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Chapter 14: Problem 79

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?

Short Answer

Expert verified
The overall reaction for the oxidation of SO₂ by O₂ to give SO₃ is given by: 2SO₂(g) + O₂(g) → 2SO₃(g). NO₂ and NO are both intermediates in this reaction, as they are formed and consumed during the course of the reactions. This is an example of homogeneous catalysis since all reactants and products are in the gas phase.

Step by step solution

01

Sum the reactions with appropriate coefficients

To find the overall reaction, we need to eliminate any elements that appear on both the reactants' and products' side. In our case, it is NO. Since there are two moles of NO in the second reaction, we'll multiply the first reaction by 2. 1. 2(NO₂(g) + SO₂(g) → NO(g) + SO₃(g)) 2. 2NO(g) + O₂(g) → 2NO₂(g) Now, we will add the reactions together and eliminate NO: Overall reaction: 2NO₂(g) + 2SO₂(g) + O₂(g) → 2(NO(g) + SO₃(g)) + 2NO₂(g) Cancel out 2NO₂(g) on both sides: Overall reaction: 2SO₂(g) + O₂(g) → 2SO₃(g)
02

Identify NO₂ as catalyst or intermediate

A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process, while an intermediate is a substance that is formed and consumed during the course of a reaction. To determine if NO₂ is a catalyst or an intermediate, let's look at both reactions: 1. NO₂(g) + SO₂(g) → NO(g) + SO₃(g) 2. 2NO(g) + O₂(g) → 2NO₂(g) NO₂ is consumed in the first reaction and formed again in the second reaction. Therefore, NO₂ is an intermediate in this reaction.
03

Identify NO as catalyst or intermediate

To determine if NO is a catalyst or an intermediate, let's again look at both reactions: 1. NO₂(g) + SO₂(g) → NO(g) + SO₃(g) 2. 2NO(g) + O₂(g) → 2NO₂(g) NO is formed in the first reaction and consumed in the second reaction. Therefore, NO is also an intermediate in this reaction.
04

Identify the type of catalysis

Homogeneous catalysis occurs when the catalyst and the reactants are in the same phase (e.g., gas or liquid), while heterogeneous catalysis occurs when the catalyst and the reactants are in different phases. In this case, all the reactants and products are in the gas phase: 1. NO₂(g) + SO₂(g) → NO(g) + SO₃(g) 2. 2NO(g) + O₂(g) → 2NO₂(g) Therefore, this is an example of homogeneous catalysis.

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Most popular questions from this chapter

For a first order reaction \(\mathrm{A} \longrightarrow \mathrm{B}+\mathrm{C},\) if the half-life of \(\mathrm{A}\) at \(25^{\circ} \mathrm{C}\) is $3.05 \times 10^{4} \mathrm{~s},\( what is the rate constant \)k$ at this temperature? What percentage of A will not have reacted after one day?

The gas-phase reaction of \(\mathrm{NO}\) with \(\mathrm{F}_{2}\) to form \(\mathrm{NOF}\) and \(\mathrm{F}\) has an activation energy of $E_{a}=6.3 \mathrm{~kJ} / \mathrm{mol}\(. and a frequency factor of \)A=6.0 \times 10^{8} M^{-1} \mathrm{~s}^{-1}$. The reaction is believed to be bimolecular: $$ \mathrm{NO}(g)+\mathrm{F}_{2}(g) \longrightarrow \mathrm{NOF}(g)+\mathrm{F}(g) $$ (a) Calculate the rate constant at \(100^{\circ} \mathrm{C}\). (b) Draw the Lewis structures for the \(\mathrm{NO}\) and the NOF molecules, given that the chemical formula for NOF is misleading because the nitrogen atom is actually the central atom in the molecule. (c) Predict the shape for the NOF molecule. (d) Draw a possible transition state for the formation of NOF, using dashed lines to indicate the weak bonds that are beginning to form. (e) Suggest a reason for the low activation energy for the reaction.

(a) Most commercial heterogeneous catalysts are extremely finely divided solid materials. Why is particle size important? (b) What role does adsorption play in the action of a heterogeneous catalyst?

As shown in Figure 14.23 , the first step in the heterogeneous hydrogenation of ethylene is adsorption of the ethylene molecule on a metal surface. One proposed explanation for the "sticking" of ethylene to a metal surface is the interaction of the electrons in the \(\mathrm{C}-\mathrm{C} \pi\) bond with vacant orbitals on the metal surface. (a) If this notion is correct, would ethane be expected to adsorb to a metal surface, and, if so, how strongly would ethane bind compared to ethylene? (b) Based on its Lewis structure, would you expect ammonia to adsorb to a metal surface using a similar explanation as for ethylene?

The following mechanism has been proposed for the gasphase reaction of \(\mathrm{H}_{2}\) with ICl: $$ \begin{array}{l} \mathrm{H}_{2}(g)+\mathrm{ICl}(g) \longrightarrow \mathrm{HI}(g)+\mathrm{HCl}(g) \\ \mathrm{HI}(g)+\mathrm{ICl}(g) \longrightarrow \mathrm{I}_{2}(g)+\mathrm{HCl}(g) \end{array} $$ (a) Write the balanced equation for the overall reaction. (b) Identify any intermediates in the mechanism. (c) If the first step is slow and the second one is fast, which rate law do you expect to be observed for the overall reaction?
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