What makes a substance undergo a phase change

A phase change is when matter changes to from one state solid, liquid, gas, plasma to another. These changes occur when sufficient energy is supplied to the system or a sufficient amount is lost , and also occur when the pressure on the system is changed.

The temperatures and pressures under which these changes happen differ depending on the chemical and physical properties of the system. The energy associated with these transitions is called latent heat. Water is a substance that has many interesting properties that influence its phase changes. The pressure affects these transition points, so for water, the boiling point actually decreases as the pressure decreases.

Water is a good substance to use as an example because many people are already familiar with it. Other substances have melting points and boiling points as well. Again, consider H 2 O as an example. Only after all of the solid has melted into liquid does the addition of heat change the temperature of the substance.

For each phase change of a substance, there is a characteristic quantity of heat needed to perform the phase change per gram or per mole of material.

Remember that a phase change depends on the direction of the heat transfer. If heat transfers in, solids become liquids, and liquids become solids at the melting and boiling points, respectively. If heat transfers out, liquids solidify, and gases condense into liquids. At these points, there are no changes in temperature as reflected in the above equations. How much heat is necessary to melt The heat of fusion of H 2 O is Note the units on these quantities; when you use these values in problem solving, make sure that the other variables in your calculation are expressed in units consistent with the units in the specific heats or the heats of fusion and vaporization.

This phase change is called sublimation. Each substance has a characteristic heat of sublimation associated with this process. We encounter sublimation in several ways. These are commonly used to visually show the relationship between phase changes and enthalpy for a given substance. In Figure At the melting point, the heat added is used to break the attractive intermolecular forces of the solid instead of increasing kinetic energy, and therefore the temperature remains constant.

After all the solid has melted, once again, the heat added goes to increasing the kinetic energy and temperature of the liquid molecules until the boiling point. At the boiling point, once again, the heat added is used to break the attractive intermolecular forces instead of supplying kinetic energy, and the temperature remains constant until all liquid has been turned to gas. Skip to content Chapter Solids and Liquids. Describe what happens during a phase change.

Calculate the energy change needed for a phase change. What is the energy change when Without a sign, the number is assumed to be positive.

Test Yourself What is the energy change when g of C 6 H 6 freeze at 5. Answer kJ. Phase changes can occur between any two phases of matter. All phase changes occur with a simultaneous change in energy. All phase changes are isothermal. Questions What is the difference between melting and solidification? What is the difference between boiling and condensation?

Describe the molecular changes when a solid becomes a liquid. Describe the molecular changes when a liquid becomes a gas. It can effuse through solids like a gas , and dissolve materials like a liquid. Carbon dioxide and water are the most commonly used supercritical fluids, as they are used for decaffeination and power generation, respectively.

In general terms, supercritical fluids have properties between those of a gas and a liquid. The critical properties of some substances used as solvents and as supercritical fluids are shown in Table 1. Table 2 shows density, diffusivity, and viscosity for typical liquids, gases, and supercritical fluids. Critical Properties of Various Solvents : Supercritical fluids have properties between those of a gas and a liquid. In addition, there is no surface tension in a supercritical fluid, as there is no liquid to gas phase boundary.

One of the most important properties of supercritical fluids is their ability to act as solvents. Solubility in a supercritical fluid tends to increase with the density of the fluid at constant temperature.

Since density increases with pressure, solubility tends to increase with pressure. The relationship with temperature is a little more complicated. At constant density, solubility will increase with temperature. However, close to the critical point, the density can drop sharply with a slight increase in temperature.

Therefore, close to the critical temperature, solubility often drops with increasing temperature, then rises again. All supercritical fluids are completely miscible with each other; therefore a single phase for a mixture can be guaranteed if the critical point is exceeded.

The critical point of a binary mixture can be estimated as the arithmetic mean of the critical temperatures and pressures of the two components,. For greater accuracy, the critical point can be calculated using equations of state, such as the Peng Robinson or group contribution methods.

Other properties, such as density, can also be calculated using equations of state. In the pressure-temperature phase diagram of CO 2 , the boiling separates the gas and liquid region and ends in the critical point, where the liquid and gas phases disappear to become a single supercritical phase.

At well below the critical temperature, e. The system consists of 2 phases in equilibrium, a dense liquid and a low density gas. As the critical temperature is approached K , the density of the gas at equilibrium becomes denser, and that of the liquid becomes lower. At the critical point, Thus, above the critical temperature a gas cannot be liquified by pressure. At slightly above the critical temperature K , in the vicinity of the critical pressure, the line is almost vertical. A small increase in pressure causes a large increase in the density of the supercritical phase.

Many other physical properties also show large gradients with pressure near the critical point, such as viscosity, the relative permittivity, and the solvent strength, which are all closely related to the density. A close look at supercritical carbon dioxide : A pressure vessel made of aluminum and acrylic is filled with pieces of dry ice. The dry ice melts under high pressure, and forms a liquid and gas phase. When the vessel is heated, the CO2 becomes supercritical — meaning the liquid and gas phases merge together into a new phase that has properties of a gas, but the density of a liquid.

Supercritical CO2 is a good solvent, and is used for decaffeinating coffee, dry cleaning clothes, and other situations where avoiding a hydrocarbon solvent is desirable for environmental or health reasons.

Freezing is a phase transition in which a liquid turns into a solid when its temperature is lowered to its freezing point. Freezing, or solidification, is a phase transition in which a liquid turns into a solid when its temperature is lowered to or below its freezing point. All known liquids, except helium, freeze when the temperature is low enough. Liquid helium remains a liquid at atmospheric pressure even at absolute zero, and can be solidified only under higher pressure.

For most substances, the melting and freezing points are the same temperature; however, certain substances possess different solid-liquid transition temperatures. Most liquids freeze by crystallization, the formation of a crystalline solid from the uniform liquid.

Crystalline Solid : Model of closely packed atoms within a crystalline solid. This is a first-order thermodynamic phase transition, which means that as long as solid and liquid coexist, the equilibrium temperature of the system remains constant and equal to the melting point.

Crystallization consists of two major events: nucleation and crystal growth. Nucleation is the step in which the molecules start to gather into clusters on the scale of nanometers , arranging themselves in the periodic pattern that defines the crystal structure.