When solute-solvent interactions are strong enough to overcome the solute-solute interactions, solutes get dissolved in solvents and a solution is formed. The dissolving process in which water is used as a solvent is called hydration, and the energy therefore released is termed hydration energy.
Enthalpy change when one mole of gaseous ions is dissolved in water at standard conditions to form an infinitely dilute solution is called hydration energy/hydration enthalpy. It is represented by ΔH(hyd).
The general chemical equation for hydration enthalpy is;
Mn+(g) + mH2O (excess) → Mn+(aq)
It is generally denoted by;
Enthalpy change = ΔH(hyd)
There are some confusing terms like lattice energy and solution enthalpy which are similar yet different from the hydration energy.
Some detailed differences between lattice energy, heat/enthalpy of solution, and hydration energy are given (one by one) in the tables below:
Hydration Energy vs. Lattice Energy
ΔH(hyd) = ΔH(sol) – ΔH(lat)
| Hydration Energy | Lattice Energy |
| The energy released when one mole of gaseous ions are dissolved in water at S.T.P to form infinite dilute solution | The energy released when two oppositely charged ions in gaseous state combine to form an ionic compound, or the energy required to break an ionic compound into its gaseous ions |
| Enthalpies of hydration are always negative i.e. hydration is exothermic process and energy is released | Lattice energy of formation of ionic compounds is negative and lattice energy of breaking of ionic compound is positive |
|
For example: ΔH(hyd) for sodium (Na(g)) is -405 kJ/mol and for chlorine (Cl(g)) is -363 kJ/mol |
For example: ΔH(lat) for the formation of sodium chloride is -787 kJ/mol, and for dissociation is +787 kJ/mol |
Enthalpy of Hydration vs. Enthalpy of Solution
ΔH(sol) = ΔH(lat) + ΔH(hyd)
| Hydration Enthalpy | Solution Enthalpy |
| It is the energy released when one mole of separated gaseous ions form hydrated ions at S.T.P | It is the enthalpy change when one mole of solid ionic compound is dissolved in excess water to form hydrated ions at S.T.P |
| Hydration enthalpy is always negative (exothermic process) | Enthalpy of solution can either be negative (exothermic) or positive (endothermic) |
|
For example: ΔH(hyd) for lithium (Li(g)) is-520 kJ/mol, and for sodium (Na(g)) is -405 kJ/mol |
For example: ΔH(sol) for the dissolution of sodium hydroxide is -42.7kJ/mol |
The enthalpy of solution is also known as enthalpy of solvation or enthalpy of dissolution.
Hydration Energy and Solubility
- The solubility of salts is inversely proportional to the atomic size of atoms except that of fluorides and hydroxides. The solubility of salts depends on lattice energy and hydration energy.
Solubility ∝ hydration energy
Solubility ∝ 1/ lattice energy
- A compound has minimum solubility if the hydration enthalpies of constituent ions are equal. If these enthalpies are different, solubility increases.
- For ionic compounds to dissolve, hydration energy must exceed the lattice energy.
- A decrease in lattice energy increases the solubility, while a decrease in hydration energy decreases the solubility.
- The size of metal ions increases down the group so, both lattice energy and hydration energy decrease.
- For most compounds, hydration energy decreases more rapidly than lattice energy. The overall effect depends on which of the two has changed the most, therefore, solubility decreases down the group.
- For fluorides and hydroxides, lattice energy decreases more rapidly than hydration energy, down the group, therefore, their solubility increases down the group.
- There are strong forces of attraction between water and ions of ionic compounds, therefore, most ionic compounds are soluble in water. However, in some ionic compounds, lattice energy is so high that water molecules pull the constituent ions apart with much effort. Such compounds are termed ‘sparingly soluble’ in water.
Hydration Energy of Ions
Order of hydration energy of group IA [Alkali Metals] ions
Li+ > Na+ > K+ > Rb+ > Cs+
Order of hydration energy of group ΙIA [Alkaline Earth Metals] ions
Be+2 > Mg+2 > Ca+2 > Sr+2 > Ba+2
If the hydration energies of the first three groups are compared, the trend comes out as increasing as we go to the right. [ΙΙΙA>ΙΙA>ΙA]. For example,
Al+3 > Mg+2 > Na+
Examples of Hydration Energies
Factors Affecting Hydration Energy
Ionic size
H.E ∝ 1 / size of ions
The smaller the size of the ions, the higher the hydration energy.
For example, the ionic size of sodium is greater than lithium (Li+<Na+), therefore, the hydration energy of Li+ (-520 kJ/mol) is greater than Na+ (-405 kJ/mol).
Ionic charge
H.E ∝ charge on ions
Ions with a higher charge tend to have higher hydration energies.
For example, the ionic charge of magnesium is greater than sodium (Na+<Mg+2), therefore hydration energy of Mg+2 (-1922 kJ/mol) is greater than Na+ (-405 kJ/mol).
Applications of Hydration Enthalpy
- Hydration energy is used to calculate the solubility of substances in aqueous solutions.
- It is used in the study of solvation processes in chemical reactions.
- The hydration enthalpy of an electrolyte can be used to determine its strength, as stronger electrolytes have higher hydration enthalpies.
Concepts Berg
What is the enthalpy of hydration?
The enthalpy of hydration is the energy released when one mole of gaseous ions dissolves in water to form an infinitely dilute solution. It is usually expressed in kJ/mol.
What is the enthalpy of a solution?
The enthalpy change when one mole of a compound is dissolved in water to form hydrated ions at standard conditions is called the enthalpy of solution.
What is hydration in simple words?
Hydration is a dissolving process in which water is used as a solvent.
What factors affect hydration enthalpy?
Charge density (charge-to-size ratio) is directly related to hydration enthalpy. Greater the charge density, the higher the hydration enthalpy.
H.E ∝ charge density
Why is hydration enthalpy important for energy?
It is important for energy because it can affect the stability and reactivity of compounds. The hydration enthalpy of a compound can be used to predict its solubility.
Which has higher hydration energy: Cu+ or Cu2+?
H.E ∝ charge density
The Cu2+ ion has higher hydration energy than the Cu+ ion due to its high charge density.
How do hydration and lattice energies differ?
Hydration energy is the energy required to hydrate a solute. It is the energy released when gaseous ions dissolve in water to form an infinitely dilute solution.
Lattice energy is the energy required to separate the ions of a crystal lattice into gaseous ions. It is the energy released when gaseous ions are formed from a crystal lattice.
Why is the enthalpy of hydration exothermic?
The enthalpy of hydration is exothermic because energy is released when water molecules surround and bond with ions. This is due to the strong hydration enthalpy, which is the energy required to break the ion–water bonds. The energy released upon the formation of these hydrogen bonds is greater than the energy needed to break them apart, resulting in an overall decrease in energy.
References Books
- Oxford University Press, 2001 – A-level examinations – by Max Parsonage
- Competition Science Vision Magazine May 2009
- Single-ion Solvation: Experimental and Theoretical Approaches to Elusive …By Philippe Hünenberger, Maria Reif
- Chemistry: The Molecular Science By John W. Moore, Conrad L. Stanitski
- Chemistry for Degree Students B.Sc. First Year By R L Madan
Reference links
- An overview (ScienceDirect)
