The enthalpy of solution can be calculated using the Born-Haber cycle concept applied to dissolution: (\Delta_{sol}H = \Delta_{lattice}H + \Delta_{hyd}H). The lattice enthalpy represents the energy required to separate the ionic lattice into gaseous ions (always endothermic, positive for ionic compounds). The enthalpy of hydration represents the energy released when gaseous ions are surrounded by water molecules (always exothermic, negative). For NaCl: (\Delta_{sol}H = +788 + (-784) = +4) kJ/mol. Since the value is positive, the dissolution is endothermic. This explains why dissolving NaCl in water produces a very slight cooling effect. Option -4 kJ/mol incorrectly takes hydration enthalpy as positive. Option +1572 kJ/mol adds both magnitudes without appropriate signs. Option +784 kJ/mol ignores the lattice enthalpy entirely. This is a direct NCERT Thermodynamics and Solutions cross-chapter concept. Plausibility check: NaCl dissolves nearly spontaneously at room temperature despite being slightly endothermic — entropy increase drives dissolution — confirming the small positive value makes physical sense.