Chemistry · Thermodynamics · NEET
Lattice enthalpy is the energy needed to pull one mole of a solid ionic compound apart into gas ions. The force between two ions gets stronger when the ions are closer together. F- is a very small anion, so the cation and F- sit very close. This short distance means a strong pull between them, so the crystal is very hard to break. That is why fluorides have a high lattice enthalpy. Cl- is bigger, ions sit farther apart, the pull is weaker, and the lattice enthalpy is lower. For NEET, remember: smaller ion = shorter distance = stronger force = higher lattice enthalpy.
High lattice enthalpy makes a salt LESS soluble (harder to dissolve). Lattice enthalpy is the energy holding the crystal together. To dissolve, water must supply enough energy (through hydration) to break this crystal apart. If the lattice enthalpy is very high, hydration often cannot overcome it, so the salt stays solid. So high lattice enthalpy fights against dissolving. Many students flip this - be careful.
NCERT gives: enthalpy of solution = lattice enthalpy + hydration enthalpy (Delta_sol H = Delta_lattice H + Delta_hyd H). Here lattice enthalpy is a positive number (energy needed to break the crystal), and hydration enthalpy is a negative number (energy released when water surrounds the ions). A salt dissolves easily when hydration enthalpy is large enough (very negative) to beat the lattice enthalpy. NCERT states the rule plainly: for a salt to dissolve, its solvation (hydration) enthalpy must be greater than its lattice enthalpy.
Good question - both do rise for smaller ions, but not by the same amount. A small F- has both a higher lattice enthalpy AND a more negative hydration enthalpy. The problem is that lattice enthalpy often rises MORE than hydration enthalpy for many fluoride salts. When the two do not fully cancel, the leftover lattice energy wins, and the salt dissolves poorly. This is why solubility depends on the BALANCE between the two, not on either one alone.
NaCl is much more soluble than NaF. NaF has the smaller F- ion, so it has a higher lattice enthalpy, and water cannot easily break it apart. NaCl has the larger Cl- ion, a lower lattice enthalpy, so it dissolves readily (its enthalpy of solution is close to zero, so NaCl dissolves with very little heat change - an NCERT point). This is a direct example of the small-ion, high-lattice, low-solubility trend.
No. The rule explains why MANY fluorides dissolve poorly, but not all. Solubility depends on the balance of lattice and hydration enthalpy for each specific salt, plus the size and charge of the cation. Some fluorides (like NaF or KF) still dissolve to some extent because hydration is enough. The concept explains a general trend for NEET, not an absolute rule for every compound.
Try the real previous-year questions from this chapter — each with the answer and a full solution.
Lattice enthalpy is the energy needed to completely separate one mole of a solid ionic compound into its gaseous ions. NCERT example: lattice enthalpy of NaCl is 788 kJ/mol.
F- is the smallest halide ion, so the cation and F- sit closest together. Closer ions feel a stronger electrostatic force, giving the highest lattice enthalpy among the halides.
Hydration enthalpy is the energy released when gaseous ions are surrounded by water molecules. It is negative because energy is given out. It is the energy that helps break the crystal and dissolve the salt.
NCERT states a salt dissolves when its solvation (hydration) enthalpy is greater than its lattice enthalpy, so the hydration energy can overcome the crystal's holding energy.
Yes. It applies the enthalpy of solution relation (Delta_sol H = Delta_lattice H + Delta_hyd H) to a real trend. NEET can ask which salt is more soluble or how lattice and hydration enthalpy control dissolving.
