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Enthalpy of Hydration - Definition, Standard Enthalpy of Hydration, Enthalpy of Hydration and Solubility, Factors Affecting Hydration, Applications, Practice Problems, FAQs

Enthalpy of Hydration - Definition, Standard Enthalpy of Hydration, Enthalpy of Hydration and Solubility, Factors Affecting Hydration, Applications, Practice Problems, FAQs

Do you know what the colour of CuSO4 is?

You might be confused if it is white or blue. CuSO4 is white in colour, whereas CuSO4.5H2O is blue in colour. The difference between the two aforementioned formulae is the five water molecules which is called the water of hydration. When CuSO4 is dissolved in water, the Cu2+ ions and SO42- ions are surrounded by water molecules. This process is called hydration and the energy released during this process is called the hydration enthalpy or the enthalpy of hydration.

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In this article, we will learn about the enthalpy of hydration and its related terms in detail.

Table of Contents

  • Hydration
  • Enthalpy of Hydration and Standard Enthalpy of Hydration
  • Enthalpy of Hydration and Solubility
  • Factors Affecting Hydration
  • Enthalpy of Hydration of Metal Ions
  • Application of Enthalpy of Hydration
  • Practice Problems
  • Frequently Asked Questions

Hydration

When substances dissolve in a solvent, it can do so on two counts. The substance- called the solute, shall be so small compared to the molecular structure in the solvent, that it can get lodged inside it without any interaction between them.

A better and higher solubility will be there when there is a favourable interaction between the solute and the solvent. For example, the water-a polar solvent is able to dissolve most of the ionic compounds as per the dictum- like dissolves like.

But what type of interactions are there?

The simplest of the positive interactions is between the cations of the solute and the partially negatively charged (oxide) ions of water. This makes the cations being surrounded by water molecules where the oxides ions oriented towards the cations.

Similarly, the anions of the solute also will be surrounded by the water, But the positively charged protons of water will be oriented towards the anions. The dipole-dipole interactions of the oppositely charged solute and water molecules stabilise the cations and anions individually. This process of stabilization of cations and anions of electrolytes by water solvent by encircling them is called hydration.

Enthalpy of Hydration and Standard Enthalpy of Hydration

The surrounding of the cations and anions by the water molecules stabilizes both the ions by releasing heat energy. Hydration is always an exothermic process and hence hydration enthalpy also called hydration energy is always negative.

Mn+(g) + aq → M++(aq) and xn-(g) → xn-(aq)

Enthalpy change = hydrationH

The amount of energy released when one mole of gaseous cations and anions are stabilised by the water molecules is called enthalpy of hydration.

The heat energy released when one mole of electrolytes in the gaseous phase are surrounded and stabilized by the water molecules at standard conditions of infinite dilution, one bar pressure and 273 K is referred to as standard enthalpy of hydration.

Enthalpy of Hydration and Solubility

When an ionic crystal is placed in water, two processes can occur.

  1. Breaking of the crystal lattice freeing the cations and anions into the solvent and
  2. Surrounding of cations and anions by the solvent molecules.

Breaking the lattice needs energy equivalent to Lattice energy, which is the energy released when gaseous cations and gaseous anions form one mole of the crystal lattice. This will be always an endothermic process and hence will not be a spontaneous process.

Hydration on the opposite is an exothermic process and spontaneous nature.

The solubility of the electrolytes will depend on the overall thermodynamic stabilization of both processes. So, the lesser the lattice energy and higher the hydration energy, the larger will be the solubility of the electrolyte in water.

Example 1.

Anhydrous calcium chloride is very deliquescent and absorbs water even from the atmosphere to get hydrated releasing a lot of energy. The hydration energy of anhydrous calcium chloride is very much higher than its lattice energy making the hydration highly spontaneous.

Example 2.

The lattice energy of NaCl is defined as the energy released when Na+ and Cl ions come close to each other to form a lattice.

Na+(g) +Cl-(g) NaCl(s)+ΔH

When an ionic compound (any salt, say NaCl) is dissolved in water the solid-state structure of the compound is destroyed and the Na+ and Cl are separated by absorbing equally to lattice energy.

Water is a polar solvent because it has positive hydrogen and negative oxygen dipoles.

In the solution, they exist as Na+ atoms surrounded by the negative ends of water molecules and similarly the Clis surrounded by the positive end of water molecules. Since additional bonds are formed between the atoms the process releases some energy equal to the hydration energy.

Na+(g) +Cl-(g)  Na+(aq.) +Cl-(aq.)  +ΔH

Therefore, the enthalpy of solution is calculated as:

solH = Enthalpy of hydration + Lattice energy­­

Where solH is the enthalpy of the solution.

Substituting the enthalpy values in the above expression.

solH = Enthalpy of hydration + Lattice energy = −783kJ mol-1 + 786kJ mol-1 = 3kJmol-1

Lattice energy of sodium chloride is relatively more than the hydration energy, such that overall enthalpy ( enthalpy os solution) is endothermic. Hence, the sodium chloride draws heat from the solvent making the solution a bit colder, while dissolving.

The favourable conditions for the formation of the solution involve a negative value for ∆H for the solution, ie, when the heat released on hydration is more than the heat required to overcome the force of attraction ie the lattice enthalpy.

Example 3.

Solvation of a solute in a solid state (Potassium chloride) is shown as follows:

KCl(s) + aq.   KCl(aq.),  ΔsolHo= 4.4 kcal

When an ionic chemical dissolves in a solvent, the ions in the crystal lattice lose their ordered position. i.e., in this case, KCl(s) will separate into its gaseous constituent ions K+ and Cl-.

KCl(s)  K+ (g) + Cl-(g)

The amount of energy required to break 1 mole of KCl((s) into its constituent gaseous ions is known

as lattice enthalpy and this will always be positive.

In the solution, these gaseous ions are now freer. Solvation of these ions (hydration in case

the solvent is water) will also occur at the same time. The hydration of K+ (g) and Cl-(g) are shown as

follows:

K+  (g) + Aqueous  K+ (aq.)

Cl- (g) + Aqueous  Cl-(aq.)

Certain amount of energy is liberated in the hydration of K+ (g) and Cl-(g) which is known as the

hydration enthalpy and this will always be negative.

Thus, the enthalpy of solution (solHo) is equal to the sum of lattice enthalpy (latticeHo) and

hydration enthalpy (hydrationHo) i.e., at 298 K and 1 bar pressure, the standard enthalpy of solution

(solHo) can be written as follows:

solHo = latticeHo +hydrationHo

For most of the ionic compounds, solHo is positive and the dissociation process is endothermic.

This is the reason why as the temperature rises, the solubility of salts rises as well.

The value of solHo will become positive when the lattice enthalpy (latticeHo) is higher than the hydration enthalpy (hydrationHo).

Example 4

The following is the solvation of a solute in a liquid state (Sulfuric acid):

H2SO4(l) + aq.    H2SO4(aq.),  ΔsolHo = -20.2kCal

Factors Affecting Hydration

Enthalpy of hydration of electrolytes will be more negative on lower lattice energy of the crystal and higher hydration energy of the cations and anions.

Higher hydration energy is favoured by

i) smaller cations and anions and

ii) higher charges on cations and anions.

Both these factors increase the charge density of the ions favouring more water molecules effectively surrounding the ions.

So, Enthalpy of hydration ∝ Effective Nuclear charge (Zeff) ∝ 1radius

Alkali metal ions have higher hydration enthalpy than alkali metal ions because of their higher charge in them.

Hydration energy decreases on going down the column.

Enthalpy of Hydration of Metal Ions

This is the energy released when one mole of metal ions is hydrated in infinite dilution.

Mn+(g) + aq → M++(aq)

M+(aq) the ion is surrounded by water molecules forming a weak bond.

Hydration enthalpy values of alkali and halide elemental ions are given below.

Ion hydrationH(kJmol-1)

Li+ -520

Na+ -405

K+ -321

Rb+ -300

Cs+ -277

F- -506

Cl- -364

Br- -337

I- -296

The hydration enthalpy directly depends on the charge density of the ions. The charge density which is a ratio of the charge to radius is more for smaller and higher charged ions. The higher the charge density of the metal ion, the higher will be the force of attraction by the metal ion on the oxygen of the polar water.. Thus the alkaline metal ions will be more hydrated than alkali metal ions. Down the group, the increase in size decreases the hydration and hence the hydration enthalpy.

Application of Enthalpy of Hydration

  1. As the hydration increases the effective radius of the ions, their mobility decreases under electrical current. This affects the electrolytic conductance of the solution and the process of electrolysis.
  2. To understand the difference in the hydration energy of the d-block ions. Hydration energy increases with the crystal field splitting energy of the ions.
  3. To predict the acidity of the aqueous metal ion coordination complexes. Higher the charge on the cations higher the acidity of the solution
  4. To nullify the negative effect of enthalpy of hydration in constructions. When cement is mixed with water it liberates heat while setting. The setting of cement is a slow process and hence liberates heat over time. During the setting time, the exposed surface gets cooled faster leaving the internal cement mix to retain the heat which may develop cracks due to the liberation of gases or a vacuum created. This will make the concrete less strong than expected.

Understanding this property shall help in either use of less heat-liberating cement or adopting methods to cool uniformly in bulk and surface.

Practice Problems

Q1. What will be the standard enthalpy of hydration of the NaCl given the following data:

ΔsolHo(NaCl)= 1246 KJ mol-1

ΔlatticeHo (NaCl)= 1446 KJmol-1

  1. -200 KJmol-1
  2. -300 KJmol-1
  3. -400 KJmol-1
  4. -500 KJmol-1

Answer: (A)

By using,

ΔsolHo(NaCl)=ΔlatticeHo(NaCl) +ΔhydrationHo(NaCl)

1246=1446 +ΔhydrationHo(NaCl)

ΔhydrationHo(NaCl) = 1246 - 1446 = -200 kJmol-1

Q2. What will be the standard enthalpy of hydration of chloride ion given the following data:

ΔhydrationHo(Na+)=- 1335 KJ mol-1

ΔhydrationHo (NaCl)= - 1446 KJmol-1

A) -111 KJmol-1
B) -222 KJmol-1
C) -333 KJmol-1
D) -444 KJmol-1

Answer: Option (A)

By using,

ΔhydrationHo(NaCl)=ΔhydrationHo(Na+) +ΔhydrationHo(Cl-)

On rearranging the above equation we get,

ΔhydrationHo(Cl-)= ΔhydrationHo(NaCl)- ΔhydrationHo(Na+)

ΔhydrationHo(Cl-)= -1446 -(-1335) = -111 KJmol-1

Q3. What will be the standard enthalpy of solvation of the NaCl given the following data:

ΔhydrationHo(NaCl)= -1333 KJ mol-1

ΔlatticeHo (NaCl)= 1446 KJmol-1

A) -200 kJmol-1
B) 113 kJmol-1
C) -113 kJmol-1
D) -204 kJmol-1

Answer: (B)

By using,

ΔsolvHo(NaCl)=ΔlatticeHo(NaCl) +ΔhydrationHo(NaCl)

ΔsolvHo(NaCl)=1446 + (-1333)

ΔsolvHo(NaCl) = 1446 - 1333 = 113 kJmol-1

Q4. One mole of anhydrous MgCl2 dissolves in water and librates 40 cal mol-1 of heat. ∆Hhydration of MgCl2= – 35 cal mol-1. Heat of dissolution of MgCl2 .7H2O(s) is:-

  1. +5 cal mol-1
  2. –5 cal mol-1
  3. 55 cal mol-1
  4. –55 cal mol-1

Answer: (A)

Heat liberated when MgCl2 dissolves in water = 40 cal mol-1

Heat of ionisation = 30 cal mol-1

Heat of hydration = - 35 cal mol-1

We have,

Heat of solution = heat of ionisation + heat of hydration

= 40 + (-35)

= 5 cal mol-1

As the heat of ionisation is more than the heat of hydration, so the heat of solution will be positive.

Frequently Asked Questions-FAQs

Q1. What is the difference between the enthalpy of solution and the enthalpy of solvation?
Answer:
The enthalpy change when one mole of a substance dissolves in a specified amount of solvent is referred to as enthalpy of the solution while the enthalpy change when one mole of a substance dissolves in a large amount of solvent(infinite dilution) is referred to as enthalpy of solvation.

Q2. Can lattice enthalpy be negative?
Answer:
No, Lattice enthalpy can never be negative as it is the energy required to break 1 mole of the ionic compound into its constituent gaseous ions.

Q3.Why can the enthalpy of the solution be positive or negative?
Answer:
Depending on the type of reaction, this enthalpy of the solution (solH) might either be positive or negative. When the reaction is endothermic, it is positive; when it is exothermic, it is negative.

Q4. Is the standard enthalpy of hydration exothermic or endothermic?
Answer:
There is no endothermic step in hydration since there are no bonds to be broken; instead, energy is only released when the bonds between the ions and the water are created. As in hydration, there is the only release of energy, so the standard enthalpy of hydration is exothermic.

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