# Question #9c357

##### 1 Answer

#### Explanation:

As you know, the equilibrium constant is calculated by taking the ratio between the *equilibrium concentrations* of the products and the *equilibrium concentrations* of the reactants, each raised to the power of their respective **stoichiometric coefficients**.

For this reaction

#color(red)(2)"SO"_text(3(g]) rightleftharpoons color(blue)(2)"SO"_2 + "O"_2#

the equilibrium constant has the following expression

#K_c = ( ["SO"_2]^color(blue)(2) * ["O"_2])/(["SO"_3]^color(red)(2))#

Rearrange the equation to solve for

#["O"_2] = (["SO"_3]^color(red)(2))/(["SO"_2]^color(blue)(2)) * K_c#

#["O"_2] = ( (4.46 * 10^(-2))^2 color(red)(cancel(color(black)("M"^2))))/((2.92 * 10^(-2))^2 color(red)(cancel(color(black)("M"^2)))) * 2.90 * 10^(-2)#

#["O"_2] = 2.333 * 2.90 * 10^(-2) = color(green)(6.76 * 10^(-2)"M")#

Now, does this result make sense?

Notice that in the expression for **equal** equilibrium concentrations for

#["O"_2] = K_c#

Any different in magnitude between the equilibrium concentrations of sulfur trioxide and sulfur dioxide will be **amplified** by their respective stoichiometric coefficients.

This tells you that anytime the equilibrium mixture contains more **greater** than the value of