What you'll be able to do: State the postulates of the kinetic molecular theory and use them to explain gas pressure and gas behaviour.
Gases look simple from the outside: they fill their container and push on the walls. The kinetic molecular theory explains both facts with a short list of assumptions about the particles.
These are recommended, not required. You can start this lesson at any time.
A gas consists of a very large number of particles whose own volume is negligible compared with the volume of the container. Those particles are in constant, random, straight-line motion. Collisions between particles, and with the walls, are elastic, so no kinetic energy is lost overall. There are no attractive or repulsive forces between the particles. Finally, the average kinetic energy of the particles is proportional to the absolute temperature in kelvin.
Each time a particle strikes a wall it reverses direction, which means the wall exerted a force on it and it exerted an equal force on the wall. Pressure is the sum of those tiny forces divided by the wall area. Anything that raises the collision rate or the force per collision raises the pressure.
The average translational kinetic energy of a gas particle is (3/2)kT, and for one mole it is (3/2)RT. Because this depends only on temperature, two different gases at the same temperature have the same average kinetic energy, even though the heavier gas moves more slowly.
KEavg = (3/2)RT per mole
Shrink the volume at fixed temperature and the particles hit the walls more often, so pressure rises. Raise the temperature at fixed volume and the particles hit harder and more often, so pressure rises. Add more particles and there are simply more collisions. Every gas law you will meet is this model in algebraic form.
The particles themselves take up almost none of the container volume, which is why gases compress so easily.
Ideal particles neither attract nor repel, so they travel in straight lines between collisions.
A collision in which total kinetic energy is conserved, so the gas does not slowly grind to a halt.
Average kinetic energy depends on absolute temperature only, never on the identity of the gas.
KEavg = (3/2)RT
A sealed rigid flask of helium is heated from 300 K to 600 K. Explain what happens to the pressure and to the average kinetic energy.
Try it first: Decide which quantities are fixed by the word rigid and by the word sealed.
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Helium and argon are in separate flasks at 25 °C. Which gas has the greater average kinetic energy, and which has the greater average speed?
Try it first: Write down what average kinetic energy depends on before you think about mass.
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Why it's wrong: Average kinetic energy is set by temperature alone, not by molar mass.
Check instead: Compare temperatures first, then use mass only to compare speeds.
Why it's wrong: Only the kelvin scale is proportional to particle energy, so Celsius ratios are meaningless.
Check instead: Convert to kelvin before any proportional reasoning.
Why it's wrong: Collisions are modelled as elastic, so kinetic energy is redistributed rather than lost.
Check instead: Remember that a sealed gas at constant temperature keeps its pressure indefinitely.
No practice questions are available for this topic yet. You can still practice the whole unit.
The kinetic molecular theory models a gas as a very large number of tiny particles in constant, random, straight-line motion, with negligible volume and no attractions, colliding elastically. Pressure is the total force of those collisions per unit area, and the average kinetic energy of the particles depends only on the absolute temperature.
This lesson is original Chem Help content. No external sources were adapted.