---
title: "Van der Waals Forces | Thermodynamics II"
description: "Van der Waals forces are weak intermolecular attractions from temporary dipoles, and in Thermodynamics II they help explain real gas behavior and mixtures."
canonical: "https://fiveable.me/thermodynamics-ii/key-terms/van-der-waals-forces"
type: "key-term"
subject: "Thermodynamics II"
unit: "Unit 8"
---

# Van der Waals Forces | Thermodynamics II

## Definition

Van der Waals forces are weak attractions between molecules caused by temporary or induced dipoles. In Thermodynamics II, they show up when real gases deviate from ideal gas behavior, especially in mixtures at higher pressure and lower temperature.

## What It Is

Van der Waals forces are the weak intermolecular attractions that make gas molecules behave less perfectly than the ideal gas law predicts. In Thermodynamics II, the term usually points to the small attractive forces that appear when molecules get close enough for their electron clouds to interact, even if they are not forming chemical bonds.

These forces come from shifting electron density. A molecule may have a momentary uneven charge distribution, which creates a temporary dipole. That dipole can induce a dipole in a nearby molecule, and the result is a weak attraction. In gas mixtures, those attractions are constantly forming and disappearing as particles move around.

This is why the effect becomes more noticeable at lower temperatures and higher pressures. Lower temperature means molecules move more slowly, so attractions have a better chance to matter. Higher pressure packs molecules closer together, which increases the number of interactions between neighboring particles.

For an ideal gas, we assume molecules do not attract or repel each other. Van der Waals forces are one of the main reasons real gases break that assumption. The gas may exert less pressure than the ideal model predicts because nearby molecules are being pulled inward a bit instead of bouncing away cleanly.

In Thermodynamics II, you usually meet this idea through real gas behavior, mixture calculations, or a discussion of why Dalton's law is an approximation when intermolecular forces are not negligible. The key move is recognizing that the gas is not just a collection of independent particles, it is a collection of particles that still feel each other at short range.

## Why It Matters

Van der Waals forces matter because Thermodynamics II is full of situations where the ideal gas model is good enough for a first pass, but not good enough for a real engineering answer. Once pressure rises, temperature drops, or the gas mixture contains molecules that attract each other more strongly, those tiny forces start changing predicted pressure, volume, and phase behavior.

That shows up in real gas calculations, compressibility discussions, and mixture analysis. If you ignore intermolecular attractions, you can overestimate how closely a gas follows ideal behavior. In a problem set, that might mean the ideal gas law gives a clean number, but the real gas model or compressibility factor tells you the actual system is less compressible than expected.

The concept also connects to mixture behavior. In gas mixtures, each component is not just contributing its own partial pressure, it is also interacting with neighboring molecules. That matters when you are comparing a simple Dalton's law setup to a real industrial gas stream, like compressed air, natural gas, or a vapor mixture near condensation.

If you are heading toward phase equilibrium, separation processes, or combustion analysis, van der Waals forces give you the physical reason why gases stop acting ideally. They are one of the first signs that molecular structure and intermolecular attraction can affect engineering calculations.

## Connections

### Ideal Gas Law

The ideal gas law assumes molecules do not attract each other, so it gives a clean baseline for pressure, volume, temperature, and moles. Van der Waals forces explain where that model starts to miss reality. When the forces matter, the gas can have a lower observed pressure or different compressibility than the ideal equation predicts.

### Compressibility Factor

The compressibility factor is one way Thermodynamics II measures how far a real gas is from ideal behavior. Van der Waals forces are a major reason that factor shifts away from 1. If attractions dominate, the gas can be more compressible than expected, especially at moderate to high pressure.

### Dipole-Dipole Interactions

Dipole-dipole interactions are a stronger, more directional type of intermolecular attraction between polar molecules. Van der Waals forces is a broader umbrella in many thermo settings, and temporary dipoles can exist even in nonpolar gases. That is why nonpolar species can still show real-gas behavior.

### London Dispersion Forces

London dispersion forces are the temporary dipole attractions that make up a big part of van der Waals behavior. If your gas molecules are nonpolar, dispersion is often the main intermolecular force you need to think about. Larger or more polarizable molecules usually show stronger dispersion effects.

## On the AP Exam

A quiz or problem-set question will usually give you a gas, a pressure, a temperature, and a clue that the ideal gas model is not perfect. Your job is to decide whether intermolecular attractions matter and explain the direction of the effect. If conditions are high pressure and low temperature, van der Waals forces are more likely to cause noticeable deviation from ideal behavior.

In a mixture problem, you may use the idea to justify why Dalton's law is only approximate for real gases or why a compressibility correction is needed. In a written response, you should connect the force to a physical effect, not just repeat the definition. Say that attractive interactions pull molecules together, reduce the ideal pressure, and become more noticeable when molecules are crowded or moving slowly.

## Key Takeaways

- Van der Waals forces are weak intermolecular attractions, not chemical bonds, so they affect physical behavior more than molecular identity.
- In Thermodynamics II, they matter because real gases do not always follow the ideal gas law, especially at high pressure and low temperature.
- These forces come from temporary or induced dipoles, which can exist even in nonpolar molecules.
- When attractions become noticeable, the measured pressure can be lower than the ideal model predicts because molecules are pulled toward each other.
- In gas mixtures, van der Waals forces help explain why mixture properties can shift away from simple Dalton's law behavior.

## FAQs

### What is van der Waals forces in Thermodynamics II?

Van der Waals forces are weak intermolecular attractions that make real gases deviate from ideal behavior. In Thermodynamics II, they show up when you analyze gas mixtures, compressibility, or phase behavior. They matter most when molecules are close together and moving slowly enough for attractions to affect the calculation.

### Are van der Waals forces the same as dipole-dipole interactions?

Not exactly. Dipole-dipole interactions are between molecules with permanent dipoles, while van der Waals forces is often used more broadly in thermo to include weak attractions like temporary dipoles and dispersion. The common idea is the same, though, because both involve intermolecular attraction that changes real-gas behavior.

### Why do van der Waals forces matter more at high pressure?

At high pressure, molecules are packed closer together, so they interact more often. That gives attractive forces more chances to affect pressure and volume. If you are solving a real-gas problem, high pressure is a clue that the ideal gas law may need a correction.

### How do van der Waals forces affect gas mixtures?

They can make the mixture behave less ideally than Dalton's law assumes. Instead of each gas acting like it is completely independent, the molecules attract one another and shift the observed pressure or compressibility. That is why real mixture problems sometimes need more than just mole fractions and partial pressures.

## Related Study Guides

- [8.1 Properties of Gas Mixtures and Dalton's Law](/thermodynamics-ii/unit-8/properties-gas-mixtures-daltons-law/study-guide/fe7FxGg6EsNKg8Vl)

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