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Chapter 5

Analytical Supercritical Fluid Extraction

for Food Applications

Tracy Doane-Weideman and Phillip B. Liescheski

Isco Incorporated, Lincoln, NE 68504

Abstract

In this review, we explore the fundamental concepts of supercritical fluids and

supercritical fluid extractions. Carbon dioxide and other solvents are discussed; the

solubility theory is introduced together with the calculation of the density of carbon

dioxide. The state-of-the-art instrumentation is presented in terms of fundamental

components. The most widely used application of analytical SFE is in the food

industry and this review includes fats, oils, vitamins, and pesticides in research and

routine applications.

Introduction

Supercritical fluid extraction (SFE) is becoming an important sample preparation

method in the chemical analysis of food products, especially for fats and fatty oils.

SFE has been used successfully for over a decade in analyses of food samples

(1,2). The most popular SFE solvent is carbon dioxide (CO2). Triglycerides, cho￾lesterol, waxes, and free fatty acids are quite soluble in supercritical CO2. The sol￾ubility of polar lipids, such as phospholipids, can be improved by augmenting the

supercritical CO2 with a small addition of ethanol or other polar modifier solvent.

Even though CO2 is considered a “green-house” gas, it is ubiquitous in nature and

can be retrieved from the environment and returned clean (3). As a result, SFE can

still contribute positively to “Green Chemistry.” CO2 has the additional advantage

of being nonflammable and less toxic than most organic solvents. For example,

petroleum ether, which is commonly used in fat extractions, can be easily detonat￾ed by static electricity, and diethyl ether can form explosive peroxides. On the

other hand, some fire extinguishers use CO2, which is also commonly found in

foods and drinks such as bread and carbonated drinks. Finally, several common

chlorinated solvents are banned by law, and supercritical CO2 can be an alternative

to these solvents. All of these factors make SFE attractive.

What Is a Supercritical Fluid?

A supercritical fluid is a dense gas (4). It is compressible and thus expands to com￾pletely fill its container. A liquid, on the other hand, takes the shape of its container

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but does not expand to fill the container. Instead it settles at the bottom. Supercritical

fluids, unlike the air we breathe, have densities comparable to liquids. As a result,

these fluids have solvating power. A supercritical fluid can be defined as a form of

matter in which the liquid and gaseous phases are indistinguishable (5).

The three most common phases of matter on earth are solid, liquid, and gas.

The phase of a pure simple substance depends on the temperature and pressure. A

plot showing a substance's phase for a given temperature and pressure is called a

phase diagram. Figure 5.1 is a phase diagram for CO2. In a phase diagram, the

solid, liquid, and gas regions are divided by branches or equilibrium curves. These

curves represent valuable information concerning the substance's melting, boiling,

or sublimation temperatures at given pressures. These curves characterize the tem￾peratures and pressures at which two phases coexist in equilibrium. For example,

the liquid-gas equilibrium curve divides the liquid and gaseous phase regions. On

this curve, the substance coexists as both a liquid and gas (vapor) in equilibrium.

When the temperature and pressure change so that the substance leaves the liquid

phase region and crosses the equilibrium curve into the gas phase region, the sub￾stance boils. As its state crosses this curve, there is an entropy change. In the case

of boiling, its entropy increases and absorbs heat, known as heat (enthalpy) of

vaporization. In the case of condensation, its entropy decreases and liberates heat,

known as heat of condensation. An obvious physical change is seen in the sub￾stance as its state crosses one of these curves.

Fig. 5.1. Phase diagram of carbon dioxide.

Temperature (°C)

Pressure (atm)

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The three equilibrium curves (solid-gas, solid-liquid, liquid-gas) intersect at a

common point, called the triple point. At this point, the substance coexists in equi￾librium with all three phases. Each substance has only one triple point. The solid￾liquid equilibrium curve radiates from the triple point to infinity. The solid-gas

equilibrium curve radiates from the triple point and eventually terminates at

absolute zero and vacuum. The liquid-gas equilibrium curve does not radiate indef￾initely from the triple point but terminates at another important point, called the

critical point. This point is the critical temperature and critical pressure of the sub￾stance. Beyond the critical point, there is no longer an equilibrium curve to divide

the liquid and gaseous regions; thus, the liquid and gas phases are no longer distin￾guishable. There are no physical changes observed as the substance's state crosses

over this region. This region of the phase diagram is sometimes called the super￾critical fluid region.

The critical temperature is the temperature above which the substance can no

longer be condensed into a liquid. Increasing the pressure will not induce conden￾sation. For a liquid that partially fills a tube, the liquid's meniscus disappears when

it is heated above the critical temperature. The critical pressure is the vapor pres￾sure of the substance at its critical temperature. It is also the maximum vapor pres￾sure of the substance because at a higher temperature, the liquid phase cannot be

distinguished from its vapor. Based on the critical point, a supercritical fluid can

also be defined as the state beyond the critical temperature and critical pressure of

the substance. The critical temperature for CO2 is 31.1°C, and its critical pressure

is 72.84 atm.

Supercritical fluids can be considered on the molecular level. Molecules have

both kinetic and potential energies. The kinetic energy is related to the motion of

the molecules, which depends on the temperature. The potential energy is related

to the Van der Waal force, the close proximity attractive interaction between mole￾cules. The potential energy of molecules depends on how close they are to each

other. This attractive force between solvent molecules and solute molecules allows

for solvation, the dissolving process. It also allows solvent molecules to “stick”

together into clusters, thus forming a liquid. The molecules aggregate locally but

there is no long-range order as observed in solid crystals. In the liquid phase, exter￾nal pressure is not required to keep the molecules close together because they

already “stick” together. However, these “sticky” molecules also give rise to high￾er surface tension, viscosity, and slower diffusion. Such properties can hinder an

extraction process. In supercritical fluids, the temperature is above the solvent's

critical temperature. At these higher temperatures, molecules move more quickly

and thus have a higher kinetic energy. This higher kinetic energy reduces the sig￾nificance of the potential energy to a point at which the molecules no longer

“stick” together. As a consequence, lowered surface tension, viscosity, and faster

diffusion allow supercritical fluids to perform better during extraction. Lower sur￾face tension allows the fluid to “wet” surfaces better and to penetrate more deeply

into small pores and features. However, higher pressure is required to keep the

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