\begin{align} \ce{CO(g) + 3H2(g) &->[Ni] CH4(g) + H2O(g)}\\ \ce{CO(g) + 2H2(g) &->[Cu/ZnO-Cr2O3] CH3OH(g) }\\ \ce{CO(g) + H2(g) &->[Cu] HCHO(g) } \end{align}

As we see here, using different catalysts in the reaction between Carbon monoxide and hydrogen yields different products. Is this in contradiction to the following description of the properties of catalysts?

However, it is very important to keep in mind that the addition of a catalyst has no effect whatsoever on the final equilibrium position of the reaction. It simply gets it there faster.


Adding a catalyst makes absolutely no difference to the position of equilibrium


If a catalyst is not supposed to affect the reaction's final equilibrium position how do we explain the catalyst selectivity seen here? I saw a similar question (Selectivity of catalysts) but it wasn't addressed directly at this principle (and unanswered still).

  • 4
    $\begingroup$ The quotes you have quoted are in contradiction with the title. Catalysts do not affect the position of equilibrium, they accelerate the respective ( and possibly selected ) forward and backward reactions. Acceleration just some reactions ( in both ways) is not against any principle. $\endgroup$
    – Poutnik
    Commented Sep 16, 2021 at 13:32
  • 2
    $\begingroup$ The question itself is probably not a bad one, but it needs a lot of work reformatting to make it more attractive. $\endgroup$ Commented Sep 16, 2021 at 19:09

2 Answers 2


Catalysts very much do affect end products because they may act differently on competing reactions. For instance, given ethylene and oxygen a suitable catalyst may promote formation of ethylene oxide and not as strongly promote oxidizing the ethylene to carbon dioxide and water. (In this particular case, a silver catalyst with carefully controlled temperature is used to maximize ethylene oxide yield.)

What a catalyst does not affect is the equilibrium product mixture for a particular reaction.


If a catalyst is not supposed to affect the reaction's final equilibrium position how do we explain the catalyst selectivity seen here?

If you wait long enough so that all three reactions attain equilibrium, the presence or absences of catalysts have no effect on the product mixtures.

In the examples, however, the reactions without catalysts are all slow. If you add one catalyst but not the other two, one reaction will reach equilibrium and you can work up the product mixtures, which will contain little or none of the other possible products.

The fancy way of describing this is to say that the product ratio is under kinetic control (rather than thermodynamic control). Here are two specific scenario where reaction conditions determine product identity: Kinetic and thermodynamic control of sulphonation of toluene and Thermodynamic vs Kinetic Sulphonation of Naphthalene


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