Question

In the reaction: \(\mathrm{NO}_{2}+\mathrm{CO} \rightleftarrows \mathrm{NO}+\mathrm{CO}_{2}\), the equilibrium state may be recognized by the constancy of colour.

Short answer

Answer: The equilibrium state in the reaction between nitrogen dioxide and carbon monoxide can be observed through the constancy of color. When the reddish-brown color of the reaction mixture stops fading, it indicates that the reaction has reached equilibrium and the rate of the forward and reverse reactions are equal.

Step-by-step solution

Understand the chemical reaction

The given chemical reaction is: \(\mathrm{NO}_{2}+\mathrm{CO} \rightleftarrows \mathrm{NO}+\mathrm{CO}_{2}\) This is a reversible reaction, which means it can proceed in both the forward and reverse directions. The forward direction is the reaction of nitrogen dioxide with carbon monoxide to produce nitric oxide and carbon dioxide. The reverse direction is the reaction of nitric oxide with carbon dioxide to produce nitrogen dioxide and carbon monoxide.

Identify the characteristics of the reactants and products

It is essential to know the properties of the reactants and products in this reaction. Nitrogen dioxide (NO2) is a reddish-brown gas, carbon monoxide (CO) is a colorless gas, nitric oxide (NO) is a colorless gas, and carbon dioxide (CO2) is a colorless gas.

Recognize the color change

As the reaction proceeds, the reddish-brown color of nitrogen dioxide will decrease and eventually disappear. This is because when NO2 reacts with CO, it forms colorless gases (NO and CO2). The decrease in the NO2 concentration causes the color to fade.

Observing the equilibrium state

When the reaction reaches equilibrium, there will be no noticeable change in the color of the reaction mixture. This is because, at equilibrium, the rate of forward reaction (from reddish-brown NO2 to colorless gases) is equal to the rate of reverse reaction (the colorless gases forming back into reddish-brown NO2). Thus, at equilibrium, the concentration of NO2 remains constant, as well as the concentrations of CO, NO, and CO2.

Conclusion

The equilibrium state of the given reaction can be recognized by the constancy of color. When the color of the reaction mixture stops changing (i.e., reddish-brown color is no longer fading), it indicates that the reaction has reached equilibrium and that the rate of the forward and reverse reactions are equal.

Key concepts

Reversible Reactions

In the fascinating world of chemistry, reversible reactions play a pivotal role. A reversible reaction, such as the interplay between nitrogen dioxide (NO2) and carbon monoxide (CO) producing nitric oxide (NO) and carbon dioxide (CO2), can proceed in both directions—forward and reverse. During the forward reaction, the reddish-brown NO2 reacts with CO, turning into the colorless gases NO and CO2. Notably, the color change from reddish-brown to colorless is a visual cue for students to understand the progression of the reaction. The reverse reaction, on the other hand, reverts NO and CO2 back to NO2 and CO. Reversible reactions reach a point of balance called chemical equilibrium, where the rate of the forward and reverse reactions are equal, keeping the concentrations of reactants and products constant over time.

Visual Signals

The distinct reddish-brown hue of NO2, in contrast to the colorlessness of CO, NO, and CO2, provides a unique opportunity for students to visually monitor the equilibrium process. This change in color is a direct indication of the reaction's shift in direction, a topic central in understanding reversible reactions and their dynamic equilibrium.

Chemical Kinetics

Chemical kinetics delves into the rate at which chemical reactions occur and the factors influencing these rates. For the reaction in question, involving NO2 and CO, understanding kinetics is crucial. The reaction rate is affected by the concentration of the reactants: as NO2 and CO react, their concentrations decrease, slowing the forward reaction rate. Conversely, as NO and CO2 are produced, their concentration increase would usually speed up the reverse reaction rate. However, at equilibrium, both rates are equal and the concentrations of all participating substances stabilize. If the coloration of the reaction remains constant, it is an indicator that equilibrium has been achieved, since the formation of NO2 (causing the color) is countervailed by its consumption in the forward reaction.

Factors Affecting Reaction Rates

Several factors can influence these rates, including temperature, pressure, and the presence of catalysts. When applied to educational content, a focus on these factors and their effects provides a comprehensive understanding of why reactions proceed at the rate that they do—a foundation of chemical kinetics.

Le Chatelier's Principle

Le Chatelier's Principle is a vital concept when exploring the behavior of reversible reactions in dynamic equilibrium. This principle states that if a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system will adjust its equilibrium position to counteract the imposed change. In the context of the NO2 and CO reaction, if we were to increase the concentration of NO2, the system would respond by favoring the forward reaction to produce more NO and CO2, thereby reducing the excess NO2. Similarly, decreasing the NO2 concentration would shift the equilibrium to favor the reverse reaction. Le Chatelier's Principle is pertinent to students' understanding of equilibrium adjustments. It also has practical implications in the production of chemicals. Illustrating how industrial processes optimize yields by manipulating conditions can provide real-world context to this abstract concept, making it not only easier to grasp but also more engaging.