A2-Level Enthalpy and Entropy
Solid ionic lattices are held together by strong attraction between oppositely charged ions.
The stronger the attraction between ions, the stronger the lattice (higher melting point)
Ion charge and size influence the strength of ionic attraction.
When two opposite charges come close together, their energies lower and energy is released.
Lattice enthalpy of formation is the enthalpy change when 1 mole’s worth of an ionic compound is formed from gaseous ions.
Lattice formation enthalpies are always negative (exothermic process).
Lattice enthalpy of dissociation is the enthalpy change to break apart 1 mole’s worth of an ionic compound into gaseous ions.
Lattice dissociation enthalpies are always positive (endothermic process).
When in solid state, ionic compounds form lattice structures. These lattice structures are held together by the strong forces of attraction between oppositely charged ions.
The stronger the attraction between ions, the stronger the ionic lattice will be and the harder it will be to break apart. The weaker the attractions between the ions, the easier I will be to break the lattice apart.
Ionic lattices are generally very stable, and when they are formed energy is released. Charged ions are high energy species. By coming together with oppositely charged ions their charges ‘cancel out’ and their energy decreases.
The closer that the oppositely charged ions can get to each other, the more stable they become and the more energy they release when they bond.
Lattice enthalpies show how much energy is released when a lattice is formed from gaseous ions. The lattice enthalpy, ∆H 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).
Lattice formation is always exothermic and lattice enthalpies should always be a negative value. This is because charged ions are higher in energy than the ionic lattice formed and this difference in energy is released when the lattice forms.
To break apart a lattice into gaseous ions energy is needed and the process is endothermic. This is called lattice dissociation and is, basically, the complete opposite of lattice enthalpy.
Size of Lattice Enthalpies
There are two factors that determine the strength of an ionic attraction. The charges of the ions involved and the sizes of the ions involved.
The greater the charge of the ions, the stronger the attraction between them. The larger the ions get, however, the weaker the attraction as their charge density gets smaller.
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 electron cloud around the negative ion. If the non-metal ion is highly electronegative (for example, fluorine) this does not matter as the fluorine will keep those electrons close and they will stay around the fluorine – not get moved to the sodium.
However, if you have a non-metal ion that is not very electronegative (for example, phosphorus) it cannot hold the cloud of electrons around it as well. If an aluminium (3+ charge) ion is next to this, it can pull the electron cloud ever so slightly towards itself.
The electrons pulled move a little bit to the aluminium ion. Each ion is no longer a ‘perfect sphere’, and the ionic structure is said to take on covalent character.