H1 Chemistry Definitions
H1 Chemistry 8873 definitions by topic, including physical, inorganic and organic distinctions.
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A topic-organised glossary for H1 Chemistry (8873). Recall the meaning, then explain the condition or distinction in the last column. These are revision definitions, not a promise of marks for fixed wording.
Atoms and bonding
| Term | Meaning | Remember |
|---|---|---|
| Isotope | An atom of the same element as another, with a different neutron number. | Proton number fixes the element; nucleon number can differ. |
| Atomic orbital | A region described by a one-electron wavefunction, associated with a probability distribution for an electron in an atom. | An orbital is not a circular path; it holds at most two electrons with opposite spins. |
| First ionisation energy | Energy required to remove one mole of electrons from one mole of isolated gaseous atoms to form one mole of gaseous singly positive ions. | X(g) → X + (g) + e⁻; kJ/mol; the gaseous condition matters. |
| Electronegativity | Ability of an atom in a bond to attract the bonding electrons towards itself. | A comparison within bonding; different from ionisation energy. |
| Ionic bond | Electrostatic attraction between oppositely charged ions. | Acts throughout an extended lattice. |
| Covalent bond | Electrostatic attraction between a shared electron pair and the nuclei of bonded atoms. | A bond may be polar even when the whole molecule is non-polar. |
| Dative covalent bond | Covalent bond whose shared pair is supplied by one atom. | Once formed, it is not a different kind of attraction. |
| Metallic bond | Attraction between positive metal ions and delocalised electrons. | Delocalisation helps explain conduction and malleability. |
| Sigma bond | Covalent bond formed by head-on orbital overlap along the internuclear axis. | A carbon–carbon single bond is a sigma bond. |
| Pi bond | Covalent bond formed by sideways overlap, with electron density on opposite sides of the internuclear axis. | A double bond contains one sigma and one pi bond. |
| Hydrogen bond | Attraction between hydrogen bonded to N, O or F and a lone pair on N, O or F in the school model. | Stronger than many other intermolecular attractions, but not an O–H covalent bond. |
| Permanent dipole | Separation of partial positive and negative charge in a molecule. | Molecular shape determines whether bond dipoles cancel. |
| Instantaneous dipole–induced dipole attraction | Attraction arising when a fluctuating electron distribution induces a dipole in a neighbouring particle. | Present in all atoms and molecules; strength depends on polarisation and contact. |
| Bond energy | Energy required to break one mole of a specified covalent bond in gaseous species. | Bond breaking is endothermic; average bond energies give estimates. |
| Bond length | Equilibrium distance between the nuclei of two bonded atoms. | Compare the specified bonds rather than atomic radii alone. |
Moles and reacting quantities
| Term | Meaning | Remember |
|---|---|---|
| Mole | Amount containing exactly 6.02214076 × 10²³ specified entities. | State whether the entities are atoms, molecules, ions or electrons. |
| Relative isotopic mass | Mass of an atom of a particular isotope compared with one twelfth of the mass of a carbon-12 atom. | A ratio, with no unit. |
| Relative atomic mass | Weighted mean mass of atoms of an element compared with one twelfth of the mass of a carbon-12 atom. | Includes isotope abundances. |
| Relative molecular / formula mass | Sum of relative atomic masses in a molecule / formula unit. | No unit; distinguish from molar mass in g/mol. |
| Empirical formula | Simplest whole-number ratio of atoms of each element in a compound. | Derived from mole ratios, not mass ratios directly. |
| Molecular formula | Actual number of atoms of each element in a molecule. | A whole-number multiple of the empirical formula. |
| Molar concentration | Amount of solute per unit volume of solution. | Common unit mol/dm³; use solution volume, not solvent volume. |
| Limiting reactant | Reactant exhausted first according to the balanced reacting ratio. | Compare moles divided by their coefficients, not masses directly. |
| Oxidation | Loss of electrons or increase in oxidation number. | Oxidising agents are themselves reduced. |
| Reduction | Gain of electrons or decrease in oxidation number. | Reducing agents are themselves oxidised. |
| Oxidation number | Formal charge assigned to an atom using electron-allocation rules. | Not necessarily the actual charge on that atom. |
| Endpoint | Experimentally observed signal used to stop a titration. | Should lie close to, but is conceptually different from, the equivalence point. |
| Equivalence point | Point at which titrant and analyte have reacted in their stoichiometric ratio. | Does not necessarily have pH 7. |
Enthalpy and energy
| Term | Meaning | Remember |
|---|---|---|
| Enthalpy change of reaction | Heat change for a reaction at constant pressure, for the quantities specified by its equation. | State the reaction equation and molar basis. |
| Standard conditions | Specified reference conditions: standard pressure 100 kPa, stated temperature and species in their standard states. | Standard does not automatically fix the temperature at 298 K; it must be stated. |
| Standard enthalpy change of formation | Enthalpy change when one mole of a substance forms from its elements in their standard states under standard conditions. | Fractional coefficients are allowed to form exactly one mole. |
| Standard enthalpy change of combustion | Enthalpy change when one mole of a substance burns completely in oxygen under standard conditions, with all substances in their standard states. | Specify the product states, especially water. |
| Standard enthalpy change of neutralisation | Enthalpy change when acid and base react to form one mole of water under standard conditions. | The one-mole basis is water, not always acid. |
| Lattice energy | Enthalpy change when one mole of an ionic solid forms from its separated gaseous ions. | Formation convention: negative; a lattice dissociation value uses the opposite sign. |
| Hess’s law | Reaction enthalpy depends only on initial and final states, not the route between them. | Reversing an equation reverses its enthalpy sign. |
| Exothermic / endothermic | A process releasing heat to / absorbing heat from its surroundings. | Negative / positive Δ H; temperature change depends on which part is measured. |
| Activation energy | Minimum energy needed by reacting particles for a successful reaction along a specified pathway. | A catalyst changes the pathway; it does not change reaction enthalpy. |
Rates and equilibrium
| Term | Meaning | Remember |
|---|---|---|
| Rate of reaction | Change in concentration of a specified reactant or product per unit time, with the sign or stoichiometric convention stated. | Product formation and reactant disappearance can have different numerical rates. |
| Rate equation | Experimentally determined relationship between rate and reactant concentrations at fixed temperature. | r = k[A]^m[B]ⁿ; powers need not equal overall equation coefficients. |
| Order with respect to a reactant | Power of that reactant’s concentration in the rate equation. | Overall order is the sum of the powers. |
| Rate constant | Proportionality constant in a rate equation at a specified temperature. | Units depend on overall order; temperature and catalyst affect its value. |
| Half-life | Time for the concentration of a specified reactant to fall to half its initial value. | Successive half-lives are constant for a first-order reaction. |
| Catalyst | Substance that increases rate through an alternative pathway and is regenerated overall. | It changes neither the equilibrium constant nor the equilibrium composition at fixed temperature. |
| Homogeneous / heterogeneous catalyst | Catalyst in the same phase as / a different phase from the reactants. | Surface adsorption matters in heterogeneous catalysis. |
| Enzyme | Biological catalyst, usually a protein, with a specific active site. | Temperature and pH can change its activity. |
| Dynamic equilibrium | State in a closed system where forward and reverse rates are equal and macroscopic composition stays constant. | Reactant and product concentrations need not be equal. |
| Le Chatelier’s principle | When an equilibrium system is disturbed, its response tends to oppose the imposed change. | A catalyst speeds approach to equilibrium without shifting it. |
| Equilibrium constant | Value of the equilibrium composition expression for a specified reaction at a specified temperature. | Use equilibrium concentrations or partial pressures; changing temperature can change the constant. |
Acids, bases and buffers
| Term | Meaning | Remember |
|---|---|---|
| Arrhenius acid / base | Substance increasing hydrogen-ion / hydroxide-ion concentration in aqueous solution. | This model is limited to aqueous systems. |
| Brønsted–Lowry acid / base | Proton donor / proton acceptor. | Identify the species that actually transfers or accepts the proton. |
| Conjugate acid–base pair | Two species differing by one proton. | Their charges differ by one unit. |
| Strong / weak acid | Acid that ionises essentially completely / partially in water under the stated conditions. | Strength and concentration describe different things. |
| Strong / weak base | Base reacting essentially completely / partially with water to generate hydroxide ions, or a hydroxide that dissociates essentially completely / partially. | Compare like species and the appropriate equilibrium. |
| pH | Negative base-10 logarithm of the numerical hydrogen-ion concentration in mol/dm³ in the school concentration model. | pH measures acidity; it is not a direct measure of total acid concentration. |
| Acid dissociation constant, Kₐ | Equilibrium constant for an acid transferring a proton to water. | For HA, Kₐ = [H⁺][A⁻]/[HA]; depends on temperature. |
| Base dissociation constant, K_b | Equilibrium constant for a base reacting with water to produce its conjugate acid and hydroxide ions. | For B, K_b = [BH⁺][OH⁻]/[B]. |
| Ionic product of water, K_w | Product of hydrogen-ion and hydroxide-ion concentrations in the school model at a specified temperature. | Neutrality means equal ion concentrations, not pH 7 at every temperature. |
| Buffer | Solution that resists changes in pH when small amounts of acid or alkali are added. | Contains appreciable amounts of a weak acid/base and its conjugate partner; capacity is finite. |
| Acid–base indicator | Weak acid/base whose conjugate forms have different colours. | Choose its transition range within the steep part of the titration curve. |
| Amphoteric substance | Substance that can react as both an acid and a base. | Identify its role in each particular reaction. |
Organic chemistry and polymers
| Term | Meaning | Remember |
|---|---|---|
| Functional group | Atom or group responsible for characteristic reactions of an organic compound. | The same molecule can have more than one functional group. |
| Homologous series | Family with the same functional group and general formula, similar chemical reactions and a gradation of physical properties. | Successive members differ by CH₂. |
| Structural isomers | Compounds with the same molecular formula but different connectivity. | Atoms are joined in different orders. |
| Stereoisomers | Compounds with the same connectivity but different spatial arrangements. | Compare with structural isomers. |
| Cis–trans isomerism | Stereoisomerism arising from restricted rotation, with groups on the same or opposite sides of a reference plane. | Each alkene carbon needs two different substituents for alkene cis–trans isomerism. |
| Addition / substitution | Reaction adding groups across a multiple bond / replacing an atom or group. | Distinguish product connectivity, not just reactant names. |
| Elimination | Reaction removing atoms or groups to create a multiple bond. | Often competes with substitution for a suitable substrate. |
| Condensation | Reaction joining molecules with elimination of a small molecule. | Ester and amide formation are examples. |
| Hydrolysis | Bond cleavage involving reaction with water. | Acidic or alkaline conditions can give different final species. |
| Addition polymerisation | Joining unsaturated monomers without eliminating small molecules. | The repeat unit is not the unchanged alkene molecule. |
| Condensation polymerisation | Joining suitable monomers with loss of small molecules. | A difunctional monomer can form a chain. |
| Denaturation | Loss of a protein’s functional three-dimensional structure. | Does not normally mean hydrolysis of its primary peptide sequence. |
Test the distinction
A weak acid can have a higher concentration than a strong acid. Strength concerns the extent of ionisation; concentration concerns the amount per solution volume. To explain a pH difference, compare the resulting hydrogen-ion concentrations rather than repeating “stronger” or “more concentrated”.
H1 boundary
H1 calculates pH for strong acids and strong bases, treats buffer action qualitatively and uses concentration-based equilibrium constants. The conjugate-pair relationship K_w = KₐK_b, weak-acid and buffer pH calculations, Kₚ, entropy/Gibbs calculations and electrochemical equations are not required. Integrated rate equations and H2 organic mechanisms are also outside this course.
Return to the course hub for the lesson behind a term or relationship, then practise without this reference.