The Haber Process

Author: Rohit Shashi


Introduction


The Haber process is an industrial method of producing ammonia from a reaction between hydrogen and nitrogen. Under the right conditions of pressure and temperature, nitrogen and hydrogen can be combined to form ammonia.
The Haber process is also known as the Haber-Bosch process because the process was invented and developed by Fritz Haber and Carl Bosch, who were both German scientists. The process has its origins in the World War I when there was a shortage of ammonia around the world. Ammonia was crucial in the production of food for people during the war, so the shortage was a major problem. In 1908, Fritz Haber discovered the ideal conditions for the formation of ammonia, and Carl Bosch studied the effects of high pressure on chemical reactions. Together, they perfected the Haber-Bosch industrial process in 1913.
Nitrogen for the reaction is obtained from the air, and hydrogen gas is acquired through mainly methane gas. The reaction between nitrogen and hydrogen is reversible, which shall be explained in a further section. The reaction is also exothermic, which means that the reaction releases energy. The Haber process also requires a catalyst, more specifically an iron (Fe) catalyst. The effects of catalysts on the reaction shall also be explained further in a later section. So, the general reaction for the Haber process is as following: habereq.gif
The nitrogen and hydrogen combine to form NH3, or ammonia.


History & Development


The Haber Process was developed during World War I to mass produce ammonia. The process was created for supplying nitrogen for explosives and ammonia for fertilizers. The Allies had blocked all trade routes going to and from Germany, so the Germans lost their sources of sodium nitrate and potassium nitrate, which were sources of nitrogen. So, the German scientist, Fritz Haber, developed the Haber process to obtain nitrogen from the air, which had 80% nitrogen. The Haber Process took nitrogen from the air and combined it with molecular hydrogen to produce ammonia.
Nitrogen was also used for fertilizers crucial in the production of food. The shortage of regular nitrogen sources created problems in the manufacturing of fertilizers. However, the Haber Process created a source of nitrogen and ammonia for making fertilizer. As a trivial fact, it takes only 1 percent of the world’s energy to make 500 tons of artificial fertilizer per year which is used to help feed 40 percent of the ’s population.

Fritz Haber

Fritz Haber (1868 – 1934) was a true German patriot. He believed that a scientist should do whatever is in his/her power in a time of war. He is the Father of Chemical Warfare. He was the first scientists to think of using poisonous gas for combat use. Haber first used chlorine gas against French troops in 1915, during which 10,000 were killed by the gas within a few minutes. Later on, he created other chemical weapons such as Phosgene and Mustard gas. Despite all the contributions Haber has done for his country, his chemist wife is against all the terrible work Haber was doing. She hated it so much that she shot herself in the head; however, Haber only continued with his patriotic work.

Later, after creating a method to attain capture nitrogen from air, he won a Nobel Prize in 1918 while he was still viewed as a war criminal to a lot of countries. He kept on creating chemical weapons, which unknowingly would become a weapon used by the Nazis during the holocaust – ironically, Haber is Jewish. 15 years later, he fled with his own will to Switzerland before they took him away. There he lived in exile for the rest of his life.


The Haber Process


In very basic terms, the Haber Process obtains nitrogen gas from air and hydrogen gas from natural gas, or methane, and combines them to produce ammonia, or NH3. This reaction is exothermic, which means that the reaction releases energy from the system into the environment. In terms of chemistry, this exothermic reaction denotes that the sum of the enthalpies of the reactants is greater 9than the enthalpies of the products. Enthalpy is defined as the energy in a system. The equation for this reaction is: N2(g) + 3H2(g) → 2NH3(g) ΔH=-92.4 kJ/mol (forward reaction)
The reaction is also a reversible reaction: 2NH3(g) → N2(g) + 3H2(g) ΔH=+ 92.4 kJ/mol (backward reaction)
The following flowchart will help explain the process in slightly more detail:

haberprocess_flowchart.gif


First, methane and steam combine to produce hydrogen and carbon monoxide. This reaction is the source of hydrogen. Then, that hydrogen combines with oxygen in the air to produce water. Nitrogen gas is also released and this serves as the source of nitrogen. Finally, the nitrogen and hydrogen react to form ammonia. This reaction takes place under high pressure and temperature, and includes an iron oxide catalyst.
haberflow.gif

The following sections will be about Le Chatelier’s Principle and reversible reactions, which are heavily incorporated in the Haber Process.

Reversible Reactions


When a reaction is referred to as being reversible, it means that the reaction can go either forwards or backwards. The forward reaction is the desirable reaction because this is where the products are created. The backward reaction is the reaction when the products turn into their original reactants. Both these reactions occur at the same time. For the Haber Process, I have mentioned above which reactions are forward and backwards.
Next, a closed system is one where none of the reactants or products has any interaction with the outside environment. So, nothing can escape and nothing can go into the system. After a while in a closed system, the equilibrium mixture is reached. Here a specific proportion of the mixture exists as products and as reactants.
When this equilibrium is reached, the reactions are still going on. It may seem like there are no reactions occurring because nothing is changing. However, the forward reaction is producing the same amount of products as the backward reaction is producing reactants. This occurrence is known as dynamic equilibrium. This is to notify that the equilibrium is still reacting because dynamic means “moving” or “changing.”

Le Chatelier's Principle



Le Chatelier’s Principle states that a change in the reaction conditions will bring on a change in the equilibrium position to counterbalance each other. For example, if a product is removed, the equilibrium position will change to make more of that product. This is usually more desirable than having to remove the reactant and the equilibrium position changing to make more of the reactant.
In reactions, heat and pressure may be treated as products or reactants. Heat may be a reactant in exothermic reactions and a product in endothermic reactions.
So, if an exothermic reaction is cooled down, heat is removed and more of the product will be produced. So, more heat and products will be produced in the equilibrium mixture. If an exothermic reaction is heated up, there will be less of the product will be produced in the equilibrium mixture.
Pressure is a factor in reversible reactions involving gases. So, increasing the pressure will increase the amount of products produced, and decreasing the pressure will decrease the amount of products produced.

The video below summarizes the Le Chatelier's Principle:


To relate the Haber Process with Le Chatelier’s Principle, one can already infer that lower temperature and higher pressure will increase the amount of ammonia that can be produced. However, this is not the true case. In the next section, he industrial conditions for the Haber Process shall be explained.



Industrial Conditions

Proportions of Hydrogen and Nitrogen

The equation for the Haber Process states that the ratio for the mixture going into the reactor should be 1 part nitrogen to 3 parts hydrogen measured in volume. This is supported by Avogadro’s Law, which simply states that an equal volume of gases at the same temperature and pressure contain an equal number of molecules. Therefore, there will be 1 molecule of nitrogen to every 3 molecules of hydrogen. If the Haber Process reaction did not follow these proportions given by the equation, there will be a waste of space in the reactor because there will be nothing for some molecules to react with.

Temperature

According to Le Chatelier’s Principle one might expect that cooler temperatures would yield the highest amount of ammonia in the equilibrium mixture. However, the Haber Process uses a temperature of 400 – 450°C in the reactors. The reason is quite simple. It is widely known that higher temperatures cause reactions to proceed at a faster rate. So, the Haber Process has to make a compromise between having a high concentration yield of ammonia and rate of the reaction. For example, a lower temperature would indeed give a high concentration of ammonia, but in a very long time. But, a higher temperature would give a lower concentration of ammonia, but in larger quantities and in shorter time. In the end, the temperature of 400 – 450°C is the best compromise.

Pressure

Again using Le Chatelier’s Principle, higher pressure would be ideal in producing high amounts of ammonia in the equilibrium mixture. The Haber Process uses a pressure of 200 atmospheres (atm). This is definitely a high pressure, but it is not the highest pressure possible. So, why is a higher pressure not used in the reactors? This can be explained in economic terms. Very high pressures require very expensive equipment. One would need very strong pipes and containment chambers to endure the high pressures. Also, it is very expensive to produce and maintain. As a result, 200 atm is the best compromise between high pressure and expensive costs.

Catalyst

A catalyst is used for the Haber Process for the only reason to speed up the reaction, not to affect the equilibrium position in any way. In fact, without the catalyst, the Haber Process proceeds so slowly that there is no sense in continuing the reaction without the catalyst. So, a catalyst works by decreasing the activation energy to speed up the reaction. The activation energy is the minimum energy required to start a reaction. For the Haber Process, the catalyst used is an iron oxide metal. Also, potassium hydroxide is added to the iron oxide catalyst to increase its efficiency. The iron catalyst is used to lower the activation energy so help break down the hydrogen and nitrogen molecules easily. The following is a graph showing the activation energy with and without the catalyst:
activationenergy_copy.gif

The purple curve shows the activation energy without the catalyst, and the blue curve shows the activation energy with the catalyst. Obviously, the activation energy is higher without the catalyst than it is with the catalyst.

Applications & Uses


The function of the Haber Process is to produce ammonia from nitrogen and hydrogen. Ammonia can then be mixed with nitric acid created from ammonia to create ammonium nitrate, which is a major fertilizer used by farmers.
Hydrogen is obtained by using methane and reacting it with steam. This creates carbon dioxide and hydrogen, which is separated and used for the Haber Process.
Nitrogen is obtained from the air, which is 80% nitrogen, through fractional distillation. Fractional distillation is a process that is based on the different boiling points of different substances. First, the mixture of air is heated up. Then, the substance with the lowest boiling point starts to boil and evaporate into a cooling jacket, which is then liquefied, and eventually separated into a container. Eventually, the nitrogen will be obtained and then later used for the Haber Process.

Below is a video of the Haber Process.




Nitric Acid

Nitric acid is created so that it can then be reacted with ammonia to create ammonium nitrate, which is an important fertilizer. First, ammonia and air, or oxygen, is reacted to create nitrogen monoxide and water.
ammonia + oxygen nitrogen monoxide + water.
4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)

Next, the nitrogen monoxide is reacted with more oxygen to create nitrogen dioxide.
nitrogen monoxide + oxygen nitrogen dioxide.
2NO(g) + O2(g) 2NO2(g)

Then, the nitrogen dioxide is reacted with oxygen and water to form nitric acid.
nitrogen dioxide + oxygen + water nitric acid.
4NO2(g) + O2(g) + 2H2O(l) 4HNO3(aq)
So, the overall equation for the reaction to obtain nitric acid is:
ammonia + oxygen nitric acid + water.
4NH3(g) + 8O2(g) 4HNO3(aq) + 4H2O(l)

Ammonium Nitrate or Fertilizer

Ammonium nitrate is one of the most widely used fertilizers. It is manufactured by reacting the nitric acid and the ammonia.
nitric acid + ammonia ammonium nitrate.
HNO3(aq) + NH3(g) NH4NO3(aq)
Ammonium nitrate has a high percentage of nitrogen, so it is very ideal for crops. Using this nitrogenous fertilizer causes crops to grow taller and healthier. So, there is a higher yield of crops, and therefore food will be more abundant and cheaper.
One of the main disadvantages of using ammonium nitrate is that it goes into water bodies. There, it takes effect by allowing water plants and algae to grow in more abundance. When they die in large numbers, there will be a lot of bacteria feeding on the dead plant material. This uses up the oxygen supply in the water, and marine life will suffer as a result.

Further Reading

For more details on Fritz Haber, the following website offers a great article on him: http://toxsci.oxfordjournals.org/content/55/1/1.full. For information on industrial processes that use the concepts of chemical equilibrium, you may go to this website: http://www.docbrown.info/page07/equilibria3.htm. The website has detailed descriptions on process such as the Haber process, the production of lime from limestone.

Sources:
  1. Fr3Education. (Producer). (2011). Haber process. Retrieved on May 5, 2011 from http://www.youtube.com/watch?v=c4BmmcuXMu8
  2. Interiano, N. (n.d.). Case Study: The Haber Process. UC Davis Chemwiki. Retrieved May 1, 2011 from: http://chemwiki.ucdavis.edu/Physical_Chemistry/Chemical_Equilibrium/Haber_Process
  3. Haber Process Research Site. (n.d.). Retrieved May 1, 2011 from: http://haberchemistry.tripod.com/
  4. Ammonia and the Haber Process. (n.d.). Retrieved May 1, 2011 from: http://www.ewart.org.uk/science/patterns/pat11.htm
  5. The Haber Process. (n.d.). Retrieved May 1, 2011 from: http://www.chemguide.co.uk/physical/equilibria/haber.html