Why do real gases deviate from ideal behavior?

Short Answer

Real gases deviate from ideal behavior because the assumptions made in the ideal gas equation are not completely true for actual gases. Real gas particles have some intermolecular forces between them and occupy a definite volume, while ideal gases assume no attraction and zero particle volume.

At high pressures and low temperatures, these effects become more noticeable. Gas particles come closer, attraction increases, and their actual volume affects their behavior, causing real gases to deviate from the predictions of the ideal gas equation.

Detailed Explanation

Reasons Real Gases Deviate from Ideal Behavior

Real gases deviate from ideal behavior because real gas particles do not follow the assumptions of the ideal gas model perfectly. The ideal gas equation PV = nRT assumes that gas particles have no volume and no intermolecular forces, and that all collisions are perfectly elastic. These assumptions make calculations simple, but they are not completely accurate for gases in the real world. Real gas molecules do have size and do experience attractions and repulsions. These factors become especially important under certain conditions, causing real gases to behave differently from ideal gases.

Understanding these deviations is important because real gases are used in industries, laboratories, and natural processes. Knowing when and why gases deviate helps scientists and engineers make accurate predictions and design safe systems involving gases.

Finite Volume of Gas Particles

One important reason for deviation is that real gas molecules occupy space. In an ideal gas, the particles are considered point-sized, meaning their volume is assumed to be zero. But in reality, every gas molecule has a definite size.

At low pressure, this volume is small compared to the total container volume, so deviation is minimal.
At high pressure, the gas is compressed and molecules are very close. Here, the space occupied by molecules becomes significant, so the actual volume of the gas is greater than the volume predicted by the ideal gas equation.

Intermolecular Forces of Attraction

Ideal gases assume that particles do not attract or repel each other. But real gas particles experience intermolecular forces. These forces are weak at ordinary conditions, but they become important under low temperature and high pressure.

  • At low temperature, particles move slowly and come closer, increasing attractions.
  • At high pressure, particles are forced close, making attractive forces stronger.

Due to these attractions, particles hit the walls with less force, causing the pressure of a real gas to be lower than predicted by ideal behavior.

Deviation at High Pressure

When a gas is highly compressed, the molecules are extremely close. Due to this:

  1. Volume of molecules cannot be ignored.
    The free space available becomes smaller than predicted.
  2. Repulsive forces increase.
    When molecules are forced too close, they start repelling each other.
  3. Pressure becomes higher than expected because the repulsions push back strongly.

This makes real gases deviate significantly at high pressure.

Deviation at Low Temperature

At low temperature:

  1. Molecules lose kinetic energy and move slowly.
  2. Attractive forces become stronger, pulling molecules closer.
  3. Gas pressure becomes lower than ideal because weak collisions occur with container walls.

This causes real gases to condense into liquids at very low temperatures, which an ideal gas would never do.

Ideal Gas Assumptions vs Real Gas Behavior

Ideal gas assumptions break down in real gases:

Ideal Gas Assumption Real Gas Behavior
Zero volume of particles Particles occupy finite volume
No intermolecular forces Real gases have attractive and repulsive forces
Perfectly elastic collisions Minor energy changes may occur
Particles far apart always Only true at low pressure and high temperature

Because these assumptions do not match real conditions completely, deviations occur.

Van der Waals Equation Correction

To correct these deviations, the van der Waals equation is used:

Where:

  • a = correction for intermolecular attraction
  • b = correction for finite molecular volume

This equation predicts real gas behavior more accurately.

Conditions When Real Gases Behave Nearly Ideally

Real gases behave almost like ideal gases when:

  • Temperature is high → particles move fast, reducing attractions
  • Pressure is low → particles are far apart, volume becomes negligible

Examples of gases behaving nearly ideally:

  • Helium
  • Hydrogen
  • Nitrogen
  • Oxygen

Under normal room conditions, these gases show very little deviation.

Examples of Deviations

  1. Carbon dioxide liquefies easily at low temperature due to strong attractions.
  2. Ammonia shows strong deviations because of hydrogen bonding.
  3. LPG is stored under high pressure, showing significant deviations.

These examples demonstrate how real gases behave differently from ideal gases.

Conclusion

Real gases deviate from ideal behavior because their molecules have finite volume and experience intermolecular forces. These factors become significant at high pressure and low temperature, causing deviations from the ideal gas law. While ideal gases are theoretical models, real gases show actual behaviour that must be corrected using equations like the van der Waals equation. Understanding these deviations helps in predicting and working with gases in real-life situations.