Problem 79 - Chemistry The Central Science In Si Units

Question

The oxidation of \(\mathrm{SO}_{2}\) to \(\mathrm{SO}_{3}\) is accelerated by \(\mathrm{NO}_{2}\). The reaction proceeds according to: $$ \begin{array}{l} \mathrm{NO}_{2}(g)+\mathrm{SO}_{2}(g) \longrightarrow \mathrm{NO}(g)+\mathrm{SO}_{3}(g) \\ 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g) \end{array} $$ (a) Show that, with appropriate coefficients, the two reactions can be summed to give the overall oxidation of \(\mathrm{SO}_{2}\) by \(\mathrm{O}_{2}\) to give \(\mathrm{SO}_{3}\). (b) Do we consider \(\mathrm{NO}_{2}\) a catalyst or an intermediate in this reaction? (c) Would you classify NO as a catalyst or as an intermediate? (d) Is this an example of homogeneous catalysis or heterogeneous catalysis?

Step-by-Step Solution

Answer

(a) Overall: \(\mathrm{SO}_2 + \frac{1}{2}\mathrm{O}_2 \rightarrow \mathrm{SO}_3\); (b) \(\mathrm{NO}_2\) is a catalyst; (c) \(\mathrm{NO}\) is an intermediate; (d) Homogeneous catalysis.

Step 1: Identify and Sum the Reactions

The given reactions are: \( \mathrm{NO}_2(g) + \mathrm{SO}_2(g) \rightarrow \mathrm{NO}(g) + \mathrm{SO}_3(g) \) and \( 2 \mathrm{NO}(g) + \mathrm{O}_2(g) \rightarrow 2 \mathrm{NO}_2(g) \). By summing the reactions, we cancel out \(\mathrm{NO}_2\) and \(\mathrm{NO}\) to obtain the overall reaction. Combine the equations:\[ \begin{aligned} \mathrm{NO}_2(g) + \mathrm{SO}_2(g) &\rightarrow \mathrm{NO}(g) + \mathrm{SO}_3(g) \2\mathrm{NO}(g) + \mathrm{O}_2(g) &\rightarrow 2\mathrm{NO}_2(g) \end{aligned} \]Add them and cancel out the intermediates:\[ \mathrm{NO}_2(g) + \mathrm{SO}_2(g) + 2\mathrm{NO}(g) + \mathrm{O}_2(g) \rightarrow \mathrm{NO}(g) + \mathrm{SO}_3(g) + 2\mathrm{NO}_2(g) \]Simplify to: \[ \mathrm{SO}_2(g) + \dfrac{1}{2}\mathrm{O}_2(g) \rightarrow \mathrm{SO}_3(g) \]

Step 2: Identify Role of NO2

\(\mathrm{NO}_2\) is regenerated at the end of the reaction mechanism, and it facilitates the conversion of \(\mathrm{SO}_2\) to \(\mathrm{SO}_3\), showing that it is not consumed in the overall process. Therefore, \(\mathrm{NO}_2\) acts as a catalyst because it participates in the reaction but is not consumed by it.

Step 3: Identify Role of NO

\(\mathrm{NO}\) is produced in the first reaction and consumed in the second, without appearing in the overall balanced reaction. Thus, \(\mathrm{NO}\) acts as an intermediate because it is formed and used up during the reaction sequence.

Step 4: Classify Type of Catalysis

Both the reaction mixture and the catalyst \(\mathrm{NO}_2\) are in the gas phase, which indicates that the process is an example of homogeneous catalysis. In homogeneous catalysis, the reactants and catalyst share the same phase.

Key Concepts

Oxidation-Reaction

Oxidation reactions are processes where a substance loses electrons, often involving the transfer of electrons to another substance. In the oxidation of \(\mathrm{SO}_2\) to \(\mathrm{SO}_3\), \(\mathrm{SO}_2\) undergoes chemical change as it gains oxygen. This reaction can be understood as an electron transfer process where the sulfur molecule in \(\mathrm{SO}_2\) increases its oxidation state by receiving oxygen from another molecule, like \(\mathrm{O}_2\).Moreover, oxidation reactions are key components in many natural and industrial processes. For instance, they are critical in processes such as combustion and metabolism.
Understanding how these reactions are balanced and how intermediates are involved are core aspects of studying chemical kinetics, the area of chemistry concerned with the speed or rate of these reactions.
An important point to remember about oxidation reactions in the context of kinetics is that they often occur in a series of steps rather than a single step. This is where intermediates and catalysts could feature prominently, aiding the overall transformation without being consumed themselves.

Catalysis

Catalysts are substances that increase the rate of a chemical reaction without themselves being consumed or permanently altered. They are crucial in both industrial and natural catalytic processes. In the exercise provided, \(\mathrm{NO}_2\) serves as a catalyst.

• It accelerates the reaction of \(\mathrm{SO}_2\) to \(\mathrm{SO}_3\).
• It participates in the reaction by temporarily forming an intermediate but is regenerated at the end of the reaction process.

Keeping that in mind, this exemplifies how catalysts do not undergo permanent change, allowing them to be used repeatedly.
Catalysis is pivotal in many fields, such as catalyzing the production of chemicals, refining petroleum, and even in biological processes involving enzymes.
The efficiency and effectiveness of a catalyst are generally measured by the rate-enhancing factor it provides to a reaction, while also considering the selectivity and the conditions under which it operates.

Reaction-Mechanisms

Reaction mechanisms offer a step-by-step description of how reactants turn into products in chemical reactions. They are essentially the blueprint of reaction pathways. Understanding them helps chemists and chemical engineers design better processes and products.In the problem's scenario, the mechanism involves a sequence of steps:

• The reaction between \(\mathrm{NO}_2\) and \(\mathrm{SO}_2\) to form \(\mathrm{NO}\) and \(\mathrm{SO}_3\).
• The regeneration of \(\mathrm{NO}_2\) from \(\mathrm{NO}\) using oxygen.

In such multi-step reactions, it's crucial to differentiate between intermediates and catalysts. \(\mathrm{NO}\) is an intermediate, as it is both formed and consumed during the reaction, whereas \(\mathrm{NO}_2\) is a catalyst as noted before.
Understanding these roles helps clarify how substances interact in complex reactions and why such a reaction does not follow a single elementary step model.
Overall, breaking down the reaction into individual steps makes it easier to study the rates and pathways of complete reaction processes, providing profound insights into chemical dynamics.