The first law of thermodynamics,
ΔU = Q + W
Where:
- ΔU is the change in internal energy as a result of work or heating done by or on the system in Joules
- Q is the thermal energy put into the system (i.e. the heating done) in Joules
- W is the work done on the system in Joules
Types of changes due to heating and working
These changes involve one of the values (ΔU, Q or W) staying constant.
| Isothermal | No change in internal energy ΔU ΔU = 0 ∴ W = -Q |
| Adiabatic | No heat lost or gained Q = 0 ∴ ΔU = W |
| Constant volume | No work is done on or by the system Only heat transfer occurs, so this is most common in solids and liquids (as they don't expand a large amount when heated). W = 0 ∴ ΔU = Q |
Q = mcΔθ
Where c = the specific heat capacity of the substance, and Δθ is the change in temperature in Kelvin.
and Q = ItV, Q = I2Rt
Q = ml
Where l = the specific latent heat of the substance in J kg-1
The area under a force-distance graph is equal to the work done
W = Fs
Gasses
Work done = pΔV
- In a pressure-volume graph:
- the area under the graph = work done,
- work is done on the gas when the volume is decreased under constant pressure
- work is done by the gas when the volume is increased under constant pressure.