Physical I Chemistry Definitions: AQA A-level

Acceleration

Second stage of TOF spectrometry where ions are accelerated by an electric field for equal kinetic energy.

Atom

Smallest part of an element; all substances are composed of atoms.

Atomic nucleus

Positively charged centre of an atom with protons and neutrons, surrounded by electrons.

Atomic number

Number of protons in an atom's nucleus.

Electron

Negatively charged subatomic particle orbiting the nucleus with a relative mass of 1/2000.

Electron configuration

Arrangement of electrons in atom orbitals, e.g., Ca: 1s²2s²2p⁶3s²3p⁶4s².

Electron impact ionisation

TOF spectrometry ionisation method using high-energy electrons to create 1+ ions.

Electrospray ionisation

TOF spectrometry ionisation method where a sample is dissolved, ionised, and turned into a gas.

First ionisation energy

Energy needed to remove electrons from gaseous atoms to form 1+ ions, e.g., O(g) → O⁺(g) + e⁻.

Ion detection

Fourth stage of TOF spectrometry where charged particles are detected for mass spectrum.

Ion drift

Third stage of TOF spectrometry where ions drift through a region with no electric field.

Ionisation

First stage of TOF spectrometry where the sample is ionised by different methods.

Isotope

Atoms of the same element with different neutron numbers, e.g., 35Cl and 37Cl.

Mass number

Total protons and neutrons in an atom's nucleus.

Mass spectrometer

Instrument providing isotopic mass and abundance information.

Mass spectrometry

Technique for identifying elements and determining relative molecular mass.

Neutron

Neutral subatomic particle in the atom's nucleus with a relative mass of 1.

Nuclear charge

Total charge of protons in the nucleus, equal to the atomic number.

Proton

Positively charged subatomic particle in the atom's nucleus with a relative mass of 1.

Second ionisation energy

Energy to remove electrons from 1+ ions to form 2+ ions, e.g., O⁻(g) → O²⁺(g) + e⁻.

Sub-shells (orbitals)

Divisions of electron shells with different energy levels, holding up to two electrons each.

Time of Flight (TOF) spectrometer

Mass spectrometry method measuring ions' mass-to-charge ratio based on time of flight; consists of four stages.

Atom economy

Measure of starting materials becoming useful products.

Avogadro’s constant

Number of entities in one mole of a substance.

Concentration

Amount of substance per unit volume

Empirical formula

Simplest ratio of elements in a compound

Limiting reactant

Reactant completely consumed, limiting product formation

Mole

Mass with the same atoms as 12g of carbon-12

Molecular formula

Actual ratio of elements in a compound

Percentage by mass

Element concentration in a compound or mixture

Percentage yield

Actual yield percentage compared to theoretical yield

Relative atomic mass

Average mass of an element's atom compared to carbon-12

Relative molecular mass

Average mass of one molecule compared to carbon-12

Co-ordinate bond

Bond with shared electrons from one atom

Covalent bond

Shared electron pair between non-metals

Dipole

Charge difference in a covalent bond

Electron pair repulsion

Electron pairs repel, positioning apart in a molecule

Electronegativity

Atom's ability to attract electrons in a bond

Electrostatic forces

Attraction between oppositely charged ions

Hydrogen bonding

Strong bond between hydrogen and electronegative atom

Intermolecular forces

Forces between molecules affecting physical properties

Ion

Atom/molecule with electric charge from electron loss/gain

Ionic bond

Bond between oppositely charged ions

Ionic compound

Compound of ions held by strong electrostatic forces

Lattice

Regular arrangement of atoms/ions in crystal structures

Macromolecular crystal structure

Giant covalent structures with high melting points

Metallic bond

Bond between metal ions and electrons

Permanent dipole-dipole forces

Stronger intermolecular forces than van der Waals

Polar bond

Covalent bond with uneven electron distribution, inducing a dipole

Simple molecular crystal structure

Structure with strong covalent bonds and low melting points

Van der Waals forces

Weak intermolecular forces due to temporary electron density fluctuations

Calorimetry

Measuring energy changes in a chemical reaction

Endothermic reaction

Reaction absorbing energy, reducing surroundings' temperature

Enthalpy change (∆H)

Heat energy change at constant pressure

Exothermic reaction

Reaction releasing energy, increasing surroundings' temperature

Hess’s law

Enthalpy change independence from reaction route

Mean bond enthalpy

Average energy to break one mole of a covalent bond

Molar enthalpy change

Enthalpy change per mole of substance

Standard conditions

100 kPa and 298K temperature

Activation energy

Minimum energy for successful reaction initiation

Catalyst

Substance increasing reaction rate without being consumed

Collision theory

Reactions require collisions with sufficient energy

Effect of concentration on reaction rate

Increased reactant concentration leads to more collisions and faster reactions

Effect of pressure on reaction rate

Higher pressure increases collisions, accelerating reaction rate

Effect of temperature on reaction rate

Increasing temperature increases kinetic energy, leading to more collisions and higher reaction rates.

Maxwell-Boltzmann distribution

Graph showing molecular energy distribution in a gas at constant temperature.

Rate of reaction

Measure of product formation or reactant use over time, expressed in g/s, cm3/s, or mol/s.

Closed system

System with only heat exchange, no matter entering or exiting.

Dynamic equilibrium

State where forward and backward reaction rates are equal, concentrations remain constant.

Equilibrium constant (KC)

Value expressing product/reactant concentrations at equilibrium in a reversible reaction.

Heterogeneous system

System with chemicals in different phases.

Homogeneous system

System with chemicals in the same phase.

Le Chatelier’s principle

States equilibrium shifts to counteract changes in concentration, temperature, or pressure.

Reversible reaction

Reaction where products can reform reactants, direction altered by conditions.

Half equation

Split of a redox equation into oxidation and reduction half-equations for balancing.

Oxidation state

Charge of an ion or theoretical charge of an atom in a covalent compound.

Oxidising agent

Elements/compounds accepting electrons, causing their reduction.

Redox reaction

Reaction involving simultaneous reduction and oxidation.

Reducing agent

Elements/compounds donating electrons, causing their oxidation.