The Haber process is used to produce ammonia. (Insert equation here) Explain the optimum conditions for this reaction and why these may differ from the conditions used in industry.
N2 (g)+3H2 (g) ⇋ 2NH3 (g) ΔH=-92kJmol-1 (this is the equation for the above question)
According to Le Chatelier's principle, the position of equilibrium will shift to minimise the change made to the conditions. In this case, to shift the equilibrium to the right hand side, the optimum conditions would be high pressure and low temperature. High pressure means that, the system will shift to minimise the pressure change, by favouring the side that produces the least amount of moles, which is the right side. Also, because the reaction is exothermic, the decrease in temperature will cause the system to shift to favour the exothermic reaction, which, is also the desired product.
However, these conditions will not be the best to use in industry. Even though high pressure will also increase the rate of reaction, it can be very hazardous and hard to control, and will also require a large amount of energy. With low temperature, the rate of reaction will be lower than at higher temperatures, because less collisions will happen in a given time. Due to these problems, in industry the conditions are more likely to be a higher temperature and a lower pressure than what is optimum for the system.
Related Chemistry A Level answers
All answers ▸Using your knowledge of periodicity and atomic structure, why does the first ionisation energy decrease moving down a group yet increase moving along a period in the periodic table?
What factors affect ionisation energy and how does each of them affect it?
Explain the relative resistance to bromination of benzene compared to phenol and compared to cyclohexene.
How can you tell if a reaction is endothermic or exothermic? Describe a way of determining if a reaction is exothermic or endothermic using simple laboratory equipment.