Usually, the word ‘real gas’ refers to a gas that does not function as an ideal gas. The interactions between gaseous molecules can explain their behaviour. Such intermolecular interactions between gas particles are the explanation of why the ideal gas law does not adhere to real gases. A real gas can therefore be characterized as a non-ideal gas whose molecules occupy a given amount of space and are capable of interacting with one another. In this article, we will study the real gas definition, real gas equation, and ideal and real gases in detail.
Real Gas Definition
A real gas is defined as a gas that at all standard pressure and temperature conditions does not obey gas laws. It deviates from its ideal behaviour as the gas becomes huge and voluminous. True gases have velocity, mass, and volume. They liquefy when cooled to their boiling point. The space filled by gas is not small when compared to the total volume of gas.
Ideal and Real Gas Equation
An ideal gas is defined as a gas that obeys gas laws at all pressure and temperature conditions. Ideal gases have velocity as well as mass. They have no volume. The volume taken up by the gas is small as compared to the overall volume of the gas. It does not condense and triple-point does not exist.
The ideal gas law is the equation of the state of a hypothetical ideal gas, also called the general gas equation. Under many conditions, it is a reasonable approximation of the behaviour of several gases, but it has many limitations. In 1834, Benoît Paul Émile Clapeyron first described it as a variation of the empirical law of Boyle, the law of Charles, the law of Avogadro, and the law of Gay-Lussac. In an empirical form, the ideal gas law is also written:
pV=nRT
Real Gas Law
By explicitly including the effects of molecular size and intermolecular forces, the Dutch physicist Johannes van der Waals modified the ideal gas law to explain the behavior of real gases. The Van der Waal real gas equation is given below-
Real gas law equation,
= [frac {(P+an^2)} {V^2} = (V-nb) nRT]
Where a and b represent the empirical constant which is unique for each gas.
[frac {n^2} {V^2}] represents the concentration of gas.
P represents pressure
R represents a universal gas constant and T is the temperature
Ideal and Real Gases
The difference below shows the properties of real gas and ideal gas, and also the ideal and real gas behaviour.
|
Ideal Gas |
Real Gas |
|
No definite volume |
Definite volume |
|
Elastic Collision of particles |
Non-elastic collisions between particles |
|
No intermolecular attraction force |
Intermolecular attraction force |
|
Does not really exist in the environment and is a hypothetical gas |
It really exists in the environment |
|
High pressure |
The pressure is less when compared to Ideal gas |
|
Independent |
Interacts with others |
|
Obeys PV = NRT |
Obeys [P + ( frac {(n^2a)} {V^2}) (V-nb) = nRT] |
Did You Know?
A factor known as compressibility factor Z is determined by the deviation of real gas from ideal gas and is defined as the ratio of the actual volume to the volume predicted by the ideal gas law at the same temperature and pressure Z = Actual volume/volume predicted by the ideal gas = v/RT/P
But the ideal gas rate, Videal, is RT/P. The compressibility factor can therefore also be defined as the ratio of specific real gas volume to specific ideal gas volume, i.e.
Compressibility factor Z= [frac {V_{real gas}} {V_{ideal gas}}]
As we all know, at very low pressures and high temperatures, all gases act as ideal gases. So when the pressures are reduced, as the gas behaves as ideal, the value of Z tends to unite.
It is to be remembered that, depending on the pressure and temperature, the value of Z can be less than unity or greater than unity. The compressibility factor chart shows the Z values corresponding to the pressure.
Liquefaction of Gases
The kinetic molecular theory of gases does neither predict nor explain the liquefaction of gases. According to both theory and the ideal gas law, gases crushed to extremely high pressures and chilled to extremely low temperatures should still behave like gases, albeit cold, dense ones. When gases are compressed and cooled, they invariably condense to become liquids, although light elements like helium require extremely low temperatures to liquefy (for He, 4.2 K at 1 atm pressure).
Liquefaction can be thought of as an extreme deviation from ideal gas behavior. When the molecules in a gas are cooled to the point that their kinetic energy is no longer adequate to resist intermolecular attraction forces, this phenomenon happens. The exact temperature and pressure combination required to liquefy a gas is highly dependent on its molar mass and structure, with heavier and more complicated molecules liquefying at higher temperatures. Because large coefficients suggest relatively strong intermolecular attractive interactions, substances with large van der Waals coefficients are generally easy to liquefy. Small molecules containing only light components, on the other hand, have low coefficients, indicating weak intermolecular interactions and making them difficult to liquefy. On a large scale, gas liquefaction is used to separate O2, N2, Ar, Ne, Kr, and Xe. After liquefying a sample of air, the mixture is warmed, and the gases are separated according to their properties.