Hydration Enthalpy

The amount of energy released when a mole of an ion dissolves in a large amount of water, forming an infinitely dilute solution in the process, Mz+(g) + mH2O↔ Mz+(aq) where Mz+(aq) represents ions surrounded by water molecules and dispersed in the solution, is referred to as the enthalpy of hydration, abbreviated Hhyd.  It is evident from the table that as the atomic number increases, the ionic size reduces, which results in a decrease in the absolute enthalpy of hydration values.

In thermodynamics, the hydration enthalpy (Hhyd) is the change in enthalpy that occurs when one mole of gaseous ion dissolves in a sufficient amount of water to make an infinitely dilute solution under standard conditions of one bar pressure (infinite dilution means a further addition of solute will not cause any heat change).   

For example, when we dissolve salt in water, the outermost ions (those on the edge of the lattice) move away from the lattice and become covered by the water molecules in the surrounding area. Water-soluble salts are those in which the hydration energy is equal to or greater than the lattice energy; this is referred to as being water-soluble. The process of solvation happens in salts where the hydration energy is known to be higher than the lattice energy, and the release of energy in the form of heat is seen. When CaCl2 (anhydrous calcium chloride) is dissolved in water, it warms the water to a certain degree. For example, the hexahydrate of calcium chloride, CaCl2 6.H2O, cools the water when it is dissolved. As a result of this inability to entirely overcome the lattice energy, water must be extracted from its surrounding environment in order to compensate for the energy lost through the hydration process.

Enthalpy Change of Solution

The enthalpy change of solution is the change in enthalpy that occurs when one mole of an ionic material dissolves in water to form a solution of infinite dilution. The enthalpies of a solution might be either negative or positive depending on the situation. To put it another way, some ionic substances dissolve endothermically (for example, NaCl), whereas others dissolve exothermically (for example, KCl, NaCl) (for example NaOH).

When there is a sufficient enough excess of water, an infinitely dilute solution can be described as one in which adding additional water does not result in any further heat being absorbed or evolved. When one mole of sodium chloride crystals is dissolved in an excess of water, the enthalpy change of the solution is found to be +3.9 kJ mol-1, which is a positive value. The shift is mildly endothermic in nature. As a result, it is reasonable to predict that the temperature of the solution will be slightly lower than the temperature of the original water.

Features affecting Hydration Enthalpy

The heat of Hydration: Ionic Charge & Radius of Hydration

It is the quantity of ions that are attracted to water molecules that has an effect on the normal enthalpy fluctuations of hydration (Hhyd). The attraction of ionic charge and radiation are two factors that have an impact on this.

Ionic Radius

Hydride becomes particularly exothermic as a result of the reduction of ionic radii.

In solution, the charge density of smaller ions is higher, which results in a larger attraction of ion-dipole interactions between water molecules and the ions in solution. As a result, more energy is released as it becomes hydrated, and the exothermic property of Hydride increases.

For example, magnesium sulphate (MgSO4) is more toxic than barium sulphate (BaSO4) when it is dissolved in water (BaSO4)

Based on the fact that both compounds contain sulphate (SO42-) ions, the variation in ion concentration must be attributed to the presence of magnesium ions (Mg2+)in MgSO4 and barium ions  (Ba2+) in BaSO4.

Magnesium is classified as a Group 2 element.

Barium is classified as a Group 2 element.

In other words, Mg2+ ions are significantly smaller in size than Ba2+ ions.

As a result, the attraction for the ion Mg2+is quite high.

Therefore, the standard enthalpy for hydration of MgSO4 is more harmful than the enthalpy for hydration of BaSO4.

CONCLUSION

 The enthalpy of hydration is defined as the amount of energy released when one mole of gaseous ions is diluted by one mole of water. It can be thought of as the enthalpy of solvation, with the solvent being water as the solvent. Hydration enthalpy, also known as hydration energy, is always negative, and its values are determined by the amount of water absorbed.

 Water is classified as a polar solvent due to the presence of both positive (H atom) and negative (O atom) poles in its structure. As soon as an ionic compound (such as NaCl) is dissolved in water, the solid-state structure of the complex is disrupted, and the ions Na+ and Cl– are separated from one another.

When this energy is stated as an enthalpy of hydration, the difference between M+(g) and M+(aq) is the fact that in M+(aq),, the ion is enclosed by water molecules, creating a weak bond with the water molecules.