Catalytic Air Pollution Control

Commercial Technology

Ronald M. Heck, Robert J. Farrauto, Suresh T. Gulati

• Publisher: Wiley
• Year: 2012
• Print ISBN: 9780470275030
• eBook ISBN: 9781118397718
• Edition: 3
• Topic: Physical Sciences
• Subtopic: Organic Chemistry

Catalytic Air Pollution Control: Commercial Technology is the primary source for commercial catalytic air pollution control technology, offering engineers a comprehensive account of all modern catalytic technology. This Third Edition covers all the new advances in technology in automotive catalyst control technology, diesel engine catalyst control technology, small engine catalyst control technology, and alternate sustainable fuels for auto and diesel.

PART I

Fundamentals

Chapter 1

Catalyst Fundamentals

1.1 INTRODUCTION

Chemical reactions occur by breaking chemical bonds of reactants and by forming new bonds and new compounds. Breaking stable bonds requires the absorption of energy, whereas making new bonds results in the liberation of energy. The combination of these energies results in either an exothermic reaction in which the conversion of reactants to products liberates energy or an endothermic process in which the conversion process requires energy. In the former case, the energy of the product is lower than that of the reactants, with the difference being the heat liberated. In the latter case, the product energy is greater by the amount that must be added to conserve the total energy of the system. Under the same reaction conditions, the heat of reaction (ΔH) being a thermodynamic function does not depend on the path or the rate by which reactants are converted to products. Similarly, the free energy of reaction (ΔG) of the reaction is not dependent on the reaction path because it too is a thermodynamic state function. This will be emphasized once we discuss catalytic reactions. The rate of reaction is determined by the slowest step in a conversion process independent of the energy content of the reactants or products.

1.2 CATALYZED VERSUS NONCATALYZED REACTIONS

A few decades ago, chlorofluorocarbons (i.e., CF₂Cl₂), emitted primarily from refrigerants, were found to catalyze the destruction of the ozone (O₃) layer in the stratosphere necessary to protect us from harmful ultraviolet (UV) radiation and its skin cancer consequences. Fortunately alternative chemicals are now used, and this problem is no longer of great concern. It does, however, serve as an excellent example of a homogeneous gas phase catalytic reaction. First let us consider the very slow noncatalytic reaction between gaseous O₃ and O atoms produced by dissociation of O₂ by solar radiation in the upper atmosphere:

(1.1)

Chlorine atoms, produced by solar radiation of chlorofluorohydrocarbons, catalyze the decomposition of ozone by reacting with it to form ClO and O₂ (1.2). The ClO then reacts with the O atoms regenerating Cl and producing more O₂ (1.3).

(1.2)

(1.3)

Adding both reactions results in Eq. (1.1) and completes the catalytic cycle since the Cl and ClO are both consumed and regenerated in the two reactions. Thus, Cl is a homogeneous catalyst for the destruction of O₃. The uncatalyzed reaction is very slow, and its reaction profile can be described kinetically by the Arrhenius profile in which reactants convert to products by surmounting the noncatalytic activation energy barrier (E_NC) as shown in Figure 1.1. The rate constant k of the reaction is inversely related to the exponential of the activation energy, where T is the absolute temperature, R is the universal gas constant, and k_o is the preexponential constant. The Arrhenius equation (1.4) indicates that the rate constant k decreases the higher the activation energy (E).

FIGURE 1.1. Catalyzed and uncatalyzed reaction energy paths for O₃ decomposition to O₂. Activation energy for catalyzed reaction E_C is lower, and the reaction is faster than the noncatalyzed E_NC.

(1.4)

Since the catalyzed reaction has a lower activation energy (E_C), its reaction rate is greater. The barrier was lowered by the Cl catalyst providing a chemical shortcut to products. Although the rate is greater for the catalyzed reaction, the enthalpy (ΔH) and free energy (ΔG) are not changed. Similarly the equilibrium constant for both catalyzed and noncatalyzed reactions is not changed since both operate under the same reaction conditions in the stratosphere. The catalyst can only influence the rate of which reactants are converted to products in accordance to the equilibrium constant and cannot make thermodynamically unfavorable reactions occur. In industrial practice, reactions conditions, such as temperature and pressure, are varied to bring the free energy to a desirable value to permit the reaction to occur.

Now we will consider the conversion of carbon monoxide (CO), a known human poison, to CO₂, a reaction of great importance to the quality of air we breathe daily. The overall rate of the noncatalytic reaction is controlled by the dissociation of the O₂ molecule to O atoms (rate-limiting step), which rapidly react with CO forming CO₂. The temperature required to initiate the dissociation of O₂ is greater than 700°C, and once provided, the reaction rapidly goes to completion with a net liberation of energy (the heat of reaction is exothermic). The requirement to bring about the O₂ dissociation and ultimately the conversion of CO to CO₂ has an activation energy (E_NC). Reaction occurs when a sufficient number of molecules (O₂) possess the energy necessary (as determined by the Boltzmann distribution) to surmount the activation energy barrier (E_NC) shown in Figure 1.2a). The rate of reaction is expressed in accordance with the Arrhenius equation (1.4). Typically the activation energy for the noncatalytic or thermal conversion of CO to CO₂ is about 40 Kcal/mole.

FIGURE 1.2. Activation energy diagram for a) thermal reaction of CO and O₂ and b) the same reaction in the presence of Pt. Activation energy for the noncatalyzed reaction is E_NC. The Pt catalyzed reaction activation energy is designated E_c. Note that the heat of reaction ΔH_R is the same for both reactions. ΔH_a = heat of adsorption; ΔH_D = heat of desorption.

Let us now discuss the effect of passing the same gaseous reactants, CO and O₂, through a reactor containing a solid catalyst. Since the process is now carried out in two separate phases, the term heterogeneous catalytic reaction is used. In the presence of a catalyst such as Pt, the O₂ and CO molecules adsorb on separate sites in a process called chemisorption in which a chemical partial bond is formed between reactants and the catalyst surface. Dissociation of chemisorbed O₂ molecules to chemisorbed O atoms is rapid, occurring essentially at room temperature. Highly reactive adsorbed O atoms react with chemisorbed CO on adjacent Pt sites producing CO₂, which desorbs from the Pt site, completing the reaction and freeing the catalytic site for another cycle. Thus, the activation energy for the Pt catalyzed reaction (E_c), shown in Figure 1.2b), is considerably smaller than that for the noncatalyzed reaction, enhancing the conversion kinetics. Typically the activation energy for Pt catalyzed CO to CO₂ is less than about 20 Kcal/mole. Figure 1.3 shows the initial lightoff of a conversion versus temperature plot for the catalyzed reaction occurring around 100°C. The noncatalyzed reaction has a considerably higher lightoff temperature (around 700°C) because of its higher activation energy. More input energy is necessary to provide the molecules the necessary energy to surmount the activation barrier so lightoff occurs at higher temperatures. It should be noted, however, that the noncatalyzed reaction has a greater sensitivity to temperature, (slope of plot). Thus, the reaction with the higher the activation energy has the greater sensitivity to temperature, making it increase to a greater extent with temperature than that with a lower activation energy. This is a serious problem for highly exothermic reactions, such as CO and hydrocarbon oxidation, where noncatalytic free radical reactions, with large activation energies, can lead to undesirable products. Thus, the temperature must be carefully controlled within the reactor.

FIGURE 1.3. Conversion of CO versus temperature for a noncatalyzed (homogeneous) and a catalyzed reaction.

Equations relating reaction rates to activation energies will be discussed in considerable detail in Chapter 4, but for now, it is sufficient to understand that an inverse relationship exists between the activation energy and the reaction rate.

The environmental significance of catalyzed reactions is now apparent; a r...

Table of contents

• Cover
• Half Title page
• Title page
• Copyright page
• Dedication
• Preface
• Acknowledgments
• Acknowledgments, First Edition
• Acknowledgments, Second Edition
• Part I: Fundamentals
• Part II: Mobile Sources
• Part III: Stationary Sources
• Part IV: New and Emerging Technologies
• Index