The Haber process of producing ammonia
In the early 1900’s, a chemist named Fritz Haber developed a laboratory procedure to form ammonia, the process which now bears his name. The Haber process involves the reacting of nitrogen and hydrogen in temperatures of about 550 degrees Celsius and pressure of 355 atm, using an iron catalyst, to form ammonia. The equation is shown below:
3H2 (g) + N2 (g) 2NH3 (g)
In his original experiment, Fritz Haber obtained hydrogen by decomposing steam over hot coals and nitrogen from liquefied air. Carl Bosch developed Haber’s experiment into an industrial size process, using methane instead of steam, and obtaining nitrogen through fractional distillation of liquid air. The process takes place in a large reaction chamber and the remaining reactants are cycled through the process again, while the product, ammonia is liquefied and drained away.
The yield of Ammonia in the Haber process is reduced at higher temperatures due to Le Chatelier’s Principle. Le Chatelier’s Principle states that “The equilibrium position will respond to oppose a change in the reaction conditions”. This means an equilibrium will change accordingly if the total pressure of the reaction is increased or decreased, or if more product or reactant is added, or if the temperature of the reaction
is changed. In the Haber process, the nitrogen, hydrogen and ammonia are in equilibrium. The reaction is also an exothermic reaction, meaning it produces heat.
The higher the temperature in the reaction chamber, the less ammonia is produced, as the equilibrium shifts to the right to counteract the heat as it is an exothermic reaction, producing more hydrogen and nitrogen gas.
The Haber process itself is a delicate balancing act because of all the factors that affect the yield of ammonia. These include the reaction energy of the reactant, the reaction rate and the resulting equilibrium conditions. The reaction energy is the amount of energy needed to synthesize the hydrogen and nitrogen gas together to form ammonia. This is achieved by raising the temperature of the reaction, shifting the equilibrium to the left and producing the ammonia product. However, if the
temperature is raised too high, the equilibrium will shift and the ammonia will
decompose back into the reactants of hydrogen and nitrogen. The reaction rate is the time it takes for the reactants to form the product. This is also increased by increasing the temperature of the reaction and by the use of a catalyst, but again, if it is raised too high the equilibrium will shift and the ammonia will decompose into the reactants.
To increase the rate of reaction, a catalyst is used. In the Haber process, the catalyst used is a finely ground porous iron powder (usually Fe3O4), with a large surface area. It absorbs the nitrogen and hydrogen gases and they react with each other on the catalyst’s surface, producing ammonia.
If the pressure of the system is increased, the hydrogen and nitrogen gas molecules
are compressed together, and the equilibrium shifts to the left, forming ammonia molecules. Increased pressure also increases the reaction rate, as the gas molecules are more concentrated and closer together
The reaction vessel in the Haber process must be closely monitored for safety and economic reasons. The reactants must be pure, as any contaminants such as oxygen may cause an explosion within the chamber and the temperature, pressure and volume of reactants and product must be monitored to ensure that they stay at the recommended level to achieve 30% yield of ammonia and do not damage the catalyst. The catalyst must also be monitored to ensure it is not contaminated by contaminants in the reactants.
At the time of Haber’s discovery, the main source for ammonia and nitrogenous compounds needed for the production of fertilizer and explosives was a large guano deposit in Chile. However this natural resource was far from where it was needed and it would not last forever. Haber solved this problem when he found he was able to produce ammonia from nitrogen and hydrogen, two gases readily available in the atmosphere. With some modifications by Carl Bosch, the process became an industrially viable process for synthesizing ammonia for the war effort by Germany and for fertilizer production. During the war, most of Germany’s supply of nitrates for explosives and gas warfare got cut off, and the production of explosive and mustard gas slowed, Haber’s discovery was a major breakthrough. Germany was able to manufacture tones of explosives and mustard gas for the war effort that advanced their tactical position in the warfare and nearly won them the war.
The Haber process itself has few implications on the environment, however the product of the process, ammonia, does. The nitrogenous products formed from ammonia used in fertilizers easily leach out of soils into waterways, polluting the water, making it unsuitable for drinking and encouraging the growth of algae and water weeds, choking water supplies. This algae and plant growth eventually die and the bacteria that feeds on them uses up any oxygen in the water, killing any fish that live and breed in waterways. Nitrates also pollute drinking water, creating a health hazard, as nitrates interfere with oxygen transport within the blood.
The Haber process is still a viable industrial process for ammonia synthesis. However it is constantly being improved with the discovery of new catalysts that increase ammonia yield and the addition of a reforming exchanger that increases the rate of reaction. These improvements make the Haber process more efficient and more economically viable.