Question

The oxidation of \(\mathrm{SO}_{2}\) to \(\mathrm{SO}_{3}\) is catalyzed 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 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) Why do we consider \(\mathrm{NO}_{2}\) a catalyst and not an intermediate in this reaction? (c) Is this an example of homogeneous catalysis or heterogeneous catalysis?

Short answer

The overall reaction for the oxidation of SO₂ to SO₃ is found by combining the given reactions: \[2\mathrm{SO}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2\mathrm{SO}_{3}(g)\]. NO₂ is considered a catalyst, not an intermediate, because it gets consumed in the first reaction and regenerated in the second, speeding up the reaction without being permanently consumed. Since all substances involved in the process are gases and in the same phase, this is an example of homogeneous catalysis.

Step-by-step solution

Find the overall reaction

To find the overall reaction, we need to sum the two given reactions.\\ Reaction 1: \[\mathrm{NO}_{2}(g)+\mathrm{SO}_{2}(g) \longrightarrow \mathrm{NO}(g)+\mathrm{SO}_{3}(g)\] Reaction 2: \[2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g)\]\\ First, multiply reaction 1 by two: \[2*\mathrm{Reaction \ 1} : 2\mathrm{NO}_{2}(g)+2\mathrm{SO}_{2}(g) \longrightarrow 2\mathrm{NO}(g)+2\mathrm{SO}_{3}(g)\] Now, sum up the multiplied reaction 1 and reaction 2: \[\begin{array}{ll} & 2\mathrm{NO}_{2}(g)+2\mathrm{SO}_{2}(g) \longrightarrow 2\mathrm{NO}(g)+2\mathrm{SO}_{3}(g)\\ + & 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g)\\ \hline \end{array}\] The obtained overall reaction is: \[2\mathrm{SO}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2\mathrm{SO}_{3}(g)\]

Determine if NO₂ is a catalyst or an intermediate

A catalyst is a substance that speeds up a chemical reaction but does not get consumed in the process. An intermediate, on the other hand, is a short-lived substance produced and consumed during a reaction. Looking at the overall reaction, we see that \(\mathrm{NO}_{2}\) is not part of the reactants or products. However, it is present in both the reactions comprising the overall reaction but gets consumed in the first reaction and regenerated in the second. This means NO₂ serves a purpose in speeding up the reaction but does not get consumed permanently. Therefore, \(\mathrm{NO}_{2}\) is a catalyst, not an intermediate.

Classify the type of catalysis

A catalytic reaction can be classified as homogeneous or heterogeneous. Homogeneous catalysis occurs when the catalyst and the reactants are in the same phase, while heterogeneous catalysis takes place when the catalyst is in a different phase than the reactants. In the given reactions, all the substances are gases (\(\mathrm{NO}_{2}, \mathrm{SO}_{2}, \mathrm{O}_{2}, \mathrm{NO},\) and \(\mathrm{SO}_{3}\)). Therefore, the catalyst \(\mathrm{NO}_{2}\) and the reactants are in the same gas phase. Hence, this is an example of homogeneous catalysis.

Key concepts

Homogeneous Catalysis

Homogeneous catalysis is an important concept in chemistry where the catalyst and the reactants are in the same phase—typically gas or liquid. This similarity in phase allows for the catalyst to interact with the reactants more easily, which can lead to more effective and efficient reactions. For instance, in the oxidation of sulfur dioxide (\textbf{SO}\(_2\)) to sulfur trioxide (\textbf{SO}\(_3\)) using nitrogen dioxide (\textbf{NO}\(_2\)) as a catalyst, all the components are in the gas phase. This ensures that \textbf{NO}\(_2\) can effectively facilitate the reaction without being used up in the end product.

It's essential to distinguish homogeneous catalysis from heterogeneous catalysis. The latter involves the catalyst being in a different phase from the reactants, like a solid catalyst in a liquid reaction mixture. Homogeneous catalysis is widely used in industrial processes due to its advantages, which include ease of mixing and often, a higher selectivity and rate of reaction.

Chemical Reaction Mechanism

Understanding the chemical reaction mechanism is crucial for grasping how reactions proceed at a molecular level. It involves a step-by-step description of how reactants transition into products, which includes the breaking and forming of chemical bonds. A pivotal concept in these mechanisms is the idea of intermediates—species that are formed and consumed during the course of the reaction—versus catalysts which are regenerated at the end of the reaction.

In the case of the reaction given in our exercise, we see that \textbf{NO}\(_2\) helps convert \textbf{SO}\(_2\) into \textbf{SO}\(_3\) by first reacting with it and then being regenerated. The mechanism shows how \textbf{NO}\(_2\) temporarily becomes \textbf{NO} before being turned back into \textbf{NO}\(_2\). By examining each step, scientists can determine the most effective way to facilitate a reaction and potentially find more efficient catalysts to accomplish a desired chemical transformation.

Role of a Catalyst

A catalyst plays a pivotal role in chemical reactions by increasing the rate of the reaction without being permanently consumed by it. Catalysts operate by providing an alternative pathway with a lower activation energy compared to the uncatalyzed reaction. Although a catalyst is involved in the intermediate steps of a reaction, by the end of the reaction, it remains unchanged, allowing it to catalyze multiple rounds of the reaction.

In our textbook example, \textbf{NO}\(_2\) serves as a catalyst for the oxidation of \textbf{SO}\(_2\) to \textbf{SO}\(_3\). It does not appear in the overall reaction equation because it is not a reactant nor a product of the overall process. Instead, it facilitates the individual steps of the mechanism, effectively speeding up the reaction without being consumed. Catalysis is a fascinating area because of its practical benefits. It not only makes reactions faster but can also make them more selective, cost-effective, and environmentally friendly by reducing the need for excessive energy and potentially harmful by-products.