Hydration energy, also known as hydration enthalpy, can be defined as the amount of energy released when one mole of ions undergo hydration which is a special case of solvation. It is a special case of dissolution of energy with the solvent being water.
For example, when we dissolve salt in water, the outermost ions (those at the edge of the lattice) move away from the lattice and become covered with the neighbouring water molecules. If the hydration energy is equal to or if the hydration energy is greater than the lattice energy, the salt is known to be water-soluble. In salts for which the hydration energy is known to be higher than the lattice energy, solvation occurs with a release of energy in the form of heat. For example,CaCl2 (anhydrous calcium chloride) heats the water when dissolving. However, for instance, the hexahydrate, CaCl2·6H2O cools the water upon dissolution. The reason why the latter happens is that the hydration energy does not completely overcome the lattice energy and the remainder has to be taken from the water to compensate for the energy loss.
Enthalpy Change of Solution
The enthalpy change of solution can be defined as the enthalpy change when 1 mole of an ionic substance dissolves in water to give a solution of infinite dilution. Enthalpies of the solution can be either negative or positive. In other words, we can say that some ionic substances dissolve endothermically (for example, NaCl); others dissolve exothermically (for example NaOH).
An infinitely dilute solution can be defined as one where there is a sufficiently large excess of water that adding any more does not cause any further heat to be absorbed or evolved. So, when 1 mole of sodium chloride crystals is dissolved in an excess of water, the enthalpy change of solution is found to be +3.9 kJ mol-1. The change is slightly endothermic. So it can be said that the temperature of the solution will be slightly lower than that of the original water.
Factors Affecting the Size of Hydration Enthalpy
Hydration enthalpy: Hydration enthalpy can be defined as the measure of the energy released when attractions are set up between positive or negative ions and water molecules.
With positive ions, there may only lose ion-dipole attractions between the δ- oxygen atoms in the water molecules and the positive ions, or there may be formal dative covalent (coordinate covalent) bonds.
With negative ions, ion-dipole attractions are formed between the negative ions and the δ+ hydrogens in water molecules.
The size of the hydration enthalpy is usually governed by the amount of attraction between the ions and the water molecules.
The attractions are generally stronger for the smaller ions. For example, the hydration enthalpies fall as and when we go down a group in the Periodic Table. In the periodic table, the small lithium-ion has the highest hydration enthalpy in Group1 and the small fluoride ion has the highest hydration enthalpy in Group 7. In both groups, hydration enthalpy falls as the ions get bigger.
The more highly charged the ion is, the attraction will be stronger. For instance, the hydration enthalpies of Group 2 ions (like Mg2+) are much higher than those of Group 1 ions (like Na+)
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Which Element has the Highest Hydration Energy? Why does Lithium have High Hydration Enthalpy?
The lithium-ion has by far the highest hydration enthalpy in Group 1 and the small fluoride ion has by far the highest hydration enthalpy in Group 7.
Lithium exerts the greatest polarizing effect out of all the alkali metals on the negative ion. Being smallest in alkali metals, it's ion Li+ is smaller, this increases the charge density for this job appreciably. Consequently, its hydration energy becomes large.
Why does Hydration Enthalpy Decrease Down the Group?
Smaller the ion, the higher the hydration enthalpy will be because smaller atoms can accommodate a large number of water molecules around it and get hydrated. Hydration enthalpy decreases down the group the size of the atom increases due to the addition of extra valence shells.
Also, the hydration enthalpy decreases since the size of the cation increases. However, due to the square factor, lattice enthalpy decreases faster than the hydration enthalpy. That's why the solubility of Group 2 hydroxides increases while it progresses down the group.
Hydration Enthalpy of Elements
Hydration enthalpy values of various elements are tabulated in the table given below:
Application of Hydration Enthalpy
One application of enthalpy of hydration is the reaction of cement with water. The reaction being exothermic releases a large amount of heat. This heat released becomes significant in mass constructions like building dams and big structures. For the construction of massive concrete blocks, large quantities of cement are used.
During the process of setting, the heat is released. The outer surfaces of the block cool relatively faster than the interior, this creates a thermal gradient in the block and can initiate cracks that lead to failure of the structure. To avoid this, low heat types of cement are preferred for massive construction; cement with pozzolanic admixtures preferably fly ash or slag and also using ice instead of water to prepare concrete.