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Chapter 13: Problem 84

The synthesis of ammonia gas from nitrogen gas and hydrogen gas represents a classic case in which a knowledge of kinetics and equilibrium was used to make a desired chemical reaction economically feasible. Explain how each of the following conditions helps to maximize the yield of ammonia. a. running the reaction at an elevated temperature b. removing the ammonia from the reaction mixture as it forms c. using a catalyst d. running the reaction at high pressure

Short Answer

Expert verified
Increasing temperature is unfavorable for the yield but favorable for reaction rate, while removing ammonia as it forms shifts the equilibrium towards more ammonia production. Using a catalyst increases the reaction rate without affecting equilibrium, and high pressure shifts the equilibrium towards ammonia formation, maximizing the yield.

Step by step solution

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a. Running the reaction at an elevated temperature

The forward reaction (formation of ammonia from nitrogen and hydrogen) is exothermic according to the reaction: \[N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) + heat\] Increasing the temperature shifts the equilibrium toward the reverse reaction (endothermic) according to Le Chatelier's principle. Thus, running the reaction at an elevated temperature is not favorable for maximizing the yield of ammonia. However, an elevated temperature increases the reaction rate, allowing the reaction to reach equilibrium faster. In practice, the temperature is maintained at a compromise level where the rate of reaction is considerable and ammonia production is still favorable.
02

b. Removing the ammonia from the reaction mixture as it forms

According to Le Chatelier's principle, removing a product from the reaction mixture helps to shift the equilibrium toward the forward reaction. In this case, ammonia gas is removed from the reaction mixture, and the equilibrium shifts towards the production of more ammonia, leading to a maximized yield. Therefore, continuous removal of ammonia is beneficial for maximizing ammonia synthesis.
03

c. Using a catalyst

A catalyst is used to increase the rate at which a reaction takes place by providing an alternative pathway with a lower activation energy. The use of a catalyst in this reaction does not alter the position of equilibrium but helps the reaction to reach equilibrium more quickly. This is beneficial in industrial processes, as a faster reaction rate allows for more efficient production. In the synthesis of ammonia, iron is used as a catalyst to ensure faster production.
04

d. Running the reaction at high pressure

Le Chatelier's principle states that increasing pressure favors the side of the reaction with fewer moles of gas. In the synthesis of ammonia, there are 4 moles of gas on the reactant side and only 2 moles of gas on the product side. Therefore, running the reaction at high pressure shifts the equilibrium toward the forward reaction, and more ammonia is produced. Hence, applying high pressure is helpful in maximizing the yield of ammonia in its synthesis.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Le Chatelier's Principle
Understanding Le Chatelier's Principle is crucial when optimizing chemical reactions, such as the synthesis of ammonia. This principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.

For instance, when ammonia is removed from the system, as in the textbook exercise, the principle predicts that the equilibrium will shift to produce more ammonia, compensating for its loss. This is exactly why continuous removal of ammonia leads to an increased yield. By applying this concept, chemical engineers can manipulate the conditions to favor the production of desired substances.
Chemical Equilibrium
Chemical equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction, and the concentrations of reactants and products remain constant over time.

In the context of ammonia synthesis, the equilibrium lies between nitrogen, hydrogen, and ammonia. Adjusting conditions such as temperature and pressure can disturb the equilibrium. For example, increasing pressure shifts the equilibrium towards producing more ammonia because there are fewer gas molecules on the product side, as the exercise solution notes. Understanding how to manipulate these conditions allows for enhanced ammonia yield while maintaining a state of equilibrium.
Reaction Kinetics
Reaction Kinetics deals with the speed at which chemical reactions occur and the factors that affect this speed. One might wonder why we don't simply run the ammonia synthesis at low temperatures to maximize yield, as it's an exothermic reaction. The reason is kinetics.

At lower temperatures, the rate at which equilibrium is reached is extremely slow. Chemical kinetics tells us that increasing temperature speeds up reactions. Thus, there is a balance to strike鈥攁 compromise temperature that optimizes both the yield and rate of production of ammonia, demonstrating the interplay between kinetics and equilibrium considerations.
Catalysis
In the arena of Catalysis, catalysts are the unsung heroes. They speed up reactions without being consumed by lowering the activation energy needed for the reaction to occur. This property is incredibly beneficial in industrial processes like the synthesis of ammonia where iron, the catalyst, accelerates the reaction rate, enabling a faster approach to equilibrium as explained in the exercise solution.

Importantly, catalysts don't affect the position of equilibrium; they merely get us there quicker. In mass production, time is money, and therefore, using catalysts like iron in ammonia production is essential for economic viability.

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Most popular questions from this chapter

At a particular temperature, \(K=1.00 \times 10^{2}\) for the reaction $$ \mathrm{H}_{2}(\mathrm{~g})+\mathrm{I}_{2}(g) \rightleftharpoons 2 \mathrm{HI}(g) $$ In an experiment, \(1.00\) mole of \(\mathrm{H}_{2}, 1.00 \mathrm{~mole}\) of \(\mathrm{I}_{2}\), and \(1.00\) mole of HI are introduced into a \(1.00-\mathrm{L}\) container. Calculate the concentrations of all species when equilibrium is reached.

Predict the shift in the equilibrium position that will occur for each of the following reactions when the volume of the reaction container is increased. a. \(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)\) b. \(\mathrm{PCl}_{5}(g) \rightleftharpoons \mathrm{PCl}_{3}(g)+\mathrm{Cl}_{2}(g)\) c. \(\mathrm{H}_{2}(g)+\mathrm{F}_{2}(g) \rightleftharpoons 2 \mathrm{HF}(g)\) d. \(\mathrm{COCl}_{2}(g) \rightleftharpoons \mathrm{CO}(g)+\mathrm{Cl}_{2}(g)\) e. \(\mathrm{CaCO}_{3}(s) \rightleftharpoons \mathrm{CaO}(s)+\mathrm{CO}_{2}(g)\)

In a given experiment, \(5.2\) moles of pure NOCl was placed in an otherwise empty \(2.0-\mathrm{L}\) container. Equilibrium was established by the following reaction: $$ 2 \mathrm{NOCl}(g) \rightleftharpoons 2 \mathrm{NO}(g)+\mathrm{Cl}_{2}(g) \quad K=1.6 \times 10^{-5} $$ a. Using numerical values for the concentrations in the Initial row and expressions containing the variable \(x\) in both the Change and Equilibrium rows, complete the following table summarizing what happens as this reaction reaches equilibrium. Let \(x=\) the concentration of \(\mathrm{Cl}_{2}\) that is present at equilibrium. b. Calculate the equilibrium concentrations for all species.

The hydrocarbon naphthalene was frequently used in mothballs until recently, when it was discovered that human inhalation of naphthalene vapors can lead to hemolytic anemia. Naphthalene is \(93.71 \%\) carbon by mass, and a \(0.256\) -mole sample of naphthalene has a mass of \(32.8 \mathrm{~g}\). What is the molecular formula of naphthalene? This compound works as a pesticide in mothballs by sublimation of the solid so that it fumigates enclosed spaces with its vapors according to the equation Naphthalene \((s) \rightleftharpoons\) naphthalene \((g)\) $$ K=4.29 \times 10^{-6}(\text { at } 298 \mathrm{~K}) $$ If \(3.00 \mathrm{~g}\) solid naphthalene is placed into an enclosed space with a volume of \(5.00 \mathrm{~L}\) at \(25^{\circ} \mathrm{C}\), what percentage of the naphthalene will have sublimed once equilibrium has been established?

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?
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