Industrial chemistry Kazem.R.Abdollah (Asiaban) The Haber Process & The Ostwald Process 1

The Haber Process In the early 1900’s a German chemist called Fritz Haber came up with his chemical process to make ammonia using the “free” very unreactive Nitrogen from the air. (N 2 is 80% of atmosphere) Fritz Haber, 1918

The Haber process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic. Nitrogen + Hydrogen Ammonia

Raw Materials N 2 (g) is taken from the air via a process of fractional distillation. H 2 (g) comes from natural gas, CH 4 (g) CH 4 (g) + H 2 O (g) 3H 2 (g) + CO (g) The carbon monoxide then reacts with more steam: CO (g) + H 2 O (g) H 2 (g) + CO 2 (g)

scheme for the Haber process

The temperature Equilibrium considerations You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of ammonia in the equilibrium mixture. The forward reaction (the production of ammonia) is exothermic. According to Le Chatelier's Principle, this will be favoured if you lower the temperature. However, °C isn't a low temperature!

The compromise The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium °C is a compromise temperature producing a reasonably high proportion of ammonia in the equilibrium mixture (even if it is only 15%), but in a very short time.

The pressure According to Le Chatelier's Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. That will cause the pressure to fall again. Increasing the pressure brings the molecules closer together. In this particular instance, it will increase their chances of hitting and sticking to the surface of the catalyst where they can react. The higher the pressure the better in terms of the rate of a gas reaction.

Economic considerations Very high pressures are very expensive to produce on two counts. You have to build extremely strong pipes and containment vessels to withstand the very high pressure. That increases your capital costs when the plant is built. High pressures cost a lot to produce and maintain. That means that the running costs of your plant are very high. 200 atmospheres is a compromise pressure chosen on economic grounds.

The catalyst The catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't produce any greater percentage of ammonia in the equilibrium mixture. Its only function is to speed up the reaction. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor. The catalyst is actually slightly more complicated than pure iron. It has potassium hydroxide added to it as a promoter - a substance that increases its efficiency.

Ostwald process The Ostwald process is a chemical process for making nitric acid (HNO 3 ). Wilhelm Ostwald developed the process, and he patented it in 1902.

Description Ammonia is converted to nitric acid in 2 stages. It is oxidized (in a sense "burnt") by heating with oxygen in the presence of a catalyst such as platinum with 10% rhodium, to form nitric oxide and water.

Stage 1 This step is strongly exothermic, making it a useful heat source once initiated: 4 NH 3 (g) + 5 O 2 (g) → 4 NO (g) + 6 H 2 O (g) (ΔH = −905.2 kJ)

Stage 2 Stage two encompasses two reactions and is carried out in an absorption apparatus containing water. Initially nitric oxide is oxidized again to yield nitrogen dioxide: This gas is then readily absorbed by the water, yielding the desired product (nitric acid, albeit in a dilute form), while reducing a portion of it back to nitric oxide: 2 NO (g) + O 2 (g) → 2 NO 2 (g) (ΔH = −114 kJ/mol) 3 NO 2 (g) + H 2 O (l) → 2 HNO 3 (aq) + NO (g) (ΔH = −117 kJ/mol) Alternatively, if the last step is carried out in air: 4 NO 2 (g) + O 2 (g) + 2 H 2 O (l) → 4 HNO 3 (aq)

The NO is recycled, and the acid is concentrated to the required strength by distillation. Typical conditions for the first stage, which contribute to an overall yield of about 98%, are: pressure between 4 and 10 atmospheres (approx kPa or psig) and Temperature is about 500 K (approx. 217 °C or °F.).°C°F