If you've ever boiled a pot of water or put ice in your drink to cool it down, you're already familiar with phase changes. A phase change is simply the change in the state of matter of a substance by the addition or subtraction of heat. What you might not know, though, is that when a substance is changing state, the temperature remains constant. This might not make any sense at first, after all, if you are adding heat or taking heat away from a substance constantly, you would expect that substance to constantly heat up or cool. However, during phase changes, the heat supplied (called the latent heat) is fully used to change the state of matter by overcoming the force of attraction between the particles (intermolecular forces).
This is best demonstrated by a heating curve (which shows the temperature against time as a substance is heated across phase changes):
All phase changes involve the absorption or release of latent heat. Endothermic phase changes absorb heat from the environment (they are cooling processes). Exothermic phase changes release heat to the environment (they are heating processes).
Exothermic (heating processes)
Endothermic (cooling processes)
The specific latent heat (l) of a substance is a measure of the heat energy (Q or Eʜ) per unit mass (m) released or absorbed during a phase change.
Specific Heat Capacity is defined like this:
Q = ml
l = Q/m
The SI Unit for Specific Latent Heat is Joules per kilogram (J/kg).
There are three basic types of latent heat each associated with a different pair of phases:
Latent heat of fusion, lғ- This is the heat energy required for the phase change between a liquid and a solid to occur without a change in temperature. This is at the temperature known as the melting point or freezing point.
Latent heat of vaporization, lᴠ- This is the heat energy required for the phase change between a liquid and a gas to occur without a change in temperature. This is at the temperature known as the boiling point or dew point.
Latent heat of sublimation- CSEC doesn't require you to know this, but it is the heat energy required for the phase change between a solid and a gas without a change in temperature. This is at the temperature known as the sublimation point or frost point.
The latent heat of vaporization of water is 2,258 kJ/kg. The latent heat of fusion of water is 334 kJ/kg
CSEC requires that you also know the difference between evaporation and boiling. After all, on the surface, they seem the same since they both result in a change from the liquid to gaseous phase.
Evaporation takes place at all temperatures, while boiling takes place only at the boiling point of the liquid.
Evaporation is a surface phenomenon (i.e. only occurs at the surface of the liquid), while boiling is a bulk phenomenon (i.e. a bubble of vapour is formed below the surface of the liquid).
Evaporation is a slow process. Boiling is a fast process.
Evaporation is affected by several factors such as surface area, wind speed, temperature and humidity. Boiling does not depend on such factors- as long as the boiling point is reached, boiling will occur.
Temperature may change during evaporation (usually a cooling effect) , whereas temperature during boiling does not change.
Boiling and Kinetic Theory
When particles in the liquid phase are heated, they gain kinetic energy and move faster and further apart. Eventually they have enough energy to escape the intermolecular forces of attraction holding them together in the liquid phase and they move very quickly and far from each other and exist in the gaseous phase. This is how boiling occurs.
Evaporation and Kinetic Theory
Within a liquid some particles have more energy than other. These "more energetic particles" may have sufficient energy to escape from the surface of the liquid as gas or vapour. This how evaporation occurs, and the result of evaporation is commonly observed when puddles or clothes dry. Evaporation is also assisted by windy conditions which help to remove the vapour particles from the liquid so that more escape.