Question #af614

1 Answer
May 4, 2017

Answer:

Here's what I got.

Explanation:

You're dealing with a chemical reaction that involves gases, so you can write two equilibrium constants, one that uses the equilibrium concentrations of the species, #K_c#, and one that uses the equilibrium partial pressures of the species, #K_p#.

In both cases, you're going to be using the ratio that exists between the products and the reactants.

For equilibrium concentrations, you will have

#K_c = (["CO"] * ["H"_2]^3)/(["CH"_4] * ["H"_2"O"])#

Notice that the equilibrium concentration of each species is raised to the power of the stoichiometric coefficient that the species has in the balanced chemical equation, i.e. #1# for #"CO"#, #"CH"_4#, and #"H"_2"O"#, and #3# for #"H"_2#.

For equilibrium partial pressures, you will have

#K_p = (("CO") * ("H"_2)^3)/(("CH"_4) * ("H"_2"O"))#

This time. the partial pressure of each species is raised to the stoichiometric coefficient that the species has in the balanced chemical equation.