⇒ The kinetic theory model of solids, liquids and gases assumes that particles are incompressible spheres

⇒ Solids have a close-packed, regular particle structure - the particles vibrate about fixed points

⇒ Liquids have a close-packed, random, irregular particle structure - the particles are free to move

⇒ Gases have a widely spaced, irregular particle structure - the particles are free to move

⇒ Thermal energy can be transferred from somewhere hot (at a high temperature) to somewhere cooler (at a lower temperature) by the processes of conduction, convection, radiation, and evaporation

⇒ Developing engines required an understanding of heat energy - this became known as thermodynamics

⇒ Thermodynamics deals with the macroscopic (large-scale) behaviour of a system, but it is complemented by the kinetic theory of matter, which deals with the microscopic, particle-scale behaviour of matter

⇒ Internal energy, U, is one of the most fundamental properties of thermodynamics - internal energy is the sum (sometimes called an ensemble in thermodynamics) of the randomly distributed kinetic energies and potential energies of the particles in a body

⇒ Consider a glass of water:

• The water particles have two types of energy associated with their movement (the faster they move, or vibrate, or rotate, the higher their kinetic energy) and potential energy associated with any forces or interactions between the particles (such as any electrostatic attraction or repulsion)

⇒ The kinetic energies of the particles depend on their temperature, and the potential energies depend on any intermolecular forces between the particles

• For ideal gases, in which there are no intermolecular forces, the internal energy is dependent on only the kinetic energies

⇒ The modern version of this law is stated as follows: the increase in internal energy of a system is equal to the heat added to the system minus the work done by the system

⇒ In terms of symbols, this can be written:

⇒ Where ΔU is the increae in internal energy of the system (usually a gas), ΔQ is the thermal energy added to the system, and ΔW is the work done by the system

⇒ When a gas expands, it exerts a force on the surroundings, causing them to move - the gas does work on the surroundings

⇒ We can use the first law of thermodynamics to determine the work done, ΔW, by an expanding gas at constant temperature (called an isothermal change)

⇒ Consider a gas enclosed in a cylinder by a frictionless piston

⇒ The gas of volume W exerts a pressure p on the walls of the cylinder. This in turn exerts a force on the frictionless piston of area A, where F = pA

⇒ This causes an increase in the volume, ΔV

• We assume that ΔV is very small and that the force moves the piston at a slow but steady rate such that the external force exerted on the piston is equal to the force exerted by the pressure p of the gas in the culinder
• This effectively makes the pressure exerted by the gas constant during the expansion

⇒ The gas does work, and so ΔW is positive

⇒ The force on the piston moves it through a distance, Δx, such that:

⇒ Also see our notes on:

• Heating up Substances and Changes of State
• The Gas Laws
• Avogadro's law, the ideal gas equation and moles
• Ideal Gases
• A molecular kinetic theory model
• Comparing two models of the behaviour of gases