The Ideal Gas Law is an equation of state that provides a good approximation of the behavior of many gases under a variety of conditions. However, it is based on certain assumptions that do not hold true under some conditions. Under these conditions, gases do not behave ideally, and the Ideal Gas Law does not accurately describe their behavior. These conditions are:
1. **High Pressure**: At high pressure, the volume of individual gas molecules is not negligible compared to the total volume of the gas. This means that the assumption that the volume of gas molecules is insignificant is no longer valid. Additionally, the gas molecules are closer together, so the intermolecular forces between them become significant. These forces can cause the gas to deviate from ideal behavior.
2. **Low Temperature**: When the temperature of a gas is low, the kinetic energy of the molecules decreases. As a result, the importance of intermolecular attractions increases since the molecules are moving slower and have more time to interact with one another. At very low temperatures, some gases may even condense into liquids or solids, which is a clear deviation from ideal gas behavior.
Therefore, the conditions where gases are considered real and the Ideal Gas Law does not apply are:
- High pressure
- Low temperature
In contrast, under low pressure and high temperature conditions, gases are more likely to behave ideally because the molecules are far apart (minimizing the effects of their volume and intermolecular forces), and they have enough kinetic energy to overcome intermolecular attractions.
