Combustion
Fourth Edition
Irvin Glassman Richard A. Yetter
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08 09 10 9 8 7 6 5 4 3 2 1
tributed so much to the atmosphere for learning and the technical contributions that emanated from Princeton’s Combustion Research Laboratory.
of your knowledge.
If he (the teacher) is wise he does not bid you to enter the house of his wisdom, but leads you to the threshold of your own mind.
The astronomer may speak to you of his understanding of space, but he cannot give you his understanding.
And he who is versed in the science of numbers can tell of the regions of weight and measures, but he cannot conduct you hither.
For the vision of one man lends not its wings to another man .
Gibran, The Prophet
The reward to the educator lies in his pride in his students ’ accomplishments. The richness of that reward is the satisfaction in knowing the frontiers of knowledge have been extended.
D. F. Othmer
CHAPTER 1. CHEMICAL THERMODYNAMICS AND
FLAME TEMPERATURES 1
A. Introduction 1
B. Heats of reaction and formation 1
C. Free energy and the equilibrium constants 8
D. Flame temperature calculations 16
1. Analysis 16
2. Practical considerations 22
E. Sub- and super sonic combustion thermodynamics 32
1. Comparisons 32
2. Stagnation pressure considerations 33
Problems 36
CHAPTER 2. CHEMICAL KINETICS 43
A. Introduction 43
B. Rates of reactions and their temperature dependence 43
1. The Arrhenius rate expression 45
2. Transition state and recombination rate theories 47 C. Simultaneous interdependent reactions 52
D. Chain reactions 53
E. Pseudo-fi rst-order reactions and the “ fall-off ” range 57
F. The partial equilibrium assumption 60
G. Pressure effect in fractional conversion 61 H. Chemical kinetics of large reaction mechanisms 62
1. Sensitivity analysis 63
2. Rate of production analysis 65
3. Coupled thermal and chemical reacting systems 66
4. Mechanism simplifi cation 68
Problems 69
Prologue xvii Preface xix
CHAPTER 3. EXPLOSIVE AND GENERAL OXIDATIVE
CHARACTERISTICS OF FUELS 75
A. Introduction 75
B. Chain branching reactions and criteria for explosion 75 C. Explosion limits and oxidation characteristics of hydrogen 83 D. Explosion limits and oxidation characteristics of carbon
monoxide 91
E. Explosion limits and oxidation characteristics of hydrocarbons 98
1. Organic nomenclature 99
2. Explosion limits 103
3. “ Low-temperature ” hydrocarbon oxidation mechanisms 106
F. The oxidation of aldehydes 110
G. The oxidation of methane 112
1. Low-temperature mechanism 112
2. High-temperature mechanism 113
H. The oxidation of higher-order hydrocarbons 117
1. Aliphatic hydrocarbons 117
2. Alcohols 127
3. Aromatic hydrocarbons 129
4. Supercritical effects 139
Problems 141
CHAPTER 4. FLAME PHENOMENA IN PREMIXED
COMBUSTIBLE GASES 147
A. Introduction 147
B. Laminar fl ame structure 151
C. The laminar fl ame speed 153
1. The theory of Mallard and Le Chatelier 156
2. The theory of Zeldovich, Frank-Kamenetskii, and Semenov 161 3. Comprehensive theory and laminar fl ame structure analysis 168 4. The laminar fl ame and the energy equation 176
5. Flame speed measurements 176
6. Experimental results: physical and chemical effects 185 D. Stability limits of laminar fl ames 191
1. Flammability limits 192
2. Quenching distance 200
3. Flame stabilization (low velocity) 201
4. Stability limits and design 207
E. Flame propagation through stratifi ed combustible mixtures 211 F. Turbulent reacting fl ows and turbulent fl ames 213 1. The rate of reaction in a turbulent fi eld 216
2. Regimes of turbulent reacting fl ows 218
3. The turbulent fl ame speed 231
G. Stirred reactor theory 235 H. Flame stabilization in high-velocity streams 240
I. Combustion in small volumes 250
Problems 254
CHAPTER 5. DETONATION 261
A. Introduction 261
1. Premixed and diffusion fl ames 261
2. Explosion, defl agration, and detonation 261
3. The onset of detonation 262
B. Detonation phenomena 264
C. Hugoniot relations and the hydrodynamic theory of
detonations 265
1. Characterization of the Hugoniot curve and the uniqueness of the
C–J point 266
2. Determination of the speed of sound in the burned gases for conditions above the C–J point 276 3. Calculation of the detonation velocity 282 D. Comparison of detonation velocity calculations with
experimental results 286
E. The ZND structure of detonation waves 293 F. The structure of the cellular detonation front and other
detonation phenomena parameters 297
1. The cellular detonation front 297
2. The dynamic detonation parameters 301
3. Detonation limits 302
G. Detonations in nongaseous media 306
Problems 307
CHAPTER 6. DIFFUSION FLAMES 311
A. Introduction 311
B. Gaseous fuel jets 311
1. Appearance 312
2. Structure 316
3. Theoretical considerations 318
4. The Burke–Schumann development 322
5. Turbulent fuel jets 329
C. Burning of condensed phases 331
1. General mass burning considerations and the evaporation coeffi cient 332 2. Single fuel droplets in quiescent atmospheres 337
D. Burning of droplet clouds 364
E. Burning in convective atmospheres 365
1. The stagnant fi lm case 365 2. The longitudinally burning surface 367
3. The fl owing droplet case 369
4. Burning rates of plastics: The small B assumption and radiation effects 372
Problems 374
CHAPTER 7. IGNITION 379
A. Concepts 379
B. Chain spontaneous ignition 382
C. Thermal spontaneous ignition 384
1. Semenov approach of thermal ignition 384 2. Frank-Kamenetskii theory of thermal ignition 389
D. Forced ignition 395
1. Spark ignition and minimum ignition energy 396 2. Ignition by adiabatic compression and shock waves 401
E. Other ignition concepts 402
1. Hypergolicity and pyrophoricity 403
2. Catalytic ignition 406
Problems 407
CHAPTER 8. ENVIRONMENTAL COMBUSTION
CONSIDERATIONS 409
A. Introduction 409
B. The nature of photochemical smog 410
1. Primary and secondary pollutants 411
2. The effect of NO x 411
3. The effect of SO x 415
C. Formation and reduction of nitrogen oxides 417
1. The structure of the nitrogen oxides 418
2. The effect of fl ame structure 419
3. Reaction mechanisms of oxides of nitrogen 420
4. The reduction of NO x 436
D. SOx emissions 441
1. The product composition and structure of sulfur compounds 442 2. Oxidative mechanisms of sulfur fuels 444
E. Particulate formation 457
1. Characteristics of soot 458
2. Soot formation processes 459
3. Experimental systems and soot formation 460
4. Sooting tendencies 462
5. Detailed structure of sooting fl ames 474
6. Chemical mechanisms of soot formation 478 7. The infl uence of physical and chemical parameters on soot formation 482
F. Stratospheric ozone 485
1. The HO x catalytic cycle 486
2. The NO x catalytic cycle 487
3. The ClO x catalytic cycle 489
Problems 491
CHAPTER 9. COMBUSTION OF NONVOLATILE FUELS 495 A. Carbon char, soot, and metal combustion 495
B. Metal combustion thermodynamics 496
1. The criterion for vapor-phase combustion 496 2. Thermodynamics of metal–oxygen systems 496 3. Thermodynamics of metal–air systems 509
4. Combustion synthesis 513
C. Diffusional kinetics 520
D. Diffusion-controlled burning rate 522
1. Burning of metals in nearly pure oxygen 524 2. Burning of small particles – diffusion versus kinetic limits 527
3. The burning of boron particles 530
4. Carbon particle combustion (C. R. Shaddix) 531 E. Practical carbonaceous fuels (C. R. Shaddix) 534
1. Devolatilization 534
2. Char combustion 539
3. Pulverized coal char oxidation 540
4. Gasifi cation and oxy-combustion 542
F. Soot oxidation (C. R. Shaddix) 545
Problems 548
APPENDIXES 551 APPENDIX A. THERMOCHEMICAL DATA AND
CONVERSION FACTORS 555
Table A1. Conversion factors and physical constants 556 Table A2. Thermochemical data for selected chemical
compounds 557
Table A3. Thermochemical data for species included in reaction list of Appendix C 646 APPENDIX B. ADIABATIC FLAME TEMPERATURES OF
HYDROCARBONS 653
Table B1. Adiabatic fl ame temperatures 653
APPENDIX C. SPECIFIC REACTION RATE CONSTANTS 659
Table C1. H 2 /O 2 mechanism 659
Table C2. CO/H 2 /O 2 mechanism 661
Table C3. CH 2 O/CO/H 2 /O 2 mechanism 662 Table C4. CH 3 OH/CH 2 O/CO/H 2 /O 2 mechanism 663 Table C5. CH 4 /CH 3 OH/CH 2 O/CO/H 2/O2 mechanism 665 Table C6. C 2 H 6 /CH 4 /CH 3 OH/CH 2 O/CO/H 2 /O 2 mechanism 668 Table C7. Selected reactions of a C 3 H 8 oxidation mechanism 673 Table C8. N x O y /CO/H 2 /O 2 mechanism 677 Table C9. HCl/N x O y /CO/H 2 /O 2 mechanism 683 Table C10. O 3/Nx O y /CO/H 2 /O 2 mechanism 684 Table C11. SO x /N xOy /CO/H 2 /O 2 mechanism 685 APPENDIX D. BOND DISSOCIATION ENERGIES OF
HYDROCARBONS 693
Table D1. Bond dissociation energies of alkanes 694 Table D2. Bond dissociation energies of alkenes, alkynes, and
aromatics 695
Table D3. Bond dissociation energies of C/H/O compounds 698 Table D4. Bond dissociation energies of sulfur-containing
compounds 699
Table D5. Bond dissociation energies of nitrogen-containing
compounds 700
Table D6. Bond dissociation energies of halocarbons 702 APPENDIX E. FLAMMABILITY LIMITS IN AIR 703 Table E1. Flammability limits of fuel gases and vapors in air at
25 ° C and 1 atm 704
APPENDIX F. LAMINAR FLAME SPEEDS 713 Table F1. Burning velocities of various fuels at 25 ° C air-fuel
temperature (0.31 mol% H 2 O in air). Burning velocity S as a function of equivalence ratio φ in cm/s 714 Table F2. Burning velocities of various fuels at 100 ° C air-fuel
temperature (0.31 mol% H 2 O in air). Burning velocity S as a function of equivalence ratio φ in cm/s 719 Table F3. Burning velocities of various fuels in air as a function
of pressure for an equivalence ratio of 1 in cm/s 720 APPENDIX G. SPONTANEOUS IGNITION TEMPERATURE
DATA 721
Table G1. Spontaneous ignition temperature data 722
APPENDIX H. MINIMUM SPARK IGNITION ENERGIES AND QUENCHING DISTANCES 743 Table H1. Minimum spark ignition energy data for fuels in air at
1 atm pressure 744
APPENDIX I. PROGRAMS FOR COMBUSTION KINETICS 747
A. Thermochemical parameters 747
B. Kinetic parameters 747
C. Transport parameters 748
D. Reaction mechanisms 748
E. Thermodynamic equilibrium 750
F. Temporal kinetics (Static and fl ow reactors) 752
G. Stirred reactors 753
H. Shock tubes 754
I. Premixed fl ames 754
J. Diffusion fl ames 756
K. Boundary layer fl ow 756
L. Detonations 756
M. Model analysis and mechanism reduction 756
Author Index 759
Subject Index 769
xvii This 4th Edition of “ Combustion ” was initiated at the request of the publisher, but it was the willingness of Prof. Richard Yetter to assume the responsibil- ity of co-author that generated the undertaking. Further, the challenge brought to mind the oversight of an acknowledgment that should have appeared in the earlier editions.
After teaching the combustion course I developed at Princeton for 25 years, I received a telephone call in 1975 from Prof. Bill Reynolds, who at the time was Chairman of the Mechanical Engineering Department at Stanford. Because Stanford was considering developing combustion research, he invited me to present my Princeton combustion course during Stanford’s summer semester that year. He asked me to take in consideration that at the present time their graduate students had little background in combustion, and, further, he wished to have the opportunity to teleconference my presentation to Berkeley, Ames, and Sandia Livermore. It was an interesting challenge and I accepted the invi- tation as the Standard Oil of California Visiting Professor of Combustion.
My early lectures seemed to receive a very favorable response from those participating in the course. Their only complaint was that there were no notes to help follow the material presented. Prof. Reynolds approached me with the request that a copy of lecture notes be given to all the attendees. He agreed it was not appropriate when he saw the handwritten copies from which I pre- sented the lectures. He then proposed that I stop all other interactions with my Stanford colleagues during my stay and devote all my time to writing these notes in the proper grammatical and structural form. Further, to encourage my writing he would assign a secretary to me who would devote her time organiz- ing and typing my newly written notes. Of course, the topic of a book became evident in the discussion. Indeed, eight of the nine chapters of the fi rst edition were completed during this stay at Stanford and it took another 2 years to fi n- ish the last chapter, indexes, problems, etc., of this fi rst edition. Thus I regret that I never acknowledged with many thanks to Prof. Reynolds while he was alive for being the spark that began the editions of “ Combustion ” that have already been published.
“ Combustion, 4th Edition ” may appear very similar in format to the 3 rd Edition. There are new sections and additions, and many brief insertions that are the core of important modifi cations. It is interesting that the content of these insertions emanated from an instance that occurred during my Stanford presentation. At one lecture, an attendee who obviously had some experience
in the combustion fi eld claimed that I had left out certain terms that usually appear in one of the simple analytical developments I was discussing. Sur- prisingly, I subconscientiously immediately responded “ You don’t swing at the baseball until you get to the baseball park! ” The response, of course, drew laughter, but everyone appeared to understand the point I was trying to make.
The reason of bringing up this incident is that it is important to develop the understanding of a phenomenon, rather than all its detailed aspects. I have always stressed to my students that there is a great difference between knowing something and understanding it. The relevant point is that in various sections there have been inserted many small, important modifi cations to give greater understanding to many elements of combustion that appear in the text. This type of material did not require extensive paragraphs in each chapter section.
Most chapters in this edition contain, where appropriate, this type of important improvement. This new material and other major additions are self-evident in the listings in the Table of Contents.
My particular thanks go to Prof. Yetter for joining me as co-author, for his analyzing and making small poignant modifi cations of the chapters that appeared in the earlier additions, for contributing new material not covered in these earlier additions and for further developing all the appendixes. Thanks also go to Dr. Chris Shaddix of Sandia Livermore who made a major contribu- tion to Chapter 9 with respect to coal combustion considerations. Our gracious thanks go to Mary Newby of Penn State who saw to the fi nal typing of the complete book and who offered a great deal of general help. We would never have made it without her. We also wish to thank our initial editor at Elsevier, Joel Stein, for convincing us to undertake this edition of “ Combustion ” and our fi nal Editor, Matthew Hart, for seeing this endeavor through.
The last acknowledgments go to all who are recognized in the Dedication.
I initiated what I called Princeton’s Combustion Research Laboratory when I was fi rst appointed to the faculty there and I am pleased that Prof. Fred Dryer now continues the philosophy of this laboratory. It is interesting to note that Profs. Dryer and Yetter and Dr. Shaddix were always partners of this laboratory from the time that they entered Princeton as graduate students. I thank them again for being excellent, thoughtful, and helpful colleagues through the years.
Speaking for Prof. Yetter as well, our hope is that “ Combustion, 4th Edition ” will be a worthwhile contributing and useful endeavor.
Irvin Glassman December 2007
xix When approached by the publisher Elsevier to consider writing a 4th Edition of Combustion, we considered the challenge was to produce a book that would extend the worthiness of the previous editions. Since the previous editions served as a basis of understanding of the combustion fi eld, and as a text to be used in many class courses, we realized that, although the fundamentals do not change, there were three factors worthy of consideration: to add and extend all chapters so that the fundamentals could be clearly seen to provide the background for helping solve challenging combustion problems; to enlarge the Appendix section to provide even more convenient data tables and com- putational programs; and to enlarge the number of typical problem sets. More important is the attempt to have these three factors interact so that there is a deeper understanding of the fundamentals and applications of each chapter.
Whether this concept has been successful is up to the judgment of the reader.
Some partial examples of this approach in each chapter are given by what follows.
Thus, Chapter 1, Chemical Thermodynamics and Flame Temperatures, is now shown to be important in understanding scramjets. Chapter 2, Chemical Kinetics, now explains how sensitivity analyses permit easier understanding in the analysis of complex reaction mechanisms that endeavor to explain environ- mental problems. There are additions and changes in Chapter 3, Explosive and General Oxidative Characteristics of Fuels, such as consideration of wet CO combustion analysis, the development procedure of reaction sensitivity analysis and the effect of supercritical conditions. Similarly the presentation in Chapter 4, Flame Phenomena in Premixed Combustible Gases, now considers fl ame propagation of stratifi ed fuel–air mixtures and fl ame spread over liquid fuel spills. A point relevant to detonation engines has been inserted in Chapter 5.
Chapter 6, Diffusion Flames, more carefully analyzes the differences between momentum and buoyant fuel jets. Ignition by pyrophoric materials, cata- lysts, and hypergolic fuels is now described in Chapter 7. The soot section in Chapter 8, Environmental Combustion Considerations, has been completely changed and also points out that most opposed jet diffusion fl ame experiments must be carefully analyzed since there is a difference between the temperature fi elds in opposed jet diffusion fl ames and simple fuel jets. Lastly, Chapter 9, Combustion of Nonvolatile Fuels, has a completely new approach to carbon combustion.
The use of the new material added to the Appendices should help students as the various new problem sets challenge them. Indeed, this approach has changed the character of the chapters that appeared in earlier editions regard- less of apparent similarity in many cases. It is the hope of the authors that the objectives of this edition have been met.
Irvin Glassman Richard A. Yetter
1
Chemical Thermodynamics and Flame Temperatures
A . INTRODUCTION
The parameters essential for the evaluation of combustion systems are the equilibrium product temperature and composition. If all the heat evolved in the reaction is employed solely to raise the product temperature, this temperature is called the adiabatic fl ame temperature. Because of the importance of the temperature and gas composition in combustion considerations, it is appropri- ate to review those aspects of the fi eld of chemical thermodynamics that deal with these subjects.
B . HEATS OF REACTION AND FORMATION
All chemical reactions are accompanied by either an absorption or evolution of energy, which usually manifests itself as heat. It is possible to determine this amount of heat—and hence the temperature and product composition—from very basic principles. Spectroscopic data and statistical calculations permit one to determine the internal energy of a substance. The internal energy of a given substance is found to be dependent upon its temperature, pressure, and state and is independent of the means by which the state is attained. Likewise, the change in internal energy, ΔE , of a system that results from any physical change or chemical reaction depends only on the initial and fi nal state of the system. Regardless of whether the energy is evolved as heat, energy, or work, the total change in internal energy will be the same.
If a fl ow reaction proceeds with negligible changes in kinetic energy and potential energy and involves no form of work beyond that required for the fl ow, the heat added is equal to the increase of enthalpy of the system
Q ΔH
where Q is the heat added and H is the enthalpy. For a nonfl ow reaction proceeding at constant pressure, the heat added is also equal to the gain in enthalpy
Q ΔH
and if heat evolved,
Q ΔH
Most thermochemical calculations are made for closed thermodynamic systems, and the stoichiometry is most conveniently represented in terms of the molar quantities as determined from statistical calculations. In dealing with compressible fl ow problems in which it is essential to work with open ther- modynamic systems, it is best to employ mass quantities. Throughout this text uppercase symbols will be used for molar quantities and lowercase symbols for mass quantities.
One of the most important thermodynamic facts to know about a given chemical reaction is the change in energy or heat content associated with the reaction at some specifi ed temperature, where each of the reactants and prod- ucts is in an appropriate standard state. This change is known either as the energy or as the heat of reaction at the specifi ed temperature.
The standard state means that for each state a reference state of the aggre- gate exists. For gases, the thermodynamic standard reference state is the ideal gaseous state at atmospheric pressure at each temperature. The ideal gaseous state is the case of isolated molecules, which give no interactions and obey the equation of state of a perfect gas. The standard reference state for pure liquids and solids at a given temperature is the real state of the substance at a pressure of 1 atm. As discussed in Chapter 9, understanding this defi nition of the stand- ard reference state is very important when considering the case of high-tem- perature combustion in which the product composition contains a substantial mole fraction of a condensed phase, such as a metal oxide.
The thermodynamic symbol that represents the property of the substance in the standard state at a given temperature is written, for example, as HT, ET, etc., where the “ degree sign ” superscript ° specifi es the standard state, and the subscript T the specifi c temperature. Statistical calculations actually permit the determination of ET E0 , which is the energy content at a given temperature referred to the energy content at 0 K. For 1 mol in the ideal gaseous state,
PV RT (1.1)
H E (PV) E RT (1.2) which at 0 K reduces to
H0 E0 (1.3)
Thus the heat content at any temperature referred to the heat or energy content at 0 K is known and
(H H0) (E E0) RT(E E0) PV (1.4)
The value (E E0) is determined from spectroscopic information and is actually the energy in the internal (rotational, vibrational, and electronic) and external (translational) degrees of freedom of the molecule. Enthalpy (H H0) has meaning only when there is a group of molecules, a mole for instance; it is thus the Ability of a group of molecules with internal energy to do PV work. In this sense, then, a single molecule can have internal energy, but not enthalpy. As stated, the use of the lowercase symbol will signify values on a mass basis. Since fl ame temperatures are calculated for a closed thermodynamic system and molar conservation is not required, working on a molar basis is most convenient. In fl ame propagation or reacting fl ows through nozzles, conserva- tion of mass is a requirement for a convenient solution; thus when these systems are considered, the per unit mass basis of the thermochemical properties is used.
From the defi nition of the heat of reaction, Qp will depend on the tempera- ture T at which the reaction and product enthalpies are evaluated. The heat of reaction at one temperature T0 can be related to that at another temperature T1 . Consider the reaction confi guration shown in Fig. 1.1 . According to the First Law of Thermodynamics, the heat changes that proceed from reactants at tem- perature T0 to products at temperature T1 , by either path A or path B must be the same. Path A raises the reactants from temperature T0 to T1 , and reacts at T1 . Path B reacts at T0 and raises the products from T0 to T1 . This energy equality, which relates the heats of reaction at the two different temperatures, is written as
nj H H H H
j
T T
,react
∑
⎡⎣⎢( ) ( )
⎤⎦⎥j⎧⎨
⎪⎪⎪
⎩⎪⎪⎪
⎫⎬
⎪⎪⎪
1 0 0 0 ⎭⎪⎪⎪
ΔHH
H n H H H H
T
T i
i
T T
i
1
0 1 0 0 0
Δ
,prod
∑
⎡⎣⎢( ) ( )
⎤⎦⎥⎧⎨
⎪⎪
⎩⎪⎪
⎫⎬
⎪⎪
⎭⎭⎪⎪
(1.5)
where n specifi es the number of moles of the i th product or j th reactant. Any phase changes can be included in the heat content terms. Thus, by knowing the difference in energy content at the different temperatures for the products and
Reactants Products
T1
T0
(2) (1)
(2)
(1)
Path A
Path B ΔHT1
ΔHT0
FIGURE 1.1 Heats of reactions at different base temperatures.
reactants, it is possible to determine the heat of reaction at one temperature from the heat of reaction at another.
If the heats of reaction at a given temperature are known for two separate reactions, the heat of reaction of a third reaction at the same temperature may be determined by simple algebraic addition. This statement is the Law of Heat Summation. For example, reactions (1.6) and (1.7) can be carried out conve- niently in a calorimeter at constant pressure:
Cgraphite O (g) CO g kJ
2 ⎯⎯⎯⎯298K → 2( ), Qp 393 52. (1.6)
CO(g) O g CO g kJ
12 2( )⎯⎯⎯⎯298K → 2( ), Qp 283 0. (1.7) Subtracting these two reactions, one obtains
Cgraphite O (g) CO g kJ
12 2 ⎯⎯⎯⎯298K → ( ), Qp 110 52. (1.8) Since some of the carbon would burn to CO 2 and not solely to CO, it is diffi - cult to determine calorimetrically the heat released by reaction (1.8).
It is, of course, not necessary to have an extensive list of heats of reaction to determine the heat absorbed or evolved in every possible chemical reaction.
A more convenient and logical procedure is to list the standard heats of forma- tion of chemical substances. The standard heat of formation is the enthalpy of a substance in its standard state referred to its elements in their standard states at the same temperature. From this defi nition it is obvious that heats of forma- tion of the elements in their standard states are zero.
The value of the heat of formation of a given substance from its elements may be the result of the determination of the heat of one reaction. Thus, from the calorimetric reaction for burning carbon to CO 2 [Eq. (1.6)], it is possible to write the heat of formation of carbon dioxide at 298 K as
(ΔHf)298,CO2 393 52. kJ/mol
The superscript to the heat of formation symbol ΔHf represents the standard state, and the subscript number represents the base or reference temperature.
From the example for the Law of Heat Summation, it is apparent that the heat of formation of carbon monoxide from Eq. (1.8) is
(ΔHf)298,CO 110 52. kJ/mol
It is evident that, by judicious choice, the number of reactions that must be measured calorimetrically will be about the same as the number of substances whose heats of formation are to be determined.
The logical consequence of the preceding discussion is that, given the heats of formation of the substances comprising any particular reaction, one can directly determine the heat of reaction or heat evolved at the reference tem- perature T0 , most generally T298 , as follows:
ΔHT ni ΔH n ΔH Q
i
T i j
j
T j p
0 0 0
prod f
react
∑
( ) ,∑
( f) , (1.9)Extensive tables of standard heats of formation are available, but they are not all at the same reference temperature. The most convenient are the compila- tions known as the JANAF [1] and NBS Tables [2] , both of which use 298 K as the reference temperature. Table 1.1 lists some values of the heat of forma- tion taken from the JANAF Thermochemical Tables. Actual JANAF tables are reproduced in Appendix A. These tables, which represent only a small selec- tion from the JANAF volume, were chosen as those commonly used in com- bustion and to aid in solving the problem sets throughout this book. Note that, although the developments throughout this book take the reference state as 298 K, the JANAF tables also list ΔHf for all temperatures.
When the products are measured at a temperature T2 different from the ref- erence temperature T0 and the reactants enter the reaction system at a tempera- ture T0 different from the reference temperature, the heat of reaction becomes
ΔH n H H H H ΔH
n
i i
T T T
i
j j
prod
f
re
∑
⎡⎣⎢⎢{ ( ) ( ) }
⎤⎦⎥
⎥
2 0 0 0 ( )0
aact
f 0
evolve
∑
⎡⎣⎢⎢{ ( ) ( ) }
⎤⎦⎥
H H H H H ⎥
Q
T T T
j p
0 0 0 0
( )
(
Δ d
d) (1.10)
The reactants in most systems are considered to enter at the standard refer- ence temperature 298 K. Consequently, the enthalpy terms in the braces for the reactants disappear. The JANAF tables tabulate, as a putative convenience,
(HT H298) instead of (HT H0) . This type of tabulation is unfortunate since the reactants for systems using cryogenic fuels and oxidizers, such as those used in rockets, can enter the system at temperatures lower than the ref- erence temperature. Indeed, the fuel and oxidizer individually could enter at different temperatures. Thus the summation in Eq. (1.10) is handled most con- veniently by realizing that T0 may vary with the substance j .
The values of heats of formation reported in Table 1.1 are ordered so that the largest positive values of the heats of formation per mole are the highest and those with negative heats of formation are the lowest. In fact, this table is similar to a potential energy chart. As species at the top react to form species at the bot- tom, heat is released, and an exothermic system exists. Even a species that has a negative heat of formation can react to form products of still lower negative heats of formation species, thereby releasing heat. Since some fuels that have
Chemical Name State ΔHf ( kJ/mol) Δhf ( kJ/g mol)
C Carbon Vapor 716.67 59.72
N Nitrogen atom Gas 472.68 33.76
O Oxygen atom Gas 249.17 15.57
C 2 H 2 Acetylene Gas 227.06 8.79
H Hydrogen atom Gas 218.00 218.00
O 3 Ozone Gas 142.67 2.97
NO Nitric oxide Gas 90.29 3.01
C 6 H 6 Benzene Gas 82.96 1.06
C 6 H 6 Benzene Liquid 49.06 0.63
C 2 H 4 Ethene Gas 52.38 1.87
N 2 H 4 Hydrazine Liquid 50.63 1.58
OH Hydroxyl
radical
Gas 38.99 2.29
O 2 Oxygen Gas 0 0
N 2 Nitrogen Gas 0 0
H 2 Hydrogen Gas 0 0
C Carbon Solid 0 0
NH 3 Ammonia Gas 45.90 2.70
C 2 H 4 O Ethylene oxide Gas 51.08 0.86
CH 4 Methane Gas 74.87 4.68
C 2 H 6 Ethane Gas 84.81 2.83
CO Carbon
monoxide
Gas 110.53 3.95
C 4 H 10 Butane Gas 124.90 2.15
CH 3 OH Methanol Gas 201.54 6.30
CH 3 OH Methanol Liquid 239.00 7.47
H 2 O Water Gas 241.83 13.44
C 8 H 18 Octane Liquid 250.31 0.46
H 2 O Water Liquid 285.10 15.84
SO 2 Sulfur dioxide Gas 296.84 4.64
C 12 H 16 Dodecane Liquid 347.77 2.17
CO 2 Carbon dioxide Gas 393.52 8.94
SO 3 Sulfur trioxide Gas 395.77 4.95
negative heats of formation form many moles of product species having nega- tive heats of formation, the heat release in such cases can be large. Equation (1.9) shows this result clearly. Indeed, the fi rst summation in Eq. (1.9) is gener- ally much greater than the second. Thus the characteristic of the reacting species or the fuel that signifi cantly determines the heat release is its chemical compo- sition and not necessarily its molar heat of formation. As explained in Section D2, the heats of formation listed on a per unit mass basis simplifi es one’s ability to estimate relative heat release and temperature of one fuel to another without the detailed calculations reported later in this chapter and in Appendix I.
The radicals listed in Table 1.1 that form their respective elements have their heat release equivalent to the radical’s heat of formation. It is then apparent that this heat release is also the bond energy of the element formed. Non-radicals such as acetylene, benzene, and hydrazine can decompose to their elements and/
or other species with negative heats of formation and release heat. Consequently, these fuels can be considered rocket monopropellants. Indeed, the same would hold for hydrogen peroxide; however, what is interesting is that ethylene oxide has a negative heat of formation, but is an actual rocket monopropellant because it essentially decomposes exothermically into carbon monoxide and methane [3] . Chemical reaction kinetics restricts benzene, which has a positive heat of forma- tion from serving as a monopropellant because its energy release is not suffi cient to continuously initiate decomposition in a volumetric reaction space such as a rocket combustion chamber. Insight into the fundamentals for understanding this point is covered in Chapter 2, Section B1. Indeed, for acetylene type and eth- ylene oxide monopropellants the decomposition process must be initiated with oxygen addition and spark ignition to then cause self-sustained decomposition.
Hydrazine and hydrogen peroxide can be ignited and self-sustained with a catalyst in a relatively small volume combustion chamber. Hydrazine is used extensively for control systems, back pack rockets, and as a bipropellant fuel.
It should be noted that in the Gordon and McBride equilibrium thermodynamic program [4] discussed in Appendix I, the actual results obtained might not be realistic because of kinetic reaction conditions that take place in the short stay times in rocket chambers. For example, in the case of hydrazine, ammonia is a product as well as hydrogen and nitrogen [5] . The overall heat release is greater than going strictly to its elements because ammonia is formed in the decom- position process and is frozen in its composition before exiting the chamber.
Ammonia has a relatively large negative heat of formation.
Referring back to Eq. (1.10), when all the heat evolved is used to raise the tem- perature of the product gases, ΔH and Qp become zero. The product temperature T2 in this case is called the adiabatic fl ame temperature and Eq. (1.10) becomes
n H H H H H
n
i i
T T T
i
j j prod
f
react
∑
⎡⎣⎢⎢{ (
2 0) (
0 0) }
0⎤⎦⎥⎥
(Δ )
∑
∑
⎡⎣⎢⎢{ (
HT 0 H0) (
HT 0 H0) }
(ΔHf)T0⎤⎦⎥⎥j (1.11)Again, note that T0 can be different for each reactant. Since the heats of for- mation throughout this text will always be considered as those evaluated at the reference temperature T0 298 K, the expression in braces becomes {(HT H0) (HT H0)} (HT HT)
0 0, which is the value listed in the
JANAF tables (see Appendix A).
If the products ni of this reaction are known, Eq. (1.11) can be solved for the fl ame temperature. For a reacting lean system whose product temperature is less than 1250 K, the products are the normal stable species CO 2 , H 2 O, N 2 , and O 2, whose molar quantities can be determined from simple mass bal- ances. However, most combustion systems reach temperatures appreciably greater than 1250 K, and dissociation of the stable species occurs. Since the dissociation reactions are quite endothermic, a small percentage of dissocia- tion can lower the fl ame temperature substantially. The stable products from a C¶H¶ O reaction system can dissociate by any of the following reactions:
CO CO O
CO H CO H O
H O H O
H O H OH
H O H OH
H
2 1
2 2
2 2 2
2 2 1
2 2
2
2 1
2 2
2 2
H H O2 2O, etc.
Each of these dissociation reactions also specifi es a defi nite equilibrium con- centration of each product at a given temperature; consequently, the reactions are written as equilibrium reactions. In the calculation of the heat of reaction of low-temperature combustion experiments the products could be specifi ed from the chemical stoichiometry; but with dissociation, the specifi cation of the product concentrations becomes much more complex and the ni ’s in the fl ame tempera- ture equation [Eq. (1.11)] are as unknown as the fl ame temperature itself. In order to solve the equation for the ni ’s and T2 , it is apparent that one needs more than mass balance equations. The necessary equations are found in the equilibrium relationships that exist among the product composition in the equilibrium system.
C . FREE ENERGY AND THE EQUILIBRIUM CONSTANTS
The condition for equilibrium is determined from the combined form of the fi rst and second laws of thermodynamics; that is ,
dETdSPdV (1.12)
where S is the entropy. This condition applies to any change affecting a system of constant mass in the absence of gravitational, electrical, and surface forces.
However, the energy content of the system can be changed by introducing more mass. Consider the contribution to the energy of the system on adding one mol- ecule i to be μi . The introduction of a small number dni of the same type contrib- utes a gain in energy of the system of μi dn i . All the possible reversible increases in the energy of the system due to each type of molecule i can be summed to give
dE TdS PdV idni
i
∑
μ (1.13)It is apparent from the defi nition of enthalpy H and the introduction of the con- cept of the Gibbs free energy G
GHTS (1.14)
that
dH TdS VdP idni
i
∑
μ (1.15)and
dG SdT VdP idni
i
∑
μ (1.16)Recall that P and T are intensive properties that are independent of the size of mass of the system, whereas E, H, G, and S (as well as V and n ) are extensive properties that increase in proportion to mass or size. By writing the general relation for the total derivative of G with respect to the variables in Eq. (1.16), one obtains
dG G
T dT G
P dP G
P ni T ni ni
∂
∂
⎛
⎝⎜⎜⎜ ⎞
⎠⎟⎟⎟⎟ ∂
∂
⎛
⎝⎜⎜⎜ ⎞
⎠⎟⎟⎟⎟ ∂
∂
⎛
⎝⎜⎜⎜
, , ⎜
⎞⎞
⎠⎟⎟⎟
⎟⎟ ( )
∑
P T n i i
j j i
dn
, ,
(1.17)
Thus,
μi
i T P n
G n
j
∂∂
⎛
⎝⎜⎜⎜
⎜
⎞
⎠⎟⎟⎟
⎟⎟, ,
(1.18)
or, more generally, from dealing with the equations for E and H μi
i T P n i S V n i
G n
E n
H n
j j
∂
∂
⎛
⎝⎜⎜⎜
⎜
⎞
⎠⎟⎟⎟
⎟⎟ ∂
∂
⎛
⎝⎜⎜⎜
⎜
⎞
⎠⎟⎟⎟
⎟⎟ ∂
, , , , ∂
⎛⎛
⎝⎜⎜⎜
⎜
⎞
⎠⎟⎟⎟
⎟⎟S P n
, , j
(1.19)
where μi is called the chemical potential or the partial molar free energy. The condition of equilibrium is that the entropy of the system have a maximum value for all possible confi gurations that are consistent with constant energy and volume. If the entropy of any system at constant volume and energy is at its maximum value, the system is at equilibrium; therefore, in any change from its equilibrium state dS is zero. It follows then from Eq. (1.13) that the condi- tion for equilibrium is
μidni 0
∑
(1.20)The concept of the chemical potential is introduced here because this property plays an important role in reacting systems. In this context, one may consider that a reaction moves in the direction of decreasing chemical potential, reaching equi- librium only when the potential of the reactants equals that of the products [3] .
Thus, from Eq. (1.16) the criterion for equilibrium for combustion products of a chemical system at constant T and P is
(dG)T P, 0 (1.21)
and it becomes possible to determine the relationship between the Gibbs free energy and the equilibrium partial pressures of a combustion product mixture.
One deals with perfect gases so that there are no forces of interactions between the molecules except at the instant of reaction; thus, each gas acts as if it were in a container alone. Let G , the total free energy of a product mix- ture, be represented by
G
∑
n Gi i, iA, B,…,R, S... (1.22) for an equilibrium reaction among arbitrary products:aAbB…rR sS … (1.23) Note that A, B, … , R, S, … represent substances in the products only and a, b, … , r, s, … are the stoichiometric coeffi cients that govern the proportions by which different substances appear in the arbitrary equilibrium system cho- sen. The ni ’s represent the instantaneous number of each compound. Under the ideal gas assumption the free energies are additive, as shown above. This assumption permits one to neglect the free energy of mixing. Thus, as stated earlier,
G P T( , )H T( )TS P T( , ) (1.24) Since the standard state pressure for a gas is P0 1 atm, one may write
G P T( 0, ) H T( )TS P T( 0, ) (1.25)
Subtracting the last two equations, one obtains
G G (H H ) T S( S ) (1.26) Since H is not a function of pressure, H H° must be zero, and then
G G T S( S ) (1.27) Equation (1.27) relates the difference in free energy for a gas at any pressure and temperature to the standard state condition at constant temperature. Here dH 0, and from Eq. (1.15) the relationship of the entropy to the pressure is found to be
S ln( / )S R p p0 (1.28)
Hence, one fi nds that
G T P( , ) G RTln( /p p0) (1.29) An expression can now be written for the total free energy of a gas mixture.
In this case P is the partial pressure Pi of a particular gaseous component and obviously has the following relationship to the total pressure P :
p n
n P
i i
i i
∑
⎛
⎝⎜⎜⎜
⎜⎜⎜
⎞
⎠
⎟⎟⎟⎟⎟⎟ (1.30)
where ( /ni
∑
ini) is the mole fraction of gaseous species i in the mixture.Equation (1.29) thus becomes
G T P ni G RT p p
i
i i
( , )
∑ {
ln( / 0)}
(1.31) As determined earlier [Eq. (1.21)], the criterion for equilibrium is ( dG ) T,P 0.Taking the derivative of G in Eq. (1.31), one obtains
G dni RT dn p p RT n dp p
i
i i i i i i
i i
∑ ∑
( ) ln( / 0)∑
( / ) 0 (1.32)Evaluating the last term of the left-hand side of Eq. (1.32), one has
n dp p
n
p dp n
p dp
i i
i i
i i i
i i
i i
∑ ∑
i ⎛∑ ∑ ∑
⎝⎜⎜⎜
⎜⎜
⎞
⎠⎟⎟⎟
⎟⎟ 0 (1.33)
since the total pressure is constant, and thus
∑
idpi 0. Now consider the fi rst term in Eq. (1.32):
G dni dn G dn G dn G dn G
i
i
∑
( A) A ( B) B ( R) R ( S) s (1.34)By the defi nition of the stoichiometric coeffi cients,
dni ~ ai, dni kai (1.35)
where k is a proportionality constant. Hence
G dni i k aG bG rG sG
i
{ A B R S }
∑
(1.36)Similarly, the proportionality constant k will appear as a multiplier in the sec- ond term of Eq. (1.32). Since Eq. (1.32) must equal zero, the third term already has been shown equal to zero, and k cannot be zero, one obtains
( ) ( ) ( )
( )
aG bG rG sG RT p p p p
p p
r s
A B R s a
R S
A
/ /
ln / 0 0
0 ((p pB/ 0)b
⎧⎨
⎪⎪
⎩⎪⎪
⎫⎬
⎪⎪
⎭⎪⎪(1.37) One then defi nes
ΔG aG A bGB rGR sGS (1.38) where ΔG° is called the standard state free energy change and p0 1 atm.
This name is reasonable sinceΔG° is the change of free energy for reaction (1.23) if it takes place at standard conditions and goes to completion to the right. Since the standard state pressure p0 is 1 atm, the condition for equilib- rium becomes
ΔG RTln(p p p prR sS/ aA bB) (1.39) where the partial pressures are measured in atmospheres. One then defi nes the equilibrium constant at constant pressure from Eq. (1.39) as
Kp p p p prR sS/ aA bB
Then
ΔG RTlnKp, Kp exp(ΔG RT/ ) (1.40)
where Kp is not a function of the total pressure, but rather a function of temper- ature alone. It is a little surprising that the free energy change at the standard state pressure (1 atm) determines the equilibrium condition at all other pres- sures. Equations (1.39) and (1.40) can be modifi ed to account for nonideality in the product state; however, because of the high temperatures reached in com- bustion systems, ideality can be assumed even under rocket chamber pressures.
The energy and mass conservation equations used in the determination of the fl ame temperature are more conveniently written in terms of moles; thus, it is best to write the partial pressure in Kp in terms of moles and the total pressure P . This conversion is accomplished through the relationship between partial pressure p and total pressure P , as given by Eq. (1.30). Substituting this expression for pi
[Eq. (1.30)] in the defi nition of the equilibrium constant [Eq. (1.40)], one obtains Kp (n n n nrR Ss/ aA Bb)(P/
∑
ni)r s a b (1.41) which is sometimes written asKp KN(P/
∑
ni)r s a b (1.42)where
KN n n n nrR Ss/ aA Bb (1.43) When
r 0 s a b (1.44)
the equilibrium reaction is said to be pressure-insensitive. Again, however, it is worth repeating that Kp is not a function of pressure; however, Eq. (1.42) shows that KN can be a function of pressure.
The equilibrium constant based on concentration (in moles per cubic cen- timeter) is sometimes used, particularly in chemical kinetic analyses (to be dis- cussed in the next chapter). This constant is found by recalling the perfect gas law, which states that
PV
∑
n RTi (1.45)or
(P/
∑
ni)(RT V/ ) (1.46) where V is the volume. Substituting for (P/∑
ni) in Eq. (1.42) givesK n n n n RT
p r s a b V
r s a b
( R S) /( A B)
⎡⎣⎢ ⎤
⎦⎥⎛
⎝⎜⎜⎜ ⎞
⎠⎟⎟⎟⎟ (1.47)
