Short Answer

Gases behave ideally when they are kept at low pressure and high temperature. Under these conditions, gas particles remain far apart and intermolecular forces become extremely weak. As a result, gases follow the ideal gas equation PV = nRT closely.

In this state, the volume of gas molecules becomes negligible and their collisions follow the assumptions of the kinetic molecular theory. Many common gases, such as oxygen, nitrogen, and hydrogen, behave almost ideally under normal room conditions.

Detailed Explanation

Conditions Under Which Gases Behave Ideally

Ideal gas behaviour is based on two assumptions: gas molecules have negligible volume and no intermolecular forces. In reality, no gas completely satisfies these assumptions, but gases can behave almost ideally under certain conditions. Ideal behaviour means the gas follows the equation PV = nRT accurately and does not require corrections like the van der Waals equation.

The conditions that make gases behave ideally are low pressure, high temperature, and sometimes small molecular size. These factors help reduce the influence of molecular attractions and volume, making the gas resemble the ideal gas model more closely.

Why Low Pressure Favors Ideal Behaviour

At low pressure, gas molecules are spread far apart. This creates a large amount of empty space between particles. As a result:

• Intermolecular attractions become extremely weak.
• Collisions between particles occur less frequently.
• The volume occupied by gas molecules becomes almost negligible.

Because the particles are far apart, their actual volume does not interfere with gas behaviour. This condition closely matches the assumption that ideal gas particles have zero volume.

Low pressure prevents molecules from coming close enough for attractive or repulsive forces to act strongly. This reduces deviation from ideal behaviour.

Why High Temperature Favors Ideal Behaviour

At high temperature, gas molecules move very fast due to high kinetic energy. This rapid motion has important effects:

• Fast-moving particles easily overcome intermolecular forces.
• Attractive forces become insignificant compared to kinetic energy.
• Collisions with container walls are strong and follow ideal gas predictions.

High temperature supports the assumption that molecules move in random motion without being influenced by attraction. The high speed keeps them apart and prevents condensation.

Thus, high temperature makes real gases resemble ideal gases.

Combined Effect of Low Pressure and High Temperature

When both conditions occur together:

• Gas molecules have large distances between them.
• Intermolecular forces become almost zero.
• Volume of molecules becomes negligible.
• Pressure and temperature follow the ideal gas equation closely.

This combination produces the most ideal behaviour among real gases.

Role of Molecular Size

Small gas molecules behave more ideally than large ones.
Examples of gases that behave very ideally:

• Helium (He)
• Hydrogen (H₂)
• Neon (Ne)

These gases have:

• Weak intermolecular forces
• Very small molecular volumes
• High mobility at ordinary temperatures

Large molecules with strong intermolecular forces (NH₃, CO₂) behave less ideally unless kept under special conditions.

Conditions Where Gases Deviate from Ideal Behaviour

To understand ideal behaviour better, it is helpful to know when gases do not behave ideally:

• At high pressure, molecules are packed closely and volume cannot be ignored.
• At low temperature, molecules slow down and attractions become strong.
• Near liquefaction, gases deviate significantly.

These are the exact opposite conditions of ideal behaviour.

Examples of Ideal Behaviour

1. Nitrogen and oxygen at room temperature:
   These gases behave nearly ideally because the temperature is moderate and pressure is low.
2. Helium balloons:
   Helium behaves almost ideally because the atoms are small and attractive forces are extremely weak.
3. Air in open environments:
   At atmospheric pressure and normal temperature, air behaves very close to an ideal gas.
4. Laboratory gas experiments:
   Most gas laws studied in schools assume ideal behaviour because deviations are small under normal conditions.

Importance of Knowing Ideal Gas Conditions

Understanding when gases behave ideally is important in chemistry and industry:

• It makes calculations using PV = nRT simpler and accurate.
• It helps predict gas behaviour in engines, refrigeration, and gas storage.
• It explains why gases resist liquefaction at high temperature.
• It provides the basis for understanding deviations through advanced models.

Knowing ideal conditions also helps chemists decide when to use the ideal gas equation and when to switch to real gas equations like van der Waals.

Conclusion

Gases behave ideally mainly under low pressure and high temperature, where intermolecular forces are minimal and molecular volume becomes negligible. Under these conditions, gas particles move freely, remain far apart, and follow the ideal gas equation closely. Many common gases behave almost ideally at ordinary conditions, making the ideal gas model a useful and accurate tool for studying gas behaviour.