Applications of Chemistry in Manufacturing of Ammonia

Thermodynamics and Equilibrium - Free Energy and Favorable Changes

N2(g) + 3H2(g)

Free energy (Gibbs Free Energy, Symbol G)- a measure of the useful work that can be obtained from a
2NH3(g) △H = -92 Kj/mol
reaction.

△G = △H -T△S, where

△H the change in enthalpy

△S the change in entropy.
△H
△S
△G
Reaction
Comment
Exothermic (-)
+
Always -
Yes
No exceptions
Exothermic (-)
-
﹣At lower temperatures
Probably
At low temperature
Endothermic (+)
+
﹣At higher temperatures
Probably
At high temperature
Endothermic (+)
-
Never -
No
No exceptions


.

exothermic reaction.

△H is negative.

△S negative

-T△S is not large enough
△G is negative


The Rate Law - Rate Constant


Relatively low temperature
Rate Law equation, Rate =k [A]m[B]n

[A]/[B] - concentrations of reactants.

m, n - Rate Law exponents

k, rate constant.

Concentration of reactant↓ → Rate ↓

Concentration of reactant↑ → Rate ↑

Reaction - at high temperatures.
Le Chatelier's Principle

If a stress is placed upon a system in equilibrium, the equilibrium is displaced in the direction that
counteracts the effect of the stress.



Temperatures

Exothermic

Heat released

↓ temperatures, yield of ammonia ↑
Pressure

Pressure↑, yield of ammonia ↑

Catalyst



Fritz Haber,

The gas mixture – condenser.

Cooling the gas mixture.

the gas → a liquid,

Nitrogen → removed, equilibrium forward

Gases are recycled
Carl Bosch


Catalyst → Rate↑
Double-wall steel vesse
Catalyst

Iron catalyst ↑ rate of the reaction
The Change in Enthalpy

Normally represented by ∆H

Indicates the change in total internal energy of a system under constant pressure

Results from chemical bonds being broken and formed
Consider the reaction involved in manufacturing of ammonia:
N2(g) + 3H2(g) → 2NH3(g)

One mole of nitrogen-nitrogen triple bonds and three moles
of hydrogen-hydrogen single bond are broken. Two moles of
nitrogen-hydrogen triple bonds are formed

More energy is released to form bonds than is used to break
bonds

Results in a net release of energy, which is called an
exothermic reaction.
Enthalpy of Reaction

Dependent on conditions such as temperature and pressure

Usually refers to the standard enthalpy of reaction, ∆H°rxn

In the reaction of ammonia synthesis, ∆H°rxn = - 92 kJ. The negative sign indicates that the forward reaction
is exothermic and it releases energy.
Sustainable Ammonia Production

Ammonia production depends on plentiful supplies of natural gas

Using electrolysis of water to provide a sustainable source of hydrogen
2 H2O(l) → 2 H2(g) + O2(g)

Reduction at cathode:
2 H+(aq) + 2e− → H2(g)
The standard reduction potential is -0.414V

Anode (oxidation):
2 H2O(l) → O2(g) + 4 H+(aq) + 4e−
The standard reduction potential is 0.815V

Overall reaction: 2 H2O(l) → 2 H2(g) + O2(g)

The Ecell for the decomposition of pure water at 25℃ by electrolysis
has a value of -1.229 V
Therefore, the value of the external voltage needed is 1.229 V.

The number of hydrogen molecules produced is twice the number of oxygen molecules.
Under equal temperature and pressure for both gases,

The produced hydrogen gas has twice the volume of the produced oxygen gas.

The number of electrons pushed through the water is twice the number of generated hydrogen molecules and
four times the number of generated oxygen molecules.
The Use of Catalyst In Manufacturing Process of Ammonia

Iron-based catalyst is used to speed up the reaction

Prepared by exposing a mass of magnetite, an iron oxide, to the hot hydrogen feedstock

Maintains most of its bulk volume during the reduction, and so the result is a highly
porous material whose large surface area aids its effectiveness as a catalyst

Other minor components of the catalyst: calcium and aluminum oxides

Lower the activation energy of the reaction by providing an alternative mechanism for
the reaction

Increases the rate of a chemical reaction without being consumed

The reaction mechanism, involving the heterogeneous catalyst, is believed to be as follows:
1.
N2 (g) → N2 (adsorbed)
2.
N2 (adsorbed) → 2 N (adsorbed)
3.
H2(g) → H2 (adsorbed)
4.
H2 (adsorbed) → 2 H (adsorbed)
5.
N (adsorbed) + 3 H(adsorbed)→ NH3 (adsorbed)
6.
NH3 (adsorbed) → NH3 (g)
Reference
Ammonia production - Wikipedia, the free encyclopedia. (n.d.). Wikipedia, the free encyclopedia. Retrieved May
14, 2011, from http://en.wikipedia.org/wiki/Ammonia_production
Encyclopedia Britannica Online School Edition. (n.d.). Encyclopedia Britannica Online School Edition. Retrieved
May 14, 2011, from http://school.eb.com/eb/article-277711
Mascetta, J. A. (2010). Barron's SAT Chemistry
. New York: Barron's Education Series. (Original work published
1969)
Mustoe, F. J. (2002). McGraw-Hill Ryerson chemistry 12
. Toronto: McGraw-Hill Ryerson.