Changes of State

Despite having previously examined a phase-change diagram in detail (here) I haven’t yet covered the particle model and how it links to different states of matter. So that topic, which is relevant for both chemistry and physics, is the focus of this article.

First let’s clarify the terminology… we will be confining ourselves to the three common states of matter, also known as phases, which are; solid, liquid and gas.

You are expected to be able to name the common phase changes; melting and boiling/vapourisation (going up in temperature) as well as condensing and solidification/freezing (coming down in temperature). You should also know that solids can turn directly into gases in a process known as sublimation. Similarly, gases can turn directly into solids through deposition (sometimes known as vapour deposition). These phase changes are shown in the diagram below.

Note: evaporation is not the same as boiling. Evaporation can happen at any temperature and is the process that causes rain puddles to disappear, whereas boiling happens only at one specific temperature that is unique to the liquid concerned (100 °C for water).

The particle model explains the properties of matter by thinking of all substances as a collection of particles that are normally represented as spheres (drawn as circles).

Particles in a solid are closely packed together in a regular pattern. Solids have their own shape because the particles cannot move.

Particles in a liquid are closely packed in a random way, with no particular pattern. Liquids take the shape of the container that holds them because they are able to flow past each other. Importantly, liquids always fill the bottom of a container owing to the effect of gravity. (Sometimes you will see “half-bubble” droplets of liquid on a surface: this is due to the force of the liquid’s surface tension partially resisting the gravitational force.)

Particles in a gas are far apart and free to move independently. Gases expand to fill the full volume of the container that holds them. Gas particles are in constant motion (see the previous post about Kinetic Theory) and their kinetic energy allows them to move in all directions, including upwards. (It is true that there are fewer “air particles” at high altitudes but if we are considering a relatively small container then we can say that the gas particles are evenly distributed within the full volume of the container.)

We can represent these facts using a familiar states-of-matter diagram, which should look something like the version shown below.

Note that the diagram doesn’t just show the distances between particles; it also shows information about the particles’ distributions (the shape of each state of matter). Equally obvious, but often overlooked, is the enormous jump in density between liquids and gases compared to liquids and solids. Most solids are slightly denser than liquids (except for ice, which is slightly less dense than water) whereas gases are much less dense than both of the other two states of matter.

Remember that all particles are in a constant state of motion, regardless of their state of matter. In solids and liquids the motion is very localised (in solids, we talk about vibrations rather than linear motion) whereas gas particles can move considerable distances (think about the smell of cooking moving from the kitchen into adjacent rooms of a house). This process, which is known as diffusion, occurs due to the existence of a concentration gradient.

Substances change state due to changes in the movement (kinetic energy) of their particles. If a solid is given more energy by being heated then its particles will initially just vibrate a bit more (causing a small amount of expansion) but if heating continues the particles will eventually have enough energy to break out of their fixed pattern and the substance will melt. But this doesn’t happen instantly: it takes time for all the particles to break free and the temperature of the substance stays the same while the substance is melting (the supplied energy is being used to break bonds).

A similar process happens when a heated liquid starts to boil and the particles break free of the final bonds that hold them close. Once again, the temperature remains constant even though energy continues to be supplied to the substance. This means that heated water cannot go above 100 °C for as long as there is still water to boil – and this fact can produce some unexpected effects. For example, if you blow up a balloon and hold it over a lit match, it will burst almost instantly. But if you put some water inside the balloon it will stay in tact because the energy supplied (heat) is being used to boil the water instead of damaging the rubber, as shown below.

With this foundation knowledge in place you should now be in a position to understand the more detailed examination of phase-change diagrams, which was posted previously and is available here.