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Chapter 18: Problem 5

How does the entropy of a system change for each of the following processes? (a) A solid melts. (b) A liquid freezes. (c) A liquid boils. (d) A vapor is converted to a solid. (e) A vapor condenses to a liquid. (f) A solid sublimes. (g) Urea dissolves in water.

Short Answer

Expert verified
Entropy increases when a solid melts, a liquid boils, a solid sublimes, and when urea dissolves in water. Entropy decreases when a liquid freezes, a vapor condenses to a liquid, and a vapor is converted to a solid.

Step by step solution

01

Understanding Entropy Change During Melting

When a solid melts to a liquid, the entropy increases. This is because the particles in a liquid are more disordered and have more freedom of movement than in a solid.
02

Understanding Entropy Change During Freezing

The reverse process, when a liquid freezes into a solid, the entropy decreases. The particles in a solid are more ordered and have less freedom of movement than in a liquid.
03

Understanding Entropy Change During Boiling

When a liquid boils to form a gas, the entropy increases. The particles in a gas are much more disordered and have a lot more freedom of movement than in a liquid.
04

Understanding Entropy Change During Vapor to Solid Conversion

When a vapor is converted to a solid, the entropy decreases significantly. The particles in a vapor are very unordered and have high freedom of movement. In a solid, the opposite is true.
05

Understanding Entropy Change During Condensation

When a vapor condenses to a liquid, the entropy decreases. The particles in a liquid are more ordered and have less freedom of movement compared to a gas.
06

Understanding Entropy Change During Sublimation

When a solid sublimes to a gas, the entropy increases greatly. The path from solid to gas bypasses the liquid state, resulting in a large increase in disorder.
07

Understanding Entropy Change During Dissolution of Urea in Water

When urea dissolves in water, the entropy increases. The dissolved particles of urea increase the disorder of the solution.

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Most popular questions from this chapter

The equilibrium constant \(\left(K_{P}\right)\) for the reaction $$ \mathrm{H}_{2}(g)+\mathrm{CO}_{2}(g) \rightleftharpoons \mathrm{H}_{2} \mathrm{O}(g)+\mathrm{CO}(g) $$ is 4.40 at \(2000 \mathrm{~K}\). (a) Calculate \(\Delta G^{\circ}\) for the reaction. (b) Calculate \(\Delta G\) for the reaction when the partial pressures are \(P_{\mathrm{H}_{2}}=0.25 \mathrm{~atm}, P_{\mathrm{CO}_{2}}=0.78 \mathrm{~atm}\) \(P_{\mathrm{H}_{2} \mathrm{O}}=0.66 \mathrm{~atm},\) and \(P_{\mathrm{CO}}=1.20 \mathrm{~atm}\)

Ammonium nitrate \(\left(\mathrm{NH}_{4} \mathrm{NO}_{3}\right)\) dissolves spontaneously and endothermically in water. What can you deduce about the sign of \(\Delta S\) for the solution process?

(a) Calculate \(\Delta G^{\circ}\) and \(K_{P}\) for the following equilibrium reaction at \(25^{\circ} \mathrm{C}\). The \(\Delta G_{f}^{\circ}\) values are 0 for \(\mathrm{Cl}_{2}(g),-286 \mathrm{~kJ} / \mathrm{mol}\) for \(\mathrm{PCl}_{3}(g),\) and \(-325 \mathrm{~kJ} / \mathrm{mol}\) for \(\mathrm{PCl}_{5}(g)\) $$ \mathrm{PCl}_{5}(g) \rightleftharpoons \mathrm{PCl}_{3}(g)+\mathrm{Cl}_{2}(g) $$. (b) Calculate \(\Delta G\) for the reaction if the partial pressures of the initial mixture are \(P_{\mathrm{PCl}_{5}}=0.0029 \mathrm{~atm}\) \(P_{\mathrm{PCl}_{3}}=0.27 \mathrm{~atm},\) and \(P_{\mathrm{Cl}_{2}}=0.40 \mathrm{~atm}\).

Consider the following Bronstead acid-base reaction at \(25^{\circ} \mathrm{C}\) : $$ \mathrm{HF}(a q)+\mathrm{Cl}^{-}(a q) \rightleftharpoons \mathrm{HCl}(a q)+\mathrm{F}^{-}(a q) $$. (a) Predict whether \(K\) will be greater or smaller than unity, (b) Does \(\Delta S^{\circ}\) or \(\Delta H^{\circ}\) make a greater contribution to \(\Delta G^{\circ} ?\) (c) Is \(\Delta H^{\circ}\) likely to be positive or negative?

Explain what is meant by a spontaneous process. Give two examples each of spontaneous and nonspontaneous processes.
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