(ii) Which side of the reaction is favored by entropy? (iii) If ∆S° = 0 for these reactions, calculate ∆G° (Assume T = 298 K) [BDE for O―H = 110 kcal /mol.]
(b)

Would you expect ∆S to be greater than, less than, or equal to zero in the following reactions?

a.

For a reaction carried out at 25 °C with an equilibrium constant of 1 × 10^-3, to increase the equilibrium constant by a factor of 10:

a. how much must ∆G° change?

b. how much must ∆H° change if ∆S° = 0 kcal mol^-1 K^-1?

c. how much must ∆S° change if ∆H° = 0 kcal mol^-1?

From the following rate constants, determined at five temperatures, calculate the experimental energy of activation and ∆G^‡, ∆H^‡, and ∆S^‡ for the reaction at 30 °C:

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a. What is the equilibrium constant for a reaction that is carried out at 25 °C (298 K) with ∆H° = 20 kcal/mol and ∆S° = 5.0 × 10^-2 kcal mol^-1 K^-1?

b. What is the equilibrium constant for the same reaction carried out at 125 °C?

For the following values of ∆H° , ∆S°, and T, tell whether the process would be favored.

(b) ∆H° = +7.34 kcal/mol ; ∆S° = +43 cal/mol•K ; T = 325 K

Considering the process described in Assessment 5.13, will it be favored or disfavored at a temperature higher than the one you calculated? How about at a temperature below what you calculated?

For the following values of ∆H° , ∆S°, and T, tell whether the process would be favored.

(d) ∆H° = -8.3 kcal/mol ; ∆S° = -12 cal/mol•K ; T = 298 K

Calculate ∆G°, ∆H°, and ∆S° for the following acid–base reactions. Rationalize the value of ∆H° based on the structure of the conjugate bases. [Assume T = 298 K.]

(b)