The term "real gases" among chemists and physicists is usually called such gases, whose properties most directly depend on their intermolecular interaction. Although in any specialized reference book it can be read that one mole of these substances under normal conditions and steady state occupies a volume of approximately 22.41108 liters. Such a statement is true only for the so-called "ideal" gases, for which, according to the Clapeyron equation, the forces of mutual attraction and repulsion of molecules do not act, and the volume occupied by the latter is negligible.
Of course, such substances do not exist in nature, so all these arguments and calculations have a purely theoretical orientation. But real gases, which to one degree or another deviate from the laws of ideality, are encountered very often. Between the molecules of such substances there are always forces of mutual attraction, from which it follows that their volume is somewhat different from the derived perfect model. Moreover, all real gases have varying degrees of deviation from ideality.
But here there is a very clear tendency: the higher the boiling point of a substance is close to zero degrees Celsius, the stronger this compound will differ from the ideal model. The equation of state of real gas, owned by the Dutch physicist Johannes Diderik Van der Waals, was derived by him in 1873. Two very significant corrections are introduced into this formula, which has the form (p + n 2 a / V 2 ) (V - nb) = nRT, compared to the Clapeyron equation (pV = nRT), which are determined experimentally. The first of them takes into account the forces of molecular interaction, which are affected not only by the type of gas, but also by its volume, density and pressure. The second amendment determines the molecular weight of the substance.
These adjustments acquire the most significant role at high gas pressures. For example, for nitrogen at a rate of 80 atm. calculations will differ from ideality by about five percent, and with an increase in pressure to four hundred atmospheres, the difference will reach one hundred percent. It follows that the laws of an ideal gas model are very approximate. Deviation from them is both quantitative and qualitative. The first is that the Clapeyron equation is observed for all real gaseous substances very approximately. Derogations of a qualitative nature are much deeper.
Real gases can very well be transformed into both liquid and solid state of aggregation, which would be impossible if they strictly followed the Clapeyron equation. The intermolecular forces acting on such substances lead to the formation of various chemical compounds. This is again impossible in a theoretical ideal gas system. The bonds thus formed are called chemical or valence bonds. In the case when the real gas is ionized, Coulomb attraction forces begin to appear in it, which determine the behavior of, for example, a plasma, which is a quasi-neutralized ionized substance. This is especially true in light of the fact that the plasma physics today is a vast, rapidly developing scientific discipline, which has an extremely wide application in astrophysics, the theory of radio wave signals, and the problem of controlled nuclear fusion reactions.
Chemical bonds in real gases by their nature practically do not differ from molecular forces. Both of them, by and large, come down to the electrical interaction between elementary charges, of which the entire atomic and molecular structure of matter is built. However, a complete understanding of molecular and chemical forces became possible only with the advent of quantum mechanics.
It is worth recognizing that not every state of matter compatible with the equation of the Dutch physicist can be implemented in practice. For this, a factor of their thermodynamic stability is also required. One of the important conditions for such stability of the substance is that in the isothermal pressure equation the tendency for a decrease in the total body volume must be strictly observed. In other words, as the value of V increases, all isotherms of the real gas must steadily drop. Meanwhile, on the isothermal graphs of Van der Waals, rising sections are observed below the critical temperature level. Points lying in such zones correspond to an unstable state of a substance, which cannot be realized in practice.