Chemistry · Thermodynamics · NEET
It splits dissolving into two steps. Step 1: pull one mole of the solid apart into free gas-phase ions. This needs energy, so Delta_lattice H is positive (for NaCl it is +788 kJ/mol). Step 2: water molecules surround those ions and stick to them. This releases energy, so Delta_hyd H is negative. The total heat of dissolving is just these two added together. NEET loves this because you only need the two numbers and their correct signs.
Lattice enthalpy is the energy REQUIRED to separate one mole of a solid ionic compound into gaseous ions. You are breaking strong ion-ion attractions, so energy goes in and the sign is positive (+). Hydration enthalpy is the energy RELEASED when gas-phase ions get surrounded by water. New ion-water attractions form, so energy comes out and the sign is negative (-). Breaking costs energy; making bonds gives energy back.
Compare the sizes. If hydration enthalpy (the negative number) is LARGER in size than lattice enthalpy (the positive number), then Delta_sol H is negative or only slightly positive, and the salt usually dissolves. If lattice enthalpy is much bigger, Delta_sol H is strongly positive and the salt is nearly insoluble. NCERT gives lithium fluoride (LiF) as an example that barely dissolves because its lattice enthalpy is too high.
No. Many salts dissolve even though Delta_sol H is positive (dissolving feels cold, endothermic). Dissolving also increases disorder (entropy goes up), and spontaneity depends on Delta G = Delta H - T*Delta S, not on enthalpy alone. So a small positive Delta_sol H can still be spontaneous. This is a classic trap: 'endothermic' does not mean 'will not happen'.
Solvation enthalpy is the general word for energy released when ions are surrounded by ANY solvent. Hydration enthalpy is the special case when the solvent is water. So hydration is a type of solvation. Both are always negative because ion-solvent attractions form and release energy. NCERT notes non-polar solvents give very small solvation enthalpy, which is why ionic salts do not dissolve in oil or petrol.
Try the real previous-year questions from this chapter — each with the answer and a full solution.
Delta_sol H = Delta_lattice H + Delta_hyd H. Lattice enthalpy is positive (energy to break the crystal into gas ions) and hydration enthalpy is negative (energy released when water surrounds the ions). Add them to get the heat of dissolving.
NCERT gives the lattice enthalpy of NaCl as +788 kJ/mol. This is the energy needed to fully separate one mole of solid NaCl into gaseous Na+ and Cl- ions.
LiF has a very high lattice enthalpy because both ions are small and strongly attracted. The hydration enthalpy released is not enough to overcome it, so Delta_sol H is strongly positive and LiF is almost insoluble.
Yes. Smaller ions and higher-charge ions attract water more strongly, so they release more energy and have a more negative hydration enthalpy. This is why small, highly charged ions are strongly hydrated.
No. Entropy also matters. Since dissolving usually increases disorder, a salt can dissolve even when Delta_sol H is positive. The real test of spontaneity is Delta G = Delta H - T*Delta S.
