Chemical Equilibrium in the Haber Process

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This Demonstration calculates the number of moles at equilibrium for the reversible, exothermic reaction that synthesizes ammonia () from hydrogen (
) and nitrogen (
), known as the Haber process. This reaction typically takes place near 200 bar and 675 to 725 K. The system starts with 1 mol
and goes to equilibrium. Use sliders to add additional moles of
,
and
at constant pressure and temperature and observe how they change equilibrium. Vary pressure and temperature with sliders. Because 4 mol of reactants form 2 mol of product, raising the pressure shifts equilibrium toward products. Gases are assumed ideal, but at the high pressures used for this reaction, significant deviation from ideal behavior is likely.
Contributed by: Benjamin L. Kee and Rachael L. Baumann (May 2014)
Additional contributions by: John L. Falconer
(University of Colorado Boulder, Department of Chemical and Biological Engineering)
Open content licensed under CC BY-NC-SA
Snapshots
Details
The reaction is used in the Haber process. The moles of each component at equilibrium is:
,
where are the moles of component
added,
is the stoichiometric coefficient and
is extent of reaction (mol). Initially only 1 mol
is present.
The mole fraction at equilibrium is:
,
,
where is the total number of moles.
The extent of reaction is found by setting the equilibrium constant equal to the equilibrium rate constant
and solving for
:
,
,
where is Gibbs free energy (J/mol),
is the heat of reaction (J/mol),
is the entropy change of reaction (J/(mol K)),
is temperature (K),
is the ideal gas constant (J/(mol K)) and
is pressure (bar).
The screencast video at [1] shows how to use this Demonstration. The screencast at [2] show solutions of an example problem on gas phase equilibrium.
References
[1] Chemical Equilibrium in the Haber Process [Video]. (Jan 20, 2017) www.colorado.edu/learncheme/thermodynamics/HaberProcess.html.
[2] Gas Phase Chemical Equilibrium [Video]. (Jan 20, 2017) www.youtube.com/watch?v=3ArBH_gbsNw.
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