A real gas is one that exists in nature with different chemical structures and that does not show idealized behavior. They can be diatomic molecules such as oxygen, nitrogen, etc., as well as monatomic molecules, including helium, neon, and others. There may even be heavier gases, such as carbon dioxide, methane, and ammonia.
Ideal gases is a model created that guides us to understand the behavior of gases under various environmental conditions. The so-called Ideal Gas Law was established by Benoit Paul Émile Clapeyron in 1834, enunciated in the following formula: PV = nRT.
The law is based on a set of assumptions, including: assuming that the molecules of a gas have no dimensions, that is, that they are punctual and that there are no attractive forces between these molecules.
Real gases do not meet these assumptions. Under certain conditions, such as high pressures and low temperatures, they escape from the ideal gas compartment by increasing intermolecular forces. Likewise, the proportion of the volume of the molecules in relation to the total space occupied by the gas increases.
Characteristics of real gases
Existence of intermolecular forces
There is an attractive force between the molecules of a gas that tends to bring them together, restricting their mobility. These intermolecular forces are known as van der Waals forces, in honor of the Dutch scientist Johannes van der Waals (1837-1923).
These intermolecular forces are the dipole-dipole interaction and the London dispersive forces. Likewise, van der Waals in 1873 introduced the effect of intermolecular forces in the equations of state of a gas.
When considering such interactions, there is a significant deviation in the behavior of real gases in relation to ideal gases; especially at high pressures and a reduction in the volume of the gas, which produces a greater interaction between the gaseous molecules.
Volume of molecules
Among the characteristics attributed to ideal gases are that their molecules are considered as point; and therefore, the volume they occupy in relation to the total space of the gas is negligible.
However, the volume occupied by the molecules of a real gas can be important when the gas is subjected to a pressure that produces a reduction in its volume, increasing the proportion that the volume of the gas molecules occupies in relation to the total space occupied. for gas.
This situation increases the magnitude of the intermolecular forces in the gas as its molecules approach, which brings about some changes in the properties of the gas. For example, there is a decrease in the theoretical pressure of the gas exerted on the walls of the container that contains it.
This is because the frequency of collisions of gas molecules against the walls of the container decreases. Meanwhile, collisions between the same molecules increase, so their mobility is decreased.
Van der Waals equation
Real gases can approach compliance with the Ideal Gas Law (PV = nRT) under specific conditions. But not under all conditions, producing the need to modify the established law.
Several authors contributed to a modification that could be adapted to real gases. Among these contributions is the van der Waals Equation:
(P + an 2 / V 2 ) (V-nb) = nRT
The expression (an 2 / V 2 ) is a correction for the decrease in pressure exerted by the gas as a result of the interaction between the gas molecules. The term ‘a’ is an empirical constant that is specific to each gas and has as a unit L 2 · atm · mol -2 .
The expression (V-nb) corrects the effect of ignoring the volume occupied by the molecules of a gas on the properties of a real gas. The term ‘b’ is obtained empirically and has as a unit: L · mol -1 , whose value varies for each gas. Also, b represents the volume occupied by the gas molecules.
When a real gas is forced through a valve, there is a reduction in its volume; but when leaving it it expands, which produces a decrease in the temperature of the gas. This feature has found application in refrigeration.
Compression factor (z) or compressibility of a gas
The compression factor (PV / nRT) is a ratio that in ideal gases has a constant value of 1, regardless of the pressure or temperature to which they are subjected.
On the contrary, real gases, such as: hydrogen (H 2 ), nitrogen (N 2 ), oxygen (O 2 ) and carbon dioxide (CO 2 ), present a value for the compression factor greater than 1 when the pressure exerted on them is greater than 400 atm.
However, carbon dioxide and oxygen can have a compression factor value of less than 1 for a lower pressure of less than 400 atmospheres. In conclusion: the compression factor is not constant in real gases.
Ideal gases, when subjected to a process of adiabatic compression and expansion, decrease their temperature and increase their density . But without a phase change occurring. In contrast, real gases do undergo a phase change: they liquefy, they go into the liquid phase.
Application of the van der Waals equation
Calculate the pressure exerted by a mole of methane gas (real gas) in a 0.5 L container at 25 ºC.
a) By applying the ideal gas equation.
b) By applying the van der Waal equation with a value for the constant ‘a’ of 2.25 L 2 · atm · mol -2 and 0.0428 for the constant ‘b’.
In part a)
PV = nRT
P = nRT / V
= (1 mol) (0.082 L atm mol -1 K -1 ) (298 K) / (0.50 L)
= 48.87 atm
And in subsection b)
(P + an 2 / V 2 ) (V-nb) = nRT
a = 225 L 2 atm mol -2
b = 0.0428 L mol -1
[P + (1 mol) 2 (2.25 L 2 · atm · mol -2 /(0.5 L) 2 )] [(0.500 L – 0.0428 L)] = (1 mol) (0.082 L · atm · mol -1 ) (298K)
(P + 9 atm) (0.4572 L) = 24.36 atm L
P = 44.28 atm
A decrease in the pressure exerted by the real gas is observed when the van der Waals equation is used instead of the ideal gas equation. This is a consequence of the existence of intermolecular forces and the volume of gas molecules.
Examples of real gases
All gases that exist in nature are real, including gases with diatomic molecules, such as oxygen, nitrogen, hydrogen, chlorine, fluorine, bromine, and iodine; and monatomic gases, such as helium, argon, krypton, neon, and radon.
In addition to chemical compounds in a gaseous state such as butane, carbon dioxide, sulfur dioxide, methane, among others.