Real gases: deviation from ideality

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 an assertion is valid only for 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 negligibly small.

Of course, such substances do not exist in nature, therefore all these arguments and calculations have a purely theoretical orientation. But real gases, which to some extent deviate from the laws of ideality, are found very often. Between the molecules of such substances there are always forces of mutual attraction, from which it follows that their volume differs somewhat from the derived perfect model. And all real gases have different degrees of deviation from ideality.

But here we can see a very clear tendency: the more the boiling point of the substance approaches zero degrees Celsius, the stronger this connection will be from the ideal model. The equation of state of real gas, belonging to the Dutch physicist Johannes Diederik van der Waals, was derived by him in 1873. In this formula, having the form (p + n ² a / V ² ) (V - nb) = nRT, two very important corrections are introduced in comparison with the Clapeyron equation (pV = nRT), determined experimentally. The first of them takes into account the forces of molecular interaction, influenced not only by the type of gas, but also by its volume, density and pressure. The second correction determines the molecular mass of the substance.

The most significant role these adjustments acquire 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. The deviation from them is both quantitative and qualitative. The first is manifested in the fact that the Clapeyron equation is satisfied for all real gaseous substances very approximately. Retreats of a qualitative nature are much more profound.

Real gases can be completely transformed into both a liquid and a solid aggregate state, which would be impossible if they strictly follow the Clapeyron equation. Intermolecular forces acting on such substances lead to the formation of various chemical compounds. This is again impossible in the theoretical ideal gas system. The bonds thus formed are called chemical bonds or valence bonds. In the case when the real gas is ionized, the forces of Coulomb attraction begin to manifest in it, which determine the behavior of, for example, the plasma, which is a quasi-neutral ionized substance. This is especially true in the light of the fact that plasma physics today is a vast, rapidly developing scientific discipline, which has an extremely wide application in astrophysics, the theory of propagation of radio wave signals, in the problem of controlled nuclear and thermonuclear reactions.

Chemical bonds in real gases by their nature practically do not differ from molecular forces. Both of them, by and large, boil down to the electric interaction between elementary charges, of which the entire atomic and molecular structure of matter is constructed. However, a complete understanding of molecular and chemical forces became possible only with the advent of quantum mechanics.

It is worth acknowledging that not every state of matter compatible with the equation of the Dutch physicist can be realized in practice. For this, a thermodynamic stability factor is also needed. One of the important conditions for such stability of the substance is that the isothermal pressure equation must strictly adhere to the tendency to reduce the total body volume. In other words, as the value of V increases, all the isotherms of the real gas must be lowered steadily. Meanwhile, on the isothermal plots of Van der Waals below the critical temperature mark, rising areas are observed. The points lying in such zones correspond to the unstable state of the substance, which in practice can not be realized.