A Level Chemistry Important Definitions
Advanced Subsidiary
Terms/ Concept Definition
Basic ideas about atoms
Atomic number The number of protons in the nucleus of an atom
Mass number The number of protons and neutrons in the nucleus of an atom
Isotope Atoms having the same number of protons but different number of
neutrons
Alpha particles Helium nuclei
Beta Particles Fast moving electron (not ONLY electron)
Gamma ray Electromagnetic waves of high energy
Half Life Time needed to half the atoms in a radioisotope to decay
Shielding effect Repulsion between electrons in different shell
First ionisation energy Energy needed to remove one mole of electrons from one mole of
atoms in GASEOUS state to form one mole of +1 ions, under standard
conditions
Successive ionisation Energy needed to remove each electron in turn until all electrons are
energy removed from an atom
Absorption Spectrum Electrons absorbs specific frequencies of light and excite to higher
energy levels, darks lines formed corresponds to light absorbed
Emission Spectrum Then they fall back to lower energy levels and emits energy as a
photon
Convergence limit Individual spectrum lines are so close to each other that cannot be
distinguished from one another
Chemical Equations
Relative atomic mass Average mass of an atom of an element relative to 1/12 of the mass of
an atom Carbon-12
Relative isotopic mass Mass of an isotope relative to 1/12 of the mass of an atom carbon-12
Relative formula mass The sum of relative atomic masses of all atoms present in its formula
Mole The amount of substance that has the same number of particles as
there are atoms in exactly 12g of carbon-12
Empirical formula Shows the simplest ratio of the atoms present in a molecule
Molar mass The mass of 1 mole of substance
Molecular formula Shows the actual number of each type of atom present in a molecule
Atom economy The percentage by mass of all the reactants that ends up in the desired
product
Avogadro’s constant The number of particles in one mole
Fragmentation Splitting molecules in a mass spectrometer into smaller parts
Bonding
Ionic bond Electrostatic attraction between the positive and negative ions/ cation
and anion (Strong, non directional)
Covalent bond Has a pair of electrons with opposed spins shared between two atoms
with each atom giving one electron (Strong, directional)
Metallic Bond Electrostatic attraction between the positive metal ion and the
negative ’sea’ of delocalised electrons
Coordinate bond / Covalent bond which both electrons come from one atom
dative covalent bond
Polar bond One end of the bond with a slightly positive charge and the other end
with a slightly negative charge
, Intermolecular force The weak bonding holding the molecules together
Electronegativity The power of an atom to attract electrons in a covalent bond
Intramolecular force The strong bonding between the atoms in the molecule and governs its
chemistry
Van der Waals’ forces Includes all types of intermolecular forces
Hydrogen Bonding Bonding occurs between oxygen, nitrogen or fluorine of 1 molecule
and hydrogen
Solid structures
Coordinate number The coordinate number of an ion gives the number of its nearest
neighbours
Volatility Describes how readily a substance vaporises
Simple equilibria
Reversible reaction One that can go in either direction depending on the conditions
Dynamic equilibrium Forward and reverse reactions occur at the same rate
La Chatter’s principle States that if a system at equilibrium is subjected to a change, the
equilibrium tends to shift to minimise the effect of the change
Acid base reaction
Acid Proton donor
Base Proton acceptor
Strong acid Fully dissociates in aqueous solution
Weak acid Partially dissociates in aqueous solution
Concentrated acid Large quantity of acid, small quantity of water
Dilute Large quantity of water
Strong base a substance that accepts H+ in a non-reversible reaction
Weak base a species that accepts H+ in a reversible reaction
Standard solution A solution with a known concentration
Thermochemistry
Exothermic reaction One that releases energy to the surroundings
(Enthalpy Change is negative)
Endothermic Reaction One that takes in energy from the surroundings
(Enthalpy Change is positive)
Enthalpy The heat content of a system in constant pressure
Enthalpy change The heat added to a system in constant pressure
Standard Enthalpy The Enthalpy Change when 1 mole of substance is formed from its
Formation element in their standard condition
Standard Enthalpy The Enthalpy Change when 1 mole of substance is completely
Combustion combusted in oxygen under standard conditions
Hess’s Law The total enthalpy change for a reaction is independent of the route
taken from the reactants to the products
Bond Enthalpies Enthalpy required to break a covalent bond X-Y bond into X atoms
and Y atoms, all in Gas Phase
Average Bond Enthalpy The average value of the energy required to break a given type of
covalent bond in the molecules of a gaseous species.
Enthalpy Change of Reaction between the number of moles of reactant shown in the
reaction equation for the reaction
Rates of reaction
The rate of reaction The change in concentration of a reactant per unit time
Activation energy The minimum energy required to start a reaction by breaking of bonds
Catalyst Substance that increases the rate of reaction by lowing the activation
energy
Homogeneous Catalyst In the same physical state as the reactants
Heterogeneous Catalyst In a different physical state from the reactants
Advanced Subsidiary
Terms/ Concept Definition
Basic ideas about atoms
Atomic number The number of protons in the nucleus of an atom
Mass number The number of protons and neutrons in the nucleus of an atom
Isotope Atoms having the same number of protons but different number of
neutrons
Alpha particles Helium nuclei
Beta Particles Fast moving electron (not ONLY electron)
Gamma ray Electromagnetic waves of high energy
Half Life Time needed to half the atoms in a radioisotope to decay
Shielding effect Repulsion between electrons in different shell
First ionisation energy Energy needed to remove one mole of electrons from one mole of
atoms in GASEOUS state to form one mole of +1 ions, under standard
conditions
Successive ionisation Energy needed to remove each electron in turn until all electrons are
energy removed from an atom
Absorption Spectrum Electrons absorbs specific frequencies of light and excite to higher
energy levels, darks lines formed corresponds to light absorbed
Emission Spectrum Then they fall back to lower energy levels and emits energy as a
photon
Convergence limit Individual spectrum lines are so close to each other that cannot be
distinguished from one another
Chemical Equations
Relative atomic mass Average mass of an atom of an element relative to 1/12 of the mass of
an atom Carbon-12
Relative isotopic mass Mass of an isotope relative to 1/12 of the mass of an atom carbon-12
Relative formula mass The sum of relative atomic masses of all atoms present in its formula
Mole The amount of substance that has the same number of particles as
there are atoms in exactly 12g of carbon-12
Empirical formula Shows the simplest ratio of the atoms present in a molecule
Molar mass The mass of 1 mole of substance
Molecular formula Shows the actual number of each type of atom present in a molecule
Atom economy The percentage by mass of all the reactants that ends up in the desired
product
Avogadro’s constant The number of particles in one mole
Fragmentation Splitting molecules in a mass spectrometer into smaller parts
Bonding
Ionic bond Electrostatic attraction between the positive and negative ions/ cation
and anion (Strong, non directional)
Covalent bond Has a pair of electrons with opposed spins shared between two atoms
with each atom giving one electron (Strong, directional)
Metallic Bond Electrostatic attraction between the positive metal ion and the
negative ’sea’ of delocalised electrons
Coordinate bond / Covalent bond which both electrons come from one atom
dative covalent bond
Polar bond One end of the bond with a slightly positive charge and the other end
with a slightly negative charge
, Intermolecular force The weak bonding holding the molecules together
Electronegativity The power of an atom to attract electrons in a covalent bond
Intramolecular force The strong bonding between the atoms in the molecule and governs its
chemistry
Van der Waals’ forces Includes all types of intermolecular forces
Hydrogen Bonding Bonding occurs between oxygen, nitrogen or fluorine of 1 molecule
and hydrogen
Solid structures
Coordinate number The coordinate number of an ion gives the number of its nearest
neighbours
Volatility Describes how readily a substance vaporises
Simple equilibria
Reversible reaction One that can go in either direction depending on the conditions
Dynamic equilibrium Forward and reverse reactions occur at the same rate
La Chatter’s principle States that if a system at equilibrium is subjected to a change, the
equilibrium tends to shift to minimise the effect of the change
Acid base reaction
Acid Proton donor
Base Proton acceptor
Strong acid Fully dissociates in aqueous solution
Weak acid Partially dissociates in aqueous solution
Concentrated acid Large quantity of acid, small quantity of water
Dilute Large quantity of water
Strong base a substance that accepts H+ in a non-reversible reaction
Weak base a species that accepts H+ in a reversible reaction
Standard solution A solution with a known concentration
Thermochemistry
Exothermic reaction One that releases energy to the surroundings
(Enthalpy Change is negative)
Endothermic Reaction One that takes in energy from the surroundings
(Enthalpy Change is positive)
Enthalpy The heat content of a system in constant pressure
Enthalpy change The heat added to a system in constant pressure
Standard Enthalpy The Enthalpy Change when 1 mole of substance is formed from its
Formation element in their standard condition
Standard Enthalpy The Enthalpy Change when 1 mole of substance is completely
Combustion combusted in oxygen under standard conditions
Hess’s Law The total enthalpy change for a reaction is independent of the route
taken from the reactants to the products
Bond Enthalpies Enthalpy required to break a covalent bond X-Y bond into X atoms
and Y atoms, all in Gas Phase
Average Bond Enthalpy The average value of the energy required to break a given type of
covalent bond in the molecules of a gaseous species.
Enthalpy Change of Reaction between the number of moles of reactant shown in the
reaction equation for the reaction
Rates of reaction
The rate of reaction The change in concentration of a reactant per unit time
Activation energy The minimum energy required to start a reaction by breaking of bonds
Catalyst Substance that increases the rate of reaction by lowing the activation
energy
Homogeneous Catalyst In the same physical state as the reactants
Heterogeneous Catalyst In a different physical state from the reactants
