Extensive properties depend on the amount of matter in the system. Volume is an extensive property as it depends on the size or amount of the substance, whereas temperature, pressure, and density are intensive properties as they do not depend on the amount of matter.

Enthalpy (H) is defined as H = U + PV. The change in enthalpy (ΔH) is equal to the heat absorbed or released at constant pressure, as under constant pressure conditions, ΔH = q_p.

For a spontaneous process, the Gibbs free energy change (ΔG) is negative. This indicates that the process can occur without external intervention under the given conditions.

In an isothermal process, the temperature remains constant. For an ideal gas, since internal energy depends only on temperature, the internal energy remains constant (ΔU = 0).

The first law of thermodynamics states that energy can neither be created nor destroyed, only transferred or transformed. Hence, it is also known as the law of conservation of energy.

In an adiabatic process, there is no heat exchange between the system and its surroundings. Therefore, q = 0. From the first law of thermodynamics, ΔU = q + w becomes ΔU = w.

Entropy is a measure of disorder. During evaporation of water, the system transitions from a liquid to a gas, increasing the disorder and thus increasing the entropy.

For a reaction involving ideal gases, the relationship between enthalpy change (ΔH) and internal energy change (ΔU) is given by ΔH = ΔU + PΔV, where PΔV accounts for the work done due to volume change at constant pressure.

A state function depends only on the current state of the system, not on the path taken to reach that state. Heat (q) is a path-dependent quantity and thus is not a state function, whereas internal energy, enthalpy, and entropy are state functions.

According to the second law of thermodynamics, the entropy of the universe always increases during a spontaneous process. This reflects the natural tendency of systems to move toward greater disorder.