1. Why Real Gases Are Not Perfect
The ideal gas model assumes that molecules have no volume and no intermolecular forces. Real gases do not satisfy these assumptions, especially under certain conditions. Their behaviour matches the ideal gas equation only approximately.
2. Assumptions of Ideal Gases (Where Real Gases Differ)
Ideal gases follow two major assumptions that real gases cannot fully satisfy:
2.1. 1. Molecules Have No Volume
In real gases, molecules occupy space. At high pressure, this volume becomes significant compared to container volume.
2.2. 2. No Intermolecular Forces
Real gas molecules attract and repel each other. These forces become important at low temperature and high pressure.
3. Conditions Where Real Gases Behave Like Ideal Gases
Real gases can behave almost ideally when:
- pressure is low
- temperature is high
Under these conditions, molecules are far apart and intermolecular forces are negligible.
3.1. Examples
Gases like nitrogen, oxygen, and helium behave nearly ideally at room temperature and normal atmospheric pressure.
4. Deviations From Ideal Behaviour
Real gases deviate from the ideal gas law due to molecular size and intermolecular forces.
4.1. 1. Effect of High Pressure
At high pressure, molecules come close together. Their finite volume becomes important, causing real gases to occupy more volume than predicted by PV = nRT.
4.2. 2. Effect of Low Temperature
At low temperature, attractive forces slow down molecules and reduce pressure compared to ideal predictions.
4.3. 3. Liquefaction
At very low temperature and sufficiently high pressure, real gases can condense into liquids—something impossible for an ideal gas.
5. van der Waals Correction
The van der Waals equation modifies the ideal gas equation to account for real gas behaviour:
\left(P + \dfrac{a}{V_m^2}\right)(V_m - b) = RT
- a accounts for intermolecular attraction
- b accounts for molecular volume
5.1. Meaning of Corrections
Attraction term: reduces pressure.
Volume term: reduces available volume because molecules occupy space.
6. Graphical View: PV vs P or Z vs P
The compressibility factor Z is used to measure deviation:
Z = \dfrac{PV}{nRT}
For an ideal gas: Z = 1. For real gases, Z can be greater or less than 1 depending on conditions.
6.1. Interpretation
Z > 1: Repulsive forces dominate (high pressure).
Z < 1: Attractive forces dominate (low pressure).
7. Everyday Examples of Real Gas Deviations
- LPG cylinders contain liquefied gas under high pressure where ideal gas law fails.
- CO2 fire extinguishers contain liquid CO2 at low temperature.
- Helium-filled balloons shrink in cold environments due to greater attraction and reduced pressure.