Deviation From Ideal Gas Behaviour

Ideal gas behaviour is a theory that expects the gases to behave in a certain way and assumes that the gases have negligible or no space at all and they have no intermolecular force of attraction. Deviation of gases from their ideal gas behaviour occurs when the molecules of a gas are cooled down to a point where they no longer possess sufficient kinetic energy to overcome attractive intermolecular forces.

Ideal and Real Gases

Ideal gases are those gases that obey the ideal equation of PV = nRT under all amounts of pressure and temperature. But there is no such gas that behaves the same in every pressure and temperature. Hence, this concept is theoretical. Real gases are those who obey the gas law if the pressure is low or the temperature is high. All gases are real gases.

Difference Between Ideal Gas and Real Gas

The differences between ideal gas and real gas are given below.

Pressure, Volume, and Temperature Relationship in Gases - Why do the Real Gases Deviate?

The graphs below represent different gases and show how they behave under high and low pressure.

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The figure shows how gases behave differently from their ideal behaviour, particularly in high pressure. At low pressure, as shown in figure(b), the real gases behave more like that of the expected ideal behaviour. For gases such as CO2 and C2H4, they deviate more than other real gases because these gases tend to liquefy at lower pressures.

Now, the graph below shows the behaviour of real gas N2 under different temperatures.

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The figure shows that the real gas nitrogen behaves more according to the ideal gas behaviour when the temperature is high. Why do gases deviate so much under high pressure and low temperature? At both the conditions, the basic assumptions that the law of the ideal gas holdsare: the volume of the molecules of the gas are negligible and intermolecular interaction is negligible – these two become invalid.

Under low pressure, the gas molecules are farther apart from each other, and the volume of molecules is the same as the volume of the container. As the pressure increases, the molecular space contracts, and their volume becomes significant as compared to the container. If more pressure is exerted, then the gas liquefies under very high pressure such as CO2.

All the molecules attract each other by a combination of forces. At high temperatures, these have enough energy, and they overcome their attractive force and predominate by the effects of the molecular volume. On the other hand, with the decrease in the temperature, the energy of the molecules also decreases. Eventually, there comes the point where it becomes impossible for the molecules to overcome the force of attraction, and it results in the liquefaction of gas and turns into a liquid state. That is why the ideal gas behaviour is a theoretical concept and does not apply in real situations.

Van der Waals Equation

In 1873, J.D. Van der Waals did some modifications with the ideal law of gas equation to explain the behaviour of real gases, in which he took into account:

• The volume of the gas molecules.

• The forces of attraction between the gas molecules.

He put forward the following equation:

\[(P+\frac{a}{v^2})(v-b)=RT\]

For n moles of the gas,

\[(P+\frac{an^2}{v^2})(v-nb)=nRT\]

The constants ‘a’ and ‘b’ represent the scale of intermolecular attraction and the excluded volume, respectively. The higher the value of 'a', the greater is the molecular attraction and the gas will easily compress. The term 'b' represents the excluded volume that is occupied by gas particles. These constants are different for different gases.

Conclusion

Hence, the article explained the important concept of gases of deviation from ideal gas behaviour. The Van der Waals equation is important for understanding the variation of temperature and pressure on gases. The article will develop a strong understanding of the behaviour of ideal gases.