Chapter 17 Thermodynamics

Standard enthalpy of atomisation, ΔatHᶱ

enthalpy change which accompanies the formation of one mole of gaseous atoms from the element in its standard state under standard conditions

First electron affinity, ΔeaHᶱ

standard enthalpy change when a mole of gaseous atoms is converted to a mole of gaseous 1- ions

Lattice enthalpy of formation, ΔLHᶱ

standard enthalpy change when one mole of solid ionic compound is formed from its gasious ions

Lattice enthalpy of dissociation, ΔdissHᶱ

standard enthalpy change when one mole of solid ionic compound dissociates into its gaseous ions

Enthalpy of hydration, ΔhydHᶱ

standard enthalpy change when water molecules surround one mole of gaseous ions

Enthalpy of solution, ΔsolHᶱ

standard enthalpy change when one mole of solute dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far enough apart not to interact with each other

what is the trend in lattice enthalpies for ionic compounds

as the ions get smaller the lattice enthalpies increase

because the ionic radius is smaller so the oppositely charged ions are closer together

resulting in stronger ionic bonding

how to calculate enthalpy of solution

enthalpy of solution = enthalpy of lattice dissociation + enthalpy of hydration

reason for discrepancy between experimental and theoretical values of lattice enthalpy

the ionic compound has covalent character

factors which increase polarisation for anions

large size

high negative charge

factors which increase polarisation for cations

small size

high positive charge

define entropy

the measure of disorder in a chemical system

how to work out change in entropy for a reaction

ΔS = ΣS_products - ΣS_reactants

units for entropy

JK^-1mol^-1

what 2 factors determine the feasibility of a chemical reaction

enthalpy change and entropy change

equation for Gibb’s free energy

ΔG = ΔH - TΔS

ΔG = Gibb’s free energy change

ΔH = enthalpy change

T = temperature in kelvin

ΔS = entropy change

what value of ΔG is required for a reaction to be feasible

ΔG negative or equal to 0

why does water have a higher entropy of vapourisation compared to other liquids

water has hydrogen bonds

so in the liquid state it is more ordered than other liquids without hydrogen bonds

how to work out temperature at which a reaction is feasible

when feasible ΔG = 0

T = ΔH/ΔS

explain in terms of molecules why entropy is 0 at 0K (2)

molecules have 0 kinetic energy so they don’t move

there is no disorder

why does entropy increase with temperature (2)

particles gain kinetic energy so start to move more

disorder increases

why is L₂ longer than L₁ (2)

L₂ corresponds to boiling, L₁ corresponds to melting

there is a greater change in disorder in boiling compared to melting

in terms of forces acting on particles why is the first electron affinity of oxygen exothermic (1)

there is a force of attraction between the nucleus of the oxygen atom and an external electron

standard free energy change for formation of MgO, ΔG = -570kJ mol^-1. Suggest one reason why a sample of MgO appears to be stable in air at room temperature despite negative value for ΔG (1)

a protective layer of MgO forms