The initial concentration for the compounds involved in the reaction shown were determined to be [CS2(g)] = 1.075 mol/L, [H2(g)] = 1.436 mol/L, [CH4(g)] = 0.6740 mol/L, [H2S(g)] = 0.3120 mol/L. Calculate the value of the equilibrium constant (Kc) at 1175 K if the equilibrium concentration of CH4(g) was 0.7910 mol/L.
CS2(g)+4H2(g) = CH4(g)+2H2S(g)
CH4 is increased by .117
H2S is increased by 2x.117
H2 is decreased by 4x.117
CS2 is decreased by .117
final:
CH4=.791
HS2=.546
H2=.968
CS2=.958
(.791)(.546)^2/(.958)(.968)^4
That looks good to me.
To calculate the value of the equilibrium constant (Kc) at 1175 K using the given information, you need to use the equilibrium concentrations of the reactants and products.
First, you need to determine the changes in concentrations for each compound based on the changes mentioned in the question.
CH4 is increased by 0.117 mol/L, so the final concentration of CH4 is given as 0.791 mol/L.
H2S is increased by 2 times 0.117 mol/L, so the final concentration of H2S is calculated as 2 * 0.117 = 0.234 mol/L.
H2 is decreased by 4 times 0.117 mol/L, so the final concentration of H2 is calculated as 1.436 - (4 * 0.117) = 0.968 mol/L.
CS2 is decreased by 0.117 mol/L, so the final concentration of CS2 is calculated as 1.075 - 0.117 = 0.958 mol/L.
Now that you have the equilibrium concentrations, you can use the balanced equation to calculate the equilibrium constant (Kc).
The balanced equation is: CS2(g) + 4H2(g) = CH4(g) + 2H2S(g)
Kc = ([CH4] * [H2S]^2) / ([CS2] * [H2]^4)
= (0.791 * (0.234)^2) / (0.958 * (0.968)^4)
= 0.0357 / 0.8107
≈ 0.0441
Therefore, the value of the equilibrium constant (Kc) at 1175 K is approximately 0.0441.