The manufacture of ammonia: the Haber processThe Haber process, in which ammonia is synthesised from nitrogen and hydrogen, is one of the most important chemical processes to have had an impact on human civilisation. The growth of plants depends on the availability of a source of nitrogen in the soil in a form that the plants can use.Useable forms of nitrogen include soluble ammonium and nitrate salts and urea. N2 in the atmosphere needs to be converted into compounds that plants can use to promote growth. This process, called nitrogen fixation, is carried out in nature by nitrogen-fixing bacteria.In earlier times, crop rotation was an essential part of regenerating soil with adequate levels of inorganic nitrogen to support the growth of most crops. However, as populations increased, greater quantities of food were required and the use of nitrogenous fertilisers became essential. At the beginning of the twentieth century, the main supply of fertilisers came from natural deposits of saltpetre (potassium nitrate) and guano (bird droppings that have accumulated over thousands of years). During the early years of the twentieth century, the worldwide demand for nitrogen-based fertilisers was far greater than available supplies.Fritz HaberIn 1912, Fritz Haber, a German chemist, developed a process for the synthesis of ammonia from nitrogen and hydrogen. So important was this process that it undoubtedly had a significant influence on world history. At the time of the First World War, Germany was highly dependent on overseas supplies of nitrate for agriculture and the manufacture of explosives. The naval blockade of Germany by allied forces blocked this supply route for nitrate and other important materials required for the war effort. The Haber synthesis of ammonia facilitated the manufacture of fertilisers for continued food production, and nitric acid, an essential component in the production of explosives and other ammunition. Carl Bosch assisted in taking Habers process from laboratory production to full-scale industrial production.In fact, the synthesis of ammonia is a classic example of the influence of society upon chemistry and of the impact of chemistry on human life.

Uses of AmmoniaAmmonia ranks second to sulphuric acid in terms of quantity produced worldwide per year. It is used to make: Fertilisers (sulfate of ammonia, ammonium nitrate, urea) Fibres and plastics (rayon, acrylics, nylon) Nitric acid, which in turn is used to make fertiliser, dyes, fibres and plastics and explosives such as ammonium nitrate, TNT, (trinitrotoluene) and nitroglycerine Household cleaners Detergents (non ionic ones)INDUSTRIAL SYNTHESIS OF AMMONIAThe synthesis of ammonia uses the simple exothermic reaction:N2(g) + 3H2(g) 2NH3(g) H = -92kJ/molThis is an equilibrium reaction which at ordinary pressures and temperatures lies well to the left.EQUILIBRIUM CONSIDERATIONSLe Chateliers principle shows how to maximise the conversion of nitrogen and hydrogen to ammonia.1. If the pressure on a reaction system is increased, the equilibrium moves in the direction which tends to reduce the pressure.Look at the reactionN2(g) + 3H2(g) 2NH3(g) 4 moles of gas 2 moles of gasIf the pressure is increased then the reaction moves to the right.2. If the temperature is lowered, the equilibrium will move in the direction which tends to increase temperature.This reaction is exothermic, so if temperature is lowered it will move towards the right.On equilibrium considerations alone the reaction should be conducted at high pressure and low temperature.RATE CONSIDERATIONSAnother consideration is how long it will take to reach equilibrium, that is, the rate of reaction.As for most reactions, the rate decreases as temperature decreases.If we lower the temperature in order to produce more ammonia, we make the reaction very slow and so it takes a very long time to reach equilibrium.The rate of reaction can be increased by using a suitable catalyst. Iron is a good catalyst for this reaction but while it speeds it up it is still to slow at room temperature.Remember catalysts speed up reactions they do not affect the position of the equilibrium as they speed up both the forward and reverse reactions.Hence have these situations: Low temperatures produces a high yield but it takes a very long time (weeks to months) even with a catalyst High temperatures causes equilibrium to be reached quickly but the yield is extremely lowCOMPROMISEA moderate temperature produces a moderate yield moderately quickly.Typical conditions for the industrial process, called the Haber Process are: A temperature of about 700 K (or about 400C) and A total pressure of about 2.5 x 104 kPa (250 times standard atmosphere) Catalyst is magnetite Fe3O4 with its surface layer reduced to free ironWith a reactant mixture having H2 and N2 in the ratio of 3:1, these conditions give an equilibrium conversion to ammonia of about 45%.

THE HABER PROCESS

Reactants pass through the catalytic reactor, The mixture is cooled to condense out the ammonia formed Unreacted gases are fed back into the catalyst chamber along with incoming fresh reactants None of the reactant mixture is wasted Energy must be managed to prevent catalyst overheating Heat released in reaction is used to partially heat reactants ( saves energy and costs)

THE SOURCE OF REACTANTSNitrogen can be obtained from the atmosphere so hydrogen is the difficult or expensive reactant to obtain.Industrially, hydrogen is generally produced by reacting methane or some other hydrocarbon with steam in the presence of a nickel catalyst at a temperature of about 750C:CH4 (g) + H20 (g) CO (g) + 3H2 (g) H = +206 kJ/molCarbon monoxide poisons the iron catalyst and must be removed.CO (g) + H2O (g) CO2 (g) + H2 (g)Which has an added advantage of producing more hydrogen.The catalyst used is either Fe3O4 at 500C or Cu at 250CThere must not be any oxygen in the gas mixture as it is explosive with hydrogen under the conditions used. Methane is used to remove oxygen from air (normal combustion to CO2 and steam)The only unwanted gas in the mixture is carbon dioxide. This is removed by reaction with a base.Use of a 3:1 mixture of hydrogen and nitrogen is the most efficient way to make ammonia.

MONITORINGBecause many different conditions must be maintained for efficient and safe operation of the Haber process, monitoring is essential. The conditions that need to be monitored include: Temperature and total pressure in the reaction vessel to keep within the optimum range Prevent damage to catalyst Ratio of H2 to N2 in the incoming gas stream to prevent build up of one reactant Concentrations of O2, CO, CO2 and sulfur compounds explosions, contamination of catalyst Concentrations of methane and argon lower efficiency of conversion Purity of product ammonia no contamination.