Understanding the Kinetic Theory

• The Kinetic Theory is a model that explains the behaviour of gases. It assumes that gases are composed of a large number of small particles, either atoms or molecules.
• These particles are in constant, random motion. The rapidly moving particles constantly collide with each other and with the walls of the container, causing pressure.
• The overall energy of these particles remains constant unless there's an external influence, such as a change in temperature or pressure.

Assumptions of the Kinetic Theory

• Gases are made up of particles with insignificant volume. The volume of the individual gas particles compared to the volume of the empty space between them is negligible.
• No forces of attraction or repulsion exist between the particles when they aren't colliding.
• The gas particles are in constant, random motion. They collide with each other and with the walls of their container, exerting pressure.
• Collisions between the gas particles themselves, or with the walls of the container, are perfectly elastic. This means the total kinetic energy is conserved - no energy is lost or gained during the collision.
• The average kinetic energy of gas particles is directly proportional to the temperature of the gas, as per the equation E_k=1/2mv², where 'E_k' is kinetic energy, 'm' is mass, and 'v' is the average velocity of the particles.

Effects of Heat on Matter

• When a substance is heated, its particles gain kinetic energy, causing them to move faster and collide more often. Depending on the type of matter, this causes expansion, a rise in temperature, or a change in state.
• In solids, increased energy makes particles vibrate more rapidly around their fixed points. This increase in vibration can cause expansion of the material.
• In liquids and gases, an increase in energy causes faster and more erratic movement among particles, leading to expansion and an increase in temperature. For gases, this can also lead to increased pressure if the gas is confined in a space.
• When sufficient energy is supplied, a substance may change its state from solid to liquid (melting), or from liquid to gas (evaporation). These phase changes require energy but do not cause a temperature increase since the energy is used to break atomic bonds rather than increase kinetic energy.

The Gas Laws

• Based on the Kinetic Theory, physicists have outlined several laws that connect the variables of pressure, volume, temperature, and the number of moles when dealing with gases.
• Boyle's Law states that the pressure and volume of a gas are inversely proportional at a fixed temperature. Mathematically, for a given amount of gas at a constant temperature, P∝1/V or PV=const.
• Charles' Law states that the volume of a gas is directly proportional to its absolute temperature at a fixed pressure. In equation form, for a given amount of gas at constant pressure, V∝T or V/T=const.
• The Ideal Gas Law combines Boyle’s Law, Charles' Law, and Avogadro’s Law into a single equation of state for an ideal gas: PV=nRT, where 'P' is the pressure, 'V' is the volume, 'n' is the number of moles, 'R' is the gas constant, and 'T' is the absolute temperature.