U9.1 · Introduction to EntropyPractice · 熵入门
Lessons covered:U9.1 Introduction to EntropyU9.2 Absolute Entropy and Entropy Change
- 1Select all
In which of these processes does the entropy of the system increase (ΔS > 0)? Select all that apply.
More than one option may be correct. Choose all that apply, then press Submit.
ExplanationSolid to gas and solid to freely moving dissolved ions both spread particles out, so ΔS > 0; heating spreads the molecules over more energy levels, so ΔS > 0. Condensation (gas to liquid) has ΔS < 0, and three gas molecules becoming liquid water has ΔS < 0.
Review: U9.1 Introduction to Entropy
- 2Multiple choice
A flask of ideal gas is connected through a closed valve to an identical evacuated flask. When the valve opens, the gas spreads through both flasks on its own (at constant temperature). Which is the best explanation?
ExplanationIdeal gas molecules don’t interact, so ΔH = 0 for expansion at constant temperature. The process is spontaneous because entropy rises: each molecule can be anywhere in either flask, and arrangements with every molecule in one flask are a vanishingly small fraction. A vacuum doesn’t “pull”; the molecules simply wander across.
Review: U9.1 Introduction to Entropy
- 3Multiple choice
For N₂(g) + O₂(g) → 2NO(g), what is the most reasonable prediction for ΔS°?
ExplanationThe number of gas molecules and the states stay the same, so the entropy change is small. From S° data: ΔS° = 2 × 210.8 − 191.6 − 205.2 = +24.8 J/K, indeed small (for comparison, two gas molecules becoming one typically gives about −150 J/K). Small, but not necessarily exactly zero.
Review: U9.1 Introduction to EntropyU9.2 Absolute Entropy and Entropy Change
- 4Multiple choice
At −10 °C, supercooled water freezes spontaneously, even though freezing lowers the system’s entropy. Does this contradict the idea that spontaneous processes increase the total entropy?
ExplanationTotal entropy = system change + surroundings change. Freezing releases heat into the cold surroundings, so ΔSsurr = −ΔH ÷ T > 0. At −10 °C the surroundings gain more entropy than the system loses, so the total rises and freezing is spontaneous. Above 0 °C the surroundings’ gain is too small to make up for it, so ice no longer forms on its own.
Review: U9.1 Introduction to Entropy
- 5Multiple choice
One mole of ice at −10 °C is heated to steam at 110 °C. In which step does the system’s entropy increase the most?
ExplanationGoing from liquid to gas, molecules go from close together to filling the whole container, the biggest jump in possible arrangements, so the entropy of vaporization far exceeds that of fusion (about 109 J/K versus 22 J/K for water). Warming within one phase raises entropy too, but much less.
Review: U9.1 Introduction to Entropy
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