Freezing and Melting

Freezing and melting

Freezing is the change that occurs when a liquid changes into a solid as the temperature decreases. Melting is the opposite change, from a solid to a liquid as the temperature increases. These are both examples of changes in the states of matter of substances.

Substances freeze at exactly the same temperature as they melt. As a consequence, the temperature at which—under a specified pressure—liquid and solid exist in equilibrium is defined as the melting or freezing point. When the pressure is one atmosphere, this temperature is known as the normal freezing (or melting) point. A change in pressure will change the temperature at which the change in the state of matter occurs. A decrease in pressure will decrease the temperature at which this occurs and an increase in pressure will increase the temperature required.

At a fundamental level freezing and melting represent changes in the energy levels of the molecules of the substance under consideration. Freezing is a change from a high energy state to one of lower energy, the molecules are moving less as their temperature falls. They become more ordered and fixed in shape. When a substance melts the average energy level of the constituent molecules increases. The molecules are moving more rapidly and in a less ordered manner in a liquid than in a solid. It is this greater freedom of movement that allows a liquid to flow to touch the walls of its container whereas a solid is fixed in a rigid shape. This consideration of the energy of the molecules is known as the kinetic molecular theory.

The temperature at which substances freeze and melt is different for different chemicals. The chemical formula of a substance is not necessarily a true indicator of what the freezing or melting point may be. Isomers of substances can have different physical properties including freezing and melting points. Similarly the presence of hydrogen bonds and other attractive forces such as van der Waals forces can influence the bonding within the substance and hence the freezing and melting points. If any intermolecular forces are present more energy must be added to the system to change from a solid to a liquid. This is because the intermolecular bonds have to be overcome to allow the molecules to move more freely. This is less of a change than occurs from the change from liquid to gas, because the molecules are still touching each other in both liquids and solids.

The purity of the compound can influence the temperature at which the solid-liquid change takes place. For example adding sodium chloride (common salt) or another salt to water depresses the freezing point, which is why salt is put on roads to stop their icing over. A pure substance has a definite melting or freezing point, the addition of an impurity lowers this temperature as well as spreads it so that there is a less definite, more diffuse melting or freezing point. This means that we can use the freezing or melting point as an indicator of the purity of a substance. When a solid is melted by heating or a liquid frozen while cooled, the temperature remains constant. Thus, if a graph of temperature is plotted against heat added a shoulder or plateau will be seen which represents the freezing or melting point. With an impure substance, this shoulder will not be so precise. A graph of this nature is known as a heating curve. The conversion between solid and liquid occurs at a constant temperature.

With most substances the solid is denser than the liquid phase. As a result of this when freezing the solid will sink to the bottom of the liquid. Water does not behave in this manner. Ice is less dense than water and consequently ice will float on water. Water has its maximum density at 39°F (4°C). This is caused by hydrogen bonding, which in the liquid phase is unordered. When the water freezes to form ice, the molecules assume an open ordered pattern that allows the maximum amount of hydrogen bonding. This characteristic has had a profound effect on life on Earth (e.g., it allows lakes and streams to freeze at the surface and provide insulation to life underneath the ice during frigid winter months) and results in an active agent of geological change. Because water expands when freezing it is able to crack rock ; the cyclic freezing and refreezing of water is an important weathering agent.

Normally, when we talk about a substance being a solid or a liquid we are referring to its appearance at standard temperature and pressure, this is a pressure of one atmosphere and a temperature of 68°F (20°C). If the melting point is below this temperature and the boiling point is above it then the chemical is a liquid at standard temperature and pressure.

It is possible to cool a liquid below its freezing point and still have it remain as a liquid. This is known as a super-cooled liquid. This represents an unstable equilibrium and in time the liquid freezes. It is very easy to supercool water down to 12°F (−11.1°C) and still have it remain a liquid. The super-cooled liquid will not start to freeze until there is a point for the ice to start to form. This may be a single piece of dust, which acts as a nucleation point for the ice to start forming. Supercooled water is not encountered in nature because there is too much particulate material in the atmosphere. If any of these particles lands in a supercooled liquid it will instantly turn into the solid form.

Some chemicals do not have a point at which they turn from solid to liquid—they can change directly from solid to gas, a property called sublimation. Dry ice, solid carbon dioxide , exhibits this. Like melting and freezing this also happens at one specific temperature.

Solids and liquids are both densely packed at a molecular level. One difference in terms of the molecules is that with a liquid the molecules are more readily capable of slipping over each other. It is this property that makes it easier to pour a liquid. The molecules in a liquid are still touching each adjacent molecule (as they do in a solid), although they are less freely held.

Ionic compounds generally have a higher melting point than covalent compounds. This is because the intermolecular forces in an ionic compound are much stronger. If the pressure is increased the molecules are forced closer together and this means that the intermolecular forces are holding the particles closer together and more tightly, so a higher temperature is required to make the material melt.

Melting is also called fusion, and the energy required to bring about this change of state is called the heat of fusion or the enthalpy of fusion. For ice to turn into liquid water the heat of fusion is 6.01 kJ/mol. Melting and sublimation are both endothermic processes and freezing is an exothermic process. Whenever a material changes from one state to another there is an energy change within the system. For melting the order of the system is decreasing, so energy must be supplied to increase the randomness of the molecules. For freezing the molecules are becoming more ordered, so energy is lost from the system.

Freezing and melting are the change of state from liquid to solid and from solid to liquid. For any given pure chemical they happen at a specific temperature, which is the same for freezing and melting.

See also Chemical bonds and physical properties; Chemical elements; Evaporation; Faults and fractures; Glacial landforms; Glaciation; Glaciers; Glass; Ice heaving and ice wedging; Ice