Complex Mechanisms

There are many intermediates between fuel and oxidizer and equilibrium products

Even a mechanism of simple molecules combustion of hydrogen

2 H2 + O2 !" 2H2O

Is a complex interaction of many intermediates

Combustion mechanisms are complex with many intermediates, stages and hierarchies between the fuel and oxidizer and the final equilibrium products.

Even a simple molecule such as hydrogen is a complex mechanism with many intermediates, reactions and submechanisms.

Hydrogen Combustion

H2 + M !" H + H + M

H2 + O2 !" HO2 + H

H + O2 !" O + OH

O + H2 !" H + OH

H2 + OH !" H2O + H

O + H2O !" OH + OH

H + H + M !" H2 + M

O + O + M !" O2 + M

H + O + M !" OH + M

H + OH !" H2O

H + O2 + M !" HO2 + M

HO2 + H !" OH + OH

HO2 + H !" H2O + O

HO2 + O !" H2O + OH

HO2 + HO2 !" H2O2 + O2

HO2 + H2 !" H2O2 + H

H2O2 + OH !" H2O + HO2

H2O2 + H !" H2O + OH

H2O2 + H !" HO2 + H2

H2O2 + M !" OH + OH + MThe hydrogen mechanism is the base of all hydrocarbon mechanisms. At the first level, the oxidation of hydrogen can be divided into two submechanisms. The rate constants for these reactions are standardized with accepted values.

The first submechanism involves combinations of simple hydrogen and oxygen species, H, H2, H2O, O, O2, OH, and H, including some pressure sensitive reactions.

The second submechanism involves the oxidation of H2O2 and involves the additional intermediate HO2.

Bonding of Hydrogen

H

H

Hydrogen Radical

Single Bonded Hydrogen

Hydrogen: Single Valence Electron

The first submechanism of the hydrogen mechanism involves essentially all combinations of simple oxygen and hydrogen radicals. To understand them, we can look at the valence structure of each.

Hydrogen has one valence electron. The one electron is either alone making the hydrogen radical or shared to bond with another atom.

Bonding of Oxygen

O

O

O

Oxygen: 6 valence electrons

Double Bonded Oxygen

Single Bonded Oxygen

Oxygen Atom (singlet or triplet)

The oxygen atom has 6 valence electrons.

The electrons can be single in a radical, paired in a lone-pair or shared with another in one or two bonds.

There are 3 types of bonding that oxygen can be a part of: - Oxygen Atom Two electron pairs and two unpaired electrons - Single Bond A single bond, a single electron and two electron pairs. - Double Bond Two bonds, to electron pairs.

Mechanism Structure

Not just a list of reactions

Structured set of submechanisms

Also different reactions are important under different conditions

The hydrogen mechanism is not just a list of reactions. It is, in fact a structured set of submechanism and regimes that are significant under di!erent conditions.

H2, O2, H, O, OH Reactions

Initiation Reactions H2 + M !" H + H + M(fuel and oxidizer) H2 + O2 !" HO2 + HChain Reaction Steps H + O2 !" O + OH(increase of radicals) O + H2 !" H + OH

H2 + OH !" H2O + HO + H2O !" OH + OH

Chain Terminating Steps H + H + M !" H2 + M(decrease of radicals) O + O + M !" O2 + M

H + O + M !" OH + MH + OH !" H2O

Viewed as a chain branching process, we can see the di!erent types of reactions based on the number of radicals created and destroyed between reactants and products.The initiation reactions are those where stable species such as H2 and O2 form radicals.The chain reaction steps are those in which one radical species in the reactants produces one or two radicals in the products. If these are prevalent, then the hydrogen mechanism moves faster toward ignition.The chain terminating steps are when two radicals in the reactants come together to form at least one stable species. When these reactions are dominant, then reaction progress is slowed or stopped.

HO2, H2O2 Reactions

Transition to HO2 H + O2 + M !" HO2 + MHO2 Reactions HO2 + H !" OH + OH

HO2 + H !" H2O + OHO2 + O !" H2O + OH

Tranisition to H2O2 HO2 + HO2 !" H2O2 + O2

HO2 + H2 !" H2O2 + HH2O2 Reactions H2O2 + OH !" H2O + HO2

H2O2 + H !" H2O + OHH2O2 + H !" HO2 + H2

H2O2 + M !" OH + OH + M

The first submechanism of simple oxygen and hydrogen radicals has a transition through recombination to the mechanism with more complex species such as HO2 and H2O2. Since HO2 is a radical, the formation of HO2 are chain branching reactions. The formation of H2O2 can be both terminating and branching.

Surface Reactions

2 HO2 !" H2O2 + O2

H2O2 !" H2O +12O2

H,O, OH !" destruction

To fully explain the ignition behavior of hydrogen, it was necessary to augment the reactions between intermediates with further catalytic reactions with the walls. All these reactions tend to slow the progress toward ignition.

Complex Behavior

Different reactivity in different regions

The three limits representthree regimes

Subsets of reactionsdominate

Under di!erent regions, di!erent reactions dominate. If the terminating reactions dominate, then ignition is inhibited, but if branching reactions dominate, ignition is promoted.

The graph shows the di!erent ‘limits’ of hydrogen ignition with respect to temperature with increasing pressure. Each regime represents a certain subset of reactions that are dominate.

First Limit

The explosion limit is dependent on the size of the vessel

The inhibitors are the destruction

of radicals, H, O, OH at the walls.

Very low pressures

The first limit is at very low pressure, even below atmospheric pressures. Experimental evidence of this is that the explosion limit is dependent on the size of the vessel. This points to the destruction of radicals at the walls.

First Limit

As the pressure is going down Something is inhibiting the chain

reaction steps

H + O2 !" O + OHO + H2 !" H + OH

This means that the number of collisions to produce the chain reaction steps is low. At lower pressures, the reactions with the walls are more prominent. A reaction at the wall acts as a termination reaction, thus lowering the explosive tendency.

Remember, if the termination reactions dominate, explosion does not occur.

Second Limit

Increase of Pressure inhibits the explosionhigher temperature needed to explode

not dependent on the size of the vessel Thus not inhibited by walls

At higher pressures, the size of the vessel does not play a significant role.

But as the pressure increases, a higher temperature is needed to create an explosion.

At high pressures, three body reactions can dominate. Thus reactions with M, which are generally dependent on pressure, can be significant.

As pressure increases, the relative rates of terniary reactions increases more than binary reactions.

Second Limit

A gas phase reaction that is strongly dependent on pressure.

Competition between

H + O2 !" O + OH (more radicals produced)

and

H + O2 + M !" HO2 + M (a different radical produced)

Higher pressure is larger M concentration

HO2 less reactive This can diffuse to wallsWithin the hydrogen mechanism, at the pressures above the second explosive limit, HO2 is relatively unreactive. Thus its formation can be thought to be a termination reaction. It is able to migrate to the walls.

The inhibition of the HO2 reaction dominating above the second explosive limit, means that it requires a higher temperature to explode. The peak pressure, which requires the highest temperature is around 300 atmospheres.

Third Limit

As the pressure increases, HO2 reacts with H2

HO2 + H2 !" H2O2 + H

H2O2 + M !" OH + OH + M

In addition, H2O, the major final product is an excellent third body

In the third limit, at higher pressures, less temperature is needed to result in explosion. This means that the inhibiting e!ect of the second limit, namely the HO2 inhibition, is not significant.

As the pressure increases even the inhibiting e!ect of HO2 formation is countered with the formation of H2O2 which in turn is a significant source of hydroxyl radicals, in fact two for every H2O2 molecule.

CO Oxidation

CO Mechanism Propagation CO + HO2 !" CO2 + OHPropagation CO + OH !" CO2 + H

H2 Mechanism Branching H + O2 !" OH + OPropagation H + O2 + M !" HO2 + M

CO Mechanism Termination CO + O + M !" CO2 + M

The following are the major steps for CO oxidation, made up of branching, propagating and termination steps.

The two propagation steps of the CO mechanism produce CO2 the final equilibrium product. But what is important is that at the same time the number of radicals, in fact important radicals such as H and OH, in the system is not diminished (hence propagation).

Further propagation and branching are found the two important reactions in the hydrogen mechanism. H reactions with the original oxidizer, O2, react to produce radicals important for the two CO oxidation reactions, namely, OH and HO2, one of which is a branching reaction bringing more radicals into the system.

This complementary set of reactions substantiates the observation that H is a catalyst for CO oxidation.

Hierarchical Structure

CO Mechanism

H2 Mechanism

CO2

H2OOxidation

Hydrocarbon

Combustion mechanisms have a clear hierarchal structure. At the base is the hydrogen mechanism, where many of the critical radicals driving the ignition are formed, such as OH and the carbonmonoxide submechanism.

On top of these mechanisms are the hydrocarbon mechanisms whose products produce precursors for these base mechanisms.

Carbon Bonding

C

C

C C

C

Single Bonded Carbon

Methylene: Radical Carbon

Double Bonded Carbon

Triple Bonded Carbon

(Singlet and Triplet)

C Double Bond with radical

C Triple Bond with radical

In order to understand the multitude of hydrocarbon reactions, it is useful to know at least the basic (lewis) bonding types of carbon. Hydrocarbon mechanisms are essentially many combinations of transformations between these states.

The carbon has four electrons available for bonding (in its valence shell). The main stable carbon bonded states have either four single bonds, a double bond with two single or a triple bond with a single bond. The simple radicals can be these same configurations, where one of the bonds is replaced by an unbonded electron.

When it has two unbonded electrons, the carbon can be in two states, a singlet or triplet, both having two single bonds.

CH4 High Temperature Oxidation

AbstractionOH,H,O

Addition of OLoss of H

CH3

CH4

CH2O

AbstractionOH,H,O

CO2

Aldehyde

HCOH Loss Third Body (M, H20)

H Abstraction with HO2

CO

CO Oxidation Mechanism

Backbone of

Methane Oxidation

Mechanisms for methane oxidation can be quite large, but the main pathways will be explained in the following slides.

Methane oxidation, in detail, can be complex and with hundreds of reaction. Many mechanisms exist in the literature, but all have the same major features.

After an initial abstraction of hydrogen, and with the addition of an oxygen (and loss of hydrogen), an aldehyde is formed. This aldehyde decomposes to carbon monooxide. Through the carbonmonoxide submechanism, the final equilibrium products are made.

This is the backbone of the methane oxidation.

CH4 High Temperature Oxidation

CH3

CH4

CH2O

CO2

CH2

CH2

CH

(S)OH,H Abstraction

OH Abstraction

OH AdditionLoss of H

H20 AdditionLoss of H

Side path withCH2 Radicals

A significant side reaction to the main path is further abstractions through triplet and singlet CH2 radicals, which, with subsequent addition of oxygen, through OH or water, form an aldehyde, within the backbone mechanism.

CH4 High Temperature Oxidation

CH3

CH4

CH2O

CO2

CH2

CH2

CH

(S)CH

2OH

OH AdditionH Loss

Third BodyH Loss

CH2OH

Formation

From CH3, this set of side reactions go through the CH2OH as an intermediate to the aldehyde. The CH2OH is formed through the addition of OH radicals with the simultaneous loss of H.

CH4 Low Temperature Oxidation

CH3

CH4

CH2O

CO2

CH3OH

CH2

CH2

CH

(S)CH

2OH

Reverse

HCO

2O

H AbstractionH,OH

H,OHH Abstraction

CH3

OThird Body

H Loss

C2H

4COCH

2Abstractions

Addition

3CH

Two significant paths that appear in the low temperature ( < 1500) oxidation of methane.

One is the recombination reaction forming C2H6. This is why the propane submechanism, although a larger product, is an important to methane oxidation.

Another submechanism involves the direct addition of OH to form CH3OH and its subsequent oxidation to aldehyde.

Aldehyde Decomposition

C

O

HH

X XH

H

C

O

H

X XH

H

C

O

R

C

O

HR

Aldehydes are an important intermediate from hydrocarbons to CO to CO2. This shows the electron structure of the conversion of an aldehyde to the radical through abstraction or loss of H. The aldehyde starts as a double bond with two single bonds, one of which is to hydrogen. The hydrogen is abstracted or lost and the aldehyde radical is produced with a doubly bonded carbon with a single electron radical and a single bond.

Competition for OH

Any competition for the radicals, OH, HO2 and O reduces efficiency

Reaction 1000K 1500K 2000KAbs: CH4 + OH !" CH3 + H2OProd: CO + OH !" CO2 + HRatio Abs

Prod 4.6 9.2 13.3

In fuel rich environments, CO may remain as a product.

The main driving force of combustion reactions is the number of radicals in the system. Two of the most important of these is OH and H.

In methane combustion, in two propagating reactions, OH can be used in abstraction such as from CH4 or with CO to produce the final equilibrium product, CO2.

When abstraction produces CH3, to oxidize to equilibrium products the mechanism has to go through the CO oxidation, which again can be inhibited or side-tracked, inhibiting the oxidation of CO.

This competition can inhibit the oxidation of CO and, for example, in fuel rich environments, CO may be a more prominent product. Meaning that methane combustion does not go to completion.