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You heat water on a stove. The temperature rises steadily — until it reaches 100°C. Then, no matter how high you turn the heat, the temperature stays at exactly 100°C while the water boils. You are pouring energy in, but the temperature is not changing. Where is that energy going?
You heat water on a stove. The temperature rises steadily — until it reaches 100°C. Then, no matter how high you turn the heat, the temperature stays at exactly 100°C while the water boils. You are pouring energy in, but the temperature is not changing. Where is that energy going?
Phase transitions — melting, boiling, freezing — require energy without any temperature change. This hidden energy is called latent heat. It breaks or forms the bonds between molecules without changing their kinetic energy.
During a phase transition, energy is absorbed or released at constant temperature. This energy (latent heat) goes into breaking or forming intermolecular bonds, not changing molecular kinetic energy. The specific latent heat L is the energy per unit mass for the transition.
There are two key latent heats: L_f (fusion/melting) and L_v (vaporisation/boiling). For water: L_f = 334,000 J/kg and L_v = 2,260,000 J/kg. The large latent heat of vaporisation explains why sweating is so effective at cooling — evaporating just 1 g of sweat removes 2,260 J from your body. The total energy to transform a substance follows three stages: (1) heating solid: Q = mcΔT using c_solid; (2) melting: Q = mL_f at constant T; (3) heating liquid: Q = mcΔT using c_liquid; (4) boiling: Q = mL_v at constant T; (5) heating gas: Q = mcΔT using c_gas. Water's uniquely high latent heats are why oceans moderate climate (they absorb enormous heat with little temperature change) and why steam burns are so dangerous (steam releases 2260 J/g when condensing on skin, on top of the temperature difference).