Quick Notes Lattice Enthalpies

  • Lattice enthalpy of formation is the enthalpy change that occurs when 1 mole’s worth of an ionic compound is formed from its constituent, gaseous ions.
    • Lattice formation enthalpies are always negative (exothermic process).
  • Lattice enthalpy of dissociation is the enthalpy change that occurs when 1 mole’s worth of an ionic lattice is broken apart into its constituent, gaseous ions.
    • Lattice dissociation enthalpies are always positive (endothermic process).
  • Solid ionic lattices are held together by strong electrostatic attraction between oppositely charged ions.
    • Ion charge and size influence the strength of ionic attraction.
    • The stronger the attraction between ions, the stronger the lattice (higher melting point).
    • When two opposite charges come close together their energies lower, releasing energy to the surroundings.

Full Notes Lattice Enthalpies

When in solid state, ionic compounds have a lattice structure. These lattice structures are held together by strong electrostatic forces of attraction between oppositely charged ions.

The stronger the attraction between ions, the stronger the ionic lattice, making it harder to break apart. The weaker the attractions between the ions, the easier it is to break the lattice apart. See Ionic Bonding.

attraction between oppositely charged ions lattice

Ionic lattices are generally very stable, and when they are formed energy is released. Charged ions are high energy species. By moving two oppositely charged ions closer, their energy decreases. This means energy is released by the charged ions when they bond to each other.

how distance between particles changes their energy

The closer the oppositely charged ions can get to each other, the more stable they become and the more energy they release.

Lattice enthalpies show how much energy is released when a lattice is formed from gaseous ions.

The lattice enthalpy of formation, ΔHf of an ionic compound is the enthalpy change that occurs when 1 mole of an ionic compound is formed from its constituent, gaseous ions (the ions that are in the compound). Standard enthalpies mean the enthalpy change is measured under standard conditions.

Lattice formation is always exothermic, meaning lattice formation enthalpies should always have a negative value. This is because the charged ions are higher in energy than the ionic lattice they form. The difference in energy between the ions and the lattice is released when the lattice forms.

lattice enthalpy of sodium chloride endothermic

To break apart a lattice into gaseous ions, energy is needed, so the process is endothermic. This is called lattice dissociation and is, basically, the complete opposite of lattice formation.

Size of Lattice Enthalpies

There are two factors that determine the strength of an ionic attraction - the charges of the ions involved and their sizes.

charge of ion and strength of ionic attraction lattice enthalpy sodium chloride magnesium chloride

The greater the difference in charge between the ions, the stronger the attraction between them. The larger the ions get, however, the weaker the attraction as their charge is now spread over a larger area, meaning their ‘charge density’ gets smaller.

size of ion and strength of ionic attraction lattice enthalpy sodium chloride potassium bromide

Charge density refers to how the charge of the ion compares to the size (or mass) of the ion. High charge densities of ions give strong ionic attraction.

Covalent Character

This is a brief outline of covalent character, for more detail see Ionic-Covalent Character.

In a ‘perfect’ ionic bond, the metal ion is a positively charged species that is a perfect sphere, and the non-metal ion is a negatively charged species that is also a perfect sphere.

The reality is a little messier than this. If you have a positive metal ion (for example, sodium) it will attract the electrons from around a negative ion. If the non-metal ion is highly electronegative (for example, fluorine) this doesn’t matter as the fluorine will keep those electrons close around itself, they won’t get moved to the sodium.

ionic and covalent character polarisation of covalent bond to form ionic character strength energy

However, if you have a non-metal ion that is not very electronegative it cannot keep hold of its electrons as well. An aluminium (3+ charge) ion, for example, would easily pull the electrons of a phosphorus ion ever so slightly towards itself.

The electron density around the negative ion is now ‘distorted’ and the negative ion is no-longer a perfect sphere. A small amount of electron density ends up being shared between the ions and the ionic structure is said to take on covalent character.