Phase Transitions

Types of Phase Transitions

The transition between different states of matter is called a phase transition.
The transition from the gaseous to the liquid state is called condensation.
When a substance goes from the liquid to the solid state, the process is called solidification (or freezing).
The reverse processes are known as melting (solid to liquid) and evaporation (liquid to gas).
If a substance transitions directly from solid to gas, this is called sublimation, while the reverse (gas to solid) is known as deposition (or resublimation).

An example is dry ice, which sublimates directly from solid to gas at room temperature and atmospheric pressure.
Normal ice, in contrast, melts into liquid water before evaporating.

Phase Diagrams

The state of a substance depends primarily on pressure and temperature.
If a gas is cooled below its condensation temperature, it becomes liquid. The same effect can be achieved by keeping the temperature constant and increasing the pressure. In both cases, the gas particles come closer together until they interact strongly enough to behave like a liquid.

This principle is applied, for example, in lighters filled with gas under high pressure, making it liquid at room temperature. When the valve is opened, the pressure drops, and the liquid evaporates. A spark then ignites the released gas.

A similar effect occurs with liquids. If either the pressure or the temperature is increased, the liquid can solidify into a solid because the particle interactions increase again.

The relationship between pressure (p) and temperature (T) can be represented in a p–T phase diagram, where pressure is on the x-axis and temperature on the y-axis. The lines between the regions in the diagram represent phase boundaries. By holding one variable constant, one can determine how much the other must be changed to induce a phase transition.

One particularly important curve is the vapor pressure curve, which separates the liquid and gas phases. In a vacuum, molecules constantly evaporate from a liquid, forming a vapor above it. As the vapor pressure increases, more molecules return to the liquid. At a certain point, determined by temperature, an equilibrium is reached. The pressure at this point is called the vapor pressure, and it can be read from the phase diagram.

Two important points also appear in phase diagrams:

• The triple point $P_T$: the unique condition at which all three phases (solid, liquid, gas) coexist in equilibrium. For water, this occurs at about 6 mbar and 0.01 °C. * The critical point $P_K$: the highest temperature and pressure at which the liquid and gas phases are distinguishable. Above this point, the substance becomes a supercritical fluid, where liquid and gas are indistinguishable.

The Density Anomaly of Water

Water is a special case because it exhibits a density anomaly.
Due to the polar nature of water molecules, water reaches its maximum density at 4 °C. Below this temperature, the density decreases—unlike most substances, where density increases as temperature decreases.
This means ice is less dense than water, which is why it floats.

As a result, lakes do not freeze solid in winter. Water colder than 4 °C rises to the top, allowing a layer of liquid water to remain at the bottom.

In the phase diagram, this anomaly is reflected by the negative slope of the boundary between solid and liquid phases—contrary to the typical behavior of other substances.