Why is enthalpy a state function?

Enthalpy (H) is considered a state function in thermodynamics. State functions are properties that depend only on the current state of a system and are independent of the path taken to reach that state. Enthalpy meets this criterion for several reasons:

Independence of Path: Enthalpy is defined as the internal energy (U) plus the product of pressure (P) and volume (V) of a system, represented by the equation H = U + PV. Since U, P, and V are all state functions, the sum H is also a state function. This means that the value of enthalpy at a specific state of a system is determined solely by the initial and final states, regardless of how the system got there (i.e., the specific process or path taken).

Measurable Property: Enthalpy can be measured experimentally. When you measure the enthalpy change (ΔH) for a chemical reaction or a physical process, such as heating or cooling, the value of ΔH depends only on the initial and final states of the system and not on the specific conditions or the route taken to change from one state to another.

Useful in Thermodynamics: Enthalpy is a particularly useful state function because it directly relates to heat exchange at constant pressure. When a chemical reaction occurs at constant pressure, the change in enthalpy, ΔH, is equal to the heat absorbed or released by the system. This relationship simplifies thermodynamic calculations and is widely used in various applications, including calorimetry and chemical engineering.

In summary, enthalpy is a state function because it satisfies the criteria of being path-independent and having a well-defined value at each state of the system. Its definition and properties make it a valuable tool for studying and understanding the thermodynamics of physical and chemical processes.