1.3
Solubility and enthalpy change
Water molecules surrounding the cations and anions hydrogen-bond to further water molecules, releasing more energy. When the ions move through the solution, the attached water molecules move with them. There is not a fixed number of these molecules, and the outer water molecules are held less strongly than those close to the ion. However, average values are sometimes given for the extent of hydration. Some examples are listed in Table 5.
K+
Rb+
5
15
26
4
4
Now imagine that these gaseous ions dissolve in water to form hydrated ions (Figure 10). The enthalpy change for this is the enthalpy of hydration, which is defined as: The enthalpy change when one mole of gaseous ions dissolves in sufficient water to produce a solution of infinite dilution.
Table 5
Enthalpy of solution
The hydration of any gaseous ion is an exothermic process since bonds are being formed:
When some compounds are put in water, there is a measurable enthalpy change. This is the enthalpy of solution. In this section we will consider only ionic compounds.
Mx+(g) + aq → Mx+(aq) ΔhydH is negative Xy−(g) + aq → Xy−(aq)
The enthalpy change may be exothermic or endothermic. For example:
In these equations, ‘aq’ is taken to mean a large excess of water.
›› Sodium chloride dissolves in an excess of water, lowering the temperature of the water. This is an endothermic change. The enthalpy change is +3.9 kJ mol−1.
›› Sodium hydroxide dissolves in an excess of water, raising the temperature of the water. This is an exothermic change. The enthalpy change is −44.5 kJ mol−1.
The enthalpy of solution is defined as:
H
Enthalpy of hydration When an ionic compound dissolves in water a solution of hydrated cations and hydrated anions forms. For
O
O
H
Al
H
H
H
O
H
O
H O
H
3+
O
H
H
O
H O
H
H
H
H
H O
O
O
H
H H
H
H
H
O
O
H
the enthalpy change when one mole of an ionic compound dissolves in sufficient water to produce a solution of infinite dilution. Enthalpy changes can be determined in the laboratory using calorimetry (see Chapter 7 of Year 1 Student Book). However, strictly speaking, to determine an enthalpy of solution would require making ‘solution of infinite dilution’, as the definition says. This is impossible, but we can get close enough because what it really means is that there is sufficient water present that adding any more does not cause a further enthalpy change; in other words there is no observable change in temperature.
ΔhydH is negative
O
O H
H
H
H
H
H
O
H
Water molecules are attracted to the cation and energy is released. H The positive charge on the O H hydrogens attracts more water molecules. More energy is released.
H
H
H
H
H
O
H
O
O H
H
H
Water molecules are held more weakly the further they are from the metal ion. H
H H
H
H H H
H O
O
O O
H
O
H
H
O
H
Eventually, there is a limit to the number of water molecules that are bound around the cation. This is termed infinite dilution.
H
H O
H H
H
O
O
Al
O
H
H
O
O
H
H
H O
H
O
O
O
H
H
H
H O H
H
H
H
H
O 3+
H
O
H
O
O
H
H
H H
H
H H
H O
O
H
O
H
H
O
Al3+
H
Average hydration number
Mg2+
Imagine it was possible to break up an ionic compound into gaseous ions. The enthalpy change for this is the enthalpy of lattice dissociation (you learned earlier in this chapter how this can be determined using a Born–Haber cycle).
H
Metal ion
Na+
some compounds this is an exothermic process; for others it is endothermic. A thought experiment helps to explain why.
H H
H
O H
O
H H
Figure 10 Hydration of a cation
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