Difference between revisions of "Enthalpy"
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Explanation
Enthalpy of a system is defined as the sum of the change in internal energy of the system and the amount of work done by that system to change the volume and pressure to accommodate itself in a state of equilibrium. As the internal energy of a system depends on different variables and cannot be accurately measured, Enthalpy was introduced as a thermodynamic potential to measure the total energy of a system. Mathematically, enthalpy H is related to the internal energy U, volume V and pressure P as ΔH = ΔU + Δ(PV), and it is measured in the SI unit, joule.
Frequently Asked Questions
Why enthalpy can't be measured directly?
As enthalpy is technically a measure of the sum of a system's internal energy and the energy required to accommodate itself, it is impractical and impossible to look for a reference point to begin the measure. Even at the absolute zero temperature scale, all systems would possess a minimum amount of energy called zero-point energy. So by convention, measuring the change in enthalpy with a reference point is easier and convenient in practice. For example, when you pour a litre of water into the ocean, you can tell that you had increased the quantity of the ocean by 1 litre, without knowing how many litres was initially in the ocean. This metaphor can help understanding why the change in enthalpy is convenient in reality.
In a system at a constant pressure, the change in enthalpy is equal to the amount of heat transferred by the system. But enthalpy is a state function, which means that the change in enthalpy does not depend on the path of the function. Only the initial and final state of the system matters. In contrary, heat is a path function that is exclusively governed by the temperature gradient of the system. Thus under constant pressure, the change in heat becomes equal to the enthalpy.