Chemistry · Thermodynamics · NEET
In a solid, particles are held together tightly by strong intermolecular forces. To melt it, you must add heat to loosen these forces so particles can move freely as a liquid. The system absorbs this heat, so the enthalpy change is positive. NCERT states it directly: 'Melting of a solid is endothermic, so all enthalpies of fusion are positive.' For water, Delta_fus H = +6.00 kJ/mol.
Both are endothermic, meaning they absorb heat, so their enthalpy values are positive. Vaporization (liquid to gas) needs even more heat than fusion because you must fully separate the particles into a gas: for water Delta_vap H = +40.79 kJ/mol. Sublimation (solid straight to gas) needs the most because you break out of a solid AND spread into a gas: dry ice (solid CO2) has Delta_sub H = +25.2 kJ/mol.
During a phase change the temperature does NOT rise. The heat you add is not raising temperature; it is used only to break the forces holding particles together. That is why melting ice stays at 273 K until all the ice is gone. The absorbed heat becomes potential energy of the separated particles, so Delta H is still positive.
Sublimation goes solid to gas in one step, so it must do the job of BOTH fusion (solid to liquid) AND vaporization (liquid to gas). By Hess's law, Delta_sub H = Delta_fus H + Delta_vap H (at the same temperature). Since you add two positive numbers, sublimation enthalpy is the largest and clearly positive.
Only the REVERSE processes are negative (exothermic): freezing, condensation and deposition release heat. So freezing water gives -6.00 kJ/mol. But the named quantities 'enthalpy of fusion, vaporization, sublimation' always refer to the forward (particles spreading apart) direction, which is always positive. NEET tests this: reverse the equation, flip the sign.
No. Positive Delta H (endothermic) does not block a process. Ice melts on its own at room temperature even though it absorbs heat, because entropy increases (solid to liquid to gas raises disorder). Spontaneity depends on Delta G = Delta H - T*Delta S, not on Delta H alone. So endothermic phase changes still happen naturally.
Consider the reaction 2H2(g) + O2(g) -> 2H2O(g), Delta_r H = -483.64 kJ. What is the enthalpy change for the decomposition of one mole of water?
In which of the following processes does entropy increase? A. A liquid evaporates to vapour. B. Temperature of a crystalline solid is lowered from 130 K to 0 K. C. 2NaHCO3(s) -> Na2CO3(s) + CO2(g) + H2O(g). D. Cl2(g) -> 2Cl(g).
Try the real previous-year questions from this chapter — each with the answer and a full solution.
Yes. NCERT states 'Melting of a solid is endothermic, so all enthalpies of fusion are positive.' Every solid needs added heat to melt, so Delta_fus H is positive for all substances. The exact value depends on how strong the intermolecular forces are.
From NCERT: Delta_fus H = +6.00 kJ/mol (ice to water), Delta_vap H = +40.79 kJ/mol at its boiling point (water to steam). Water's high vaporization value comes from strong hydrogen bonds that must be broken.
At the same temperature, Delta_sub H = Delta_fus H + Delta_vap H. Sublimation equals melting plus vaporizing because solid to gas is the same as solid to liquid to gas. This is a direct application of Hess's law and is a common NEET link.
Stronger forces need more heat to break. Water (hydrogen bonds) needs about 40.79 kJ/mol to vaporize, while acetone (weaker dipole-dipole forces) needs less. So the magnitude of the enthalpy tells you how strongly particles are held together.
Yes, exactly. Water evaporating from your skin ABSORBS heat from your body (positive Delta_vap H). That absorbed heat leaves your skin, so you feel cooler. This is a real-life example NEET uses of an endothermic phase change.
