Case Study: The Haber Process
The Haber process is the production of ammonia from a reaction between nitrogen and hydrogen, because of an iron substitute. This process is known for the commercial synthesis of ammonia. There is great abundance of nitrogen in the air when it is combined with hydrogen under extreme pressure and high temperature. This process is a great example of chemical equilibrium.
The Haber process, also known as the Haber-Bosch Process, was founded by Fritz Haber and Carl Bosch, both who were German Chemists. Haber discovered the conditions for the formation of ammonia, and Bosch discovered the work of high-pressure on chemical reactions (developed into industrial process). Both were awarded the Nobel Prize. During the 1920’s, there was a shortage of the world's supply for fixed nitrogen. Nitrogen was mainly used for fertilizer. Fertilizer was used in order to produce food, so that in WWI people could continue to fight. It only requires 1 percent of the world's energy to make 500 million tons of artificial fertilizer per year, which, in turn, helps feed 40 percent of the world's population.
The Haber process takes nitrogen gas from air and combines it with molecular hydrogen gas to form ammonia gas. This is an exothermic reaction, meaning it releases energy so that the sum of the enthalpies of N2 and H2 (the reactants) is greater than the enthalpy of NH3 (the products).
N2(g) + 3H2(g) → 2NH3(g) ΔH=-92.4 kJ
which is a reversible reaction:
2NH3(g) → N2(g) + 3H2(g) ΔH=+ 92.4 kJ mol-1
Here's a visual to help convey the process:
From the flow chart above, we can see that methane and steam combine to form hydrogen and carbon monoxide, which in turn releases hydrogen. The hydrogen then combines with oxygen from the air to produce water. Finally, nitrogen gas is released which combines with hydrogen gas to form ammonia. This takes place under high pressure and temperature and with an iron catalyst*** (mentioned later on).
Le Châtelier's Principle
The Haber process incorporates Le Chatlier's Principle, which is a good example of equilibrium principles. Uses of Le Chatlier's Principle are reversible reactions and reversible reactions involving gases. Chemical equilibrium is when a reaction has no tendency to change the quantity of the products and reactants, so the reaction can go both ways.
Examples of Le Châtelier's Principle
N2 (g) + 3H2 (g) ↔ 2NH3 (g)
When you increase the pressure so that the least amount of molecules will be formed, there won't be an increase in collisions. However, if more gas molecules are formed, there will be an increase in collisions, thus moving the way that will produce the least amount of molecules.
PCl3(g) + Cl2(g) ↔ PCl5(g) + energy
H2+ I2 ↔ 2HI
A catalyst is used to speed up a reaction by lowering the activation energy. So, in this reaction, the iron catalyst is used to lower the activation energy so that the N2 and H2 can be easier to break down.
From the above diagram the blue curve represents what the activation energy would look like with an iron catalyst. The purple curve shows what the activation energy would look like if a catalyst was not involved. Note that without the catalyst, the activation energy is much bigger. A catalyst is needed to break down the nitrogen and hydrogen gases.
By mixing one part ammonia to nine parts air with the use of a catalyst, the ammonia will get oxidized to nitric acid.
4 NH3 + 5 O2 → 4 NO + 6 H2O
2 NO + O2 → 2 NO2
2 NO2 + 2 H2O → 2 HNO3 + H2
Ammonia is mixed with irrigation water to form a solution for fertilizer ingredients. Fertilizer is mainly used in fields that grow crops such as corn, barley, sorghum, rapeseed, soybean, and sunflower.
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