Real gases are non-hypothetical gases whose molecules occupy space and have interactions; consequently, they adhere to gas laws. An attempt to produce an equation that describes the behavior of gases more realistically is represented by the equation of real gases.
The corrections made to the equation of perfect gases are two: the proper volume of the molecules is taken into account, which is therefore no longer considered point-like, and the interactions between molecules that were neglected in the case of perfect gases are considered.
The first correction has the effect of making the gas not indefinitely compressible; its empirical finding is the liquefaction to which real gases are subjected if compressed (and cooled) sufficiently. The second correction ensures that the real gases do not expand infinitely but reach a point where they cannot occupy more volume (this is because a very small force is established between the atoms, due to the random variation of the electrostatic charges in the individual molecules, called Strength of van der Waals).
For this reason, the ideal gas law does not provide accurate results in the case of real gases, especially in low temperature and/or high-pressure conditions, while it becomes more accurate in case of rarefied gases, at high temperature and low pressure, that is when intermolecular forces and molecular volume become negligible. So, to understand the behavior of real gases, the following must be taken into account:
- compressibility effects;
- variable specific heat capacity;
- van der Waals forces;
- non-equilibrium thermodynamic effects;
- issues with molecular dissociation and elementary reactions with variable composition.
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