Question

For each of the following, which shows the greater deviation from ideal behavior at the same set of conditions? Explain. (a) Argon or xenon (b) Water vapor or neon (c) Mercury vapor or radon (d) Water vapor or methane

Short answer

a) Xenon, b) Water vapor, c) Mercury vapor, d) Water vapor

Step-by-step solution

Understand Ideal Behavior

Ideal behavior refers to a gas behaving according to the Ideal Gas Law, expressed as \( PV = nRT \). Real gases deviate from this behavior due to intermolecular forces and the volume occupied by gas molecules.

Consider Molecular Size and Intermolecular Forces

Larger molecules and those with stronger intermolecular forces generally show greater deviations from ideal behavior because they experience significant attractions or repulsions not accounted for in the Ideal Gas Law.

Evaluate Argon vs. Xenon

Xenon is larger and has stronger London dispersion forces compared to Argon. Thus, Xenon shows a greater deviation from ideal behavior.

Evaluate Water Vapor vs. Neon

Water vapor (H₂O) has hydrogen bonding, which is a strong intermolecular force, whereas Neon is a noble gas with weak London dispersion forces. Therefore, water vapor shows greater deviation from ideal behavior.

Evaluate Mercury Vapor vs. Radon

Mercury vapor is composed of larger and more polarizable atoms compared to the noble gas Radon. Consequently, Mercury vapor shows a greater deviation from ideal behavior.

Evaluate Water Vapor vs. Methane

Water vapor has hydrogen bonding and is more polar compared to Methane, which has weaker intermolecular forces. Hence, water vapor shows greater deviation from ideal behavior.

Key concepts

ideal gas law

The Ideal Gas Law is a fundamental principle in chemistry, represented by the equation: \( PV = nRT \). In this equation, \( P \) stands for pressure, \( V \) for volume, \( n \) for the number of moles of the gas, \( R \) for the universal gas constant, and \( T \) for temperature. This law assumes that gas molecules do not interact with each other and that they occupy no volume.
This means it's incredibly useful for predicting and understanding the behavior of gases under hypothetical 'ideal' conditions.
However, real gases often don't behave ideally.
Understanding where and why these deviations from ideal behavior occur is key to grasping more complex gas laws.

intermolecular forces

Intermolecular forces are the attractions or repulsions that occur between molecules.
These include London dispersion forces, dipole-dipole interactions, and hydrogen bonds.
In the context of gases, these forces become important because they influence how molecules interact with each other.
For instance, gases with strong intermolecular forces, like hydrogen bonds in water vapor, exhibit significant deviations from the Ideal Gas Law.
In contrast, noble gases like Neon, which mainly exhibit weak London dispersion forces, come closer to ideal gas behavior.

molecular size

The size of gas molecules also plays a crucial role in their deviation from ideal behavior.
Larger molecules occupy more space and can have stronger intermolecular forces.
For example, Xenon is a much larger molecule than Argon and thus has stronger London dispersion forces.
This results in Xenon deviating more from ideal gas behavior compared to Argon.
Similarly, Mercury vapor, composed of larger atomic structures, deviates from ideal behavior more than Radon.

hydrogen bonding

Hydrogen bonding is a specific type of strong intermolecular force that occurs when hydrogen is bonded to highly electronegative atoms like oxygen, nitrogen, or fluorine.
This leads to a significant attractive force between molecules.
Water vapor (H₂O) is a prime example of a gas with strong hydrogen bonding.
These strong interactions cause water vapor to deviate significantly from ideal gas behavior.
On the other hand, gases like Methane, which lack hydrogen bonding, show much less deviation and behave more ideally.