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📘 Lesson 6 of 10 · Thermodynamics

⚛️ Kinetic Theory

Kinetic theory explains the behavior of gases by imagining them as vast numbers of tiny particles in constant, random motion — connecting microscopic particle behavior to the macroscopic properties we measure.

Course progress: 60%

01 Key Concepts

Core Assumptions

Gas particles are in constant, random motion; collisions between particles (and with container walls) are perfectly elastic; particles have negligible volume compared to the container; there are no attractive forces between particles (for an ideal gas).

Pressure from a Molecular View

Gas pressure results from countless particle collisions against the container walls -- more frequent or more forceful collisions mean higher pressure.

Temperature and Average Kinetic Energy

Average kinetic energy of gas particles is directly proportional to absolute (Kelvin) temperature: KE_avg = (3/2)*k*T, where k is Boltzmann's constant.

Root-Mean-Square Speed

A statistical measure of the typical speed of gas particles, accounting for the wide range of individual particle speeds in a gas sample.

Real Gases vs. Ideal Gases

Real gases deviate from ideal behavior at high pressure or low temperature, where particle volume and intermolecular forces become significant.

02 Key Formulas

03 Solved Examples

Example 1 If the temperature of a gas doubles (in Kelvin), what happens to the average kinetic energy of its particles?
  1. Average KE is directly proportional to Kelvin temperature.
  2. Doubling temperature doubles average KE.
Answer: It also doubles
Example 2 Explain, using kinetic theory, why heating a sealed gas container increases its pressure.
  1. Heating increases the average kinetic energy (and speed) of the gas particles.
  2. Faster particles collide with the container walls more frequently and with greater force.
  3. More frequent, forceful collisions produce higher pressure.
Answer: Higher temperature increases particle speed, leading to more frequent and forceful wall collisions, raising pressure
Example 3 Why do real gases deviate from ideal gas behavior at very high pressure?
  1. At high pressure, gas particles are squeezed much closer together.
  2. The assumption of negligible particle volume (used for ideal gases) becomes less accurate as particles take up a more significant fraction of the total volume.
Answer: The particles' own volume becomes significant, violating one of kinetic theory's core assumptions for an ideal gas

04 Practice Questions

1If temperature triples, what happens to average KE of gas particles?
It also triples
2What causes gas pressure, according to kinetic theory?
Particle collisions with the container walls
3What assumption is made about collisions between gas particles in kinetic theory?
They are perfectly elastic (no kinetic energy lost overall)
4Under what condition might real gases deviate most from ideal gas predictions?
At high pressure or low temperature
5What does 'root-mean-square speed' represent?
A statistical measure of the typical particle speed in a gas sample

📄 Kinetic Theory — Downloadable Worksheet

10 questions with a full answer key. Grab the PDF to print, or try the interactive version in your browser.

🗒 Cheat Sheet ⬇ Download PDF 🖥 Interactive Worksheet