NO and NO2 react at 25°C to give N2O3 as shown in the equation. NO(g) + NO2(g) ⇌ N2O3(g) ΔΗ = −7.2 kJ mol-1 The reaction is reversible and reaches equilibrium in a closed system.

1. (a)Fig. 2.1 shows how the rate of the forward reaction changes with time. Initially, the rate of the reverse reaction is zero. Complete Fig. 2.1 to sketch how the rate of the reverse reaction changes with time. [Figure 2.1][1]
2. (b)State how the position of equilibrium changes, if at all, when the reaction takes place at 100°C. Explain your answer. Assume the pressure remains constant.[2]
3. (c)Table 2.1 shows the composition of an equilibrium mixture of NO(g), NO2(g) and N2O3(g) at 101 kPa. [Table 2.1] gas | number of moles at equilibrium/mol ---|--- NO | 0.605 NO2 | 0.605 N2O3 | 0.390 Calculate Kp, the equilibrium constant with respect to partial pressures. Deduce the units of Kp.[3]
4. (d)Identify one natural process and one man-made process that cause the formation of atmospheric NO and NO2.[2]
5. (e)(i)NO2 is a free radical. Define free radical.[1]
6. (e)(ii)NO2 has a catalytic role in the oxidation of atmospheric sulfur dioxide. Write equations to show the catalytic role of NO2 in this oxidation.[2]
7. (e)(iii)State one environmental consequence of the oxidation of atmospheric sulfur dioxide.[1]
8. (f)(i)A student titrates nitric acid with a base to form a solution containing aqueous magnesium nitrate. Identify a base that the student could use.[1]
9. (f)(ii)The student evaporates the water to obtain magnesium nitrate solid. When this solid is heated it decomposes. Write an equation for the decomposition of magnesium nitrate.[1]
10. (f)(iii)State how the thermal stability of Group 2 nitrates changes down the group.[1]

This question uses a diagram from the original paper, which isn't shown here. Open the full paper from the past papers archive to see it.