- Decrease temperature - this will decrease kinetic energy. Energy is on the right side, so the system will shift to make energy. Since the reaction is exothermic, this will favour ammonia (right side) - more ammonia will be produced. If the temperature is increased, the system will shift away from energy and more reactants will be produced and less ammonia will be produced.
-
Increase pressure / decrease the volume - the system will shift to reduce pressure. Since less particles = less pressure and knowing that the reactant side has 4 particle and the product side has 2, the system will shift right producing more NH3 (g). If the pressure is decreased or the volume is increase, the system will shift to increase pressure. Since more particles = more pressure and knowing that the reactant side has 4 particle and the product side has 2, the system will shift left producing more H2 (g) and N2 (g) and less NH3 (g).
Reversible Reactions
- The reaction of nitrogen and hydrogen is a reversible reaction meaning the reaction can proceed in either forward or the reverse direction depending on conditions.
- The forward reaction is exothermic reaction (produces heat) and is favored at low temperatures. Increasing the temperature tends to drive the reaction in the reverse direction.
- High pressure favours the forward reaction because there are fewer molecules on the right side. So the only compromise in pressure is the economical situation trying to increase the pressure as much as possible.
- The catalyst has no effect on the position of equilibrium, however it does increase the reaction rate. This allows the process to be operated at lower temperatures, which favors the forward reaction.
Factors Affecting the Process:
1. Temperature
-
At low temperatures, the reaction of N2 (g) and H2 (g) is not economical
- Adding heat increases the rate of reaction
- The higher the temperature, the lower the ammonia yield
The reaction of nitrogen and hydrogen at low temperatures is so slow that the process becomes uneconomical. Adding heat increases the rate of the reaction. However, the higher the temperature, the lower the yield of ammonia
For the process to be optimum it is important to balance the rate of the reaction (increased by increasing temperature) against the equilibrium of the reaction (pushed to the right by decreasing temperatures)
2. Catalyst:
It was observed by Haber that using an iron oxide catalyst eliminates the need for excessively high temperatures, allowing the equilibrium position to move quickly to the right at lower temperatures
Therefore for this process a catalyst (iron or osmium) is used. This speeds up the reaction by lowering the activation energy. This helps by breaking the N2 bonds and H2 bonds more readily.
3. Pressure:
According to Le Chetalier's Principle Increasing the pressure causes the equilibrium position to move to the right resulting in a higher yield of ammonia since there are more gas molecules on the left hand side of the equation (4 in total) than there are on the right hand side of the equation (2). Increasing the pressure means the system adjusts to reduce the effect of the change, that is, to reduce the pressure by having fewer gas molecules.
The equilibrium expression for this reaction is:
Keq = [NH3] 2 / [N2][H2] 3
Rate considerations:
- Increasing the temperature gives more reactant molecules having sufficient energy activation energy to overcome the energy barrier to reacting. Therefore the reaction is faster at higher temperatures (but a low yield of ammonia).
-
A temperature range of 400-500oC is a compromise designed to achieve an acceptable yield of ammonia (10-20%) within an acceptable time period.
-
At 200oC and pressures above 750atm there is an almost 100% conversion of reactants to the ammonia product.
- Since there are difficulties associated with containing larger amounts of materials at this high pressure, lower pressures of around 200 atm are used industrially.
-
By using a pressure of around 200atm and a temperature of about 500oC, the yield of ammonia is 10-20%, while costs and safety concerns in the building and during operation of the plant are minimized.
Industrial Production of Ammonia:
- For industrial production of ammonia the temperature must be 450ºC to 500ºC. Cooling the reaction down will result in the equilibrium mixture making richer ammonia but the rate of production will decrease since the temperature is lower.
- It is better to have 10 % ammonia fast rather then 90 % ammonia slow therefore the atmospheric pressure is raised. Industrial ammonia is produced at the atmospheric pressure of 100 atm because it is too expensive to make a high-pressure chemical plant. Running the reaction at 200 atm is the highest pressure with the greatest return value.
- With a reversible reaction, a catalyst, which increases the rate, will increase the rate of both the forward and the backward reaction. This is useful because the catalyst will, cause the reaction mixture to reach its equilibrium composition more quickly. The catalyst will not change the equilibrium composition of the substance.
Economic issues for industrial production:
For Haber process reaction vessel is made from reinforced steel and is usually 20 meters high with a mass up to 200 tons. The cost of making the reaction vessel is high. If even higher pressures have to be used then the height of the vessel will be even higher and would have to be even thicker and stronger thereby increasing the costs. Further running costs too will be higher for high pressure since reacting gases would have to be pumped at higher pressure. Therefore economies can be maintained at the optimum pressure of 200 atmosphere.
Yield of ammonia is higher when the temperature of reaction is lower but this also causes a slow rate of reaction. Slow rate of reaction increases the cost of manufacturing. The catalyst is used to increase the rate in both forward as well as backward reaction.
Uses:
The Haber process produces 500 million tons of artificial fertilizer per year. 1% of the world's energy supply is consumed in the manufacturing of that fertilizer which is responsible for sustaining 40% of the Earth's population. Ammonia is used for the following:
1. To manufacture Fertilizers:
-
Ammonium sulfate, (NH4)2SO4
Ammonia + Sulphuric acid ➔ Ammonium Sulphate
2NH3(aq) + H2SO4(aq) ➔ (NH4)2SO4(aq)
-
Ammonium phosphate, (NH4)3PO4
-
Ammonium nitrate, NH4NO3
Ammonia + nitric acid ➔ammonium nitrate
NH3 (aq) + HNO3 (aq) ➔ NH4NO3(aq)
-
Urea, (NH2)2CO.
2. Chemicals
-
Synthesis of nitric acid, HNO3,
-
Sodium hydrogen carbonate (sodium bicarbonate), NaHCO3
-
Sodium carbonate, Na2CO3
- Hydrogen cyanide (hydrocyanic acid), HCN
-
Hydrazine, N2H4 (used in rocket propulsion systems)
3. Explosives: such as ammonium nitrate, NH4NO3
4. Fibres & Plastics: such as nylon, [(CH2) 4-CO-NH-(CH2)6-NH-CO]-,and other polyamides
5. Refrigeration: for making ice, large scale refrigeration plants, air-conditioning units in buildings and plants
6. Pharmaceuticals: used in the manufacture of drugs such as sulfonamide, anti-malarials and vitamins such as nicotinamide (niacinamide) and thiamine. Production of barbiturates (sedatives), is made by the reaction of ammonia with carbon dioxide
7. Pulp & Paper: Ammonium hydrogen sulfite, NH4HSO3, enables some hardwoods to be used
8. Mining & Metallurgy: Ammonia is used in nitriding (bright annealing) steel,
used in zinc and nickel extraction.
9. Cleaning: Ammonia in solution is used as a cleaning agent such as in 'cloudy ammonia'
Environmental Hazards of Ammonia:
The benefit of using nitrogenous fertilizers is that it increases the growth of crops, and turn out to be healthier. In addition to more food, the crops can be sold at a cheaper rate. However, every good discovery has its disadvantages, and the Haber process is no exception.
- Overuse of ammonia fertilisers on fields can cause major environmental problems.
- Ammonium salts are water-soluble and get washed into the groundwater, rivers and streams by rain contaminating them with ammonium ions and nitrate ions. This contamination causes several problems.
- Excess fertilisers in streams and rivers cause eutrophication.
- Overuse of fertilisers causes increasing quantities dissolving in rainwater.
- This increases levels of nitrate or phosphate in rivers and lakes.
- This causes 'algal bloom' i.e. too much rapid growth of water plants on the surface where the sunlight is the strongest.
- This prevents light from reaching plants lower in the water.
- These lower plants decay and the active aerobic bacteria use up any dissolved oxygen.
- This means any microorganisms or higher life forms relying on oxygen cannot respire.
- All the eco-cycles are affected and fish and other respiring aquatic animals die.
- The river or stream becomes 'dead' below the surface as all the food webs are disrupted.
- Nitrates are potentially carcinogenic (cancer or tumor forming).
- The presence in drinking water is a health hazard. Rivers and lakes contaminated can be used as initial sources for domestic water supply.
Haber Process Summary:
References:
Wikipedia: Catalyst: Retrieved from www on 21st Jan 2006 URL:
Uses of Ammonia: Retrieved from www on 20th Jan 2006 URL:
The Effect Of Catalysts On Reaction Rates: Retrieved from www on 18th Jan 2006 URL:
Haber Process: http: Retrieved from www on 18th Jan 2006 URL:
Haber Process: http: Retrieved from www on 18th Jan 2006 URL:
Haber Process: http: Retrieved from www on 17th Jan 2006 URL: