Hydration Enthalpy

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.

What is the hydration energy?

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 called the enthalpy of hydration, H_hyd. Here are the approximate hydration energies of some common ions. The table shows that as the atomic number grows, the ionic size decreases, resulting in a fall in absolute enthalpy of hydration values.

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⁺)

Solution enthalpy change

The enthalpy change of solution is the enthalpy change that occurs when one mole of an ionic compound dissolves in water to produce an infinitely diluted solution. Enthalpies in solution can be positive or negative, meaning that some ionic substances dissolve endothermically (such as NaCl) while others dissolve exothermically (for example NaOH)

An infinitely dilute solution has a large enough excess of water that adding more does not cause any further heat to be absorbed or released. The enthalpy change of solution is found to be +3.9 kJ mol^-1 when 1 mole of sodium chloride crystals are dissolved in an excess of water. Because the reaction is slightly endothermic, the temperature of the solution will be lower than that of the original water.

Which Element has the Highest Hydration Energy? Why does Lithium have High Hydration Enthalpy?

• Lithium-ion

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:

Hydration power

Enthalpy of hydration, H_hyd, ion is the amount of energy released when ion molecules dissolve in large amounts of water to form a soluble solution that does not end up in this process.

Mz + (g) + mH₂O Mz + (aq)

Mz + (aq) ions are covered by H₂O molecules and dispersed in solution. Hydration power is limited by some of the common ions listed here. The table shows the fact that as the number of atoms increases, so does the ionic size, which leads to a decrease in the total enthalpy values ​​of hydration.

Features affecting Hydration Enthalpy

Enthalpy of Hydration: Ionic Charge & Radius

• The typical enthalpy fluctuations of hydration (ΔH_hydθ) are affected by the number of ions attracted to water molecules.

• Factors affecting this are the attraction of ionic charging and radiation

Ionic Radius

ΔH_hydθ becomes extremely exothermic by reducing ionic radii

• Smaller ions have a higher charge density which leads to a stronger attraction of ion-dipole interactions between water molecules and ions in solution.

• Therefore, more energy is released when it becomes hydrated and ΔH_hydθ

• becomes more exothermic

For example, ΔH_hydθ of magnesium sulfate (MgSO₄) is more harmful than ΔH_hydθ of barium sulfate (BaSO₄)

• Since both compounds contain sulfate (SO₄^2-) ions, the difference in ΔH_hydθ must be due to magnesium ion (Mg²⁺) in MgSO₄ and barium (Ba²⁺) ions in BaSO₄.

• Magnesium is of Group 2 

• Barium is of Group 2 

• This means that Mg²⁺ ions are smaller than Ba²⁺ ions

• The attraction is therefore very strong for Mg²⁺ ions

• As a result, standard enthalpy for hydration of MgSO₄ is more dangerous than BaSO₄

Ionic Charging

ΔH_hyd θ is very harmful to ions with high ionic ions

• Ions with high ionic ions have a high charge capacity leading to a strong ion-dipole attraction between water molecules and ions in solution.

• Therefore, more energy is released when it becomes hydrated and ΔHhydꝋ becomes more exothermic

For example, ΔH_hyd θ of calcium oxide (CaO) is more harmful than ΔHhydθ of potassium chloride (KCl)

• Calcium oxide is an ionic compound that contains calcium (Ca²⁺) and oxide (O₂^-) ions.

• Potassium chloride is made up of potassium (K⁺) and chloride (Cl^-) ions

• Both ions in calcium oxide have a higher ionic charge than ions in potassium chloride

• This means that the attraction is strong between water molecules and Ca²⁺ and O₂^- ions over CaO hydration.

• Attractive properties are weak between water molecules and K⁺ and Cl^- ions over KCl water infiltration

Therefore, ΔH_hyd θ of calcium oxide is very dangerous as more energy is released from its absorption.

Hydration Enthalpy and Solubility

The ions in the solute are bound together by the coulombic force of attraction, in order to dissolve this soluble substance (here water) the water molecule must overcome these gravitational forces. The force required to cross this line of attraction is called lattice enthalpy.

Most ionic compounds do not dissolve in waterless solutions but show high solubility in water. A factor determining the dissolution of salts is the interaction of ions with the solvent. As previously described, water is a polar molecule with a positive charge, part hydrogen and part negative charge in oxygen, interacts with ions and forms a strong bond that releases energy.

The termination process can be considered as a combination of two processes.

The first is,

M + (S) → M + (g) △ = △H_hyd θ Lattice enthalpy

The second process is hydration,

M + (g) + aq → M + (aq) △ = △H_hyd θ Hydration enthalpy

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.