and collide perfectly elastically. Obeys \(PV = nRT\). → Properties of Gases
  • Intermolecular forces — Attractions between separate molecules (dispersion, dipole–dipole, hydrogen bonding). Much weaker than chemical bonds. → Chemical Bonding
  • Ionic bond — The electrostatic attraction between oppositely charged ions, formed when a metal transfers electrons to a non-metal. → Chemical Bonding
  • Ionic lattice — A 3D arrangement where every ion is surrounded by ions of the opposite charge (e.g. \(\ce{NaCl}\)). → Chemical Bonding
  • Isotopes — Atoms of the same element (same number of protons) with different numbers of neutrons. Same chemical properties, different masses and densities. → Atomic Structure

K

  • Kelvin (K) — the absolute temperature scale. \(T(\text{K}) = T(^\circ\text{C}) + 273.15\). Always used in gas calculations. → Properties of Gases

L

  • Limiting reagent — The reactant that runs out first. It decides how much product forms. → Mole Concept
  • Lone pair — A pair of valence electrons that isn't shared in a bond. Affects a molecule's shape. → Chemical Bonding

M

  • Mass number (\(A\)) — protons + neutrons in an atom's nucleus. → Atomic Structure
  • Metallic bonding — The electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons. → Chemical Bonding
  • Mixture — Two or more substances mixed but not chemically bonded. Each part keeps its own properties, so it can be separated physically. → Separation of Mixtures
  • Molar concentration (\(C\), mol L⁻¹) — moles of solute per litre of solution, \(C = \dfrac{n}{V}\). → Measures of Concentration
  • Molar mass (\(MM\), g mol⁻¹) — the mass of one mole of a substance. Add up the relative atomic masses in the formula. → Mole Concept
  • Molar volume — The volume of one mole of gas. At STP (0 °C, 100 kPa) it's 22.71 L. → Properties of Gases
  • Mole (mol) — the chemist's counting unit: \(6.022 \times 10^{23}\) particles. \(n = \dfrac{m}{MM}\). → Mole Concept

N

O

  • Open system — A system where matter (like a gas) and energy can get in or out, so the measured mass can change. → Conservation of Mass
  • Orbital — A region where electrons are found. Each orbital holds up to 2 electrons. → Atomic Structure
  • Oxidation — Loss of electrons; the oxidation state goes up. (OIL RIG: Oxidation Is Loss.) → Classifying Chemical Reactions
  • Oxidation state — A number that tracks electrons: 0 for elements, the charge for simple ions, usually −2 for O and +1 for H. Goes up in oxidation, down in reduction. → Classifying Chemical Reactions

P

  • Percentage composition — The percentage of a compound's (or mixture's) mass that comes from one element or component. → Mole Concept, Separation of Mixtures
  • Period — A row of the periodic table. Elements in a period have their outer electrons in the same energy level. → Periodic Trends
  • Polar bond — A covalent bond between atoms with different electronegativities; the electrons are pulled towards one atom (\(\delta^-\)). → Chemical Bonding
  • Polar molecule — A molecule with polar bonds arranged unevenly, so the charges don't cancel (e.g. \(\ce{H2O}\), \(\ce{NH3}\)). → Chemical Bonding
  • Polyatomic ion — An ion made of several atoms bonded together, e.g. sulfate \(\ce{SO4^{2-}}\), ammonium \(\ce{NH4+}\). → Chemical Bonding
  • ppm / ppb — Parts per million / parts per billion. For water solutions, 1 ppm = 1 mg L⁻¹ and 1 ppb = 1 µg L⁻¹. Used for very small concentrations. → Measures of Concentration
  • Precipitate — An insoluble solid that forms when two solutions are mixed. → Classifying Chemical Reactions
  • Primary standard — A substance pure and stable enough to weigh out an exact number of moles (e.g. anhydrous \(\ce{Na2CO3}\)). → Measures of Concentration
  • Pure substance — A substance with only one kind of particle and fixed properties — an element or a compound. → Separation of Mixtures

R

  • Radioisotope — An isotope with an unstable nucleus that decays and gives off radiation. → Atomic Structure
  • Rate of reaction — How fast reactants are used up or products are made — shown by how steep a product-vs-time graph is. → Rates of Reactions
  • Redox reaction — A reaction where electrons are transferred: one substance is oxidised while another is reduced. → Classifying Chemical Reactions
  • Reduction — Gain of electrons; the oxidation state goes down. (OIL RIG: Reduction Is Gain.) → Classifying Chemical Reactions
  • Relative atomic mass — The weighted average mass of an element's isotopes, on a scale where carbon-12 is exactly 12. → Atomic Structure, Mole Concept

S

  • Solubility rules — Rules for predicting which ionic compounds dissolve in water — and so which combinations form a precipitate. → Classifying Chemical Reactions
  • Specific heat capacity (\(c\)) — the energy needed to raise 1 g of a substance by 1 °C. Water: 4.18 J g⁻¹ K⁻¹. → Energy Changes
  • Spectator ions — Ions present in a solution that don't take part in the reaction — they appear unchanged on both sides. → Classifying Chemical Reactions
  • Standard temperature and pressure (STP) — 0 °C (273.15 K) and 100 kPa. → Properties of Gases
  • Stoichiometry — The number relationships between the amounts of reactants and products in a balanced equation. → Conservation of Mass, Mole Concept
  • Sublevel — A division of an energy level: s (holds 2), p (6), d (10) or f (14) electrons. → Atomic Structure
  • Successful collision — A collision with at least the activation energy and the correct orientation, so the particles react. → Rates of Reactions
  • Synthesis — Two or more substances joining to make one product, e.g. \(\ce{2Mg + O2 -> 2MgO}\). → Classifying Chemical Reactions

V

  • Valence electrons — The electrons in an atom's outer energy level — the ones involved in bonding. → Periodic Trends
  • VSEPR theory — Valence shell electron pair repulsion: electron pairs around a central atom spread as far apart as possible, which sets the molecule's shape. → Chemical Bonding