Understanding the First Law of Thermodynamics
The First Law of Thermodynamics is a fundamental principle of physics that states energy cannot be created or destroyed in an isolated system; it can only be transferred or changed from one form to another. It is essentially the law of conservation of energy applied to heat and thermodynamic processes.
Section 2: The Core Equation
The law is most commonly expressed by the equation: ΔU = Q - W. In this formula, ΔU represents the change in the internal energy of a system, Q is the net heat transferred into the system, and W is the net work done by the system on its surroundings. This equation shows that a system's internal energy can be changed by adding heat or by doing work.
Section 3: A Practical Example
Consider a gas in a cylinder with a movable piston. If you add heat (Q) to the gas, its internal energy (U) increases, causing the gas molecules to move faster. This increased pressure pushes the piston outward, causing the gas to do work (W) on the surroundings. According to the first law, the heat you added is accounted for by the increase in the gas's internal energy and the work it performed.
Section 4: Importance of the First Law
This law is critically important because it governs all energy transformations in the universe. It is the foundational principle behind engines, refrigerators, power plants, and even metabolic processes in living organisms. It confirms that a perpetual motion machine that creates energy out of nothing is impossible, as energy input must always equal energy output plus any change in stored energy.