3.6 ENTHALPY CHANGES FOR SOLIDS AND SOLUTIONS

Enthalpy of atomisation ΔHₐₜ

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

• Na(s)→Na(g)

Standard conditions

298K

1 moldm^-3 solutions

Pressure 1atm

Standard state

Physical state of a substance under standard conditions t

Enthalpy change

Enthalpy change = Enthalpy of formation (products) - Enthalpy change of formation(reactants)

Hess’s law

The energy change of any chemical reaction is the same regardless route taken

Lattice formation enthalpy ΔHₗₐₜₜ (formation)

The enthalpy change when one mole of an ionic solid is formed from its gaseous ions

• Na⁺ (g) + Cl^- (g) → NaCl(s)

ALWAYS EXOTHERMIC

Lattice breaking enthalpy ΔHₗₐₜₜ (dissociation)

The enthalpy change required to separate one mole of an ionic solid into gaseous ions

• NaCL(s) → Na⁺ (g) + Cl^- (g)

ALWAYS ENDOTHERMIC

Hydration enthalpy ΔHₕyd

Enthalpy change when one mole of gaseous ions dissolved in water to form aqueous ions

• Na⁺ (g) + → Na⁺ (aq)

EXOTHERMIC

Enthalpy of solution ΔHₛₒₗ

The enthalpy change when one mole of an ionic solid dissolves in water to form aqueous ions

• NaCl (s) → Na⁺(aq) + Cl^- (aq)

Enthalpy change of solution equation

ΔH_solution​=ΔH_{lattice breaking​}+ΔH_hydration​

Electron affinity

Enthalpy change when one mole of gaseous negative ions formed from gaseous atoms of a substance by gaining an electron

• Cl (g) + e^- → Cl^- (g)

Ionisation energy

Enthalpy change when one mole of gaseous positive formed from gaseous atoms a substance by losing an electron

• Na (g) → Na⁺ (g) + e^-

Ionic substances consist of

Positive and negative ions held by electrostatic forces

Ions must be broke to dissolve

What determines solubility?

Balance between

• Lattice breaking enthalpy (endothermic)

• Hydration enthalpy (exothermic)

Soluble salt from

Hydration > lattice breaking

Exothermic > Endothermic

Insoluble salt from

Hydration < Lattice breaking

Exothermic < Endothermic

• Hydration can’t compensate

Lattice enthalpy and hydration increase when

• Charges are higher

• Small ion radius

Factors affecting enthalpy of lattice breaking and of solution

• Increased charge on ions → increase

• Decreasing ion size → increase

Born haber cycle shows

All steps to form ionic compound from its elements

Haber cycle of NaCl step by step

1. Atomisation Na (s) → Na (g) form gaseous

2. Ionisation energy Na(g) → Na+ (g) → Na⁺ (g) + e^-

3. Atomisation of non metal ½ Cl₂ (g) → Cl(g) break bonds

4. Electron affinity Cl (g) + e- → Cl^- (g) gain electron

5. Lattice formation Na⁺ (g) + Cl^- (g) → Nacl(s)
   

ΔHf​= atomisation+IE+EA+lattice

Constructing born-haber

1. Elements in standard states

2. To gaseous atoms

3. Gaseous ions

4. Drop down → Ionic solid

Enthalpy of formation ΔH°f

Enthalpy change when 1 mol of a compound forms from elements in standard states under standard conditions

• more negative → more stable

Why is a more negative enthalpy of formation more stable

More energy is released → stronger bonding → more stable

Why is MgO more stable than NaCl

Mg²⁺ and O^2- have higher than Na⁺ and Cl^-

Stronger electrostatic attraction

Larger lattice enthalpy

More negative ΔH°f

More stable

Why does increasing ionic charge increase lattice enthalpy

Higher charges

Stronger electrostatic attraction between ions

Why do smaller ions increase lattice enthalpy

Have higher charge density

Increase attraction between ions

Why is the second electron affinity positive

The negative electron is repelled from negative species