A reaction mixture initially contains 2.8 M #H_2O# and 2.6 M #SO_2#. How do you determine the equilibrium concentration of #H_2S# if Kc for the reaction at this temperature is #1.3 xx 10^-6#?

Question

A reaction mixture initially contains 2.8 M #H_2O# and 2.6 M #SO_2#.  How do you determine the equilibrium concentration of #H_2S# if Kc for the reaction at this temperature is #1.3 xx 10^-6#?

Luna Chen · Accepted Answer

#["H"_2"S"] = "0.12 M"#

The first thing to do here is write the equilibrium reaction

#2"SO"_ (2(g)) + 2"H"_ 2"O"_ ((g)) rightleftharpoons 2"H"_ 2"S"_ ((g)) + 3"O"_ (2(g))#

Now, you know that at a certain temperature, the equilibrium constant for this reaction is equal to

#K_c = 1.3 * 10^(-6)#

Right from the start, you can tell just by looking at the value of #K_c# that the equilibrium concentration of hydrogen sulfide, #"H"_2"S"#, will be lower than the equilibrium concentration of the two reactants.

This is the case because you have #K_c < 1#, which means that at this temperature you can expect the equilibrium mixture to contain more reactants than products.

The next thing to do here is use an ICE table to find the equilibrium concentration of hydrogen sulfide

#" "2"SO"_ (2(g)) " "+" " 2"H"_ 2"O"_ ((g)) rightleftharpoons 2"H"_ 2"S"_ ((g)) " "+" " 3"O"_ (2(g))#

#color(purple)("I")color(white)(aaaaacolor(black)(2.6)aaaaaaaaaaaacolor(black)(2.8)aaaaaaaacolor(black)(0)aaaaaaaaaaacolor(black)(0)# #color(purple)("C")color(white)(aaacolor(black)((-2x))aaaaaaaacolor(black)((-2x))aaaacolor(black)((+2x))aaaaaaacolor(black)((+3x))# #color(purple)("E")color(white)(aaacolor(black)(2.6-2x)aaaaaaacolor(black)(2.8-2x)aaaaaacolor(black)(2x)aaaaaaaaaacolor(black)(3x)#

By definition, the equilibrium constant for the reaction will be

#K_c = (["H"_2"S"]^2 * ["O"_2]^3)/(["SO"_2]^2 * ["H"_2"O"]^2)#

In your case, this expression is equivalent to

#K_c = ( (2x)^2 * (3x)^3)/( (2.6 - 2x)^2 * (2.8 - 2x)^2)#

#K_c = (4x^2 * 27x^3)/( (2.6 - 2x)^2 * (2.8 - 2x)^2)#

#K_c = (108x^5)/( (2.6 - 2x)^2 * (2.8 - 2x)^2) = 1.3 * 10^(-6)#

Now, because the value of #K_c# is so small compared with the initial concentrations of water vapor and sulfur dioxide, you can use the approximations

#2.8 - 2x ~~ 2.8" "# and #" "2.6 - 2x ~~ 2.6#

This will give you

#1.3 * 10^(-6) = (108x^5)/(2.6^2 * 2.8^2)#

which allows you to calculate #x# by

#x = root(5)( (1.3 * 2.6^2 * 2.8^2 * 10^(-3))/108) = 0.05767#

Keep in mind that because the equilibrium concentration of hydrogen sulfide is #2x#, you will have

#["H"_2"S"] = 2 xx "0.05767 M" = "0.11534 M"#

Rounded to two sig figs, the number of sig figs you have for the initial concentrations of sulfur dioxide and water vapor, the answer will be

#["H"_2"S"] = color(green)(|bar(ul(color(white)(a/a)color(black)("0.12 M")color(white)(a/a)|)))#

As predicted, the equilibrium concentration of hydrogen sulfide is lower than the equilibrium concentrations of the two reactants, which are

#["SO"_2] = 2.6 - 2 * "0.05767 M" = "2.5 M"#

#["H"_2"O"] = 2.8 - 2 * "0.05767 M" = "2.7 M"#

Kai Bowman · Answer

undefined To determine the equilibrium concentration of H₂S, you need additional information such as the balanced chemical equation for the reaction involving H₂O, SO₂, and H₂S. With this information, you can set up an ICE (Initial, Change, Equilibrium) table and use the given Kc value to calculate the equilibrium concentration of H₂S.