Melting (fusion) is the endothermic phase transition in which a solid becomes a liquid; the system absorbs energy equal to n × ΔH_fus while its temperature stays constant, because all incoming heat goes into overcoming intermolecular attractions, not raising kinetic energy (AP Chem 6.5.A).
Melting is the solid-to-liquid phase change, and on the AP exam it's all about energy bookkeeping. Per essential knowledge 6.5.A.1, energy must be transferred into a system to make a substance melt, so melting is always endothermic and the system's energy increases. The signature detail the exam loves is this: while a pure substance melts, its temperature does not change. Every joule of heat goes into breaking down the ordered solid structure (overcoming intermolecular forces) instead of speeding up the particles. That's why a heating curve has a flat plateau at the melting point.
The math is clean. The heat absorbed is q = n × ΔH_fus, where n is moles and ΔH_fus is the molar enthalpy (heat) of fusion. Melting and freezing are complementary processes (6.5.A.2), so the energy a substance absorbs while melting is exactly the energy it releases while freezing. Same number, opposite sign. Notice there's no temperature change and no specific heat in this equation. Mixing up q = nΔH_fus with q = mcΔT is one of the most common Unit 6 errors.
Melting lives in Unit 6 (Thermochemistry), mainly Topic 6.5, where learning objective 6.5.A asks you to explain the heat absorbed or released during a phase transition using moles and molar enthalpy. It also shows up in Topic 6.2 (6.2.A), because melting is a physical process you can represent on an energy diagram showing its endothermic nature. Then it reaches into Unit 7 through Topic 7.14, where melting becomes a go-to example of an endothermic process that still happens spontaneously above a certain temperature because the entropy increase wins out. If you understand why ice melts at room temperature even though ΔH is positive, you understand the heart of thermodynamic favorability.
Keep studying AP Chemistry Unit 6
Heat of Fusion (Unit 6)
ΔH_fus is the price tag for melting. It tells you how much energy one mole of a substance needs to go from solid to liquid, and it's the constant you plug into q = nΔH_fus. A bigger heat of fusion means a longer flat plateau on a heating curve, because more energy is needed before the temperature can start rising again.
Freezing Point (Unit 6)
Freezing is melting run in reverse, and the energy is a perfect mirror image (6.5.A.2). If melting 1 mol of silver absorbs 11.3 kJ, then freezing 1 mol of liquid silver releases 11.3 kJ to the surroundings. Same magnitude, flipped sign, same temperature.
Energy Diagram (Unit 6)
Under LO 6.2.A, you can draw melting as an energy diagram where the products (liquid) sit higher than the reactants (solid). That visual instantly tells you the process is endothermic, no calculation required.
Free Energy of Dissolution (Unit 7)
Melting is the classic case of an endothermic process driven by entropy. ΔH is positive, but ΔS is also positive because a liquid is more disordered than a solid. Above the melting point, the TΔS term outweighs ΔH, so ΔG goes negative and melting becomes thermodynamically favored. Topic 7.14 applies this exact logic to dissolving.
Melting shows up most often in calculation and heating-curve questions. Expect multiple-choice stems like 'how much energy is released when 53.9 g of liquid silver solidifies, given ΔH_fus = 11.3 kJ/mol' (convert grams to moles, multiply by ΔH_fus, watch the sign) or 'which segment of the heating curve represents the phase change where internal energy increases but temperature stays constant' (the flat plateau at the melting point). Ice-and-water calorimetry problems are also classic, where the heat lost by warm water equals the heat absorbed to warm and then melt the ice. You need to do three things: identify melting as endothermic, hold the temperature constant during the transition, and use q = nΔH_fus rather than q = mcΔT for the phase-change step. In FRQ settings, melting points also appear in lab contexts, like the 2022 free response involving recovered salicylic acid crystals, where phase behavior connects to identifying and purifying a solid product.
Melting and dissolving both turn a solid into something fluid, but they are different processes. Melting needs only heat. The substance stays chemically alone and changes phase at its melting point. Dissolving needs a solvent, and the particles get pulled apart and surrounded by solvent molecules instead. On the exam, melting uses ΔH_fus, while dissolution uses ΔH_solution (Topic 7.14), and dissolution can be endothermic or exothermic depending on the solute-solvent interactions. Melting is always endothermic.
Melting is always endothermic, so the system absorbs energy and its internal energy increases during the solid-to-liquid transition (6.5.A.1).
The temperature of a pure substance stays constant while it melts, which is why heating curves show a flat plateau at the melting point.
Calculate the heat of melting with q = n × ΔH_fus using moles, not q = mcΔT, which only applies when temperature is actually changing.
Melting and freezing are complementary, so the energy absorbed when a substance melts exactly equals the energy released when it freezes (6.5.A.2).
On an energy diagram, the liquid sits at higher energy than the solid, which is how you represent melting as an endothermic physical process (6.2.A).
Melting above the melting point is spontaneous despite a positive ΔH because the entropy increase makes ΔG negative, the same logic Unit 7 applies to dissolution.
Melting is the endothermic phase change from solid to liquid. The system absorbs heat equal to n × ΔH_fus while its temperature holds constant, because the energy goes into overcoming intermolecular forces rather than raising particle speed.
No. For a pure substance, temperature stays constant during melting even though heat keeps flowing in. All of that energy increases the system's potential energy by breaking down the solid structure, which is why heating curves flatten out at the melting point.
No. Melting is a phase change driven by heat alone and uses ΔH_fus, while dissolving requires a solvent and uses ΔH_solution (Topic 7.14). Dissolving can be endothermic or exothermic, but melting is always endothermic.
Use q = n × ΔH_fus for the melting step itself, since there is no temperature change during the phase transition. Use q = mcΔT only for the segments where the solid or liquid is actually warming up, like ice going from -10°C to 0°C.
Entropy. Melting increases disorder (ΔS > 0), so above the melting point the TΔS term outweighs the positive ΔH and ΔG becomes negative. That's why ice melts on its own at room temperature, and it's the same reasoning AP Chem applies to endothermic dissolution in Topic 7.14.