ΔU = Q - W
Where:
* ΔU represents the change in internal energy of a system.
* Q represents the heat added to the system.
* W represents the work done by the system.
Explanation:
* Internal Energy (U): This is the total energy contained within a system, including kinetic and potential energy of its molecules.
* Heat (Q): This is the energy transferred between the system and its surroundings due to a temperature difference.
* Work (W): This is the energy transferred between the system and its surroundings due to a force acting over a distance.
Key Points:
* The first law of thermodynamics states that energy cannot be created or destroyed, only transferred or transformed.
* The equation expresses 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.
* If heat is added to the system (Q > 0), the internal energy increases (ΔU > 0).
* If work is done by the system (W > 0), the internal energy decreases (ΔU < 0).
Note:
* The sign convention for work can be reversed depending on the definition used. Some sources define work done *on* the system as positive, while others define work done *by* the system as positive.
* The first law of thermodynamics is a fundamental principle in physics and has wide-ranging applications in various fields, including engineering, chemistry, and biology.
