The reaction for the formation of gaseous hydrogen fluoride (HF) from molecular hydrogen (H2) and fluorine (F2) has an equilibrium constant of 1.15 x 102 at a certain temperature. H2 (g) + F2 (g) ⇋ 2 HF (g) In a particular experiment, 3.00 mol H2 and 6.00 mol F2 are mixed in a 3.00 L flask at this temperature and allowed to come to equilibrium. Given this information, calculate the concentrations of H2, F2, and HF present in the flask at equilibrium.

[H2]eq = 0.0129 M

[F2]eq = 1.0129 M

[HF]eq = 0.9871 M

Explanation:

H2 (g) + F2 (g) ↔ 2HF (g)

∴ Ke = [HF]² / [H2]*[F2] = 1.15 E2

experiment:

∴ n H2 = 3.00 mol

∴ n F2 = 6.00 mol

∴ V sln = 3.00 L

⇒ [H2]i = 3.00 mol / 3.00 L = 1 M

⇒ [F2]i = 6.00 mol / 3.00 L = 2 M

[ ]i change [ ]eq

H2 1 1 - x 1 - x

F2 2 2 - x 2 - x

HF - x x

⇒ K = (x) ² / (1 - x) * (2 - x) = 1.15 E2

⇒ x² / (2 - 3x + x²) = 1.15 E2 = 115

⇒ x² = (2 - 3x + x²) (115)

⇒ x² = 230 - 345x + 115x²

⇒ 0 = 230 - 345x + 114x²

⇒ x = 0.9871

equilibrium:

⇒ [H2] = 1 - x = 1 - 0.9871 = 0.0129 M

⇒ [F2] = 2 - x = 2 - 0.9871 = 1.0129 M

⇒ [HF] = x = 0.9871 M