Problem 89 - Chemistry The Central Science In Si Units

Question

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

Step-by-Step Solution

Answer

Gases deviate from ideal behavior at high pressures and low temperatures due to intermolecular forces and molecular volume.

Step 1: Identify Experimental Conditions

When studying gases, they generally deviate from ideal behavior under conditions of very high pressures and very low temperatures. At high pressures, gas molecules are forced closer together than the ideal gas model assumes, and at low temperatures, gas molecules lose kinetic energy and attract more strongly to one another.

Step 2: Understand the Causes of Deviation

Gases deviate from ideal behavior predominantly due to two main reasons: intermolecular forces and molecular volume. Ideal gas laws assume no intermolecular forces between gas molecules and that the volume of individual gas molecules is negligible. However, at high pressures and low temperatures, these assumptions fail as molecules interact more and their finite volume becomes significant.

Key Concepts

Experimental Conditions

Gases tend to behave ideally under moderate conditions of temperature and pressure. However, this ideal behavior starts to deviate markedly in specific experimental conditions.

One critical condition that causes deviation is **high pressure**. When pressure increases, gas molecules are pushed closer together. This proximity leads to interactions that do not align with the assumptions of ideal gas behavior. As a result, gases act differently than predicted by the ideal gas laws.

Another significant experimental condition is **low temperature**. At lower temperatures, gas molecules move more slowly and have less kinetic energy. This reduced speed allows the attractive forces between molecules to have a more pronounced effect. Consequently, gases tend to deviate more from ideal behavior when they are in cooler surroundings.

Intermolecular Forces

Despite the ideal gas assumption that there are no interactions between molecules, real gases exhibit attractive and repulsive forces. These forces significantly influence gas behavior, especially under non-ideal conditions.

In ideal gas calculations, it is assumed that the potential energy due to intermolecular forces is negligible. However, when molecules are nearby, as in high-pressure situations, attractive forces cause them to gather more closely. This gathering leads to deviations from predicted volumes and pressures based on the ideal gas law.

Attractive forces are also more effective when temperatures drop. Gas molecules slow down, allowing these forces to act over a prolonged period. Thus, intermolecular forces are a primary reason for significant deviations from ideal behavior in gases under real-world conditions.

Molecular Volume

The ideal gas law assumes that gas molecules have an insignificant volume and do not occupy space. However, in reality, each molecule occupies a finite volume, which can become significant under certain conditions.

At high pressures, the volume of individual gas molecules cannot be ignored. As pressure increases, the space available for molecules to move is reduced, causing their actual volume to impact the overall behavior of the gas. This congestion leads to deviations from predictions made using the ideal gas equation.

Moreover, in low-temperature settings, the molecules' available energy to overcome the effects of their own volume decreases. Therefore, the size of the gas molecules plays a more pronounced role. Recognizing the finite volume of molecules is key to understanding why gases deviate from ideal behavior in varying conditions.