Topic 4.4: Entropy and spontaneity Questions

Calculate the standard entropy change for the reaction N2(g) + 3H2(g) -> 2NH3(g) using the values below.

Identify which of the following changes has the most negative entropy change: (A) 2NO2(g) -> N2O4(g) (B) CaCO3(s) -> CaO(s) + CO2(g) (C) H2O(l) -> H2O(g) (D) NaCl(s) -> Na+(aq) + Cl-(aq)

Outline what happens to the entropy of a system when a solid dissolves in water.

Consider the reaction below, used to make quicklime:

CaCO₃(s) → CaO(s) + CO₂(g)

Using the standard molar entropies S°: CaCO₃(s) = 93, CaO(s) = 40, CO₂(g) = 214 (all in J K⁻¹ mol⁻¹), calculate ΔS° for the reaction and state, with a reason, whether the sign is expected.

Without using any data, predict and explain the sign of ΔS for each of the following changes:

(a) the sublimation of solid iodine, I₂(s) → I₂(g)

(b) 2NO(g) + O₂

(g) → 2NO₂(g)

Calculate the standard Gibbs energy change at 298 K for a reaction with enthalpy change -92 kJ mol-1 and entropy change -199 J K-1 mol-1.

Explain the effect of increasing temperature on the spontaneity of a reaction that is exothermic with a negative entropy change.

Dissolving ammonium chloride in water is endothermic (ΔH = +14.8 kJ mol⁻¹) yet the solid dissolves readily at room temperature (298 K).

Using the value ΔS = +75.0 J K⁻¹ mol⁻¹, explain, with a calculation of ΔG, why the dissolving is spontaneous. [3]

Which change has the largest positive entropy change?

A reaction has ΔH = −100 kJ mol⁻¹ and ΔS = −150 J K⁻¹ mol⁻¹. At which temperatures is it spontaneous?

For the reaction 2SO₂(g) + O₂(g) → 2SO₃(g), ΔH = −198 kJ mol⁻¹ and ΔS = −188 J K⁻¹ mol⁻¹.

(a) Calculate ΔG at 298 K and state whether the reaction is spontaneous at this temperature.

(b) Calculate the temperature above which the reaction becomes non-spontaneous. [4]