What does the first law of thermodynamics state in terms of internal energy?

The first law of thermodynamics fundamentally states that the change in internal energy of a system is equal to the heat added to the system minus the work done by the system. This can be articulated mathematically as:

ΔU = Q - W

In this context, ΔU represents the change in internal energy, Q signifies the heat added to the system, and W stands for the work done by the system. If the system does work on its surroundings, that work is subtracted from the heat added because it detracts from the energy remaining within the system, which relates to its internal energy. Therefore, when heat is added or work is done, the internal energy changes accordingly.

In summary, understanding that internal energy changes based on the interactions of heat and work allows one to grasp the essence of the first law of thermodynamics. This principle emphasizes the conservation of energy within a closed system, reinforcing that energy cannot be created or destroyed but can only change forms or be transferred in various ways.