Question #af614

May 4, 2017

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

${K}_{c} = \left(\left[\text{CO"] * ["H"_2]^3)/(["CH"_4] * ["H"_2"O}\right]\right)$

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 $\text{CO}$, ${\text{CH}}_{4}$, and $\text{H"_2"O}$, and $3$ for ${\text{H}}_{2}$.

For equilibrium partial pressures, you will have

${K}_{p} = \left(\left(\text{CO") * ("H"_2)^3)/(("CH"_4) * ("H"_2"O}\right)\right)$

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