What does the first law of thermodynamics state about the internal energy of a system?

The first law of thermodynamics is fundamentally a statement about the conservation of energy. It asserts that the internal energy of a system changes as energy is added to or removed from the system through work or heat transfer. Specifically, the law can be mathematically expressed as:

ΔU = Q - W

where ΔU represents the change in internal energy, Q is the heat added to the system, and W is the work done by the system.

Choosing the correct answer emphasizes the relationship described by the first law, which indicates that the internal energy of a system is equal to the heat added to the system minus the work done by the system. This principle allows us to analyze and predict how different processes affect a system's energy state, making it a foundational concept in thermodynamics.

In contrast, the other choices misinterpret aspects of thermodynamics. For instance, stating that internal energy remains constant during all processes does not hold true, as energy can be transferred in and out of the system, affecting internal energy. The claim that internal energy increases with every cycle of energy transfer overlooks the specifics of energy conservation in each cycle, where energy can be lost as work or heat. Lastly, the idea that internal energy only depends on temperature changes simplifies the concept too