Hydrocarbon Combustion Chemistry-Reaction Mechanisms

This chapter provides a brief overview of the gas phase chemical kinetics associated with the oxidation of hydrocarbon fuels, mainly n-alkane fuels, i.e., methane, propane, heptane, etc. There has been a significant progress in the development of detailed reaction mechanisms, and, consequently, there is extensive literature concerning the global and detailed chemistry models for a variety of hydrocarbon fuels for a wide range of pressure and temperature, as well as for a variety of flame configurations. Therefore, the present overview is quite limited in scope, and intended to be introductory in nature. It mainly considers high-temperature oxidation chemistry of few representative fuels.

Since the combustion of hydrocarbon fuels invariably involves the formation of H2 and CO as intermediate fuels, and their subsequent oxidation to H2O and CO2, we will first discuss the oxidation chemistry of H2 and CO. Of course, the oxidation chemistry of these two fuels is also relevant otherwise.

     1.  h+o2=o+oh                                                                                    2.  o+h2=h+oh                                                                                    3.  oh+h2=h+h2o                                                                                4.  o+h2o=oh+oh        5.  h2+m=h+h+m        6.  o2+m=o+o+m        7.  oh+m=o+h+m                                                                                  8.  h2o+m=h+oh+m                                                                              9.  h+o2(+m)=ho2(+m)                                                                    10.  ho2+h=h2+o2                                                                              11.  ho2+h=oh+oh                                                                              12.  ho2+o=oh+o2                                                                              13.  ho2+oh=h2o+o2                                                                          14.  h2o2+o2=ho2+ho2                                                                      15.  h2o2+o2=ho2+ho2                                                                      16.  h2o2(+m)=oh+oh(+m)                                                                17.  h2o2+h=h2o+oh                                                                          18.  h2o2+h=h2+ho2                                                                          19.  h2o2+o=oh+ho2                                                                          20.  h2o2+oh=h2o+ho2                                                                      21.  h2o2+oh=h2o+ho2                                                                  

Reaction Mechanism for H2-O2

Reaction, NR=21 Species: Ns=7

H, O2, O, OH, H2, H2O, H2O2

€

υ i, j#

i=1

NS

∑ Mi = υ i, j##

i=1

NS

∑ Mi

j=1, 2, … NR

€

ω i =dcidt

= (υ i, j$$ −υ i, j

$)ω jj=1

NR

∑

ω j = k fj.ΠiNSci

"υi, j − kbj.ΠiNSci

""υi, j

For practice, write equations for the reaction rate of reaction # 3 and 5, and for the net rate of production of H2O

Hydrogen Combustion Chemistry

Conaire et al. mechanism A units mole-cm-sec-K, E units cal/mole

Hydrogen Combustion Chemistry

Hydrogen-Oxygen Chemistry and Explosion Limits

€

2H2 +O2 ⇒ 2H2O− 483.7kJ

The global reaction describing the oxidation of H2 to H2O is

implying that 2 moles of H2 reacts with one mole of O2 to produce 2 moles of H2O and releases 483.7 kJ of heat. Such reactions rarely occur in reality, instead there are several intermediate steps involving chain reactions. Here we discuss some important steps associated with the H2 oxidation.

Initiation Reactions:

€

H2 + M⇒ H + H + M

€

O2 + M⇒ O+O+ M

€

H2 +O2 ⇒ HO2 + H

(R5) (-R6)

(-R10)

Hydrogen-Oxygen Chemistry and Explosion Limits •  Endothermicities of these three reactions are 104, 118, and 55 kcal/mol, respectively, and

activation energies are 96, 115, and 57.8 kcal/mol respectively.

•  Thus reaction -R10 is the more important initiation reaction at low temperatures, producing relatively stable hydroperoxy radical and H, while reactions R5 and –R6 are important at higher temperatures.

•  Once H radical is produced (even in small concentration), it initiates the following chain reactions.

Chain Reactions:

€

H +O2 ⇒ OH +O

€

OH + H2 ⇒ H2O+ H

(R1) (R2) (R3)

€

O+ H2 ⇒ OH + H

•  R1 leads to R2 and R3 chain reactions. Reverse reactions (-R1, -R2, -R3) are not important initially because of the low concentrations of radical species (H, OH, O) and H2O.

•  In the context of explosion limits, the explosion is not possible for sufficiently low pressures and temperatures, since the key reaction R1 is strongly endothermic (by 66 kJ). Thus at low T, radical species (H) rapidly diffuse to the chamber wall, due to high diffusivity at low pressures, and get destroyed there.

•  Surface (wall) reactions involve heterogeneous chemistry and are not discussed here.

Hydrogen-Oxygen Chemistry and Explosion Limits

Z-shaped curve illustrating the three explosion limits

First limit: Competition between chain branching and chain termination (radical destruction at the vessel wall) reactions

Second limit: Competition between chain-branching (main:H+O2→OH+O) and chain-terminating (H+O2 +M →HO2+M) reactions. As the temperature is increased, this limit shifts to higher pressures, since the three-body reactions are more significant at higher pressure,

Third limit: At still higher pressures, the third limit is reached, as the HO2 concentration becomes significant, it leads to reactions HO2+ H2 → H2O2+H and H2O2+M→2OH+M, consuming HO2 and producing radical species.

Hydrogen-Oxygen Chemistry and Explosion Limits

€

H +O2 + M⇒ HO2 + M (R9)

First Explosion Limit: As pressure and/or temperature is increased, collision and reaction rates increase, and the rate of branching becomes more dominant compared to diffusion rate, rate of removal at the wall, etc. Consequently, the first explosion limit is crossed.

Second Limit: As the pressure is increased further, the second explosion limit is crossed. This occurs because the following three body reaction becomes more dominant compared to R1.

Since HO2 radical is relatively nonreactive, reaction R9 is effectively a chain termination reaction. This breaks the chain sequence (R1-R3). Thus, the second limit is determined by the competition between the growth (R1-R3) and destruction (R9) of the H atom. As discussed by Law*, by considering R1, R2, R3, and R9, and using steady state assumption for O and OH, the net production rate of H can be obtained as

d[H ]dt

= 2ω1 −ω9 (8)

*C. K. Law, Combustion Physics, Cambridge University Press, 2006

Hydrogen-Oxygen Chemistry and Explosion Limits This implies the second limit is given by 2k1=k9M. Using p=[M]RuT yields p=(2k1/k9).RuT. This temperature is defined as the crossover temperature. The second limit is well approximated by plotting this equation in the p-T diagram. Also rate of R1 increases with temperature, while that of R9 decreases.

Third Limit: As the pressure is increased further, the third explosion limit is crossed. This occurs because as the concentration of HO2 becomes higher, the following H2O2 reactions become important.

This again enriches the radical pool to induce explosion.

At still higher temperatures (say above 900 K), radical (H, O) concentrations are relatively high, and HO2 is consumed through reactions R11 and R12. Then explosion will always occur.

€

HO2 + H2 ⇒ H2O2 + H (-R18)

€

H2O2 + M⇒ OH +OH + M

(-R14)

€

HO2 + HO2 ⇒ H2O2 +O2

(R16)

€

HO2 + H⇒ OH +OH

€

HO2 +O⇒ OH +O2(R11) (R12)

Explosion Limits in Terms of Mass Burning Rate

Laminar Premixed Flames: Predicted mass-burning rate versus pressure Predictions based on Mueller et al. and GRI-Mech 3.0 mechanisms

Briones et al., Effect of pressure on counterflow H2–air partially premixed flames, Combust. Flame, Vol. 140, 2005.

•  First transition is due to the second limit, or the competition between chain-branching (H+O2→OH+O) and chain-terminating (H+O2 +M →HO2+M) reactions.

•  Second transition is due to the third limit.

Explosion Limits in Terms of Mass Burning Rate

Laminar flame mass burning rate versus pressure for: a) H2/O2/He mixture (phi=0.85) with dilution adjusted such that the adiabatic flame temperature is near 1600 K b) H2/O2/He mixture (phi=0.30) with dilution adjusted such that the adiabatic flame temperature is near 1400 K. Symbols represent experimental data and solid lines the present model; dashed lines the model of Li et al.

Burke et al., Comprehensive H2/O2 kinetic model for high-pressure combustion, Int. J. Chemical Kinetics, Dec. 2011.

Explosion Limits in terms of Ignition Temperature

Ignition in a counterflow configuration (strain rate:100) Fuel stream H2/CO (5%H2-95%CO). Air stream temperature is adjusted to obtain the minimum ignition temperature, which is plotted versus pressure Symbols: experimental data

Sung and Law, CST,, Vol. 180, 1097-1116, 2008

CO Oxidation Chemistry CO+O2 ⇒CO2 +O (C1)

O+H2O⇒OH +OH (C3)

CO+OH ⇒CO2 +H (C2)

CO+HO2 ⇒CO2 +OH (C4)

Reaction C1 represents the initiation step, but is a very slow reaction with high activation energy (48 kcal/mol) and does not contribute much to the formation of CO2. However, if a small amount of H2 is present, even as small as 20 ppm, it leads to the formation of OH through reactions R2, R3, and R1

€

O+ H2 ⇒ OH + HOH +H2 ⇒ H2O+H

H +O2 ⇒OH +O

R2 R3

R1

Reaction C2 is the major CO oxidation reaction. Note that with H2 present, the entire H2-O2 mechanism needs to be included in the CO oxidation mechanism

If there is a little bit of moisture in the system, reaction C3 produces OH which then feeds into reaction C2

If HO2 is present, then another route opens up through reaction C4 at high pressure

Methane Oxidation Chemistry: Major Pathways

GRI 2.11 Mechanism 277 reactions of 49 species

Dark arrows indicate additional pathways at low temperatures (<1500K)

GRI 3.0 Mechanism: See Turns 325 reactions of 53 species

• S.R. Turns, An Int. to Combustion, 3rd Ed, McGraw Hill, 2012

• C. K. Law, Combustion Physics, Cambridge University Press, 2006

Methane Oxidation Important steps involved in methane oxidation:

1. Initiation steps:

CH4 + M => CH3 + H + M (M1)

CH4 + O2=> CH3 + HO2 (M2)

2a. If M1 is the dominant initiation reaction (favored at high T due to its large activation energy), then the chain reactions are:

H + O2 => OH + O (R1 for H2 oxidation)

CH4 + (H, O, OH) => CH3 + (H2 , OH, H2O) (M3, M4, M5)

•  Reaction M3 also plays an inhibiting role as it competes with chain branching step (R1), and converts active H atoms to less active radicals CH3. This decreases mixture reactivity ; for instance the ignition delay time increases at high CH4 concentrations.

•  Reverse of M1 also plays inhibiting role as the radical pool builds up

2b. If M2 is the dominant initiation reaction, the following reactions occur

CH4 + HO2 => CH3 + H2O2 H2O2 + M => OH + OH + M (M6, M7)

Methane Oxidation 3. Now we have significant CH3 and radical pool. Conversion of CH3 to CO occurs via three main paths:

Path 1: CH3 => CH2O => HCO => CO (only main path shown)

CH3 + O => CH2O + H (formaldehyde)

CH2O + H => HCO + H2 and CH2O + OH => HCO + H2O (formation of formyl)

CHO + H => CO + H2 and CHO + O2 => CO + HO2

Path 2: CH3 => CH2(S) => CH2 => HCO (and CH)

CH2(S) has no unpaired electrons but has an empty orbital, which makes it highly energetic and active species, It becomes more stable CH2 (with two unpaired electrons) by collisions with other molecules. CH2 forms HCO and CH

CH3 + OH => CH2(S) +H2O and CH2(S) + M => CH2 + M

CH2 + O2 => HCO + OH CH2 + H => CH + H2

CH reacts with H2O and O2 to form CH2O and HCO, which form CO

CH + H2O => CH2O + H and CH + O2 => HCO + O

Path 3: CH3 + CH3 => C2H6 =>C2H5 =>C2H4 =>C2H3=>C2H2=>C2H=>CO and CH2

4. CO and H2 are oxidized to CO2 and H2O respectively

Oxidation of Higher Hydrocarbon Fuels Higher Paraffins (CnH2n+2) Oxidation (n>2)

Important steps involved in the oxidation of propane and larger alkanes are:

1. Breaking of C-C bonds and formation of alkyls: C3H8=> C2H5+ CH3

C-C bond energy is 80 kcal/mol while C-H bond energy is 98.

n-C7H16=> C6H13+ CH3, C5H11+ C2H5, C4H9+ C3H7

2. Alkyl radicals undergo hydrogen abstraction reactions, generating H radicals:

C2H5 +M=> C2H4 + H + M and CH3 +M=> CH2 + H + M

Similarly, C6H13 +M=> C6H12 + H + M and other alkyls

3. H atoms start the creation of radical pool through reaction: H +O2=> OH + O

4. Radicals (H, O, OH) react with fuel molecules through H-abstraction reactions: RH + (H, O, OH)=>R+ (H2 , OH, H2O), and produce alkyl radicals.

Rate of H-abstraction reaction depends on the type of CH bonds, i.e., primary (p), secondary (s), or tertiary (t) bonds; Primary C-H bond has 1C atom bonded with 1C and 3H, Secondary C-H bond has 1C atom bonded with 2C and 2H, and Tertiary C-H bond has 1C atom bonded with 3C atoms.

Larger (n>2) Alkane Oxidation Thus propane has 6 p and 2 s bonds, while n-heptane has 6 p and 10 s bonds, and 2 methyl butane has 9 p, 4 s, and 1 t bond. Generally, s and t bonds are relatively weaker, and breaking of these bonds have lower activation energy. However at high temperatures, breaking of p, s, t bonds occur at about the same rate. However, H abstraction reaction leads to different products, depending on the type of C-H bond. For example, breaking of p and s bond in propane produces n-propyl (n-C3H7) and iso-propyl (i-C3H7)radicals, respectively.

5a. Alkyl radical species (C7H13, C5H11, C3H7, C2H5 etc.) produce olefins, and smaller hydrocarbon radicals via H abstraction and breaking of C-C bonds: C3H7 + M=> C3H6+ H +M, C3H7 + H=> C3H6+ H2 and C3H7 + M=> C2H4+ CH3

These reactions follow the β scission rule: The C-C or C-H bond (β bond) broken is the one that is one place removed from radical site (site of unpaired electron), which strengthens the adjacent bonds.

Note: The C-H and C-C bonds adjacent to the radical site have bond energies of 99 kcal/mol and >100 kcal/mol, while the C-H and C-C bonds one place away have bond energies of 34 and 24 kcal/mol, respectively. Thus

C7H15=> C2H4+ C5H11, => C7H14+ H, => C3H6+ C4H9, etc.

Larger (n>2) Alkane Oxidation 5b. Larger alkyl radicals then dissociate to smaller hydrocarbons (olifins and alkyls) through similar β-scission reactions.

6. Oxidation of olifins through reactions with O producing formyl (CHO) and formaldehyde (CH2O). For example: C3H6 +O => C2H5+ CHO and C3H6 +O => C2H4+ CH2O

Propene also undergoes H abstraction to produce allyl radical (C3H5), which is then oxidized by O. OH, O2,and HO2 producing oxygenated or nonoxygenated C1 and C2, and H2 species

7. Now the system contains species that involve reactions of H2, CO, CH4 and C2Hx mechanisms.

8. CO oxidizes to CO2 through CO oxidation reactions (discussed earlier), with the dominant reaction: CO + OH => CO2 + H. Similarly H2 formed through various steps oxidizes to form H2O through H2 oxidation reactions.

Notes •  Basic steps are similar for the oxidation of higher aliphatic hydrocarbons •  Sub-mechanisms for the oxidation of H2, CO, CH4, and C2Hx are the building

blocks for the oxidation of higher aliphatic hydrocarbons

Propane Combustion in a Steady Flow Reactor

Note the reactant mixture properties at the reactor entrance boundary

NTC Ignition Behavior for Alkanes (N>2) The ignition chemistry of alkanes at intermediate temperatures is characterized by the existence of a NTC (negative temperature coefficient) region, or cool flame behavior, where the ignition delay increases as T is increased. This region is important in many practical systems, such as autoignition in CI engines. The NTC region is related to the fuel oxidation chemistry at low to intermediate temperatures.

Ignition delay versus inverse of temperature for different blends of n-dodecane, H2, CO, and syngas

Aggarwal, Helma & Li, Int J Adv Eng Sci Appl Math (2014), 6:49-64