Chemistry · Thermodynamics · NEET
It is always negative. When gaseous ions enter water, the water molecules turn their opposite charge toward the ion and pull on it. This pulling forms new ion-water attractions and gives out heat. Any process that releases heat has a negative enthalpy sign. So ΔhydH is negative for both the positive ion and the negative ion. NEET often gives a wrong option saying it is positive to trap you.
They are different steps. Hydration only measures the heat released when free gaseous ions get wrapped in water, so it is negative. Enthalpy of solution is the total heat when a solid dissolves. To dissolve, you must first break the crystal (lattice enthalpy, which needs energy and is positive), then hydrate the ions (negative). The total ΔsolH = lattice enthalpy + hydration enthalpy. If breaking the crystal costs more than hydration gives back, the total is positive (endothermic). For NaCl it is +788 − 784 = +4 kJ/mol, only slightly positive.
Lattice enthalpy is the energy needed to pull an ionic solid apart into free gaseous ions. It is positive (energy in) when written as breaking the lattice. Hydration enthalpy is the energy released when those free gaseous ions get surrounded by water. It is negative (energy out). One breaks the crystal, the other builds ion-water bonds. Both together decide if a salt dissolves easily.
A smaller ion has MORE (more negative) hydration enthalpy. Charge sits on a smaller surface, so the charge density is higher and water is pulled more strongly. So Li+ is hydrated more strongly than Na+ or K+. Higher charge also increases it: Mg2+ (small and 2+) is hydrated far more strongly than Na+ (large and 1+). Remember: smaller size and higher charge → more negative ΔhydH.
It is the heat change when 1 mole of a gas-phase ion, like Na+(g) or Cl−(g), dissolves in a large amount of water to give the ion in solution, like Na+(aq). Example: Na+(g) + water → Na+(aq), ΔhydH is negative. It tells you how strongly water holds that ion. NEET uses it inside the equation ΔsolH = lattice enthalpy + hydration enthalpy.
Water is polar. Its oxygen end is slightly negative and its hydrogen end is slightly positive. Near a positive ion, the oxygen ends point toward the ion; near a negative ion, the hydrogen ends point toward it. These new attractions (ion-dipole bonds) are stable, and forming stable bonds always gives out heat. That released heat is the negative hydration enthalpy.
Try the real previous-year questions from this chapter — each with the answer and a full solution.
Yes. Forming ion-water attractions always releases heat, so ΔhydH is negative for every ion. There is no ordinary ion with a positive hydration enthalpy.
NCERT uses ΔhydH = −784 kJ/mol for NaCl (sum of Na+ and Cl− hydration). With lattice enthalpy +788 kJ/mol, the enthalpy of solution is +788 − 784 = +4 kJ/mol.
Li+ has the higher (more negative) hydration enthalpy. It is the smallest, so its charge is concentrated and water is pulled in more strongly. Hydration enthalpy decreases down a group as ion size grows.
A salt dissolves easily when hydration enthalpy is large enough to overcome the lattice enthalpy. If hydration energy released is close to or more than the lattice energy needed, ΔsolH is near zero or negative and the salt is soluble.
Yes. Higher charge makes it more negative. Mg2+ is hydrated much more strongly than Na+ because a 2+ charge pulls water molecules more strongly than a 1+ charge of similar size.
