Topic 3: Thermal Physics

 

3.1 Thermodynamic properties (thermal concepts)

Temperature. What is it? Please, none of that ‘how cold or hot something is’. Temperature is a measure of the average internal energy of a system. The internal energy is the sum of the potential and kinetic energy. Specifically, the ‘microscopic’ potential and kinetic energy. So if you put a box of temperature T on a really fast plane, and then fly the plane, the macroscopic kinetic energy increases, but the microscopic kinetic energy associated with the molecules in the box stays the same. The temperature thus stays the same, as expected.

What about pressure? The molecules, usually of gas, whizz around in a container. When they bounce off walls, the wall has to exert a force on them in order to change their velocities. Similarly, they exert a force on the wall. If there are many particles, then the average force on the wall at any time is constant. Divide the force by the area of the wall, and you get pressure.

How much energy do you need to transfer to a system to increase its temperature by a fixed amount? This quantity is giving by the heat capacity. The heat capacity of an object is the energy required to raise its temperature by 1K. The specific heat capacity is the energy required to raise 1kg of an object by 1K.

During changes of phase, the intermolecular separation changes. That is, the potential energy part of the internal energy changes, while the kinetic energy part stays the same (alright, not in all cases, but let’s keep things simple for this syllabus). The energy required to change the phase of a substance is called the latent heat of of vapourisation/fusion (liquid to gas and liquid to solid respectively). Of course, that depends on ‘how much substance’ there is. So we introduce the specific latent heat, which is the energy required to change the phase of 1kg of a substance.

3.2 Ideal gases (modelling a gas)

Physical phenomena might be difficult to understand. So we make approximations. Simplify things. Try to understand the key properties of a system first. Followed by slowly making it more complex, taking into consideration more things. To develop an understanding of gases, we will begin with perhaps the simplest model of a gas: the ideal gas.

There are a few assumptions we make. The two most important ones are as follows. That the volume of the gas particles is negligible compared to the volume the gas occupies (ie the ‘container’), and that the interaction between the gas particles is negligible except during short, infrequent collisions.

Equation of state: pV = nRT

Deviations from ideality

 

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