A supercritical fluid (SCF) is any compound at a temperature and pressure above the critical values. The critical temperature of a compound is defined as the temperature above which a pure, gaseous component cannot be liquefied regardless of the pressure applied. The critical pressure is then defined as the vapour pressure of the gas at the critical temperature. The temperature and pressure at which the gas and liquid phases become identical is the critical point. In the supercritical environment only one phase exists. The fluid, as it is termed, is neither a gas nor a liquid and is best described as intermediate to the two extremes. This phase retains the solvent power common to liquids as well as the transport properties common to gases.

 

 

 

As carbon dioxide is the most commonly used SCF, the pressure-temperature diagram for carbon dioxide is presented in Figure 1 to illustrate the differences between the gas, liquid and supercritical states. The critical point (CP) is marked at the end of the gas-liquid equilibrium curve, and the supercritical fluid region is indicated by the shaded area. It can be shown that by using a combination of isobaric changes in temperature with isothermal changes in pressure, it is possible to convert a pure component from liquid to gas (and vice versa) via the supercritical region without incurring a phase transition. 

 

 

Table 2. Critical conditions for various supercritical fluids.
 

More information regarding SCF technology will come in the following weeks.