Question

Why does high, but not low, pressure cause a deviation from the ideal gas law? A. Higher pressure decreases the interatomic distance to the point where intermolecular forces reduce the volume below that predicted by the ideal gas equation. B. Low pressure increases the atomic radius of a gas making it more stable whereas high pressure compresses the gas particles decreasing their stability. C. Low pressure does cause a significant deviation from the ideal gas law because the increased interatomic distance means that no particles ever collide. D. Low pressure does cause a significant deviation because a low pressure implies a reduction in temperature via Charles' law, which increases the power of intermolecular forces.

Short answer

A: Higher pressure decreases interatomic distance causing intermolecular forces to reduce the volume below prediction.

Step-by-step solution

Recall the Ideal Gas Law

The Ideal Gas Law is given by the formula: \[ PV = nRT \] where \(P\) is pressure, \(V\) is volume, \(n\) is the number of moles, \(R\) is the gas constant, and \(T\) is temperature. The law assumes no intermolecular forces and infinitely small gas molecules.

Understand the Effect of High Pressure

At high pressure, the molecules of the gas are forced closer together. This decreases the interatomic distance to the point where intermolecular forces (such as Van der Waals forces) become significant.

Impact of Intermolecular Forces

When these forces become significant, they start to reduce the volume of the gas below that predicted by the Ideal Gas Law. This is because the attractive forces pull the gas molecules closer together than predicted.

Evaluate the Given Options

Option A states that high pressure decreases the interatomic distance causing intermolecular forces to reduce the volume below prediction, which aligns with our understanding.Option B incorrectly claims low pressure increases atomic radius and high pressure decreases stability.Option C incorrectly states low pressure leads to no particle collisions.Option D incorrectly links low pressure with a significant deviation due to temperature reduction, neglecting actual intermolecular forces.

Conclusion

Based on our evaluation, the correct explanation is provided by Option A: Higher pressure decreases the interatomic distance to the point where intermolecular forces reduce the volume below that predicted by the ideal gas equation.

Key concepts

Ideal Gas Law

The Ideal Gas Law is a fundamental principle in chemistry. It is expressed by the equation: Ideal Gas Law states: PV = nRT Here, **P** is the pressure exerted by the gas, **V** is the volume occupied, **n** is the number of moles, **R** is the ideal gas constant, and **T** is the temperature in Kelvin. This equation assumes gases have no intermolecular forces and that molecules occupy zero volume. However, in reality, these assumptions don't always hold true.

Intermolecular Forces

Intermolecular forces are attractions or repulsions that act between neighboring particles (atoms, molecules, or ions). They are much weaker than the strong bonds inside molecules (covalent, ionic, or metallic).
Types of intermolecular forces include:

• Van der Waals Forces: These are the weakest intermolecular forces, and they can be divided into London dispersion forces, dipole-dipole interactions, and hydrogen bonds.
• Dipole-Dipole Interactions: These occur between neutral polar molecules due to the attraction between positive and negative charges.
• Hydrogen Bonds: A special case of dipole-dipole interactions, they occur when hydrogen is bonded to an electronegative atom such as oxygen, nitrogen, or fluorine.

When dealing with gases at low pressure, these forces become negligible. But under high pressure, the gas molecules are compressed together, making these forces significant.

High Pressure Effects on Gases

As pressure on a gas increases, its molecules are forced closer together. This decreases the interatomic distances and makes the assumptions of the Ideal Gas Law less valid.
Important points to consider:

• **Decreased Volume:** The actual volume occupied by gas molecules becomes significant when they are closely packed.
• **Intermolecular Forces Become Significant:** The closer proximity allows intermolecular attractions (or repulsions) to affect gas behavior.

Under high pressure conditions, these effects cause deviations from the Ideal Gas Law, as the real volume and intermolecular attractions reduce the predicted volume or pressure, depending on the context of the problem.

Van der Waals Forces

Van der Waals forces are intermolecular forces that include attractions and repulsions between atoms, molecules, and surfaces, and are important when dealing with gases at high pressure.
Components of Van der Waals Forces:

• **London Dispersion Forces:** Caused by temporary fluctuations in electron density within atoms or molecules, leading to a short-lived dipole that attracts other dipoles.
• **Dipole-Dipole Forces:** Permanent interactions between polar molecules due to attraction between positive and negative ends.
• **Hydrogen Bonds:** Although a type of dipole-dipole interaction, hydrogen bonds are stronger and occur specifically when hydrogen is directly bonded to N, O, or F.

These forces significantly influence a gas's behavior, especially under high pressure scenarios. They cause the gas volume to deviate from what is predicted by the Ideal Gas Law due to intermolecular attractions pulling molecules closer together.