Kinetic Theory of Matter and Behaviour of Gases (LC 2027) (Leaving Cert Chemistry): Revision Notes
The Kinetic Theory of Matter Applied to Gases
The kinetic theory of matter helps us understand how gas particles behave at the molecular level. When we apply this theory specifically to gases, we can explain many of the properties and behaviours we observe in real-world gas samples.
Understanding kinetic theory for gases
The kinetic theory describes gases as being made up of tiny particles (molecules or atoms) that are in constant, random motion. These particles are incredibly small compared to the spaces between them, and they move very rapidly in straight lines until they collide with other particles or the container walls.
When we examine gas behaviour using kinetic theory, we need to understand both what the theory assumes about ideal gases and how real gases sometimes behave differently from these assumptions.
Understanding the distinction between ideal and real gas behaviour is crucial for applying kinetic theory correctly in different situations. This theory provides the foundation for explaining gas laws and predicting gas behaviour under various conditions.
Assumptions of the kinetic theory of matter as applied to gases
The kinetic theory makes five key assumptions about how gas particles behave:
These five assumptions form the foundation of kinetic theory for gases. While they don't perfectly describe real gases, they provide an excellent framework for understanding gas behaviour under most conditions.
1. Continuous random motion Gas particles (molecules or atoms) move continuously and randomly with each other and with the walls of their container. This constant motion is what creates gas pressure when particles hit the container walls.
2. Negligible attractive or repulsive forces
There are no attractive or repulsive forces between the molecules of a gas. This means the molecules are completely independent of each other during their movement.
3. Negligible molecular volume The gas molecules are so small and widely separated that the actual volume of all the molecules is negligible compared with the total volume of the container they occupy.
4. Elastic collisions When molecules collide with each other or the container walls, the collisions are perfectly elastic. This means there is no loss of kinetic energy in these collisions, although there may be a transfer of energy between the colliding particles.
5. Temperature-dependent kinetic energy The average kinetic energy of the molecules in a gas sample is proportional to the temperature measured on the Kelvin (absolute) scale. Higher temperatures mean faster-moving molecules.
Limitations to the kinetic theory of matter as applied to gases
While the kinetic theory is very useful, it has some important limitations when we examine real gases more closely. Understanding these limitations helps us recognise when the theory's predictions might be less accurate.
Common Misconception Alert: Many students think kinetic theory perfectly describes all gas behaviour. However, real gases deviate from ideal behaviour, especially under extreme conditions of temperature and pressure.
Intermolecular forces do exist Contrary to assumption 2, real gases actually do have tiny attractive or repulsive forces between molecules. For example, we studied polar molecules like NH₃ and HCl, which have attractive intermolecular forces because the partially negative end of one molecule will be attracted to the partially positive end of another molecule.
Molecules occupy space Real gas molecules aren't actually negligible in size compared to the space they occupy, especially under very high pressure when molecules are crowded closer together. Under these conditions, the volume that the molecules themselves occupy becomes significant compared with the distances between them.
Temperature and pressure effects At very low temperatures and high pressures, the assumptions of kinetic theory become less accurate because molecules move more slowly and are forced closer together.
Real gases vs ideal gases
An ideal gas is one that perfectly obeys all the assumptions of the kinetic theory of gases under all conditions of temperature and pressure. However, ideal gases don't actually exist in reality.
Real gases do come very close to ideal gas behaviour under two specific conditions:
- At low pressure - when the molecules are widely spaced
- At high temperature - when the molecules are moving rapidly, preventing the forces between molecules from exerting significant influence
At very high pressures and low temperatures, molecules are close together and moving slowly. Under these conditions, the attractive forces between molecules can be significant, and real gases behave quite differently from ideal gases.
Under normal atmospheric conditions (around 1 atmosphere pressure and 20°C), most real gases behave very similarly to ideal gases. This is why the kinetic theory remains extremely useful for everyday applications and calculations.
When real gases behave most like ideal gases
Real gases differ from ideal gases because:
- Forces of attraction and repulsion do exist between the molecules
- The volume of the molecules is not negligible, especially at high pressure
Under normal atmospheric conditions (around 1 atmosphere pressure and 20°C), the behaviour of real gases comes very close to that of ideal gases. This means we can reasonably apply the assumptions of kinetic theory to real gases under everyday conditions.
However, at extremely low temperatures like -273°C (absolute zero), real gases become liquids long before they reach this temperature, showing that the kinetic theory assumptions break down under extreme conditions.
Worked Example: Explaining Gas Behaviour
Question: Give two reasons why real gases depart from ideal behaviour. Under what conditions of temperature and pressure do real gases come closest to ideal behaviour?
Answer:
Real gases depart from ideal behaviour because:
(i) Intermolecular forces exist - In an ideal gas, it is assumed that there are no forces of attraction between the molecules. In real gases, attractive forces exist between the molecules.
(ii) Molecular volume is significant - In an ideal gas, it is assumed that the molecules occupy negligible space. However, in real gases, the molecules occupy a small but definite space, which is not negligible at high pressure.
Real gases come closest to ideal behaviour at high temperature and low pressure.
Key Points to Remember:
- The kinetic theory makes five key assumptions about gas particle behaviour: continuous motion, no intermolecular forces, negligible size, elastic collisions, and temperature-dependent energy
- Real gases don't perfectly follow these assumptions because molecules do have forces between them and do occupy space
- Real gases behave most like ideal gases at high temperatures and low pressures
- Under normal atmospheric conditions, real gases behave very similarly to ideal gases
- The kinetic theory helps explain gas laws like Boyle's Law, Charles' Law, and Avogadro's Law