Explain solvation energy and solubility of ionic solids.
Solvation Energy
Solvation energy is the energy released when ions or molecules (usually in the gas phase) are surrounded by solvent molecules and becomes stable in solution. This is an exothermic process, that means energy released as the ion interacts and form bonds with the solvent molecules.
When ionic solids are dissolve in the water (a good polar solvent having high dielectric constant), then the ions (cations and anions) interact with water molecules to stabilize the ions and release energy called hydration energy and is represented as ΔHhyd with the unit of K-Cal/mol or K-J/mol.
M2+ + nH2O → [M(H2O)n]2+ or Maqn+
Cu2+ + 4H2O → [Cu(H2O)4]2+ or Cuaq2+
Hydration energy is a subset of solvation energy and is generally expressed as a negative value since it releases energy.
When a solid dissolves into ions the entropy increases (ΔS = +ve) as the disorder increases with the change from solid to solution and the number of particles increases. This will make ΔG negative. In that cases magnitude of solvent-solute interaction is much higher than the individual solvent-solvent interaction & solute-solute interaction. This shows that the solvated ions are more stable than gaseous ions
Solvation involves two key thermodynamic steps:
Breaking the Ionic Lattice: Energy is required to overcome the strong electrostatic forces holding the ions together in the crystal lattice. This is the lattice energy (ΔHlattice), which is always an endothermic process (positive value).
Solvating the Ions: The separated gaseous ions are surrounded by the polar solvent molecules (e.g., water). This attraction releases energy called solvation energy (ΔHsolvation). This process is always exothermic (negative value).
The overall solvation energy is the sum of the energy needed to separate ions from the lattice and the energy released when those ions are surrounded by solvent molecules. The process is only thermodynamically favoured (spontaneous) if the overall Gibbs free energy (ΔG) is negative, meaning the system loses energy and becomes more stable.
The overall energy change for the dissolution process is called the enthalpy of solution (ΔHsolution) which is the sum of these two energies:
ΔHsolution = ΔHlattice + ΔHsolvation
Solubility of Ionic Solids
Solubility is the maximum amount of a solute that can dissolve in a given amount of solvent at a particular temperature. The solubility of an ionic solid is determined by a thermodynamic balance between the lattice energy and the solvation energy.
For an ionic solid to dissolve spontaneously, the Gibbs free energy change (ΔG) for the process must be negative.
ΔG = ΔH − TΔS
Here, ΔH is the enthalpy of solution (ΔHsolution), and ΔS is the change in entropy (disorder) of the system.
Generally, the solubility of an ionic solid depends on these two factors:
If solvation energy is greater than lattice energy, the overall process is exothermic and the compound is likely to be soluble as the energy released from the solvent-ion attraction is enough to overcome the energy needed to break the lattice.
If lattice energy is greater than solvation energy, the overall process is endothermic and the compound is likely to be insoluble as the energy neede to break the lattice is too high.
Solubility increases with:
☛ High solvation/hydration energy.
☛ Lower lattice energy.
☛ Greater disorder/entropy upon dissolution.
☛ Strong interactions between ions and a polar solvent (such as water).
Factors Affecting Solvation Energy and Solubility
Several factors affects the magnitudes of lattice energy and solvation energy, thereby affecting the solubility of an ionic solid:
Ionic Charge:
Higher charges on ions lead to stronger electrostatic attractions within the crystal lattice, resulting in a higher lattice energy. Due to high charges, ion-solvent attractions are also strong which increases the solvation energy.
However, the increase in lattice energy due to higher charges is often much greater than the increase in solvation energy, which is why compounds with highly charged ions (e.g., CaCO3, BaSO4) are generally insoluble.
Ionic Size:
Lattice Energy: As ionic size increases, the distance between the ions in the lattice also increases and weakens the electrostatic forces causes the lower lattice energy.
Solvation Energy: As ionic size increases, the charge density of the ion decreases (charge is spread over a larger volume). This weakens the attraction between the ion and the solvent molecules, leading to a lower solvation energy.
The relative change in these two energies play a very important role. For example, for Group IA halides, on moving down the group, both lattice and hydration energies decrease. The solubility trend depends on which energy decreases faster.
Nature of the Solvent:
The solubility of an ionic solid is greatest in a polar solvent, like water. Polar solvents have a permanent dipole that can effectively surround and stabilize the positive and negative ions through strong ion-dipole interactions, leading to a high solvation energy.
Non-polar solvents lack this ability, so they cannot dissolve ionic solids. This is the basis of the like dissolves like principle.
Test Your Knowledge
1. What is solvation energy?
a) Energy required to break the lattice of an ionic solid.
b) Energy released when ions are surrounded by solvent molecules.
c) Energy absorbed when solvent molecules interact with each other.
d) Energy needed to remove an electron from a neutral atom.
View Answer
b) Energy released when ions are surrounded by solvent molecules.
2. What is hydration energy?
a) A measure of lattice energy in ionic solids.
b) Solvation energy when the solvent is water.
c) The entropy change during dissolution.
d) The enthalpy of solution of water.
View Answer
b) Solvation energy when the solvent is water.
3. Which of the following best describes lattice energy?
a) Energy released during ion–solvent interaction.
b) Energy required to separate gaseous ions from the lattice.
c) Energy change when a solute dissolves spontaneously.
d) Energy absorbed when entropy increases.
View Answer
b) Energy required to separate gaseous ions from the lattice.
4. Which is the correct relation for the enthalpy of solution?
a) ΔHsolution = ΔS + ΔHsolvation
b) ΔHsolution = ΔHlattice + ΔHsolvation
c) ΔHsolution = ΔG − TΔS
d) ΔHsolution = ΔHlattice − ΔHsolvation
View Answer
b) ΔHsolution = ΔHlattice + ΔHsolvation
5. Which thermodynamic parameter decides spontaneity of dissolution?
a) ΔHsolution
b) ΔS
c) ΔG
d) ΔHlattice
View Answer
c) ΔG
6. What happens to entropy (ΔS) when an ionic lattice dissolves?
a) Decreases because ions are ordered.
b) Stays constant.
c) Increases because disorder increases.
d) Becomes zero.
View Answer
c) Increases because disorder increases.
7. Why are salts with highly charged ions (e.g., BaSO4) often insoluble?
a) Their solvation energy is very high.
b) Their lattice energy is much larger than solvation energy.
c) Their entropy change (ΔS) is always negative.
d) They cannot form ionic bonds with water.
View Answer
b) Their lattice energy is much larger than solvation energy.
8. How does increasing ionic size affect solubility trends?
a) Lattice energy increases, solvation energy decreases.
b) Lattice energy decreases, solvation energy decreases.
c) Both lattice and solvation energies increase.
d) Both lattice and solvation energies remain constant.
View Answer
b) Lattice energy decreases, solvation energy decreases.
9. Why are ionic solids more soluble in water than in non-polar solvents?
a) Water molecules form covalent bonds with ions.
b) Water has a high dielectric constant and strong ion–dipole interactions.
c) Ionic solids cannot exist in non-polar solvents due to lattice energy.
d) Water decreases entropy during dissolution.
View Answer
b) Water has a high dielectric constant and strong ion–dipole interactions.
10. Which of the following factors generally increases solubility of an ionic solid?
a) High lattice energy and low solvation energy.
b) Low lattice energy and high solvation energy.
c) Low entropy change and high lattice energy.
d) Non-polar solvent medium.
View Answer
b) Low lattice energy and high solvation energy.
11. Why does NaCl dissolve in water even though its ΔHsolution is slightly positive?
a) Because ΔG is positive.
b) Because the entropy change (ΔS) is highly negative.
c) Because the entropy change (ΔS) is sufficiently positive to make ΔG negative.
d) Because lattice energy is higher than hydration energy.
View Answer
c) Because the entropy change (ΔS) is sufficiently positive to make ΔG negative.
12. Which of the following is most soluble in water?
a) NaCl
b) KCl
c) CsCl
d) LiCl
View Answer
b) KCl — solubility first increases down the group as lattice energy decreases faster than hydration energy, but decreases again for heavier halides.
13. Which salt is sparingly soluble due to extremely high lattice energy?
a) BaSO4
b) NaNO3
c) KCl
d) NH4Cl
View Answer
a) BaSO4
14. Hydration energy of small highly charged ions (e.g., Al3+) is:
a) Very low
b) Very high
c) Always zero
d) Independent of ionic charge
View Answer
b) Very high
15. Which thermodynamic condition makes dissolution always spontaneous?
a) ΔH negative, ΔS negative
b) ΔH positive, ΔS negative
c) ΔH negative, ΔS positive
d) ΔH positive, ΔS zero
View Answer
c) ΔH negative, ΔS positive
16. In Group 1 halides (NaCl, KCl, CsCl), which factor decreases faster with increasing ionic size?
a) Lattice energy
b) Solvation (hydration) energy
c) Both decrease equally
d) Neither changes
View Answer
b) Solvation (hydration) energy decreases faster than lattice energy.
17. Why is CaCO3 poorly soluble in water?
a) Because hydration energy strongly exceeds lattice energy.
b) Because hydration energy is much smaller than lattice energy.
c) Because ΔS is very large and positive.
d) Because it reacts with water.
View Answer
b) Because hydration energy is much smaller than lattice energy.
18. What is the role of a solvent’s dielectric constant in ionic solubility?
a) Higher dielectric constant reduces ion–ion attraction, increasing solubility.
b) Higher dielectric constant increases lattice energy.
c) A high dielectric constant weakens solvation energy.
d) Dielectric constant has no effect on solubility.
View Answer
a) Higher dielectric constant reduces ion–ion attraction, increasing solubility.
19. Which combination generally favours solubility?
a) Low charge density ions and nonpolar solvent
b) High lattice energy and low hydration energy
c) Low lattice energy and high hydration energy
d) High dielectric constant solvent with low entropy change
View Answer
c) Low lattice energy and high hydration energy
20. Which of the following best explains “like dissolves like” principle?
a) Ionic solutes dissolve best in non-polar solvents.
b) Solvents dissolve solutes with similar polarity nature through favourable interactions.
c) Only covalent solutes dissolve in polar solvents.
d) Solubility is independent of solvent type.
View Answer
b) Solvents dissolve solutes with similar polarity nature through favourable interactions.