The Haber process is used to manufacture ammonia. Explain the optimum conditions for this reaction and why these conditions may not be used in industry

N2(g)+3H2(g) ⇋ 2NH3(g) ΔH=-92kJmol-1 (equation for question above)
According to Le Chatelier's principle, a system in dynamic equilibrium will shift the position of equilibrium to minimise the change made to it. In this case, optimum conditions would be high pressure and low temperature. High pressure would mean that the system would favour the side that produces the least amount of moles, thereby decreasing the pressure, which in this case is the right hand side.Also, due to the fact the forward reaction is exothermic, the system would favour this side with a lower temperature, because the system wants to negate the change by favouring the exothermic reaction.
However, in industry, these conditions may not be the best options. A high pressure is not only hazardous, but it also requires a large amount of energy to create and maintain. Whilst, a low temperature will cause a lower rate of reaction compared to a higher temperature, due to less collisions occurring. This is why, in industry, they would most likely favour a lower pressure and a higher temperature compared to the optimum conditions.

LA
Answered by Luke A. Chemistry tutor

5301 Views

See similar Chemistry A Level tutors

Related Chemistry A Level answers

All answers ▸

What are 3 characteristics of Benzene that go against the proposed Kekule model?


Briefly describe the nature of three types of intramolecular bonding and two types of intermolecular bonding (drawings encouraged)


Which is more reactive, an alkane or an alkene and why?


What affects the boiling point of an alkane and why?


We're here to help

contact us iconContact ustelephone icon+44 (0) 203 773 6020
Facebook logoInstagram logoLinkedIn logo

© MyTutorWeb Ltd 2013–2025

Terms & Conditions|Privacy Policy
Cookie Preferences