1.8 Thermodynamics - AQA Chemistry A-level

Hess’s Law

Enthalpy change is independent of the reaction route

Standard enthalpy of formation

Enthalpy change forming one mole of compound from elements in standard states

Standard enthalpy of an element

Zero enthalpy by definition

Standard enthalpy of combustion

Enthalpy change when one mole of substance is burnt in excess oxygen

Standard enthalpy of atomisation

Enthalpy change forming one mole of gaseous atoms from a compound in standard state

First ionisation energy

Enthalpy change removing electrons to form 1+ ions

Second ionisation energy

Enthalpy change removing electrons to form 2+ ions

First electron affinity

Enthalpy change when gaseous atoms gain electrons to form 1- ions

Second electron affinity

Enthalpy change when 1- ions gain electrons to form 2- ions

Lattice enthalpy of formation

Enthalpy change forming one mole of solid ionic lattice from gaseous ions

Lattice enthalpy of dissociation

Enthalpy change when solid ionic lattice breaks into gaseous ions

Enthalpy of hydration

Enthalpy change when gaseous ions dissolve in water to infinite dilution

Enthalpy of solution

Enthalpy change when solute dissolves completely in a solvent

Mean Bond Dissociation Enthalpy

Enthalpy change when breaking one mole of covalent bonds in gaseous state

Standard Enthalpy of Formation

Enthalpy change forming one mole of a compound from elements

Standard Enthalpy of Combustion

Enthalpy change when one mole of a substance combusts completely

First Ionisation Energy

Energy needed to remove one electron from a mole of gaseous atoms

Second Ionisation Energy

Energy needed to remove a second electron from a mole of gaseous ions

First Electron Affinity

Energy released when one mole of gaseous atoms gains an electron

Lattice Enthalpy of Formation

Enthalpy change forming one mole of an ionic compound from gaseous ions

Enthalpy of Hydration

Enthalpy change when one mole of gaseous ions dissolve in water

Born-Haber Cycle

Thermochemical cycle showing enthalpy changes in ionic compound formation

Factors Affecting Lattice Enthalpy

Size and charge of ions influence lattice enthalpy of ionic compounds

Perfect Ionic Model

Assumption of perfectly spherical ions with even charge distribution

Interactions of Ions with Water Molecules

Solid breaks into gaseous ions, dissolves in water surrounded by dipolar water molecules

Polarisation

Occurs with small positive or large negative ions, adding covalent character to ionic bonds.

Crystal Structure

May differ in irregular lattices.

Most Ionic Bonds

Between large positive and small negative ions, e.g., CsF.

Spontaneous

Reaction occurs on its own; disregards reaction rate.

Feasible

Reaction that can occur.

Enthalpy Change

Negative enthalpy (exothermic) favors spontaneity.

Entropy

Measure of system randomness/disorder.

Units for Entropy

JK-1mol-1.

Second Law of Thermodynamics

Entropy in an isolated system always increases.

General Entropy Values

Solids < liquids < gases.

Entropy Change Calculation

ΔS = ΣS(products) - ΣS(reactants).

Gibbs Free Energy Equation

ΔG = ΔH - TΔS.

Gibbs Free Energy Significance

G < 0: feasible, G = 0: just feasible, G > 0: not feasible.

Temperature at G=0

Temperature at which reaction becomes feasible.

Temperature for Feasibility

T = (ΔH)/(ΔS) when G = 0.

G Limitations

Doesn't consider reaction rate; only indicates feasibility.

Exothermic with Entropy Increase

G is negative, reaction is feasible and product-favored.

Endothermic with Entropy Decrease

G is positive, reaction is not feasible and reactant-favored.

Exothermic with Entropy Decrease

Temperature-dependent.

Endothermic with Entropy Increase

Temperature-dependent.

Entropy at 0K

Zero due to maximum order.

Entropy Changes

Increase with more moles; solid to liquid/gas increases entropy.

Constant Temperature in Endothermic Reaction

Heat released escapes to surroundings.