shock tube measurements of oxygenated fuel … · ocho), methyl acetate (ch. 3. oc(o)ch. 3), methyl...

SHOCK TUBE MEASUREMENTS OF OXYGENATED FUEL

COMBUSTION USING LASER ABSORPTION SPECTROSCOPY

A DISSERTATION

SUBMITTED TO THE DEPARTMENT OF MECHANICAL ENGINEERING

AND THE COMMITTEE ON GRADUATE STUDIES

OF STANFORD UNIVERSITY

IN PARTIAL FULFILLMENT OF THE REQUIREMENTS

FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY

King Yiu Lam June 2013

http://creativecommons.org/licenses/by-nc/3.0/us/

This dissertation is online at: http://purl.stanford.edu/ty513tt0976

© 2013 by King Yiu Lam. All Rights Reserved.

Re-distributed by Stanford University under license with the author.

This work is licensed under a Creative Commons Attribution-Noncommercial 3.0 United States License.

ii



http://purl.stanford.edu/ty513tt0976

I certify that I have read this dissertation and that, in my opinion, it is fully adequatein scope and quality as a dissertation for the degree of Doctor of Philosophy.

Ronald Hanson, Primary Adviser


Craig Bowman


David Davidson

Approved for the Stanford University Committee on Graduate Studies.

Patricia J. Gumport, Vice Provost Graduate Education

This signature page was generated electronically upon submission of this dissertation in electronic format. An original signed hard copy of the signature page is on file inUniversity Archives.

iii

v

Abstract In the current engine development, fuel reformulation is considered as one of the

potential strategies to improve fuel efficiency, reduce petroleum consumption, and

minimize pollutant formation. Oxygenated fuels can be used as neat fuels or additives in

spark-ignition and diesel engines to allow for more complete combustion. To understand

the influence of oxygenated fuels on engine performance, accurate comprehensive kinetic

mechanisms, which can consist of hundreds to thousands of elementary reactions, are

needed to describe the chemistry of the combustion events, such as autoignition and

pollutant formation.

The primary objective of the research presented in this dissertation is to provide

reliable experimental kinetic targets, such as ignition delay times, species time histories,

and direct reaction rate constant measurements, using shock tube and laser absorption

techniques in order to evaluate and refine the existing kinetic mechanisms for two

different types of oxygenated fuels (i.e., ketones and methyl esters) and to reexamine the

kinetics of the H2 + OH reaction. The topics of this work are mainly divided into three

sections: (1) H2 + OH kinetics, (2) ketone combustion chemistry, and (3) methyl ester +

OH kinetics.

The reaction of OH with molecular hydrogen (H2)

H2 + OH → H2O + H (1)

is an important chain-propagating reaction in all combustion systems, particularly in

hydrogen combustion, and its direct rate constant measurements are discussed in the first

part of this dissertation. The rate constant for reaction (1) was measured behind reflected

shock waves over the temperature range of 902-1518 K at pressures of 1.15-1.52 atm.

OH radicals were produced by rapid thermal decomposition of tert-butyl hydroperoxide

vi

(TBHP) at high temperatures, and were monitored using the narrow-linewidth ring dye

laser absorption of the well-characterized R1(5) line in the OH A–X (0, 0) band near

306.69 nm. Consequently, this work aims to report the rate constant for reaction (1) with

a much lower experimental scatter and overall uncertainty (as compared to the data

available in the literature).

Ketones are important to a variety of modern combustion processes. They are

widely used as fuel tracers in planar laser-induced fluorescence (PLIF) imaging of

combustion processes due to their physical similarity to gasoline surrogate components.

Additionally, they are often formed as intermediate products during oxidation of large

oxygenated fuels, such as alcohols and methyl esters. In the second part of this

dissertation, the combustion characteristics of acetone (CH3COCH3), 2-butanone

(C2H5COCH3), and 3-pentanone (C2H5COC2H5) are discussed in the context of the

reflected shock wave experiments. These experiments were performed using different

laser absorption methods to monitor species concentration time histories (i.e., ketones,

CH3, CO, C2H4, CH4, OH, and H2O) over the temperature range of 1100-1650 K at

pressures near 1.6 atm. These speciation data were then compared with the simulations

from the detailed mechanisms of Pichon et al. (2009) and Serinyel et al. (2010).

Consequently, the overall rate constants for the thermal decomposition reactions of

acetone, 2-butanone, and 3-pentanone

CH3COCH3 (+ M) → CH3 + CH3CO (+ M) (2)

C2H5COCH3 (+ M) → Products (+ M) (3)

C2H5COC2H5 (+ M) → Products (+ M) (4)

were inferred by matching the species profiles with the simulations from the detailed

mechanisms at pressures near 1.6 atm. In addition, an O-atom balance analysis from the

speciation data revealed the absence of a methyl ketene removal pathway in the original

models. Furthermore, the overall rate constants for the reactions of OH with a series of

ketones

CH3COCH3 + OH → CH3COCH2 + H2O (5)

C2H5COCH3 + OH → Products (6)

C2H5COC2H5 + OH → Products (7)


vii

were determined using UV laser absorption of OH over the temperature range of 870-

1360 K at pressures of 1-2 atm. These measurements included the first direct high-

temperature measurements of the overall rate constants for reactions (6)-(8), and were

compared with the theoretical calculations from Zhou et al. (2011) and the estimates

using the structure-activity relationship (SAR) (1995).

Biodiesel, which consists of fatty acid methyl esters (FAMES), is a promising

alternative to fossil fuels. The four simplest methyl esters include methyl formate

(CH3OCHO), methyl acetate (CH3OC(O)CH3), methyl propanoate (CH3OC(O)C2H5),

and methyl butanoate (CH3OC(O)C3H7), and their combustion chemistry is a building

block for the chemistry of large methyl esters. In the third part of this dissertation, the

rate constant measurements for the reactions of OH with four small methyl esters are

discussed:

CH3OCHO + OH → Products (9)

CH3OC(O)CH3 + OH → Products (10)

CH3OC(O)C2H5 + OH → Products (11)


These reactions were studied behind reflected shock waves using UV laser absorption of

OH over 876-1371 K at pressures near 1.5 atm. This study presented the first direct high-

temperature rate constant measurements of reactions (9)-(12). These measurements were

also compared with the estimated values from different detailed mechanisms and from

the structure-activity relationship (SAR).

ix

Acknowledgements First, I would like to thank my advisor, Prof. Ronald Hanson, for the opportunities

and support he has offered throughout my graduate studies at Stanford. His critical

thinking, carefulness, and wisdom have shaped me into a better and more careful

researcher. I would like to thank Dr. David Davidson for offering numerous advice and

guidance throughout my Ph.D. career. His willingness to help students at the lab has

made my time at Stanford much smoother. I am also thankful to Prof. Bowman for

serving on my qualification exam committee and my reading committee.

I have been very fortunate to work with many outstanding students in the Hanson

group. In particular, I am immensely grateful to Zekai Hong, Genny Pang, and Robert

Cook for teaching me how to operate shock tubes and different laser equipment, allowing

me to accomplish the work presented in this dissertation. I am very grateful to Wei Ren

for his friendship and the collaborative efforts we have made together in research and

course works. I am also grateful to many students in the Hanson group who have made

my life at Stanford more meaningful and joyful. Finally, I am sincerely grateful to my

parents for their support and encouragement in times of trouble and frustration.

This work was supported by the U.S. Department of Energy, Basic Energy

Sciences (DE-FG02-88ER13857) with Dr. Wade Sisk as program manager.

xi

Table of Contents Abstract .................................................................................................................... v

Acknowledgements .......................................................................................................... ix

Table of Contents ............................................................................................................. xi

List of Tables .................................................................................................................. xv

List of Figures ................................................................................................................ xvii

Chapter 1 Background and Motivation ......................................................................... 1

1.1 Introduction ............................................................................................................1

1.2 Background and Motivation ..................................................................................2

1.2.1 H2 + OH Kinetics ........................................................................................ 2

1.2.2 Ketone Combustion Chemistry ................................................................... 5

1.2.3 Methyl Ester + OH Kinetics ....................................................................... 6

1.3 Scope and Organization of Thesis .........................................................................8

Chapter 2 Experimental Methods ................................................................................ 11

2.1 Shock Tube Facility .............................................................................................11

2.2 Laser Absorption Methods ..................................................................................12

2.2.1 UV Laser Absorption of OH ..................................................................... 12

2.2.2 UV Laser Absorption of Ketones ............................................................. 13

2.2.3 UV Laser Absorption of CH3 .................................................................... 14

2.2.4 IR Laser Absorption of CO ....................................................................... 16

2.2.5 CO2 Laser Absorption of C2H4 ................................................................. 17

2.2.6 IR Laser Absorption of H2O ..................................................................... 18

xii

2.2.7 IR Laser Absorption of CH4 ..................................................................... 19

2.3 Summary ..............................................................................................................20

Chapter 3 A Shock Tube Study of H2 + OH → H2O + H using OH Laser

Absorption .................................................................................................... 21

3.1 Introduction ..........................................................................................................21

3.2 Experimental Details ...........................................................................................21

3.3 Kinetic Measurements .........................................................................................22

3.3.1 Kinetic Mechanism Description ............................................................... 22

3.3.2 H2 + OH Kinetics ...................................................................................... 24

3.4 Comparison with Earlier Work ............................................................................29

3.5 Summary ..............................................................................................................32

Chapter 4 Multi-Species Time History Measurements during High-Temperature

Acetone and 2-Butanone Pyrolysis ............................................................. 33

4.1 Introduction ..........................................................................................................33


4.2.1 Mixture Preparation .................................................................................. 34

4.2.2 Species Absorption Coefficient Evaluations ............................................ 35

4.3 Results and Discussion ........................................................................................35

4.3.1 Acetone Pyrolysis ..................................................................................... 35

4.3.2 2-Butanone Pyrolysis ................................................................................ 45

4.4 Summary ..............................................................................................................53

Chapter 5 Shock Tube Measurements of 3-Pentanone Pyrolysis and Oxidation .... 55

5.1 Introduction ..........................................................................................................55


5.2.1 Mixture Preparation .................................................................................. 56

xiii

5.2.2 Species Absorption Coefficient Evaluations ............................................ 57

5.3 Results and Discussion ........................................................................................60

5.3.1 3-Pentanone Pyrolysis ............................................................................... 60

5.3.2 3-Pentanone Oxidation.............................................................................. 72

5.3.3 Comparisons of Ketone Oxidation Characteristics ................................... 84

5.4 Summary ..............................................................................................................87

5.5 Possible Future Work ..........................................................................................87

Chapter 6 High-Temperature Measurements of the Reactions of OH with a Series

of Ketones: Acetone, 2-Butanone, 3-Pentanone, and 2-Pentanone ......... 89

6.1 Introduction ..........................................................................................................89


6.3 Kinetic Measurements .........................................................................................91

6.3.1 Choice of Kinetic Mechanisms ................................................................. 91

6.3.2 Acetone + OH Kinetics ............................................................................. 93

6.3.3 2-Butanone + OH Kinetics........................................................................ 98

6.3.4 3-Pentanone + OH Kinetics .................................................................... 101

6.3.5 2-Pentanone + OH Kinetics .................................................................... 105

6.3.6 Comparison of Ketone + OH Kinetics .................................................... 109

6.4 Comparison with Low Temperature Data .........................................................110

6.5 Comparison with Structure-Activity Relationship ............................................114

6.6 Summary ............................................................................................................116

Chapter 7 High-Temperature Measurements of the Reactions of OH with Small

Methyl Esters: Methyl Formate, Methyl Acetate, Methyl Propanoate,

and Methyl Butanoate ............................................................................... 117

7.1 Introduction ........................................................................................................117

7.2 Experimental Details .........................................................................................118

xiv

7.3 Kinetic Measurements .......................................................................................118

7.3.1 Choice of Kinetic Mechanisms ............................................................... 119

7.3.2 Methyl Formate (MF) + OH Kinetics ..................................................... 120

7.3.3 Methyl Acetate (MA) + OH Kinetics ..................................................... 125

7.3.4 Methyl Propanoate (MP) + OH Kinetics ................................................ 130

7.3.5 Methyl Butanoate (MB) + OH Kinetics ................................................. 135

7.4 Comparison with Low Temperature Data .........................................................141

7.5 Comparison with Structure-Activity Relationship ............................................143

7.6 Summary ............................................................................................................145

Chapter 8 Conclusions and Future Work ................................................................. 147

8.1 Summary of Results ...........................................................................................147

8.1.1 H2 + OH Kinetics .................................................................................... 147

8.1.2 Ketone Combustion Chemistry ............................................................... 148

8.1.3 Methyl Ester + OH Kinetics ................................................................... 150

8.2 Publications ........................................................................................................151

8.3 Recommendations for Future Work ..................................................................152

8.3.1 Ethyl Radical Diagnostics and Decomposition Pathway ........................ 152

8.3.2 Methyl Ester Kinetics ............................................................................. 153

Appendix A Shock Tube Ignition Delay Time Measurements in Propane/O2/Argon

Mixtures at Near-Constant-Volume Conditions ..................................... 155

Appendix B Ignition Delay Time Measurements of Normal Alkanes and

Cycloalkanes ............................................................................................... 173

Appendix C Multi-Species Time History Measurements during the Oxidation of n-

Decane, iso-Octane, and Toluene ............................................................. 179

Bibliography ................................................................................................................ 191

xv

List of Tables Table 3.1: Reactions Describing H2 + OH Experiments at P = 1.3 atm. ........................................ 23

Table 3.2: Rate Constant Data for H2 + OH → H2O + H. ............................................................. 27

Table 4.1: Summary of acetone unimolecular dissociation rate constant data. ............................. 41

Table 4.2: Summary of overall 2-butanone decomposition rate constant data. ............................. 48

Table 5.1: Summary of test gas mixture compositions and measured species. ............................. 56

Table 5.2: Kinetic parameters employed in the Serinyel et al. mechanism. .................................. 64

Table 5.3: Summary of overall 3-pentanone decomposition rate constant data. ........................... 65

Table 6.1: CH3COCH3 + OH → Products: Rate Constant Data. ................................................... 95

Table 6.2: C2H5COCH3 + OH → Products: Rate Constant Data. ................................................ 100

Table 6.3: C2H5COC2H5 + OH → Products: Rate Constant Data. ............................................... 104

Table 6.4: C3H7COCH3 + OH → Products: Rate Constant Data. ................................................ 108

Table 7.1: CH3OCHO + OH → Products: Rate Constant Data. .................................................. 122

Table 7.2: CH3OC(O)CH3 + OH → Products: Rate Constant Data. ............................................ 128

Table 7.3: CH3OC(O)C2H5 + OH → Products: Rate Constant Data. .......................................... 133

Table 7.4: CH3OC(O)C3H7 + OH → Products: Rate Constant Data. .......................................... 138

Table 7.5: Comparison of the rate constants for channels (12a)-(12d) from Fisher et al.

[48], Dooley et al. [54], and Hakka et al. [55] at 1133 K and 1300 K. ................... 140

xvii

List of Figures Figure 1.1: (a) Arrhenius plot for the reaction of OH with H2 at all temperatures. (b)

Arrhenius plot for the reaction of OH with H2 at high temperatures (1000-

2500 K). ...................................................................................................................... 4

Figure 1.2: Impact of the rate constant for the reaction of H2 + OH → H2O + H and its

uncertainty (within a factor of 2) on the laminar flame speed predictions of

H2-air mixtures at 298 K and 1 atm. Laminar flame speed simulations were

performed using the preliminary CEFRC foundational fuel model [33] with

the H2 + OH reaction rate constant from Michael and Sutherland [18]. ..................... 5

Figure 2.1: Schematic of OH detection using a narrow-linewidth ring dye laser near

306.69 nm. Figure adapted from refs. [23-24]. ........................................................ 13

Figure 2.2: High-temperature ketone absorption cross-sections at 306.65 nm. ............................. 14

Figure 2.3: Schematic of CH3 detection near 216.6 nm using the frequency-quadrupled

output of near-infrared radiation from a pulsed Ti:Sapphire laser. Figure

adapted from ref. [63]. .............................................................................................. 15

Figure 2.4: Schematic of CO detection using cw quantum cascade laser near 4.56 µm.

Figure adapted from ref. [66]. ................................................................................... 16

Figure 2.5: Schematic of C2H4 detection using cw CO2 laser near 10.5 µm. Figure

adapted from ref. [67]. .............................................................................................. 18

Figure 2.6: Schematic of H2O detection using cw DFB laser near 2.55 µm. Figure

adapted from ref. [69]. .............................................................................................. 18

Figure 2.7: Schematic of CH4 detection using a scanned-wavelength mid-IR laser near

3.4 µm. Figure adapted from ref. [70]. ..................................................................... 19

Figure 3.1: OH sensitivity plot for the rate constant measurement of H2 + OH at 1228 K

and 1.29 atm. ............................................................................................................. 25

xviii

Figure 3.2: Sample H2 + OH rate constant measurement using the mixture of 1001 ppm

H2 with ~26 ppm TBHP (and 99 ppm water) in Ar at 1228 K and 1.29 atm.

Simulation from USC-Mech v2.0 [72] for the best-fit rate constant, along

with variations of ±20%, is also shown. ................................................................... 26

Figure 3.3: H2 + OH rate constant measurements at various temperatures, along with the

simulations from USC-Mech v2.0 for the best-fit rate constants. ............................. 26

Figure 3.4: Uncertainty analysis for the rate constant of H2 + OH → H2O + H at 1228 K

and 1.29 atm. ............................................................................................................. 28

Figure 3.5: Arrhenius plot for H2 + OH (k1) at temperatures above 833 K. ................................... 29

Figure 3.6: Comparison with previous studies at temperatures above 833 K. ............................... 32

Figure 4.1: CO sensitivity for 0.25% acetone in Ar using the Pichon et al. mechanism

[89]. ........................................................................................................................... 36

Figure 4.2: CO rate of production (ROP) plot for 0.25% acetone in Ar using the Pichon et

al. mechanism [89]. ................................................................................................... 37

Figure 4.3: Sample CO time histories: measured and calculated values. ...................................... 37

Figure 4.4: Summary of acetone dissociation rate constant (k2). ................................................... 38

Figure 4.5: Acetone sensitivity for 1% acetone in Ar using the Pichon et al. mechanism

[89]. ........................................................................................................................... 39

Figure 4.6: Acetone time histories for 1% acetone in Ar: measured and simulated values. .......... 40

Figure 4.7: CH3 time histories for 0.25% acetone in Ar: measured and calculated values. ........... 42

Figure 4.8: C2H4 time histories for 1% acetone in Ar: measured and calculated values. .............. 43

Figure 4.9: CH4 time histories for 1.5% acetone in Ar: measured and calculated values. ............. 44

Figure 4.10: Arrhenius plot of the rate constants for the reaction of CH3COCH3 + CH3 →

CH3COCH2 + CH4 from Saxena et al. [88], Sato and Hidaka [87], and

Pichon et al. [89]. ...................................................................................................... 45

Figure 4.11: 2-Butanone time histories for 1% 2-butanone in Ar: measured and simulated

values. ....................................................................................................................... 46

Figure 4.12: 2-Butanone sensitivity for 1% 2-butanone in Ar. ...................................................... 46

Figure 4.13: Arrhenius plot for overall 2-butanone decomposition rate constant (k3). .................. 47

Figure 4.14: CH3 time histories for 0.25% 2-butanone in Ar: measured and calculated

values. ....................................................................................................................... 48

xix

Figure 4.15: CO rate of production (ROP) plot for 1% 2-butanone in Ar using the original

Serinyel et al. mechanism (with the revised k3). ....................................................... 49

Figure 4.16: CO time histories for 1% 2-butanone in Ar: measured and calculated values. ......... 50

Figure 4.17: C2H4 time histories for 1% 2-butanone in Ar: measured and calculated

values. ....................................................................................................................... 51

Figure 4.18: CH4 time histories for 1.5% 2-butanone in Ar: measured and calculated

values. ....................................................................................................................... 52

Figure 5.1: Comparison of CO mole fraction time histories at 1325 K and 1.60 atm with

different absorption coefficients in Beer’s law. ........................................................ 58

Figure 5.2: Comparison of OH mole fraction time histories at 1486 K and 1.52 atm with

different absorption coefficients in Beer’s law. The OH mole fractions by

constant U, V and constant H, P are virtually indistinguishable for OH. ................. 59

Figure 5.3: Measured and simulated 3-pentanone time histories for 1% 3-pentanone in

Ar. Simulations used the Serinyel et al. mechanism. ................................................ 61

Figure 5.4: 3-pentanone sensitivity analysis for 1% 3-pentanone in Ar at 1323 K and 1.6

atm. ............................................................................................................................ 61

Figure 5.5: CH3 time histories for 0.1% 3-pentanone in Ar. Simulations were done using

the Serinyel et al. mechanism. .................................................................................. 62

Figure 5.6: CH3 sensitivity analysis for 0.1% 3-pentanone in Ar at 1433 K and 1.6 atm. ............. 63

Figure 5.7: (a) Best-fit 3-pentanone time histories and (b) best-fit CH3 time histories

using the Serinyel et al. mechanism with revised overall 3-pentanone

decomposition rate constant (k4). .............................................................................. 66

Figure 5.8: Arrhenius plot for the overall 3-pentanone decomposition rate constant (k4) at

1.6 atm. ...................................................................................................................... 67

Figure 5.9: Measured 3-pentanone and CO time histories during 3-pentanone pyrolysis at

1248 K and 1.6 atm. .................................................................................................. 68

Figure 5.10: CO sensitivity analysis for 1% 3-pentanone in Ar at 1248 K and 1.6 atm. ............... 68

Figure 5.11: CO time histories for 0.25% 3-pentanone in Ar: measured and calculated

values from the (a) original and (b) modified Serinyel et al. mechanisms. .............. 70

Figure 5.12: C2H4 time histories for 0.25% 3-pentanone in Ar: measured and calculated

values. ....................................................................................................................... 71

xx

Figure 5.13: Sample sidewall pressure and endwall OH* emission time histories recorded

during an experiment of 3-pentanone ignition at 1113 K and 1.1 atm (3-

pentanone/ 4.0% O2/ Ar, Φ = 0.5). A tailored gas mixture of 60% helium/

40% nitrogen was used as driver gas to achieve a long test time. For high

fuel concentration mixtures, the definition of the endwall ignition delay time

is shown in the figure. ............................................................................................... 72

Figure 5.14: Measured and simulated 3-pentanone ignition delay times at (a) Φ = 1.0 and

(b) Φ = 0.5 and P5 = 1.0 atm. .................................................................................... 74

Figure 5.15: Comparison of model predictions between (a) the Serinyel et al. mechanism

of NUI Galway [98] and (b) the modified mechanism on ignition delay time

measurements from Serinyel et al. ............................................................................ 75

Figure 5.16: OH sensitivity analysis for 400 ppm 3-pentanone with 0.28% O2 in Ar (Φ =

1.0) at 1486 K and 1.52 atm. ..................................................................................... 76

Figure 5.17: H2O sensitivity analysis for 400 ppm 3-pentanone with 0.28% O2 in Ar (Φ =

1.0) at 1486 K and 1.52 atm. ..................................................................................... 77

Figure 5.18: Comparisons of measured and simulated H2O time histories from the (a)

original and (b) modified Serinyel et al. mechanisms for 400 ppm 3-

pentanone with 0.28% O2 in Ar (Φ = 1.0). ............................................................... 79

Figure 5.19: Comparisons of measured and simulated H2O time histories from the (a)


pentanone with 0.56% O2 in Ar (Φ = 0.5). ............................................................... 80

Figure 5.20: Comparisons of measured and simulated OH time histories from the (a)


pentanone with 0.28% O2 in Ar (Φ = 1.0). Inset figures are provided to

show the early-time features over 0-400 µs. ............................................................. 83

Figure 5.21: Comparisons of measured and simulated OH time histories from the (a)


pentanone with 0.56% O2 in Ar (Φ = 0.5). Inset figures are provided to

show the early-time features over 0-400 µs. ............................................................. 84

Figure 5.22: Comparison of ignition delay times for different ketones (acetone, 2-

pentanone and 3-pentanone). .................................................................................... 85

xxi

Figure 5.23: Comparison of OH time histories for the mixtures of ketone (i.e., acetone, 2-

pentanone and 3-pentanone) with 0.525% O2 in Ar at a pressure of 2.6 atm

and an equivalence ratio of 1.0. An inset figure is provided to show the

early-time features over 0-400 µs. ............................................................................ 86

Figure 6.1: OH sensitivity plot for the rate constant measurement of acetone + OH at

1148 K and 1.95 atm. ................................................................................................ 94

Figure 6.2: Sample acetone + OH rate constant measurement using the mixture of 304

ppm acetone with ~28 ppm TBHP (and 73 ppm water) in Ar at 1148 K and

1.95 atm. Simulation from the modified Pichon et al. mechanism for the

best-fit rate constant, along with perturbations of ±50%, is also shown. .................. 95

Figure 6.3: Uncertainty analysis for the rate constant of CH3COCH3 + OH → products at

1148 K and 1.95 atm. ................................................................................................ 96

Figure 6.4: Arrhenius plot for acetone + OH (k5) at temperatures above 833 K. .......................... 97

Figure 6.5: OH sensitivity plot for the rate constant measurement of 2-butanone + OH at

1039 K and 1.41 atm. ................................................................................................ 98

Figure 6.6: Sample 2-butanone + OH rate constant measurement using the mixture of 152

ppm 2-butanone with ~14 ppm TBHP (and 41 ppm water) in Ar at 1039 K

and 1.41 atm. Simulation from the modified Serinyel et al. mechanism for

the best-fit rate constant, along with perturbations of ±50%, is also shown. ............ 99

Figure 6.7: Arrhenius plot for 2-butanone + OH (k6) at temperatures above 833 K. ................... 101

Figure 6.8: OH sensitivity plot for the rate constant measurement of 3-pentanone + OH at

1188 K and 1.94 atm. .............................................................................................. 102

Figure 6.9: Sample 3-pentanone + OH rate constant measurement using the mixture of

213 ppm 3-pentanone with ~17 ppm TBHP (and 59 ppm water) in Ar at

1188 K and 1.94 atm. Simulation from the modified Serinyel et al.

mechanism for the best-fit rate constant, along with perturbations of ±50%,

is also shown. .......................................................................................................... 103

Figure 6.10: Arrhenius plot for 3-pentanone + OH (k7) at temperatures above 833 K. ............... 105

Figure 6.11: Sample 2-pentanone + OH rate constant measurement using the mixture of

161 ppm 2-pentanone with ~15 ppm TBHP (and 45 ppm water) in Ar at

1186 K and 1.30 atm. Simulation from the modified Serinyel et al.

xxii

mechanism for the best-fit rate constant, along with perturbations of ±50%,

is also shown. .......................................................................................................... 107

Figure 6.12: Arrhenius plot for 2-pentanone + OH (k8) at temperatures above 900 K. ............... 108

Figure 6.13: Arrhenius plot of the measured rate constants for reactions (5)-(8) at

temperatures above 870 K. ...................................................................................... 109

Figure 6.14: Arrhenius plot for acetone + OH → products (k5) at all temperatures. ................... 111

Figure 6.15: Arrhenius plot for 2-butanone + OH → products (k6) at all temperatures. ............. 112

Figure 6.16: Arrhenius plot for 3-pentanone + OH → products (k7) at all temperatures. ............ 113

Figure 6.17: Arrhenius plot for 2-pentanone + OH → products (k8) at all temperatures. ............ 114

Figure 7.1: OH sensitivity plot for the rate constant measurement of methyl formate +

OH at 1168 K and 1.40 atm. ................................................................................... 121

Figure 7.2: Sample methyl formate + OH rate constant measurement using the mixture of

322 ppm methyl formate with ~26 ppm TBHP (and 70 ppm water) in Ar at

1168 K and 1.40 atm. Simulation from the Dooley et al. mechanism [49] for

the best-fit rate constant, along with perturbations of ±50%, is also shown. .......... 122

Figure 7.3: Uncertainty analysis for the rate constant of methyl formate + OH → products

at 1168 K and 1.40 atm. .......................................................................................... 123

Figure 7.4: Arrhenius plot for methyl formate + OH (k9) at temperatures above 833 K. ............ 125

Figure 7.5: OH sensitivity plot for the rate constant measurement of methyl acetate + OH

at 1091 K and 1.37 atm. .......................................................................................... 126

Figure 7.6: Sample methyl acetate + OH rate constant measurement using the mixture of

384 ppm methyl acetate with ~28.5 ppm TBHP (and 73.5 ppm water) in Ar

at 1091 K and 1.37 atm. Simulation from the Dooley et al. mechanism [49]

for the best-fit rate constant, along with perturbations of ±50%, is also

shown. ..................................................................................................................... 127

Figure 7.7: Arrhenius plot for methyl acetate + OH (k10) at temperatures above 833 K. ............ 129

Figure 7.8: OH sensitivity plot for the rate constant measurement of methyl propanoate +

OH at 1208 K and 1.33 atm. ................................................................................... 132

Figure 7.9: Sample methyl propanoate + OH rate constant measurement using the

mixture of 281 ppm methyl propanoate with ~22 ppm TBHP (and 68 ppm

water) in Ar at 1208 K and 1.33 atm. Simulation from the Dooley et al.

xxiii

mechanism [54] for the best-fit rate constant, along with variations of ±50%,

is also shown. .......................................................................................................... 133

Figure 7.10: Arrhenius plot for methyl propanoate + OH (k11) at temperatures above 870

K. ........................................................................................................................... 135

Figure 7.11: Chemical notations for fuel radicals from MCH + OH reactions used by

Orme et al. [133]. .................................................................................................... 136

Figure 7.12: OH sensitivity plot for the rate constant measurement of methyl butanoate +

OH at 1133 K and 1.37 atm. ................................................................................... 137

Figure 7.13: Sample methyl butanoate + OH rate constant measurement using the mixture

of 241 ppm methyl butanoate with ~20 ppm TBHP (and 60 ppm water) in

Ar at 1133 K and 1.37 atm. Simulation from the Dooley et al. mechanism

[54] for the best-fit rate constant, along with variations of ±50%, is also

shown. ..................................................................................................................... 138

Figure 7.14: Arrhenius plot for methyl butanoate + OH (k12) at temperatures above 870

K. ........................................................................................................................... 140

Figure 7.15: Arrhenius plots for methyl ester + OH reactions at temperatures above 250

K. ........................................................................................................................... 143

Figure 7.16: Comparison of the present rate constant measurements with the modified

SAR estimations. ..................................................................................................... 145

Figure A.1: Previous ignition delay time measurements for propane oxidation in air at Φ

= 0.5. The constant U, V model calculations utilize the Curran et al.

mechanism [100]. .................................................................................................... 157

Figure A.2: Comparison of pressure profiles for a mixture of 0.8% C3H8/ 8% N2/ Ar

obtained with and without driver insert in the Stanford 14.13 cm diameter

shock tube. The fractional pressure rise without driver insert (over 20 ms) is

approximately 20%, compared to ±3.0% local pressure variations with

driver insert. The decay beginning at 25 ms is due to arrival of the

rarefaction wave from the driver section. ............................................................... 162

Figure A.3: Comparison of pressure profiles for reactive mixture with and without LPST

driver insert. Pressure rise without driver insert (over 10 ms) is 20%,

compared to ±3.0% pressure variations with driver insert. Initial reflected

xxiv

shock conditions: T5 = 1034 K and P5 = 7.1 atm (with dP5/dt ~ 2%/ms), T5 =

1044 K and P5 = 6.7 atm (with dP5/dt ~ 0%/ms). .................................................... 163

Figure A.4: Comparison of pressure profiles for reactive mixture with and without HPST

driver insert. Pressure rise without driver insert (over 2.5 ms) is 17.5%,

compared to ±1.0% pressure variations with driver insert. Initial reflected

shock conditions: T5 = 996 K and P5 = 54.7 atm (with dP5/dt ~ 7%/ms), T5 =

1008 K and P5 = 53.7 atm (with dP5/dt ~ 0%/ms). .................................................. 164

Figure A.5: Ignition delay times for 0.8% C3H8/ 8% O2/ Ar mixture at P5 = 6 atm, plotted

at the initial post-shock T5. Experimental data and calculated values from

JetSurF v1.0 mechanism [155] and Curran et al. mechanism [100]. ...................... 166

Figure A.6: Low-pressure experimental data and CHEMSHOCK modeling using JetSurF

v1.0 mechanism. ..................................................................................................... 168

Figure A.7: High-pressure experimental data (at P5 = 24 and 60 atm), along with

CHEMKIN and CHEMSHOCK modeling using JetSurF v1.0 and Curran et

al. mechanisms. ....................................................................................................... 170

Figure B.1: n-Pentane ignition delay time measurements at pressures of 1.8 and 3.6 atm

and equivalence ratios of 1.0 and 0.5. ..................................................................... 175

Figure B.2: n-Hexane ignition delay time measurements at pressures of 1.8 and 3.6 atm


Figure B.3: n-Octane ignition delay time measurements at pressures of 1.8 and 3.6 atm


Figure B.4: n-Nonane ignition delay time measurements at pressures of 1.8 and 3.6 atm


Figure B.5: Cyclohexane (CH) ignition delay time measurements at pressures of 1.5 and

3.0 atm and equivalence ratios of 1.0 and 0.5. ........................................................ 177

Figure B.6: Methylcyclohexane (MCH) ignition delay time measurements at pressures of

1.5 and 3.0 atm and equivalence ratios of 1.0 and 0.5. ........................................... 177

Figure B.7: n-Butylcyclohexane (BCH) ignition delay time measurements at pressures of

1.5 and 3.0 atm and equivalence ratios of 0.88 and 0.45. ....................................... 178

Figure C.1: OH and C2H4 time history measurements for the mixture of 424 ppm JP-8

with 0.813% O2 in Ar. Two JP-8 proposed surrogate models were

employed. Simulations were done using the Dooley et al. mechanism [172]. ....... 181

xxv

Figure C.2: OH time histories for the mixture of ~360 ppm n-decane with 0.813% O2 in

Ar. Simulations were done using JetSurF v1.1 mechanism. An inset figure

is also shown to provide the early-time features. .................................................... 183

Figure C.3: C2H4 time histories for the mixture of ~360 ppm n-decane with 0.813% O2 in

Ar. Simulations were done using JetSurF v1.1 mechanism. .................................. 184

Figure C.4: CO time histories for the mixture of ~360 ppm n-decane with 0.813% O2 in

Ar. Simulations were done using JetSurF v1.1 mechanism. .................................. 185

Figure C.5: OH time histories for the mixture of 511 ppm iso-octane with 0.813% O2 in

Ar. Simulations were done using LLNL mechanism (iso-octane mech. v3).

An inset figure is also shown to provide the early-time features. ........................... 186

Figure C.6: CO time histories for the mixture of 511 ppm iso-octane with 0.813% O2 in

Ar. Simulations were done using LLNL mechanism (iso-octane mech. v3). ........ 187

Figure C.7: OH time histories for the mixture of 640 ppm toluene with 0.813% O2 in Ar.

Simulations were done using JetSurF v1.1 mechanism. An inset figure is

also shown to provide the early-time features. ........................................................ 188

Figure C.8: CO time histories for the mixture of 640 ppm toluene with 0.813% O2 in Ar.

Simulations were done using JetSurF v1.1 mechanism. ......................................... 189

Figure C.9: Comparison of ignition delay time measurements for JP-8, n-decane, iso-

octane, and toluene at 1.6 atm. ................................................................................ 190

1

Chapter 1 Background and Motivation

1.1 Introduction

The path towards cleaner and more efficient fuel burning is highly desirable

worldwide. To mitigate CO and particulate matter (PM) emissions in the United States,

strict federal regulations have been implemented into the engine industries. For instance,

PM and NOx emissions from heavy-duty diesel engines produced after 2010 must be

reduced by 90% of the emission levels that were recorded in 2004. Owing to these strict

regulations, further improvements in combustion chamber and fuel injection systems are

required. In conjunction with these improvements, the use of oxygenated fuels to

supplement petroleum-based fuels is also considered as one of the potential strategies in

achieving higher fuel efficiency and reducing pollutant emissions. In particular,

oxygenated fuels are known to assist in more complete combustion by adding oxygen as

part of the fuel, thereby lowering CO and hydrocarbon emissions. Hence, optimizing the

use of oxygenated fuels as neat fuels or additives is pertinent to the current engine

development.

To further understand the influence of oxygenated fuels on engine performance, a

comprehensive model, that can predict important observables such as efficiency and

pollutant emissions, is needed. This type of predictive model generally requires intensive

knowledge in the areas of heat transport, fluid dynamics, and chemistry. Additionally, in

some advanced combustion systems (e.g., homogeneous charge compression ignition

(HCCI) engines), chemistry plays a critical role in governing the overall system

performance, such as autoignition of the fuel-oxidizer mixtures, exhaust gas compositions

2

(i.e., CO, CO2, and NOx), and heat release rates. A detailed kinetic mechanism is

typically used to describe the chemistry of these particular combustion events, and is

comprised of hundreds to thousands of elementary reactions specified by rate constants,

which are strong functions of temperature and pressure. Unfortunately, most of the

existing mechanisms are likely to be error-prone, and require some sort of experimental

validation via several kinetic targets. Typical kinetic targets for these mechanisms are

ignition delay times, species concentration time histories, and direct reaction rate constant

measurements [1]. In this dissertation, the kinetics of the H2 + OH reaction is

reexamined, and the combustion characteristics of two types of oxygenated fuels (i.e.,

ketones and methyl esters) are investigated. Several important kinetic targets are

provided in order to evaluate and refine the existing kinetic mechanisms, and some

definite conclusions regarding these mechanisms can be drawn from these experimental

observations.

1.2 Background and Motivation

1.2.1 H2 + OH Kinetics

The reaction of hydroxyl radicals with molecular hydrogen

H2 + OH → H2O + H (1)


hydrogen combustion. Its reverse reaction plays a critical role in the establishment of

partial equilibrium in the post-combustion regime.

Because of its important role in combustion, numerous direct rate constant

measurements for reaction (1) have been conducted across a broad range of temperatures

[2-21]. Figure 1.1 demonstrates some of the previous experimental determinations for

reaction (1). Ravishankara and co-workers [8-9] measured the rate constant for the title

reaction (under pseudo-first-order kinetic conditions) using the flash photolysis–

resonance fluorescence technique to monitor the temporal profiles of OH decays in a

heated quartz cell over 250-1050 K. Their measurements confirmed the nonlinearity of

the Arrhenius plot for reaction (1) over this wide temperature range. Recently, Orkin et

3

al. [15] reexamined the rate constant for reaction (1) using the flash photolysis–resonance

fluorescence technique in a Pyrex reactor over a narrower temperature range of 200-479

K, and their results are in excellent agreement with the measured values from

Ravishankara and co-workers [8-9]. At combustion-relevant conditions (T > 1000 K),

the measurements for reaction (1) were generally carried out using shock tubes, and these

high-temperature data have a much larger scatter (within a factor of 2) than the low-

temperature data, as illustrated in Figure 1.1. Frank and Just [17] measured the rate

constants for the reactions of H + O2 → OH + O and H2 + OH → H2O + H by employing

atomic resonance absorption spectrometry (ARAS) to monitor H- and O-atom

concentration profiles behind reflected shock waves over 1700-2500 K. Michael and

Sutherland [18] studied the rate constant for the reaction of H + H2O → H 2 + OH

(reaction (-1)) using flash photolysis of H2O to generate H-atoms and using ARAS to

monitor the temporal profiles of H-atom decays behind reflected shock waves over 1246-

2297 K. The rate constant for reaction (1) was calculated from the measured reverse rate

constant and the equilibrium constant, and the resulting data were then compiled with the

earlier experimental work from Ravishankara and co-workers [8-9] and Frank and Just

[17] to form a three-parameter least-squares fit: k1(T) = 2.16 × 108 T1.51 exp(-3430

[cal/mol]/RT) cm3 mol-1 s-1 over 250-2500 K. This expression is currently adopted in

GRI-Mech 3.0 [22]. It is pertinent to note that the recent revised standard enthalpy of

formation for OH at 298 K [23-24] seems to suggest a lower rate constant expression for

reaction (1) (~15% lower) if the rate constant is evaluated from the reverse reaction and

the revised equilibrium constant. Figure 1.1 also shows the revised evaluations from

Michael and Sutherland [18] using the revised equilibrium constants. Similarly,

Davidson et al. [19] studied the rate constant for the reverse reaction by using laser

photolysis of H2O and UV laser absorption of OH near 307 nm behind reflected shock

waves over 1600-2500 K, and the rate constant for reaction (1) can be evaluated from

their measured values and the revised equilibrium constants, as shown in Figure 1.1.

Oldenborg et al. [20] conducted direct rate constant measurements for reaction (1) (under

pseudo-first-order kinetic conditions) using the laser photolysis / laser-induced

fluorescence technique to monitor OH radical concentration profiles in a heated cell over

4

800-1550 K. Moreover, Krasnoperov and Michael [21] reexamined the rate constant for

reaction (1) using a novel multi-pass absorption spectrometric detection technique to

monitor OH species profiles at 308 nm in reflected shock wave experiments over 832-

1359 K.

Figure 1.1: (a) Arrhenius plot for the reaction of OH with H2 at all temperatures. (b) Arrhenius plot for the reaction of OH with H2 at high temperatures (1000-2500 K).

An accurate knowledge of the rate constant for reaction (1) is important in

interpreting laminar flame speed measurements of H2-O2 mixtures. Figure 1.2 presents

the unstretched laminar flame speed measurements of H2-air mixtures as a function of

equivalence ratio at standard initial temperature (298 K) and ambient pressure (1 atm)

5

[25-32], along with the simulations from the preliminary CEFRC foundational fuel model

[33] with the rate constant for reaction (1) from Michael and Sutherland [18]. Note that

the previous rate constant measurements of reaction (1) have a relatively large scatter

(within a factor of 2) at elevated temperatures, and such large uncertainty in reaction (1)

can affect the laminar flame speed predictions of H2-air mixtures from the model by

approximately ±11%, as demonstrated in Figure 1.2.

Figure 1.2: Impact of the rate constant for the reaction of H2 + OH → H2O + H and its uncertainty (within a factor of 2) on the laminar flame speed predictions of H2-air mixtures at 298 K and 1 atm. Laminar flame speed simulations were performed using the preliminary CEFRC foundational fuel model [33] with the H2 + OH reaction rate constant from Michael and Sutherland [18].

1.2.2 Ketone Combustion Chemistry

There has been an increased interest in studying bio-derived oxygenated fuels

because of their potential to help minimize fossil fuel consumption. Among these

oxygenated fuels, methyl esters and alcohols have attracted a great deal of attention in the

form of both theoretical and experimental studies. Other oxygenates, such as ketones,

though not used as fuels, play a large role in the oxidation of the hydrocarbon and

6

oxygenated fuels. Fewer experimental studies at combustion conditions, however, have

been carried out for ketones (e.g., acetone, 2-butanone, and 3-pentanone).

Ketones are also used as fuel tracers for quantitative planar laser-induced

fluorescence measurements (PLIF) of temperature and species concentration distributions

in internal combustion engine research [34-37]. They are chosen for this purpose because

they exhibit broad absorption spectra in the ultraviolet region (π* ← n transition) and

sufficient quantum yields and strong fluorescence spectra in the visible region to be

easily monitored. 3-pentanone, in particular, is a popular ketone fuel tracer due to its

similar physical characteristics (e.g., boiling point) to that of the gasoline primary

reference fuels (n-heptane and iso-octane). The impact of 3-pentanone in internal

combustion engine studies requires more detailed information about its oxidation and

pyrolysis chemistry in order to predict its influence as an additive on the ignition

processes of the main fuel.

Moreover, ketones are listed as a class of volatile organic compounds (VOCs),

and are massively produced and used as solvents or polymer precursors in industries. As

one of the common pollutants, some amounts of ketones are emitted into the atmosphere

from a variety of natural and anthropogenic sources. The reactions with OH radicals are

the primary removal pathways for ketones in the atmosphere, which may result in the

formation of ozone and other components of the photochemical smog in urban areas [38].

In addition, these reactions of OH with ketones are one of the primary fuel consumption

pathways during oxidation, and are poorly understood at high temperatures. Hence, an

accurate knowledge of these H-atom abstraction reactions is needed in the development

of successful detailed mechanisms suitable for high-temperature applications.

1.2.3 Methyl Ester + OH Kinetics

Biodiesel is a promising alternative fuel because it has physical properties similar

to conventional crude-oil-derived fossil fuels, and it provides the opportunity to reduce

overall emissions of atmospheric pollutants [39]. Biodiesel is generally comprised of a

mixture of extended alkyl chain methyl esters 16-18 carbon atoms long [40], that are

7

typically derived from soybean oil in U.S. or rapeseed oil in Europe. Despite the

complexity of these molecules, there has been a growing effort to develop comprehensive

reaction mechanisms that can be used to describe the combustion of these large methyl

esters [41-44]. In these detailed mechanisms, the reaction rate constants for large methyl

esters are primarily based on the kinetic parameters of smaller methyl esters (e.g., methyl

formate and methyl butanoate) [41, 42, 45-47]. Thus, accurate knowledge of the kinetic

parameters for smaller methyl esters is crucial to the development of the detailed

mechanism for practical biodiesel fuels.

The combustion chemistry of small methyl esters has been a subject of interest for

the past decade. Fisher et al. [48] developed the first comprehensive chemical kinetic

mechanisms for the oxidation of methyl formate and methyl butanoate. However, the

mechanisms were validated against only a limited set of low-temperature experimental

data. Recently, Dooley et al. [49] have compiled a detailed mechanism for methyl

formate oxidation, and the mechanism has been validated against a wide variety of

experimental data, including shock tube ignition delay times, speciation data from a

variable-pressure flow reactor, and laminar burning velocities of outwardly propagating

spherical flames. Similarly, Ren et al. [50] conducted direct rate constant measurements

of the initial dissociation pathways of methyl formate over 1202-1607 K at pressures near

1.6 atm using shock tube/laser absorption techniques, and their measurements are in close

accord with the estimated values adopted in the Dooley et al. mechanism [49].

Concurrently, Peukert et al. [51-52] investigated the high-temperature thermal

decomposition and the H-atom abstraction reactions by H-atoms for methyl formate and

methyl acetate over 1194-1371 K at pressures around 0.5 atm using shock tube/atomic

resonance absorption spectrometry technique. In addition, Westbrook et al. [53]

developed a detailed mechanism for a group of four small alkyl esters, including methyl

formate, methyl acetate, ethyl formate, and ethyl acetate. The mechanism was validated

against the speciation data from fuel-rich, low-pressure, premixed laminar flames.

Similarly, Dooley et al. [54] and Hakka et al. [55] developed two separate detailed

mechanisms for methyl butanoate oxidation, and the mechanisms were tested against

different sets of experimental data, including shock tube and rapid compression machine

8

ignition delay times and speciation data from a flow reactor, a jet-stirred reactor, and an

opposed-flow diffusion flame. Moreover, numerous experimental studies [56-59] for

methyl butanoate pyrolysis and oxidation were performed in order to improve the global

performance of the existing detailed mechanisms. However, among most of these

studies, little attention has been given to a better understanding of the elementary kinetics

of these methyl esters. In particular, the H-atom abstraction reactions by OH radicals for

methyl esters, which are one of the major fuel consumption pathways during oxidation,

are not well known at combustion-relevant conditions.

1.3 Scope and Organization of Thesis

The dissertation is organized as follows:

1) Chapter 2 describes the shock tube facility and different laser absorption techniques,

which were utilized in this work to monitor key combustion radicals and intermediate

species (i.e., OH, ketones, CH3, CO, C2H4, H2O, and CH4).

2) Chapter 3 presents the high-temperature experimental determination of the important

chain-propagating reaction in all combustion systems.

H2 + OH → H2O + H (1)

The present high-temperature measurements were also compared with the previous

experimental determinations, the values employed in several detailed kinetic

mechanisms, and the theoretical calculation using semi-classical transition state

theory (SCTST).

3) Chapter 4 discusses the multi-species time history measurements during high-

temperature acetone and 2-butanone pyrolysis. In this chapter, five different species,

namely ketone, CH3, CO, C2H4, and CH4, were presented and compared with the

simulations from the detailed kinetic mechanisms. During acetone pyrolysis, the CO

and acetone concentration time histories were used to infer the rate constant for

acetone unimolecular decomposition reaction at pressures of 1.23-1.66 atm:

CH3COCH3 (+ M) → CH3 + CH3CO (+ M) (2)

9

Similarly, during 2-butanone pyrolysis, the measured 2-butanone time histories were

used to determine the overall rate constant (k3 = k3a + k3b + k3c) for 2-butanone

decomposition at pressures of 1.39-1.62 atm:

C2H5COCH3 (+ M) → C2H5 + CH3CO (+ M) (3a)

C2H5COCH3 (+ M) → CH3 + C2H5CO (+ M) (3b)

C2H5COCH3 (+ M) → CH3 + CH3COCH2 (+ M) (3c)

In addition, using the measured 2-butanone and CO time histories and an O-atom

balance analysis, a missing removal pathway for methyl ketene (one of the major

products predicted by the model) was identified.

4) Chapter 5 discusses the shock tube measurements of 3-pentanone pyrolysis and

oxidation. In this chapter, we provided six species time history measurements (i.e., 3-

pentanone, CH3, CO, C2H4, OH, and H2O), along with 3-pentanone ignition delay

time measurements. These measurements were also compared with the simulations

from the detailed kinetic mechanism. More importantly, during 3-pentanone

pyrolysis, the measured 3-pentanone and CH3 time histories were used to determine

the overall rate constant (k4 = k4a + k4b) for 3-pentanone decomposition at pressures of

1.32-1.75 atm:

C2H5COC2H5 (+ M) → C2H5 + C2H5CO (+ M) (4a)

C2H5COC2H5 (+ M) → CH3 + C2H5COCH2 (+ M) (4b)

Similar to 2-butanone pyrolysis, an O-atom balance analysis from the measured 3-

pentanone and CO time histories identified the absence of the methyl ketene

decomposition pathway in the detailed mechanism.

5) Chapter 6 provides the direct high-temperature overall rate constant measurements of

acetone + OH, 2-butanone + OH, 3-pentanone + OH, and 2-pentanone + OH

reactions.





10

6) Chapter 7 presents the first direct high-temperature overall rate constant

measurements of methyl formate + OH, methyl acetate + OH, methyl propanoate +

OH, and methyl butanoate + OH reactions.





7) Chapter 8 summarizes the present rate constant determinations of reactions (1)-(12),

and proposes some future plans, which include multi-species time history

experiments and direct rate constant measurements of H-atom abstraction reactions

for large methyl esters (i.e., methyl decanoate).

8) At long shock tube test times, as are needed at low reaction temperatures, small

gradual increases in pressure that result from incident shock wave attenuation and

boundary layer growth can significantly shorten measured ignition delay times. In

Appendix A, we investigated such pressure effects on propane ignition delay times at

pressures of 6, 24, and 60 atm.

9) Appendix B presents the ignition delay time measurements of four n-alkanes (e.g., n-

pentane, n-hexane, n-octane, and n-nonane) and three cycloalkanes (e.g.,

cyclohexane, methylcyclohexane, and n-butylcyclohexane) at various reflected shock

temperatures and pressures (between 1240 and 1500 K and 1.5 and 3.8 atm) and at

two equivalence ratios, namely Φ = 1.0 and Φ = 0.5.

10) Appendix C presents the species time history measurements of OH, C2H4, and CO

during the high-temperature oxidation of n-decane, iso-octane, and toluene, which are

the proposed surrogate fuel components for JP-8. Concurrently, the measured

ignition delay times of these surrogate fuel components were compared with the

measured JP-8 ignition delay times.

11

Chapter 2 Experimental Methods This chapter discusses the shock tube facility and different laser diagnostic

systems employed in this work.

2.1 Shock Tube Facility

Experiments were performed in a stainless-steel, high-purity, low-pressure shock

tube at Stanford. The shock tube is comprised of a 3.7-m driver section and a 10-m

driven section, with an inner diameter of 15.24 cm. The shock tube driver and driven

sections are separated by a polycarbonate diaphragm of 0.005-0.08” in thickness.

Incident shock velocity measurements were made using a series of five piezoelectric

pressure transducers (PCB 113A26 transducer, PCB 483B08 amplifier) over the last 1.5

m of the shock tube and linearly extrapolated to the endwall. Average shock velocity

attenuation rates were between 0.5-1.5% per meter. Reflected shock temperatures and

pressures were determined from the incident shock velocity at the endwall using standard

normal shock relations, with uncertainties of approximately ±0.7% and ±1%,

respectively, mainly due to the uncertainty in the measured shock velocity (±0.2%) [23].

For all experiments presented in this dissertation, vibrational equilibrium can be assumed

immediately behind the incident and reflected shock waves. In addition to the five

piezoelectric pressure transducers, a Kistler™ pressure transducer was utilized to

measure the pressure time histories upon shock-heating. All laser absorption diagnostics,

along with the Kistler™ pressure transducer, were located at a test section 2 cm from the

driven section endwall. Concurrently, prior to every experiment, the shock tube and

12

mixing assembly were routinely turbomolecular pumped down to ~5 µtorr to ensure

purity of the test mixtures, with a typical subsequent leak-plus-outgassing rate of less

than 50 µtorr/min. Further details of the shock tube facility can be found elsewhere [60-

62].

2.2 Laser Absorption Methods

2.2.1 UV Laser Absorption of OH

OH radical concentration was measured using the frequency-doubled output of a

narrow-linewidth ring dye laser near 306.69 nm, as illustrated in Figure 2.1. The laser

wavelength was tuned to the peak of the well-characterized R1(5) absorption line in the

OH A–X (0, 0) band. Visible light near 613.4 nm was generated by pumping Rhodamine

6G dye in a Spectra Physics 380A laser cavity with the 5 W, cw output of a Coherent

Verdi laser at 532 nm. The visible light was then intracavity frequency-doubled using a

temperature-tuned AD*A nonlinear crystal to generate ~1 mW of light near 306.69 nm.

Using a common-mode-rejection detection scheme, a minimum absorbance of 0.1% can

be detected, which resulted in the current experiments in a minimum detection sensitivity

of ~0.2 ppm at 1400 K and 1.5 atm (with kλ = 199.3 cm-1 atm-1). Further details of the

OH laser diagnostic setup are discussed elsewhere [23-24]. The overall estimated

uncertainty in the measured OH mole fraction (XOH) is approximately ±3%, mainly due

to the uncertainty in temperature (±0.7%). To check for the interference absorption, the

laser was also tuned away from the narrow OH absorption line by approximately 4 cm-1.

If the interference absorption was found, an OH off-line measurement was required for

each OH on-line measurement. Under the assumption that the interfering species has

wavelength-independent absorption near 306.69 nm, the interference absorbance of the

off-line measurement can be directly subtracted from the total absorbance of the OH

online measurement. OH species concentration can then be calculated from Beer’s law:

-ln(I/Io)corrected = -ln(I/Io)online + ln(I/Io)offline

-ln(I/Io)corrected = kOHXOHPL

13

where I and Io are the transmitted and incident laser intensities, kOH is the OH absorption

coefficient, XOH is the OH mole fraction, P is the total pressure, and L is the path length

(15.24 cm).

Figure 2.1: Schematic of OH detection using a narrow-linewidth ring dye laser near 306.69 nm. Figure adapted from refs. [23-24].

2.2.2 UV Laser Absorption of Ketones

Ketones (i.e., acetone, 2-butanone, 2-pentanone, and 3-pentanone) are known to

have a near-UV absorption spectrum that corresponds to the symmetry forbidden

electronic π* ← n transition where an electron from a non-binding orbital localized near

the oxygen atom is excited to an anti-bonding orbital around the CO group [34-35]. At

current experimental conditions, the spectrum is broad, lacks any fine structure, and

varies gradually from 220 to 340 nm with peak absorption at around 295 nm. This is

completely consistent with the interfering absorption seen in the OH measurements

during ketone combustion studies. Taking advantage of this fact, we have measured

ketones (off-line of OH) at 306.65 nm using the same laser system as was used for OH

measurements. Figure 2.2 shows the absorption cross-sections (at 306.65 nm) of acetone,

2-butanone, and 3-pentanone over 1050-1650 K at pressures near 1.5 atm, which were

determined by measuring the absorption (from the mixtures of 1% ketone in Ar)

immediately behind reflected shock waves when only ketone existed. The uncertainties

14

in ketone cross-sections were estimated to be ±5%. Using a common-mode-rejection

detection scheme, a minimum ketone detection sensitivity of ~300 ppm at 1300 K and 1.5

atm can be achieved (with kλ ≈ 0.52 cm-1 atm-1).

Figure 2.2: High-temperature ketone absorption cross-sections at 306.65 nm.

2.2.3 UV Laser Absorption of CH3

Methyl radical has a wide predissociatively broadened absorption feature (B2A′1–

X2A″2) near 216 nm, with peak absorption at 216.6 nm. In this work, CH3 was measured

using the frequency-quadrupled output of near-infrared radiation from a pulsed

Ti:Sapphire laser, as illustrated in Figure 2.3. Stable mode-locking of the Ti:Sapphire

laser (MIRA HP, Coherent Inc.) was obtained at a wavelength of 866.4 nm with a peak

output of 1 W, a repetition rate of approximately 76 MHz, and a pulse duration of

approximately 2 picoseconds. Deep UV light was generated by frequency conversion,

using fourth harmonic generation (FHG, Coherent Inc.), to obtain output at 216.6 nm.

Further details of this laser setup can be found elsewhere [63]. Similar to the other laser

absorption methods described here, using a common-mode-rejection detection scheme, a

minimum absorbance of ~0.1% can be detected, resulting in a minimum detection

sensitivity of ~1 ppm at 1300 K and 1.5 atm (with kλ = 55.8 cm-1 atm-1).

15

It is also known that there is some interference absorption near 216 nm from the

intermediate products of pyrolysis, such as ethylene, higher olefins and conjugated

olefins. To account for the interference absorption, another wavelength at 218.7 nm was

also employed in this study. Here we assume that the interference absorbances at 216.6

nm and 218.7 nm are identical, based on the fact that these two wavelengths are very

close to each other, and at high temperatures, these hydrocarbons show wide broad

absorption features near 216 nm. CH3 absorption coefficients at 216.6 nm and 218.7 nm

were previously determined in our laboratory [64-65]. Thus, methyl radical

concentration can then be calculated as follows:

-ln(I/Io)217 = kCH3,217XCH3PL + kintXintPL

-ln(I/Io)219 = kCH3,219XCH3PL + kintXintPL

-ln(I/Io)217 + ln(I/Io)219 = (kCH3,217 – kCH3,219)XCH3PL

where kCH3 is the CH3 absorption coefficient, XCH3 is the CH3 mole fraction, P is the total

pressure, and L is the path length. Based on the two-wavelength subtraction scheme, the

interference absorption contributes about 25% to the total absorption signal at 216.6 nm.

Figure 2.3: Schematic of CH3 detection near 216.6 nm using the frequency-quadrupled output of near-infrared radiation from a pulsed Ti:Sapphire laser. Figure adapted from ref. [63].

16

2.2.4 IR Laser Absorption of CO

A quantum cascade laser (QCL) operating in cw mode has recently become an

important diagnostic tool for many combustion products, such as CO, CO2 and H2O.

This mid-IR CO laser allows access to the R(13) transition line in the CO fundamental

rovibrational band at 4.56 µm, where H2O and CO2 absorption interference is minimal.

As compared to previous CO diagnostics near 1.2, 1.5 and 2.3 µm, this new diagnostic

scheme offers orders-of-magnitude greater sensitivity, resulting in ppm-level CO

detectivity in shock tube measurements.

A fixed-wavelength direct-absorption strategy was employed to monitor the peak

intensity of the R(13) absorption line at 2193.359 cm-1 (with kλ = 12.1 cm-1 atm-1 at 1300

K and 1.5 atm), as illustrated in Figure 2.4. The spectroscopic parameters for the R(13)

transition, including the line-strength and self-broadening coefficient, were taken directly

from the HITRAN database. The collisional broadening coefficient for CO with argon

(not available in HITRAN) was measured in the shock tube over the temperature range of

1000-1800 K. Further details regarding the CO diagnostic setup are described elsewhere

[66].

Figure 2.4: Schematic of CO detection using cw quantum cascade laser near 4.56 µm. Figure adapted from ref. [66].

17

2.2.5 CO2 Laser Absorption of C2H4

C2H4 was measured using CO2 laser absorption at 10.532 µm, as illustrated in

Figure 2.5. This diagnostic takes advantage of the strong overlap of the P(14) line of the

CO2 laser transition with the strong Q-branch of the ν7 ethylene band. It is capable of

detecting 100 ppm levels of C2H4 over a path length of 15 cm at 1200 K and 3 atm (with

σλ = 9.69 m2/mol or kλ = 0.98 cm-1 atm-1). Details regarding this diagnostic and the C2H4

absorption cross-sections are described elsewhere [67]. There is some weak interference

absorption directly from acetone, 2-butanone, and 3-pentanone at this wavelength. In

addition, higher alkenes, such as propene and butene, have weak absorption features at

10.532 µm. However, the pyrolysis of acetone, 2-butanone, and 3-pentanone generate

negligible amounts of these higher alkenes, and thus, the primary interfering species are

ketones. In the present study, C2H4 was also measured at 10.675 µm to account for

ketone interference absorbance; C2H4 absorption cross-sections at 10.675 µm were found

to be 4.0 ± 0.1 m2/mol over 1100-1500 K at 1-6 atm [68]. Moreover, ketone absorption

cross-sections at 10.532 µm and 10.675 µm were determined by measuring the

absorption immediately behind reflected shock waves when only ketone existed. Thus,

using these two wavelengths, C2H4 species concentration can be found by solving the

following two equations:

-ln(I/Io)10.532 = nC2H4σC2H4,10.532L + nketσket,10.532L

-ln(I/Io)10.675 = nC2H4σC2H4,10.675L + nketσket,10.675L

where σC2H4 and σket are the absorption cross-sections of C2H4 and ketone, nC2H4 and nket

are the number densities of C2H4 and ketone, and L is the path length.

18

Figure 2.5: Schematic of C2H4 detection using cw CO2 laser near 10.5 µm. Figure adapted from ref. [67].

2.2.6 IR Laser Absorption of H2O

H2O concentration was measured using a DFB (distributed feedback) diode laser

at 2550.96 nm (3920.09 cm−1) within the ν3 fundamental vibrational band, as illustrated

in Figure 2.6. This absorption feature has been well-characterized previously in our

laboratory [69]. During experiments, the beam path (outside the shock tube) was

continuously purged with pure N2 to minimize the laser attenuation due to ambient water.

A minimum H2O detection sensitivity of ~40 ppm can be achieved at 1400 K and 1.5 atm

(with kλ = 2.05 cm-1 atm-1).

Figure 2.6: Schematic of H2O detection using cw DFB laser near 2.55 µm. Figure adapted from ref. [69].

19

2.2.7 IR Laser Absorption of CH4

CH4 was measured near 3.4 µm using a scanned-wavelength mid-IR laser

absorption diagnostic developed by Pyun et al. [70], as illustrated in Figure 2.7. Mid-IR

light near 3.4 µm was generated using difference-frequency-generation (DFG) of a near-

IR signal laser and a near-IR pump laser combined in a PPLN crystal. Due to the

structural differences of the absorption spectrum of methane and other hydrocarbons near

3.4 µm, a differential absorption (peak minus valley) scheme was employed to obtain

interference-free CH4 concentration. To attain this scheme, the signal laser was current

modulated at a 50 kHz scanning frequency and a 0.5 Vp-p scanning amplitude by a

function generator. Then the modulated signal laser was combined with the pump laser

through the PPLN crystal to create a modulated mid-IR laser light near 3.4 µm that

included the peak and valley wavelengths in each scan. To maximize the signal-to-noise

ratio of CH4 and minimize that of the interfering species, 2917.64 cm-1 and 2917.45 cm-1

were selected as the optimal peak and valley wavelength pair in this work. Using this

method, CH4 concentration time histories with a time resolution of 20 µs and a minimum

detection sensitivity of ~250 ppm at 1300 K and 1.5 atm were obtained (with a

differential absorption coefficient of kpeak-valley = 0.73 cm-1 atm-1).

Figure 2.7: Schematic of CH4 detection using a scanned-wavelength mid-IR laser near 3.4 µm. Figure adapted from ref. [70].

20

2.3 Summary

In this dissertation, a low-pressure shock tube facility and several laser absorption

diagnostics were utilized to monitor key combustion radicals and intermediate species,

including OH radicals near 306.69 nm, ketones at 306.65 nm, CH3 radicals near 216 nm,

CO near 4.56 µm, C2H4 near 10.5 µm, H2O near 2.55 µm, and CH4 near 3.4 µm.

21

Chapter 3 A Shock Tube Study of H2 + OH → H2O + H using OH Laser Absorption

3.1 Introduction

As introduced in Chapter 1, the reaction of OH with molecular hydrogen

H2 + OH → H2O + H (1)


hydrogen combustion. Its reverse reaction plays a critical role in the establishment of

partial equilibrium in the post-combustion regime. At elevated temperatures, the

previous rate constant evaluations have an uncertainty factor of 2, and this relatively large

uncertainty has a significant impact on the laminar flame speed predictions of H2-air

mixtures from the detailed kinetic mechanism. Thus, there is motivation to reduce the

existing experimental uncertainty in k1 at combustion-relevant conditions. In this

chapter, we aim to report the rate constant for reaction (1) with a much lower

experimental scatter and overall uncertainty over the temperature range of 902-1518 K.

3.2 Experimental Details

Test mixtures were prepared manometrically in a 40 liter stainless-steel tank

heated uniformly to 50 oC and mixed with a magnetically-driven stirring vane. A double-

dilution process was employed to allow for more accurate pressure measurements in the

22

manometrical preparation of a highly dilute mixture. A highly concentrated mixture was

first prepared and mixed for at least 2 hours to ensure homogeneity and consistency, and

the mixture was then further diluted with argon and mixed for additional 2 hours prior to

the experiments. The gases utilized in this work were hydrogen (Research Grade)

99.999% and argon (Research Grade) 99.999%, which were supplied by Praxair and used

without further purification. The liquid chemical was commercially available 70% tert-

butyl hydroperoxide (TBHP) in water from Sigma-Aldrich, and was purified using a

freeze-pump-thaw procedure to remove dissolved volatiles and air prior to mixture

preparation.

3.3 Kinetic Measurements

3.3.1 Kinetic Mechanism Description

A series of 21 reflected shock wave experiments were conducted to determine the

rate constant for the reaction of H2 + OH → H2O + H over the temperature range of 902-

1518 K at pressures of 1.15-1.52 atm. Dilute test mixtures with 94-125 ppm TBHP (and

water) and 1001-1516 ppm H2 in argon were used to minimize the temperature drop

caused by the chemistry effects, and the temperature profile behind the reflected shock

wave was nearly constant (<1 K change) over the time frame of the experiment. In the

present study, the CHEMKIN PRO package [71] was used to simulate the consumption

of OH radicals by molecular hydrogen under the standard constant energy and volume

assumption, and a comprehensive reaction mechanism of Wang et al. (USC-Mech v2.0)

[72] was selected as the base mechanism. This mechanism consists of 111 species and

784 elementary reactions, and has been validated against a series of shock tube ignition

delay times, laminar flame speeds, and speciation data from a shock tube and a flow

reactor during high-temperature oxidation of H2, CO, and C1-C4 hydrocarbons. It is

pertinent to note that the conclusions of the present study are effectively independent of

the mechanism used, and near-identical results could be obtained using the GRI-Mech 3.0

[22].

23

Tert-butyl hydroperoxide (TBHP or (CH3)3−CO−OH) was chosen as an OH

radical precursor in the present study, because it decomposes very rapidly to form an OH

radical and a tert-butoxy radical, (CH3)3CO, at temperatures greater than 1000 K [73].

(Note that the weakest bond in TBHP is the O−O bond with the bond dissociation energy

of ~47 kcal/mol at 298 K.) The tert-butoxy radical further decomposes to form acetone

and a methyl radical. Additionally, TBHP reacts with OH to form other products, and

hence a TBHP sub-mechanism was also incorporated into the base mechanism, i.e.

(CH3)3−CO−OH → (CH3)3CO + OH (13)

(CH3)3CO → CH3COCH3 + CH3 (14)

(CH3)3−CO−OH + OH → H2O + O2 + tert-C4H9 (15)

(CH3)3−CO−OH + OH → H2O + HO2 + iso-C4H8 (16)

The rate constants for reactions (13), (15) and (16) were adopted from Pang et al.

[74], and the rate constant for reaction (14) was obtained from Choo and Benson [75].

The rate constants for reactions (13)-(16) are listed in Table 3.1. In addition, the

thermodynamic parameters for TBHP and tert-butoxy radical were taken from the

thermodynamic database from Goos et al. [76], and the standard enthalpy of formation

for OH radical was updated with the measured value from Herbon et al. [23-24].

Table 3.1: Reactions Describing H2 + OH Experiments at P = 1.3 atm.

Rate Constant Reaction A [†] b E [cal/mol] No. Reference

H2 + OH → H2O + H see text 1 this work TBHP → (CH3)3CO + OH 3.57E+13 0 3.575E+04 13 [74]

(CH3)3CO → CH3COCH3 + CH3 1.26E+14 0 1.530E+04 14 [75] TBHP + OH → H2O + O2 + tert-C4H9 2.30E+13 0 5.223E+03 15 [74]

TBHP + OH → H2O + HO2 + iso-C4H8 2.49E+13 0 2.655E+03 16 [74] CH3 + OH → CH2(s) + H2O 1.65E+13 0 0 17 [74]

C2H6 (+ M) → CH3 + CH3 (+ M) 1.88E+50 -9.72 1.073E+05 18 [80] Low-Pressure Limit: 3.72E+65 -13.14 1.015E+05 Troe centering: 0.39 100 1900 6000

CH3COCH3 + OH → CH3COCH2 + H2O 3.30E+13 0 4.840E+03 5 [82] CH3OH + M → CH3 + OH + M 5.62E+15 0 6.128E+04 19 [77]

† Units of A are in s-1 for unimolecular reactions and cm3 mol-1 s-1 for bimolecular reactions.

24


A local OH sensitivity analysis for the mixture of 1001 ppm H2 with ~26 ppm

TBHP (and 99 ppm H2O) in Ar at 1228 K and 1.29 atm is shown in Figure 3.1. The OH

sensitivity is calculated as SOH = (∂XOH/∂ki)×(ki/XOH), where XOH is the local OH mole

fraction and ki is the rate constant for reaction i. The analysis reveals that the OH time

history is predominantly sensitive to reaction (1) over the time frame of the experiment,

with some minor interference from the following secondary reactions:

CH3 + OH → CH2(s) + H2O (17)

C2H6 (+ M) → CH3 + CH3 (+ M) (18)


The rate constant for reaction (17) was updated with the value of 1.65×1013 cm3

mol-1 s-1 recently inferred by Pang et al. [74], which is in good agreement with the

measurements from Srinivasan et al. [77] and Vasudevan et al. [78] and the theoretical

calculation from Jasper et al. [79] (within ±35%). The rate constant for reaction (18) was

updated with the measured values from Oehlschlaeger et al. [80], and the measurements

from Oehlschlaeger et al. are in close accord with another experimental study from Kiefer

et al. [81]. Recently, Lam et al. [82] have measured the rate constant for reaction (5)

using UV laser absorption of OH near 306.69 nm behind reflected shock waves over 872-

1355 K at pressures near 2 atm, and their measured rate constant was adopted for reaction

(5) in the present study. (For more details, please read Chapter 6.) In addition, the rate

constant for the reaction of CH3OH + M → CH 3 + OH + M (reaction (19)) was updated

with the measured values from Srinivasan et al. [77] at ~0.3-1.1 atm, and their values

agree well with the theoretical calculation from Jasper et al. [79] and the measurements

from Vasudevan et al. [78] at 1.3 atm. The rate constants for reactions (5) and (17)-(19)

are also provided in Table 3.1.

25

Figure 3.1: OH sensitivity plot for the rate constant measurement of H2 + OH at 1228 K and 1.29 atm.

Figure 3.2 shows an OH time history measurement for the mixture of 1001 ppm

H2 in argon at 1228 K and 1.29 atm, and the measured peak OH mole fraction is

approximately 26 ppm. Due to wall adsorption and condensation of TBHP, the initial

TBHP mole fraction was assumed to be the same as the measured peak OH mole fraction.

This assumption is valid because the OH radical is formed very rapidly after the thermal

decomposition of TBHP at T > 1000 K. More importantly, the presence of H2O in the

mixture has a negligible influence on the simulated OH time history; hence, its presence

would not affect our rate constant evaluation. As illustrated in Figure 3.2, a best-fit rate

constant for reaction (1) of 2.45×1012 cm3 mol-1 s-1 was obtained between the

experimental data and the simulation at 1228 K and 1.29 atm. Simulations for the

variations of ±20% in the inferred rate constant are also shown in Figure 3.2. Similarly,

Figure 3.3 shows the measured OH time histories for different test mixtures at different

temperatures, along with the simulations from USC-Mech v2.0 [72] for the best-fit rate

constants. Interestingly, our measured values are identical to or very close to the values

originally proposed by Michael and Sutherland [18], within ±4% at most temperatures,

and their expression is currently adopted in GRI-Mech 3.0 [22]. Additionally, Table 3.2

summarizes the rate constant measurements of reaction (1) at T = 902-1518 K and P =

1.15-1.52 atm.

26

Figure 3.2: Sample H2 + OH rate constant measurement using the mixture of 1001 ppm H2 with ~26 ppm TBHP (and 99 ppm water) in Ar at 1228 K and 1.29 atm. Simulation from USC-Mech v2.0 [72] for the best-fit rate constant, along with variations of ±20%, is also shown.

Figure 3.3: H2 + OH rate constant measurements at various temperatures, along with the simulations from USC-Mech v2.0 for the best-fit rate constants.

27

Table 3.2: Rate Constant Data for H2 + OH → H2O + H.

T5 [K] P5 [atm] k1 [cm3 mol-1 s-1]

94 ppm TBHP (and water), 1516 ppm H2, Ar 1343 1.16 3.12E+12 1053 1.36 1.48E+12 984 1.34 1.31E+12 972 1.46 1.19E+12 933 1.52 1.03E+12

95 ppm TBHP (and water), 1500 ppm H2, Ar

1466 1.15 4.10E+12 1405 1.18 3.68E+12 1279 1.23 2.83E+12 1225 1.25 2.38E+12 1217 1.19 2.33E+12 1157 1.30 2.05E+12 1152 1.24 2.03E+12 1140 1.33 1.96E+12 1098 1.35 1.75E+12 981 1.47 1.25E+12 902 1.49 9.23E+11

125 ppm TBHP (and water), 1001 ppm H2, Ar

1518 1.15 4.71E+12 1228 1.29 2.45E+12 1189 1.32 2.20E+12 1058 1.43 1.56E+12 1032 1.42 1.40E+12

A detailed error analysis was performed to estimate the overall uncertainty in the

measured rate constant for reaction (1) at 1228 K. The following contributions were

considered: (a) temperature (±1%), (b) OH absorption coefficient (±3%), (c) wavemeter

reading in the UV (±0.01 cm-1), (d) fitting the data to computed profiles (±5%), (e)

locating time-zero (±0.5 µs), (f) the rate constant for CH3 + OH → CH2(s) + H2O

28

(uncert. factor = 2), (g) the rate constant for C2H6 (+ M) → CH3 + CH3 (+ M) (±20%),

and (h) the rate constant for CH3COCH3 + OH → CH3COCH2 + H2O (±28%). As

demonstrated in Figure 3.4, the individual error sources were introduced independently

(within the estimated positive and negative bounds of their 2σ uncertainties) and their

effects on the rate constant for the title reaction were studied. These uncertainties were

combined in a root-sum-squared method to give an overall (2σ) uncertainty of ±17% at

1228 K. Similar error analysis was conducted for k1 at 972 K, and the overall (2σ)

uncertainty was also estimated to be ±17%.

Figure 3.4: Uncertainty analysis for the rate constant of H2 + OH → H2O + H at 1228 K and 1.29 atm.

Figure 3.5 shows the Arrhenius plot for the present rate constant measurements of

reaction (1) at T = 902-1518 K and P = 1.15-1.52 atm, along with the non-Arrhenius

expression originally proposed by Michael and Sutherland [18]. The measured values

from the present study can be expressed in Arrhenius form as k1(T) = 4.38×1013 exp(-

3518/T) cm3 mol-1 s-1 over 902-1518 K. As illustrated in Figure 3.5, the current data

have a relatively low scatter (<7%). In the present study, three different mixture

compositions were utilized to demonstrate that the inferred rate constants are not strongly

29

dependent on any secondary chemistry contributions, and the measured values from these

mixtures are consistent with each other. It is interesting to note that the present

measurements are in excellent agreement with the non-Arrhenius expression proposed by

Michael and Sutherland (within ±6%).

Figure 3.5: Arrhenius plot for H2 + OH (k1) at temperatures above 833 K.

3.4 Comparison with Earlier Work

Figure 3.6 presents the current data along with some earlier measurements of

reaction (1) at temperatures above 833 K. Note that the present measurements have a

much lower scatter (<7%) than the previous work. Frank and Just [17] investigated the

rate constant for reaction (1) using the test mixtures with a few ppm N2O and 100-500

ppm H2 and O2 in Ar behind reflected shock waves over 1700-2500 K. In their

experiments, the thermal dissociation of N2O upon shock-heating was used to generate

O-atoms, followed by the reaction of O + H2 → OH + H to produce H-atoms and OH

radicals. These OH radicals would then react with H2 through reaction (1), and atomic

resonance absorption spectrometry (ARAS) was used to monitor H- and O-atom

concentration profiles simultaneously. These measured species time histories were also

30

quite sensitive to another major reaction (H + O2 → OH + O), which was one of the

targeted reactions in their study. Their estimated uncertainty limit for k1 was found to be

less than ±40%. Michael and Sutherland [18] examined the reverse rate constant using

the test mixtures of 0.1-1% H2O in Ar behind reflected shock waves over 1246-2297 K.

H-atoms were generated in the post-shock regions using the flash photolysis of H2O and

were monitored using ARAS. In the present study, their measured values were also

converted to the forward rate constants using the revised equilibrium constants (with the

revised standard enthalpy of formation for OH from Herbon et al. [23-24]), as shown in

Figure 3.6. The new calculated values for reaction (1) are approximately 15% lower than

the old values evaluated by Michael and Sutherland. Such differences are solely due to

the equilibrium constant evaluations. Nevertheless, their values for reaction (1) have a

relatively large experimental scatter (within a factor of 2). Similarly, Davidson et al. [19]

measured the reverse rate using the test mixtures of 0.915-1.83% H2O in Ar behind

reflected shock waves over 1600-2500 K. An ArF excimer laser at 193.3 nm was

employed to photo-dissociate a small amount of H2O in order to generate H-atoms and

OH radicals, and a cw, narrow-linewidth ring dye laser at 306.6 nm was then used to

monitor the temporal evolution of OH radicals. The errors of their measurements varied

from ±40% at 1600 K to ±12% at 2500 K. Their measured values were also converted to

the forward rates using the revised equilibrium constants.

As illustrated in Figure 3.6, the data from Frank and Just [17], Michael and

Sutherland [18], and Davidson et al. [19] are in good agreement with each other (within

their experimental scatter). Oldenborg et al. [20] also conducted direct rate constant

measurements of reaction (1) using the laser photolysis / heated flow cell technique and

using laser-induced fluorescence method to monitor the temporal profiles of OH decays

at 800-1550 K. Their experiments were performed under pseudo-first-order kinetic

conditions with an excess of H2. In addition, Ravishankara et al. [9] studied reaction (1)

using the flash photolysis–resonance fluorescence technique (under pseudo-first-order

kinetic conditions) in a heated quartz cell over 250-1050 K, and their high-temperature

measurements (at 960 and 1050 K) are ~20% higher than the measurements from

Oldenborg et al. [20]. Moreover, Krasnoperov and Michael [21] reexamined reaction (1)

31

using a test mixture of ~5 ppm TBHP with 404 ppm H2 in krypton behind reflected shock

waves at lower temperatures (832-1359 K). TBHP was utilized as the OH precursor in

their study, and a novel multi-pass absorption spectrometric detection method (with a

MW discharge driven resonance OH lamp) was employed to measure OH species

profiles. Although their experimental scatter is slightly high (±25%), their measured

values are generally in good agreement with Oldenborg et al. [20] and Ravishankara et al.

[9]. Concurrently, the present measurements are in excellent agreement with all previous

studies.

Figure 3.6 also shows the values of k1 employed in three different comprehensive

reaction mechanisms: GRI-Mech 3.0 [22], USC-Mech v2.0 [72], and Hong et al. [83].

The present measurements agree well with the values from GRI-Mech 3.0 and Hong et al.

(within ±6%). However, the values of k1 from USC-Mech v2.0 are ~20% lower than the

present measurements. Concurrently, Nguyen et al. [84] computed the rate constant for

the title reaction with semi-classical transition state theory (SCTST), which implemented

non-separable coupling among all degrees of freedom (including the reaction coordinate)

in the transition state region and multi-dimensional quantum mechanical tunneling along

the curved reaction path. Their theoretical calculation (the black dashed line in Figure

3.6) is also in excellent agreement with the present and previous experimental studies,

and their calculation is nearly indistinguishable from the rate constant adopted in GRI-

Mech 3.0.

32

Figure 3.6: Comparison with previous studies at temperatures above 833 K.

3.5 Summary

The rate constant for the reaction of H2 + OH → H 2O + H was studied behind

reflected shock waves over the temperature range of 902-1518 K at pressures of 1.15-

1.52 atm using OH laser absorption. The current high-temperature data can be expressed

in Arrhenius equation as k1(T) = 4.38×1013 exp(-3518/T) cm3 mol-1 s-1 over the

temperature range studied. A detailed error analysis was carried out with the

consideration of both experimental and secondary chemistry contributions, and the

overall (2σ) uncertainties in k1 were found to be ±17% at 972 and 1228 K. Note that the

experimental scatter from the present study is less than 7%, which is much lower than in

previous work [17-21]. The present data are consistent with the previous measurements

from Frank and Just [17], Michael and Sutherland [18], Davidson et al. [19], Oldenborg

et al. [20], and Krasnoperov and Michael [21]. Additionally, the present measurements

are in excellent agreement with the non-Arrhenius expression from GRI-Mech 3.0 [22]

and the recent theoretical calculation using semi-classical transition state theory (SCTST)

from Nguyen et al. [84].

33

Chapter 4 Multi-Species Time History Measurements during High-Temperature Acetone and 2-Butanone Pyrolysis

4.1 Introduction

The pyrolysis of acetone (IUPAC name: propanone) has been studied by many

researchers, particularly at temperatures below 1000 K. At temperatures above 1000 K,

five recent studies have been performed. Capelin et al. [85] examined acetone pyrolysis

utilizing flash vaporization into a heated reaction chamber, and suggested a pyrolysis

mechanism with CH3COCH3 (+ M) → CH3COCH2 + H (+ M) as the initiation reaction,

based on the product distributions from gas chromatography. Ernst et al. [86] studied

acetone pyrolysis using a shock tube and UV laser absorption technique. They

recommended a similar mechanism with a different initiation reaction:

CH3COCH3 (+ M) → CH3 + CH3CO (+ M) (2)

Similarly, Sato and Hidaka [87] investigated acetone pyrolysis and oxidation in a shock

tube; they evaluated the rate constant (k2) for reaction (2) by monitoring acetone

concentrations using laser absorption at 200 nm and 3.39 µm. Saxena et al. [88]

performed direct rate constant measurements of reaction (2) using laser-schlieren

technique. And finally, Pichon et al. [89] developed a detailed kinetic mechanism for

acetone, which was validated against their flame speed and ignition delay time

measurements.

34

In contrast to acetone, fewer 2-butanone studies have been performed. Early low-

temperature 2-butanone oxidation static reactor studies were conducted by Bardwell and

Hinshelwood [90-93]. Decottignies et al. [94] investigated 2-butanone oxidation using

laminar premixed methane/air flames doped with different amounts of 2-butanone.

Based on the product distributions from gas chromatography, they postulated a kinetic

mechanism and suggested three initial decomposition pathways:

2-Butanone (+ M) → C2H5 + CH3CO (+ M) (3a)

2-Butanone (+ M) → CH3 + C2H5CO (+ M) (3b)

2-Butanone (+ M) → CH3 + CH3COCH2 (+ M) (3c)

with channel (3b) as the dominant pathway. Similarly, Serinyel et al. [95] developed a

comprehensive kinetic mechanism, which was validated against their shock tube ignition

delay times. However, they suggested that channel (3a) is the primary initial 2-butanone

decomposition channel.

This chapter presents high-temperature pyrolysis studies of acetone and 2-

butanone behind reflected shock waves using laser absorption methods to measure time

histories of five species: ketone, CO, CH3, CH4, and C2H4. These measurements were

used to determine the decomposition rate constants for both acetone and 2-butanone.


4.2.1 Mixture Preparation

Test mixtures were prepared manometrically in a 40 liter stainless steel tank

heated uniformly to 50 oC and mixed with a magnetically-driven stirring vane for at least

2 hours prior to the experiments. The mixture compositions were 0.25%, 1%, and 1.5%

ketone in argon. Research grade argon (99.999%) from Praxair was used with the ACS

spectrophotometric grade acetone (≥ 99.5%) and the CHROMASOLV® grade 2-butanone

(≥ 99.7%) from Sigma Aldrich. All liquid chemicals were purified using a freeze-pump-

thaw procedure to remove dissolved volatiles and air.

35

4.2.2 Species Absorption Coefficient Evaluations

Individual species time histories for dilute fuel mixtures (0.25% acetone or 2-

butanone in Ar) were determined from Beer’s law using a constant absorption coefficient

value for each species, evaluated at the initial reflected shock temperatures and pressures.

On the other hand, for the higher fuel concentration mixtures (1% and 1.5% ketone in

Ar), the changes in temperature and pressure that occur during pyrolysis slightly

perturbed the absorption coefficients of CO, C2H4 and CH4 up to 10%, 7% and 15%,

respectively, for experiments with initial temperatures higher than 1400 K, and the use of

constant absorption coefficients in determining the species mole fractions was not valid.

To determine the experimental species time histories more quantitatively, the temperature

and pressure profiles were first calculated by solving the energy equation under the

standard constant energy (U) and volume (V) assumption (using CHEMKIN PRO [71]).

The species mole fraction time histories were then inferred from the measured absorption

data using values of the absorption coefficients evaluated at the simulated T and P.

4.3 Results and Discussion

4.3.1 Acetone Pyrolysis

High-temperature acetone pyrolysis was investigated using five species time

history measurements over the temperature range of 1200-1650 K at pressures around 1.6

atm. In the present study, the CHEMKIN PRO package [71] was used to simulate all

species time histories under the standard constant energy and volume assumption. Fig.

4.1 shows the CO sensitivity analysis for the mixture of 0.25% acetone in argon at 1393

K and 1.55 atm simulated using the Pichon et al. mechanism of NUI Galway [89]. The

CO time history is predominantly sensitive to reaction (2), with some minor interference

from the reactions of C2H6 (+ M) → CH3 + CH3 (+ M), C2H6 + H → C2H5 + H2, and CH3

+ CH3 → C2H5 + H. Similar result for the CO sensitivity analysis was obtained for the

mixture of 1% acetone in Ar.

36

Figure 4.1: CO sensitivity for 0.25% acetone in Ar using the Pichon et al. mechanism [89].

Fig. 4.2 shows the CO rate of production (ROP) analysis for the mixture of 0.25%

acetone in Ar at 1393 K and 1.55 atm simulated using the Pichon et al. mechanism [89].

The ROP analysis reveals that the primary CO formation pathway is via the reaction of

CH3CO (+ M) → CH3 + CO (+ M) over the time frame of the experiment, and the acetyl

(CH3CO) radical is directly formed from reaction (2). In particular, the CH3CO radical is

rather short-lived at T > 1100 K; once it is formed, it will decompose near-

instantaneously to form a CH3 radical and a CO molecule. According to the CO

sensitivity and ROP analysis, the measured CO time histories can then be used to infer

the acetone dissociation rate constant (k2) over the temperature range of 1150-1600 K at

pressures of 0.54-5.28 atm. The measured CO mole fractions were best-fit with the

simulated profiles by varying the value of k2 in the detailed kinetic mechanism of Pichon

et al.

37

Figure 4.2: CO rate of production (ROP) plot for 0.25% acetone in Ar using the Pichon et al. mechanism [89]. Fig. 4.3 shows a sample measured CO concentration time history during acetone

pyrolysis for the mixture of 0.25% acetone in argon at 1393 K and 1.55 atm, and the best-

fit rate constant (k2) of 3516 s-1, along with perturbations of ±30%, using the Pichon et al.

mechanism. The best-fit simulation curve is virtually indistinguishable from the data.

Figure 4.3: Sample CO time histories: measured and calculated values.

38

The present rate constant measurements (k2) of reaction (2) (inferred from the CO

profiles) at pressures near 1.6 atm are summarized in an Arrhenius diagram, Fig. 4.4,

along with three recent evaluations [87-89]. It should be noted that the acetone

concentration was varied from 0.25% to 1% in order to confirm that the current rate

constant measurements were weakly dependent on any secondary reaction effects, and

the measured values from these two mixtures were consistent with each other. Using a

least-squares fit, the acetone dissociation rate constant can be expressed in Arrhenius

form as k2 = 2.46×1014 exp(-69.3 [kcal/mol]/RT) s-1 over the temperature range of 1200-

1600 K at pressures around 1.6 atm (see the dashed line in Fig. 4.4). The primary

contributions to uncertainties in the rate constant were: temperature (±10%), CO

absorption cross-section (±5%), fitting the data to computed profiles (±5%), and

uncertainties resulting from secondary reactions (±5%), giving an overall uncertainty in

k2 of ±25%. Similar experiments performed at different pressures (0.5-0.7 atm and 5

atm) show only a very weak pressure-dependence for k2, confirming the assumption that

these measurements were performed close to the high-pressure limit. In addition, Table

4.1 summarizes the rate constant measurements of acetone dissociation reaction that are

inferred from the measured CO profiles over 1150-1600 K at pressures of 0.54-5.28 atm.

Figure 4.4: Summary of acetone dissociation rate constant (k2).

39

Similarly, as is evident in Fig. 4.5, the acetone sensitivity analysis shows that the

acetone concentration time history is strongly sensitive to reaction (2), with some minor

interference from the reactions of C2H6 (+ M) → CH3 + CH3 (+ M), CH3 + CH3 → C2H5

+ H, and CH3COCH3 + H → CH 3COCH2 + H2. Hence, the measured acetone time

histories can also be used to determine the acetone dissociation rate constant (k2).

Figure 4.5: Acetone sensitivity for 1% acetone in Ar using the Pichon et al. mechanism [89].

Fig. 4.6 shows the measured acetone time histories, along with the simulations

from the original Pichon et al. mechanism. Similar to the previous analysis based on the

CO profiles, the experiments suggest much faster acetone removal rates than the original

Pichon et al. mechanism. To infer the acetone dissociation rate constant, the measured

acetone concentration time histories were also best-fit with the simulated profiles by

varying the value of k2 in the detailed kinetic mechanism of Pichon et al. [89], as depicted

in Fig. 4.6. Table 4.1 also includes the rate constant measurements of reaction (2) that

are inferred from the acetone profiles over 1340-1650 K at pressures of 1.23-1.41 atm.

The inferred values based on the acetone profiles are consistent with those from the CO

profiles. Note that the measurements seem to experience some slight non-Arrhenius

curvature at higher temperatures, as illustrated in Fig. 4.4. By compiling the rate constant

40

measurements based on the CO and acetone time histories, the acetone dissociation rate

constant can be represented by the following three-parameter equation:

k2(1.23-1.66 atm) = 9.38 × 1041 T-7.85 exp(-44,236/T) s-1

over the temperature range of 1200-1650 K.

Figure 4.6: Acetone time histories for 1% acetone in Ar: measured and simulated values. As illustrated in Fig. 4.4, the measurements of Saxena et al. [88] and the high-

temperature measurements of Sato and Hidaka [87] are in good agreement with the

current results (within 30% at T > 1450 K) at pressures near 1.6 atm. However, at lower

temperatures, the determination from Sato and Hidaka departs significantly from the

current study. For instance, their determination is at least three times slower than the

measured rate constant from the current study at T < 1250 K. In addition, the theoretical

estimate by Pichon et al. [89] using a chemical activation formulation based on Quantum

Rice-Ramsperger-Kassel theory

k2, theoretical(1.6 atm) = 2.88 × 1041 T-8 exp(-43,400/T) s-1

recovers an activation energy similar to the current measurements, but its A-factor is

approximately three times lower than the measured A-factor from the current study, as

shown in Fig. 4.4.

41

Table 4.1: Summary of acetone unimolecular dissociation rate constant data.

T5 [K] P5 [atm] k2 [s-1] Initial Mixture: 0.25% Acetone / Ar (from CO)

1273 1.65 3.61E+02 1351 1.62 1.61E+03 1393 1.55 3.52E+03 1469 1.48 1.29E+04 1545 1.42 3.76E+04 1166 0.73 1.70E+01 1226 0.71 8.90E+01 1324 0.63 8.22E+02 1404 0.62 3.31E+03 1458 0.59 7.82E+03 1567 0.54 4.14E+04 1260 5.28 2.72E+02 1350 4.98 2.30E+03

Initial Mixture: 1% Acetone / Ar (from CO)

1213 1.66 8.53E+01 1332 1.59 1.01E+03 1423 1.53 5.42E+03 1507 1.47 2.06E+04

Initial Mixture: 1% Acetone / Ar (from CH3COCH3)

1340 1.41 1.19E+03 1352 1.24 1.27E+03 1474 1.23 9.76E+03 1561 1.26 3.51E+04 1650 1.28 1.17E+05

Other species (CH3, C2H4, and CH4) were also measured during pyrolysis, and

comparisons with the simulations demonstrate how the current evaluated k2 significantly

improves the overall performance of the detailed mechanism. Fig. 4.7 displays the

measured CH3 time histories during acetone pyrolysis, along with the computed values

from the original Pichon et al. mechanism and the modified mechanism (with our revised

value for k2). The computed CH3 peak values from the original Pichon et al. mechanism

are approximately half of the measured values. Additionally, the CH3 sensitivity analysis

reveals that CH3 time histories are strongly sensitive to two reactions: CH3COCH3 (+ M)

42

→ CH3 + CH3CO (+ M) (reaction (2)) and the relatively well-established CH3 + CH3 (+

M) → C2H6 (+ M) [80-81]. Consequently, using the revised (higher) values for k2 brings

the simulations from the Pichon et al. mechanism into closer agreement with the

measured CH3 time histories. Note that the uncertainty in the CH3 concentration during

the first 100 µs was approximately ±20%, which was mainly attributed to the

uncertainties in the absorption coefficient and the interference subtraction scheme. The

CH3 plateau levels (after 500 µs) had a much larger uncertainty of ±30%, owing to larger

interference absorption (from other intermediate products, such as C2H4).

Figure 4.7: CH3 time histories for 0.25% acetone in Ar: measured and calculated values.

Ethylene is an important major product formed during acetone pyrolysis. Fig. 4.8

shows the measured C2H4 time histories, along with the simulations from the original and

modified Pichon et al. mechanisms. The original Pichon et al. mechanism clearly fails to

predict the formation rates and the ultimate yields of ethylene. During acetone pyrolysis,

C2H4 is mainly formed from the direct competition between two reactions: CH3 + CH3 (+

M) → C2H6 (+ M) and CH3 + CH3 → C2H5 + H, immediately followed by C2H5 (+ M) →

C2H4 + H (+ M). The higher formation rates of CH3 in the modified mechanism (via the

higher values for k2) significantly improve the simulations at early times.

43

Figure 4.8: C2H4 time histories for 1% acetone in Ar: measured and calculated values.

Methane is another major product formed during acetone pyrolysis. Fig. 4.9

shows the measured CH4 time histories, along with the simulations from the modified

Pichon et al. mechanisms (i.e., with the revised values for k2) with two different

CH3COCH3 + CH3 reaction rate constants adopted by Pichon et al. [89] and Saxena et al.

[88]. The original Pichon et al. mechanism significantly underpredicted the CH4 time

histories by at least a factor of 4, and the modified Pichon et al. mechanism v1 (with

revised k2 and original acetone + CH3 rate constant used by Pichon et al.) was still not

able to capture the CH4 formation rates at T < 1400 K. However, at T = 1556 K, the

modified mechanism v1 predicted the CH4 concentration reasonably well. Additionally,

the CH4 sensitivity analysis (not presented here) shows the importance of reaction (2),

and the reactions of CH3 plus CH3, C2H6, and CH3COCH3. There are only limited

uncertainties in the CH3 + CH3 and the CH3 + C2H6 reaction rate constants as given by

Baulch et al. [96] (±20%). Thus, the rate constant for the reaction of CH3COCH3 + CH3

→ CH3COCH2 + CH4 in the original Pichon et al. mechanism

kPichon et al. = 7.96 × 1011 exp(-9741 [cal/mol]/RT) cm3 mol-1 s-1

is likely to be too slow at T < 1400 K. In particular, the Saxena et al. mechanism uses a

significantly different value for the rate constant of CH3COCH3 + CH3 (see Fig. 4.10),

which is:

44

kSaxena et al. = 0.550 T4 exp(-8290 [cal/mol]/RT) cm3 mol-1 s-1

The rate constant used by Saxena et al. [88] is at least three times faster than the value of

the rate constant adopted by Pichon et al. [89] at 1250 K, and is approximately six times

faster at 1538 K. When the rate constant from Saxena et al. is used in the modified

Pichon et al. mechanism (labeled as modified mechanism v2), the simulated time

histories show excellent agreement with the measured time histories at T < 1400 K.

However, the modified mechanism v2 seems to overpredict the ultimate methane yield at

T = 1556 K. This reveals that the activation energy of the acetone + CH3 reaction likely

still requires some fine adjustment.

Figure 4.9: CH4 time histories for 1.5% acetone in Ar: measured and calculated values.

45

Figure 4.10: Arrhenius plot of the rate constants for the reaction of CH3COCH3 + CH3 → CH3COCH2 + CH4 from Saxena et al. [88], Sato and Hidaka [87], and Pichon et al. [89].

4.3.2 2-Butanone Pyrolysis

High-temperature 2-butanone pyrolysis was studied using multi-species time

history measurements over the temperature range of 1100-1500 K at pressures around 1.5

atm. Fig. 4.11 shows the measured 2-butanone time histories during 2-butanone

pyrolysis, along with the simulations from the original Serinyel et al. mechanism [95].

Serinyel et al. postulated three initial 2-butanone decomposition pathways:

2-Butanone (+ M) → C2H5 + CH3CO (+ M) (3a)

2-Butanone (+ M) → CH3 + C2H5CO (+ M) (3b)

2-Butanone (+ M) → CH3 + CH3COCH2 (+ M) (3c)

As illustrated in Fig. 4.12, the 2-butanone sensitivity analysis reveals that the 2-

butanone time history is predominantly sensitive to its three initial decomposition

pathways, with channel (3a) as the primary decomposition channel. In addition, there is

some minor interference from the reactions of 2-butanone + H → CH3CHCOCH3 + H2

and CH3CHCOCH3 → CH3CHCO + CH3. To determine the overall initial 2-butanone

decomposition rates, the measured 2-butanone time histories were best-fit with the

simulated profiles by adjusting the rate constants for channels (3a)-(3c), without

46

modifying their branching ratios. The best-fit simulated 2-butanone time histories are

also shown on Fig. 4.11. At all temperatures, the modified Serinyel et al. mechanism

captures the initial 2-butanone decomposition rates very accurately, at least for the first

500 µs. (Note that some simulations are nearly indistinguishable from the data traces.)

Figure 4.11: 2-Butanone time histories for 1% 2-butanone in Ar: measured and simulated values.

Figure 4.12: 2-Butanone sensitivity for 1% 2-butanone in Ar.

The overall 2-butanone decomposition rate constant measurements (k3 = k3a + k3b

+ k3c) are plotted on Fig. 4.13, along with the estimated values from the Serinyel et al.

mechanism. Using a least-squares fit, the overall 2-butanone decomposition rate constant

47

was found to be k3 = 6.08×1013 exp(-63.1 [kcal/mol]/RT) s-1 over the temperature range

of 1119-1412 K at pressures around 1.5 atm (see the dashed line in Fig. 4.13). (Note that

the current data are the first direct high-temperature rate constant measurements for the

initial 2-butanone decomposition.) The major contributions to uncertainties in the rate

constants were: temperature (±10%), 2-butanone absorption cross-section (±5%), fitting

the data to computed profiles (±5%), and uncertainties resulting from secondary reactions

(±15%), giving an overall uncertainty in k3 of ±35%. In addition, Table 4.2 summarizes

the overall 2-butanone decomposition rate constant measurements over 1119-1412 K at

pressures near 1.5 atm. The effect of the branching ratios of the initial 2-butanone

decomposition pathways on the measured rate constant was investigated by perturbing

the branching ratio of channel (3a) from 0.70 to 0.50 at 1361 K, and no significant

difference on the simulated 2-butanone time histories was found. In the following

discussion, the original branching ratios of the 2-butanone decomposition pathways used

by Serinyel et al. were retained. Our measured value for k3 is approximately 30% faster

than Serinyel et al. at 1119 K, and is approximately 100% faster at 1412 K. Our inferred

rate constants for channels (3a)-(3c) at 1.5 atm can be expressed as follows:

k3a = 1.31×1013 exp(-59.9 [kcal/mol]/RT) s-1

k3b = 8.47×1014 exp(-75.3 [kcal/mol]/RT) s-1

k3c = 3.49×1014 exp(-72.9 [kcal/mol]/RT) s-1

Figure 4.13: Arrhenius plot for overall 2-butanone decomposition rate constant (k3).

48

Table 4.2: Summary of overall 2-butanone decomposition rate constant data.

T5 [K] P5 [atm] k3 [s-1] Initial Mixture: 1% 2-Butanone / Ar

1119 1.62 3.00E+01 1170 1.61 1.00E+02 1214 1.53 2.60E+02 1252 1.48 5.45E+02 1299 1.47 1.42E+03 1361 1.48 4.11E+03 1412 1.39 1.17E+04

Fig. 4.14 shows the measured CH3 time histories during 2-butanone pyrolysis,

along with the simulations from the original and modified Serinyel et al. mechanisms.

The modified Serinyel et al. mechanism captures the initial CH3 formation rates more

closely, but the simulated CH3 peak values are still underpredicted. According to the

CH3 sensitivity analysis, the CH3 time histories are mainly sensitive to channels (3a)-(3c)

and the relatively well-established reaction of CH3 + CH3 (+ M) → C 2H6 (+ M) [80-81].

Thus, discrepancies between the measured and simulated CH3 peak values may be

attributed, at least partially, to uncertainties in the relative branching ratios of the initial

2-butanone decomposition pathways. (Note that the agreement here is not sacrificed by

the subsequent addition of a methyl ketene decomposition channel, which will be

discussed in the following section.)

Figure 4.14: CH3 time histories for 0.25% 2-butanone in Ar: measured and calculated values.

49

Similar to acetone pyrolysis, CO is another important stable species formed

during 2-butanone pyrolysis. Fig. 4.15 shows the CO rate of production (ROP) analysis

for the mixture of 1% 2-butanone in Ar at 1292 K and 1.58 atm simulated using the

Serinyel et al. mechanism with the revised k3, and the ROP analysis reveals that CO is

mainly generated via two reaction pathways over the time frame of the experiment, which

are CH3CO (+ M) → CH3 + CO (+ M) and C2H5CO → C2H5 + CO. The acetyl (CH3CO)

radical can be formed via 2-butanone (+ M) → C 2H5 + CH3CO (+ M) (channel (3a)) and

2-butanone + H → CH 2CH2COCH3 + H2, followed by the fuel radical decomposition

(CH2CH2COCH3 → C2H4 + CH3CO). Additionally, the propionyl (C2H5CO) radical can

be formed via 2-butanone (+ M) → CH3 + C2H5CO (+ M) (channel (3b)).

Figure 4.15: CO rate of production (ROP) plot for 1% 2-butanone in Ar using the original Serinyel et al. mechanism (with the revised k3). Fig. 4.16 shows the measured and simulated CO time histories during 2-butanone

pyrolysis. Based on the measured 2-butanone and CO time histories at T ≈ 1292 K, the

measurements suggest ~87% conversion to CO from 2-butanone, while the original

Serinyel et al. mechanism [95] predicts only about 57% conversion to CO. The model

with the revised overall 2-butanone decomposition rate constant (k3) generates slightly

higher CO concentration (~67%), but still well below that measured. According to the

simulations, the major species containing O atoms are 2-butanone, CO, and CH3CHCO

50

(methyl ketene), and the model seems to predict significant amounts of methyl ketene

formed during 2-butanone pyrolysis. This implies that some of the CO formation

pathways might be incomplete in the model, particularly the methyl ketene sub-

mechanism. Methyl ketene is produced through the reaction of 2-butanone + H →

CH3CHCOCH3 + H2, followed by CH3CHCOCH3 → CH3CHCO + CH3. According to

the model, the removal pathway of methyl ketene is only through the H-abstraction

reaction from methyl ketene (CH3CHCO + H → C2H5 + CO). Since methyl ketene is not

a stable species, it should undergo a unimolecular decomposition process, which is not

included in the original Serinyel et al. mechanism [95]. In the present analysis, a methyl

ketene decomposition pathway (CH3CHCO (+ M) → C2H4 + CO (+ M)) was

incorporated, and the corresponding rate constant was assumed to be the same as the

value for ketene decomposition (CH2CO (+ M) → CH2 + CO (+ M)). The modified

mechanism (with revised k3 and added methyl ketene decomposition reaction) simulates

the CO concentrations rather accurately, as illustrated in Fig. 4.16.

Figure 4.16: CO time histories for 1% 2-butanone in Ar: measured and calculated values.

Fig. 4.17 displays the measured C2H4 time histories, along with the simulations

from the original and modified Serinyel et al. mechanisms. The measurement suggests

the ethylene yield (defined as the ratio of the long-time C2H4 concentration to the initial

51

2-butanone concentration) to be ~0.88 at 1412 K, while the original Serinyel et al.

mechanism predicts the yield to be ~0.73 and the modified mechanism predicts the yield

of ~0.87. In general, the modified mechanism is able to accurately simulate the ultimate

yields of C2H4. (Note that the simulated C2H4 time history from the modified mechanism

lies exactly on top of the measured time history at 1252 K.) The C2H4 sensitivity analysis

shows the significance of the initial 2-butanone decomposition pathways (channels (3a)-

(3c)) and the H-abstraction reactions from 2-butanone. C2H4 is initially formed through

the following processes:

(i) 2-Butanone (+ M) → C2H5 + CH3CO (+ M), followed by C2H5 (+ M) → C2H4 +

H (+ M);

(ii) 2-Butanone + H → CH2CH2COCH3 + H2, followed by CH2CH2COCH3 → C2H4

+ CH3CO.

Hence, fine refinement on the branching ratios of the initial 2-butanone decomposition

pathways and the H-abstraction reaction rate constants appears needed to perfectly match

the initial formation rates of C2H4.

Figure 4.17: C2H4 time histories for 1% 2-butanone in Ar: measured and calculated values.

In addition to CO and C2H4, methane is another major product formed during 2-

butanone pyrolysis. A plot of the measured methane time histories, along with the

52

computed values from the modified Serinyel et al. mechanism, is illustrated in Fig. 4.18.

As compared to the measurements, the modified mechanism underpredicts the methane

concentrations by at least a factor of 2 at temperatures less than 1400 K. However, at

temperatures higher than 1400 K, the modified mechanism is able to capture the methane

formations reasonably well. Methane is mainly formed through the H-abstraction

reactions from 2-butanone by CH3 radicals, which are:

2-Butanone + CH3 → CH2CH2COCH3 + CH4

2-Butanone + CH3 → CH3CHCOCH3 + CH4

2-Butanone + CH3 → C2H5COCH2 + CH4

Such huge differences between the measurements and simulations may be caused by the

inaccurate model predictions for CH3 concentrations and the uncertainties in the rate

constants for 2-butanone + CH3 reactions, particularly their activation energies. Further

increases in the rate constants for 2-butanone + CH3 reactions at temperatures less than

1400 K appear needed in order to improve the model predictions on methane.

Figure 4.18: CH4 time histories for 1.5% 2-butanone in Ar: measured and calculated values.

53

4.4 Summary

High-temperature acetone and 2-butanone pyrolysis was investigated behind

reflected shock waves using multi-species time history measurements (acetone/2-

butanone, CO, CH3, C2H4, and CH4). Direct determinations of the acetone dissociation

rate constant (k2) and the overall 2-butanone dissociation rate constant (k3 = k3a + k3b +

k3c) were made by taking advantage of the measured species time histories for CO (and

acetone) and 2-butanone, respectively.

In the 2-butanone pyrolysis system, an analysis of the O-atom balance based on

the simulated and measured 2-butanone and CO time history measurements revealed

pooling of methyl ketene in the simulations. The addition of the methyl ketene

decomposition pathway to remove this pooling significantly improved the mechanism’s

performance. Further improvement in the 2-butanone mechanism will require a better

understanding and refinement of the branching ratios of the initial 2-butanone

decomposition pathways and the rate constants for the H-abstraction reactions from 2-

butanone.

55

Chapter 5 Shock Tube Measurements of 3-Pentanone Pyrolysis and Oxidation

5.1 Introduction

In contrast to acetone and 2-butanone, very little experimental data are available

for high-temperature 3-pentanone combustion studies. Three studies of this fuel are of

note. Davidson et al. [97] measured shock tube ignition delay times for a series of

oxygenated fuels, including 3-pentanone, over temperatures of 1150-1550 K and a

pressure of ~1.8 atm. From their experiments, they concluded that 3-pentanone has much

faster ignition delay times than found for acetone, n-pentane, methyl butanoate, and

butanal. Similarly, Serinyel et al. [98] performed shock tube ignition delay time

measurements in the temperature range 1250-1850 K, pressures near 1 atm, and

equivalence ratios of 0.5-2.0 for mixtures of 0.875-1.31% 3-pentanone in O2/argon. They

also conducted laminar flame speed measurements in a spherical bomb for mixtures of 3-

pentanone in air with various equivalence ratios at an initial temperature of ~305 K and

an initial pressure of 1 atm. Through their flame speed measurements, they concluded

that 3-pentanone has higher reactivity than acetone and 2-butanone. Finally, Hong et al.

[99] examined the influence of oxygenates (such as 3-pentanone) on soot formation

during fuel rich n-heptane oxidation at temperatures of 1600-1900 K and pressures of 20-

30 atm. A significant reduction in the overall soot yield was discovered with the addition

of small quantities of oxygenates.

56

In this chapter, we present high-temperature pyrolysis and oxidation studies of 3-

pentanone behind reflected shock waves using laser absorption methods to measure time

histories of six species: 3-pentanone, CH3, CO, C2H4, H2O, and OH. In addition, 3-

pentanone oxidation behavior was compared with the oxidation behavior of two other

ketones, acetone and 2-pentanone, by examining their ignition delay times and OH time

histories.


5.2.1 Mixture Preparation


heated uniformly to 50 oC and mixed with a magnetically driven stirring vane for at least

2 hours prior to the experiments. Research grade (99.999%) gases (from Praxair) and

ReagentPlus® grade (≥99%) 3-pentanone (from Sigma-Aldrich), which was further

treated using a freeze-pump-thaw procedure, were used in mixture preparation. The

mixture compositions from this study are summarized in Table 5.1, along with the

measured species for the corresponding mixtures. For the lower fuel concentration

mixtures (<0.25% 3-pentanone), a double-dilution method was used in mixture

preparation to allow for more accurate mixture compositions.

Table 5.1: Summary of test gas mixture compositions and measured species.

Mix #

Gas Compositions Species Time histories τign

3-Pent. O2 Ar 3-Pent. CH3 CO C2H4 OH H2O A 1.00% -- 99.00% x x x B 0.25% -- 99.75% x x C 0.10% -- 99.90% x D 0.040% 0.280% 99.68% x x x E 0.040% 0.560% 99.40% x x x F 0.075% 0.525% 99.40% x x x G 0.571% 4.00% 95.43% x H 0.286% 4.00% 95.71% x I 0.875% 12.25% 86.88% x

57

5.2.2 Species Absorption Coefficient Evaluations

Because of the endothermic nature of the pyrolysis reaction, there is a temperature

drop in the reacting test gas mixture during the experiment, which increases with the

initial 3-pentanone mole fraction. This change in temperature can perturb (generally

increase) the absorption coefficients of individual species, and hence perturb the

conversion of measured absorbance to mole fraction. More accurate species mole

fraction time histories are obtained by accounting for this effect rather than assuming a

constant coefficient evaluated at the initial temperature. To determine these approximate

time-varying absorption coefficients, the temperature and pressure profiles were

calculated using the Serinyel et al. mechanism of NUI Galway [98] under either constant

energy (U) and volume (V) constraints or constant enthalpy (H) and pressure (P)

constraints (using CHEMKIN PRO [71]). The species mole fraction time histories were

then inferred from the measured absorption data using known values of the absorption

coefficients evaluated at the simulated T and P.

The 3-pentanone mole fraction time history for a 1% 3-pentanone/Ar mixture and

an initial temperature of 1248 K and an initial pressure of 1.58 atm was calculated from

the measured absorption data using three different approaches: (1) constant absorption

coefficient for Beer’s law (as has been common in the past), (2) T- and P-dependent

absorption coefficients based on constant U, V calculation, and (3) T- and P-dependent

absorption coefficients based on constant H, P calculation. The 3-pentanone mole

fraction time histories from these three approaches are nearly indistinguishable. Hence,

the measured 3-pentanone mole fraction time histories are effectively insensitive to any

small variation in temperature and pressure change that is a result of the gasdynamic

model used. In this chapter, constant absorption coefficients for Beer’s law are thus

employed for 3-pentanone time history measurements.

On the other hand, the absorption coefficients of CO and C2H4 can increase by up

to 10% and 7%, respectively, during the pyrolysis of 1% 3-pentanone in Ar. Clearly, to

minimize the effects of temperature drop during pyrolysis, lower fuel concentration

mixtures (0.1% or 0.25% 3-pentanone in Ar) are preferred. The drop in temperature for

58

the mixture of 0.25% 3-pentanone in Ar is ~30 K (approximately 4 times less than 1% 3-

pentanone mixture) at an initial temperature of 1325 K and an initial pressure of 1.60

atm. As shown in Fig. 5.1, the CO mole fraction time history for the mixture of 0.25% 3-

pentanone in Ar at 1325 K and 1.60 atm was calculated using absorption coefficients

based on three different gas dynamic models. The initial CO formation rates (for the first

400 µs) from these three approaches are very nearly identical, but the ultimate yield from

the constant absorption coefficient approach (labeled as method 1 and generally

employed in most past studies) is ~2% higher than the yields from the T- and P-

dependent absorption coefficient approaches that are based on constant U, V (method 2)

and constant H, P (method 3) calculations. To correct for this slight difference at later

times, all the CO mole fraction time histories in the present work were calculated using

the T- and P-dependent absorption coefficients based on constant U, V calculations; the

estimated uncertainty in the measured yields using this approach is ±2-3%. Similarly, all

the C2H4 time histories were also calculated from the measured absorption data using the

T- and P-dependent absorption coefficients based on constant U, V calculations.

Figure 5.1: Comparison of CO mole fraction time histories at 1325 K and 1.60 atm with different absorption coefficients in Beer’s law.

During 3-pentanone oxidation, very dilute mixtures (e.g., 400 ppm 3-pentanone)

are used for OH and H2O species time history measurements in order to minimize the

59

rapid energy release at the time of ignition, which increases the temperature (and reduces

the absorption coefficient). As illustrated in Fig. 5.2, for the mixture of 400 ppm 3-

pentanone with 0.28% O2 in Ar, the early-time features of OH obtained from the

measured absorption data using three different absorption coefficient approaches are

effectively identical. Similar results can be observed from the H2O time history profiles.

This is particularly important because these early-time features are unique to individual

fuels, as will be discussed in the later section. However, the final plateau levels of OH

and H2O from method 1 are lower than the levels from methods 2 and 3 by 4% and 2%,

respectively. This is mainly due to the fact that there is a temperature rise of ~50 K at the

time of ignition, and the use of a constant absorption coefficient for Beer’s law is not

strictly valid at these later times. Interestingly, the final plateau levels of OH and H2O

obtained from methods 2 and 3 are indistinguishable. Hence, the temperature corrections

on the measured species are independent of the specific gasdynamic model (const. U, V

or const. H, P). Similar to CO and C2H4 measurements (as described above), all the

measured OH and H2O time histories are corrected using the T- and P-dependent

absorption coefficients based on constant U, V calculations.

Figure 5.2: Comparison of OH mole fraction time histories at 1486 K and 1.52 atm with different absorption coefficients in Beer’s law. The OH mole fractions by constant U, V and constant H, P are virtually indistinguishable for OH.

60

5.3 Results and Discussion

5.3.1 3-Pentanone Pyrolysis

A high-temperature 3-pentanone pyrolysis study was performed behind reflected

shock waves using four species time history measurements (fuel, CH3, CO, and C2H4)

over 1070-1530 K at a pressure of ~1.6 atm. The test mixtures were 0.1% to 1% 3-

pentanone in balance argon. In the present study, the CHEMKIN PRO package [71] was

used to simulate all species time histories under the standard constant energy and volume

assumption (constant U, V), and the Serinyel et al. mechanism of NUI Galway [98] was

chosen as the base mechanism. To the best of our knowledge, the Serinyel et al.

mechanism is the only available detailed mechanism in the literature that is suitable for

high temperature 3-pentanone combustion. The sub-mechanism of 3-pentanone was

developed by Serinyel et al. and implemented into the well-established C4 mechanism of

NUI Galway [100]. In particular, the rate constants for 3-pentanone unimolecular

decomposition reactions were estimated in the reverse direction, and their high-pressure

limit values were further assumed and treated using Quantum Rice-Ramsperger Kassel

(QRRK) theory with a master equation analysis to include the pressure fall-off effects. In

addition, the detailed mechanism of Serinyel et al. was then validated against their

ignition delay time and laminar flame speed measurements.

Fig. 5.3 shows the measured 3-pentanone time histories during pyrolysis of 1% 3-

pentanone in argon, along with the simulations from the Serinyel et al. mechanism of

NUI Galway [98]. The measured fuel time histories are inconsistent with the simulated

profiles from the model, and the model significantly underpredicts the fuel removal rates

at current experimental conditions. As illustrated in Fig. 5.4, the 3-pentanone sensitivity

analysis was performed to determine which reactions are pertinent to 3-pentanone time

histories. As expected, 3-pentanone time histories are primarily sensitive to the initial

fuel decomposition pathways:



61

In addition, there is some minor interference from the reactions of C2H4 + H (+ M) →

C2H5 (+ M), CH3 + CH3 → C2H5 + H, and the H-atom abstraction reactions from 3-

pentanone by H radicals. The branching fraction of the initial fuel decomposition

through reaction (4a) ranges from 0.59 to 0.53 over 1070-1530 K at 1.6 atm [98]. This

indicates that 3-pentanone undergoes unimolecular decomposition through these two

reaction pathways at similar rates.

Figure 5.3: Measured and simulated 3-pentanone time histories for 1% 3-pentanone in Ar. Simulations used the Serinyel et al. mechanism.

Figure 5.4: 3-pentanone sensitivity analysis for 1% 3-pentanone in Ar at 1323 K and 1.6 atm.

62

Methyl radical (CH3) is an important transient species during 3-pentanone

pyrolysis. CH3 radicals are first formed through reaction (4b) of the initial fuel

decomposition pathways, and hence the initial CH3 formation rates are a good measure of

the initial fuel decomposition rates. Fig. 5.5 shows the measured CH3 time histories

during pyrolysis of 0.1% 3-pentanone in argon, along with the computed profiles from

the Serinyel et al. mechanism. It should be noted that the uncertainty of the CH3

concentration during the first 100 µs was approximately ±10%, which was mainly

contributed from the uncertainties in the absorption coefficient and the interference

subtraction scheme. At later times, the CH3 plateau levels had a slightly larger

uncertainty of ±20% due to larger interference absorption (primarily from C2H4). Similar

to the fuel time histories, the model fails to capture the initial CH3 formation rates, and

the predicted CH3 peak values are approximately 30% lower than the measured values.

However, the model does simulate the CH3 removal rates reasonably well from which we

can infer that the reaction rate constants for these CH3 removal channels are reasonable.

Figure 5.5: CH3 time histories for 0.1% 3-pentanone in Ar. Simulations were done using the Serinyel et al. mechanism.

As shown in Fig. 5.6, the CH3 sensitivity analysis reveals that the CH3 time

histories are mainly sensitive to the initial fuel decomposition pathways (reactions (4a)

63

and (4b)) at early times. At later times, there is some interference from the secondary

reactions, which are described as follows:

C2H4 + H (+ M) → C2H5 (+ M) (20)

CH3 + CH3 → C2H5 + H (21)

C2H6 (+ M) → CH3 + CH3 (+ M) (18)

Of note is that the rate constants for reactions (18), (20), and (21) are relatively

well-established. In particular, the reverse of reaction (18) is a primary CH3 removal

channel at high temperatures. In the present analysis, we updated the rate constant for

reaction (18) with the values measured by Oehlschlaeger et al. [80], whose measured

values are consistent with another recent study from Kiefer et al. [81]. The rate constants

for reactions (18), (20), and (21) (and the H-atom abstraction reactions from 3-pentanone

by H radicals) are also provided in Table 5.2. Additionally, the most uncertain rate

constants among reactions (4a), (4b), (18), (20), and (21) are the initial fuel

decomposition pathways, reactions (4a) and (4b). Hence, the measured 3-pentanone and

CH3 time histories can be used to infer the overall fuel decomposition rate constant (k4 =

k4a + k4b) at the measured pressure.

Figure 5.6: CH3 sensitivity analysis for 0.1% 3-pentanone in Ar at 1433 K and 1.6 atm.

64

Table 5.2: Kinetic parameters employed in the Serinyel et al. mechanism.

Rate Constant

Reaction A [†] b E [cal/mol] No. Reference

C2H4 + H (+ M) → C2H5 (+ M) 1.081E+12 0.45 1.822E+03 20 [98] Low-Pressure Limit: 1.200E+42 -7.62 6.970E+03 Troe centering: 0.975 210 984 4374

CH3 + CH3 → C2H5 + H 4.990E+12 0.10 1.060E+04 21 [98] C2H6 (+ M) → CH3 + CH3 (+ M) 1.880E+50 -9.72 1.073E+05 18 [80]

Low-Pressure Limit: 3.720E+65 -13.14 1.015E+05 Troe centering: 0.39 100 1900 6000

C2H5COC2H5 + H → C2H5COC2H4p + H2

1.332E+06 2.54 6.756E+03 22a [98]

C2H5COC2H5 + H → C2H5COC2H4s + H2

1.913E+06 2.29 2.875E+03 22b [98]

C2H5COC2H4p → C2H5CO + C2H4 7.170E+13 -1.93 2.626E+04 23 [98]

C2H5COC2H4s → CH3CHCO + C2H5 1.210E+19 0.42 4.272E+04 24 [98] CH3CHCO + H → C2H5 + CO 4.400E+12 0 1.459E+03 25 [98]

C2H4 + CO (+ M) → CH3CHCO (+ M) 8.100E+11 0 0 26 this study

Low-Pressure Limit: 2.690E+33 -5.11 7.095E+03 Troe centering: 5.907E-01 275 1226 5185

H + O2 → OH + O 3.547E+15 -0.41 1.660E+04 27 [98] C2H4 + OH → C2H3 + H2O 1.800E+06 2.00 2.500E+03 28 [98]

† Units of A are in s-1 for unimolecular reactions, cm3 mol-1 s-1 for bimolecular reactions, and cm6 mol-2 s-1 for termolecular reactions.

Fig. 5.7 shows the measured 3-pentanone and CH3 time histories, along with their

best-fit profiles simulated from the Serinyel et al. mechanism by revising the overall 3-

pentanone decomposition rates. As a result, the measured 3-pentanone time histories

provided the values for k4 over 1070-1330 K at 1.6 atm, and the measured CH3 time

histories provided the values for k4 over 1230-1530 K at 1.6 atm, as are shown in Fig. 5.8.

A best-fit expression for the overall 3-pentanone decomposition rate constant

measurements can be given as k4 = 4.383×1049 T-10 exp(-44,780/T) s-1 over 1070-1530 K

(see the dashed line in Fig. 5.8). (In addition, the measured values can be expressed in

Arrhenius form as k4 = 1.501×1015 exp(-34,640/T) s-1 over 1070-1330 K.) The major

contributions to the uncertainties in the overall rate constant were: temperature (±10%),

65

3-pentanone absorption coefficient (±5%) (or CH3 absorption coefficient and interference

subtraction scheme (±10%)), fitting the data to computed profiles (±5%), and

uncertainties resulting from secondary reactions (±15%), giving an overall uncertainty in

k4 of ±35% over 1070-1330 K from the measured 3-pentanone time histories (or ±40%

over 1230-1530 K from the measured CH3 time histories). The values for k4 are also

summarized in Table 5.3. The influence of the branching ratio (k4a/k4) of reaction (4a) on

the overall fuel decomposition rate constant was examined by perturbing the branching

ratio from 0.4 to 0.7 and keeping k4 constant, and the changes in the computed profiles

were negligible. Hence, the measured 3-pentanone and CH3 time histories are insensitive

to the branching ratios of the initial fuel decomposition pathways. In the present analysis,

the original branching ratios postulated from Serinyel et al. are utilized.

As illustrated in Fig. 5.8, the measured overall 3-pentanone decomposition rate

constant is approximately 3.5 times those of Serinyel et al. over 1070-1330 K at 1.6 atm,

and as such, the high-pressure limit rate constants for reactions (4a) and (4b) as given by

Serinyel et al. are in need of revision. Note that the measured and predicted overall 3-

pentanone decomposition rate constants both experience severe non-Arrhenius curvature,

which explains the large values inferred for the A-factor and pre-exponential temperature

dependence.

Table 5.3: Summary of overall 3-pentanone decomposition rate constant data.

T5 [K] P5 [atm] k4 [s-1] Initial Mixture: 0.1% 3-Pentanone / Ar 1237 1.75 1.29E+03 1346 1.67 8.97E+03 1433 1.60 2.86E+04 1524 1.52 9.08E+04

Initial Mixture: 1% 3-Pentanone / Ar 1071 1.70 1.39E+01 1113 1.64 4.24E+01 1164 1.64 1.81E+02 1248 1.58 1.41E+03 1323 1.32 6.10E+03

66

Figure 5.7: (a) Best-fit 3-pentanone time histories and (b) best-fit CH3 time histories using the Serinyel et al. mechanism with revised overall 3-pentanone decomposition rate constant (k4).

67

Figure 5.8: Arrhenius plot for the overall 3-pentanone decomposition rate constant (k4) at 1.6 atm.

Fig. 5.9 shows the measured 3-pentanone and CO time histories during pyrolysis

of 1% 3-pentanone in argon at 1248 K and 1.6 atm, along with the simulations from the

Serinyel et al. mechanism with the revised k4. Note that initially, all O atoms (100

percent) are present in 3-pentanone. At 750 µs, approximately 35% of the total O atoms

remains in 3-pentanone, with about 55% of the O atoms in CO; together the O atoms

from these species add up to ~90% of the total available O atoms. Similarly, at 1500 µs,

there are approximately 23% of the O atoms in 3-pentanone and 67% of the O atoms in

CO, and these O atoms also sum up to ~90%. Therefore, the measurements suggest

approximately 90% conversion of 3-pentanone to CO. However, the model with the

revised k4 only predicts ~57% conversion of 3-pentanone to CO.

68

Figure 5.9: Measured 3-pentanone and CO time histories during 3-pentanone pyrolysis at 1248 K and 1.6 atm.

Based on the CO sensitivity analysis (see Fig. 5.10), the CO time histories, and

particularly the initial formation rates, are mainly sensitive to the initial 3-pentanone

decomposition pathways (reactions (4a) and (4b)), the ethyl radical decomposition

(reaction (20)), and the H-atom abstraction reactions from 3-pentanone, which are:

C2H5COC2H5 + H → C2H5COC2H4p + H2 (22a)

C2H5COC2H5 + H → C2H5COC2H4s + H2 (22b)

Figure 5.10: CO sensitivity analysis for 1% 3-pentanone in Ar at 1248 K and 1.6 atm.

69

Abstraction of hydrogen atom from the fuel molecule yields the formation of fuel

radicals C2H5COC2H4p and C2H5COC2H4s, where p and s denote the primary and

secondary sites, respectively. However, these reactions do not seem to significantly

affect the final CO plateau values, and only minor changes in the computed CO plateaus

are found if the rate constants for these reactions are increased by a factor of 3. One

possible explanation for the discrepancy between the measurements and simulations is

the incomplete CO formation pathways in the model. Based on the Serinyel et al.

mechanism with the revised k4, the major species that contain O-atoms are 3-pentanone,

CO, and methyl ketene (CH3CHCO). Hence, we infer that the model overpredicts the

concentration of methyl ketene. Methyl ketene is formed through the H-atom abstraction

reaction (C2H5COC2H5 + H → C 2H5COC2H4s + H2), immediately followed by the

decomposition of the fuel radical.

C2H5COC2H4s → CH3CHCO + C2H5 (24)

Methyl ketene is then removed via one pathway only in the original Serinyel et al.

mechanism [98], as described in the following:

CH3CHCO + H → C2H5 + CO (25)

As discussed in Chapter 4, methyl ketene should also undergo thermal

decomposition to form other stable species, as suggested in one previous study [101].

Here we also incorporate the methyl ketene unimolecular decomposition pathway into the

Serinyel et al. mechanism:

CH3CHCO (+ M) → C2H4 + CO (+ M) (-26)

The rate constant for methyl ketene unimolecular decomposition was assumed to have the

same value as the rate constant for ketene unimolecular decomposition (CH2CO (+ M) →

CH2 + CO (+ M)). With this modification, the model can now predict higher CO

concentrations and lower methyl ketene concentrations. In addition, based on the model,

the remaining O-atom containing species is primarily ketene and an accurate knowledge

of the ketene sub-mechanism becomes particularly important in predicting the CO

plateau levels during 3-pentanone pyrolysis.

Fig. 5.11 shows the measured CO time histories during 0.25% 3-pentanone in Ar,

along with the simulations from the (a) original and (b) modified Serinyel et al.

70

mechanisms. As expected, the original Serinyel et al. mechanism underpredicts the CO

time histories by at least 30% at current experimental conditions. With the revised

overall 3-pentanone decomposition rate constant and the addition of the methyl ketene

decomposition reaction, the computed profiles from the modified mechanism show much

better agreement with the measurements at all temperatures. Similar agreement between

the measurements and the simulations from the modified mechanism can also be found

for the mixture of 1% 3-pentanone in Ar. As mentioned previously, the initial CO

formation rates are sensitive to the H-atom abstraction reactions (reactions (22a) and

(22b)). Hence, the current study supports the use of the Serinyel et al. values for these

reaction rate constants.

Figure 5.11: CO time histories for 0.25% 3-pentanone in Ar: measured and calculated values from the (a) original and (b) modified Serinyel et al. mechanisms.

71

During 3-pentanone pyrolysis, ethylene is mainly formed through either

unimolecular decomposition pathways or H-atom abstraction reactions, and these H-atom

abstraction reactions are particularly important at later times when H atoms are abundant

in the system. The primary H-atom abstraction reactions are C2H5COC2H5 + H →

C2H5COC2H4p + H2 and C2H5COC2H5 + H → C 2H5COC2H4s + H2. These fuel radicals

(C2H5COC2H4p and C2H5COC2H4s) can then form ethylene through the following

processes:

(i) C2H5COC2H4p → C2H5CO + C2H4 (23)

(ii) C2H5COC2H4s → CH 3CHCO + C2H5, followed by methyl ketene and ethyl

radical decompositions.

Thus, in addition to CO, ethylene is another major product during 3-pentanone

pyrolysis, and each 3-pentanone molecule eventually turns into at least one C2H4

molecule. As illustrated in Fig. 5.12, the measured ethylene yields (defined as the ratio

of the long-time C2H4 concentration to the initial 3-pentanone concentration) for the

mixture of 0.25% 3-pentanone in Ar at 1248-1491 K are approximately 1.4. When

compared to the measurements, the Serinyel et al. mechanism fails to capture the initial

C2H4 formation rates, and the simulated ethylene yield at 1248 K is ~30% lower than the

measured value. On the other hand, the modified mechanism simulates the initial C2H4

formation rates and ultimate yields quite accurately at the measured temperatures.

Figure 5.12: C2H4 time histories for 0.25% 3-pentanone in Ar: measured and calculated values.

72

5.3.2 3-Pentanone Oxidation

Ignition delay times were measured with mixtures varying in concentration from

0.040% to 0.875% 3-pentanone in O2/balance argon over 1150-1550 K at a pressure of

~1 atm and equivalence ratios of 1.0 and 0.5. For high fuel concentration mixtures (X3-

pentanone > 0.1%), the endwall ignition delay time is defined as the time interval between

the arrival of the incident shock and the initial rise in the OH* emission

chemiluminescence trace at the endwall. The initial rise is located by linear extrapolation

of the signal at the time of maximum rate of rise to the baseline. A representative

ignition delay time plot is also provided in Fig. 5.13. On the other hand, for lower fuel

concentration mixtures, the emission signal is rather weak, and a different definition of

ignition delay time is employed. It is defined as the time to reach 50% of the peak OH

concentration (measured using OH absorption method), with time zero being defined as

the arrival of the reflected shock at the sidewall measurement location.

Figure 5.13: Sample sidewall pressure and endwall OH* emission time histories recorded during an experiment of 3-pentanone ignition at 1113 K and 1.1 atm (3-pentanone/ 4.0% O2/ Ar, Φ = 0.5). A tailored gas mixture of 60% helium/ 40% nitrogen was used as driver gas to achieve a long test time. For high fuel concentration mixtures, the definition of the endwall ignition delay time is shown in the figure.

73

Fig. 5.14 shows the measured ignition delay times from the current study (see the

solid points) at Φ = 1.0 and 0.5, along with the simulations from the original and

modified Serinyel et al. mechanisms under the assumption of constant internal energy

and constant volume. In addition, Davidson et al. [97] performed ignition delay time

measurements for the mixture of 0.571% 3-pentanone with 4% O2 in Ar (Φ = 1.0) over

1173-1306 K at pressures of 1.68-1.85 atm. Their ignition delay time data were

normalized to 1 atm using an overall correlation pressure dependence of P-0.52 (see the

hollow points), and their data are in good agreement with the current measurements

within 7%. As is evident in Fig. 5.14, fuel-lean mixtures ignite much faster than

stoichiometric mixtures, and this general trend is consistent with other hydrocarbon

studies at similar temperatures and pressures [62, 97, 102-103]. At high temperatures (T

> 1100 K), the overall reactivity is mainly controlled by the chain branching reaction (H

+ O2 → OH + O), so that reactivity is very sensitive to molecular oxygen concentration.

Therefore, at Φ = 0.5, the mixture of 0.875% 3-pentanone (with the highest O2 content)

ignites much faster than other lower fuel concentration mixtures. Interestingly, for the

mixture of 0.875% 3-pentanone with 12.25% O2 in Ar, the current measurements are

approximately 30% faster than the measurements from Serinyel et al. [98].

74

Figure 5.14: Measured and simulated 3-pentanone ignition delay times at (a) Φ = 1.0 and (b) Φ = 0.5 and P5 = 1.0 atm.

A linear regression analysis was performed on the current ignition delay time

measurements, and these measurements can be expressed in a correlation with an R2

value of 0.974:

τ [µs] = 1.232×10-5 P-0.52 Φ0.89 XO2-0.61 exp(20,450/T)

where P is the total pressure [atm], Φ is the equivalence ratio, and XO2 is the oxygen mole

fraction. As shown in Fig. 5.14, the simulated ignition times from the Serinyel et al.

mechanism are in agreement with the measured values for low fuel concentration

mixtures (400 ppm 3-pentanone) within ~35%. However, for higher fuel concentration

mixtures, the computed ignition times are approximately twice those of the

measurements. On the other hand, the modified Serinyel et al. mechanism shows much

better agreement with the measurements at all concentrations (within 30%), and the

revision of the overall 3-pentanone decomposition rate constant and the addition of the

methyl ketene decomposition pathway significantly improve the general performance of

the 3-pentanone oxidation chemistry model. When compared to the measurements from

Serinyel et al. [98], the modified mechanism also shows improved agreement with their

data (within 15%) over 1200-1550 K at 1 atm. The simulations from the original and

75

modified mechanisms at the test conditions of Serinyel et al. are also provided in Fig.

5.15.

Figure 5.15: Comparison of model predictions between (a) the Serinyel et al. mechanism of NUI Galway [98] and (b) the modified mechanism on ignition delay time measurements from Serinyel et al.

In addition to ignition delay time measurements, OH and H2O time histories were

acquired behind reflected shock waves using the mixtures of 400 ppm 3-pentanone with

O2 in balance argon over 1200-1550 K at pressures around 1.6 atm and equivalence ratios

of 1.0 and 0.5 (see Figs. 5.18-5.21). Low fuel concentration mixtures are preferred in

species time history measurements during hydrocarbon oxidation in order to minimize

76

the rapid energy release at the time of ignition, which increases the reflected shock

temperature. This would affect the analysis of species mole fractions, particularly if

constant absorption coefficients for Beer’s law are used, as has been common in the past.

For the mixture of 400 ppm 3-pentanone with 0.28% O2 in argon, there is approximately

a 50 K increase in temperature after ignition for the case at 1377 K and 1.54 atm (under

constant energy and volume constraints), which slightly perturbs (reduces) the OH and

H2O absorption coefficients by up to 4% and 2%, respectively, an effect which has been

accounted for in our data processing.

Sensitivity analysis reveals that both OH and H2O time histories (for the mixture

of 400 ppm 3-pentanone / 0.28% O2 / Ar) are mainly sensitive to the following set of

reactions (see Figs. 5.16 and 5.17):

H + O2 → OH + O (27)



C2H4 + H (+ M) → C2H5 (+ M) (20)

C2H4 + OH → C2H3 + H2O (28)

CH3 + CH3 → C2H5 + H (21)

Figure 5.16: OH sensitivity analysis for 400 ppm 3-pentanone with 0.28% O2 in Ar (Φ = 1.0) at 1486 K and 1.52 atm.

77

Figure 5.17: H2O sensitivity analysis for 400 ppm 3-pentanone with 0.28% O2 in Ar (Φ = 1.0) at 1486 K and 1.52 atm.

During hydrocarbon oxidation, H2O is regarded as an important combustion

progress marker, which gives nearly identical information to that of another combustion

progress marker, CO. The H2O profiles for typical hydrocarbons [60-61, 102, 104]

exhibit sequential features: an initial gradual formation of H2O is followed by a rapid

increase indicating ignition, which is in turn succeeded by a very slow rise in H2O

concentration. During 3-pentanone oxidation, the H2O profiles are slightly different from

those of common hydrocarbons (n-alkanes and cycloalkanes). There is no obvious

distinction between the initial gradual H2O formation and the rapid increase in H2O

concentration during ignition. The formation of H2O is fast and steady starting from time

zero, followed by a nearly constant H2O concentration. After this post-ignition H2O

plateau level has been reached, all volatile hydrocarbons have been depleted and only

small intermediates (i.e., H, O, OH, CO, CO2, and H2) would remain, gradually

approaching chemical equilibrium. The kinetics that controls these small intermediates

(after ignition) is well-established and subject to small uncertainties [22, 33, 72].

Additionally, this nearly constant H2O plateau level (after ignition) is primarily sensitive

to the relatively well-established thermodynamic parameters, as suggested by Hong et al.

[102]. This observation was also validated through the perturbation of the rate constants

78

for the above important reactions, following which the post-ignition H2O concentration

level effectively remained the same. Thus, the post-ignition H2O plateau level can be

used to confirm pre-shock fuel concentration.

Fig. 5.18 shows the measured H2O time histories for the mixture of 400 ppm 3-

pentanone with 0.28% O2 in argon (Φ = 1.0), along with the computed profiles from the

(a) original (top) and (b) modified (bottom) Serinyel et al. mechanisms. At a first glance,

the measured and computed H2O plateau levels are consistent with each other, and this

agreement justifies the accuracy of the initial fuel loading in our experiments. The

original Serinyel et al. mechanism captures the general shape of the measured H2O

profiles, but it does not predict the initial H2O formation rates accurately. In addition, the

modified mechanism performs slightly better at higher temperatures (1486 K and 1542

K), and the computed profiles from the modified mechanism show much better

agreement with the measurements at lower temperatures (1343 K and 1377 K).

Similarly, Fig. 5.19 illustrates the measured H2O time histories for the mixture of

400 ppm 3-pentanone with 0.56% O2 in argon (Φ = 0.5), along with the simulations from

the (a) original (top) and (b) modified (bottom) Serinyel et al. mechanisms. Here also,

the measured and computed H2O plateau levels (after ignition) are consistent with each

other. For this fuel-lean mixture, the computed profiles from the original Serinyel et al.

mechanism are quite different from the measured profiles in terms of the initial formation

rates and the ignition delay times, especially when compared to the stoichiometric

mixture. At 1217 K and 1.74 atm, the model significantly underpredicts the formation

rate of H2O, and the computed ignition delay time based on the H2O time history is at

least twice that of the measured value. On the other hand, the modified mechanism is

able to capture the H2O formation rates reasonably well at all temperatures, particularly at

1217 K. Hence, the revision of the overall 3-pentanone decomposition rate constant and

the addition of the methyl ketene decomposition pathway greatly improve the model

predictions of H2O concentrations under oxidizing conditions.

79

Figure 5.18: Comparisons of measured and simulated H2O time histories from the (a) original and (b) modified Serinyel et al. mechanisms for 400 ppm 3-pentanone with 0.28% O2 in Ar (Φ = 1.0).

80

Figure 5.19: Comparisons of measured and simulated H2O time histories from the (a) original and (b) modified Serinyel et al. mechanisms for 400 ppm 3-pentanone with 0.56% O2 in Ar (Φ = 0.5).

81

OH species time histories provide another important kinetic target for the

validation of detailed kinetic mechanisms during hydrocarbon oxidation. At early times,

there is the rapid formation of OH simultaneous with the initial fuel decomposition.

Based on the temperature, the OH species develops a well-defined plateau level during

the induction period, especially at lower temperatures. At the time of ignition, the OH

mole fraction rapidly increases to its post-ignition plateau level. More importantly, the

early-time feature of OH appears to provide important information about the breakdown

of the fuel, because this feature is governed by the fuel unimolecular decomposition and

the H-atom abstraction reactions from the fuel (mainly by H radicals). At later times, the

rapid OH rise at ignition is not unique to this fuel. At the time of ignition, fuel molecules

have mostly decomposed into small intermediate radicals or molecules, such as H, H2,

C2H4, C3H6, etc., and these small fragments tend to control the rate of ignition (the rapid

OH rise at ignition) during hydrocarbon oxidation, as suggested by Warnatz et al. [105].

Fig. 5.20 illustrates the measured OH time histories for the mixture of 400 ppm 3-

pentanone with 0.28% O2 in argon (Φ = 1.0), along with the simulated profiles from the

(a) original (top) and (b) modified (bottom) Serinyel et al. mechanisms. At higher

temperatures (1486 K and 1542 K), the original Serinyel et al. mechanism captures the

rapid OH rise at the time of ignition reasonably well, but it underpredicts the initial OH

formation rates and the initial OH plateau levels (see inset on Fig. 5.20). At lower

temperatures (1343 K and 1377 K), the model underpredicts both the initial and final OH

formation rates, but it is able to simulate the initial plateau levels quite well. When

compared to the original Serinyel et al. mechanism, the modified model generally

provides a much better agreement with the measurements, in terms of the initial plateau

levels and both the initial and final formation rates. However, the modified model still

cannot capture the slight overshoot prior to the formation of the initial plateau. The

overshoot seems to be more obvious at higher temperatures, and such overshoot is quite

sensitive to the H-atom abstraction reactions from 3-pentanone by H radicals and these

rate constants may require some fine adjustment at higher temperatures.

Fig. 5.21 shows the measured OH time histories for the mixture of 400 ppm 3-

pentanone with 0.56% O2 in argon (Φ = 0.5), along with the simulations from the (a)

82

original (top) and (b) modified (bottom) Serinyel et al. mechanisms. It should be noted

that there is now no noticeable overshoot in OH concentration prior to the formation of

the initial plateau for the fuel-lean mixture. Instead, there is a smooth transition from the

initial OH rise to the first plateau level. In addition, with more O2 molecules in the

system, the initial and post-ignition plateau levels are much higher than those of the

stoichoimetric mixture, and this observation is well-simulated by the model. This is

mainly due to the fact that more oxygen molecules are available to undergo the chain

branching reaction (H + O2 → OH + O). The Serinyel et al. mechanism underpredicts

both the initial and final OH formation rates by at least 30% at all temperatures.

Similarly, it does not simulate the initial plateau levels properly at temperatures greater

than 1300 K. On the other hand, the modified model is able to simulate the OH time

histories very accurately, in terms of the initial plateau levels and the initial and final

formation rates, particularly at 1217 K. Hence, the revision of the overall 3-pentanone

decomposition rate constant greatly improves the model predictions on the first OH

plateau levels and its initial formation rates, and the branching ratio of reaction (4a) for

the initial fuel decomposition pathways predicted by Serinyel et al. seems to be quite

reasonable. More importantly, the methyl ketene decomposition pathway seems to

promote the ignition behavior. One possible explanation is that more CO molecules are

now introduced to the system, and each CO molecule reacts with an OH radical through

the reaction of CO + OH → CO2 + H, which is an important exothermic reaction in

combustion chemistry. As a result, more heat is generated during 3-pentanone oxidation,

thereby promoting the ignition behavior.

83

Figure 5.20: Comparisons of measured and simulated OH time histories from the (a) original and (b) modified Serinyel et al. mechanisms for 400 ppm 3-pentanone with 0.28% O2 in Ar (Φ = 1.0). Inset figures are provided to show the early-time features over 0-400 µs.

84

Figure 5.21: Comparisons of measured and simulated OH time histories from the (a) original and (b) modified Serinyel et al. mechanisms for 400 ppm 3-pentanone with 0.56% O2 in Ar (Φ = 0.5). Inset figures are provided to show the early-time features over 0-400 µs.

5.3.3 Comparisons of Ketone Oxidation Characteristics

In addition to 3-pentanone oxidation, the oxidation characteristics of acetone and

2-pentanone were examined and compared with that of 3-pentanone using ignition delay

85

time and OH time history measurements for the mixtures of ketone with 0.525% O2 in

balance argon at Φ = 1.0. Shown in Fig. 5.22 is a plot of the measured ignition delay

times for these three ketone mixtures. Strikingly, the ignition delay times of 3-pentanone

mixtures are only about half those of acetone and 2-pentanone mixtures. However, the

apparent activation energies of these ketone mixtures are quite similar.

Figure 5.22: Comparison of ignition delay times for different ketones (acetone, 2-pentanone and 3-pentanone).

Shown in Fig. 5.23 is a plot of the measured OH time histories for the above

ketone mixtures at T ≈ 1370 K and P = 2.6 atm. Note that the OH time history for the 3-

pentanone mixture is at a slightly lower temperature (1354 K). Interestingly, the 3-

pentanone mixture has a much higher initial OH plateau level than the acetone and 2-

pentanone mixtures, while the initial OH plateau levels for the acetone and 2-pentanone

mixtures are approximately the same. As discussed above, this early-time OH feature is

unique to the individual fuel, and is primarily controlled by the initial fuel decomposition

pathways and the H-atom abstraction reactions from the fuel by H radicals. In addition,

as suggested by Davidson et al. [62], the initial OH plateau level is quite sensitive to the

branching ratios of the fuel decomposition pathways. For instance, there is often a direct

86

competition between the C2H5 and CH3 formation channels during 3-pentanone

decomposition. A slight increase in the rate constant for the C2H5 channel can further

increase the first OH plateau level, and a slight increase in the rate constant for the CH3

channel can reduce the plateau level. Once the C2H5 radical is formed, it is quickly

followed by its decomposition to form an H radical. The H radical can then undergo the

chain branching reaction (with O2) to form an OH radical. On the other hand, if the CH3

radical is formed, it tends to form ethane through the methyl-methyl recombination

reaction and fewer H radicals are formed. In the case of acetone and 2-pentanone

decomposition, CH3 radicals are generally formed, and insufficient amounts of C2H5

radicals are developed to give H radicals. Thus, the initial OH plateau levels of acetone

and 2-pentanone are much less than that of 3-pentanone. More importantly, there seems

to be a strong positive correlation between the initial OH plateau level and the ignition

delay time. This correlation further explains why 3-pentanone has the fastest ignition

delay times among these ketone mixtures at current experimental conditions.

Figure 5.23: Comparison of OH time histories for the mixtures of ketone (i.e., acetone, 2-pentanone and 3-pentanone) with 0.525% O2 in Ar at a pressure of 2.6 atm and an equivalence ratio of 1.0. An inset figure is provided to show the early-time features over 0-400 µs.

87

5.4 Summary

High-temperature 3-pentanone pyrolysis and oxidation studies were investigated

using laser-based species time history measurements for 3-pentanone, CH3, CO, C2H4,

H2O and OH. To our knowledge, these measurements are the first laser-based species

time history measurements for high-temperature 3-pentanone pyrolysis and oxidation.

Using these time histories and the Serinyel et al. 3-pentanone mechanism [98],

improved determinations of the initial 3-pentanone unimolecular decomposition reactions

were possible. As well, a comparison of the measured and modeled CO time history

pathways identified the need to include the methyl ketene decomposition pathway to

improve the simulations. These two modifications to the Serinyel et al. mechanism also

significantly improved the agreement with ignition delay times and OH and H2O time

histories during 3-pentanone oxidation. Finally, a comparison of OH time histories

during the oxidation of 3-pentanone, acetone, and 2-pentanone showed that the initial OH

plateau level of 3-pentanone was higher than that of acetone and 2-pentanone and this

was consistent with the shorter ignition delay times seen with this fuel.

5.5 Possible Future Work

More work is definitely needed to further improve the model predictions under 3-

pentanone pyrolytic and oxidizing conditions. As discussed in Section 5.3.1, the original

Serinyel et al. mechanism [98] appears to overpredict the methyl ketene concentration

during 3-pentanone pyrolysis, and such effect cannot currently be observed

experimentally (based on the O-atom balance). In the present analysis, the possible

solution to such discrepancy is to introduce a unimolecular decomposition reaction for

methyl ketene (reaction (26)) in order to remove the pooling of methyl ketene in the

original model, and reasonable agreement can be obtained between the current

measurements and the simulations from the modified model. Despite its improved

predictive capability, the modified model is very likely to suffer from other deficiencies

that have not been addressed in this dissertation. For instance, the rate constant for

88

reaction (24) (C2H5COC2H4s → C2H5 + CH3CHCO) is possibly too fast, resulting in the

pooling of methyl ketene and prohibiting the CO formation. Therefore, a slower rate

constant for reaction (24) is recommended. Additionally, as demonstrated in Chapters 4

and 5, the kinetics of methyl ketene is poorly understood, and more experimental and

theoretical studies for methyl ketene are definitely required. In particular, the rate

constant for reaction (26) (CH3CHCO (+ M) → C2H4 + CO (+ M)) was only estimated by

analogy with the rate constant for ketene unimolecular decomposition in the present

analysis, and there is a need for a more accurate rate constant expression.

One of the major weaknesses in the modified mechanism is the fact that it is not

able to accurately predict the ignition delay times for high fuel concentration mixtures

(X3-pentanone > 0.2%), as shown in Fig. 5.14. Ignition delay time is an important global

kinetic target that is commonly used for the validation of the detailed models, but this

global target is quite sensitive to many secondary chemistry reactions. In particular, the

C2 chemistry, such as the ethyl radical decomposition reaction, is very crucial to the

development of the successful 3-pentanone kinetic model suitable for high-temperature

application. In the future, different base mechanisms, which might consist of different

C2 chemistry sets, should also be considered.

89

Chapter 6 High-Temperature Measurements of the Reactions of OH with a Series of Ketones: Acetone, 2-Butanone, 3-Pentanone, and 2-Pentanone

6.1 Introduction

Due to their significant roles in atmospheric chemistry, the rate constants for the

reactions of OH radicals with a series of ketones, including acetone, 2-butanone, 3-

pentanone and 2-pentanone, have been extensively studied by many researchers [38, 106-

116] over the temperature range of 240-400 K. However, the kinetic data on ketones +

OH at combustion-relevant conditions are generally scarce. There were a few

experimental studies for the acetone + OH reaction rate constant over 500-1300 K.

Yamada et al. [117] utilized two different OH precursors (HONO and N2O/H2O) and

measured the rate constants for OH + CH3COCH3 and CD3COCD3 in a reactor over 298-

832 K using the pulsed laser photolysis/pulsed laser-induced fluorescence technique.

Bott and Cohen [118] pioneered the use of tert-butyl hydroperoxide as an OH precursor

and monitored the OH decay in a shock tube using the UV lamp absorption method at

309 nm in order to study the rate constant for acetone + OH reaction near 1200 K and 1

atm. Similarly, Vasudevan et al. [119] and Srinivasan et al. [77] both measured the

acetone + OH rate constant using shock tubes and UV absorption methods over the

combustion-relevant temperature range of 980-1300 K. These measurements are in good

90

agreement with each other. In contrast to acetone, there was only one experimental study

available for larger ketone + OH kinetic data. Tranter and Walker [120] added small

amounts of ketones (acetone, 2-butanone and 3-pentanone) individually to slowly

reacting mixtures of H2 + O2 at 753 K, and measured the consumption of ketones and H2

with the use of gas chromatography. This method allowed them to study the relative rate

constants for the reactions of H and OH with ketones at 753 K. Furthermore, Zhou et al.

[121] recently performed a theoretical study on the mechanism and kinetics of the

reactions of OH with three methyl ketones: acetone, 2-butanone and isopropyl methyl

ketone. They employed the computationally less expensive G3 and G3MP2BH&H

methods to calculate the energy barriers, and utilized the Variflex code including Eckart

tunneling corrections to compute the total rate constants over 500-2000 K. In addition,

all possible abstraction channels have been accounted for in their calculation. However,

except for acetone, their theoretical calculations have not been validated against any

high-temperature experimental data.

The overall rate constants for the reactions of OH with four ketones, namely

acetone (CH3COCH3), 2-butanone (C2H5COCH3), 3-pentanone (C2H5COC2H5) and 2-

pentanone (C3H7COCH3), were determined behind reflected shock waves over the

temperature range of 870-1360 K at pressures of 1-2 atm:

CH3COCH3 + OH → Products (5)




These measurements include the first direct high-temperature measurements of the

overall rate constants for reactions (6)-(8). These high-temperature kinetic data, along

with the earlier work [38, 77, 106-120], are compared with the theoretical calculations

(from Zhou et al. [121]) and the estimates using the group-additivity model.

91





manometrical preparation of a highly dilute mixture. A highly concentrated mixture was



the experiments. The gas utilized in this study was argon (Research Grade) 99.999%,

which was supplied by Praxair and used without further purification. The liquid

chemicals were 70% tert-butyl hydroperoxide (TBHP) in water, CHROMASOLV® grade

acetone (≥99.9%), CHROMASOLV® grade 2-butanone (≥99.7%), ReagentPlus® grade 3-

pentanone (≥99%), and ReagentPlus® grade 2-pentanone (≥99%) from Sigma-Aldrich,

and were purified using a freeze-pump-thaw procedure to remove dissolved volatiles and

air prior to mixture preparation.

The mixture composition was confirmed by sampling a portion of the mixture

(from near the endwall) into an external multi-pass absorption cell with a path length of

29.9 m and monitoring the fuel concentration in the cell with a Jodon™ Helium-Neon

laser at 3.39 µm. The details of the laser diagnostic set-up are discussed elsewhere [122].

Beer’s law was used to convert the measured absorption data into the fuel mole fraction.

The absorption cross-sections of ketones for Beer’s law were directly obtained from the

PNNL database [123], and the measured fuel concentrations were consistent with the

values expected from the manometrical preparation within ±5%.


6.3.1 Choice of Kinetic Mechanisms

A total of 58 reflected shock wave experiments were performed to determine the

overall rate constants for the reactions of OH with four ketones (acetone, 2-butanone, 3-

92

pentanone and 2-pentanone) at near-pseudo-first-order conditions. Experiments were

carried out over the temperature range of 870-1360 K at pressures of 1-2 atm using

different initial fuel concentrations: acetone (304 ppm), 2-butanone (152 ppm, 161 ppm

and 206 ppm), 3-pentanone (151 ppm and 211 ppm), and 2-pentanone (161 ppm). These

ketones were prepared with 53-101 ppm TBHP/water and diluted in argon. To properly

simulate the consumption of OH radicals by ketones, the Pichon et al. mechanism of NUI

Galway [89] with the revised k2 was chosen as the base mechanism for acetone, and the

Serinyel et al. mechanism of NUI Galway [95, 98] with the revised k3 and k4 and the

addition of the methyl ketene decomposition pathway was utilized as the base mechanism

for 2-butanone, 3-pentanone, and 2-pentanone. In addition, the tert-butyl hydroperoxide

(TBHP) sub-mechanism was incorporated in these base mechanisms. (Please read

Chapter 3 for more details on the TBHP chemistry.) Similarly, the thermodynamic

parameters for TBHP and tert-butoxy radical were taken from the thermodynamic

database from Goos et al. [76], and the thermodynamic parameters for OH were updated

with the values from Herbon et al. [23-24].

As discussed in Chapters 4 and 5, the initial decomposition pathways for acetone,

2-butanone, and 3-pentanone can be described as follows:

CH3COCH3 (+ M) → CH3CO + CH3 (+ M) (2)

C2H5COCH3 (+ M) → C2H5 + CH3CO (+ M) (3a)

C2H5COCH3 (+ M) → CH3 + C2H5CO (+ M) (3b)

C2H5COCH3 (+ M) → CH3 + CH3COCH2 (+ M) (3c)



For 2-butanone and 3-pentanone (and 2-pentanone), their initial decomposition pathways

consist of multiple channels. High-temperature decomposition pathways for 2-butanone

and 3-pentanone were first investigated by Serinyel et al. [95, 98]. Recently, Lam et al.

[124-125] have performed experimental studies during high-temperature acetone, 2-

butanone, and 3-pentanone pyrolysis, as were discussed in Chapters 4 and 5. In their

studies, they measured the rate constant for reaction (2) and the overall values for

reactions (3) and (4) at pressures near 1.6 atm. At T > 1300 K, the consumption of

93

ketones in the present study is mainly controlled by the H-atom abstraction reactions by

OH radicals and the ketone decomposition pathways. Hence, reactions (2)-(4) are

pertinent to the determinations of the overall rate constants for reactions (5)-(7) at higher

temperatures, and the rate constants for reactions (2)-(4) were updated with the values

from Lam et al. [124-125] (the rate constants from Chapters 4 and 5). In addition, a

review of the literature shows that there is currently no experimental or theoretical study

for high-temperature 2-pentanone pyrolysis. Thus, 2-pentanone decomposition

pathways, along with the corresponding rate constants, are not known in this study, and

the overall rate constant for 2-pentanone + OH reaction (reaction (8)) cannot be inferred

accurately at T > 1300 K. A thorough theoretical or experimental study for 2-pentanone

decomposition pathways is required. Nevertheless, at T < 1300 K, the consumption of 2-

pentanone in the present study is predominantly controlled by the H-atom abstraction

reactions by OH radicals, and the overall rate constant for reaction (8) was determined

over a narrower temperature range 900-1300 K. In addition, all simulations were

performed using the CHEMKIN PRO package [71] under the standard constant internal

energy and volume assumption.

6.3.2 Acetone + OH Kinetics

An OH radical sensitivity analysis for the mixture of 304 ppm acetone with 28

ppm TBHP (and 73 ppm H2O) in Ar at 1148 K and 1.95 atm is provided in Figure 6.1.

The analysis reveals that the reaction of OH with acetone (reaction (5)) is the dominant

reaction over the time frame of the experiment, with some minor interference from the

secondary reactions:

CH3 + OH → CH2(s) + H2O (17)

C2H6 (+ M) → CH3 + CH3 (+ M) (18)

CH3OH (+ M) → CH3 + OH (+ M) (19)

The rate constants for reactions (17)-(19) were updated with the values from Table 3.1 (in

Chapter 3).

94

Figure 6.1: OH sensitivity plot for the rate constant measurement of acetone + OH at 1148 K and 1.95 atm.

Figure 6.2 shows a sample measured OH concentration time history for the

mixture of 304 ppm acetone in Ar at 1148 K and 1.95 atm, and the measured peak OH

mole fraction is ~28 ppm. Due to wall adsorption and condensation of TBHP, the initial

TBHP mole fraction was assumed to be the same as the measured peak OH mole fraction,

which was formed immediately after the decomposition of TBHP behind the reflected

shock wave at T > 1000 K. Note that a 70%, by weight, solution of TBHP in water in the

liquid phase corresponds, initially, to 69% water and 31% TBHP in the vapor phase,

based on Raoult’s law [126]. Therefore, a 101 ppm TBHP/water mixture should have at

most 31.3 ppm TBHP. In the present study, the mixtures of 101 ppm TBHP/water

consist of ~28-30 ppm TBHP, based on the measured peak OH yields, hence suggesting

very little loss to the walls. In addition, the test mixtures were chosen such that the ratio

of the initial acetone concentration to the initial TBHP concentration is ~10, thereby

achieving near-pseudo-first-order conditions. For the conditions described in Figure 6.2,

a best-fit overall rate constant for reaction (5) of 3.83×1012 cm3 mol-1 s-1 was obtained

between the experimental data and the simulation. Simulations for the perturbations of

±50% in the inferred rate constant are also illustrated in Figure 6.2. In addition, Table 6.1

95

summarizes the rate constant measurements of reaction (5) at 872-1355 K and 1.69-2.12

atm.

Figure 6.2: Sample acetone + OH rate constant measurement using the mixture of 304 ppm acetone with ~28 ppm TBHP (and 73 ppm water) in Ar at 1148 K and 1.95 atm. Simulation from the modified Pichon et al. mechanism for the best-fit rate constant, along with perturbations of ±50%, is also shown.

Table 6.1: CH3COCH3 + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k5 [cm3 mol-1 s-1]

101 ppm TBHP (and water), 304 ppm CH3COCH3, Ar 934 2.01 2.52E+12 1008 2.09 2.72E+12 1011 1.85 2.83E+12 1111 1.98 3.47E+12 1148 1.95 3.83E+12 1221 1.86 4.50E+12 1247 1.82 4.96E+12 1280 1.78 4.90E+12 1307 1.69 5.29E+12 1355 1.75 5.63E+12

92 ppm TBHP (and water), 304 ppm CH3COCH3, Ar 872 1.96 2.14E+12 996 2.12 2.79E+12

96

A detailed error analysis was performed to estimate the uncertainty limits of the

measured rate constant for reaction (5) at 1148 K. The primary contributions to the

uncertainties in the rate constant are: (a) temperature (±1%), (b) mixture composition

(±5%), (c) OH absorption coefficient (±3%), (d) wavemeter reading in the UV (±0.01 cm-

1), (e) fitting the data to computed profiles (±5%), (f) locating time-zero (±0.5 µs), (g) the

rate constant for CH3 + OH → CH2(s) + H2O (uncert. factor = 2), (h) the rate constant

for CH3OH (+ M) → CH3 + OH (+ M) (uncert. factor = 2), and (i) the rate constant for

C2H6 (+ M) → CH3 + CH3 (+ M) (±20%). As shown in Figure 6.3, the individual error

sources were introduced separately and their effects on the rate constant for reaction (5)

were determined. These uncertainties were combined in a root-sum-squared method to

give an overall uncertainty estimate of ±28% at 1148 K.

Figure 6.3: Uncertainty analysis for the rate constant of CH3COCH3 + OH → products at 1148 K and 1.95 atm.

Figure 6.4 shows the Arrhenius plot for the present rate constant measurements of

reaction (5) at T = 872-1355 K, along with the previous measurements of Vasudevan et

al. [119] from the same laboratory. The current measurements agree well with the

previous values (within ±5%). These measured values can then be expressed in

Arrhenius form as k5 = 3.30×1013 exp(-2437/T) cm3 mol-1 s-1 over 872-1355 K. Bott and

Cohen [118] also utilized TBHP as the OH precursor and employed both the shock tube

97

and UV lamp absorption method at 309 nm to monitor the OH decay and study reaction

(5) near 1200 K and 1 atm. The current measurements are consistent with Bott and

Cohen’s measured value within 20%. In addition, Srinivasan et al. [77] used a similar

method to investigate the rate constant for reaction (5) and determined a rate constant of

4.40×1012 cm3 mol-1 s-1 over 1178-1299 K. Their value is in close accord with our

previous and current measurements. Figure 6.4 also shows the rate constants for reaction

(5) adopted by two different detailed mechanisms: Pichon et al. [89] and Herbinet et al.

[45]. The values of k5 from the original Pichon et al. mechanism are approximately 24%

and 43% faster than the current measurements at 1000 K and 1250 K, respectively; the

values employed from the Herbinet et al. mechanism are in excellent agreement with the

current measured values (within ±11%).

Additionally, a theoretical calculation from Zhou et al. [121], which modeled all

possible abstraction channels, was performed using the computationally less expensive

G3 and G3MP2BH&H methods to calculate the energy barriers and using the Variflex

code including Eckart tunneling corrections to compute the total rate constants for the

reactions of OH with ketones (acetone, 2-butanone, and isopropyl methyl ketone) over

500-2000 K. As shown in Figure 6.4, the computed values from Zhou et al. are

consistently lower than all high-temperature experimental data by ~55%.

Figure 6.4: Arrhenius plot for acetone + OH (k5) at temperatures above 833 K.

98

6.3.3 2-Butanone + OH Kinetics

The OH sensitivity analysis was also carried out for the rate constant

determination of 2-butanone + OH → products (reaction (6)) using the mixture of 152

ppm 2-butanone with 14 ppm TBHP (and 41 ppm water) in Ar at 1039 K and 1.41 atm,

as shown in Figure 6.5. Note that reaction (6) consists of 3 different abstraction channels,

as described in the original Serinyel et al. mechanism [95, 98]:

C2H5COCH3 + OH → CH2CH2COCH3 + H2O (6a)

C2H5COCH3 + OH → CH3CHCOCH3 + H2O (6b)

C2H5COCH3 + OH → C2H5COCH2 + H2O (6c)

At 1100 K, channel (6b) is the dominant pathway with a branching ratio of 0.53 due to

the weaker C-H bond energy at the secondary site, and channel (6a) is the next most

important pathway with a branching ratio of 0.41. However, channel (6c) is nearly

insignificant with a branching ratio of 0.06. More importantly, reaction (6) is the most

sensitive reaction at the conditions depicted in Figure 6.5, with some minor interference

from the secondary reactions (reactions (13), (17) and (18)). As shown here, as

temperature decreases, the reaction for TBHP decomposition becomes more important at

the early times.

Figure 6.5: OH sensitivity plot for the rate constant measurement of 2-butanone + OH at 1039 K and 1.41 atm.

99

Figure 6.6 shows an example of the overall rate constant measurement (k6 = k6a +

k6b + k6c) for reaction (6) at 1039 K and 1.41 atm. The mixture is 152 ppm 2-butanone in

Ar, with the measured peak OH yield to be ~14 ppm. Thus, we infer that the initial

TBHP mole fraction is 14 ppm. The model predictions from the modified Serinyel et al.

mechanism with the best-fit overall rate constant of k6 = 6.82×1012 cm3 mol-1 s-1 and the

variations of ±50% in the inferred rate constant are also shown in Figure 6.6. Due to the

near-pseudo-first-order conditions, the measured overall rate constant should be

insensitive to the branching ratios of the individual channels. The effect of the branching

ratios on the rate constant determination was also investigated at 1039 K by

interchanging the branching ratios of channels (6a) and (6b) while maintaining the total

value, and a negligible change in the inferred rate constant was found. In addition, a

detailed error analysis (similar to the analysis for reaction (5)) was performed for the rate

constant measurement of reaction (6) at 1039 K and 1.41 atm, and the overall uncertainty

was estimated to be ±22%. Table 6.2 summarizes the rate constant measurements of

reaction (6) at 879-1364 K and 1.21-1.63 atm. Three different mixture compositions

were employed to confirm that the inferred rate constants are independent of any

secondary chemistry effects.

Figure 6.6: Sample 2-butanone + OH rate constant measurement using the mixture of 152 ppm 2-butanone with ~14 ppm TBHP (and 41 ppm water) in Ar at 1039 K and 1.41 atm. Simulation from the modified Serinyel et al. mechanism for the best-fit rate constant, along with perturbations of ±50%, is also shown.

100

Table 6.2: C2H5COCH3 + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k6 [cm3 mol-1 s-1]

55 ppm TBHP (and water), 152 ppm C2H5COCH3, Ar 962 1.51 6.10E+12 999 1.45 6.42E+12

1039 1.41 6.82E+12 1119 1.27 8.25E+12 1174 1.24 9.01E+12 1247 1.26 1.05E+13

53 ppm TBHP (and water), 161 ppm C2H5COCH3, Ar

879 1.57 5.06E+12 955 1.63 5.93E+12 1110 1.39 8.00E+12 1282 1.26 1.09E+13 1297 1.23 1.14E+13

65 ppm TBHP (and water), 206 ppm C2H5COCH3, Ar

905 1.55 5.30E+12 1088 1.42 7.51E+12 1104 1.34 7.80E+12 1320 1.21 1.17E+13 1364 1.25 1.24E+13

Figure 6.7 shows the Arrhenius plot for the overall rate constant measurements of

reaction (6) at T > 833 K, along with the estimated values adopted in the original Serinyel

et al. mechanism [95, 98] and the theoretical values from Zhou et al. [121]. The

measured values can be expressed in Arrhenius form as k6 = 6.35×1013 exp(-2270/T) cm3

mol-1 s-1 over 879-1364 K. The values used in the original Serinyel et al. mechanism are

~40% lower than the measurements. Interestingly, the theoretical values from Zhou et al.

are in excellent agreement with the measurements within 10%. Note that the

measurements and the theoretical calculations both exhibit some slight non-Arrhenius

curvature at the present test conditions.

101

Figure 6.7: Arrhenius plot for 2-butanone + OH (k6) at temperatures above 833 K.

6.3.4 3-Pentanone + OH Kinetics

As illustrated in Figure 6.8, the OH sensitivity reveals that the reactions of OH

with 3-pentanone are the dominant pathways for the consumption of OH at 1188 K and

1.94 atm. In particular, reaction (7) consists of two channels, in which the OH radical

can abstract the H-atom from 3-pentanone at the primary or secondary site.

C2H5COC2H5 + OH → CH2CH2COC2H5 + H2O (7a)

C2H5COC2H5 + OH → CH3CHCOC2H5 + H2O (7b)

Based on the original Serinyel et al. mechanism [95, 98], the branching ratios of channels

(7a) and (7b) are 0.42 and 0.58, respectively, at 1188 K. In addition, there is some minor

interference from the following reactions at later times:

CH3 + OH → CH2(s) + H2O (17)

C2H4 + H (+ M) → C2H5 (+ M) (20)


In the current analysis, the rate constant for reaction (5) was updated with the Arrhenius

expression from Section 6.3.2, with an uncertainty of approximately ±28%. In addition,

102

the rate constant for reaction (20) adopted by the original Serinyel et al. mechanism was

used, and we assumed that its uncertainty is approximately a factor of 2.

Figure 6.8: OH sensitivity plot for the rate constant measurement of 3-pentanone + OH at 1188 K and 1.94 atm. Figure 6.9 shows a representative OH time history trace at 1188 K and 1.94 atm

using the mixture of 213 ppm 3-pentanone with 17 ppm TBHP (and 59 ppm H2O) in Ar.

The model predictions from the modified Serinyel et al. mechanism with the best-fit rate

constant of k7 = 1.23×1013 cm3 mol-1 s-1 and the variations of ±50% in k7 are also shown

in Figure 6.9. Note that the overall rate constant for reaction (7) is insensitive to the

branching ratios of its individual channels due to the near-pseudo-first-order conditions.

A detailed error analysis was then conducted for k7 at 1188 K and 1.94 atm, and the

overall uncertainty was estimated to be ±20%.

103

Figure 6.9: Sample 3-pentanone + OH rate constant measurement using the mixture of 213 ppm 3-pentanone with ~17 ppm TBHP (and 59 ppm water) in Ar at 1188 K and 1.94 atm. Simulation from the modified Serinyel et al. mechanism for the best-fit rate constant, along with perturbations of ±50%, is also shown.

Table 6.3 summarizes the overall rate constant determinations of reaction (7) at T

= 878-1353 K and P = 1.21-2.20 atm. Note that two different mixture compositions were

used to confirm that the inferred rate constants are free of any secondary chemistry

effects, and the values determined from these two mixtures are consistent with each

other. Figure 6.10 shows the Arrhenius plot for our measured values, along with the

estimated values in the original Serinyel et al. mechanism [95, 98], at temperatures above

833 K. The measured values can be expressed in Arrhenius form as k7 = 9.29×1013 exp(-

2361/T) cm3 mol-1 s-1 over 878-1353 K. Interestingly, the values for reaction (7) from the

original Serinyel et al. mechanism were estimated by analogy with the H-atom

abstraction rate constants from alkanes [98], and their values are in close accord with the

present measurements within ±5%. As is evident in Figure 6.10, the measurements and

the estimated values from Serinyel et al. both experience slight non-Arrhenius curvature

at the current experimental conditions.

Additionally, based on the theoretical study of the reactions of OH with ketones

from Zhou et al. [121], the expressions of the group rate constants (on a per H-atom

104

basis) for different carbon atom sites (primary, secondary, and tertiary carbon atoms)

were provided. In the present analysis, we can estimate the overall rate constant for

reaction (7) using these group rate constants, and the estimated rate constant is k7 = 6 ×

k(CH3CH2C(O)) + 4 × k(–CH2C(O)), where k(CH3CH2C(O)) and k(–CH2C(O)) are the

group rate constants (per H-atom) for the primary carbon atom adjacent to the –

CH2C(O)– group and for the secondary carbon atom adjacent to the –C(O)– group,

respectively. As shown in Figure 6.10, the estimated values are in good agreement with

the measurements within 15%.

Table 6.3: C2H5COC2H5 + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k7 [cm3 mol-1 s-1]

76 ppm TBHP (and water), 213 ppm C2H5COC2H5, Ar 1140 1.61 1.12E+13 1188 1.94 1.23E+13 1248 1.76 1.33E+13 1296 1.79 1.57E+13 1310 1.74 1.61E+13

76 ppm TBHP (and water), 211 ppm C2H5COC2H5, Ar 936 2.02 7.38E+12 1025 2.13 8.90E+12 1068 1.94 9.54E+12

75 ppm TBHP (and water), 211 ppm C2H5COC2H5, Ar 878 1.98 6.61E+12 955 2.20 7.86E+12 981 2.13 8.09E+12 1093 1.91 1.06E+13 1339 1.71 1.67E+13

53 ppm TBHP (and water), 151 ppm C2H5COC2H5, Ar 1091 1.46 1.03E+13 1157 1.34 1.16E+13 1192 1.29 1.26E+13 1258 1.27 1.43E+13 1301 1.21 1.57E+13 1353 1.23 1.69E+13

105

Figure 6.10: Arrhenius plot for 3-pentanone + OH (k7) at temperatures above 833 K.

6.3.5 2-Pentanone + OH Kinetics

As mentioned previously, a comprehensive mechanism for high-temperature 2-

pentanone kinetics is not available in the literature. In the present work, we have

assumed the pathways for the reactions of OH with 2-pentanone to be similar to those of

methyl butanoate [54].

C3H7COCH3 + OH → C2H4 + CH3COCH2 + H2O (8a)

C3H7COCH3 + OH → C3H6 + CH3CO + H2O (8b)

C3H7COCH3 + OH → C2H5CHCO + CH3 + H2O (8c)

C3H7COCH3 + OH → n-C3H7 + CH2CO + H2O (8d)

Channel (8a) describes the H-atom abstraction from 2-pentanone at the γ site to form a

CH2CH2CH2COCH3 radical and a H2O molecule. Through β-scission, the fuel radical

decomposes very rapidly to form a C2H4 molecule and a CH3COCH2 radical. Due to the

rapid decomposition of the fuel radical, we assumed that the products from the fuel

radical are formed immediately after the H-atom abstraction. Similarly, channels (8b)

and (8c) describe the H-atom abstraction from 2-pentanone at the β and α sites,

respectively. Due to its similar structure to methyl butanoate, the rate constants for

106

channels (8a)-(8c) were approximated to be the same as the rate constants for methyl

butanoate (MB) + OH reactions at the α, β and γ sites, and these values for MB + OH

reactions were obtained from the Dooley et al. mechanism [54]. Among these three

channels, channel (8b) (the H-atom abstraction at the β position) should be the fastest

route for the removal of OH, which was also suggested in previous experimental studies

[38, 107]. In addition, the rate constant for channel (8d) was assumed to be the same as

that of channel (6c) (C2H5COCH3 + OH → C2H5COCH2 + H2O). The resulting

branching ratios of channels (8a)-(8d) at 1186 K are 0.23, 0.38, 0.37 and 0.02. These

four channels were then incorporated in the modified Serinyel et al. mechanism. As

expected, the estimated branching ratios of channels (8a)-(8d) have no discernible effect

on the determinations of the overall rate constant at near-pseudo-first-order conditions.

In the present analysis, we also included the pathways for the reactions of H with

2-pentanone in the modified Serinyel et al. mechanism, which can be described as

follows:

C3H7COCH3 + H → C2H4 + CH3COCH2 + H2 (29a)

C3H7COCH3 + H → C3H6 + CH3CO + H2 (29b)

C3H7COCH3 + H → C2H5CHCO + CH3 + H2 (29c)

C3H7COCH3 + H → n-C3H7 + CH2CO + H2 (29d)

In a similar way, the rate constants for channels (29a)-(29c) were assumed to be the same

as the rate constants for the reactions of H with methyl butanoate at the α, β and γ sites,

and these values were also adopted from the Dooley et al. mechanism [54]. Additionally,

the rate constant for channel (29d) was assumed to be the same as that of 2-butanone

(C2H5COCH3 + H → C2H5COCH2 + H2). Interestingly, the addition of reactions (29a)-

(29d) has negligible influence on the overall rate constant determinations of reaction (8).

Figure 6.11 shows a representative OH time history trace at 1186 K and 1.30 atm

using the mixture of 161 ppm 2-pentanone with 15 ppm TBHP (and 45 ppm H2O) in Ar.

The simulations from the modified Serinyel et al. mechanism with the best-fit rate

constant of k8 = 1.24×1013 cm3 mol-1 s-1 and the variations of ±50% in k8 were also

illustrated. A detailed error analysis was then conducted to estimate the overall

uncertainty in k8 at 1186 K, and the uncertainty was found to be ±24%.

107

Figure 6.11: Sample 2-pentanone + OH rate constant measurement using the mixture of 161 ppm 2-pentanone with ~15 ppm TBHP (and 45 ppm water) in Ar at 1186 K and 1.30 atm. Simulation from the modified Serinyel et al. mechanism for the best-fit rate constant, along with perturbations of ±50%, is also shown.

Table 6.4 summarizes the overall rate constant measurements of reaction (8) at

902-1302 K and 1.23-1.59 atm, and Figure 6.12 also presents the Arrhenius plot for these

measured values. These measured values are expressed in Arrhenius form as k8 =

7.06×1013 exp(-2020/T) cm3 mol-1 s-1 over 902-1302 K. It should be noted that the

overall rate constant measurements for 2-pentanone + OH reaction are quite similar to the

values obtained for 3-pentanone + OH reaction at our experimental conditions.

Additionally, Figure 6.12 presents the estimated overall rate constant for reaction (8)

using the group rate constants for the reactions of OH with ketones developed by Zhou et

al. [121], and the estimated values are in excellent agreement with the measurements

within 7%.

108

Table 6.4: C3H7COCH3 + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k8 [cm3 mol-1 s-1]

60 ppm TBHP (and water), 161 ppm C3H7COCH3, Ar 902 1.59 7.58E+12 1104 1.37 1.14E+13 1186 1.30 1.24E+13 1216 1.23 1.31E+13 1302 1.26 1.51E+13

63 ppm TBHP (and water), 161 ppm C3H7COCH3, Ar 955 1.50 8.45E+12 1009 1.46 9.38E+12 1042 1.43 1.04E+13 1093 1.37 1.12E+13 1125 1.30 1.17E+13 1264 1.25 1.48E+13

Figure 6.12: Arrhenius plot for 2-pentanone + OH (k8) at temperatures above 900 K.

109

6.3.6 Comparison of Ketone + OH Kinetics

Figure 6.13 presents the Arrhenius plot of the measured rate constants for

reactions (5)-(8) at temperatures above 870 K. Among all four ketone + OH reactions,

the reaction of OH with acetone has the lowest reactivity due to the fact that there are less

C–H bonds available for H-atom abstraction in acetone. As the number of C–H bonds in

the fuel molecule increases, the fuel + OH reaction becomes more reactive. As

demonstrated in Figure 6.13, the rate constant for 3-pentanone + OH is much faster than

that for 2-butanone + OH, because 3-pentanone has one more secondary carbon atom site

than 2-butanone. In addition, the rate constant for 2-pentanone + OH is about the same as

that for 3-pentanone + OH due to the fact that both molecules have the same number of

primary and secondary carbon atom sites. Note that the rate constant measurements of

the 2-pentanone + OH reaction are slightly higher than the values of the 3-pentanone +

OH reaction at T < 1100 K. This can be explained by the fact that the secondary carbon

atom site at the β position from 2-pentanone is supposed to be more reactive than that at

the α position from 3-pentanone, as suggested in previous experimental studies [38, 107].

Figure 6.13: Arrhenius plot of the measured rate constants for reactions (5)-(8) at temperatures above 870 K.

110

6.4 Comparison with Low Temperature Data

Figure 6.14 presents the current data along with some earlier measurements of

reaction (5) at temperatures greater than 250 K. In the study from Wollenhaupt et al.

[106], the rate constant for reaction (5) was measured over 202-395 K at 20-100 torr of

Ar or N2 bath gas using the pulsed laser photolysis technique to generate OH radicals

from the sequential two-photon dissociation of NO2 in the presence of H2 at 439 nm, or

from the photolysis of HONO at 351 nm. They monitored the OH radicals using either

resonance fluorescence or laser-induced fluorescence detection scheme, and they also

concluded that their measurements are independent of pressure. Similarly, Le Calvé et al.

[38] and Gierczak et al. [109] studied k5 over 199-383 K by generating OH via pulsed

laser photolysis and detecting it via laser-induced fluorescence, while Wallington and

Kurylo [107] investigated k5 over 240-440 K using the flash photolysis/resonance

fluorescence measurement technique. Moreover, Yamada et al. [117] examined k5 over a

wide temperature range 298-832 K using the pulsed laser photolysis/pulsed laser-induced

fluorescence technique. They then performed a detailed analysis using Variational

Transition State theory and suggested that the dominant products of reaction (5) are

CH3COCH2 and H2O through direct abstraction at all temperatures (particularly above

450 K). Additionally, Tranter and Walker [120] added small amounts of acetone to

slowly reacting mixtures of H2 + O2 at 753 K, and monitored the consumption of acetone

and H2 with the use of gas chromatography. This method allowed them to infer the

relative rate constant for reaction (5) at 753 K. It is pertinent to note that these low-

temperature measurements are in excellent agreement with each other. As is evident in

Figure 6.14, the rate constants employed in the comprehensive mechanisms of Pichon et

al. [89] and Herbinet et al. [45] are able to predict the low-temperature data reasonably

well over 298-832 K, but not for T < 298 K. In particular, the rate constant from

Herbinet et al. provides much better agreement with the existing high-temperature data.

In addition to the values from the detailed mechanisms, the theoretical calculation from

Zhou et al. [121] agrees well with earlier low-temperature measurements (at T < 500 K),

but the calculated values are at least 40% lower than the measurements at T > 500 K.

111

Figure 6.14: Arrhenius plot for acetone + OH → products (k5) at all temperatures.

Figure 6.15 shows the Arrhenius plot for the overall rate constant measurements

of reaction (6) at temperatures greater than 250 K. Kinetic measurements of reaction (6)

were performed at room temperature by different researchers using both relative [112-

115] and absolute [38, 107, 110-111] methods. In general, these room temperature

measurements are in close accord with each other, except for the value obtained from

Atkinson et al. [112]. Concurrently, the rate constant for reaction (6) was examined as a

function of temperature (213-598 K) by Wallington and Kurylo [107] using the flash

photolysis/resonance fluorescence technique and by Le Calvé et al. [38], Carr et al. [110]

and Jimenez et al. [111] using the pulsed laser photolysis/laser-induced fluorescence

technique. Their measurements are in excellent agreement with each other, and no

pressure dependence can be found at their experimental conditions. Based on these low-

temperature data, k6 exhibits only slight positive temperature dependence over 250-400

K. In addition to the rate constant determination for acetone + OH reaction, Tranter and

Walker [120] measured the relative rate constant for reaction (6) at 753 K. Figure 6.15

also presents the estimated values of k6 from Serinyel et al. [95] and the theoretical values

from Zhou et al. [121]. As described previously, the calculated values from Zhou et al.

are consistent with the present high-temperature data (at T > 879 K) within 10%.

However, the calculated values are faster than the earlier low-temperature data by a factor

112

of 2 at 500 K and by a factor of 6 at 250 K. Consequently, the theoretical study predicts a

pronounced negative temperature dependence of k6 over 250-500 K, and this effect does

not appear in the existing data. On the other hand, the estimated values from Serinyel et

al. are ~40% lower than the current high-temperature data, and are in good agreement

with the low-temperature data over 345-600 K. In addition, the overall rate constant from

Serinyel et al. does not exhibit any negative temperature dependence over 250-500 K.

Figure 6.15: Arrhenius plot for 2-butanone + OH → products (k6) at all temperatures.

Figure 6.16 also shows the current high-temperature data (at T > 878 K) and three

previous low-temperature measurements (at T < 800 K) for the reaction of OH with 3-

pentanone, along with the rate constant from the original Serinyel et al. mechanism [98].

As compared to acetone and 2-butanone, fewer experimental and theoretical studies are

available in the literature. Tranter and Walker [120] measured the relative rate constant

for reaction (7) at 753 K (using the same approach as the one they employed for reactions

(5) and (6)). Atkinson et al. [116] also measured the relative rate constant for reaction (7)

at 299 K using methyl nitrite (CH3ONO) photolysis in air as a source of OH radicals.

They monitored the organic reactants using gas chromatography with flame ionization

detection. In their work, they took advantage of their previous knowledge on the rate

constant for cyclohexane + OH reaction and inferred the rate constant for reaction (7)

113

from the ratio of k7/kcyclohexane+OH at 299 K. Moreover, Wallington and Kurylo [107]

determined the absolute rate constant for reaction (7) over 240-440 K using the flash

photolysis/resonance fluorescence measurement technique, and they suggested that k7 did

not exhibit any temperature dependence at their test conditions. Furthermore, the rate

constant from the original Serinyel et al. mechanism is able to predict the existing data

rather accurately over 440-1353 K.

Figure 6.16: Arrhenius plot for 3-pentanone + OH → products (k7) at all temperatures.

Similarly, Figure 6.17 illustrates the current high-temperature data and previous

low-temperature measurements [107, 111, 116] of reaction (8) over 250-1302 K. More

importantly, the rate constant provided by Jimenez et al. [111] exhibits a pronounced

negative temperature dependence over 248-388 K, and this trend does not appear in the

kinetic measurements of reactions (5)-(7). This pronounced negative temperature

dependence of k8 is mainly attributed to the H-atom abstraction from the CH2 group in

the β position, which is the predominant reaction pathway at low temperatures. For

instance, Jimenez et al. [111] determined that the branching ratios of channels (8a)-(8d)

are 0.04, 0.76, 0.18, and 0.02, respectively, at 298 K. Some researchers [107, 111] also

postulated that the H-abstraction reaction could proceed via an OH-addition complex,

114

resulting in a six- or seven-membered ring complex which enhances the abstraction of an

H atom from the CH2 group in the β position.

Figure 6.17: Arrhenius plot for 2-pentanone + OH → products (k8) at all temperatures.

6.5 Comparison with Structure-Activity Relationship

The measured overall rate constants for reactions (5)-(8) over 250-1360 K can be

compared with the estimated values using the structure-activity relationship (SAR)

developed by Atkinson and his co-workers [127-129]. Their method of calculating the

rate constants for the reactions of OH with organic compounds is based on the estimation

of primary (–CH3), secondary (–CH2–), and tertiary (–CH<) group rate constants, and

these group rate constants depend on the nature of the neighboring atoms (substituents

bound to the groups). The group rate constants can be expressed as k(CH3–X) = kprim

F(X), k(Y–CH2–X) = ksec F(X) F(Y), and k((Z)CH(X)(Y)) = ktert F(X) F(Y) F(Z), where

kprim, ksec and ktert are the rate constants for the H-atom abstraction from –CH3, –CH2– and

–CH< groups, and F(X), F(Y) and F(Z) are the substituent factors. Recently, Pang et al.

[74] have demonstrated that the SAR estimation accurately predicts the measured rate

constants for n-alkane + OH reactions (i.e., n-pentane + OH, n-heptane + OH, and n-

115

nonane + OH) over 250-1364 K. In particular, the SAR estimation captures the

temperature dependence of their measurements reasonably well. In addition, Kwok and

Atkinson [129] provided a revised list of the substituent factors F(X) at 298 K, and they

assumed that the temperature dependence of the substituent factors can be expressed in

the form of F(X) = exp(Ex/T). In the present analysis, the Ex term in the preceding

expression was calculated from the substituent factor at 298 K, thereby allowing us to

determine the substituent factors at different temperatures.

The estimated rate constants for reactions (5)-(8) based on the SAR approach are

provided in Figures 6.14-6.17. For instance, the estimated rate constant for 2-pentanone

+ OH reaction can be evaluated as k8 = kprim F(–CH2–) + ksec F(CH3–) F(–CH2C(O)R) +

ksec F(–CH2–) F(–C(O)–) + kprim F(–CO–). It is pertinent to note that the substituent

factor F(–CH2C(O)R) is 3.9 at 298 K and is much higher than F(–CH2–), which is 1.23 at

298 K. This confirms the increased reactivity of the H-atom abstraction at the β position,

as observed by Atkinson et al. [116]. As expected, the SAR estimation shows good

agreement with the kinetic measurements of reactions (5)-(8) at 298 K, but the estimated

values are higher than the measured values over 333-1360 K. In particular, the estimated

values are ~25% faster than the present high-temperature data over 870-1360 K.

Nevertheless, the SAR estimation can capture the temperature dependence of reactions

(5)-(8) reasonably well, implying that the pre-exponential factors for the group rate

constants kprim and ksec should be reduced by ~25%, particularly for ketone + OH

reactions. With this modification, the estimated rate constants for reactions (5)-(7) show

excellent agreement with the measurements over a wide temperature range 298-1360 K

(see Figures 6.14-6.16). In addition, the modified SAR estimation for reaction (8)

precisely predicts the present high-temperature data (at 902-1302 K), and is in close

accord with the previous data from Wallington and Kurylo [107] and Jimenez et al. [111]

at temperatures near 298 K, as seen in Figure 6.17.

116

6.6 Summary

The overall rate constants for the reactions of OH with acetone (k5), 2-butanone

(k6), 3-pentanone (k7) and 2-pentanone (k8) were studied behind reflected shock waves

over 870-1360 K at pressures of 1-2 atm using OH laser absorption. The present high-

temperature measurements can be expressed in Arrhenius form as:

k5 = 3.30×1013 exp(-2437/T) cm3 mol-1 s-1

k6 = 6.35×1013 exp(-2270/T) cm3 mol-1 s-1

k7 = 9.29×1013 exp(-2361/T) cm3 mol-1 s-1

k8 = 7.06×1013 exp(-2020/T) cm3 mol-1 s-1

Detailed error analyses, which account for both experimental and secondary chemistry

contributions, yielded the uncertainty estimates of ±28% at 1148 K for k5, ±22% at 1039

K for k6, ±20% at 1188 K for k7, and ±24% at 1186 K for k8. In addition, the structure-

activity relationship (SAR) from Atkinson and his co-workers [127-129] was used to

estimate the rate constants for reactions (5)-(8), and the estimated values are in good

agreement with the present high-temperature data (within ~25%).

117

Chapter 7 High-Temperature Measurements of the Reactions of OH with Small Methyl Esters: Methyl Formate, Methyl Acetate, Methyl Propanoate, and Methyl Butanoate

7.1 Introduction

In this chapter, the overall rate constants for the reactions of OH with four small

methyl esters, namely methyl formate (CH3OCHO), methyl acetate (CH3OC(O)CH3),

methyl propanoate (CH3OC(O)C2H5), and methyl butanoate (CH3OC(O)C3H7), were

determined behind reflected shock waves over the temperature range of 876-1371 K at

pressures near 1.5 atm:





We believe these are the first direct high-temperature measurements of the overall rate

constants for reactions (9)-(12). These kinetic data were compared with the values

adopted in several detailed kinetic mechanisms and the estimates using the structure-

activity relationship (SAR) developed by Atkinson and co-workers [127-129].

118


Test mixtures were prepared manometrically in a 40 liter stainless-steel tank



manometrical preparation of a highly dilute mixture. A more concentrated mixture was



the experiments. The gas utilized in this study was argon (Research Grade) 99.999%,

which was supplied by Praxair and used without further purification. The liquid

chemicals were commercially available 70% tert-butyl hydroperoxide (TBHP) in water,

methyl formate (≥99%), methyl acetate (≥99%), methyl propanoate (≥99%), and methyl

butanoate (≥99%) from Sigma-Aldrich, and were purified using a freeze-pump-thaw

procedure to remove dissolved volatiles and air prior to mixture preparation.

The mixture composition was confirmed by sampling a portion of the mixture



laser at 3.39 µm [82, 122]. Beer’s law was then used to convert the measured absorption

data into the fuel mole fraction. The absorption cross-sections of methyl esters for Beer’s

law were directly obtained from the PNNL database [123], and the measured fuel

concentrations were consistent with the values expected from the manometrical

preparation within ±5%.


A total of 52 reflected shock wave experiments were performed to determine the

overall rate constants for the reactions of OH with four methyl esters (methyl formate,

methyl acetate, methyl propanoate, and methyl butanoate) over 876-1371 K at pressures

near 1.5 atm. Experiments were carried out using different initial fuel concentrations:

methyl formate (322 ppm, 404 ppm), methyl acetate (323 ppm, 384 ppm), methyl

119

propanoate (~281 ppm), and methyl butanoate (241 ppm, 270 ppm). Test mixtures with

individual methyl esters and 80-102 ppm TBHP (and water) diluted in argon were

utilized in the present study. Note that dilute mixtures were preferred in order to

minimize the temperature change resulted from the chemistry effects, and the temperature

profile behind the reflected shock wave (from the present study) was nearly constant (less

than 1 K change based on the calculation from CHEMKIN PRO [71]) over the time

frame of the experiment (the first 100 µs).

7.3.1 Choice of Kinetic Mechanisms

The CHEMKIN PRO package [71] was used to simulate the OH time histories

under the standard constant energy and volume assumption. A comprehensive chemical

kinetic mechanism of Dooley et al. [49] was chosen as the base mechanism for methyl

formate and methyl acetate. This mechanism can successfully simulate shock tube

ignition delay times, laminar burning velocities of outwardly propagating spherical

flames, and speciation data from a shock tube and a variable-pressure flow reactor [49-

50] during methyl formate pyrolysis and oxidation. Additionally, this kinetic mechanism

incorporates the sub-mechanism for methyl acetate, which was previously developed by

Westbrook et al. [53]. The sub-mechanism for methyl acetate consists of the

unimolecular decomposition pathways and the H-atom abstraction reactions by H, OH,

and CH3 radicals, and was validated against speciation data from fuel-rich, low-pressure,

premixed laminar flames. A detailed kinetic mechanism of Dooley et al. [54] was also

selected as the base mechanism for methyl propanoate and methyl butanoate. This

mechanism was originally developed to predict the autoignition of methyl butanoate in a

shock tube and a rapid compression machine over a wide range of experimental

conditions, and was further validated against speciation data available in the literature

from a flow reactor, a jet-stirred reactor, and an opposed-flow diffusion flame. As were

done in Chapters 3 and 6, tert-butyl hydroperoxide (TBHP or (CH3)3−CO−OH) was used

as an OH radical precursor at the present experimental conditions, and the TBHP sub-

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mechanism was also implemented into the base mechanisms for these methyl ester + OH

studies. (Please read Chapter 3 for more details on the TBHP chemistry.)

7.3.2 Methyl Formate (MF) + OH Kinetics

The reaction of OH with methyl formate consists of 2 different channels:

CH3OCHO + OH → CH3OCO + H2O (9a)

CH3OCHO + OH → CH2OCHO + H2O (9b)

The branching ratios of channels (9a) and (9b) are 0.32 and 0.68, respectively, at 1168 K,

based on the Dooley et al. mechanism [49]. In their analysis, the estimated rate constant

for channel (9a) was assumed to be an intermediate value between typical primary and

secondary C–H bonds (as in propane) due to the weaker bond strength of the CH3OCO–H

position (100.1 kcal/mol at 298 K). Similarly, the estimated rate constant for channel

(9b) (per H-atom) was assumed to be 5% faster than the value for a typical primary C–H

bond, and the corresponding bond strength was estimated to be 100.9 kcal/mol at 298 K.

An OH radical sensitivity analysis for the mixture of 322 ppm methyl formate

with 26 ppm TBHP (and 70 ppm H2O) in Ar at 1168 K and 1.40 atm is shown in Figure

7.1. The analysis reveals that the reaction of OH with methyl formate (reaction (9)) is the

dominant reaction over the time frame of the experiment, with some minor interference

from the secondary reactions:

CH3 + OH → CH2(s) + H2O (17)

C2H6 (+ M) → CH3 + CH3 (+ M) (18)

CH2O + OH → HCO + H2O (30)

In this modeling, the rate constants for reactions (5) and (17)-(19) in the Dooley et al.

mechanism [49] were updated with the values in Table 3.1. The rate constant for reaction

(30) was previously measured using UV laser absorption of OH near 307 nm behind

reflected shock waves over 934-1670 K at pressures near 1.6 atm by Vasudevan et al.

[126], and their measured rate constant was also adopted in the present study.

121

Figure 7.1: OH sensitivity plot for the rate constant measurement of methyl formate + OH at 1168 K and 1.40 atm.

Figure 7.2 illustrates a sample measured OH concentration time history for the

mixture of 322 ppm methyl formate in Ar at 1168 K and 1.40 atm, and the measured peak

OH concentration is approximately 26 ppm. According to the measured peak OH yields,

the mixtures with 96 ppm TBHP/water are comprised of ~25-28 ppm TBHP in the

present study. It should also be noted that the presence of H2O in the test mixtures does

not have any significant influence on the computed OH profiles. As shown in Figure 7.2,

a best-fit overall rate constant for reaction (9) of 4.26×1012 cm3 mol-1 s-1 was obtained

between the experimental data and the simulation at 1168 K and 1.40 atm. The

simulations for the perturbations of ±50% in the inferred rate constant are also shown in

Figure 7.2. Additionally, the effect of the branching ratios for reaction (9) on the overall

rate constant determination was tested at 1168 K by interchanging the branching ratios of

channels (9a) and (9b) while maintaining the overall value, and no discernible effect

could be observed from the simulated OH profiles. Hence, the original branching ratios

proposed by Dooley et al. [49] were kept in our simulations. In addition, Table 7.1

summarizes the overall rate constant measurements (k9 = k9a + k9b) of reaction (9) at T =

880-1344 K and P = 1.24-1.63 atm.

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Figure 7.2: Sample methyl formate + OH rate constant measurement using the mixture of 322 ppm methyl formate with ~26 ppm TBHP (and 70 ppm water) in Ar at 1168 K and 1.40 atm. Simulation from the Dooley et al. mechanism [49] for the best-fit rate constant, along with perturbations of ±50%, is also shown.

Table 7.1: CH3OCHO + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k9 [cm3 mol-1 s-1]

96 ppm TBHP (and water), 322 ppm CH3OCHO, Ar 1337 1.36 5.87E+12 1315 1.25 5.60E+12 1264 1.28 4.98E+12 1229 1.37 4.81E+12 1168 1.40 4.26E+12 1114 1.39 3.99E+12 1024 1.51 3.50E+12 965 1.55 3.10E+12 913 1.63 2.81E+12 904 1.55 2.91E+12

101 ppm TBHP (and water), 404 ppm CH3OCHO, Ar 1344 1.27 5.85E+12 1289 1.24 5.62E+12 1124 1.37 4.04E+12 1060 1.44 3.66E+12 880 1.62 2.57E+12

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A detailed error analysis was conducted to estimate the overall uncertainty of the

measured rate constant for reaction (9) at 1168 K. The primary contributions to the

overall uncertainty in k9 were considered: (a) temperature (±1%), (b) mixture

composition (±5%), (c) OH absorption coefficient (±3%), (d) wavemeter reading in the

UV (±0.01 cm-1), (e) fitting the data to the simulated profiles (±5%), (f) locating time-

zero (±0.5 µs), (g) the rate constant for CH3 + OH → CH2(s) + H2O (uncert. factor = 2),

(h) the rate constant for CH2O + OH → HCO + H2O (uncert. factor = 2), and (i) the rate

constant for C2H6 (+ M) → CH3 + CH3 (+ M) (±20%). As demonstrated in Figure 7.3,

the individual error sources were introduced separately (within the positive and negative

bounds of their 2σ uncertainties) and their effects on the overall rate constant for reaction

(9) were studied. These uncertainties were combined in a root-sum-squared method to

give an overall (2σ) uncertainty of ±24% at 1168 K. Similar error analyses were

performed for k9 at 913 K and 1289 K, and the overall uncertainties were estimated to be

±29% and ±18%, respectively.

Figure 7.3: Uncertainty analysis for the rate constant of methyl formate + OH → products at 1168 K and 1.40 atm.

Figure 7.4 presents the Arrhenius plot for the overall rate constant measurements

of reaction (9) at T = 880-1344 K, along with the estimated values proposed by Fisher et

124

al. [48] and Dooley et al. [49]. Note that two different mixture compositions (322 ppm

and 404 ppm methyl formate) were used to confirm that the current measurements are

weakly dependent on the secondary chemistry effects from the model, and the measured

values from these two mixtures are consistent with each other. The measured values can

be expressed in Arrhenius form as k9 = 2.56×1013 exp(-2026/T) cm3 mol-1 s-1 over 880-

1344 K. As is evident in Figure 7.4, the present measurements are in good agreement

with the estimated values from Fisher et al. and Dooley et al. within 10%, and the

activation energy from the present measurements seems to be slightly higher.

Interestingly, the estimated values from Fisher et al. and Dooley et al. are nearly identical

over 833-1150 K and start to deviate at higher temperatures (T > 1150 K). It also appears

that the estimated overall rate constant from Fisher et al. is in better agreement with the

present measurements at T > 1150 K.

Recently, Tan et al. [130] performed a systematic ab initio quantum mechanical

investigation of the H-atom abstraction reactions for methyl formate by five radicals: H,

CH3, O, HO2, and OH. They employed a multi-reference correlated wave function

method (the CBS-MRSDCI composite scheme) including size-extensivity corrections to

calculate the barrier heights and reaction enthalpies of these H-atom abstraction reactions.

The rate constants for these H-atom abstraction reactions were computed using transition

state theory within the separable-hindered-rotor approximation for torsions and the

harmonic oscillator approximation for other vibrational modes [130]. As illustrated in

Figure 7.4, the calculated overall rate constant for reaction (9) from Tan et al. [130] is

substantially lower than the present measurements and the estimated values from Fisher

et al. [48] and Dooley et al. [49] (by a factor of 2.2 at 1250 K and a factor of 4.8 at 850

K). To investigate this large discrepancy between the present measurements and the

theoretical calculation, the rate constants for the reactions of H, CH3, O, and HO2 with

methyl formate in the Dooley et al. mechanism [49] were first updated with the

expressions provided by Tan et al. [130]. The overall rate constant for reaction (9) was

then reexamined by matching the measured OH time histories with the simulated profiles

from the detailed mechanism, and the same measured rate constant expression (as the one

provided previously) was obtained. This indicates that our rate constant measurements

125

are insensitive to the H-atom abstraction reactions for methyl formate by other species

(i.e., H, CH3, O, and HO2), and also that the theoretical calculations for these H-atom

abstraction reactions may require further review.

Figure 7.4: Arrhenius plot for methyl formate + OH (k9) at temperatures above 833 K.

7.3.3 Methyl Acetate (MA) + OH Kinetics

The reaction of OH with methyl acetate consists of 2 different channels:

CH3OC(O)CH3 + OH → CH3OC(O)CH2 + H2O (10a)

CH3OC(O)CH3 + OH → CH2OC(O)CH3 + H2O (10b)

The branching ratios of channels (10a) and (10b) are 0.14 and 0.86, respectively, at 1091

K, based on the estimated values from the Dooley et al. mechanism [49]. Note that the

sub-mechanism for methyl acetate adopted in the Dooley et al. mechanism was

previously developed by Westbrook et al. [53]. In the development of the methyl acetate

sub-mechanism from Westbrook et al., the rate constant for channel (10b) (the H-atom

abstraction from the methyl group bound to the O-atom in the ester group) was taken

directly from that of the structurally similar methyl group in methyl butanoate (developed

by Fisher et al. [48]). In addition, the bond strength of the C–H bond adjacent to the

carbonyl group is 97.7 kcal/mol (at 298 K), which is similar to that of a tertiary C–H

126

bond in methylcyclohexane. Thus, the rate constant for channel (10a) (per H-atom) was

first assumed to be the same as the rate constant for the tertiary C–H bond in

methylcyclohexane. Channel (10a) produces the CH3OC(O)CH2 radical, followed by the

formation of ketene (CH2CO) and methoxy radical (CH3O) via β-scission. Concurrently,

the methoxy radical can react with CH3 to form dimethyl ether. When compared with

their flame measurements [53], the model predicted excessively high levels of ketene and

dimethyl ether, which suggested the need to reduce the rate constant for channel (10a).

Consequently, the rate constant for channel (10a) was reduced by a factor of 10 to match

their experimental data.

The OH sensitivity analysis was performed for the overall rate constant

determination (k10 = k10a + k10b) of reaction (10) using the mixture of 384 ppm methyl

acetate with 28.5 ppm TBHP (and 73.5 ppm water) diluted in argon at 1091 K and 1.37

atm. As illustrated in Figure 7.5, the analysis shows that reaction (10) is the dominant

reaction over the time frame of the experiment, with some minor interference from

reactions (13), (17), (18), and (30). Note that the TBHP decomposition reaction (reaction

(13)) becomes more important at the early times as temperature decreases.

Figure 7.5: OH sensitivity plot for the rate constant measurement of methyl acetate + OH at 1091 K and 1.37 atm.

127

Figure 7.6 shows a sample OH time history measurement for the mixture of 384

ppm methyl acetate in argon at 1091 K and 1.37 atm, and the measured peak OH mole

fraction is ~28.5 ppm. Thus, we inferred that the initial TBHP mole fraction was around

28.5 ppm, and there was approximately 73.5 ppm H2O. As illustrated in Figure 7.6, a

best-fit overall rate constant for reaction (10) of 3.93×1012 cm3 mol-1 s-1 was obtained

between the experiment and the simulation using the Dooley et al. mechanism [49]. In

addition, the simulations with the variations of ±50% in the inferred rate constant are

shown in Figure 7.6. The same test (as the test for methyl formate) was performed at

1091 K to confirm that the branching ratios have negligible influence on the overall rate

constant determination of reaction (10). Thus, the original branching ratios from

Westbrook et al. [53] were maintained in our simulations. Table 7.2 summarizes the

present overall rate constant measurements of reaction (10) at T = 876-1371 K and P =

1.25-1.60 atm.

Figure 7.6: Sample methyl acetate + OH rate constant measurement using the mixture of 384 ppm methyl acetate with ~28.5 ppm TBHP (and 73.5 ppm water) in Ar at 1091 K and 1.37 atm. Simulation from the Dooley et al. mechanism [49] for the best-fit rate constant, along with perturbations of ±50%, is also shown.

128

Table 7.2: CH3OC(O)CH3 + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k10 [cm3 mol-1 s-1]

100 ppm TBHP (and water), 323 ppm CH3OC(O)CH3, Ar 1371 1.25 5.88E+12 1258 1.29 5.09E+12 1160 1.36 4.33E+12 1078 1.44 3.74E+12 1028 1.50 3.41E+12

102 ppm TBHP (and water), 384 ppm CH3OC(O)CH3, Ar

1299 1.27 5.39E+12 1215 1.34 4.91E+12 1126 1.36 4.40E+12 1091 1.37 3.93E+12 1017 1.43 3.29E+12 961 1.54 2.82E+12 930 1.58 2.63E+12 907 1.59 2.38E+12 876 1.60 2.18E+12

Figure 7.7 presents the Arrhenius plot for the current overall rate constant

measurements of reaction (10) over the temperature range of 876-1371 K, along with the

estimated values proposed by Westbrook et al. [53]. Note that the measured values from

two different mixture compositions (323 ppm and 384 ppm methyl acetate) were

compared and were found to be consistent with each other. These measured values can

be expressed in Arrhenius form as k10 = 3.59×1013 exp(-2438/T) cm3 mol-1 s-1 over 876-

1371 K. Detailed error analyses were carried out with the consideration of experimental

and mechanism-induced contributions, and the overall (2σ) uncertainties in k10 were

estimated to be ±29% at 930 K, ±23% at 1091 K, and ±17% at 1299 K. As illustrated in

Figure 7.7, the activation energy of reaction (10) inferred from the present measurements

is higher than that of Westbrook et al. [53]. The estimated value from Westbrook et al. is

approximately 30% lower than the current data at 1371 K, while the estimated values are

in good agreement with the current data (within 8%) over a limited temperature range of

876-960 K.

129

Moreover, theoretical studies of the reactions of OH with ethers (dimethyl, ethyl

methyl, and isopropyl methyl ethers) and ketones (dimethyl, ethyl methyl, and isopropyl

methyl ketones) were performed by Zhou et al. [121, 131] using the computationally less-

expensive methods of G3 and G3MP2BH&H to calculate the energy barriers and using

the Variflex code including Eckart tunneling corrections to compute the total rate

constants over 500-2000 K. They also provided the expressions of the group rate

constants (per H-atom) for three different carbon sites (primary, secondary, and tertiary

carbon atoms) adjacent to the ether group (–O–) and the carbonyl group (–C(O)–). In the

present analysis, we can estimate the overall rate constant for reaction (10) using the

group rate constants provided by Zhou et al. [121, 131], and the estimated overall rate

constant is k10 = 3 × k(CH3O) + 3 × k(CH3C(O)), where k(CH3O) and k(CH3C(O)) are

the group rate constants (per H-atom) for primary carbon sites adjacent to the ether group

and the carbonyl group, respectively. As illustrated in Figure 7.7, the estimated values

are at least 60% higher than the present measurements, but the estimation seems to

capture the temperature dependence of reaction (10) reasonably well.

Figure 7.7: Arrhenius plot for methyl acetate + OH (k10) at temperatures above 833 K.

130

7.3.4 Methyl Propanoate (MP) + OH Kinetics

The reaction of OH with methyl propanoate is comprised of 3 different channels:

CH3OC(O)C2H5 + OH → C2H4 + CH3OCO + H2O (11a)

CH3OC(O)C2H5 + OH → CH3CHCO + CH3O + H2O (11b)

CH3OC(O)C2H5 + OH → C2H5CO + CH2O + H2O (11c)

Channel (11a) is the H-atom abstraction reaction from methyl propanoate at the β

position (on the same side of the carbonyl group), and channel (11b) is the H-atom

abstraction reaction from methyl propanoate at the α position. In addition, channel (11c)

is the H-atom abstraction reaction from methyl propanoate at the methyl group bound to

the O-atom in the ester group. Note that the Dooley et al. mechanism of NUI Galway

[54], which was originally developed for methyl butanoate oxidation, was used to model

the OH consumption from methyl propanoate. Unfortunately, the mechanism does not

contain a sub-mechanism for methyl propanoate. In the present analysis, we assumed

that the fuel radicals formed right after the H-atom abstraction reactions would

decompose immediately into the (relatively) stable products through β-scission. For

instance, channel (11a) forms a CH3OC(O)CH2CH2 radical and a H2O molecule, and the

CH3OC(O)CH2CH2 radical is rather short-lived and will further decompose to form C2H4

and CH3OCO through β-scission. Similar treatments were applied to channels (11b) and

(11c). Due to the structure similarity between methyl propanoate and ethyl propanoate,

the rate constants for channels (11a) and (11b) were first assumed to be the same as the

rate constants for the reactions of OH with ethyl propanoate at the β and α sites,

respectively, which were taken from the Metcalfe et al. mechanism of NUI Galway [132].

In addition, the rate constant for channel (11c) was first assumed to be the same as the

rate constant for the reaction of OH with methyl butanoate at the methyl group bound to

the O-atom in the ester group, which was taken directly from the Dooley et al.

mechanism [54]. The resulting branching ratios of channels (11a)-(11c) at 1208 K are

0.20, 0.31, and 0.49, respectively. These 3 channels, along with their corresponding rate

constants, were then incorporated into the Dooley et al. mechanism [54].

131

The reaction of H-atom with methyl propanoate was also considered in the

present study. Similar to the reaction of OH with methyl propanoate, it consists of 3

different channels:

CH3OC(O)C2H5 + H → C2H4 + CH3OCO + H2 (31a)

CH3OC(O)C2H5 + H → CH3CHCO + CH3O + H2 (31b)

CH3OC(O)C2H5 + H → C2H5CO + CH2O + H2 (31c)

Channel (31a) describes the H-atom abstraction at the β position, and channel (31b)

describes the H-atom abstraction at the α position. In addition, channel (31c) describes

the H-atom abstraction at the methyl group bound to the O-atom in the ester group.

Similarly, the rate constants for channels (31a) and (31b) were assumed to be the same as

the rate constants for the reactions of H with ethyl propanoate at the β and α sites,

respectively, which were also taken from Metcalfe et al. [132]. Additionally, the rate

constant for channel (31c) was assumed to be the same as the rate constant for the

reaction of H with methyl butanoate at the methyl group bound to the O-atom in the ester

group, which was also taken from Dooley et al. [54]. It is important to note that the

simulated OH profiles of the present study are effectively insensitive to channels (31a)-

(31c); hence, the computed OH profiles are nearly identical with and without the addition

of these 3 channels. This conclusion is expected as there are very few H-atoms available

in the initial test mixtures. Nevertheless, channels (31a)-(31c) were included in the

Dooley et al. mechanism [54] for completeness.

The OH sensitivity analysis was performed for the overall rate constant

determination (k11 = k11a + k11b + k11c) of reaction (11) using the mixture of 281 ppm

methyl propanoate with 22 ppm TBHP (and 68 ppm H2O) in Ar at 1208 K and 1.33 atm.

As demonstrated in Figure 7.8, the analysis reveals that the OH time history is

predominantly sensitive to reaction (11) over the time frame of the experiment. There is

also some minor interference from the secondary reactions (reactions (17), (19), and

(30)).

132

Figure 7.8: OH sensitivity plot for the rate constant measurement of methyl propanoate + OH at 1208 K and 1.33 atm.

Figure 7.9 shows a representative OH time history measurement for the mixture

of 281 ppm methyl propanoate in Ar at 1208 K and 1.33 atm, and the measured peak OH

mole fraction is ~22 ppm. Thus the initial TBHP mole fraction was approximately 22

ppm and the initial water mole fraction was around 68 ppm. As is evident in Figure 7.9,

a best-fit overall rate constant for reaction (11) of 8.01×1012 cm3 mol-1 s-1 was used to

match the experimental data with the computed profile, and the simulations for the

variations of ±50% in the inferred rate constant are also shown. Concurrently, the effect

of the branching ratios on k11 was found to be negligible at 1208 K (by interchanging the

branching ratios of channels (11b) and (11c) while maintaining the total value). Thus, the

original branching ratios based on the structure similarity were kept in our simulations.

Moreover, Diévart et al. [46] have recently developed a methyl propanoate sub-

mechanism, which includes the unimolecular decomposition reactions and the H-atom

abstraction reactions for methyl propanoate. This sub-mechanism was also implemented

into the Dooley et al. mechanism [54], and the thermodynamic parameters for the fuel

radicals from methyl propanoate (provided by Diévart et al.) were also added to the

thermo database of Dooley et al. Interestingly, near-identical results were found with and

without the use of the detailed sub-mechanism for methyl propanoate. Hence, the present

measurements are insensitive to the secondary chemistry effects strictly from methyl

133

propanoate. In addition, Table 7.3 summarizes the overall rate constant measurements of

reaction (11) over the temperature range of 909-1341 K at pressures of 1.23-1.58 atm.

Figure 7.9: Sample methyl propanoate + OH rate constant measurement using the mixture of 281 ppm methyl propanoate with ~22 ppm TBHP (and 68 ppm water) in Ar at 1208 K and 1.33 atm. Simulation from the Dooley et al. mechanism [54] for the best-fit rate constant, along with variations of ±50%, is also shown.

Table 7.3: CH3OC(O)C2H5 + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k11 [cm3 mol-1 s-1] 90 ppm TBHP (and water), 281 ppm CH3OC(O)C2H5, Ar

1341 1.26 1.01E+13 1208 1.33 8.01E+12 1124 1.39 6.70E+12 1049 1.36 5.81E+12 954 1.58 4.76E+12

91 ppm TBHP (and water), 283 ppm CH3OC(O)C2H5, Ar 1289 1.26 9.58E+12 1252 1.27 8.67E+12 1200 1.23 8.08E+12 1181 1.30 7.79E+12 1016 1.44 5.37E+12 909 1.52 4.11E+12

134

Figure 7.10 presents the Arrhenius plot for the present overall rate constant


estimated values used by Diévart et al. [46]. The measured values can be expressed in

Arrhenius form as k11 = 6.65×1013 exp(-2539/T) cm3 mol-1 s-1 over 909-1341 K. Detailed

error analyses were conducted with the consideration of experimental and mechanism-

induced contributions, and the overall (2σ) uncertainties in k11 were estimated to be

±25% at 909 K, ±21% at 1208 K, and ±17% at 1341 K. It is also interesting to note that

the estimated rate constants for channels (11a)-(11c) provided by Diévart et al. are

exactly identical to our initial approximations for these three rate constants. Thus,

Diévart et al. employed the same type of approximation to estimate the rate constants for

the H-atom abstraction reactions based on the structure similarity between methyl

propanoate and ethyl propanoate. However, the estimated value is ~53% higher than the

measured value at 1341 K, and is higher than the data by at least a factor of 2 at 909 K.

Additionally, the activation energy of reaction (11) inferred from the present

measurements is higher than that of the estimation from Diévart et al. This demonstrates

the importance of direct rate constant measurements in validating the current estimation

methods in the literature. Moreover, Figure 7.10 presents the estimated overall rate

constant for reaction (11) using the group rate constants for the reactions of OH with

ethers and ketones provided by Zhou et al. [121, 131]. Interestingly, the estimated values

are at least 65% higher than the current data, but the estimation appears to capture the

temperature dependence of reaction (11) reasonably well.

135

Figure 7.10: Arrhenius plot for methyl propanoate + OH (k11) at temperatures above 870 K.

7.3.5 Methyl Butanoate (MB) + OH Kinetics

The reaction of OH with methyl butanoate consists of 4 different channels, which

are:

CH3OC(O)C3H7 + OH → CH3OC(O)CH2CH2CH2 + H2O (12a)

CH3OC(O)C3H7 + OH → CH3OC(O)CH2CHCH3 + H2O (12b)

CH3OC(O)C3H7 + OH → CH3OC(O)CHCH2CH3 + H2O (12c)

CH3OC(O)C3H7 + OH → CH2OC(O)C3H7 + H2O (12d)

Channels (12a)-(12c) describe the H-atom abstraction from methyl butanoate at the γ, β,

and α positions, respectively. In addition, channel (12d) describes the H-atom

abstraction from methyl butanoate at the methyl group bound to the O-atom in the ester

group. On the basis of the Dooley et al. mechanism [54], the resulting branching ratios of

channels (12a)-(12d) are 0.15, 0.24, 0.25, and 0.36, respectively, at 1133 K. In their

analysis, the rate constants for channels (12a)-(12d) were estimated based on the nature

of the C–H bonds (primary, secondary, or tertiary carbon atoms), and were strictly based

on a study of methylcyclohexane (MCH) oxidation from Orme et al. [133]. The rate

constant for channel (12a) was assumed to be the same as the rate constant for the

136

reaction of MCH + OH → CYCHEXCH2 + H 2O, and the rate constant for channel (12b)

was assumed to be the same as the rate constant for the reaction of MCH + OH → MCH -

R4 + H2O. The chemical notations for the fuel radicals formed from the reactions of OH

with MCH are taken directly from Orme et al. [133], as illustrated in Figure 7.11. In

addition, due to the weaker C–H bond enthalpy, the carbon atom attached to the carbonyl

group (at the α position) was treated as a tertiary carbon atom, and the corresponding rate

constant for channel (12c) (per H-atom) was taken from the rate constant for the reaction

of MCH + OH → MCH -R1 + H2O. Similarly, the carbon atom at the methyl group

bound to the O-atom in the ester group was treated as a secondary carbon atom, and the

corresponding rate constant for channel (12d) (per H-atom) was taken from the rate

constant for the reaction of MCH + OH → MCH-R4 + H2O (per H-atom).

Figure 7.11: Chemical notations for fuel radicals from MCH + OH reactions used by Orme et al. [133].

Figure 7.12 shows the OH sensitivity analysis for the mixture of 241 ppm methyl

butanoate with 20 ppm TBHP (and 60 ppm H2O) in argon at 1133 K and 1.37 atm.

Similarly, the analysis reveals that reaction (12) is the dominant reaction pathway over

the time frame of the experiment, with some minor interference from the secondary

reactions (reactions (17)-(19)).

137

Figure 7.12: OH sensitivity plot for the rate constant measurement of methyl butanoate + OH at 1133 K and 1.37 atm.

Figure 7.13 shows a sample OH time history measurement for the mixture of 241

ppm methyl butanoate in Ar at 1133 K and 1.37 atm, and the measured peak OH mole

fraction is ~20 ppm. Based on the measured peak OH yield, the initial TBHP mole

fraction was ~20 ppm and the initial H2O mole fraction was around 60 ppm. As

illustrated in Figure 7.13, a best-fit overall rate constant (k12 = k12a + k12b + k12c + k12d) of

1.17×1013 cm3 mol-1 s-1 was used to match the experimental data with the simulated

profile from the Dooley et al. mechanism [54]. Additionally, the simulations for the

variations of ±50% in the inferred rate constant are also shown. Similar to the previous

methyl esters, the branching ratios of reaction (12) have negligible influence on the

overall rate constant determination at the present experimental conditions. This was also

confirmed by interchanging the branching ratios of channels (12a) and (12d) while

maintaining the total value at 1133 K. Moreover, Table 7.4 summarizes the present

overall rate constant measurements of reaction (12) over the temperature range of 897-

1355 K at pressures of 1.23-1.59 atm.

138

Figure 7.13: Sample methyl butanoate + OH rate constant measurement using the mixture of 241 ppm methyl butanoate with ~20 ppm TBHP (and 60 ppm water) in Ar at 1133 K and 1.37 atm. Simulation from the Dooley et al. mechanism [54] for the best-fit rate constant, along with variations of ±50%, is also shown.

Table 7.4: CH3OC(O)C3H7 + OH → Products: Rate Constant Data.

T5 [K] P5 [atm] k12 [cm3 mol-1 s-1] 80 ppm TBHP (and water), 241 ppm CH3OC(O)C3H7, Ar

1303 1.27 1.65E+13 1225 1.30 1.40E+13 1181 1.32 1.30E+13 1133 1.37 1.17E+13 1016 1.48 9.38E+12 897 1.59 6.85E+12

85 ppm TBHP (and water), 270 ppm CH3OC(O)C3H7, Ar

1355 1.23 1.87E+13 1320 1.23 1.71E+13 1262 1.27 1.56E+13 1062 1.44 1.02E+13 961 1.53 8.50E+12 925 1.57 7.74E+12

139

Figure 7.14 shows the Arrhenius plot for the present overall rate constant


estimated values from Fisher et al. [48], Dooley et al. [54], and Hakka et al. [55]. Note

that two different mixture compositions were employed to verify that the current rate

constant evaluations are weakly dependent on the secondary chemistry effects, and the

measured values from these two mixtures agree well with each other. These measured

values can be expressed in Arrhenius form as k12 = 1.13×1014 exp(-2515/T) cm3 mol-1 s-1

over 897-1355 K. Similar detailed error analyses were also carried out with the

consideration of experimental and mechanism-induced contributions, and the overall (2σ)

uncertainties in k12 were estimated to be ±24% at 925 K, ±20% at 1133 K, and ±16% at

1320 K. As is evident in Figure 7.14, the estimated rate constants adopted in three

different detailed mechanisms are quite different from each other. In particular, the

estimated value from Fisher et al. [48] is lower than the values from Dooley et al. [54]

and Hakka et al. [55] by approximately 40% and 83%, respectively, at 1133 K.

Interestingly, the temperature dependence of these rate constants seems to be consistent

with each other.

Table 7.5 shows a comparison of the rate constants for channels (12a)-(12d)

employed in these three mechanisms at 1133 K and 1300 K. The rate constants for

channels (12a)-(12c) proposed by Fisher et al. and Dooley et al. are nearly identical, but

the rate constant for channel (12d) from Fisher et al. is lower than that of Dooley et al. by

a factor of 2.48 at 1133 K. Fisher et al. treated the rate constant for channel (12d) the

same as the rate constant for channel (12a). On the other hand, Dooley et al. proposed

that channel (12d) should be more reactive than channel (12a) due to the weaker C–H

bond enthalpy at the methyl group bound to the O-atom. Interestingly, the rate constants

for channels (12a)-(12d) proposed by Hakka et al. are consistently higher than those of

Fisher et al. As illustrated in Figure 7.14, the present measurements display somewhat

higher activation energy than the previous estimations. Of all three estimations, the

values from Fisher et al. seem to be in closer agreement with the present measurements.

The measured rate constant is ~17% higher than the value from Fisher et al. at 1355 K,

and is ~35% lower at 897 K. Furthermore, Figure 7.14 presents the estimated overall rate

140

constant for reaction (12) using the group rate constants for the reactions of OH with

ethers and ketones developed by Zhou et al. [121, 131], as was done for reactions (10)

and (11). The estimated values are at least 27% higher than the current data, but it

appears that the estimation captures the temperature dependence of reaction (12) very

well.

Figure 7.14: Arrhenius plot for methyl butanoate + OH (k12) at temperatures above 870 K.

Table 7.5: Comparison of the rate constants for channels (12a)-(12d) from Fisher et al. [48], Dooley et al. [54], and Hakka et al. [55] at 1133 K and 1300 K.

Authors A [cm3 mol-1 s-1] b EA

[cal/mol] 1133 K rate

1300 K rate

Ratio (1133 K) to

Fisher et al. CH3OC(O)C3H7 + OH → CH3OC(O)CH2CH2CH2 + H2O Fisher et al. 5.250E+09 0.97 1590 2.376E+12 2.973E+12 1.00 Dooley et al. 5.280E+09 0.97 1586 2.394E+12 2.995E+12 1.01 Hakka et al. 2.700E+06 2.00 450 2.838E+12 3.833E+12 1.19

CH3OC(O)C3H7 + OH → CH3OC(O)CH2CHCH3 + H2O Fisher et al. 4.680E+07 1.61 -35 3.929E+12 4.893E+12 1.00 Dooley et al. 4.680E+07 1.61 -35 3.929E+12 4.893E+12 1.00 Hakka et al. 2.600E+06 2.00 -765 4.689E+12 5.909E+12 1.19

CH3OC(O)C3H7 + OH → CH3OC(O)CHCH2CH3 + H2O Fisher et al. 1.146E+11 0.51 63 4.024E+12 4.332E+12 1.00

141

Dooley et al. 1.146E+11 0.51 63 4.024E+12 4.332E+12 1.00 Hakka et al. 2.400E+06 2.00 -2450 9.153E+12 1.048E+13 2.27

CH3OC(O)C3H7 + OH → CH2OC(O)C3H7 + H2O Fisher et al. 5.250E+09 0.97 1590 2.376E+12 2.973E+12 1.00 Dooley et al. 7.020E+07 1.61 -35 5.894E+12 7.340E+12 2.48 Hakka et al. 3.600E+06 2.00 -100 4.831E+12 6.324E+12 2.03

7.4 Comparison with Low Temperature Data

Figure 7.15 presents the current high-temperature data for the reactions of OH

with four small methyl esters, along with some earlier experimental work [134-138] at

low temperatures (250-440 K). At a first glance, these kinetic data cannot be described

accurately by simple Arrhenius expressions over a wide range of temperatures. Le Calvé

et al. [134] measured the rate constants for the reactions of OH with a series of formates

(including methyl formate) under pseudo-first-order kinetic conditions using the pulsed

laser photolysis–laser induced fluorescence technique in a reaction cell over 233-372 K,

as illustrated in Figure 7.15a. Surprisingly, they suggested that the H-atom abstraction

from the –OC(O)H group (channel (9a)) is negligible over their temperature range

studied, and their suggestion is very different from the observation based on the high-

temperature measurements. Similarly, Wallington et al. [135] measured the rate constant

for reaction (9) at room temperature (296 K) under pseudo-first-order kinetic conditions

using the flash photolysis–resonance fluorescence technique in a reaction cell, and their

room temperature rate constant is ~24% higher than that of Le Calvé et al. [134]. As

mentioned previously, the estimated rate constants for reaction (9) from Fisher et al. [48]

and Dooley et al. [54] are in good agreement with the present high-temperature data.

However, neither of them can predict the low-temperature data from Le Calvé et al. and

Wallington et al. For instance, the value from Fisher et al. is higher than the measured

room temperature data by a factor of 2.8, and the value from Dooley et al. is lower than

the data by a factor of 2.2. Additionally, the estimated rate constant from Fisher et al.

seems to exhibit more non-Arrhenius curvature than that of Dooley et al., as illustrated in

Figure 7.15a.

142

As shown in Figure 7.15b, Wallington et al. [135] studied reaction (10) under

pseudo-first-order kinetic conditions over the temperature range of 240-440 K. El

Boudali et al. [136] measured the absolute rate constants for the reactions of OH with a

series of acetates (including methyl acetate) using the pulsed laser photolysis–laser

induced fluorescence technique in the cell over 243-372 K, and their measurements are

consistent with the data from Wallington et al. [135]. Similarly, the estimated rate

constant for reaction (10) from Westbrook et al. [53] cannot accurately predict the present

high-temperature data and the previous low-temperature data. The value from Westbrook

et al. is ~30% lower than the present data at 1371 K, and is ~77% higher than the

previous data at 333 K.

Figures 7.15c and 7.15d present the experimental results for the rate constant

measurements of reactions (11) and (12), respectively, at temperatures above 250 K. Le

Calvé et al. [138] measured the absolute rate constants for the reactions of OH with

methyl propanoate, methyl butanoate, methyl valerate, and methyl caproate using the

pulsed laser photolysis–laser induced fluorescence technique in the cell over 253-372 K.

They concluded that the reaction of OH with methyl caproate was the most reactive one

among those 4 methyl esters due to more –CH2– groups available in the molecule. In

addition, their rate constant measurements exhibited slight negative temperature

dependence, except for methyl propanoate. Wallington et al. [135] also measured the

room temperature rate constants for reactions (11) and (12), and their results agree well

with the data from Le Calvé et al. [138] at 296 K. Furthermore, the estimated rate

constants for reactions (11) and (12) from several detailed kinetic mechanisms [46, 48,

54-55] are rather different from the measurements over the temperature range of 250-

1355 K.

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Figure 7.15: Arrhenius plots for methyl ester + OH reactions at temperatures above 250 K.

7.5 Comparison with Structure-Activity Relationship

Similar to the reactions of OH with ketones, the present high-temperature overall

rate constant measurements for reactions (9)-(12) can also be compared with the

estimations using the structure-activity relationship (SAR) developed by Atkinson and

co-workers [127-129]. For instance, the estimated rate constant for the reaction of OH

with methyl propanoate using SAR can be expressed as k11 = kprimF(–OC(O)R) + ksecF(–

C(O)OR)F(–CH3) + kprimF(–CH2–), where R is defined as the alkyl group. As

demonstrated in Figure 7.15, the SAR estimations for reactions (9)-(12) cannot accurately

predict the present high-temperature measurements. More importantly, the SAR

estimation for the reaction of OH with methyl formate requires some additional attention.

Le Calvé et al. [134] suggested that the H-atom abstraction from the –OC(O)H group

144

(channel (9a)) is negligible for the methyl formate + OH reaction over 233-372 K. The

SAR estimation without the consideration of channel (9a) seems to agree well with the

room temperature measurements from Le Calvé et al. [134] and Wallington et al. [135],

but the estimated values are ~40% lower than the present high-temperature data. Hence,

there is a need to consider the effect of channel (9a) in the SAR estimation at high

temperatures. In the present analysis, we could treat the C–H bond in the –OC(O)H

group as a tertiary site, and the estimated rate constant for reaction (9) can be expressed

as k9 = ktertF(=O)F(–OR) + kprimF(–OC(O)H), where R is the alkyl group. The present

SAR estimation with the consideration of channel (9a) is ~25% higher than the current

high-temperature data, but the estimated values are quite different from the previous low-

temperature measurements [134-135]. Similarly, the SAR estimations show good

agreement with the kinetic measurements of reactions (10)-(12) at 298 K, but the

estimated values are higher than the measured values over 333-1371 K. In particular, the

estimated values are ~25% faster than the current rate constant measurements of reactions

(10)-(12) over 876-1371 K. Interestingly, the SAR estimations seem to capture the

temperature dependence of reactions (9)-(12) reasonably well, implying that the pre-

exponential factors for the group rate constants kprim, ksec, and ktert should be reduced by

25%, particularly for these methyl ester + OH reactions. The modified SAR estimations

are in excellent agreement with the present high-temperature measurements, as shown in

Figure 7.16. Note that the measured rate constants for reactions (9) and (10) are nearly

identical at T > 1000 K, but they start to deviate at lower temperatures. The data for

reaction (9) is ~16% higher than the data for reaction (10) at T = 880 K. This trend is

also well-captured by the modified SAR estimations, as demonstrated in Figure 7.16.

145

Figure 7.16: Comparison of the present rate constant measurements with the modified SAR estimations.

7.6 Summary

The overall rate constants for the reactions of OH with methyl formate (k9),

methyl acetate (k10), methyl propanoate (k11), and methyl butanoate (k12) were measured

using OH laser absorption near 306.69 nm behind reflected shock waves over 876-1371

K at pressures near 1.5 atm. These measured rate constants can be expressed in

Arrhenius form as:

k9 = 2.56×1013 exp(-2026/T) cm3 mol-1 s-1

k10 = 3.59×1013 exp(-2438/T) cm3 mol-1 s-1

k11 = 6.65×1013 exp(-2539/T) cm3 mol-1 s-1

k12 = 1.13×1014 exp(-2515/T) cm3 mol-1 s-1

over the temperature ranges studied. Detailed error analyses were conducted with the

consideration of both experimental and secondary chemistry contributions, and the

overall (2σ) uncertainties were estimated to be ±29% at 913 K and ±18% at 1289 K for

k9, ±29% at 930 K and ±17% at 1299 K for k10, ±25% at 909 K and ±17% at 1341 K for

k11, and ±24% at 925 K and ±16% at 1320 K for k12. Additionally, the structure-activity

relationship (SAR) developed by Atkinson and co-workers [127-129] was utilized to

146

estimate the overall rate constants for reactions (9)-(12), and the estimated values are

consistent with the current data (within ~25%).

147

Chapter 8 Conclusions and Future Work

8.1 Summary of Results

The objective of the research presented in this dissertation is to provide reliable

experimental kinetic targets, including ignition delay times, species time histories, and

direct reaction rate constant measurements, using shock tube and laser absorption

methods in order to validate and refine the comprehensive reaction mechanisms for two

different types of oxygenated fuels (i.e., ketones and methyl esters) and to reexamine the

kinetics of the H2 + OH reaction. The topics of this work are mainly divided into three

sections: (1) H2 + OH kinetics, (2) ketone combustion chemistry, and (3) methyl ester +

OH kinetics. Consequently, this work provides accurate rate constant measurements of

reactions (1)-(12), including H2 + OH, ketone decomposition, ketone + OH, and methyl

ester + OH reactions, that are practically important to oxygenated fuel combustion.

These measured rate constants are summarized here.


The rate constant for the reaction of H2 + OH → H 2O + H was measured behind

reflected shock waves using UV laser absorption of OH radicals near 306.69 nm over the

temperature range of 902-1518 K at pressures of 1.15-1.52 atm. The rate constant for

reaction (1) is given by:

k1(T) = 4.38 × 1013 exp(-3518/T) cm3 mol-1 s-1

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over the temperature range studied. The overall uncertainties of reaction (1) were

estimated to be ±17% at 972 and 1228 K. The present measurements are in excellent

agreement with the previous experimental studies [17-21]. In addition, the measured rate

constant is in close accord with the non-Arrhenius expression from GRI-Mech 3.0 [22]

and the theoretical calculation using semi-classical transition state theory (SCTST) from

Nguyen et al. [84].

8.1.2 Ketone Combustion Chemistry

High-temperature acetone and 2-butanone pyrolysis studies were performed

individually behind reflected shock waves using five species time history measurements

(ketone, CO, CH3, CH4 and C2H4). Experimental conditions covered temperatures of

1100-1650 K at pressures around 1.6 atm, for mixtures of 0.25-1.5% acetone or 2-

butanone in argon. During acetone pyrolysis, the CO and acetone time histories were

found to be strongly sensitive to the rate constant for acetone unimolecular

decomposition reaction, and this could be directly determined from the measured CO and

acetone time histories, yielding:

k2(1.23-1.66 atm) = 9.38 × 1041 T-7.85 exp(-44,236/T) s-1

with an uncertainty of ±25%. The inferred rate constant is in good agreement with the

previous shock tube studies from Sato and Hidaka [87] and Saxena et al. [88] (within

30%) at temperatures above 1450 K, but is at least three times faster than the evaluation

from Sato and Hidaka at temperatures below 1250 K. Using the revised values for k2

with the detailed mechanism of Pichon et al. [89], the simulated profiles during acetone

pyrolysis showed excellent agreement with all five species time history measurements.

Similarly, the overall 2-butanone decomposition rate constant was inferred from

the measured 2-butanone time histories, yielding:

k3(1.39-1.62 atm) = 6.08×1013 exp(-31,762/T) s-1

with an uncertainty of ±35%. This rate constant is ~30% faster than that proposed by

Serinyel et al. [95] at 1119 K, and ~100% faster at 1412 K. Using the measured 2-

butanone and CO time histories and an O-atom balance analysis, a missing removal

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pathway for methyl ketene was identified. The rate constant for the decomposition of

methyl ketene was assumed to be the same as the value for the ketene decomposition

reaction. Using the revised values for k3 and adding the methyl ketene decomposition

reaction to the Serinyel et al. mechanism [95], the simulated profiles during 2-butanone

pyrolysis were in better agreement with the measurements for all five species.

Moreover, high-temperature 3-pentanone pyrolysis and oxidation studies were

conducted behind reflected shock waves using laser-based species time history

measurements (3-pentanone, CH3, CO, C2H4, OH and H2O) and ignition delay time

measurements. The measured species time histories and ignition delay times were

compared to the simulations from the detailed kinetic mechanism of Serinyel et al. [98].

In particular, the overall 3-pentanone decomposition rate constant was determined from

the measured 3-pentanone and CH3 time histories during pyrolysis at temperatures of

1070-1530 K and pressures around 1.6 atm, yielding:

k4(1.32-1.75 atm) = 4.383×1049 T-10 exp(-44,780/T) s-1

with an uncertainty of ±35% over 1070-1330 K. The measured k4 was approximately 3.5

times faster than the value used by Serinyel et al. [98]. Using this revised overall 3-

pentanone decomposition rate constant and the additional methyl ketene decomposition

pathway, the modified mechanism was able to successfully simulate all six species time

histories, and showed a significant improvement in the predictions of ignition delay

times.

In addition to the thermal decomposition reactions, another important class of

reactions, which is pertinent to ketone combustion, is the H-atom abstraction reactions by

OH radicals. In this dissertation, the overall rate constants for the reactions of OH with a

series of ketones, namely acetone (k5), 2-butanone (k6), 3-pentanone (k7), and 2-

pentanone (k8), were measured using UV laser absorption of OH over the temperature

range of 870-1360 K at pressures of 1-2 atm. The measured rate constants for reactions

(5)-(8) are given by:

k5 = 3.30×1013 exp(-2437/T) cm3 mol-1 s-1

k6 = 6.35×1013 exp(-2270/T) cm3 mol-1 s-1

k7 = 9.29×1013 exp(-2361/T) cm3 mol-1 s-1

150

k8 = 7.06×1013 exp(-2020/T) cm3 mol-1 s-1

The overall uncertainties were estimated to be ±28% at 1148 K for k5, ±22% at 1039 K

for k6, ±20% at 1188 K for k7, and ±24% at 1186 K for k8. The measured rate constant

for acetone + OH reaction from the current study is consistent with three previous

experimental studies [77, 118-119] within ±20%. In this dissertation, we also presented

the first direct high-temperature rate constant measurements of 2-butanone + OH, 3-

pentanone + OH, and 2-pentanone + OH reactions. The measured values for 2-butanone

+ OH reaction are in close accord with the theoretical calculation from Zhou et al. [121],

and the measured values for 3-pentanone + OH reaction are in excellent agreement with

the estimates (by analogy with the H-atom abstraction rate constants from alkanes) from

Serinyel et al. [98].

8.1.3 Methyl Ester + OH Kinetics

The overall rate constants for the reactions of OH with methyl esters, namely

methyl formate (k9), methyl acetate (k10), methyl propanoate (k11), and methyl butanoate

(k12), were studied by monitoring the OH decays over the temperature range of 876-1371

K at pressures around 1.5 atm. These measured rate constants for reactions (9)-(12) are

given by:

k9 = 2.56×1013 exp(-2026/T) cm3 mol-1 s-1

k10 = 3.59×1013 exp(-2438/T) cm3 mol-1 s-1

k11 = 6.65×1013 exp(-2539/T) cm3 mol-1 s-1

k12 = 1.13×1014 exp(-2515/T) cm3 mol-1 s-1

over the temperature ranges studied. The overall uncertainties were found to be ±29% at

913 K and ±18% at 1289 K for k9, ±29% at 930 K and ±17% at 1299 K for k10, ±25% at

909 K and ±17% at 1341 K for k11, and ±24% at 925 K and ±16% at 1320 K for k12. We

believe these are the first direct high-temperature rate constant measurements for the

reactions of OH with these small methyl esters.

151

8.2 Publications

The work detailed in this dissertation has been published in the following papers:

1) K.-Y. Lam, D.F. Davidson, R.K. Hanson, “A shock tube study of H2 + OH → H2O +

H using OH laser absorption,” Int. J. Chem. Kinet. (2013), doi: 10.1002/kin.20771.

2) K.-Y. Lam, W. Ren, S.H. Pyun, A. Farooq, D.F. Davidson, R.K. Hanson, “Multi-

species time history measurements during high-temperature acetone and 2-butanone

pyrolysis,” Proc. Combust. Inst. 34 (2013) 607-615.

3) K.-Y. Lam, W. Ren, Z. Hong, D.F. Davidson, R.K. Hanson, “Shock tube

measurements of 3-pentanone pyrolysis and oxidation,” Combust. Flame 159 (2012)

3251-3263.

4) K.-Y. Lam, D.F. Davidson, R.K. Hanson, “High-temperature measurements of the

reactions of OH with a series of ketones: acetone, 2-butanone, 3-pentanone, and 2-

pentanone,” J. Phys. Chem. A 116 (2012) 5549-5559.

5) K.-Y. Lam, D.F. Davidson, R.K. Hanson, “High-temperature measurements of the

reactions of OH with small methyl esters: methyl formate, methyl acetate, methyl

propanoate, and methyl butanoate,” J. Phys. Chem. A 116 (2012) 12229-12241.

6) K.-Y. Lam, Z. Hong, D.F. Davidson, R.K. Hanson, “Shock tube ignition delay time

measurements in propane/O2/argon mixtures at near-constant-volume conditions,”

Proc. Combust. Inst. 33 (2011) 251-258.

The additional work that is not discussed in this dissertation has been published

elsewhere:

7) Z. Hong, D.F. Davidson, K.-Y. Lam, R.K. Hanson, “A shock tube study of the rate

constants of HO2 and CH3 reactions,” Combust. Flame 159 (2012) 3007-3013.

8) Z. Hong, K.-Y. Lam, R. Sur, S. Wang, D.F. Davidson, R.K. Hanson, “On the rate

constants of OH + HO2 and HO2 + HO2: A comprehensive study of H2O2 thermal

decomposition using multi-species laser absorption,” Proc. Combust. Inst. 34 (2013)

565-571.

152

9) Z. Hong, K.-Y. Lam, D.F. Davidson, R.K. Hanson, “Broad-linewidth laser

absorption measurements of oxygen between 211 and 235 nm at high temperatures,”

JQSRT 112 (2011) 2698-2703.

10) Z. Hong, K.-Y. Lam, D.F. Davidson, R.K. Hanson, “A comparative study of the

oxidation characteristics of cyclohexane, methylcyclohexane, and n-butylcyclohexane

at high temperatures,” Combust. Flame 158 (2011) 1456-1468.

11) W. Ren, K.-Y. Lam, S.H. Pyun, A. Farooq, D.F. Davidson, R.K. Hanson, “Shock

tube/laser absorption studies of the decomposition of methyl formate,” Proc.

Combust. Inst. 34 (2013) 453-461.

12) A. Farooq, W. Ren, K.-Y. Lam, D.F. Davidson, R.K. Hanson, C.K. Westbrook,

“Shock tube studies of methyl butanoate pyrolysis with relevance to biodiesel,”

Combust. Flame 159 (2012) 3235-3241.

13) S.H. Pyun, W. Ren, K.-Y. Lam, D.F. Davidson, R.K. Hanson, “Shock tube

measurements of methane, ethylene and carbon monoxide time-histories in DME

pyrolysis,” Combust. Flame (2013), doi: 10.1016/j.combustflame.2012.12.004.

14) D.F. Davidson, S.C. Ranganath, K.-Y. Lam, M. Liaw, Z. Hong, R.K. Hanson,

“Ignition delay time measurements of normal alkanes and simple oxygenates,” J.

Propul. Power 26 (2010) 280-287.

8.3 Recommendations for Future Work

8.3.1 Ethyl Radical Diagnostics and Decomposition Pathway

Ethyl radical (C2H5) is an important intermediate species formed during

oxygenated fuel combustion (e.g., 3-pentanone, n-butanol, and ethyl esters). It is rather

short-lived, and decomposes near-instantaneously to form an ethylene molecule (C2H4)

and an H atom. In particular, this decomposition reaction, C2H5 (+ M) → C2H4 + H (+

M), is of practical significance in all hydrocarbon combustion systems, and the

corresponding rate constant has an uncertainty factor of 2-3 (especially in the fall-off

range). This reaction has been primarily studied by researchers at relatively low

153

temperatures (<1100 K) [139-145], and the high-temperature values for this reaction were

poorly known and were extrapolated from the low-temperature data [146]. Hence,

reliable rate constant measurements at combustion-relevant conditions are definitely

needed in order to enhance our understanding of oxygenated fuel combustion.

To study the rate constant for the ethyl radical decomposition reaction, we need a

C2H5 radical precursor and a C2H5 diagnostic system to monitor its decay rates. C2H5

radical can be produced instantaneously upon shock-heating of ethyl iodide (C2H5I) or

azoethane (C2H5N2C2H5). Additionally, C2H5 radical is known to have a broad spectrum

over 200-260 nm, with peak absorption at around 216 and 245 nm [147-148]. During the

thermal decomposition of C2H5 radical, CH3 radical can be formed through a bimolecular

reaction of C2H5 + H → CH3 + CH3. More importantly, CH3 radical has a strong

absorption feature near 200-225 nm [65], and this feature is approximately four times

stronger than the absorption feature of C2H5 radical. To avoid the interference absorption

from CH3 radicals, the wavelength of 245 nm is preferred to monitor C2H5 species time

histories. C2H5 can be measured at 245 nm using the frequency-tripled output of near-

infrared radiation from the pulsed Ti:Sapphire laser (using third harmonic generation).

Unfortunately, the C2H4 molecule (a primary product formed from C2H5 decomposition)

also has a weak absorption feature over 200-250 nm. To account for this interference

absorption, C2H4 should be monitored simultaneously using cw fixed-wavelength laser

absorption at 10.532 µm with a CO2 laser source, and the interference absorption signal

from C2H4 can then be subtracted from the total absorption signal at 245 nm.

8.3.2 Methyl Ester Kinetics

Owing to the practical importance of biodiesel fuels, more experimental studies

are still needed for methyl ester combustion. Many of the reaction rate constants adopted

in the existing kinetic models are poorly known. In particular, the rate constants for the

H-atom abstraction reactions by H, CH3, and HO2 species, which are crucial to the

oxidation chemistry of methyl esters, have not been well studied experimentally and

theoretically. These estimated rate constants were strictly based on the rate constants

154

from other hydrocarbons. Therefore, accurate rate constant measurements of these

reactions are of great interest in the combustion community. Similar to the work

presented in Chapters 6 and 7, fast sources of H, CH3, and HO2 species at elevated

temperatures are required in order to perform these measurements. For instance, ethyl

iodide (C2H5I) and azomethane (CH3N2CH3) can be chosen as H and CH3 species

precursors, respectively. H atoms can be measured using atomic resonance absorption

spectrometry (ARAS) at the Lyman-α wavelength [51-52], while CH3 species can be

easily monitored at 216.6 nm using the frequency-quadrupled output of near-infrared

radiation from the pulsed Ti:Sapphire laser [63-65]. In this dissertation, we have

measured the overall rate constants for the reactions of OH with four small methyl esters.

This work can be further extended to larger methyl esters (i.e., methyl decanoate, a

proposed surrogate fuel for biodiesel).

Moreover, biodiesel fuels are comprised of 5 major components, which are

methyl palmitate (C17H34O2), methyl stearate (C19H38O2), methyl oleate (C19H36O2),

methyl linoleate (C19H34O2), and methyl linolenate (C19H32O2). Owing to the

complexities of these fuel molecules, there are very few experimental studies available in

the literature [43-44, 149-153], and only a few comprehensive kinetic mechanisms [41-

42] have been developed and validated against the existing data. The resulting

performance of these mechanisms is rather poor, and more accurate experimental data are

needed over a wide range of temperatures and pressures in order to evaluate and refine

these models. In particular, speciation (e.g., OH, CH3, CO, CO2, H2O, and CH2O) and

ignition delay time data can be acquired using different laser absorption methods at

Stanford, and can be used as the kinetic targets to provide very strong constraints on the

internal workings of the detailed mechanisms. Additionally, practical methyl esters have

very low vapor pressures, necessitating the use of the aerosol shock tube at Stanford. The

AST works by creating tiny aerosols from the fuels, drawing them into the shock tube,

evaporating them, and shock-heating them into ignition [43, 104].

155

Appendix A Shock Tube Ignition Delay Time Measurements in Propane/O2/Argon Mixtures at Near-Constant-Volume Conditions

Shock tube measurements of ignition delay times with high activation energies

are strongly sensitive to variations in reflected shock temperatures. At longer shock tube

test times, as are needed at low reaction temperatures, small gradual increases in pressure

(and simultaneous increases in temperature) that result from incident shock wave

attenuation and boundary layer growth can significantly shorten measured ignition delay

times. To obviate this pressure increase, we made use of a recently developed driver-

insert method of Hong et al. [154] that allows generation of near-constant-volume test

conditions for reflected-shock measurements. Using this method, we have measured

propane ignition delay times in a lean mixture (0.8% C3H8/ 8% O2/ Ar) over temperatures

between 980 and 1400 K and nominal pressures of 6, 24, and 60 atm, under both

conventional shock tube operation (with post-shock fractional pressure variation dP5/dt ~

1-7 %/ms) and near-constant-volume operation (with dP5/dt ~ 0%/ms). The near-

constant-volume ignition delay times provide a database for low-temperature propane

model development that is independent of non-ideal fluid flow and heat transfer effects.

Comparisons of these near-constant-volume measurements with predictions using the

JetSurF v1.0 mechanism of Sirjean et al. [155] and the Curran et al. mechanism of NUI

Galway [100] were performed. Ignition delay times measured with dP5/dt ~ 1-7 %/ms

156

were found to be significantly shorter (about 1/3 of the near-constant-volume values) at

the lowest temperatures and highest pressures studied. However, these ignition times are

successfully simulated using the JetSurF v1.0 mechanism when an appropriate

gasdynamic model that accounts for changes in pressure and temperature is used.

A.1 Introduction

Recently, there has been an increased interest in propane ignition data at high

pressures and low temperatures due to their significance in the validation of detailed

chemical kinetic mechanisms at these conditions. Both shock tubes and rapid

compression machines (RCM) have been used for these studies. Cadman et al. [156]

examined auto-ignition delay times in shock-heated lean propane-air mixtures (Φ = 0.5)

in the low temperature region of 835-1400 K at pressures of 5-40 bars. Similar shock

tube studies, performed by Herzler et al. [157] and Petersen et al. [158] at temperatures of

750-1300 K and pressures of 10 and 30 bars, are in agreement with the measurements by

Cadman et al. These shock tube ignition delay times start to deviate (roll off) from

detailed model simulations at temperatures around 1000 K. More recently, Gallagher et

al. [159] studied propane oxidation in air using an RCM over temperatures between 680

and 970 K at different compressed gas pressures and equivalence ratios. The measured

RCM ignition delay times are approximately two orders of magnitude longer than the

shock tube data, and a characteristic negative coefficient behavior was observed. Of

critical importance is the fact that modeling of the ignition delay times using a detailed

propane oxidation mechanism (Curran et al. [100]) under commonly employed constant

internal energy (U) and volume (V) constraints cannot reproduce the measurements at

lower temperatures for either shock tube or RCM data (see Figure A.1).

157

Figure A.1: Previous ignition delay time measurements for propane oxidation in air at Φ = 0.5. The constant U, V model calculations utilize the Curran et al. mechanism [100].

A similar discrepancy has been observed in recent reflected-shock studies of

hydrogen-oxygen-argon mixtures. Pang et al. [160] found that hydrogen ignition delay

times are at least an order of magnitude shorter than the constant U, V predictions of

detailed kinetic mechanisms at temperatures less than 930 K. For these low temperature

experiments, a gradual pressure increase (dP5/dt) was observed. That study concluded

that the post-shock pressure variation contributed significantly to shorter ignition delay

times in long-duration (low temperature) experiments. In an ideal shock tube experiment,

the incident shock travels at a constant speed and is reflected from the driven-section

endwall. The reflected shock further compresses the incident-shocked test gas, yielding a

stagnant and uniform high-temperature region. Computationally, this reflected shock

region is often approximated as a constant volume reactor. Additionally, gases in the

reflected shock region are often assumed to have constant internal energy (U), as would

be appropriate for adiabatic systems without work addition or chemical energy change.

However, due to viscous effects, sidewall boundary layers are developed behind incident

shock waves, causing incident shock attenuation and leaving a non-uniform flow field in

the incident shock region. As the reflected shock propagates into this flow field and

interacts with the growing boundary layers, changes in the post-reflected-shock pressure

0.6 0.8 1.0 1.2 1.4 1.60.01

0.1

1

10

100

1000

Petersen (2007) Herzler (2004) (Ar) Herzler (2004) Gallagher & Curran

(2008) (RCM)

2.1% C3H8 / O2 / N2Φ = 0.5, P = 30 atm

Igni

tion

Tim

e [m

s]

1000/T [1/K]

1250K 1000K 833K 714K

Const.U,VModel

158

may occur that can penetrate into the reflected shock region [161]. Consequently, an

increase in pressure (and a concomitant increase in temperature) is typically observed in

the reflected shock region, introducing uncertainties in the ignition delay time

measurements. In such experiments, the constant-volume assumption is not strictly valid.

Furthermore, the undiluted mixtures at elevated pressures exhibit an accelerated

ignition caused by some localized pre-ignition energy release, resulting in a dramatic and

sudden pressure rise prior to the primary ignition event (as discussed by Petersen et al.

[158]). Such dramatic pressure effect is different from the normal, gradual pressure

increase (dP5/dt) observed in dilute mixtures. Both pressure effects cannot be simply

reproduced by a detailed kinetic mechanism under constant U, V constraints. In the

current study, we investigate only the effects of the gradual pressure increase on ignition

delay time measurements using a dilute mixture. Hence, we measured propane ignition

delay times in a lean mixture (0.8% C3H8/ 8% O2/ Ar) at nominal pressures of 6, 24 and

60 atm, first using a conventional shock tube, to confirm the discrepancies between

experimental data and simulations based on two current chemical mechanisms using

constant U and V constraints. We then employed a driver-insert method (developed in

our laboratory [154]) to minimize the post-shock pressure variations and generate near-

constant-volume test conditions for the propane oxidation study. These near-constant-

volume measurements were compared to constant U, V predictions using the JetSurF

v1.0 mechanism of Sirjean et al. [155] and the Curran et al. mechanism of NUI Galway

[100]. Finally, a thermodynamic-gasdynamic model, CHEMSHOCK [162], was used to

simulate the conventional shock tube ignition delay data (i.e., with dP5/dt > 0) by

incorporating the effects of the measured non-ideal growth of pressure (caused by

incident shock wave attenuation and boundary layer growth).

A.2 Experimental Setup

A.2.1 Low-Pressure Shock Tube Experiments

Ignition delay times at pressures near 6 atm were measured in the Stanford

stainless-steel, high-purity low-pressure shock tube (LPST) previously described by

159

Oehlschlaeger et al. [163]. The shock tube inner diameter is 14.13 cm, and the driven

section is 8.54 m in length, separated from the driver section by polycarbonate

diaphragms (0.5 and 1.0 mm in thickness). Under normal shock tube conditions, the

driver section is 3.35-m long, and helium is used as the driver gas to obtain a 2-ms high-

quality test time. To extend the test time for low temperature ignition measurements, the

driver section was lengthened to 7.12 m to allow more time before the arrival of the

rarefaction fan at the test location. The stainless-steel driver-section extension has an

inner diameter of 14.13 cm and a U-shaped structure with a 23-cm radius of curvature on

the centerline. In addition, tailored driver gas mixtures of 30-40% nitrogen in helium

were used to minimize reflected shock-contact surface interactions [164]. These

modifications provide an available test time of approximately 25 ms at a temperature of

about 1000 K.

The gases used, propane (Instrument Grade) 99.5%, oxygen (Research Grade)

99.999%, and argon (Research Grade) 99.999%, were supplied by Praxair and used

without further purification. A test gas mixture (0.8% C3H8, 8% O2, with balance argon)

was prepared manometrically in a 12-liter stainless-steel mixing cylinder and

mechanically mixed with a magnetically-driven stirrer for at least 2 hours prior to the

experiments.

All measurements were performed behind reflected shock waves at a test location

2 cm from the driven-section endwall. Incident shock velocities were measured with five

piezo-electric pressure transducers (PCB 113A) spaced axially over the last 1.5 m of the

tube and extrapolated to the endwall. The signals of the transducers triggered by the

incident shock were delivered to four Philips PM6666 counter timers (with resolution of

0.1 μs), which determined the shock-passage time intervals. The incident shock velocity

at the endwall can then be estimated by a linear extrapolation of the incident shock

velocities determined from the shock-passage time intervals. Typical shock attenuation

rates ranged from 0.8 to 1.5% per meter. Reflected shock conditions were calculated

using one-dimensional shock relations and assuming vibrational equilibrium and frozen

chemistry. Uncertainty in the initial temperature and pressure was approximately ±0.7%

and ±1%, respectively, mainly due to the uncertainty in the measured shock velocity

160

(±0.2%). A Kistler piezo-electric pressure transducer (603B1) coated with RTV silicone

rubber was utilized to measure pressure-time history in the test section at 2 cm from the

endwall. In addition, OH* emission chemiluminescence was monitored at the test section

using a vertical slit, lens, Schott Glass UG5 filter (with >95% transmission at 306 nm),

and modified UV-enhanced Thorlabs PDA36A photodiode detector. The temporal

resolution of this emission set-up is typically 10 μs or better.

A.2.2 High-Pressure Shock Tube Measurements

All experiments at pressures higher than 20 atm were performed in the Stanford

stainless-steel, high-purity, high-pressure shock tube (HPST). A complete description of

this shock tube is provided by Petersen et al. [165]. The driver section is 3-m long with a

7.5-cm inner diameter, and the driven section is 5-m long with a 5-cm inner diameter,

separated by an aluminum diaphragm (1.27-3.18 mm in thickness) with cross-scribing.

Helium was used for the driver gas to provide approximately 2-3 ms of high-quality test

time. At lower temperatures, longer test time was achieved by using tailored driver gas

mixtures of 40-50% nitrogen in helium [164], and the available test time was

approximately 10 ms at a temperature of about 1000 K. All measurements were made

behind reflected shock waves, at a test location 1 cm from the driven-section endwall.

The incident shock velocities were measured using six piezo-electric pressure transducers

(PCB 113A), with five corresponding Philips PM6666 counter timers (with resolution of

0.1 μs), spaced over the last 2 m of the shock tube and extrapolated to the endwall.

Shock attenuation rates varied from 1.0 to 3.5% per meter. Endwall shock velocity,

incident and reflected shock conditions were determined using the same method as

described above. Uncertainty in the initial temperature and pressure is less than ±1%,

with the primary contribution being uncertainty in the measured shock velocity.

Pressure-time history in the test section was monitored by a Kistler pressure transducer

(603B1) coated with RTV. Additionally, emission from OH* chemiluminescence was

measured at the test location using the similar set-up as described above.

161

A.2.3 Driver-Insert-Method

Non-ideal gasdynamic effects of shock waves (primarily incident shock wave

attenuation and sidewall boundary layer growth) can cause a gradual increase in pressure

(and a simultaneous increase in temperature) in the reflected shock region [165-167].

Hong et al. [154] have developed a semi-analytical model for designing shock tube driver

inserts to obviate non-ideal pressure variations in the reflected shock region. Optimal

configuration for the driver insert uses a parabolic shape, which is dependent on the

reflected shock temperature, pressure, and the composition of the driver and driven gases.

The driver insert is used to reflect part of the rarefaction fan back to the test section,

superimposing a pressure decrease on the non-ideal pressure growth due to the boundary

layers. If both effects have the same order of magnitude, the non-ideal pressure rise can

be effectively eliminated, yielding near-constant-volume test conditions (with a fractional

rate of pressure change dP5/dt ~ 0%/ms) in the reflected shock region (see Figure A.2).

With a proper combination of driver insert and tailored driver gas mixtures, long uniform

test times of 25 ms and 10 ms can be acquired at temperatures around 1000 K for the

low-pressure and high-pressure shock tubes, respectively. Additionally, constant

pressure implies a constant temperature profile by assuming isentropic behavior of the

gas at the test section. This has been confirmed with an in situ two-line thermometry

diagnostic using two distributed feedback (DFB) diode lasers near 2.7 μm [168].

162

Figure A.2: Comparison of pressure profiles for a mixture of 0.8% C3H8/ 8% N2/ Ar obtained with and without driver insert in the Stanford 14.13 cm diameter shock tube. The fractional pressure rise without driver insert (over 20 ms) is approximately 20%, compared to ±3.0% local pressure variations with driver insert. The decay beginning at 25 ms is due to arrival of the rarefaction wave from the driver section.

A.3 Results and Discussion

Lean propane ignition delay times were measured at temperatures between 980

and 1400 K and nominal pressures of 6, 24, and 60 atm. The ignition delay time is

defined as the time interval between the arrival of the reflected shock and the initial rise

in the pressure and excited OH (OH*) emission traces. Ignition data obtained from the

pressure and emission traces are consistent within ±1%. The overall uncertainty in

ignition delay time measurements is approximately ±10%, with the primary contribution

from the uncertainty in reflected shock temperatures. The typical facility-related rates of

pressure change in the LPST and the HPST without the use of driver inserts were 1-

3%/ms and 6-7%/ms, respectively. The non-ideal pressure growth is larger in the HPST,

primarily due to the fact that its driven section has a smaller inner diameter. Clearly, the

inner diameter of the shock tube can have a significant impact on post-shock pressure

variations.

0 10 20 300

4

8

12

T5 = 1038 KdP5/dt ~ 1.0%/ms

Pres

sure

[atm

]

Time [ms]

0.8% C3H8/ 8% N2/ ArP5 = 6.7 atm

T5 = 1020 KdP5/dt ~ 0%/ms

163

Representative LPST propane ignition data are illustrated in Figure A.3. In the

reflected shock wave experiment without the driver insert, a gradual pressure rise with a

rate of 2%/ms is observed after the passing of the reflected shock. This non-ideal

behavior causes approximately a 20% increase (over 10 ms) in pressure prior to ignition,

and concomitantly (assuming isentropic compression) approximately an 8% increase in

temperature. Such pressure rise is caused by the incident shock attenuation and boundary

layer growth. In the reflected shock wave experiment with a properly designed driver

insert (also shown in Fig. A.3), the pressure rise behind the reflected shock was

eliminated. The pressure profile obtained using the driver insert is nearly flat with local

pressure variations of ±3% prior to ignition, relatively insignificant when compared to the

20% pressure rise in the uncompensated conventional experiments. The ignition delay

time measured under conventional operating conditions (without the driver insert) at the

initial temperature T5 = 1034 K is approximately 7 ms shorter than that obtained under

near-constant-volume conditions (with the driver insert) at T5 = 1044 K, owing to the

increase in reaction temperature occurring in the former.

Figure A.3: Comparison of pressure profiles for reactive mixture with and without LPST driver insert. Pressure rise without driver insert (over 10 ms) is 20%, compared to ±3.0% pressure variations with driver insert. Initial reflected shock conditions: T5 = 1034 K and P5 = 7.1 atm (with dP5/dt ~ 2%/ms), T5 = 1044 K and P5 = 6.7 atm (with dP5/dt ~ 0%/ms).

0 5 10 150

10

20

τign

T5 = 1044 K

Pres

sure

[atm

]

Time [ms]

T5 = 1034 KdP5/dt ~ 2%/ms

0.8% C3H8/ O2/ Ar Φ = 0.5

τign

OH* Emission (1044 K)

164

Examples of HPST ignition time data with and without driver inserts are shown in

Figure A.4. The initial reflected shock pressure of both experiments is approximately the

same (~54 atm), and the ignition delay time of the measurement taken without the use of

a driver insert (dP5/dt ~ 7%/ms) at a slightly colder initial temperature of 996 K is at least

one-third shorter than the near-constant-volume measurement taken with the driver insert

(dP5/dt ~ 0%/ms) at T5 = 1008 K.

Figure A.4: Comparison of pressure profiles for reactive mixture with and without HPST driver insert. Pressure rise without driver insert (over 2.5 ms) is 17.5%, compared to ±1.0% pressure variations with driver insert. Initial reflected shock conditions: T5 = 996 K and P5 = 54.7 atm (with dP5/dt ~ 7%/ms), T5 = 1008 K and P5 = 53.7 atm (with dP5/dt ~ 0%/ms).

At temperatures less than 1100 K (or test times longer than 5 ms), a small

pressure bump is observed in the LPST experiments (extending from about 5-8 ms in Fig.

A.3), even though a near-perfect tailoring condition has been achieved. (Note that a

small pressure bump is also present in the HPST experiments extending from about 4-6

ms.) This is primarily due to the fact that the contact surface is not actually a thin

interface. Instead, it is a mixing zone, which has a finite width and is caused by non-ideal

diaphragm bursting and boundary layer growth. Several strategies were attempted to

eliminate such a pressure bump (see Davidson et al. [167]) with varying degrees of

success, and eventually relatively small “bumps” were achieved typified by Fig. A.3. To

0 1 2 3 40

50

100

150

τign

T5 = 1008 K

Pres

sure

[atm

]

Time [ms]

T5 = 996 KdP5/dt ~ 7%/ms

0.8% C3H8/ O2/ Ar Φ = 0.5

τign

165

examine the effect of such pressure bump on ignition delay time, CHEMSHOCK (will be

discussed later) with the JetSurF v1.0 mechanism [155] and the actual experimental

pressure profile (with the pressure bump and dP5/dt ~ 0%/ms in Fig. A.3) was first used

to compute the ignition delay time, and the calculated ignition delay time was 17.9 ms at

T5 = 1044 K. After that, the pressure bump present in the experimental pressure profile

was artificially removed, and CHEMSHOCK with the artificial pressure profile gave the

ignition delay time of 18.5 ms at T5 = 1044 K, which is ~3% longer than the former

computed value (with the pressure bump). Hence, such small pressure bumps do not

affect the ignition delay time measurements significantly in the current study.

The low-pressure ignition delay times under both near-constant-volume

conditions and conventional operation conditions are plotted at the initial reflected shock

temperatures in Figure A.5, along with the calculated values from the JetSurF v1.0

mechanism of Sirjean et al. [155] and the Curran et al. mechanism of NUI Galway [100]

under constant U, V constraints. The calculated values from both mechanisms are in

excellent agreement at a pressure of 6 atm. At reflected shock temperatures less than

1250 K and pressures of 5.3-68 atm, an ignition time correlation for the current mixture

can also be determined from a total of 50 near-constant-volume measurements with a R2

value of 0.990: τ = 2.49×10 -5P-0.89exp(15674/T), where the ignition time is in

milliseconds, the pressure is in atmospheres, and the temperature is in kelvins. All data

were then scaled by P-0.89 to their nominal pressures. Note that both sets of measurements

in the current study agree with each other at temperatures above 1110 K (i.e. at short test

times). The discrepancy between the two measurement sets increases with decreasing

temperature, but the constant U, V predictions remain in good agreement with the near-

constant-volume ignition delay time measurements (dP5/dt = 0%/ms) at all temperatures.

In contrast, the current data under conventional operation conditions (dP5/dt = 1-

3%/ms) are shorter than the results from both models, starting at T5 = 1110 K. For

instance, the uncompensated conventional data is approximately 15 milliseconds faster

than the constant U, V models at T5 = 990 K. Figure A.5 also includes the previous

ignition data from Cadman et al. [156] using the same test mixture. In the study

performed by Cadman et al., the data were obtained at a reflected shock pressure of 5

166

atm. To perform a quantitative comparison, those ignition data were also scaled by P-0.89

from P5 = 5 to 6 atm. Notice that those previous measurements are substantially shorter

than the current measurements and the constant U, V model predictions by at least a

factor of 3. As mentioned previously, the inner diameter of the shock tube can have a

significant impact on post-shock pressure variations. The Cadman et al. shock tube

facility has an inner diameter of about 6 cm, which is much smaller than that of the

Stanford low-pressure shock tube (14.13 cm). Based on their published pressure profiles

(Fig. 5 from their publication), a relatively large post-shock pressure variation of

approximately 15%/ms or higher was identified. Such large dP5/dt can contribute to the

substantial differences between the previous and current measurements. For kinetic

modeling of reflected-shock experiments, it is now clear that it is critically important to

account for the post-shock pressure variations in different shock tube facilities and that

pressure growth rates be documented to allow proper modeling of experiments. Incorrect

reaction pathways and rate coefficients may be inferred if uncompensated ignition data

are fit by detailed mechanisms.

Figure A.5: Ignition delay times for 0.8% C3H8/ 8% O2/ Ar mixture at P5 = 6 atm, plotted at the initial post-shock T5. Experimental data and calculated values from JetSurF v1.0 mechanism [155] and Curran et al. mechanism [100].

0.6 0.7 0.8 0.9 1.0 1.1

0.1

1

10

Current Study (dP5/dt~0%/ms) Current Study (dP5/dt~1-3%/ms) Cadman (2000) (dP5/dt~15%/ms) JetSurF v1.0 (Const. U,V) Curran et al. (Const. U,V)

1250K 1000K1111K

Igni

tion

Tim

e [m

s]

1000/T5 [1/K]

1429K

0.8% C3H8/ O2/ ArΦ = 0.5, P5 = 6 atm

167

As discussed by Pang et al. [160], constant U, V modeling of low-temperature

shock tube ignition delay times with finite dP5/dt gives an unsatisfactory and incorrect

representation of the actual chemical kinetics. A gasdynamic solver developed in our

laboratory, named CHEMSHOCK, can be used to reasonably account for the non-ideal

growth of pressure and temperature in the reflected shock region, up to the time of

ignition, at least for simple reaction systems without significant pre-ignition heat release.

CHEMSHOCK performs a two-step simulation process over each small time step: (1)

constant U, V calculations, followed by (2) an assumed isentropic adjustment in

conditions to recover the measured value of pressure at this time. This code has been

validated against simulations from a one-dimensional reacting computational fluid

dynamics code for a heptane/ O2/ Ar mixture using a reduced mechanism, and against

experimental data (such as gas temperature and water vapor concentration) in a hydrogen/

O2/ Ar mixture by Li et al. [162]. The calculated ignition delay times from

CHEMSHOCK using the JetSurF v1.0 mechanism [155] and different pressure

constraints (dP5/dt = 1.5%/ms and 15%/ms) are displayed in Figure A.6. Here, the

computed values are much closer to the uncompensated ignition delay times. By setting

the post-shock pressure variation to 1.5%/ms, the model agrees with the conventional

data in the current study within ±10% at lower temperatures. Additionally, the

CHEMSHOCK modeling is able to capture the early roll-off behavior displayed by the

previous ignition data from Cadman et al. (within 20%) at low temperatures after

incorporating the published post-shock pressure increase of 15%/ms.

168

Figure A.6: Low-pressure experimental data and CHEMSHOCK modeling using JetSurF v1.0 mechanism.

Figure A.7 shows the Arrhenius plots of lean propane ignition delay time

measurements at elevated pressures of 24 and 60 atm, respectively. All experimental

data were scaled to either 24 or 60 atm using P-0.89. The measured facility-related rate-of-

pressure change in the high-pressure shock tube was typically 6-7% per millisecond,

which is much larger than observed in the low-pressure shock tube. Good agreement is

found between the ignition delay data with dP5/dt of ~0%/ms and ~6-7%/ms, at

temperatures higher than 1110 K or when the ignition delay times are shorter than 1

millisecond. Notice, however, that the discrepancy in ignition delay data between

conventional operation and near-constant-volume conditions is larger in the 24 atm

experiments in the high-pressure shock tube than in the 60 atm experiments. For

instance, at the initial reflected shock condition of T5 = 990 K and P5 = 24 atm, the

uncompensated data are 50% shorter than the near-constant-volume data. This is because

the longer ignition delay times for the 24 atm data are more sensitive to changes in the

temperature gradient than the shorter-ignition-time 60 atm data.

The compensated high-pressure experimental data in Fig. A.7 are compared to

predictions using both JetSurF [155] and Curran et al. [100] mechanisms under constant

U, V constraints. At a nominal pressure of 24 atm, the computed ignition delay times

0.8 0.9 1.0

1

10

JetSurF v1.0: CHEMKIN (Const. U,V) CHEMSHOCK (1.5%/ms) CHEMSHOCK (15%/ms)

Curran et al.: CHEMKIN (Const. U,V)

1250K 1000K1111K

Igni

tion

Tim

e [m

s]

1000/T5 [1/K]

0.8% C3H8/ O2/ ArΦ = 0.5, P5 = 6 atm

169

from the JetSurF mechanism show very good agreement with the near-constant-volume

measurements. At a nominal pressure of 60 atm, the constant U, V model predictions

from the JetSurF mechanism agree reasonably well with the compensated data, but are

consistently 15% longer than the measurements. However, the predicted values from the

Curran et al. mechanism are at least 30% shorter than the near-constant-volume

measurements at both pressures. Referring now to the uncompensated (conventional

operation) results, all the experimental data are much shorter than the constant U, V

model predictions at lower temperatures. The uncompensated data are at least three times

shorter than the constant U, V model calculations (JetSurF) at T5 = 950 K and P5 = 24

atm, while they are 66% shorter at T5 = 996 K and P5 = 60 atm. Hence, the

CHEMSHOCK model with the JetSurF mechanism [155] was utilized to account for the

non-ideal pressure rise in the reflected shock region. After incorporating a proper

pressure constraint (dP5/dt = 7%/ms), the uncompensated data are successfully simulated

at both pressures, as illustrated in Figure A.7. However, there does appear to be a small

residual systematic difference between the measured and simulated ignition delay times

in both the compensated and uncompensated 60 atm data. This may indicate a deficiency

in the pressure-dependent chemistry in the JetSurF mechanism. A preliminary sensitivity

analysis indicates that a 50% increase in the pressure-dependent hydrogen peroxide

decomposition rate at 60 atm would eliminate this systematic difference.

170

Figure A.7: High-pressure experimental data (at P5 = 24 and 60 atm), along with CHEMKIN and CHEMSHOCK modeling using JetSurF v1.0 and Curran et al. mechanisms.

A.4 Concluding Remarks

Propane ignition delay times in a lean mixture (0.8% C3H8/ 8% O2/ Ar) were

measured over the temperature range of 980-1400 K at nominal pressures of 6, 24, and 60

atm, under conventional shock tube operating conditions and under near-constant-volume

conditions, using both the low-pressure and high-pressure shock tubes. Under

0.8 0.9 1.0 1.10.1

1

10

Current Study (dP5/dt~0%/ms) Current Study (dP5/dt~6-7%/ms) JetSurF v1.0 (Const. U,V) JetSurF v1.0 (dP5/dt=7%/ms) Curran et al. (Const. U,V)

909K1250K 1000K1111K

Igni

tion

Tim

e [m

s]

1000/T5 [1/K]

0.8% C3H8/ O2/ ArΦ = 0.5, P5 = 24 atm

0.8 0.9 1.0 1.10.1

1

10 Current Study (dP5/dt~0%/ms) Current Study (dP5/dt~6-7%/ms) JetSurF v1.0 (Const. U,V) JetSurF v1.0 (dP5/dt=7%/ms) Curran et al. (Const. U,V)

0.8% C3H8/ O2/ ArΦ = 0.5, P5 = 60 atm

909K1250K 1000K1111K

Igni

tion

Tim

e [m

s]

1000/T5 [1/K]

171

conventional operating conditions, a facility-related rate-of-pressure rise (dP5/dt ~ 1-

7%/ms) is observed in the reflected shock region, introducing strong variations in the

shock tube ignition delay time measurements. Under near-constant-volume test

conditions, a flat pressure profile with small local variations was obtained using a driver-

insert method and tailored driver gas mixtures. Good agreement in both types of

measurements is found only at temperatures higher than 1110 K. As temperature is

decreased, the discrepancy between the two types of measurements becomes larger at all

pressures.

The near-constant-volume measurements of ignition delay times were found to be

in good agreement with the predictions of the JetSurF v1.0 mechanism [155], using

commonly employed constant U, V constraints. The ignition delay time measurements

using conventional shock tube operation (with dP5/dt ~ 1-7%/ms) were also found to be

in good agreement with detailed modeling, when CHEMSHOCK [162] is used to include

the effects of pressure and temperature variations. The new near-constant-volume data

provide a valuable database for low-temperature propane mechanism development that is

independent of fluid flow and heat transfer effects.

173

Appendix B Ignition Delay Time Measurements of Normal Alkanes and Cycloalkanes

B.1 Objectives

Ignition delay time experiments were performed on a series of straight-chain

alkanes and cycloalkanes using the Stanford Low Pressure Shock Tube (LPST). These

experiments were performed on n-pentane (n-C5H12), n-hexane (n-C6H14), n-octane (n-

C8H18), n-nonane (n-C9H20), cyclohexane (CH or cC6H12), methylcyclohexane (MCH or

CH3cC6H11), and n-butylcyclohexane (BCH or n-C4H9cC6H11) at various reflected shock

temperatures and pressures (between 1240 and 1500 K and 1.5 and 3.8 atm) and at two

equivalence ratios, namely Φ = 1.0 and Φ = 0.5. These ignition delay time measurements

were then compared with the modeling results obtained using the comprehensive kinetic

mechanism of Sirjean et al. (JetSurF v1.1) [169].

B.2 Experimental Details

Research grade O2 and argon were used in all mixtures, along with ≥ 99.5% pure

fuels from Sigma-Aldrich that were degassed by vacuum pumping prior to mixture

preparation. The fuels are n-pentane (n-C5H12), n-hexane (n-C6H14), n-octane (n-C8H18),

n-nonane (n-C9H20), cyclohexane (C6H12), methylcyclohexane (MCH), and

butylcyclohexane (BCH). Test mixtures (fuel, 4% O2, with balance Ar) were prepared in

174

a 14-liter stainless steel tank. The desired mixture ratio of the reactants was obtained on

the basis of partial pressures as measured with a 10,000 Torr Baratron pressure gauge.

The reactants were mixed by an electrically-driven stirring rod for 2 hours. For all test

mixtures, the mixing tank (with a heating jacket) and the manifold (wrapped with a

flexible constant-wattage silicone rubber heating tape) were heated such that their surface

temperature remained at around 60 °C.

The mixture compositions were validated by monitoring the fuel concentrations in

the shock tube (from near the endwall) with a Jodon™ Helium-Neon laser at 3.39 µm.

The absorption cross-sections of n-alkanes for Beer’s law were directly obtained from the

PNNL database [123], and the values from the PNNL database are in good agreement

with the previous measurements from Klingbeil et al. [122]. The measured fuel

concentrations for these n-alkane mixtures were consistent with the values expected from

the manometrical preparation within ±5%. On the other hand, the absorption cross-

sections of those three cycloalkanes were measured in the shock tube using HeNe laser

absorption at 3.39 µm, and the measured absorption cross-sections for CH, MCH, and

BCH are 75.1, 50.2, and 56.6 m2/mol, respectively, at 25 oC, with uncertainties of ±5%.

By employing these measured absorption cross-sections, the measured fuel

concentrations for the cyclohexane and methylcyclohexane mixtures were in good

agreement with the values expected from the manometrical preparation within ±5%. On

the contrary, the measured fuel concentrations for the n-butylcyclohexane mixtures were

found to be 30% less than the expected values, and these fuel losses were properly

accounted for in the test mixtures when compared to the simulations using the JetSurF

v1.1 mechanism (under constant energy and volume constraints).

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B.3 Ignition Delay Time Plots

B.3.1 Normal Alkane Ignition

Figure B.1: n-Pentane ignition delay time measurements at pressures of 1.8 and 3.6 atm and equivalence ratios of 1.0 and 0.5.

Figure B.2: n-Hexane ignition delay time measurements at pressures of 1.8 and 3.6 atm and equivalence ratios of 1.0 and 0.5.

176

Figure B.3: n-Octane ignition delay time measurements at pressures of 1.8 and 3.6 atm and equivalence ratios of 1.0 and 0.5.

Figure B.4: n-Nonane ignition delay time measurements at pressures of 1.8 and 3.6 atm and equivalence ratios of 1.0 and 0.5.

177

B.3.2 Cycloalkane Ignition

Figure B.5: Cyclohexane (CH) ignition delay time measurements at pressures of 1.5 and 3.0 atm and equivalence ratios of 1.0 and 0.5.

Figure B.6: Methylcyclohexane (MCH) ignition delay time measurements at pressures of 1.5 and 3.0 atm and equivalence ratios of 1.0 and 0.5.

178

Figure B.7: n-Butylcyclohexane (BCH) ignition delay time measurements at pressures of 1.5 and 3.0 atm and equivalence ratios of 0.88 and 0.45.

B.4 Summary

Ignition delay time measurements were performed on a series of straight-chain

alkanes and cycloalkanes over 1240-1500 K at pressures of 1.5 and 3.8 atm and at two

equivalence ratios of Φ = 1.0 and 0.5. Further details regarding this work can be found

from the following papers:

• D.F. Davidson, S.C. Ranganath, K.-Y. Lam, M. Liaw, Z. Hong, R.K. Hanson,

“Ignition delay time measurements of normal alkanes and simple oxygenates,” J.

Propul. Power 26 (2010) 280-287.

• Z. Hong, K.-Y. Lam, D.F. Davidson, R.K. Hanson, “A comparative study of the

oxidation characteristics of cyclohexane, methylcyclohexane, and n-

butylcyclohexane at high temperatures,” Combust. Flame 158 (2011) 1456-1468.

179

Appendix C Multi-Species Time History Measurements during the Oxidation of n-Decane, iso-Octane, and Toluene

The high-temperature oxidation of n-decane, iso-octane, and toluene was studied

behind reflected shock waves using laser absorption methods to measure time histories of

three species: OH, C2H4, and CO. These species time histories were then compared to

the simulations from the comprehensive reaction mechanisms (JetSurF v1.1 [169] and

LLNL [170] mechanisms). Additionally, ignition delay times for the mixtures of n-

decane, iso-octane, and toluene were compared with the measurements of the JP-8

mixtures, and the ignition times for the iso-octane and JP-8 mixtures were found to be in

good agreement.

C.1 Introduction

Jet aviation fuels such as Jet-A and JP-8 are complicated mixtures of hundreds,

even thousands of different chemical components [171]. To model the chemistry of these

fuels, surrogate fuel models consisting of a small number of representative chemical

components and a limited accompanying detailed reaction mechanism can be used.

These surrogate fuel models are generally able to successfully imitate the gas phase

combustion characteristics of the real fuel being investigated [172]. The major advantage

180

of these surrogate fuel models, of course, is that the size of the detailed kinetic

mechanism is more manageable and requires less computational power.

These surrogate models can be used to predict auto-ignition, heat release rate,

adiabatic flame temperature, extinction, sooting behavior, and other important

combustion property targets. However, it is often difficult to find a surrogate that can

successfully match both the physical properties and the chemical kinetics properties of

the fuel. Several surrogate fuel components are currently being considered, including: n-

decane, n-dodecane, iso-octane, iso-cetane, methyl cyclohexane, n-butyl cyclohexane,

toluene, n-propyl benzene, 1,3,5 tri-methylbenzene, and 1-methyl naphthalene [172].

Some idea of the current state of performance for surrogate models can be gained by

examining two surrogate fuel models for JP-8: the Lindstedt model (89% n-decane/ 11%

toluene) [173] and the CSE model (52.93% n-decane/ 12.96% iso-octane/ 34.11%

toluene) [174]. Fig. C.1 shows the measured OH and C2H4 time histories during high-

temperature JP-8 oxidation, along with the simulations based on the Lindstedt model and

the CSE model using a detailed kinetic mechanism of Dooley et al. [172]. As can be

seen, both surrogate fuel models can capture the general shapes of the OH and C2H4 time

histories. However, neither of these models can accurately predict the ignition delay

times for this JP-8 mixture, nor can they capture the initial formation rates and initial

plateau levels of the important chain-branching radical OH. Interestingly, the Lindstedt

model [173] is able to capture the initial C2H4 plateau levels reasonably well, but the CSE

model [174] underpredicts the plateau levels at the experimental conditions. Clearly,

further experimental and theoretical studies are needed to improve the current surrogate

fuel models.

The validation of multi-component surrogate kinetics is predicated on the

successful validation of each individual surrogate component. Here we present

experimental kinetic targets for the high-temperature oxidation of three individual

surrogate components: n-decane, iso-octane, and toluene. These kinetic targets are

species time histories of OH, C2H4, and CO derived from laser absorption measurements

behind reflected shock waves. These measured species time histories are also compared

with the simulations from the existing detailed mechanisms.

181

Figure C.1: OH and C2H4 time history measurements for the mixture of 424 ppm JP-8 with 0.813% O2 in Ar. Two JP-8 proposed surrogate models were employed. Simulations were done using the Dooley et al. mechanism [172].

C.2 Experimental Details

C.2.1 Mixture Preparation


heated uniformly to 60 oC and mixed with a magnetically-driven stirring vane for at least

182

2 hours prior to the experiments. Research grade (99.999%) gases from Praxair were

used in mixture preparation without further purification. In addition, liquid chemicals

were all obtained from Sigma-Aldrich. Anhydrous grade (≥99%) n-decane, anhydrous

grade (≥99.8%) iso-octane, and ACS spectrophotometric grade (≥99.5%) toluene were

further treated using a freeze-pump-thaw procedure to remove dissolved volatiles and air

prior to mixture preparation. The mixture compositions in this study were ~360 ppm n-

decane with 0.813% O2 in Ar (Φ ~ 0.7), 511 ppm iso-octane with 0.813% O2 in Ar (Φ ~

0.8), and 640 ppm toluene with 0.813% O2 in Ar (Φ ~ 0.71).

C.2.2 Helium-Neon Laser Measurement of Fuel

The mixture compositions were confirmed by sampling a portion of the mixture



laser at 3.39 µm [82, 122]. Beer’s law was used to convert the measured absorption data

into the fuel mole fraction. The absorption cross-sections of n-decane, iso-octane, and

toluene for Beer’s law were directly obtained from the PNNL database [123], and the

measured fuel concentrations were found to be lower than the values expected from the

manometrical preparation by ~30% for n-decane mixtures and ~12% for toluene

mixtures. On the other hand, the measured iso-octane concentrations were found to be

consistent with the values from the manometrical preparation within 3%.

C.3 Results and Discussion

C.3.1 n-Decane Oxidation

OH time histories offer an important kinetic target for the validation of detailed

reaction mechanisms during hydrocarbon oxidation. In particular, the early-time feature

of OH appears to provide important information about the breakdown of the fuel, because

this feature is controlled by the unimolecular decomposition and the H-atom abstraction

183

reactions. Fig. C.2 shows the measured OH time histories for the mixture of ~360 ppm

n-decane with 0.813% O2 in Ar (Φ ~ 0.8), along with the simulations from the detailed

mechanism of Sirjean et al. (JetSurF v1.1) [169]. On the basis of these OH time histories,

the ignition delay times simulated from the JetSurF v1.1 mechanism are slightly shorter

than the measured values at the present experimental conditions. Nevertheless, the model

can accurately predict the initial and final OH plateau levels. More importantly, the

model is able to capture the initial OH formation rates reasonably well. This indicates

that the rate constants for the n-decane unimolecular decomposition pathways and the H-

atom abstraction reactions from n-decane are quite reasonable.

Figure C.2: OH time histories for the mixture of ~360 ppm n-decane with 0.813% O2 in Ar. Simulations were done using JetSurF v1.1 mechanism. An inset figure is also shown to provide the early-time features.

C2H4 species is an important intermediate species during hydrocarbon oxidation.

Fig. C.3 shows the measured C2H4 time histories for the mixture of ~360 ppm n-decane

with 0.813% O2 in Ar, along with the simulations from the JetSurF v1.1 mechanism. The

measured C2H4 plateau levels are around 1050 ppm, which indicates that each n-decane

molecule forms approximately three C2H4 molecules prior to ignition. Similar to the OH

time histories, the model is able to capture the initial formation rates and plateau levels of

C2H4 reasonably well.

184

Figure C.3: C2H4 time histories for the mixture of ~360 ppm n-decane with 0.813% O2 in Ar. Simulations were done using JetSurF v1.1 mechanism.

During hydrocarbon oxidation, CO is considered as an important combustion

progress marker, which gives similar information to that of another combustion progress

marker, H2O. Fig. C.4 shows the measured CO time histories for the mixture of ~360

ppm n-decane with 0.813% O2 in Ar, along with the simulations from the JetSurF v1.1

mechanism. As illustrated in Fig. C.4, the peak CO mole fraction decreases with

temperature. When compared to the model predictions, the measured peak CO mole

fractions are consistently lower than the predicted values. This discrepancy between the

measurement and the model prediction is much more significant at T =1327 K. In

general, the model overpredicts the CO mole fractions prior to ignition at the present

experimental conditions. This might explain why the present model underpredicts the

ignition delay times at the current test conditions.

185

Figure C.4: CO time histories for the mixture of ~360 ppm n-decane with 0.813% O2 in Ar. Simulations were done using JetSurF v1.1 mechanism.

C.3.2 iso-Octane Oxidation

Fig. C.5 illustrates the measured OH time histories for the mixture of 511 ppm

iso-octane with 0.813% O2 in Ar, along with the simulations from the LLNL mechanism

(iso-octane mechanism v3) [170]. The measured OH time histories during iso-octane

oxidation are quite different from the measured time histories during n-decane oxidation,

particularly at early times. During iso-octane oxidation, the OH concentration first

develops a local maximum during the first 50 µs. However, shortly thereafter, the OH

concentration declines to an intermediate minimum prior to ignition, where the OH

concentration rises rapidly. It should be noted that this early-time feature of OH during

iso-octane oxidation resembles the OH feature during JP-8 oxidation. As illustrated in

Fig. C.5, the model can capture the general shape of OH, including the local maxima and

minima prior to ignition. However, at the present experimental conditions, the predicted

local maxima are higher than the measured values by at least a factor of 2. In addition,

the model seems to overpredict the intermediate minimum by a factor of ~2 at 1610 K,

but it is able to simulate the intermediate minimum at 1474 K. Similarly, the model

186

slightly underpredicts the ignition delay time at 1610 K, but the model prediction is in

good agreement with the measurement at 1474 K.

Figure C.5: OH time histories for the mixture of 511 ppm iso-octane with 0.813% O2 in Ar. Simulations were done using LLNL mechanism (iso-octane mech. v3). An inset figure is also shown to provide the early-time features.

As discussed by Davidson et al. [64], during high-temperature pyrolysis, iso-

octane primarily decomposes to form H, CH3, C3H6, and i-C4H8 species. These

intermediate species are relatively stable. In contrast to n-decane pyrolysis, less C2H4 is

formed during iso-octane pyrolysis. Hence, we would not expect to observe any

significant amounts of C2H4 during iso-octane oxidation. As expected, we are not able to

detect any noticeable levels of C2H4 for the mixture of 511 ppm iso-octane with 0.813%

O2 in Ar at the present experimental conditions.

Fig. C.6 also illustrates the measured CO time histories for the mixture of 511

ppm iso-octane with 0.813% O2 in Ar, along with the simulations from the LLNL

mechanism. Note that the measured peak CO mole fractions decrease with decreasing

temperatures, and are slightly higher than the model predictions. When compared to the

measured CO mole fractions, the model shows decent agreement with the measured

187

values at 1474 K; however, the discrepancy between the measurement and the model

becomes larger at higher temperatures, as can be seen from the OH time histories.

Figure C.6: CO time histories for the mixture of 511 ppm iso-octane with 0.813% O2 in Ar. Simulations were done using LLNL mechanism (iso-octane mech. v3).

C.3.3 Toluene Oxidation

Fig. C.7 shows the measured OH time histories for the mixture of 640 ppm

toluene with 0.813% O2 in Ar, along with the simulations from the JetSurF v1.1

mechanism. As mentioned previously, the early-time feature of OH is unique to the

individual fuel. As can be seen, the early-time features of OH during toluene oxidation

are quite distinct from those of n-decane and iso-octane. In particular, during toluene

oxidation, neither initial plateau level nor local maximum can be observed at early times,

but there is a rather gradual OH rise prior to ignition. In addition, the JetSurF v1.1

mechanism [169] seems to underpredict the ignition delay times by ~30% at 1524 and

1602 K. Nevertheless, the model is able to simulate the general shape of OH species, but

it fails to predict the post-ignition OH plateau levels.

188

Figure C.7: OH time histories for the mixture of 640 ppm toluene with 0.813% O2 in Ar. Simulations were done using JetSurF v1.1 mechanism. An inset figure is also shown to provide the early-time features.

Similar to iso-octane oxidation, we cannot detect any noticeable levels of C2H4

during toluene oxidation. This is primarily due to the fact that toluene undergoes

unimolecular decomposition to form H, CH3, C6H5CH2, and C6H5 radicals. There are no

direct reaction pathways from these radicals to produce C2H4 molecules. Instead, these

radicals will eventually form C2H2, C3H3, and C4H4 species. In addition, Fig. C.8 shows

the measured CO time histories during toluene oxidation, along with the simulations from

the JetSurF v1.1 mechanism. The model seems to underpredict the CO mole fractions at

early times; shortly thereafter, the model overpredicts the CO mole fractions prior to

ignition.

189

Figure C.8: CO time histories for the mixture of 640 ppm toluene with 0.813% O2 in Ar. Simulations were done using JetSurF v1.1 mechanism.

C.3.4 Comparison of Ignition Delay Times

Fig. C.9 shows the ignition delay time measurements for the mixtures of 424 ppm

JP-8/ 0.813% O2/ Ar (Φ ~ 0.85), ~360 ppm n-decane/ 0.813% O2/ Ar (Φ ~ 0.7), 511 ppm

iso-octane/ 0.813% O2/ Ar (Φ ~ 0.8), and 640 ppm toluene/ 0.813% O2/ Ar (Φ ~ 0.71).

In the present study, the ignition delay time is defined as the time to reach 50% of the

peak OH concentration, with time zero being defined as the arrival of the reflected shock

at the sidewall measurement location. As is evident in Fig. C.9, the n-decane mixture has

the fastest ignition delay times among all four mixtures. Interestingly, the ignition delay

time measurements for the JP-8 and iso-octane mixtures are in good agreement with each

other. When compared to the other three mixtures, the toluene mixture is the least

reactive one and has the slowest ignition delay times.

190

Figure C.9: Comparison of ignition delay time measurements for JP-8, n-decane, iso-octane, and toluene at 1.6 atm.

C.4 Conclusions

The high-temperature oxidation of n-decane, iso-octane and toluene was

investigated behind reflected shock waves using laser absorption methods to measure

time histories of three species: OH, C2H4, and CO. These species time histories reveal

that refinements in current surrogate/reaction mechanism models are still needed to

improve the predictions of the detailed mechanisms, particularly in the case for toluene

oxidation. Ignition delay time measurements for mixtures of n-decane, iso-octane, and

toluene were compared to measurements for mixtures of JP-8. Ignition delay times for

the iso-octane and JP-8 mixtures were found to be in good agreement with each other.

191

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shock tube measurements of oxygenated fuel … · ocho), methyl acetate (ch. 3. oc(o)ch. 3), methyl...

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