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 compound from
its constituent elements, with all substances in their standard states at 1
atmosphere. Its symbol is ΔHfO or ΔfHO. The superscript theta 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 standard state is a pressure of exactly 1 atm
For a solute present in an ideal solution: a concentration of exactly one
mole/liter at a pressure of 1 atm 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 atm For an element: the form in which the
element is most stable under 1 atm of pressure. One exception is phosphorus,
for which the most stable form at 1 atm 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 conditions above: C(s,graphite) + O2(g) → CO2(g)
All elements are written in their standard states, and 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, usually stated in kilojoule per mole, but also in calorie
per mole, joule per mole or kilocalorie per gram. In physics the energy per
particle is often expressed in electronvolts which corresponds to about
100 kJ mol−1. 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 atm, 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 ΔHf298O . Calculation
The standard enthalpy of formation is equivalent to the sum of many separate
processes included in the Born-Haber cycle of synthesis reactions. For
example, to calculate the standard enthalpy of formation of sodium
chloride, we use the following reaction: Na(s) +Cl2(g) → NaCl(s)
This process is made of many separate sub-processes, each with its own
enthalpy. Therefore, we must take into account:
The standard enthalpy of atomization of solid sodium
The first ionization energy of gaseous sodium
The standard enthalpy of atomization of chlorine gas
The electron affinity of chlorine atoms The lattice enthalpy of sodium chloride
The sum of all these values will give the standard enthalpy of formation of
sodium chloride. Additionally, applying Hess's Law shows
that the sum of the individual reactions corresponding to the enthalpy change of
formation for each substance in the reaction is equal to the enthalpy change
of the overall reaction, regardless of the number of steps or intermediate
reactions involved. This is because enthalpy is a state function. In the
example above the standard enthalpy change of formation for sodium chloride
is equal to the sum of the standard enthalpy change of formation for each of
the steps involved in the process. This is especially useful for very long
reactions with many intermediate steps and compounds.
Chemists may use standard enthalpies of formation for a reaction that is
hypothetical. For instance carbon and hydrogen will not directly react to form
methane, yet the standard enthalpy of formation for methane is determined to
be −74.8 kJ mol−1 from using other known standard enthalpies of reaction with
Hess's law. That it is negative shows that the reaction, if it were to
proceed, would be exothermic; that is, it is enthalpically more stable than
hydrogen gas and carbon. It is possible to predict heat of
formations for simple unstrained organic compounds with the Heat of formation
group additivity method. Standard enthalpy of reaction
Standard enthalpies of formation are used in thermochemistry to find the
standard enthalpy change of any reaction. This is done by subtracting
the sum of the standard enthalpies of formation of the reactants from the sum
of the standard enthalpies of formation of the products, as shown in the
equation below: ΔH° = Σ(ν × ΔHf°) - Σ(ν × ΔHf°)
This calculation has a tacit assumption of ideal solution between reactants and
products where the enthalpy of mixing is zero.
For example, for the reaction CH4 + 2 O2 → CO2 + 2 H2O:
ΔHr° = [(1 × ΔHf°(CO2)) +)] - [(1 × ΔHf°(CH4)) +)]
If the standard enthalpy of the products is less than the standard enthalpy of
the reactants, the standard enthalpy of reaction will be negative. This implies
that the reaction is exothermic. The converse is also true; the standard
enthalpy of reaction will be positive for an endothermic reaction.
Key concepts for doing enthalpy calculations
When a reaction is reversed, the magnitude of ΔH stays the same, but the
sign changes. When the balanced equation for a
reaction is multiplied by an integer, the corresponding value of ΔH must be
multiplied by that integer as well. The change in enthalpy for a reaction
can be calculated from the enthalpies of formation of the reactants and the
products Elements in their standard states make
no contribution to the enthalpy calculations for the reaction since the
enthalpy of an element in its standard state is zero. Allotropes of an element
other than the standard state generally have non-zero standard enthalpies of
formation. Subcategories
Standard enthalpy of neutralization is the change in enthalpy that occurs when
an acid and base undergo a neutralization reaction to form one mole
of water under standard conditions, as previously defined.
Standard enthalpy of sublimation, or heat of sublimation, is defined as the
enthalpy required to sublime one mole of the substance under standard conditions,
as previously defined. Standard enthalpy of solution is the
enthalpy change associated with the dissolution of a substance in a solvent
at constant pressure under standard conditions, as previously defined.
Standard enthalpy of hydrogenation is defined as the enthalpy change observed
when one mole of an unsaturated compound reacts with an excess of hydrogen to
become fully saturated under standard conditions, as previously defined.
Examples: Standard Enthalpies of Formation
(State: g = gaseous; l = liquid; s = solid; aq = aqueous)
See also Thermochemistry
Enthalpy Heat of combustion
Calorimetry Standard enthalpy change of formation
External links NIST Chemistry WebBook
References