Delta H fusion (ΔH_fus) is the amount of energy needed to convert a solid into a liquid at its melting point. For water, that value is 6.02 kJ per mole, meaning it takes 6.02 kilojoules of energy to melt 18 grams of ice into liquid water. The term “fusion” is simply the scientific word for melting, and the “delta H” refers to the change in enthalpy, which is a measure of energy in a system.
What Happens at the Molecular Level
In a solid, molecules are locked into a rigid structure, held in place by attractive forces between them. When you add heat to a solid at its melting point, that energy doesn’t make the molecules move faster. Instead, it works against those intermolecular attractions, pulling molecules apart just enough to let them slide past one another. This is why the substance transitions from a fixed, ordered solid into a flowing liquid.
The key insight here is that temperature stays completely flat during the entire melting process. If you’re heating a block of ice at 0°C, the thermometer won’t budge above 0°C until every bit of ice has melted. All the energy you’re adding goes into breaking those molecular bonds rather than increasing kinetic energy. Only after the phase change is complete does the temperature start rising again. This flat region on a heating curve is sometimes called a “plateau,” and it’s a direct reflection of the enthalpy of fusion at work.
The Formula and Its Units
The basic calculation for total heat absorbed during melting is straightforward:
ΔH = n × ΔH_fus
Here, “n” is the number of moles of the substance, and ΔH_fus is the molar enthalpy of fusion for that material. The result gives you the total energy in kilojoules required to melt that amount of solid.
The standard unit for molar enthalpy of fusion is kilojoules per mole (kJ/mol). You’ll also see it reported as joules per gram (J/g), which is more practical when you’re working with a known mass rather than counting moles. For water, the molar value is 6.02 kJ/mol, while the per-gram value is 334 J/g. Both describe the same property, just scaled differently.
Why the Value Is Always Positive for Melting
Melting requires energy input, so ΔH_fus is always a positive number. The process is endothermic: heat flows into the substance from its surroundings. The reverse process, freezing (sometimes called crystallization), releases exactly the same amount of energy back into the surroundings. Freezing is exothermic, so its enthalpy change carries a negative sign. If ice absorbs 6.02 kJ/mol to melt, then liquid water releases 6.02 kJ/mol when it freezes.
How Values Vary Across Substances
The enthalpy of fusion depends directly on the strength of the forces holding a substance’s molecules or atoms together in the solid state. Substances with strong intermolecular attractions need more energy to melt and therefore have higher ΔH_fus values. Substances with weak attractions between their particles melt easily and have low values.
Water’s enthalpy of fusion (334 J/g) is unusually high for a small molecule because of the extensive hydrogen bonding network in ice. Each water molecule forms hydrogen bonds with up to four neighbors, creating a rigid crystalline lattice that demands significant energy to dismantle. Metals, by contrast, are held together by metallic bonding, and their fusion values span a wide range depending on how tightly packed and strongly bonded their atoms are. Nonpolar substances like methane, which rely on very weak attractions between molecules, have comparatively low enthalpies of fusion.
Connection to Entropy
Melting increases disorder. A solid’s molecules are arranged in a neat, repeating pattern, while a liquid’s molecules are jumbled and free to move. This increase in disorder is measured as entropy (ΔS), and it connects to the enthalpy of fusion through a clean relationship:
ΔS = ΔH_fus / T
Here, T is the melting point expressed in Kelvin. This equation comes from the second law of thermodynamics and applies because melting at the exact melting point is a reversible process, meaning the solid and liquid are in perfect equilibrium. A substance with a high enthalpy of fusion relative to its melting temperature undergoes a large entropy increase when it melts.
Practical Applications
The enthalpy of fusion isn’t just a textbook number. It plays a central role in thermal energy storage, where engineers use phase change materials (PCMs) to capture and release heat on demand. The idea is simple: a material absorbs a large amount of energy as it melts, stores that energy without a temperature change, and then releases it later when it solidifies. The melting enthalpy of the material determines how much energy it can store per gram.
This technology shows up across industries. Food manufacturing uses phase change storage in sterilization and pasteurization processes. Agriculture and textile production use it in drying operations. The approach helps bridge gaps between when energy is available (peak collection hours) and when it’s actually needed, reducing waste and improving efficiency. Water itself, with its exceptionally high heat of fusion, is one of the most effective and cheapest phase change materials available, which is why ice-based cooling systems remain common in building climate control.