Chemical reaction:
2H2S (g) ⇄ 2H2(g) + S2(g)
The equilibrium constant is given by: Kc = [H2]^2 * [S2] / [H2S]^2 = 1.67 * 10^ -7
The initial concentration is 0.0125 mol / 0.500 L = 0.0250 M
Make a table showing the initial concentrations, the change and the final concentrations of each species
2H2S (g) ⇄ 2H2(g) + S2(g)
start 0.0250M 0 0
change - 2x +2x + x
end 0.0250 - 2x 2x x
Kc = (2x)^2 (x) / (0.0250 - 2x)^2
Kc = 4x^3 / (0.0250 - 2x)^2
To solve that equation in an easy way you can assume that 2x is << 0.0250, which leads to
Kc = 4x^3 / (0.0250)^2 = 1.67 * 10^ -7
=> x^3 = 1.67 * 10^ -7 * 0.0250 / 4 = 2.6 * 10 ^-11
=> x = 2.97 * 10^ -4 M
With this you can check that your assumption that x << 0.0250 is good and continue.
From the table you know that the concentrations at equilibrium are:
[H2] = 2x = 2 * 2.97 * 10 ^ -4 M = 5.94 * 10 ^ -4
[S2] = x = 2.97 * 10^ -4 M