Lesson Notes By Weeks and Term v5 - Grade 12

Lesson Video

This page supports the lesson note with a companion video and a short classroom-ready summary.

For class groups and homework, share this lesson page so learners also get the summary, objectives, and full lesson context.

Performance objectives

• Describe the Haber process for ammonia synthesis.
• Explain the Ostwald process for nitric acid production.
• Apply equilibrium principles to optimise industrial processes.
• Understand the role of catalysts in these reactions.

Lesson summary

Fertilizers are vital to South Africa's agricultural sector and play a crucial role in food security. Understanding the chemical processes involved in fertilizer production, especially the Haber process and the Ostwald process, allows us to appreciate the intricate balance between industrial chemistry and environmental sustainability.

Furthermore, analyzing the economic impact of fertilizer production and its environmental consequences equips us to make informed decisions about resource management and sustainable agriculture practices within our country.

Lesson notes

2.1 The Haber Process: Ammonia Synthesis The Haber process is an industrial process for the production of ammonia (NH 3 ) from nitrogen (N 2 ) and hydrogen (H 2 ). It is named after the German chemist Fritz Haber, who developed the process in the early 20th century. Ammonia is a crucial ingredient in the production of nitrogenous fertilizers.

Chemical Equation: N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g) ΔH = -92 kJ/mol This reaction is exothermic, meaning it releases heat. It is also a reversible reaction, meaning it can proceed in both directions (forward and reverse).

Reaction Conditions: Temperature: A moderate temperature of around 400-450°C is used. Lower temperatures favor the formation of ammonia because the reaction is exothermic.

However, the rate of reaction is too slow at lower temperatures. A compromise is therefore necessary.

Pressure: High pressure, typically 200-300 atmospheres, is used. High pressure favors the formation of ammonia because there are fewer moles of gas on the product side (2 moles) than on the reactant side (4 moles) according to Le Chatelier's principle. Increasing pressure shifts the equilibrium to the side with fewer moles of gas.

Catalyst: An iron catalyst (usually with potassium and aluminum oxide as promoters) is used to speed up the reaction. Catalysts lower the activation energy of a reaction without being consumed themselves. The catalyst allows the equilibrium to be reached faster.

Le Chatelier's Principle: Le Chatelier's principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. The "stress" can be a change in temperature, pressure, or concentration.

Temperature: As the reaction is exothermic, decreasing the temperature favours the forward reaction, leading to more ammonia.

Pressure: Increasing the pressure favours the forward reaction as there are fewer moles of gas on the product side.

Concentration: Increasing the concentration of reactants (N 2 and H 2 ) favours the forward reaction, producing more ammonia. Similarly, decreasing the concentration of ammonia favours the forward reaction. In practice, ammonia is continuously removed from the reactor to shift the equilibrium to the right.