A'Level Chemistry 9701 Definitions

Comprehensive list of vocabulary and definitions for A'Level Chemistry 9701 covering stoichiometry, atomic structure, bonding, energetics, equilibria, kinetics, and organic chemistry.

Relative atomic mass ($A_r$)

The average mass of one atom of an element compared to $1/12$ the mass of a $^{12}C$ atom, or the mass of one mole of atoms compared to $1/12$ the mass of one mole of $^{12}C$ atoms.

Relative isotopic mass

The mass of one isotope (of a particular element) compared to $1/12$ the mass of a $^{12}C$ atom.

Relative molecular mass ($M_r$)

The average mass of one molecule compared to $1/12$ the mass of a $^{12}C$ atom.

Relative formula mass

The average mass of one formula unit of an ionic compound compared to $1/12$ the mass of a $^{12}C$ atom.

Mole

The amount of substance that contains as many entities as there are atoms in $12\,g$ of carbon-12. This value is $6.022 \times 10^{23}$, known as the Avogadro constant ($L$ or $N_A$).

Oxidation

A process where a chemical species loses electrons (Oxidation Is Loss).

Reduction

A process where a chemical species gains electrons (Reduction Is Gain).

Redox reaction

A reaction where oxidation and reduction occur simultaneously.

Oxidising agent

A species that accepts or gains electrons in a reaction and is itself reduced.

Reducing agent

A species that donates or loses electrons in a reaction and is itself oxidised.

Disproportionation reaction

A redox reaction in which one species is simultaneously oxidised and reduced.

Atomic number

The number of protons contained in the nucleus of an element's atom.

Mass number (nucleon number)

The sum of the protons and neutrons contained in an atom's nucleus.

Isotopes

Atoms of the same element with the same number of protons but a different number of neutrons.

Atomic orbital

A region of three-dimensional space around the nucleus where there is a $95\%$ chance of locating a particular electron.

VSEPR theory

Valence-shell electron pair repulsion theory; a model used to predict molecule shape based on electron-pair electrostatic repulsion, where lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

Metallic bond

The electrostatic attraction between positively charged cations and the 'sea' of delocalised electrons.

Electrovalent (Ionic) bond

The electrostatic attraction between oppositely charged ions formed by the transfer of electrons.

Coordination number (Ionic)

The number of neighboring oppositely charged ions surrounding an ion in an ionic compound.

Covalent bond

The electrostatic force of attraction of the nuclei of two atoms for the shared pair(s) of electrons between them.

Dative / Co-ordinate Covalent bond

A covalent bond in which both electrons in the shared pair are provided by only one of the two atoms.

Electronegativity

The ability or tendency of an atom to attract electrons in a bond towards itself.

Permanent dipole-permanent dipole interactions

Intermolecular forces between polar molecules that possess a net dipole moment.

Temporary dipole-induced dipole interactions

Intermolecular forces between non-polar molecules with no net dipole moment.

Hydrogen bond

A special case of permanent dipole-permanent dipole interaction between a hydrogen atom and an electronegative atom like nitrogen, oxygen, or fluorine.

Ideal Gas equation

$PV = nRT$, where Pressure is in $Pa$, volume in $m^3$, and temperature in $K$. Constants include $1\,atm = 1.01 \times 10^5\,Pa$ and $1\,dm^3 = 10^{-3}\,m^3$.

Hess’ law

States that the enthalpy change accompanying a reaction is independent of the path taken between the initial and final states.

Standard enthalpy change of reaction ($\Delta H^{\theta}_{rxn}$)

The enthalpy change when molar quantities of reactants (as specified by the equation) form products under standard conditions ($25^{\circ}C$ and $1\,atm$).

Standard enthalpy change of formation ($\Delta H^{\theta}_f$)

The enthalpy change when $1\,mole$ of a pure compound is formed from its constituent elements in their standard states under standard conditions.

Standard enthalpy change of combustion ($\Delta H^{\theta}_c$)

The enthalpy change when $1\,mole$ of a compound is completely burnt in oxygen under standard conditions.

Standard enthalpy change of neutralisation ($\Delta H^{\theta}_{neu}$)

The enthalpy change when an acid and base react to form $1\,mole$ of water under standard conditions.

Standard enthalpy change of atomisation ($\Delta H^{\theta}_{atom}$)

The enthalpy change when $1\,mole$ of gaseous atoms is formed from the element in its normal physical state under standard conditions.

Bond dissociation energy

The energy required to break $1\,mole$ of chemical bonds between two atoms in a molecule in the gaseous phase.

First ionisation energy ($\Delta H^{\theta}_{1st\,I.E.}$)

The energy required to remove $1\,mole$ of electrons from $1\,mole$ of gaseous atoms to form $1\,mole$ of gaseous singly charged cations ($M(g) \rightarrow M^+(g) + e^-$).

First electron affinity ($\Delta H^{\theta}_{1st\,E.A.}$)

The enthalpy change when $1\,mole$ of electrons are added to $1\,mole$ of gaseous atoms to form $1\,mole$ of gaseous singly charged anions.

Lattice energy

The energy evolved when $1\,mole$ of an ionic solid is formed from its constituent gaseous ions under standard conditions.

Standard enthalpy change of hydration ($\Delta H^{\theta}_{hyd}$)

The enthalpy change when $1\,mole$ of gaseous ions form hydrated aqueous ions under standard conditions.

Standard enthalpy change of solution ($\Delta H^{\theta}_{soln}$)

The enthalpy change when $1\,mole$ of an ionic compound is dissolved in a large excess of water ($\Delta H_{soln} = \sum \Delta H_{hyd} - \text{Lattice energy}$).

Entropy ($S$)

A measure of the degree of disorder or randomness in a system; it increases when matter or energy becomes more random.

Gibbs Free Energy change ($\Delta G$)

The limiting maximum useful work obtainable from a reaction at constant pressure ($\Delta G^{\theta} = \Delta H^{\theta} - T\Delta S^{\theta}$). Spontaneous reactions have $\Delta G < 0$.

Standard electrode potential

The potential difference between an element and its $1.00\,mol\,dm^{-3}$ aqueous ion relative to the standard hydrogen electrode at $1\,atm$ and $298\,K$.

Standard cell potential ($E^{\theta}_{cell}$)

The potential difference between two standard half cells measured under standard conditions ($E^{\theta}_{cell} = E^{\theta}_{cathode} - E^{\theta}_{anode}$).

Dynamic Equilibrium

A reversible reaction where the forward and backward reactions occur at the same rate, and concentrations of reactants and products remain constant.

Le Chatelier’s Principle

States that if a system in equilibrium is subjected to a change, the system will respond to reduce or counteract the effect of that change.

Bronsted acid

A proton ($H^+$) donor.

Bronsted base

A proton ($H^+$) acceptor.

Strong acid

An acid that dissociates completely in aqueous solution to give $H_3O^+$ ions.

pH

The negative logarithm to base 10 of the hydronium ion concentration: $pH = -\log_{10}[H_3O^+]$.

Ionic product of water ($K_w$)

$K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}\,mol^2\,dm^{-6}$ at $25^{\circ}C$; it increases with temperature as auto-ionisation is endothermic.

Buffer solution

An aqueous solution consisting of a weak acid and its conjugate base (or vice versa) that resists changes in pH when small amounts of acid or base are added.

Reaction rate

The increase in concentration of a product per unit time or the decrease in concentration of a reactant per unit time.

Activation energy ($E_a$)

The minimum energy which colliding molecules must possess for a successful collision or reaction to occur.

Order of reaction

The power to which the concentration of a particular reactant is raised in the experimentally determined rate law.

Half-life ($t_{1/2}$)

The time taken for the concentration of a reactant to fall to half its initial value; it is constant only for first-order reactions.

Catalyst

A substance that increases the reaction rate by providing an alternative pathway with a lower activation energy.

Transition element

A d-block element able to form one or more stable ions with a partially filled d-subshell.

Ligand

An anion or neutral molecule with at least one lone pair of electrons available for dative bonding to a central metal atom or ion.

Complex ion

An ion containing a central metal atom or ion bonded to ligands by coordinate (dative) bonds.

Empirical formula

The simplest formula showing the ratio of each kind of atom in a molecule.

Molecular formula

A formula showing the actual number of each kind of atom in a molecule.

Homologous series

A group of compounds with the same general formula and functional group, where each member differs by a fixed group (e.g., $-CH_2-$).

Structural isomerism

Compounds with the same molecular formula but different structural formulas.

Stereoisomerism

Compounds with the same molecular formula but different spatial arrangements (includes geometric and optical isomers).

Geometric isomers

Isomers with restricted rotation (usually $C=C$ bonds) where two different groups are bonded to each side of the bond.

Optical isomers (enantiomers)

Non-superimposable mirror images of each other that typically contain at least one chiral carbon atom.

Primary structure (Protein)

The exact order or unique sequence of $\alpha$-amino acids held by peptide/amide linkages along the polypeptide chain.

Secondary structure (Protein)

Detailed configurations (like $\alpha$-helix or $\beta$-pleated sheet) stabilized by hydrogen bonds between N-H and C=O groups along the main chain.

Tertiary structure (Protein)

The overall 3D shape of a protein held by R-group interactions (van der Waals', hydrogen bonds, ionic bonds, and disulfide linkages).

Quaternary structure (Protein)

The spatial arrangement of multiple individually folded protein subunits packed together.

Denaturation

The loss of biological activity of a native protein due to the disruption of secondary and tertiary structures; the primary structure remains unaffected.