The term "real gases" among chemists and physicists is used to call such gases, the properties of which most directly depend on their intermolecular interaction. Although in any specialized reference book one can 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, in accordance with the Clapeyron equation, the forces of mutual attraction and repulsion of molecules do not act, and the volume occupied by the latter is a negligible value.
Of course, such substances do not exist in nature, so all these arguments and calculations are purely theoretical. But real gases, which deviate to one degree or another from the laws of ideality, are found all the time. Between the molecules of such substances there are always forces of mutual attraction, which implies that their volume is somewhat different fromderived perfect model. Moreover, all real gases have different degrees of deviation from ideality.
But there is a very clear trend here: the more the boiling point of a substance is close to zero degrees Celsius, the more this compound will differ from the ideal model. The equation of state for a real gas, owned by the Dutch physicist Johannes Diederik van der Waals, was derived by him in 1873. This formula, which has the form (p + n2a/V2) (V – nb)=nRT, has been compared with the Clapeyron equation (pV=nRT), determined experimentally. The first of these takes into account the forces of molecular interaction, which are influenced not only by the type of gas, but also by its volume, density and pressure. The second amendment determines the molecular weight of a substance.
These corrections acquire the most significant role at high gas pressure. For example, for nitrogen at an indicator of 80 atm. calculations will differ from ideal by about five percent, and with an increase in pressure to four hundred atmospheres, the difference will already 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 observed for all real gaseous substances very approximately. Qualitative deviations are much deeper.
Real gases may well be converted andinto a liquid, and into a solid state of aggregation, which would be impossible if they strictly followed the Clapeyron equation. Intermolecular forces acting on such substances lead to the formation of various chemical compounds. Again, this is not possible in a theoretical ideal gas system. The bonds formed in this way are called chemical or valence bonds. In the case when a real gas is ionized, the Coulomb attraction forces begin to appear in it, which determine the behavior, for example, of a plasma, which is a quasi-neutral ionized substance. This is especially relevant in 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 radio wave signal propagation, and the problem of controlled nuclear and thermonuclear reactions.
Chemical bonds in real gases by their nature practically do not differ from molecular forces. Both those and others, by and large, are reduced to the electrical interaction between elementary charges, from which the entire atomic and molecular structure of matter is built. However, a full 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. This also requires the factor of their thermodynamic stability. One of the important conditions for such stability of a substance is that inIn the isothermal pressure equation, a tendency to a decrease in the total volume of the body must be strictly observed. In other words, as the value of V increases, all isotherms of the real gas must steadily fall. Meanwhile, on the isothermal van der Waals plots below the critical temperature mark, rising sections are observed. Points lying in such zones correspond to an unstable state of matter, which cannot be realized in practice.
