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Chapter 13: Problem 13

Suppose a reaction has the equilibrium constant \(K=1.3 \times 10^{8}\). What does the magnitude of this constant tell you about the relative concentrations of products and reactants that will be present once equilibrium is reached? Is this reaction likely to be a good source of the products?

Short Answer

Expert verified
The given equilibrium constant, \(K = 1.3 \times 10^8\), is a large value, indicating that the reaction favors the formation of products. At equilibrium, the concentration of products will be much higher than that of reactants. Therefore, this reaction is likely to be a good source of the products.

Step by step solution

01

Understanding the equilibrium constant

The equilibrium constant (K) is a number that describes the ratio of the concentrations of products to reactants at equilibrium. A large K value indicates that the reaction favors the formation of products, while a small K value indicates that the reaction favors the formation of reactants.
02

Analyzing the given K value

In this problem, the equilibrium constant K is given as \(1.3 \times 10^{8}\). This is a large K value, which means that the reaction favors the formation of products.
03

Determining the relative concentrations of products and reactants at equilibrium

As the magnitude of the equilibrium constant is large, it indicates that the concentration of products in the reaction mixture will be much higher than that of the reactants when the equilibrium is reached. In other words, the products are being formed to a far greater extent compared to the amount of reactants remaining at equilibrium.
04

Evaluating if the reaction is a good source of products

Since the reaction favors the formation of products (due to large K value), it is likely to be a good source of products. In fact, with a value of \(1.3 \times 10^{8}\), the reaction is highly product-favorable, and most reactants will be converted into products at equilibrium. In conclusion, the magnitude of the equilibrium constant in this problem indicates that the reaction strongly favors the formation of products, and that the concentration of products will be much higher than that of reactants at equilibrium. This reaction is, therefore, likely to be a good source of the products.

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Most popular questions from this chapter

At a particular temperature, \(K=2.0 \times 10^{-6}\) for the reaction $$ 2 \mathrm{CO}_{2}(g) \rightleftharpoons 2 \mathrm{CO}(g)+\mathrm{O}_{2}(g) $$ If \(2.0\) moles of \(\mathrm{CO}_{2}\) is initially placed into a \(5.0-\mathrm{L}\) vessel, calculate the equilibrium concentrations of all species.

At \(25^{\circ} \mathrm{C}, K_{\mathrm{p}}=5.3 \times 10^{5}\) for the reaction $$ \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g) $$ When a certain partial pressure of \(\mathrm{NH}_{3}(g)\) is put into an otherwise empty rigid vessel at \(25^{\circ} \mathrm{C}\), equilibrium is reached when \(50.0 \%\) of the original ammonia has decomposed. What was the original partial pressure of ammonia before any decomposition occurred?

Old-fashioned "smelling salts" consist of ammonium carbonate, \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3} .\) The reaction for the decomposition of ammonium carbonate $$ \left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3}(s) \rightleftharpoons 2 \mathrm{NH}_{3}(g)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g) $$ is endothermic. Would the smell of ammonia increase or decrease as the temperature is increased?

Consider the following statements: "Consider the reaction \(\mathrm{A}(g)+\mathrm{B}(g) \rightleftharpoons \mathrm{C}(g)\), for which at equilibrium \([\mathrm{A}]=2 M\) \([\mathrm{B}]=1 M\), and \([\mathrm{C}]=4 M .\) To a \(1-\mathrm{L}\) container of the system at equilibrium, you add 3 moles of \(\mathrm{B}\). A possible equilibrium condition is \([\mathrm{A}]=1 M,[\mathrm{~B}]=3 M\), and \([\mathrm{C}]=6 M\) because in both cases \(K=2 . "\) Indicate everything that is correct in these statements and everything that is incorrect. Correct the incorrect statements, and explain.

At a particular temperature, \(K_{\mathrm{p}}=0.25\) for the reaction $$ \mathrm{N}_{2} \mathrm{O}_{4}(g) \rightleftharpoons 2 \mathrm{NO}_{2}(g) $$ a. A flask containing only \(\mathrm{N}_{2} \mathrm{O}_{4}\) at an initial pressure of \(4.5\) atm is allowed to reach equilibrium. Calculate the equilibrium partial pressures of the gases. b. A flask containing only \(\mathrm{NO}_{2}\) at an initial pressure of \(9.0\) atm is allowed to reach equilibrium. Calculate the equilibrium partial pressures of the gases. c. From your answers to parts a and \(\mathrm{b}\), does it matter from which direction an equilibrium position is reached?
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