Real Gases. Joule-Thomson Effect

The Joule-Thomson effect is a measure of the deviation of the behavior of a real gas from what is defined to be ideal-gas behavior. In this experiment a simple technique for measuring this effect will be applied to a few common gases. 

Joule-Thomson effect -> When attractive forces dominate, a real gas cools as it expands. 

The Joule-Thomson effect is the cooling of a real gas by an adiabatic expansion (without additional work). 

The solution to Problem 2.33 shows that the Joule-Thomson coefficient can be expressed in terms of the parameters representing the attractive and repulsive interactions in a real gas. If the attractive forces predominate, then expanding the gas will reduce its energy and hence its temperature. This reduction in temperature could continue until the temperature of the gas falls below its condensation point. This is the principle underlying the liquefaction of gases with the Linde refrigerator, which utilizes the Joule-Thomson effect. See Section 2.12 for a more complete discussion. 

Ideal and real gas, heat capacity, Joule-Thomson effect, physical transformations of pure substances, and transport properties. 

During adiabatic expansion of a real gas the temperature may decrease (Joule-Thomson effect), which is used for liquefaction of gases, for example, for air separation. 

The enthalpy H of an ideal gas depends only on temperature and not on pressure because of the absence of intramolecular forces. In a real gas these forces cannot be neglected, and H depends on pressure. In most cases this leads to a decrease in temperature if a real gas is adiabatically expanded (Joule-Thomson effect, named after James Prescott Joule and William Thomson, see boxes). This property is relevant for many practical applications, for example, for refrigerators, heat pumps, and for the coohng and liquefaction of gases (Linde process). 

The Joule-Thomson effect (or Joule-Kelvin effect or Kelvin-Joule effect) [3-6] describes the temperature increase or decrease of a liquid or a real gas such as natural gas, CO2 or N2 when it expands freely from high pressure to low pressure at a constant enthalpy condition (i.e., adiabatic expansion) where no heat is transferred to or from the fluid and no external mechanical woric is extracted from the fluid. 

Equations of state (EOS) offer many rich enhancements to the simple pV = nRT ideal gas law. Obviously, EOS were developed to better calculate p, V, and T, values for real gases. The point here is such equations are excellent vehicles with which to introduce the fact that gases cannot be really treated as point spheres without mutual interactions. Perhaps the best demonstration of the existence of intermolecular forces that can also be quantified is the Joule-Thomson experiment. Too often this experiment is not discussed in the physical chemistry course. It should be. The effect could not exist if intermolecular forces were not real. The practical realization of the effect is the liquefaction of gases, nitrogen and oxygen, especially. 

Later experiments, notably the Joule-Thomson experiment, have shown that Joule s law is not precisely correct for real gases. In Joule s apparatus the large heat capacity of the vat of water and the small heat capacity of the gas reduced the magnitude of the effect below the limits of observation. For real gases, the derivative (dU/dV) is a very small quantity, usually positive. The ideal gas obeys Joule s law exactly. 

The pressure dependence of the thermodynamic property enthalpy leads to interesting phenomena in the unrestrained, free expansion of real gases. Figure 5.9 shows a schematic of a gas flowing through a porous plug. It enters the system in state 1 at Pi and Pi and it exits at a significantly lower pressure, Pg. We wish to study the effect of this so-called Joule-Thomson expansion on the temperature of the gas at the exit, Tg-... 

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