Though this may sound complex, it's really a very simple idea. If you add heat to a system, there are only two things that can be done -- change the internal energy of the system or cause the system to do work (or, of course, some combination of the two). All of the heat energy must go into doing these things.
First Law of ThermodynamicsThe change in a system's internal energy is equal to the difference between heat added to the system from its surroundings and work done by the system on its surroundings.
Mathematical Representation of the First LawPhysicists typically use uniform conventions for representing the quantities in the first law of thermodynamics. They are:
- U1 (or Ui) = initial internal energy at the start of the process
- U2 (or Uf) = final internal energy at the end of the process
- delta-U = U2 - U1 = Change in internal energy (used in cases where the specifics of beginning and ending internal energies are irrelevant)
- Q = heat transferred into (Q > 0) or out of (Q < 0) the system
- W = work performed by the system (W > 0) or on the system (W < 0).
U2 - U1 = delta-U = Q - WThe analysis of a thermodynamic process, at least within a physics classroom situation, generally involves analyzing a situation where one of these quantities is either 0 or at least controllable in a reasonable manner. For example, in an adiabatic process, the heat transfer (Q) is equal to 0 while in an isochoric process the work (W) is equal to 0.
Q = delta-U + W
The First Law & Conservation of EnergyThe first law of thermodynamics is seen by many as the foundation of the concept of conservation of energy. It basically says that the energy that goes into a system cannot be lost along the way, but has to be used to do something ... in this case, either change internal energy or perform work.
Taken in this view, the first law of thermodynamics is one of the most far-reaching scientific concepts ever discovered.
Laws of Thermodynamics