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Chapter 8: Problem 1

In what ways do real gases differ from ideal gases?

Short Answer

Expert verified
Real gases have intermolecular forces and finite volumes, causing deviations from the Ideal Gas Law, especially under high pressure and low temperature.

Step by step solution

01

Introduction to Real and Ideal Gases

Understand that real gases diverge from the concept of ideal gases, which are hypothetical and obey the Ideal Gas Law perfectly: PV = nRT.
02

Intermolecular Forces

Real gases experience intermolecular forces (attractive and repulsive) which are not considered in ideal gases. This causes deviations in their behavior, especially at high pressures and low temperatures.
03

Volume of Gas Particles

In an ideal gas, the volume of individual gas particles is considered negligible. However, in real gases, particles occupy a finite volume which affects the behavior of the gas.
04

Deviations at High Pressure

At high pressures, the volume of the gas particles becomes significant compared to the total volume, leading to deviations from the Ideal Gas Law.
05

Deviations at Low Temperature

At low temperatures, intermolecular attractions cause gases to condense to liquids, thus they do not behave ideally.
06

Corrections to Ideal Gas Law

The Van der Waals equation introduces correction factors for pressure and volume to account for the deviations: del Real gases do not follow the Ideal Gas Law exactly; they show deviations due to intermolecular forces and finite volume of molecules. These deviations are particularly evident under high pressure and low temperature conditions.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Ideal Gas Law
The Ideal Gas Law is a mathematical relationship that explains how gases ideally behave. The formula is expressed as \( PV = nRT \). Here, \( P \) stands for pressure, \( V \) for volume, \( n \) for the number of moles of the gas, \( R \) for the ideal gas constant, and \( T \) for temperature in Kelvin. This law assumes that gas particles do not interact with each other and that the volume of gas particles themselves is negligible.

These assumptions, however, often do not hold true for real gases due to the presence of intermolecular forces and the actual volume of gas particles. Real gases deviate from the Ideal Gas Law under conditions of high pressure and low temperature.
Intermolecular Forces
Intermolecular forces refer to the attractive or repulsive forces between molecules. These forces are not considered in the Ideal Gas Law.

There are two main types of intermolecular forces:
  • Attractive Forces: These pull molecules towards each other, making the gas less expansive.
  • Repulsive Forces: These push molecules apart, affecting how close they can get to each other.

These forces become particularly significant under conditions of low temperature and high pressure.

At low temperatures, the attractive forces can cause gas molecules to come together and condense into a liquid.

In high-pressure situations, the repulsive forces cause the gas to occupy more volume than predicted by the Ideal Gas Law.
Van der Waals Equation
To address the limitations of the Ideal Gas Law, the Van der Waals equation introduces correction factors. The equation is:
\[ \bigg(P + \frac{an^2}{V^2}\bigg)(V - nb) = nRT \]
Here,
  • \

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