Chemistry · Thermodynamics · NEET
It is the heat given out or taken in when 1 mole of a substance dissolves in a stated amount of solvent (usually water) at constant pressure. The symbol is Δsol H. If heat is absorbed, Δsol H is positive (endothermic, like most salts). If heat is released, Δsol H is negative (exothermic). For NEET, remember the amount of solvent must be stated, because the value changes with how much water you use.
Δsol H = Δlattice H + Δhyd H. When an ionic salt dissolves, first the crystal must be pulled apart into free gaseous ions (this needs energy, so Δlattice H used this way is positive). Then water surrounds the ions and releases energy (Δhyd H is negative). The final heat you measure is the sum of these two steps. This is the single most tested idea for this concept in NEET.
For most ionic salts the lattice enthalpy (energy needed to break the crystal) is slightly larger than the energy released by hydration. So the break-up step costs a little more than the hydration payback, and the net Δsol H comes out positive (endothermic). That is why dissolving salts like KNO3 or NH4Cl makes the water feel colder. NCERT states that for most ionic compounds Δsol H is positive and dissolution is endothermic.
Enthalpy of hydration (Δhyd H) is only ONE step: the energy released when gaseous ions get surrounded by water molecules. It is always negative. Enthalpy of solution (Δsol H) is the WHOLE process: crystal breaking PLUS hydration together. So Δhyd H is part of Δsol H, not the same thing. Students often mix these up in NEET options.
Because for most salts dissolution is endothermic (Δsol H is positive, it absorbs heat). By Le Chatelier's principle, adding heat pushes an endothermic process forward, so more salt dissolves at higher temperature. If a salt's dissolution were exothermic (Δsol H negative), heating would lower its solubility. NCERT links this directly to the sign of Δsol H.
It means you dissolve the substance in so much water that the dissolved ions are too far apart to affect each other. At this point adding even more water no longer changes the heat, so Δsol H reaches a fixed limiting value. NCERT defines the enthalpy of solution at infinite dilution as the value when interactions between the ions are negligible.
Yes, watch the direction. Lattice enthalpy defined as crystal → gaseous ions is POSITIVE (energy needed, e.g. NaCl +788 kJ/mol). Hydration enthalpy (gaseous ions → hydrated ions) is NEGATIVE. In Δsol H = Δlattice H + Δhyd H you add a positive and a negative number, so the net can be small positive, small negative, or near zero. For NaCl the two nearly cancel, giving very little heat change.
Try the real previous-year questions from this chapter — each with the answer and a full solution.
As stated per mole (Δsol H in kJ/mol) it behaves like a molar property, so its per-mole value does not depend on how much salt you take. The total heat, however, is extensive because it scales with the number of moles dissolved.
It is very small. NaCl has lattice enthalpy about +788 kJ/mol and a nearly equal negative hydration enthalpy, so they almost cancel and dissolving NaCl gives very little heat change. This is a classic NCERT example.
By calorimetry at constant pressure. You dissolve a known mole amount and measure the temperature change of the solution, then use heat = mass × specific heat × ΔT to find the heat exchanged per mole.
When the hydration enthalpy released is larger than the lattice enthalpy needed to break the crystal, making Δsol H negative. Anhydrous CaCl2 is a common exothermic example that warms the water.
Small F− ions sit close to cations, giving a very high lattice enthalpy. This large break-up cost makes Δsol H strongly positive, so many fluorides are less soluble than the corresponding chlorides. NCERT raises this exact question.
