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Chapter 15: Problem 18

Which of the following reactions lies to the right, favoring the formation of products, and which lies to the left, favoring formation of reactants? (a) \(2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{NO}_{2}(g) ; K_{p}=5.0 \times 10^{12}\) (b) \(2 \mathrm{HBr}(g) \rightleftharpoons \mathrm{H}_{2}(g)+\mathrm{Br}_{2}(g) ; K_{c}=5.8 \times 10^{-18}\)

Short Answer

Expert verified
Reaction (a) lies to the right, favoring the formation of products, since \(K_p = 5.0 \times 10^{12}\) which is greater than 1. Reaction (b) lies to the left, favoring the formation of reactants, since \(K_c = 5.8 \times 10^{-18}\) which is less than 1.

Step by step solution

01

Reaction (a) Direction

For reaction (a), we have: \(2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{NO}_{2}(g)\) The equilibrium constant Kp is given as: \(K_{p} = 5.0 \times 10^{12}\) Since the value of \(K_p\) is greater than 1, this means that the reaction favors the formation of products. Therefore, reaction (a) lies to the \(\boxed{right}\).
02

Reaction (b) Direction

For reaction (b), we have: \(2 \mathrm{HBr}(g) \rightleftharpoons \mathrm{H}_{2}(g)+\mathrm{Br}_{2}(g)\) The equilibrium constant Kc is given as: \(K_{c} = 5.8 \times 10^{-18}\) Since the value of \(K_c\) is less than 1, this means that the reaction favors the formation of reactants. Therefore, reaction (b) lies to the \(\boxed{left}\).

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Most popular questions from this chapter

At \(373 \mathrm{~K}, K_{p}=0.416\) for the equilibrium $$ 2 \mathrm{NOBr}(g) \rightleftharpoons 2 \mathrm{NO}(g)+\mathrm{Br}_{2}(g) $$ If the pressures of \(\operatorname{NOBr}(g)\) and \(\mathrm{NO}(g)\) are equal, what is the equilibrium pressure of \(\mathrm{Br}_{2}(g)\) ?

At \(900 \mathrm{~K}\), the following reaction has \(K_{p}=0.345\) : $$ 2 \mathrm{SO}_{2}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{SO}_{3}(g) $$ In an equilibrium mixture the partial pressures of \(\mathrm{SO}_{2}\) and \(\mathrm{O}_{2}\) are \(0.135 \mathrm{~atm}\) and \(0.455 \mathrm{~atm}\), respectively. What is the equilibrium partial pressure of \(\mathrm{SO}_{3}\) in the mixture?

The equilibrium constant \(K_{c}\) for \(\mathrm{C}(s)+\mathrm{CO}_{2}(g) \rightleftharpoons\) \(2 \mathrm{CO}(g)\) is \(1.9\) at \(1000 \mathrm{~K}\) and \(0.133\) at \(298 \mathrm{~K}\). (a) If excess \(\mathrm{C}\) is allowed to react with \(25.0 \mathrm{~g}\) of \(\mathrm{CO}_{2}\) in a \(3.00\) - \(\mathrm{L}\) vessel at 1000 \(\mathrm{K}\), how many grams of \(\mathrm{CO}\) are produced? (b) How many grams of \(\mathrm{C}\) are consumed? (c) If a smaller vessel is used for the reaction, will the yield of CO be greater or smaller? (d) Is the reaction endothermic or exothermic?

(a) If \(Q_{c}>K_{c}\), how must the reaction proceed to reach equilibrium? (b) At the start of a certain reaction, only reactants are present; no products have been formed. What is the value of \(Q_{c}\) at this point in the reaction?

Consider the reaction $$ 4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \underset{4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g), \Delta H=-904.4 \mathrm{~kJ}}{\rightleftharpoons} $$ Does each of the following increase, decrease, or leave unchanged the yield of NO at equilibrium? (a) increase \(\left[\mathrm{NH}_{3}\right] ;\) (b) increase \(\left[\mathrm{H}_{2} \mathrm{O}\right] ;(\) c \()\) decrease \(\left[\mathrm{O}_{2}\right]\); (d) decrease the volume of the container in which the reaction occurs: (e) add a catalyst: (f) increase temperature.
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