Question

Gases deviate from the ideal gas behaviour because their molecules (a) possess negligible volume (b) have forces of attraction between them (c) are polyatomic (d) are not attracted to one another

Short answer

The correct option amongst the stated is (b) because real gases do exhibit forces of attraction between them, unlike the assumption in the ideal gas law.

Step-by-step solution

Examine Choice A

Evaluting statement A: Real gases do possess volume, this feature is neglected in the ideal gas law which assumes the gas molecules to be point particles. Therefore, choice (a) is incorrect.

Examine Choice B

Evaluating statement B: Real gases do have intermolecular forces of attraction, opposite to the ideal gas behavior, where the molecules are assumed not to attract or repel each other. Therefore, choice (b) is correct.

Examine Choice C

Analyzing statement C: The number of atoms that forms a molecule (monoatomic, diatomic, or polyatomic), does not cause deviation from ideal behaviour. Therefore, choice (c) is incorrect.

Examine Choice D

Evaluating statement D: Real gases do experience attraction, which is non-existent in the ideal gas behaviour. Therefore, choice (d) is incorrect, as it's opposite of the real gas property.

Key concepts

Ideal Gas Law

The ideal gas law is a fundamental principle in chemistry that describes the behavior of gases under certain conditions. It's encapsulated by the equation $PV=nRT$, where:

• $P$ stands for pressure,
• $V$ is the volume,
• $n$ represents the number of moles,
• $R$ is the ideal gas constant, and
• $T$ is the temperature in Kelvin.

The ideal gas law assumes gases consist of point particles with no volume and experience no intermolecular forces. In reality, however, gases deviate from this behavior due to forces present between particles and their actual volume.
These simplifications allow scientists to predict how a given amount of gas will behave under various temperatures and pressures, providing a useful approximation in many practical scenarios.

Intermolecular Forces

Intermolecular forces are the attractions that exist between molecules. These forces are significant in real gases and are one of the reasons gases deviate from ideal behavior.
There are several types of intermolecular forces, including:

• Dipole-dipole forces: Occur between molecules with permanent dipoles.
• London dispersion forces: Present in all molecules due to temporary dipoles and more noticeable in larger or heavier molecules.
• Hydrogen bonds: A special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like oxygen or nitrogen.

In an ideal gas, these forces are assumed to be zero, meaning there's no attractive or repulsive forces between the particles. However, in the real world, these intermolecular interactions can cause gases to deviate from expected behavior described by the ideal gas equation.

Point Particles Assumption

The point particles assumption is a fundamental component of the ideal gas law. It assumes that gas molecules are point particles, meaning they have negligible volume.
In reality, molecules do have volume and occupy space. This means that when they are compressed, their volume can no longer be ignored.
The assumption works well under low pressure and high temperature conditions. In such environments, the volume of the individual gas molecules is small compared to the total volume of the gas, leading to behavior that closely approximates the ideal gas law.
However, at high pressures or low temperatures, the volume of the gas molecules becomes significant when compared to the total volume of the gas, causing deviations from ideal gas behavior.