Question 15.5: Finding Equilibrium Constants from Experimental Concentratio...

1. Using the balanced equation as a guide, prepare an ICE table showing the known initial concentrations and equilibrium concentrations of the reactants and products. Leave space in the middle of the table for determining the changes in concentration that occur during the reaction. If initial concentrations of some reactants or products are not given, you may assume they are zero.

CO(g) + 2 H_2(g) \rightleftharpoons CH_3OH(g)

2. For the reactant or product whose concentration is known both initially and at equilibrium, calculate the change in concentration that occurs.

CO(g) + 2 H_2(g) \rightleftharpoons CH_3OH(g)

3. Use the change calculated in Step 2 and the stoichiometric relationships from the balanced chemical equation to determine the changes in concentration of all other reactants and products. Since reactants are consumed during the reaction, the changes in their concentrations are negative.
Since products are formed, the changes in their concentrations are
positive.

CO(g) + 2 H_2(g) \rightleftharpoons CH_3OH(g)

4. Sum each column for each reactant and product to determine the equilibrium concentrations.

5. Use the balanced equation to write an expression for the equilibrium
constant and substitute the equilibrium concentrations to calculate K.

K_c=\frac{[CH_3OH]}{[CO][H_2]^2}

=\frac{0.35}{(0.15)(0.30)^2}

= 26