Fundamental aspects of supercritical fluids

Abstract

The supercritical state of a particular substance can be demonstrated using a general pressure–temperature (PT) diagram of pure substance, which is shown in Figure 2.1 for CO2. The three common states of matter—solid, liquid, and gas—are divided by distinct phase boundaries as shown by solid lines. Along these lines, two phases are in equilibrium and the three states coexist at the triple point. The vapor–liquid equilibrium curve, also called the boiling point or the vapor pressure curve, terminates at the critical point with a critical temperature, Tc, and a critical pressure, Pc. A supercritical fluid (SCF) is a fluid with a temperature and pressure higher than Tc and Pc. In the supercritical region shown in Figure 2.1, no phase boundaries exist and therefore only a single homogeneous phase exists regardless of pressure and temperature. Hence, it is possible for a substance to cross from a liquid state to the gas state as displayed by the PT path A to B without any phase transition (in this case boiling) by passing through the SCF region. In practice, the term SCF is used to denote fluids in the approximate reduced temperature and pressure range: Tr = 0.95–1.10 and Pr = 1.01–1.5 (Tr = T/Tc, Pr = P/Pc).