Ideal Gases:
* No attraction: In the ideal gas model, gas molecules are assumed to have no attractive forces between them. This is a simplification, but it works well for many gases at relatively low pressures and high temperatures.
Real Gases:
* Weak attractions: Real gases, however, do experience weak intermolecular forces. These forces arise from temporary fluctuations in electron distribution around the molecules, leading to temporary dipoles. These forces are called London dispersion forces and are present in all gases.
* Strength of attraction: The strength of these forces depends on factors like:
* Molecular size: Larger molecules have more electrons and stronger London dispersion forces.
* Polarity: Polar molecules have permanent dipoles and experience dipole-dipole interactions, which are stronger than London dispersion forces.
* Temperature and pressure: At lower temperatures and higher pressures, molecules are closer together and the intermolecular forces become more significant.
Examples:
* Noble gases: Helium, neon, and argon are examples of gases with very weak intermolecular forces. They behave almost like ideal gases at room temperature and pressure.
* Diatomic gases: Nitrogen, oxygen, and hydrogen are also relatively nonpolar and have weak intermolecular forces.
* Polar gases: Water vapor (H2O) and ammonia (NH3) have stronger intermolecular forces due to their polarity. They deviate more from ideal gas behavior.
Conclusion:
While gas molecules are not "attracted" in the same way as solids or liquids, they do experience weak attractive forces due to temporary fluctuations in their electron distribution. These forces become more significant at lower temperatures and higher pressures, causing real gases to deviate from ideal gas behavior.
