Determine the value of ΔG^θ, in kJ, for the above reaction at 761 K using section 1 of the data booklet.

1.0 mol of N₂(g), 1.0 mol of H₂(g) and 1.0 mol of NH₃(g) are placed in a 1.0 dm³ sealed flask and left to reach equilibrium. At equilibrium the concentration of N₂(g) is 0.8 mol dm⁻³.

N₂(g) + 3H₂(g) $\rightleftharpoons$ 2NH₃(g)

What are the equilibrium concentration of H₂(g) and NH₃(g) in mol dm⁻³?

Calculate the standard Gibbs free energy change, ΔG^Θ, in kJ, for this reaction at 1000 K. Use sections 1 and 2 of the data booklet.

A mixture of 0.40 mol of CO (g) and 0.40 mol of H₂ (g) was placed in a 1.00 dm³ vessel. The following equilibrium was established.

CO (g) + 2H₂ (g)  CH₃OH (g)

At equilibrium, the mixture contained 0.25 mol of CO (g). How many moles of H₂ (g) and CH₃OH (g) were present at equilibrium?

Comment on the value of ΔG when the reaction quotient equals the equilibrium constant, Q = K.

Determine an approximate order of magnitude for K_c, using sections 1 and 2 of the data booklet. Assume ΔG^Θ for the forward reaction is approximately +50 kJ at 298 K.

A two-step mechanism is proposed for the formation of NO₂(g) from NO(g) that involves an exothermic equilibrium process.

First step: 2NO(g) $\rightleftharpoons$ N₂O₂(g)     fast

Second step: N₂O₂(g) + O₂ (g) → 2NO₂(g)     slow

Deduce the rate expression for the mechanism.

At 700 ºC, the equilibrium constant, K_c, for the reaction is 1.075 × 10⁸.

2H₂ (g) + S₂ (g) $\rightleftharpoons$ 2H₂S (g)

Which relationship is always correct for the equilibrium at this temperature?

A. [H₂S]² < [H₂]² [S₂]

B. [S₂] = 2[H₂S]

C. [H₂S] < [S₂]

D. [H₂S]² > [H₂]²[S₂]

Calculate the equilibrium constant, K_c, for this reaction at 298 K. Use your answer to (d)(iii) and sections 1 and 2 of the data booklet.

(If you did not obtain an answer to (d)(iii) use a value of 2.0 kJ mol⁻¹, although this is not the correct answer).

Calculate a value for the equilibrium constant, K_c, at 298 K, giving your answer to two significant figures. Use your answer to (f)(ii) and section 1 of the data booklet. 

(If you did not obtain an answer to (f)(ii), use −140 kJ mol⁻¹, but this is not the correct value.)

1.0 mol each of sulfur dioxide, oxygen, and sulfur trioxide are in equilibrium.

$2{\mathrm{SO}}_2 \left(\mathrm{g}\right)+{\mathrm{O}}_2 \left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{SO}}_3 \left(\mathrm{g}\right)$

Which change in the molar ratio of reactants will cause the greatest increase in the amount of sulfur trioxide?

Assume volume and temperature of the reaction mixture remain constant.

The equilibrium constant, K_c, has a value of 1.01 at 298 K.

Calculate ΔG^⦵, in kJ mol^–1, for this reaction. Use sections 1 and 2 of the data booklet.

SO₂ (g), O₂(g) and SO₃(g) are mixed and allowed to reach equilibrium at 600 °C.

Determine the value of K_c at 600 °C.

Phenylethene is manufactured from benzene and ethene in a two-stage process. The overall reaction can be represented as follows with ΔG^θ = +10.0 kJ mol⁻¹ at 298 K.

Calculate the equilibrium constant for the overall conversion at 298 K, using section 1 of the data booklet.

At a given time, the concentration of NO₂(g) and N₂O₄(g) were 0.52 and $0.10\text{ mol}\text{d}{\text{m}}^{-3}$ respectively.

Deduce, showing your reasoning, if the forward or the reverse reaction is favoured at this time.

Calculate the value of the equilibrium constant, K, at 298 K. Use sections 1 and 2 of the data booklet.

The graph shows Gibbs free energy of a mixture of N₂O₄(g) and NO₂(g) in different proportions.

N₂O₄(g) $\rightleftharpoons$ 2NO₂(g)

Which point shows the system at equilibrium?

Determine the equilibrium constant, K, for this reaction at 25 °C, referring to section 1 of the data booklet.

If you did not obtain an answer in (c)(iii), use ΔG = –43.5 kJ mol⁻¹, but this is not the correct answer.