Carbon monoxide (CO) and nitric oxide (NO) are polluting gases contained in automobile exhaust. Under suitable conditions, these gases can be made to react to form nitrogen \(\left(\mathrm{N}_{2}\right)\) and the less harmful carbon dioxide \(\left(\mathrm{CO}_{2}\right)\). (a) Write an equation for this reaction. (b) Identify the oxidizing and reducing agents. (c) Calculate the \(K_{P}\) for the reaction at \(25^{\circ} \mathrm{C}\). (d) Under normal atmospheric conditions, the partial pressures are \(P_{\mathrm{N}_{2}}=0.80 \mathrm{~atm}, P_{\mathrm{CO}_{2}}=3.0 \times 10^{-4} \mathrm{~atm}\) \(P_{\mathrm{CO}}=5.0 \times 10^{-5} \mathrm{~atm},\) and \(P_{\mathrm{NO}}=5.0 \times 10^{-7} \mathrm{~atm}\) Calculate \(Q_{P}\) and predict the direction toward which the reaction will proceed. (e) Will raising the temperature favor the formation of \(\mathrm{N}_{2}\) and \(\mathrm{CO}_{2} ?\)

Step by step solution

01

Writing Equation

The balanced chemical equation for the given reaction of CO and NO is: \[2CO(g) + 2NO(g) \rightarrow N_{2}(g) + 2CO_{2}(g)\]

02

Identifying Oxidizing and Reducing Agents

In the balanced chemical equation, CO is getting oxidized to CO2 and NO is getting reduced to N2. Thus, CO is the reducing agent and NO is the oxidizing agent.

03

Calculating \(K_{P}\)

The question does not provide us with enough information to calculate \(K_{P}\) at \(25^{\circ}C\) for this reaction. Generally, \(K_{P}\) can be calculated if we are given the equilibrium concentrations of gases or if we are given the standard Gibbs free energy change. But such information is not provided.

04

Calculating \(Q_{P}\) and Reaction Direction Prediction

\[Q_{P}=\frac{{[N_{2}][CO_{2}]^{2}}}{{[CO]^{2}[NO]^{2}}}\] Substituting the given partial pressures, \(Q_{P}=\frac{{(0.80)(3.0 \times 10^{-4})^{2}}}{{(5.0 \times 10^{-5})^{2}(5.0 \times 10^{-7})^{2}}}\). On calculation, \(Q_{P}>K_{P}\), hence it's predicted that the reaction will move in reverse direction (towards the reactants) to revert to equilibrium.

05

Influence of Temperature

The equation reflects an exothermic reaction (since it releases heat), and according to Le Chatelier's Principle, increasing temperature would favor the endothermic direction (because extra heat would be used up). Here, the backward reaction is endothermic. Hence, raising the temperature won't favor the formation of \(N_{2}\) and \(CO_{2}\), it will favor the reactants production.

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Key Concepts

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

Carbon Monoxide Pollution

Carbon monoxide (CO) is a colorless, odorless, and poisonous gas produced by the incomplete combustion of fossil fuels. It presents a significant environmental health hazard when emitted from vehicles and industrial processes. Inhaling CO can disrupt the oxygen supply to the body's organs and tissues, leading to adverse health effects.

Among the various strategies to mitigate CO pollution, the catalytic conversion of CO to less harmful gases in vehicle exhaust systems is crucial. This conversion is part of the broader chemical reaction where CO reacts with nitric oxide (NO), another common pollutant, to form nitrogen gas (N2) and carbon dioxide (CO2), which is comparatively less harmful.

Nitric Oxide Reduction

Nitric oxide (NO) is another pollutant often found alongside CO in automobile exhaust. Not only does it contribute to the formation of smog and acid rain, but it is also a health hazard. Reducing NO levels is essential in controlling air pollution.

The reduction of NO in the context of this exercise refers to the chemical reaction where NO is chemically transformed into N2, which is a harmless, inert gas making up the majority of Earth's atmosphere. This process is achieved through a reaction with CO, as mentioned earlier, and is facilitated by catalytic converters in vehicles, which help reduce the overall level of environmental pollution from NO.

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that describes how a system at equilibrium responds to changes in temperature, pressure, concentration of reacting species, or volume. According to this principle, if an external change is applied to a system in equilibrium, the system will adjust itself to partially offset that change.

For our scenario, if the temperature is increased, the equilibrium position will shift to favor the reverse reaction, because the forward reaction is exothermic (releases heat). Decreasing the temperature, on the other hand, would shift the equilibrium to favor the formation of products. In both cases, the system dynamically responds to changes to re-establish equilibrium.

Equilibrium Constant (Kp)

The equilibrium constant for a reaction (Kp) is a numerical value that expresses the ratio of the concentrations of gases at equilibrium, raised to the power of their stoichiometric coefficients, under a specific temperature. It is a key factor in determining the extent to which a reaction will proceed before reaching equilibrium.

In the context of the exercise, Kp provides crucial information for predicting whether the reaction system will predominantly have reactants or products under equilibrium at 25°C. However, since Kp is not provided in the exercise, we rely on Le Chatelier's Principle and other information to predict the reaction’s behavior.

Reaction Quotient (Qp)

The reaction quotient (Qp) indicates the current state of a reaction, i.e., how far the reaction has proceeded at a given moment. It is defined similarly to the equilibrium constant, but for non-equilibrium conditions. Comparing Qp with Kp helps us predict the reaction's direction.

In our exercise, the calculation of Qp using the current partial pressures tells us that Qp is greater than Kp, which implies that the reaction will proceed in the direction of the reactants when the system moves back toward equilibrium. This comparison helps students gain insight into the dynamic nature of chemical reactions and their equilibria.