The standard enthalpy of formation or standard heat of formation of a compound is the change of enthalpy during the formation of 1 mole of the substance from its constituent elements, with all substances in their standard states. The standard pressure value p⦵ = 105 Pa is recommended by IUPAC, although prior to 1982 the value 1.00 atm was used. There is no standard temperature, its symbol is ΔfH⦵. The superscript Plimsoll on this symbol indicates that the process has occurred under standard conditions at the specified temperature. Standard states are as follows: For a gas: the hypothetical state it would have assuming it obeyed the ideal gas equation at a pressure of 1 bar For a solute present in an ideal solution: a concentration of one mole per liter at a pressure of 1 bar For a pure substance or a solvent in a condensed state: the standard state is the pure liquid or solid under a pressure of 1 bar For an element: the form in which the element is most stable under 1 bar of pressure. One exception is phosphorus, for which the most stable form at 1 bar is black phosphorus, but white phosphorus is chosen as the standard reference state for zero enthalpy of formation.

For example, the standard enthalpy of formation of carbon dioxide would be the enthalpy of the following reaction under the above conditions: C + O2 → CO2All elements are written in their standard states, one mole of product is formed. This is true for all enthalpies of formation; the standard enthalpy of formation is measured in units of energy per amount of substance stated in kilojoule per mole, but in kilocalorie per mole, joule per mole or kilocalorie per gram. All elements in their standard states have a standard enthalpy of formation of zero, as there is no change involved in their formation; the formation reaction is a constant pressure and constant temperature process. Since the pressure of the standard formation reaction is fixed at 1 bar, the standard formation enthalpy or reaction heat is a function of temperature. For tabulation purposes, standard formation enthalpies are all given at a single temperature: 298 K, represented by the symbol ΔfH⦵298 K. For many substances, the formation reaction may be considered as the sum of a number of simpler reactions, either real or fictitious.

The enthalpy of reaction can be analyzed by applying Hess's Law, which states that the sum of the enthalpy changes for a number of individual reaction steps equals the enthalpy change of the overall reaction. This is true because enthalpy is a state function, whose value for an overall process depends only on the initial and final states and not on any intermediate states. Examples are given in the following sections. For ionic compounds, the standard enthalpy of formation is equivalent to the sum of several terms included in the Born–Haber cycle. For example, the formation of lithium fluoride, Li + ​1⁄2 F2 → LiFmay be considered as the sum of several steps, each with its own enthalpy: The standard enthalpy of atomization of solid lithium; the first ionization energy of gaseous lithium. The standard enthalpy of atomization of fluorine gas; the electron affinity of a fluorine atom. The lattice energy of lithium fluoride; the sum of all these enthalpies will give the standard enthalpy of formation of lithium fluoride.

Δ H f = Δ H sub + IE Li + 1 2 B − EA F − U L. In practice, the enthalpy of formation of lithium fluoride can be determined experimentally, but the lattice energy cannot be measured directly; the equation is therefore rearranged in order to evaluate the lattice energy. U L = Δ H sub + IE Li + 1 2 B − EA F + Δ H f; the formation reactions for most organic compounds are hypothetical. For instance and hydrogen will not directly react to form methane, so that the standard enthalpy of formation cannot be measured directly; however the standard enthalpy of combustion is mesurable using bomb calorimetry. The standard enthalpy of formation is determined using Hess's law; the combustion of methane is equivalent to the sum of the hypothetical decomposition into elements followed by the combustion of the elements to form carbon dioxide and water: CH4 → C + 2 H2 C + O2 → CO2 2 H2 + O2 → 2 H2OApplying Hess's law, ΔcombH⦵ = − ΔfH⦵. Solving for the standard of enthalpy of formation, ΔfH⦵ = − ΔcombH⦵; the value of ΔfH⦵ is determined to be −74.8 kJ/mol.

The negative sign shows. It is possible to predict heats of formation for simple unstrained organic compounds with the heat of formation group additivity method; the standard enthalpy change of any reaction can be calculated from t

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