Energetics definitions

Enthalpy change, standard conditions and the definitions

Enthalpy change, ΔH

The heat energy change measured under conditions of constant pressure.

Exothermic

Describes a reaction in which heat energy is transferred to the surroundings, so ΔH is negative.

Endothermic

Describes a reaction in which heat energy is taken in from the surroundings, so ΔH is positive.

Activation energy

The minimum energy that particles must have for a reaction to take place when they collide.

Kinetically stable

Describes a mixture that does not react at a measurable rate because its activation energy is high, even though ΔH for its reaction is negative.

Standard conditions

A pressure of 100 kPa and a stated temperature, usually 298 K, with any solutions at 1.00 mol dm⁻³.

Standard state

The physical state of a substance at 100 kPa and 298 K, such as H₂O(l), O₂(g) and C(graphite).

Standard enthalpy of formation, ΔfH°

The enthalpy change when one mole of a compound is formed from its constituent elements, with all reactants and products in their standard states under standard conditions.

Standard enthalpy of combustion, ΔcH°

The enthalpy change when one mole of a substance is burned completely in oxygen, with all reactants and products in their standard states under standard conditions.

Standard enthalpy of neutralisation, ΔneutH°

The enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce one mole of water.

Standard enthalpy of reaction, ΔrH°

The enthalpy change when the amounts of reactants shown in the equation react under standard conditions, with all substances in their standard states.

Measuring an enthalpy change: calorimetry

Calorimetry

The measurement of the heat energy transferred in a reaction from the temperature change of a known mass of water or solution.

Specific heat capacity, c

The energy needed to raise the temperature of 1 g of a substance by 1 K.

Limiting reagent

The reactant that is completely used up, which fixes the amount of product formed and the heat released.

Extrapolated temperature rise

The temperature change read at the time of mixing from a cooling line extended backwards, which corrects for heat lost to the surroundings during the reaction.

Strong acid

An acid that is completely dissociated into ions in aqueous solution.

Weak acid

An acid that is only partially dissociated into ions in aqueous solution.

Incomplete combustion

Combustion in a limited supply of oxygen, giving carbon monoxide or carbon as well as or instead of carbon dioxide, and releasing less energy.

Bomb calorimeter

A sealed steel vessel in which a sample burns completely in pure oxygen, surrounded by a known mass of water, so that almost no heat is lost.

Hess's law and enthalpy cycles

Hess's law

The enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same.

Enthalpy cycle

A diagram showing two or more routes between the same reactants and products, with the enthalpy change for each step, used to calculate an unknown enthalpy change.

Enthalpy of formation of an element

Zero for an element in its standard state, because forming the element from itself involves no change.

Water of crystallisation

Water molecules held in fixed proportion within the crystal lattice of a hydrated salt.

Anhydrous

Describes a salt with no water of crystallisation, such as white CuSO₄(s).

Enthalpy of solution

The enthalpy change when one mole of a solute dissolves in enough solvent that further dilution causes no further heat change.

Thermal decomposition

The breaking down of a compound into simpler substances by heating it.

Indirect determination

Finding an enthalpy change that cannot be measured directly by combining, in a Hess cycle, enthalpy changes that can be measured.

Mean bond enthalpies

Bond dissociation enthalpy

The enthalpy change when one mole of a specified covalent bond is broken, with reactants and products in the gaseous state.

Covalent bond

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the two bonded atoms.

Homolytic fission

The breaking of a covalent bond so that each atom takes one electron from the shared pair, forming two radicals.

Mean bond enthalpy

The enthalpy needed to break one mole of a given type of covalent bond in the gaseous state, averaged over a range of different compounds.

Enthalpy of atomisation

The enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.

Bond length

The distance between the nuclei of two atoms joined by a covalent bond.

Radical

A species with an unpaired electron, such as a hydrogen or chlorine atom formed by homolytic fission.

Enthalpy of vaporisation

The enthalpy change when one mole of a liquid is converted into a gas at the same temperature.