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- Herausgeber: John Wiley & Sons
- Kategorie: Fachliteratur
- Serie: Essential Engineering Calculations Series
- Sprache: Englisch
This books format follows an applications-oriented text and serves as a training tool for individuals in education and industry involved directly, or indirectly, with chemical reactors. It addresses both technical and calculational problems in this field. While this text can be complimented with texts on chemical kinetics and/or reactor design, it also stands alone as a self-teaching aid. The first part serves as an introduction to the subject title and contains chapters dealing with history, process variables, basic operations, kinetic principles, and conversion variables. The second part of the book addresses traditional reactor analysis; chapter topics include batch, CSTRs, tubular flow reactors, plus a comparison of these classes of reactors. Part 3 keys on reactor applications that include non-ideal reactors: thermal effects, interpretation of kinetic data, and reactor design. The book concludes with other reactor topics; chapter titles include catalysis, catalytic reactors, other reactions and reactors, and ABET-related topics. An extensive Appendix is also included
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Veröffentlichungsjahr: 2012
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CONTENTS
Cover
Half Title page
Title page
Copyright page
Dedication
Preface
Overview
References
Part I: Introduction
Chapter 1: History of Chemical Reactions
Introduction
Early History
Recent History
The Chemical Industry Today
Microscopic vs Macroscopic Approach
References
Chapter 2: The Field of Chemistry
Introduction
Inorganic Chemistry
Organic Chemistry
Physical Chemistry
Other Chemistry Topics
Analysis Procedures
References
Chapter 3: Process Variables
Introduction
Temperature
Pressure
Moles and Molecular Weights
Mass and Volume
Viscosity
Heat Capacity
Thermal Conductivity
Reynolds Number
pH
Vapor Pressure
The Ideal Gas Law
Latent Enthalpy Effects
Property Estimation
References
Chapter 4: Kinetic Principles
Introduction
Reaction Rates
Rate vs Equilibrium Considerations
Representation of Rate Expressions
Solutions to Rate Expressions
Reaction Rate Theories
References
Chapter 5: Stoichiometry and Conversion Variables
Introduction
Stoichiometry
Conversion Variables
Volume Correction Factor
Yield and Selectivity
References
Part II: Traditional Reactor Analysis
Chapter 6: Reaction and Reactor Classification
Introduction
Classification of Reactions
Classification of Reactors
Other Industrial Chemical Reactors
Ancillary Equipment
References
Chapter 7: The Conservation Laws
Introduction
Conservation of Mass
Conservation of Energy
Conservation of Momentum
References
Chapter 8: Batch Reactors
Introduction
Equipment Description and Operation
Describing Equations
Specific Reactions
Applications
References
Chapter 9: Continuous Stirred Tank Reactors
Equipment Description and Operation
Describing Equations
Applications
References
Chapter 10: Tubular Flow Reactors
Introduction
Equipment Description and Operation
Describing Equations
Applications
References
Chapter 11: Reactor Comparisons
Introduction
Specific Comparisons: Batch, CSTR, and TF
Graphical Analysis
Applications
References
Part III: Reactor Applications
Chapter 12: Thermal Effects
Introduction
Thermal Fundamentals and Principles
Batch Reactors
CSTR Reactors
Tubular Flow Reactors
References
Chapter 13: Interpretation of Kinetic Data
Introduction
Experimental Methods and Analysis of Kinetic Data
Method of Least Squares
Application to Specific Reactors
Reactions of Complex Mechanism
References
Chapter 14: Nonideal Reactors
Introduction
Nonideal Approaches
Definitions
Estimation of Mean and Variance
Residence Time Distributions
Residence Time Distribution Functions
Experimental Tracer Techniques
References
Chapter 15: Reactor Design Considerations
Introduction
Design Principles
Specific Design Considerations
Operation and Maintenance and Improving Performance
Reactor Selection
Applications
References
Part IV: Other Reactor Topics
Chapter 16: Catalysts
Introduction
Key Definitions and Testing Procedures
Chemical and Formulated Catalysts
Catalytic Processes
Catalyst Selection and Evaluation
References
Chapter 17: Catalytic Reactions
Introduction
The Overall Process
Convective Transfer
Molecular Diffusion
Adsorption/Desorption
Chemical Reaction
The Controlling Steps(s)
References
Chapter 18: Fluidized and Fixed Bed Reactors
Introduction
Fluidized Bed Reactors
Fixed Bed Reactors
Pressure Prop Calculations
Catalytic Reactor Design Considerations
References
Chapter 19: Biochemical Reactors
Introduction
Basic Operations
Design Considerations
Bio Environmental Applications
References
Chapter 20: Open-Ended Problems
Introduction
Applications
References
Chapter 21: Abet-Related Topics
Introduction
Environmental Management
Accident and Emergency Management
Ethics
Numerical Methods
Economics and Finance
References
Appendix: SI Units
The Metric System
The SI System
Seven Basic Units
Two Supplementary Units
SI Multiples and Prefixes
Conversion Constants (SI)
Selected Common Abbreviations
Index
CHEMICAL REACTOR ANALYSIS AND APPLICATIONS FOR THE PRACTICING ENGINEER
Copyright © 2012 by John Wiley & Sons, Inc. All rights reserved
Published by John Wiley & Sons, Inc., Hoboken, New JerseyPublished simultaneously in Canada
No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording, scanning, or otherwise, except as permitted under Section 107 or 108 of the 1976 United States Copyright Act, without either the prior written permission of the Publisher, or authorization through payment of the appropriate per-copy fee to the Copyright Clearance Center, Inc., 222 Rosewood Drive, Danvers, MA 01923, (978) 750-8400, fax (978) 750-4470, or on the web at www.copyright.com. Requests to the Publisher for permission should be addressed to the Permissions Department, John Wiley & Sons, Inc., 111 River Street, Hoboken, NJ 07030, (201) 748-6011, fax (201) 748-6008, or online at http://www.wiley.com/go/permission.
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Library of Congress Cataloging-in-Publication Data:
Theodore, Louis.Chemical reactor analysis and applications for the practicing engineer/Louis Theodore.p. cm.Includes bibliographical references and index.ISBN 978-0-470-91535-6 (hardback)1. Chemical reactors. I. Title.TP157.T46 2012660′.2832—dc232011051081
ToRupert MurdochA remarkable individual who, in his own way, has savedour great nation from the liberal media.
PREFACE
Despite the quality of the consultants and reviewers on this project, the author did not follow their advice/suggestions in every case, which explains whatever confusion and errors remain in the text.
The design and analysis of chemical reactors is one of the most important and complicated tasks facing the practicing engineer. Chemical reaction engineering is that branch of engineering that is concerned with the application of chemical reactions on a commercial scale. Topics typically addressed include:
This project, as is the case with other texts in this series, was a unique undertaking. Rather than preparing a textbook on chemical reaction engineering in the usual format, the author decided to write an applications-oriented text. The text would hopefully serve as a training tool for those individuals in education and industry involved directly, or indirectly, with chemical reactors. It addresses both technical and calculational problems in this field. While this text can be complemented with texts on chemical kinetics and/or reactor design, it also stands alone as a self-teaching aid.
This book is divided into four parts.
The first part serves as an introduction to the subject title and contains chapters dealing with history, process variables, basic operations, chemical kinetic principles, and stoichometry and conversion variables. The second part of the book addresses traditional reactor analysis; chapter topics include batch, CSTRs, and tubular flow reactors, plus a comparison of these classes of reactors. Part III keys on reactor applications that include thermal effects, interpretation of kinetic data, non-ideal reactors, and reactor design. The book concludes with other reactor topics; chapter titles include catalysis, catalytic reactions, fluidized and fixed bed reactors, biochemical reactors, open-ended questions, and ABET-related topics. An Appendix is also included.
The text is supplemented with numerous (nearly 300) illustrative examples. These range in difficulty from simple substitution into equations presented in the essay portion of the chapter to detailed design analysis and open-ended problems. It should be noted that the author cannot claim sole authorship to all the illustrative examples, problems, and essay material in this text. The present book has evolved from a host of sources, including: notes, homework problems and exam problems prepared by L. Theodore for a one-semester, three-credit “Chemical Reaction Kinetics” required undergraduate chemical engineering course at Manhattan College; J. Reynolds, J. Jeris, and L. Theodore’s John Wiley & Sons’ text, Handbook of Chemical and Environmental Engineering Calculations; J. Santoleri, J. Reynolds, and L. Theodore’s Wiley-Interscience text, Introduction to Hazardous Waste Incineration, 2nd edition; S. Fogler’s Prentice-Hall text, Elements of Chemical Reaction Engineering (required for the course at Manhattan College); and, to a much lesser extent, Smith’s McGraw-Hill text, Chemical Engineering Kinetics.
Although the bulk of the problems are original and/or taken from sources that the author has been directly involved with, every effort has been made to acknowledge material drawn from other sources. In particular, Fogler’s book has had a significant impact on the approach employed by the author; it is highly recommended for outside reading. As indicated above, Fogler’s text has been adopted at Manhattan College for the Chemical Reaction Kinetics course and will probably be supplemented in the future with this text.
It is hoped that this writing will place in the hands of government, industrial and academic personnel a text covering the principles and applications of chemical reactors in a thorough and clear manner. Upon completion of the text, the reader should have acquired not only a working knowledge of the principles of chemical reactors but also experience in their application; and, the reader should find themself approaching advanced texts, engineering literature and industrial applications (even unique ones) with more confidence.
Thanks are due to Peter Forzaglia for proofing the manuscript and to Karen Tschinkel for both proofing the manuscript and preparing the index.
Last but not least, we believe that this modest work will help the majority of individuals working in this field to obtain a reasonably complete understanding of chemical reactors. If you have come this far, and read through the Preface and Overview to follow, you have more than just a passing interest in chemical reactors. I strongly suggest you try this book; I think you will like it.
LOUIS THEODOREAugust 2012
Note: You may contact the author at [email protected] for over 200 additional problems and 15 hours of exams; solutions for the problems and exams are available for those who adopt the book for training and/or academic purposes.
OVERVIEW
A major goal of the practicing engineer is to increase the supply and minimize the cost of useful materials to society. Every practicing engineer should be familiar with both the processing of chemicals and the design and operation of the equipment needed to carry out the processes. This equipment can be divided into three main groups. The first group consists of large-scale equipment used to purify or separate raw materials for further processing. These separations are normally physical in nature; the design and operation of this class of equipment are studied in mass transfer operation courses(1). The second group consists of other unit operations, including heat transfer(2) and fluid flow(3) plus chemical reactors in which the processed raw materials react to create products which possess new physical and chemical properties. The reactors may be viewed as a chemical treatment step. The design and operation of chemical reactors are the main subjects of this book. The third class of equipment is similar to the first group but (generally) on a smaller scale. The sequence of steps is pictorially represented in the figure below, and it is the “chemical reactor” box that will receive treatment in the material to follow.
Process flow diagram.
The problem of predicting the performance and determining the size of a reactor can be divided into two sequential steps. The first is the study of the rate at which the chemical reaction occurs and the variables which affect this rate. This is the subject of chemical kinetics (see Chapter 4). The second is the problem of using reaction rate data to predict performance and/or to determine the size of the equipment to obtain the required quantity and quality of product. This is the subject of Parts II and III.
This book was prepared as both a professional book and an undergraduate text for the study of the principles and fundamentals of chemical reactors. Some of the introductory material is presented in the first part of the book. Understandably, more extensive coverage is given in the remainder of the book to applications and design. Furthermore, several additional topics were included in the last part of the book—Other Reactor Topics. Some of the these topics are now required by ABET (Accreditation Board for Engineering and Technology) to be emphasized in course offerings.
The policy of most technical societies and publications is to use SI (metric) units or to list both the common British engineering unit and its SI equivalent. However, British units are primarily used in this book for the convenience of the majority of the reading audience. Readers who are more familiar and at ease with SI units are advised to refer to the Appendix.
Ultimately the practicing engineer must consider the following eight questions relative to the reaction/reactor.
REFERENCES
1. L. THEODORE and F. RICCI, “Mass Transfer Operations for the Practicing Engineer,” John Wiley & Sons, Inc., Hoboken, NJ, 2010.
2. L. THEODORE, “Heat Transfer for the Practicing Engineer,” John Wiley & Sons, Inc., Hoboken, NJ, 2012.
3. P. ABULENCIA and L. THEODORE, “Fluid Flow for the Practicing Engineer,” John Wiley & Sons, Inc., Hoboken, NJ, 2009.
PART I
INTRODUCTION
We must view with profound respect the infinite capacity of the human mind to resist the introduction of useful knowledge.
—Thomas Raynesford Lounsbury (1838–1915): Quoted in The Freshman and His College (1913), by Francis Cummins Lockwood.
THE PURPOSE of Part I can be found in its title. The book itself offers the reader the fundamentals of chemical reactor analysis with appropriate practical applications, and serves as an introduction to the specialized and more sophisticated texts in this area. The reader should realize that the contents are geared towards practitioners in this field, as well as students of science and engineering, not chemical engineers per se. Simply put, topics of interest to all practicing engineers have been included.
Part I serves as the introductory section to this book. It reviews engineering and science fundamentals that are an integral part of the field of chemical kinetics and chemical reactors. It consists of five chapters, as noted below:
Those individuals with a strong background in the above area(s) may choose to bypass Part I. Parts II and III are concerned with describing and designing the various classes of reactors.
CHAPTER 1
HISTORY OF CHEMICAL REACTIONS
INTRODUCTION
Until the last century, most chemicals were discovered more or less by accident. Their potential uses were based on short-term observations and their syntheses based on sketchy and simple theoretical ideas. Much of the recent progress in chemical syntheses occurred because of an increasing ability of chemists to determine the detailed molecular structure of substances and also to better understand the correlations between structure and properties. A review of how this industry arrived at its present state is presented below. Sections to follow include:
Early History
Recent History
The Chemical Industry Today
Microscopic vs Macroscopic Approach
EARLY HISTORY
As noted in the Introduction, most chemicals were discovered by accident. No one can assign with certainty a birth date to what one would classify as a “chemical reaction.” However, some have claimed that the first known chemical processes were carried out by the artisans of Egypt and China. These individuals worked with metals such as gold or copper, which often occur in nature in a pure state, but they learned how to “smelt” metallic ores by heating them with carbon-bearing materials. In addition, a primitive chemical technology arose in these cultures as dyes, potting glazes and glass making were discovered. Most of these inventors also developed astronomical, mathematical, and cosmological ideas that were used to explain some of the changes that are today considered chemical.
The first to consider such ideas scientifically were the Greeks at about 600 BC. They assumed that all matter was derived from water, which could solidify to earth or evaporate to air. This theory was later expanded into the idea that the world was composed from four elements: earth, water, air, and fire. It was Democritus who proposed that these elements combined to form atoms.
Aristotle believed that the elements formed a continuum of mass. He became the most influential of the Greek philosophers, and his ideas dominated science for nearly 1500 years. He believed that four qualities were found in nature: heat, cold, moisture, and dryness. He proposed that elements were made up of these with each element containing variable amounts of these qualities. These, in turn, combined to form materials that are visible. Because it was possible for each element to change, the elements could be combined because it was possible that material substances could be built up from the elements.
At approximately the same time a similar alchemy arose in China. The aim was to make gold, since it was believed to be a medicine that could offer long life or even immortality on anyone who consumed it. Nevertheless, the Chinese gained much practical chemical knowledge from incorrect theories.
After the decline of the Roman Empire, Greek writings were no longer studied in western Europe and the eastern Mediterranean. However, in the 7th and 8th centuries Arab conquerors spread Islam over Asia Minor, North Africa, and Spain. The Greek texts were translated into Arabic, and along with the rest of Greek learning, the ideas and practice of alchemy once again flourished.
A great intellectual reawakening began in western Europe in the 11th century. This occurred due to the cultural exchanges between Arab and Western scholars. Later, knowledge of Greek science was disseminated into Latin and ultimately reached all of Europe. Many of the manuscripts concerned alchemy.
Among the important substances discovered were alcohol and mineral acids such as hydrochloric, nitric, and sulfuric. The Chinese discovery of nitrates and the manufacture of gunpowder also came to the West through the Arabs. Gunpowder soon became a part of warfare. Thus, an effective chemical technology existed in Europe by the end of the 13th century.
During the 13th and 14th centuries the principles of Aristotle on scientific thought began to decline. The actual behavior of matter cast doubt on the relatively simple explanation Aristotle had prescribed. These doubts spread further after the invention of printing in 1450; these doubts increased into the 16th century.
It was during the first half of the 17th century that scientists began to study chemical reactions experimentally. Jan Baptista van Helmont laid the foundations of the law of conservation of mass. Van Helmont showed that in a number of reactions an “aerial” fluid was liberated which he defined as a “gas.” A new class of substances with their own physical properties was shown to exist. A kinetic–molecular theory of gases began to develop. Notable in this field were the experiments of Robert Boyle whose studies, later known as Boyle’s law, provided an equation describing the inverse relation between pressure and volume of gas (see the ideal gas law in Chapter 3).
During the 18th century, chemists noted that certain substances combined more easily with, or had a greater affinity for, a given chemical than did others. Tables were developed showing the relative affinities of different chemicals. The use of these tables made it possible to predict many chemical reactions before testing them in the laboratory.
It was Joseph Priestley who discovered oxygen. He realized that this gas was the component of ordinary air that was responsible for combustion and made animal respiration possible. Priestley told the chemist Antoine Laurent Lavoisier about his discovery of oxygen. He at once saw the significance of this substance and the door was opened for the chemical revolution that established modern chemistry.
RECENT HISTORY
Lavoisier showed by a series of unique experiments that combustion was due to the combination of a burning substance with oxygen and that when carbon was burned, fixed air (carbon dioxide) was produced. An earlier proposed substance phlogiston therefore did not exist, and the phlogiston theory soon disappeared to be replaced by the carbon cycle. Lavoisier used the laboratory balance to give quantitative support to his work and he used chemical equations in his papers. He further defined elements as substances that could not be decomposed by chemical means and firmly established the law of the conversation of mass. He developed a chemical nomenclature that is still used today and founded the first chemical journal.
By the beginning of the 19th century, it was shown that more man one compound could be formed between the same elements. Joseph Gay-Lussac demonstrated that the volume ratios of reacting gases were small whole numbers, implying the presence of atoms. Dalton assumed that when two elements combined, the resulting compound contained one atom of each. He arbitrarily assigned to hydrogen the atomic weight of 1 and could then calculate the relative atomic weight of oxygen. Applying this principle to other compounds, he calculated the atomic weights of other elements and actually drew up a table of the relative atomic weights of all the known elements.
In the early 19th century (1803), Dalton proposed his atomic theory. In 1811, Amedeo Avogadro made clear the distinction between atoms and molecules of elementary substances. In addition, the concepts of heat, energy, work, and temperature were developed. The first law of thermodynamics was set forth by Julius Robert von Mayer and the second law of thermodynamics was postulated by Rudolf Julius Emanuel Clausius and William Thomson (Lord Kelvin). Later in the century, Clausius, Ludwig Boltzmann, and James Clerk Maxwell related the ideal gas law in terms of a kinetic theory of matter. This led to the kinetics of reactions and the laws of chemical equilibrium.
It was Carnot who proposed the correlation between heat and work. Josiah Willard Gibbs discovered the phase rule and provided the theoretical basis of physical chemistry. And, it was Walther Hermann Nernst who proposed the third law of thermodynamics and contributed to the study of physical properties, molecular structures, and reaction rates. Jacobus Hendricus van’t Hoff related thermodynamics to chemical reactions and developed a method for establishing the order of reactions. Nearing the end of this century, Syante August Arrhenius investigated the increase in the rate of chemical reactions with an increase in temperature.
The development of chemical kinetics continued into the 20th century with the contributions to the study of molecular structures, reaction rates, and chain reactions by Irving Langmuir. Another advance in chemistry in the 20th century was the foundation of biochemistry, which began with the simple analysis of body fluids; methods were then rapidly developed for determining the nature and function of the most complex cell constituents. Biochemists later unraveled the genetic code and explained the function of the gene, the basis of all life. The field has now grown so vast that its study has become a new science—molecular biology.
THE CHEMICAL INDUSTRY TODAY
The growth of chemical industries and the training of professional chemists are intertwined. In the early 19th century during the Industrial Revolution, a number of universities were established in Germany. They drew students from all over the world and other universities soon followed suit. A large group of young chemists were thus trained just at the time when the chemical industry was beginning to exploit new discoveries. This interaction between the universities and the chemical industry resulted in the rapid growth of the organic chemical industry and provided Germany with scientific predominance in the field until World War I. Following the war, the German system was introduced into all industrial nations of the world, and chemistry and chemical industries progressed rapidly.
This scientific explosion has had an enormous influence on society. Processes were developed for synthesizing completely new substances that were either better than the natural ones or could replace them more cheaply. As the complexity of synthesized compounds increased, wholly new products appeared. Plastics and new textiles were developed, energy usage increased, and new drugs conquered whole classes of disease.
The progress of chemistry in recent years has been spectacular although the benefits of this progress have included corresponding liabilities. The most obvious dangers have come from nuclear weapons and radioactive materials, with their potential for producing cancer(s) in exposed individuals and mutations in their children. In addition, some pesticides have potential damaging effects. This led to the emergence of a new industry—environmental engineering. Mitigating these negative effects is one of the challenges the science community will have to meet in the future.(1,2)
MICROSCOPIC vs MACROSCOPIC APPROACH
The history of Unit Operations is interesting. Chemical engineering courses were originally based on the study of unit processes and/or industrial technologies; however, it soon became apparent that the changes produced in equipment from different industries were similar in nature, i.e., there was a commonality in the operations in the petroleum industry as with the utility industry. These similar operations became known as Unit Operations. This approach to chemical engineering was promulgated in the 1922 A. D. Little report (1922) submitted to the American Institute of Chemical Engineers (AIChE), and has, with varying degrees and emphasis, dominated the profession to this day.
The Unit Operations approach was adopted by the profession soon after its inception. During the 130+ years (since 1880) that the profession has been in existence as a branch of engineering, society’s needs have changed tremendously and so has chemical engineering.
The teaching of Unit Operations at the undergraduate level has remained relatively unchanged since the publication of several early- to mid-19th century texts; however, by the middle of the 20th century, there was a slow movement from the unit operation concept to a more theoretical treatment called transport phenomena or, more simply, engineering science. The focal point of this science is the rigorous mathematical description of all physical rate processes in terms of mass, heat, or momentum crossing phase boundaries. This approach took hold of the education/curriculum of the profession with the publication of the first edition of the Bird et al. book.(3) Some, including the author of this text, feel that this concept set the profession back several decades since graduating chemical engineers, in terms of training, were more applied physicists than traditional chemical engineers. There has fortunately been a return to the traditional approach to chemical engineering, primarily as a result of the efforts of ABET (Accreditation Board for Engineering and Technology—see also Chapter 21). Detractors to this pragmatic approach argue that this type of theoretical education experience provides answers to what and how, but not necessarily why, i.e., it provides a greater understanding of both fundamental physical and chemical processes. However, in terms of reality, nearly all chemical engineers are now presently involved with the why questions. Therefore, material normally covered earlier has been replaced, in part, with a new emphasis on solving design and open-ended problems; this approach is emphasized in this text.
The following paragraphs attempt to qualitatively describe the differences between the above two approaches. Both deal with the transfer of certain quantities (momentum, energy, and mass) from one point in a system to another. There are three basic transport mechanisms which potentially can be involved in a process. They are:
The first mechanism, radiative transfer, arises as a result of wave motion and is not considered, since it may be justifiably neglected in most engineering applications. The second mechanism, convective transfer, occurs simply because of bulk motion. The final mechanism, molecular diffusion, can be defined as the transport mechanism arising as a result of gradients. For example, momentum is transferred in the presence of a velocity gradient; energy in the form of heat is transferred because of a temperature gradient; and, mass is transferred in the presence of a concentration gradient. These molecular diffusion effects are described by phenomenological laws.(3)
Momentum, energy, and mass are all conserved. As such, each quantity obeys the conservation law within a system (including a chemical reactor) as provided in Equations (1.1) and (1.2):
(1.1)
This equation may also be written on a time rate basis
(1.2)
The conservation law may be applied at the macroscopic, microscopic, or molecular level.
One can best illustrate the differences in these methods with an example. Consider a system in which a fluid is flowing through a cylindrical tube reactor (see Figure 1.1) and define the system as the fluid contained within the reactor between points 1 and 2 at any time. If one is interested in determining changes occurring at the inlet and outlet of a reactor, the conservation law is applied on a “macroscopic” level to the entire system. The resultant equation (usually algebraic) describes the overall changes occurring to the system (or equipment). This approach is usually applied in the Unit Operation (or its equivalent) courses, an approach that is, as noted above, highlighted in this text and its three companion texts.(4–6)
Figure 1.1 Flow reactor system.
In the microscopic/transport phenomena approach, detailed information concerning the behavior within a system is required; this is occasionally requested of and by the engineer. The conservation law is then applied to a differential element within the system that is large compared to an individual molecule, but small compared to the entire system. The resulting equation is differential and can then be expanded via an integration in order to describe the behavior of the entire system.
The molecular approach involves the application of the conservation laws to individual molecules. This leads to a study of statistical and quantum mechanics—both of which are beyond the scope of this text. In any case, the description at the molecular level is of little value to the practicing engineer; however, the statistical averaging of molecular quantities in either a differential or finite element within a system can lead to a more meaningful description of the behavior of a system.
Both the microscopic and molecular approaches shed light on the physical reasons for the observed macroscopic phenomena. Ultimately, however, for the practicing engineer, these approaches may be justified but are akin to attempting to kill a fly with a machine gun. Developing and solving these equations (in spite of the advent of computer software packages) is typically not worth the trouble.
ILLUSTRATIVE EXAMPLE 1.1 Explain why the practicing engineer/scientist invariably employs the macroscopic approach in the solution of real world chemical reactor problems.
Solution. The macroscopic approach involves examining the relationship between changes occurring at the inlet and the outlet of a reacting system. This approach attempts to identify and solve problems found in the real world, and is more straightforward than, and preferable to, the more involved microscopic approach. The microscopic approach, which requires an understanding of all internal variations taking place within a reacting system that can lead up to an overall system result, simply may not be necessary.
REFERENCES
1. L. STANDER and L. THEODORE, “Environmental Regulatory Calculation Handbook,” John Wiley & Sons, Hoboken, NJ, 2008.
2. M. K. THEODORE and L. THEODORE, “Introduction to Environmental Management,” CRC Press/Taylor & Francis Group, Boca Raton, FL, 2010.
3. R. BIRD, W. STEWART, and E. LIGHTFOOT, “Transport Phenomena,” John Wiley & Sons, Hoboken, NJ, 1960.
4. P. ABULENCIA and L. THEODORE, “Fluid Flow for the Practicing Engineer,” John Wiley & Sons, Hoboken, NJ, 2009.
5. L. THEODORE and F. RICCI, “Mass Transfer Operations for the Practicing Engineer,” John Wiley & Sons, Hoboken, NJ, 2010.
6. L. THEODORE and F. RICCI, “Thermodynamics for the Practicing Engineer,” John Wiley & Sons, Hoboken, NJ, 2009.
CHAPTER 2
THE FIELD OF CHEMISTRY
INTRODUCTION
No text concerned with chemical reactions and chemical reactors would be complete without an introduction to the field of chemistry. In a general sense chemistry involves:
There is general consensus that chemistry deals with the combination of atoms, and physics with the forces between atoms. Atomic combination involves atomic forces, and it is one of the objects of physical chemistry to see how far the chemical interactions observed between atoms and molecules can be interpreted by means of the forces existing within and between atoms. The study of atomic structure provides information of why atoms combine. For many reasons, the development of modern views concerning the structure of the atom is bound up very closely with several peripheral subjects including organic chemistry, physical chemistry, and thermodynamics.
In addition to the atomic or molecular approach to describing chemical reactions, which is treated superficially in this text, another important area involves the applications of thermodynamics, a topic which deals primarily with energy changes.(1) The treatment of the first and second laws of thermodynamics, including thermochemistry(1) provides an adequate basis for a consideration of the chemical change associated with gaseous, liquid, and solid states of matter.
A chemical reaction is a process by which atoms or groups of atoms are combined and/or redistributed, resulting in a change in the molecular composition and properties. The products obtained from reactants depend on the condition under which a chemical reaction occurs. The scientist and engineer have shown that although products may vary with changing conditions, some properties remain constant during any chemical reaction. These constant properties, called “conserved” properties, include the number of each kind of atom present, the electrical charge, and the total mass (conservation laws for mass).
Chemical symbols must also be understood. Almost all substances are made up of some combination of the aforementioned atoms of the elements found in nature. Rather than full names, scientists identify elements with one- or two-letter symbols. Some common elements and their symbols are carbon, C; oxygen, O; nitrogen, N; hydrogen, H; chlorine, Cl; and, sulfur, S. These chemical symbols are derived from the letters of the name of the element. The first letter of the symbol is capitalized, and the second (if applicable) is lowercase. Symbols for some elements known from ancient times come from earlier, usually Latin, names: for example, Cu from cuprum (copper), Ag from aurum (gold), and Fe from ferrum (iron). This set of symbols, in referring to elements, is used universally.
Section topics following this introduction to the field of chemistry include:
Inorganic Chemistry
Organic Chemistry
Physical Chemistry
Other Chemistry Topics
Analysis Procedures
INORGANIC CHEMISTRY
Inorganic chemistry is that field of chemistry which is concerned with chemical reactions and properties of all the chemical elements and their compounds, with the exception of hydrocarbons (compounds composed of carbon and hydrogen) and their derivatives. The subject of carbon–hydrogen compounds is defined as organic chemistry, a topic discussed in the next section.
Inorganic chemistry is too vast a subject to form a convenient unit of study; the term would be of little importance except for the tendency in most engineering and science schools to call courses “Inorganic Chemistry” when a better title might be “Elementary Chemistry.” The subject matter of such courses includes the elementary laws of chemistry and its symbols and nomenclature, and an introduction to the experimental methods that are important in experimental chemistry. The student is introduced to such fundamental chemical reaction topics as valence, ionization, reactivity, atomic theory, and the kinetic theory of gases. The properties and reactions of substances in aqueous solution also receive attention. Modern inorganic chemistry overlaps parts of many other scientific fields, including biochemistry, metallurgy, mineralogy, and solid-state physics. Finally, an increased understanding of the chemical behavior of the elements and of inorganic compounds has led to the discovery of a wide variety of new synthesizing techniques and the discovery of many new classes of inorganic substances.
ORGANIC CHEMISTRY
The branch of chemistry in which carbon compounds and their reactions are studied is defined as organic chemistry. A wide variety of classes of substances—such as drugs, vitamins, plastics, natural and synthetic fibers, as well as carbohydrates, proteins, and fats—consist of organic molecules. This subject involves:
Organic chemistry has had a profound effect on society: it has improved natural materials and has synthesized natural and artificial materials that have, in turn, improved health, increased comfort, and added to the convenience associated with nearly every product manufactured today.
As noted earlier, the molecular formula of a compound indicates the number of each kind of atom in a molecule of that substance. Fructose, (C6H12O6) consists of molecules containing six carbon atoms, 12 hydrogen atoms, and six oxygen atoms. Because at least 15 other compounds have this same molecular formula, one may distinguish one molecule from another by employing a structural formula to show the spatial arrangement of the atom. Note that an analytical analysis that indicates the percentage of carbon, hydrogen, and oxygen cannot distinguish fructose from ribose (C5H10O5), another sugar in which the ratios of elements are the same, namely 1:2:1.
The ability of carbon to form covalent bonds with other carbon atoms in long chains and rings distinguishes carbon from all other elements. Other elements are not known to form chains of greater than eight similar atoms. This property of carbon, and the fact that carbon nearly always forms four bonds to other atoms, accounts for the large number of known compounds. At least 80 percent of the 5 million reported chemical compounds contain carbon.
There are various classes of organic compounds, including:
Additional “classes” of organics include:
Other atoms, such as chlorine, oxygen, and nitrogen, may be substituted for hydrogen in an alkane, providing the correct number of chemical bonds is allowed—chlorine forming one bond to other atoms, oxygen forming two bonds to other atoms, and nitrogen three bonds. The chlorine atom in ethyl chloride, the –OH group in ethyl alcohol, and the –NH2 group in ethyl amine are called functional groups. Functional groups determine many of the chemical properties of compounds. Many of the chlorine bearing compounds are known to be carcinogenic and/or toxic.(2)
Regarding sources of organic compounds, coal tar was once the only source of aromatic and some heterocyclic compounds. Petroleum was the source of aliphatic compounds since it contains such substances as gasoline, kerosene, and lubricating oil. Natural gas provides (primarily) methane and ethane. These three categories of natural fossil compounds are still the major sources of organic compounds. When petroleum is not available, a chemical industry can be based on acetylene (if available), which in turn can be synthesized from limestone and coal. During World War II, Germany was forced into employing this process when it lost reliable petroleum and natural-gas supplies in Africa.
Covalent organic compounds are distinguished from inorganic salts by low melting points and boiling points. Hydrocarbons have low specific gravities (see also next chapter)—about 0.8 compared to water, 1.0—but functional groups may increase the densities of organic compounds. Only a few organic compounds possess specific gravities in excess of unity, e.g., carbon tetrachloride.
The practicing engineer usually designs organic reactions to be carried out at optimum conditions to produce maximum conversion or yields (terms to be defined later). One often resorts to catalysts, whether or not the reaction is reversible, and attempts to take advantage of equilibrium positions.(1) Also, catalysts are frequently essential for rapid chemical reactions.
PHYSICAL CHEMISTRY
Although chemistry deals with the combination of atoms and physics with the forces between atoms, the object of physical chemistry is to see how far the chemical interactions observed between atoms and molecules can be interpreted by means of the forces existing within and between atoms. Thus, one of the objectives of physical chemistry is to apply measurements of physical properties, such as density, surface tension, refractive index, dielectric constant, magnetism, and optical activity, to the description of chemical structure. Much of the knowledge provided in the earlier topics of this chapter can now be turned to use in this connection.
Others have defined physical chemistry as that field of science that applies the laws of physics to elucidate the properties of chemical substances and clarify the characteristics of chemical phenomena. The term physical chemistry is usually applied to the study of the physical properties of substances, such as vapor pressure, surface tension, viscosity, refractive index, density, and crystallography, as well as to the study of the so-called classical aspects of the behavior of chemical systems, such as thermal properties, equilibria, rates of reactions, mechanisms of reactions, and ionization phenomena. In its more theoretical aspects, physical chemistry attempts to explain spectral properties of substances in terms of fundamental quantum theory, the interaction of energy with matter, the nature of chemical bonding, the relationships correlating the number of energy states of electrons in atoms and molecules with the observable properties shown by these systems, and the electrical, thermal, and mechanical effects of individual electrons and protons on solids and liquids.(3)
Interestingly, physical chemistry can be subdivided into the study of chemical thermodynamics, chemical kinetics, the gaseous state, the liquid state, solutions, the solid state, electrochemistry, colloid chemistry, photochemistry, and statistical thermodynamics. Although providing details on each of these topics is beyond the scope of this text, a brief introductory comment on chemical kinetics is warranted at this time. This field, as the reader shall find out, is concerned with the rates of chemical reactions as a function of the concentration of the reacting species, of the products of the reaction, of any catalysts and inhibitors, of various solvent media, of temperature, and of all other variables that can affect the rate of reaction. It also seeks to relate the manner in which the reaction rate varies with time (and position) to the molecular nature of the rate-controlling intermolecular collisions involved in generating the products of reaction. Most reactions involve a series of stepwise processes, the sum of which corresponds to the overall observed reaction proportions (or stoichiometry) in which the reactants combine and the products form. Fortunately, only one of these steps is generally the rate-controlling one. By determining the nature of the rate-controlling process from a mathematical analysis of reaction kinetics data and by investigating how the reaction conditions affect the step, one can often deduce the mechanism of a reaction (see also Chapter 17).
OTHER CHEMISTRY TOPICS
There are other topics that fall under the chemistry umbrella. Some details are provided below.
ANALYSIS PROCEDURES
Chemical analysis is concerned with the procedures and techniques used to identify and quantify the chemical composition of a sample of a substance. A chemist executing a qualitative analysis seeks to identify the substances in the sample. A quantitative analysis is an attempt to determine the quantity or concentration of a specific substance in the sample.
The measurement of chemical composition is necessary throughout the chemical industry, environmental regulatory government, and many other fields of science.
The practicing engineer is often required to analyze such diverse materials as stainless steel, beer, a fingernail, a rose petal, smoke, aspirin, paper, etc. The determination of the identity or quantity of a constituent in such materials is preceded by a sampling step—the selection of the amount and uniformity of material required for the analysis—and by the separation from the sample of either the desired constituent or the undesired, interfering constituents. Some typical analytical techniques are presented in Table 2.1.
TABLE 2.1Classes of Specialized Analysis
Clinical
Environmental
Forensic
Geochemical
Inorganic
Oceanographic
Organic
Petroleum
Pharmaceutical
Polymer
Spatial
Surface
Trace
Although chromatography is the most generally applicable of the separation methods available to the practicing engineer, mere are a host of other procedures. These are detailed in Table 2.2.
TABLE 2.2Separation Methods
Chromatography
Fractional distillation
(5)
Precipitation
(5)
Solvent extraction
(5)
REFERENCES
1. L. THEODORE, F. RICCI, and T. VAN VLIET, “Thermodynamics for the Practicing Engineer,” John Wiley & Sons, Hoboken, NJ, 2009.
2. J. SANTOLERI, J. REYNOLDS, and L. THEODORE, “Introduction to Hazardous Waste Incineration,” 2nd edition, John Wiley & Sons, Hoboken, NJ, 2000.
3. Author, title, source and date unknown.
4. L. THEODORE, “Air Pollution Control Equipment Calculations,” John Wiley & Sons, Hoboken, NJ, 2008.
5. L. THEODORE and F. RICCI, “Mass Transfer Operation for the Practicing Engineer,” John Wiley & Sons, Hoboken, NJ, 2010.
CHAPTER 3
PROCESS VARIABLES*
INTRODUCTION
The author originally considered the title “State, Physical, and Chemical Properties” for this chapter; however, since these three properties have been used interchangeably and have come to mean different things to different people, it was decided to simply employ the title “Process Variables.” The three aforementioned properties were therefore integrated into this all-purpose title eliminating the need for differentiating between the three.
This chapter provides a review of some basic concepts from physics, chemistry, and engineering in preparation for material that is covered in later chapters. All of these topics are vital to chemical kinetics and reactor applications. Because many of these topics are unrelated to each other, this chapter admittedly lacks the cohesiveness that chapters covering a single topic might have. This is usually the case when basic material from such widely differing areas of knowledge such as physics, chemistry, and engineering is surveyed. Though these topics are widely divergent and covered with varying degrees of thoroughness, all of them will find later use in this text. If additional information on these review topics is needed, the reader is directed to the literature in the reference section of this chapter.
Topics to be addressed include: temperature, pressure, moles and molecular weights, mass and volume, viscosity, heat capacity, thermal conductivity, Reynolds number, pH, vapor pressure, ideal gas law, latent enthalpy effects, and chemical reaction velocity constant. The chapter concludes with a section on property estimation.
ILLUSTRATIVE EXAMPLE 3.1 Discuss the traditional difference between chemical and physical properties.
Solution. Every compound has a unique set of properties that allows one to recognize and distinguish it from other compounds. These properties can be grouped into two main categories: physical and chemical. Physical properties are defined as those that can be measured without changing the identity and composition of the substance. Key properties include viscosity, density, surface tension, melting point, boiling point, etc. Chemical properties are defined as those that may be altered via chemical reaction to form other compounds or substances. Key chemical properties include upper and lower flammability limits, enthalpy of reaction, autoignition temperature, and others.
These properties may be further divided into two categories—intensive and extensive. Intensive properties are not a function of the quantity of the substance, while extensive properties depend on the quantity of the substance.
TEMPERATURE
Whether in the gaseous, liquid, or solid state, all molecules possess some degree of kinetic energy, i.e., they are in constant motion—vibrating, rotating, or translating. The kinetic energies of individual molecules cannot be measured, but the combined effect of these energies in a very large number of molecules can. This measurable quantity is known as temperature; it is a macroscopic concept only and as such does not exist on the molecular level.
Temperature can be measured in many ways; the most common method makes use of the expansion of mercury (usually encased inside a glass capillary tube) with increasing temperature. (In many thermal applications, however, thermocouples or thermistors are more commonly employed.) The two most commonly used temperature scales are the Celsius (or Centigrade) and Fahrenheit scales. The Celsius scale is based on the boiling and freezing points of water at 1-atm (atmosphere) pressure; to the former, a value of 100°C is assigned, and to the latter, a value of 0°C. On the older Fahrenheit scale, these temperatures correspond to 212°F and 32°F, respectively. Equations (3.1) and (3.2) show the conversion from one scale to the other:
(3.1)
(3.2)
where
a temperature on the Fahrenheit scale
a temperature on the Celsius scale
Experiments with gases at low-to-moderate pressures (up to a few atmospheres) have shown that, if the pressure is kept constant, the volume of a gas and its temperature are linearly related via Charles’ law (see development later in chapter) and that a decrease of 0.3663% or (1/273) of the initial volume is experienced for every temperature drop of 1°C. These experiments were not extended to very low temperatures, but if the linear relationship were extrapolated, the volume of the gas would theoretically be zero at a temperature of approximately -273°C or -460°F. This temperature has become known as absolute zero and is the basis for the definition of two absolute temperature scales. (An absolute scale is one which does not allow negative quantities.) These absolute temperature scales are the Kelvin (K) and Rankine (°R) scales; the former is defined by shifting the Celsius scale by 273°C so that 0 K is equal to -273°C; Equation (3.3) shows this relationship:
(3.3)
The Rankine scale is defined by shifting the Fahrenheit scale 460°, so that
(3.4)
The relationships among the various temperature scales are shown in Figure 3.1.
Figure 3.1 Temperature scales.
ILLUSTRATIVE EXAMPLE 3.2 Perform the following temperature conversions:
Solution Employ Equations (3.3) and (3.4).
PRESSURE
Molecules in the gaseous state possess a high degree of translational kinetic energy, which means they are able to move quite freely throughout the body of the gas. If the gas is in a container of some type, the molecules are constantly bombarding the walls of the container. The macroscopic effect of this bombardment by a tremendous number of molecules—enough to make the effect measurable—is called pressure. The natural units of pressure are force per unit area. In the example of the gas in a container, the unit area is a portion of the inside solid surface of the container wall and the force, measured perpendicularly to the unit area, is the result of the molecules hitting the unit area and giving up momentum during the sudden change of direction.
There are a number of different methods used to express a pressure measurement. Some of them are natural units, i.e., based on a force per unit area, as with pound (force) per square inch (abbreviated lbf/in2 or psi) or dyne per square centimeter (dyn/cm2). Others are based on a fluid height, such as inches of water (in H2O) or millimeters of mercury (mm Hg); units such as these are convenient when the pressure is indicated by a difference between two levels of a liquid as in a manometer or barometer. This measurement is based on the pressure at the base of a column (height) of fluid. Barometric pressure and atmospheric pressure are synonymous and measure the ambient air pressure. Standard barometric pressure is the average atmospheric pressure at sea level, 45° north latitude at 32°F. It is used to define another unit of pressure called the atmosphere (atm). Standard barometric pressure is 1 atm and is equivalent to 14.696 psi and 29.921 in Hg. As one might expect, barometric pressure varies with weather and altitude.
Measurements of pressure by most gauges indicate the difference in pressure either above or below that of the atmosphere surrounding the gauge. Gauge pressure is the pressure indicated by such a device. If the pressure in the system measured by the gauge is greater than the pressure prevailing in the atmosphere, the gauge pressure is expressed positively. If lower than atmospheric pressure, the gauge pressure is a negative quantity; the term vacuum designates a negative gauge pressure. Gauge pressures are often identified by the letter g after the pressure unit; for example, psig (pounds per square inch gauge) is a gauge pressure in psi units.
Since gauge pressure is the pressure relative to the prevailing atmospheric pressure, the sum of the two gives the absolute pressure, indicated by the letter a after the unit [e.g., psia (pounds per square inch absolute)]:
(3.5)
The absolute pressure scale is absolute in the same sense that the absolute temperature scale is absolute, i.e., a pressure of zero psia is the lowest possible pressure theoretically achievable—a perfect vacuum.
ILLUSTRATIVE EXAMPLE 3.3 Consider the following pressure calculations.
Solution
(3.5)
MOLES AND MOLECULAR WEIGHTS
An atom consists of protons and neutrons in a nucleus surrounded by electrons. An electron has such a small mass relative to that of the proton and neutron that the weight of the atom (called the atomic weight) is approximately equal to the sum of the weights of the particles in its nucleus. Atomic weight may be expressed in atomic mass units (amu) per atom or in grams per gram · atom. One gram · atom contains 6.02 × 1023 atoms (Avogadro’s number). The atomic weights of the elements are listed in Table 3.1.
TABLE 3.1Atomic Weights of the Elementsa,b
The molecular weight (MW) of a compound is the sum of the atomic weights of the atoms that make up the molecule. Atomic mass units per molecule (amu/molecule) or grams per gram · mole (g/gmol) are used for molecular weight. One gram · mole (gmol) contains an Avogadro number of molecules. For the English system, a pound · mole (lbmol) contains 454 × 6.023 × 1023 molecules.
Molal units are used extensively in chemistry calculations as they greatly simplify material balances where chemical reactions are occurring. For mixtures of substances (gases, liquids, or solids), it is also convenient to express compositions in mole fractions or mole percentages instead of mass fractions. The mole fraction is the ratio of the number of moles of one component to the total number of moles in the mixture. Equations (3.6)–(3.9) express these relationships:
(3.6)
(3.7)
(3.8)
(3.9)
The reader should note that, in general, mass fraction (or percent) is not equal to mole fraction (or percent).
ILLUSTRATIVE EAMPLE 3.4 If a 55-gal tank contains 20.0 lb of water, (1) how many pound · moles of water does it contain? (2) how many gram · moles does it contain? and, (3) how many molecules does it contain?
Solution. The molecular weight of the water (H2O) is
Therefore
MASS AND VOLUME
The density
