M - FROM MINERALS TO ELEMENTS

Group 7 Elements: The Halogens

Theory Of Solutions

Extracting Copper from its Ore

Ions in Solids and Solutions

In ionic solids, e.g. sodium chloride, the ions are held together by their opposite electrical charges. Each cation (+) attracts several anions (-) and vice versa. The ions build up into a giant ionic lattice structure. In sodium chloride each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions which forms a cubic shaped lattice. This simple cubic shape is found in other Group 1 halides, e.g. KF and LiCl.

It is possible to fit two different sorts of cations into an anion lattice provided the + and - charges balance. An example is potassium magnesium chloride KCl.MgCl₂.H₂O which exists as a double salt. It is a single substance and not a mixture of two salts. The Cations are arranged in a regular pattern throughout the lattice. The water molecule in this salt is called water of crystallisation. The crystals are called hydrated salts.

Many ionic substances dissolve in water. When they do, the ions become surrounded by water molecules. They are said to be hydrated. The hydrated ions are spread throughout the water at random and they behave independently of each other.

Energy Changes in Solution

Some ionic substances dissolve in water, but others do not, so what decides whether an ionic substance will dissolve?

Lattice Enthalpy

Before an ionic solid can dissolve, the ions need to be separated from the lattice. This requires energy, the lattice enthalpy, ΔH_LE. This is the enthalpy change when 1 mole of a solid is formed from its gaseous ions.

E.g. Na⁺(g) + Cl^-(g) NaCl(s) ΔH_LE = -789 kJ mol^-1

All lattice enthalpies are large negative values. If we want to break down a lattice, we have to put in energy equal to -ΔH_LE. This tends to stop substances dissolving. Lattice enthalpies depend on size and charge of the ions. Lattice enthalpies become more negative when:
- the ionic charge increases
- the ionic radius decreases
Ions attract more strongly when they are small and highly charged.
Enthalpy of Hydration

Water molecules possess a small dipole. The tiny charges on the water molecules are attracted to the charges on the ions. The strength of attractions between ions and water molecules is measured by the enthalpy of hydration.

This is the enthalpy change for the production of a solution of ions from 1 mole of gaseous ions.

Na⁺(g) + aq Na⁺(aq) ΔH_hyd = -390 kJ mol^-1
Br^-(g) + aq Br^-(aq) ΔH_hyd = -337 kJ mol^-1

Enthalpies of hydration are always negative. They depend on the charge and size of an ion. They become more negative when:
- ionic charge increases
- ionic radius decreases
When we deal with solvents other than water we use the term enthalpy of solvation, ΔH_solv.
Enthalpy of Solution

The difference between the enthalpies of hydration of the ions and the lattice enthalpy gives the enthalpy of solution, ΔH_solution. It is the enthalpy change when one mole of a solute dissolves to form a solution of concentration 1 mol dm^-3.

ΔH_solution = ΔH_{hyd (cation)} + ΔH_{hyd (anion)} - ΔH_LE

A solute will dissolve if ΔH_solution is slightly negative and there will be a small temperature rise.

A solute will not dissolve if ΔH_solution is large and positive, but it will dissolve if it is slightly positive.

An enthalpy level diagram for a solute with ΔH_solution slightly negative is shewn below.