For the reaction $\mathrm{H}_{2} \mathrm{O}(g)+\mathrm{Cl}_{2} \mathrm{O}(g) \longrightarrow 2 \mathrm{HClO}(g),$ you $\mathrm{know} \Delta S_{\mathrm{rxn}}^{\circ}$ and $S^{\circ}$ of $\mathrm{HClO}(g)$ and of $\mathrm{H}_{2} \mathrm{O}(g) .$ Write an expression that can be used to determine $S^{\circ}$ of $\mathrm{Cl}_{2} \mathrm{O}(g) .$

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(a) Does the entropy of the surroundings increase for spontaneous processes? (b) In a particular spontaneous process the entropy of the system decreases. What can you conclude about the sign and magnitude of $\Delta S_{\text { surr. }} ?(\mathbf{c})$ During a certain reversible process, the surroundings undergo an entropy change, $\Delta S_{\text { surr }}=-78 \mathrm{J} / \mathrm{K}$ . What is the entropy change of the system for this process?

Indicate whether each statement is true or false. (a) $\Delta S$ is a state function. ( b) If a system undergoes a reversible change, the entropy of the universe increases.(c) If a system undergoes a reversible process, the change in entropy of the system is exactly matched by an equal and opposite

change in the entropy of the surroundings. (d) If a system undergoes a reversible process, the entropy change of the system must be zero.

Indicate whether each statement is true or false. (a) The entropy of the universe increases for any spontaneous process. (b) The entropy change of the system is equal and opposite that of the surroundings for any irreversible process. (c) The entropy of the system must increase in any spontaneous process. (a) The entropy change for an isothermal process depends on both the absolute temperature and the amount of heat reversibly transferred.

Indicate whether each statement is true or false. (a) The second law of thermodynamics says that entropy is conserved. (b) If the entropy of the system increases during a reversible process, the entropy change of the surroundings must decrease by the same amount. (c) In a certain spontaneous process the system undergoes an entropy change of $4.2 \mathrm{J} / \mathrm{K} ;$ therefore, the entropy change of the surroundings must be $-4.2 \mathrm{J} / \mathrm{K}$

Entropy

An insulated Thermos contains 130 g of water at $80.0^{\circ} \mathrm{C}$ You put in a 12.0 g ice cube at $0^{\circ} \mathrm{C}$ to form a system of $i c e+$ original water. (a) What is the equilibrium temperature of the system? What are the entropy changes of the water that was originally the ice cube (b) as it melts and (c) as it warms to the equilibrium temperature? (d) What is the entropy change of the original water as it cools to the equilibrium temperature? (e) What is the net entropy change of the $i c e+$ original water system as it reaches the equilibrium temperature?