Enthalpy of reaction

The enthalpy of reaction indicates the change in enthalpy in the course of a reaction, i.e. the energy consumption of a reaction carried out at constant pressure . According to Hess's law of heat , it does not matter in which way the reaction takes place or in which form ( heat , work ) energy is absorbed or given off during the reaction. The enthalpy of reaction is always the difference between the enthalpies of formation of the products and the reactants : ${\displaystyle \Delta _{\mathrm {R} }H}$

${\displaystyle \Delta _{\mathrm {R} }H=H_{\mathrm {Produkte} }-H_{\mathrm {Reaktante} }}$

Since substances have different energies depending on temperature and pressure (to understand: a gas has more energy stored under high pressure than under low pressure), energy balances of different reactions can only be compared directly with one another if one refers to the same external conditions. Usually standard conditions are used for this , less often normal conditions . The enthalpy of reaction under standard conditions is called the standard enthalpy of reaction . ${\displaystyle \Delta _{\mathrm {R} }H^{0}}$

In chemistry, the molar enthalpy of reaction is mostly used, in which the enthalpy of reaction is related to the amount of substance in the underlying reaction equation (cf. conversion variable and formula conversion ). The unit of the molar enthalpy of reaction is accordingly joules per mole . Enthalpies of formation of organic substances can be calculated very well with the Benson method . ${\displaystyle {\bigl (}\mathrm {\tfrac {J}{mol}} {\bigr )}}$

sign

Negative: exothermic

exothermic reaction; in the figure corresponds to in the text
${\displaystyle \Delta E_{i}}$ ${\displaystyle \Delta _{\mathrm {R} }H}$

Since the energy balance is specified for the system, it is negative if the products are energetically lower than the starting materials and thus energy is given off in total: ${\displaystyle \Delta _{\mathrm {R} }H}$

${\displaystyle {\begin{aligned}\Delta _{\mathrm {R} }H&<0\\\Leftrightarrow H_{\mathrm {Produkte} }&<H_{\mathrm {Edukte} }\end{aligned}}}$

If the released energy is not converted, heat is released and the sample heats up. The reaction is therefore exothermic .

Positive: endothermic

endothermic reaction; in the figure corresponds to in the text
${\displaystyle \Delta E_{i}}$ ${\displaystyle \Delta _{\mathrm {R} }H}$

However, if energy has to be absorbed, since the products have a higher energy than the starting materials, it becomes positive: ${\displaystyle \Delta _{\mathrm {R} }H}$

${\displaystyle {\begin{aligned}\Delta _{\mathrm {R} }H&>0\\\Leftrightarrow H_{\mathrm {Produkte} }&>H_{\mathrm {Edukte} }\end{aligned}}}$

The necessary energy is often taken from the ambient heat, the environment becomes colder. Processes in which heat is absorbed are called endothermic .

quantification

In the case of reactions of molecular substances, the sign of the enthalpy of reaction can be estimated on the basis of the broken and newly formed bonds during the reaction. This is based on the observation that polar bonds are more stable, i.e. lower in energy, than non-polar bonds. If there are more polar bonds in the product molecules than in the reactant molecules, then it is an exothermic reaction , in the opposite case it is an endothermic reaction . ${\displaystyle \Delta H_{\mathrm {R} }}$

A second, more precise way is based on the differential calculation of the enthalpy of formation of educts and products.

Temperature dependence

The enthalpy of reaction, like the enthalpy of formation, is temperature-dependent. With the proviso that in the considered temperature interval (from to ) is not a phase transition occurs, the enthalpy results in as follows: ${\displaystyle T_{1}}$ ${\displaystyle T_{2}}$ ${\displaystyle T_{2}}$

${\displaystyle H(T_{2})=H(T_{1})+\int _{T_{1}}^{T_{2}}\mathrm {C} _{\mathrm {p} }\,\cdot \mathrm {d} T}$.

If the heat capacity remains approximately constant within the selected temperature range, it can be approximated before the integral. ${\displaystyle C_{p}}$

If a reaction is now considered, Kirchhoff's law results for the enthalpy of reaction :

${\displaystyle \Rightarrow \Delta _{\mathrm {R} }H(T_{2})=\Delta _{\mathrm {R} }H(T_{1})+\int _{T_{1}}^{T_{2}}\Delta _{\mathrm {R} }C_{\mathrm {p} }\cdot \mathrm {d} T}$.

Here arises from the molar heat capacities of the substances involved in the reaction their respective stoichiometric factors : ${\displaystyle \Delta _{\mathrm {R} }C_{\mathrm {p} }}$ ${\displaystyle \nu }$

${\displaystyle \Delta _{\mathrm {R} }C_{\mathrm {p} }=\sum _{k=1}^{N}\nu _{\mathrm {k} }\cdot C_{\mathrm {p,k} }}$

Related sizes

At constant pressure:

• Enthalpy
• Gibbs energy (free enthalpy)
• Chemical potential

With constant volume:

• Inner energy
• Helmholtz energy

Individual evidence

1. ↑ Patrik Good: 5 Thermodynamics (lecture notes) ( Memento from January 17, 2012 in the Internet Archive ) (PDF; 1.1 MB). P. 6 (19 p.).