A student ran the following reaction in the laboratory at 751 K: N2(g) + 3H2(g) 2NH3(g) When she introduced 3.47×10-2 moles of N2(g) and 6.38×10-2 moles of H2(g) into a 1.00 liter container, she found the equilibrium concentration of H2(g) to be 6.25×10-2 M. Calculate the equilibrium constant, Kc, she obtained for this reaction.

Kc = 4.86×10⁻⁶

Explanation:

We begin from the equation:

N₂  +  2H₂ ⇄  2NH₃

We start from 3.47×10⁻² moles of N₂(g) and 6.38×10⁻² moles of H₂(g), so when we reach the equilibrium, we get 6.25×10⁻² moles of H₂.

This data means, that in the reaction we made react:

6.38×10⁻² - x = 6.25×10⁻²

x = 1.3×10⁻³ moles of H₂

As stoichiometry is 1:3, we will know that the moles of N₂ that have been reacted were:

1.3×10⁻³ moles / 3 = 4.33×10⁻⁴ moles of N₂

So, in the equilibrium we would have:

3.47×10⁻² moles of N₂ - 4.33×10⁻⁴ moles of N₂ = 0.0343 moles of N₂

How many ammonia, would we have in the equilibrium?

4.33×10⁻⁴ mol . 2 = 8.66×10⁻⁴ moles (from stoichiometry with N₂, 1:2)

(1.3×10⁻³ mol . 2) / 3 = 8.66×10⁻⁴ moles (from stoichiometry with H₂, 2:3)

Let's make the expression for Kc

Kc = [NH₃]³ / [N₂] . [H₂]²

(8.66×10⁻⁴ )³ / (0.0343 . (6.25×10⁻²)² = 4.86×10⁻⁶

c is the answer because when something is left after something evaporates that shows the chemical change.