Unit 9 Explanation 4

Unit 9: Applications of Thermodynamics

🔬 Real-World Chemistry Examples (with Numbers)

Example 1: Electrochemical Cell – Calculate ΔG from E°

Given:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
E°cell = +1.10 V
n = 2 mol e⁻
F = 96,485 C/mol

Solution:
ΔG° = –nFE° = –(2)(96,485)(1.10) = –212,267 J = –212.3 kJ

✅ This means the reaction is spontaneous under standard conditions.

Example 2: Predict Spontaneity Using ΔH and ΔS

Given:
ΔH = –92.0 kJ/mol
ΔS = –198 J/mol·K
T = 298 K

Convert: ΔS = –0.198 kJ/mol·K
ΔG = ΔH – TΔS = –92.0 – (298)(–0.198) = –92.0 + 59.0 = –33.0 kJ/mol

✅ Spontaneous at room temperature.

Example 3: Using ΔG to Estimate K

Given: ΔG° = –34.2 kJ/mol, T = 298 K, R = 8.314 J/mol·K
Convert ΔG: –34,200 J
ln K = –ΔG / RT = 34200 / (8.314 × 298) ≈ 13.8
K ≈ e^13.8 ≈ ~9.8 × 10⁵

✅ Product-favored reaction.

Example 4: Entropy Change in a Reaction

Reaction: N₂(g) + 3H₂(g) → 2NH₃(g)
ΔS° = ΣS°(products) – ΣS°(reactants)
= [2 × 193] – [1 × 191 + 3 × 131] = 386 – 584 = –198 J/mol·K

❗ Even though ΔS is negative, this reaction is still spontaneous because ΔH is very negative (exothermic).

Example 5: ATP Hydrolysis in Biology

ATP → ADP + Pi
ΔG° = –30.5 kJ/mol (biochemical standard conditions)

🔁 This energy is used to power non-spontaneous reactions like protein synthesis or muscle contraction by reaction coupling.

Takeaway:

• Thermodynamics is everywhere: batteries, metabolism, chemical production
• Practice reading values, converting units, and interpreting signs
• Expect math-based MCQs and verbal FRQs based on these concepts