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Our text book states use of catalyst such as iron oxide with small amounts of $\ce{K2O}$ and $\ce{Al2O3}$ ( these act as promoters) increases the rate of attainment of equilibrium in Haber's Process.

Why do we want to attain equilibrium? We want more production of ammonia so attaining equilibrium will not be efficient right?

Haber’s process is done at high pressure and low temperature (Le Chatelier’s Principle) to produce more output and increase forward rate of reaction. Here we do not wish to attain equilibrium.

When equilibrium is attained the rate of forward and backward rate becomes equal. If we want to maximise the output then the reaction should not reach equilibrium. According to my reasoning I feel that the reaction should go in forward direction and produce more of the products and simultaneously the product must be removed so as to facilitate more production otherwise it will result in backward reaction and when equilibrium is attained no more product shall be formed.

Why should we add catalyst and help the reaction attain equilibrium faster?

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    $\begingroup$ Catalysis is about reaching equilibrium faster. Catalysis does not provide going behind the equilibrium. The equilibrium is chosen by reaction conditions, like p, T, composition. The chosen T is a trade off between reaction thermodynamics (prefers lower T) and kinetics ( prefers higher T).// Your own reasoning – based on searching, reading and thinking – is supposed to be present to avoid the question closure for lack of own explicit effort. How do I ask a good question. $\endgroup$
    – Poutnik
    Commented Aug 11, 2021 at 7:10
  • $\begingroup$ IF the reaction doesn't approach equilibrium, the products will be hydrogen and nitrogen, the same as the inputs. $\endgroup$
    – matt_black
    Commented Aug 11, 2021 at 8:41
  • $\begingroup$ In fact, the reaction does not reach equilibrium even with the catalyst due short contact time. NH3 is separated after the catalyst passage and H2+N2 returns by a loop to the reaction for another catalyst path. For chemical engineering is very applicable the Pareto "80-20' rule, saying the 80% of result is achieved by 20% of effort. This is applied repeatedly, instead going of behind 80% result by enormous effort increase. $\endgroup$
    – Poutnik
    Commented Aug 11, 2021 at 9:09
  • $\begingroup$ Why should we add catalyst and help the reaction attain equilibrium faster? Because without catalyst, the reaction is very slow, At very high temperature, where it is not slow, equilibrium is shifted too much toward reactants. So you would need more energy with less yield. $\endgroup$
    – Poutnik
    Commented Aug 11, 2021 at 9:14
  • $\begingroup$ Please correct me if I’m wrong. Does this mean that we add catalysts to allow this specific reaction to take place at higher temperatures with better efficiency and also adding of catalyst does not affect this reaction’s equilibrium point as it is never reached ? $\endgroup$
    – curiosity
    Commented Aug 11, 2021 at 9:19

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Generally, catalysts never affect the equilibrium point, they just speed up reaching it from either side.

For this specific reaction,

$$\ce{N2(g) + 3 H2(g) <=> 2 NH3(g)}$$

it is an exothermic process, so higher temperatures speed up reactions, but provide lower yield at equilibrium.

Considering the reaction yield, it is advantageous:

  • Reaching equilibrium as close as possible to get a high yield at the single catalyst pass.
  • Having the equilibrium point shifted toward products.

The latter needs relatively low temperature, where reaction goes very slowly, but is accelerated by the catalyst.

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