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Enthalpy (/ ˈ ɛ n θ əl p i / ⓘ) is the sum of a thermodynamic system's internal energy and the product of its pressure and volume. [1] It is a state function in thermodynamics used in many measurements in chemical, biological, and physical systems at a constant external pressure, which is conveniently provided by the large ambient atmosphere.
The law states that the total enthalpy change during the complete course of a chemical reaction is independent of the sequence of steps taken. [2] [3] Hess's law is now understood as an expression of the fact that the enthalpy of a chemical process is independent of the path taken from the initial to the final state (i.e. enthalpy is a state ...
The Van 't Hoff equation relates the change in the equilibrium constant, K eq, of a chemical reaction to the change in temperature, T, given the standard enthalpy change, Δ r H ⊖, for the process. The subscript r {\displaystyle r} means "reaction" and the superscript ⊖ {\displaystyle \ominus } means "standard".
The enthalpy of solution is most often expressed in kJ/mol at constant temperature. The energy change can be regarded as being made up of three parts: the endothermic breaking of bonds within the solute and within the solvent, and the formation of attractions between the solute and the solvent. An ideal solution has a null enthalpy of mixing.
In an isenthalpic process, system enthalpy (H) is constant. In the case of free expansion for an ideal gas, there are no molecular interactions, and the temperature remains constant. For real gasses, the molecules do interact via attraction or repulsion depending on temperature and pressure, and heating or cooling does occur.
Thus, they are essentially equations of state, and using the fundamental equations, experimental data can be used to determine sought-after quantities like G (Gibbs free energy) or H . [1] The relation is generally expressed as a microscopic change in internal energy in terms of microscopic changes in entropy , and volume for a closed system in ...
The Clapeyron equation allows us to use pressure, temperature, and specific volume to determine an enthalpy change that is connected to a phase change. It is significant to any phase change process that happens at a constant pressure and temperature.
In solution chemistry, on the other hand, most chemical reactions are kept at constant pressure. Under this condition, the heat q {\displaystyle q} of the reaction is equal to the enthalpy change Δ H {\displaystyle \Delta H} of the system.