Energy must be supplied to raise the temperature of a solid, to melt it, to heat the liquid and to boil it.
Where does this energy go? It is worth taking a close look at a single change of state and thinking about what is happening on the atomic scale. Figure 19.3a shows a suitable arrangement. A test tube containing octadecanoic acid (a white, waxy substance at room temperature) is warmed in a water bath. At ${80^ \circ }C$, the substance is a clear liquid. The tube is then placed in a rack and allowed to cool. Its temperature is monitored, either with a thermometer or with a temperature probe and datalogger. Figure 19.3b shows typical results.

Figure 19.3: a Apparatus for obtaining a cooling curve, and b typical results.

The temperature drops rapidly at first, then more slowly as it approaches room temperature. The important section of the graph is the region BC. The temperature remains steady for some time. The clear liquid is gradually returning to its white, waxy solid state. It is essential to note that energy is still being lost even though the temperature is not decreasing. When no liquid remains, the temperature starts to drop again.
From the graph, we can deduce the melting point of octadecanoic acid. This is a technique used to help identify substances by finding their melting points.

Heating ice

In some ways, it is easier to think of the experiment in reverse. What happens when we heat a substance?
Imagine taking some ice from the deep freeze. Put the ice in a well-insulated container and heat it at a steady rate. Its temperature will rise; eventually, we will have a container of water vapour.
Water vapour and steam mean the same thing–an invisible gas. The ‘steam’ that you see when a kettle boils is not a gas; it is ‘wet steam’ – a cloud of tiny droplets of liquid water.
Figure 19.4 shows the results we might expect if we could carry out this idealised experiment. Energy is supplied to the ice at a constant rate. We will consider the different sections of this graph in some detail, in order to describe where the energy is going at each stage.

We need to think about the kinetic and potential energies of the molecules. If they move around more freely and faster, their kinetic energy has increased. If they break free of their neighbours and become more disordered, their electrical potential energy has increased.

You know that the kinetic energy of a particle is the energy it has due to its motion. Figure 19.5 shows how the electrical potential energy of two isolated atoms depends on their separation. Work must be done (energy must be put in) to separate neighbouring atoms–think about the work you must do to snap a piece of plastic or to tear a sheet of paper. The graph shows that:
The electrical potential energy of two atoms very close together is large and negative.
As the separation of the atoms increases, their potential energy also increases.
When the atoms are completely separated, their potential energy is maximum and has a value of zero.
It may seem strange that the potential energy is negative and you will see in Chapter 21 why this is so. At the moment, just notice that, as atoms or molecules become further apart, their potential energy becomes less negative and so they have more potential energy.
Now look back at the graph shown in Figure 19.4.

Section AB

The ice starts below ${0^ \circ }C$; its temperature rises. The molecules gain energy and vibrate more and more.
Their vibrational kinetic energy is increasing. There is very little change in the mean separation between the molecules and hence very little change in their electrical potential energy.

Section BC

The ice melts at ${0^ \circ }C$. The molecules become more disordered. There is a modest increase in their electrical potential energy.

Section CD

The ice has become water. Its temperature rises towards ${100^ \circ }C$. The molecules move increasingly rapidly. Their kinetic energy is increasing. There is very little change in the mean separation between the molecules and therefore very little change in their electrical potential energy.

Section DE

The water is boiling. The molecules are becoming completely separate from one another. There is a large increase in the separation between the molecules and hence their electrical potential energy has increased greatly. Their movement becomes very disorderly.

Section EF

The steam is being heated above ${100^ \circ }C$. The molecules move even faster. Their kinetic energy is increasing. The molecules have maximum electrical potential energy of zero.
You should see that, when water is heated, each change of state (melting, boiling) involves the following:
- there must be an input of energy
- the temperature does not change
- the molecules are breaking free of one another
- their potential energy is increasing.
In between the changes of state:
- the input of energy raises the temperature of the substance
- the molecules move faster
- their kinetic energy is increasing.
The hardest point to appreciate is that you can put energy into the system without its temperature rising.
This happens during any change of state; the energy goes to breaking the bonds between neighbouring molecules. The energy that must be supplied to cause a change of state is sometimes called ‘latent heat’.
The word ‘latent’ means ‘hidden’ and refers to the fact that, when you melt something, its temperature does not rise and the energy that you have put in seems to have disappeared. 
It may help to think of temperature as a measure of the average kinetic energy of the molecules. When you put a thermometer in some water to measure its temperature, the water molecules collide with the thermometer and share their kinetic energy with it. At a change of state, there is no change in kinetic energy, so there is no change in temperature.
Notice that melting the ice (section BC) takes much less energy than boiling the same amount of water (section DE). This is because, when a solid melts, the molecules are still bonded to most of their immediate neighbours. When a liquid boils, each molecule breaks free of all of its neighbours. Melting may involve the breaking of one or two bonds per molecule, whereas boiling involves breaking eight or nine.

Evaporation

A liquid does not have to boil to change into a gas. A puddle of rain-water dries up without having to be heated to ${100^ \circ }C$. When a liquid changes to a gas without boiling, we call this evaporation. 
Any liquid has some vapour associated with it. If we think about the microscopic picture of this, we can see why (Figure 19.6). Within the liquid, molecules are moving about. Some move faster than others, and can break free from the bulk of the liquid. They form the vapour above the liquid. Some molecules from the vapour may come back into contact with the surface of the liquid, and return to the liquid. However, there is a net outflow of energetic molecules from the liquid, and eventually it will evaporate away completely.
You may have had your skin swabbed with alcohol or ether before an injection. You will have noticed how cold your skin becomes as the volatile liquid evaporates. Similarly, you can become very cold if you get wet and stand around in a windy place. This cooling of a liquid is a very important aspect of evaporation.

When a liquid evaporates, it is the most energetic molecules that are most likely to escape. This leaves molecules with a below-average kinetic energy. Since temperature is a measure of the average kinetic energy of the molecules, it follows that the temperature of the evaporating liquid must fall.