Question

(a) List two experimental conditions under which gases deviate from ideal behavior. (b) List two reasons why the gases deviate from ideal behavior.

Short answer

(a) Two experimental conditions under which gases deviate from ideal behavior are (1) High Pressure and (2) Low Temperature. (b) Two reasons why gases deviate from ideal behavior are (1) Intermolecular forces and (2) Finite volume of gas particles.

Step-by-step solution

(a) Listing experimental conditions for non-ideal behavior

(1) High Pressure: Under high pressure, the volume occupied by the gas becomes comparable to the volume of the gas particles themselves. This causes the ideal gas law (PV=nRT) to give inaccurate predictions for the behavior of the gas. (2) Low Temperature: At low temperatures, intermolecular forces between gas particles become significant. These forces cause the gas particles to experience attractive or repulsive forces, which lead to deviations from the ideal gas law predictions.

(b) Listing reasons for deviation from ideal behavior

(1) Intermolecular forces: The ideal gas law is based on the assumption that gas particles do not interact with each other. However, in reality, gas particles do experience attractive and repulsive forces (e.g., Van der Waals forces) that can affect their behavior and lead to deviations from the ideal gas law predictions. (2) Finite volume of gas particles: Another assumption of the ideal gas law is that gas particles have negligible volume compared to the volume of the container they occupy. However, under conditions such as high pressure or low temperature, the volume occupied by the gas particles can become significant compared to the container volume. This will also lead to deviations from the ideal gas law's predictions.

Key concepts

Non-ideal gas behavior

Gases are often modeled using the ideal gas law, which assumes they behave perfectly. However, in real-world scenarios, gases often deviate from this ideal behavior. These deviations occur due to certain conditions where the assumptions of the ideal gas law no longer hold true. For instance, at high pressures and low temperatures, the behavior of gases starts to differ from what's predicted by the ideal gas equation.
To understand this better, it's crucial to look at the factors like intermolecular forces and the volume occupied by gas particles. These elements become significant when gases are compressed or cooled, thus behaving non-ideally. This knowledge helps in predicting the behavior of real gases more accurately in various practical applications.

Intermolecular forces

Intermolecular forces are interactions that occur between gas molecules. In the context of gases, these forces can be attractive or repulsive, and are usually weak under normal atmospheric conditions. However, under certain conditions such as low temperature, these forces become substantial.
At lower temperatures, gas molecules move slower and spend more time near each other, allowing intermolecular forces to take effect. These interactions cause the gas particles to deviate from the ideal gas behavior, where such forces are assumed to be negligible. Understanding these forces is important as it helps explain why gas behavior changes under different temperature conditions.

Van der Waals forces

Van der Waals forces are a specific type of intermolecular force that contributes to the non-ideal behavior in gases. Named after the Dutch scientist Johannes Diderik van der Waals, these forces account for the attraction and repulsion between gas molecules.
They are essential for adjusting the ideal gas law to better fit real gas behavior, especially under high pressures or low temperatures. Van der Waals corrected the ideal gas law by adding parameters that account for these attraction and volume effects:
- The 'a' factor adjusts for the attraction between particles. - The 'b' factor adjusts for the volume occupied by the particles.
Understanding Van der Waals forces is crucial for accurately describing the behavior of real gases.

Gas particle volume

In the ideal gas law, gas particles are assumed to have no volume, allowing them to move freely without occupying space. This simplification works well under low pressure and high temperature conditions. Yet, in scenarios where pressure increases, the volume of the gas particles themselves becomes significant relative to the container.
This reality leads to non-ideal behavior as the space available for gas movement is less than what the ideal gas law considers. The effect of the gas particles' finite volume is one of the key reasons why real gases deviate from ideal behavior under high-pressure conditions.

High pressure effects

At high pressures, the space between gas particles decreases, making the volume of the particles themselves influential. Under such conditions, the gas molecules are pushed closer together, making the assumption of negligible volume in the ideal gas law invalid.
The interactions between particles become more frequent and significant, leading to deviations from predicted behavior. This also explains why gases do not compress indefinitely under high-pressure conditions. Recognizing the effects of high pressure helps in adjusting our models to predict real gas behavior more accurately.

Low temperature effects

Low temperatures slow down the movement of gas molecules, allowing intermolecular forces to manifest. These forces, which are negligible at higher temperatures, can cause gas molecules to behave less ideally. As temperature decreases, the attraction between molecules becomes prominent, potentially leading to condensation into liquids.
This non-ideal behavior is accounted for in real gas models, like the Van der Waals equation, which considers the reduced kinetic energy and increased influence of intermolecular forces. Understanding low temperature effects is crucial for studying gases under various thermal conditions and predicting their behavior accurately.