Table of Contents

1. Introduction............................................................5

1.1. Options to Conventional Transport Fuels..................................................6

1.2. Reasons for Promoting Non Fossil Fuel Derivatives...................................8

1.3. Comparison of Non Fossil Fuels.................................................................9

1.4. Hydrogen................................................................................................10

1.5. Main Aims of Study.................................................................................12

2. Chemical Kinetics and Combustion Modelling..........14

2.1. Kinetic Reactions: Chemical and Thermal Events...................................15

2.2. Chemical Modelling Principles.................................................................16

2.3. Levels of Modelling..................................................................................17

2.3.1. Macroscopic Level.............................................................................17

2.3.2. Microscopic Level..............................................................................17

2.4. Kinetic Mechanisms.............................................................................18

3. Chemical Kinetics of Hydrogen-Air Combustion.......19

3.1. Overview of Mechanism..........................................................................19

3.2. Bonding Structures.................................................................................21

3.2.1. Bonding of Hydrogen........................................................................21

3.2.2. Bonding of Oxygen...........................................................................21

3.3. Mechanism Structure..............................................................................22

3.3.1. H2 O2 H OH Reactions......................................................................22

3.3.2. HO2 H2O2 Reactions........................................................................23

3.4. Complex Behaviour.................................................................................24

3.4.1. First Limit..........................................................................................26

3.4.2. Second Limit.....................................................................................26

3.4.3. Third Limit.........................................................................................27

3.5. Validation of Mechanisms.......................................................................27

3.5.1. Gas Research Institute.....................................................................27

3.5.2. Yetter et al.......................................................................................28

3.5.3. Mueller et al.....................................................................................28

3.5.4. Li et al..............................................................................................28

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Table of Contents

4. Chemical Kinetics of Methane Oxidation.................30

4.1. Carbon Bonding......................................................................................30

4.2. Hydrocarbon Chemistry – Explosion Limits.............................................32

4.3. Methane High Temperature Oxidation....................................................34

4.4. Side Path with CH2 Radicals....................................................................36

4.5. CH2OH Formation....................................................................................37

4.6. CH4 Low Temperature Oxidation............................................................38

4.7. Aldehyde Decomposition........................................................................39

4.8. CO Oxidation...........................................................................................40

5. Laminar Burning Velocity.......................................41

5.1. Case Study..............................................................................................42

5.2. Results....................................................................................................45

5.3. Discussion...............................................................................................47

6. Conclusions...........................................................52

7. References............................................................53

Appendices.............................Error! Bookmark not defined.

GOVERNING EQUATIONS FOR COMPUTATIONS.Error! Bookmark not defined.

INVESTIGATIVE WORK-ASSESSING MESH SIZE DEPENDANCE ON COMPUTATED RESULTS...........................................................Error! Bookmark not defined.

MESH SIZE – ON FIXED 0.01m.......................Error! Bookmark not defined.

MESH SIZE – ON FIXED 0.02m.......................Error! Bookmark not defined.

MESH SIZE – ON FIXED 0.03m.......................Error! Bookmark not defined.

List of Figures

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Figure 1 (a): UK emissions by sector...........................................................................................3

Figure 1.3 (a): Heating content of various fuels...........................................................................7

Figure 1.4 (a): Comparison of properties of hydrogen and conventional fossil fuels..........8

Figure 1.4 (b): Carbon dioxide release: Gas vs.

Hydrogen...................................................................9

Figure 1.5 (a): Laminar burning velocity for methane-, ethylene- and hydrogen-air............10

Figure 3.1(a): Simplified version of H2/O2 Mechanism.............................................................20

Figure 3.4 (a): Explosion Limits of H2/O2 mechanism................................................................23

Figure 3.5 (a): Detailed H2/O2 mechanism....................................................................................27

Figure 4 (a): Hierarchical structure of hydrocarbon oxidation...............................................28

Figure 4.2 (a) Explosion limits of Methane, Ethane and Propane..............................................30

Figure 4.3 (a): Backbone of methane oxidation under high temperature.................................33

Figure 4.4 (a): Side path reaction with CH2 radicals.....................................................................34

Figure 4.5 (a): Aldehyde formation...................................................................................................35

Figure 4.6 (a): CH4 low temperature oxidation..............................................................................36

Figure 4.7(a): Aldehyde decomposition to radical.........................................................................37

Figure 5.1 (a): Geometry for computation.......................................................................................41

Figure 5.1 (b): Geometry in x-y plane: 0.03m mesh size...............................................................42

Figure 5.3 (a): Laminar burning velocity of hydrogen–methane/air mixtures........................45

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Figure 5.3 (b): The normalised laminar burning velocity and maximum H concentration...48

List of Graphs

Graph 1.1 (a): Global fossil carbon emissions by type.................................................................5

Graph 5.2 (a): Mass Fraction of H2 vs. Curved Length..................................................................43

Graph 5.2 (b): Mass Fraction of CH4 vs. Curved Length...............................................................44

List of Tables

Table 1.1 (a): Major fossil fuel/non fossil fuel alternatives........................................................4

Table 4.2 (a): Different mechanisms for methane oxidation from various sources..............32

Table 5.3 (a): Reactant mole fraction of calculated flames.........................................................47

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1. Introduction

Securities of supply and climate change are high on the global energy agenda. Transport emissions are the fastest rising cause of greenhouse gases. The emission by sectors in the U.K is shown below:

Figure 1(a): UK emissions by sector [Wish, V. (2009)]

From the above [Figure 1(a)], transport contributes about 20% of the overall U.K emissions. The numbers of cars in the world are growing rapidly.

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In general, four options are considered to reduce these emissions:

Alternative fuels Reducing the need to travel Promotion of public transport Improving fuel efficiency

A major influence in road congestion and accidents currently is convenience. The ease of mobility that personal transport offers is favoured by the majority, as opposed to public transport or less driving, albeit facts about health risks and fuel prices.

Extreme difficulty is faced when trying to promote mileage reduction strategies and the automobile industry does not play a part in the influence for fear of harming its consumers. For these reasons, most transportation emission reduction programs focus on changing vehicle and fuel type rather than the amount people drive.

This drives the need to find alternates to conventional fuels, which could consequently aid the U.K in abiding to stringent emission regulations and reduce its dependence on oil from politically unstable regions.

However, modern society is driven by its dependence on oil to fuel its transport needs. With increasing concern over fossil fuel shortages, researches on improving thermal efficiency in order for vehicles to run on cleaner fuels to produce fewer harmful emissions are being conducted, as they offer some savings on fuel costs and counteract the issue high power demand.

1.1. Options to Conventional Transport Fuels

At present, about 80% of the world’s demand for transportation fuels, either by road, rail, air or sea; are met by derivatives from the fossil fuel, petroleum [Day, S. (2012)].

Society is currently bombarded with information on peak oil, alternative fuels and fuel and greenhouse gas saving vehicle technologies. All have implications for fuel price, the environment, global warming and ultimately, the structure of our society. The perennial dictum however, is deciding between these options.

Vehicular emissions can cause environmental harm and have been linked to atmospheric pollution which contributes to climate change via the

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Greenhouse Effect. Present greenhouse effect is caused mainly by increasing concentrations of CO2.

The table below summarises the major fossil fuel and non fossil fuel alternatives:

Table 1.1(a): Major fossil fuel/non fossil fuel alternatives

FOSSIL FUEL NON FOSSIL FUEL

Diesel Ethanol

Liquid Petroleum Gas (LPG) Biodiesel

Natural Gas Hydrogen

Considering fossil fuels, ultra-low sulphur versions of petrol and diesel represent options to the traditional versions, as they are around 97% cleaner, which imply they are better for the environment and operate at higher efficiencies in engines [Brevitt, B. (2002)].

However, to encourage their adoption, governing bodies have to cut excise duty rates relative to the standard product. Also, dwindling supplies of fossils is driving the quest for sustainable or non fossil fuel derivatives.

Natural gas may prove to be the most important short term option because it produces fewer emissions than other conventional fossil fuels. This is depicted below:

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Graph 1.1(a): Global fossil carbon emissions by type [Heinzerling, A. (2010)]

On the other hand, successful implementation of natural gas gives rise to some problems. These are discussed below:

Natural gas is a finite fossil fuel meaning its supply is limited Natural gas in itself is mainly methane, which is a greenhouse gas.

Although there is less of it in the atmosphere, it is far more powerful in trapping heat than carbon dioxide (CO2).

When natural gas is burnt, it produces CO2 and water. On the contrary, natural gas industry pumps methane into the atmosphere when mining and through huge pipelines.

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1.2. Reasons for Promoting Non Fossil Fuel Derivatives

Non fossil fuels are considered renewable energy sources hence contribute to sustainable development. The derivatives possess certain characteristics that are beneficial in principle:

Reduction in external energy dependence, as they can be produced domestically

Stabilisation of fossil fuel prices, as they help constrain the growth of petroleum prices

Reduction of greenhouse gas emissions, by recycling carbon from the atmosphere and having relatively cleaner emissions

Additional source of income, hence potentially boosting the economy.

Non fossil fuels have the potential to be an environmentally friendly option to traditional fossil fuels, and even prove to have higher efficiencies when combusted.

Ongoing research by the combustion and engine development society are currently being conducted into making non fossil fuel derivatives less expensive and more readily available, in order to make it a viable competitor in the fuel industry.

1.3. Comparison of Non Fossil Fuels

Majority of non-fossil fuel derivatives are seen as ideal fuels for transportation purposes due to their simple chemical structure and no fuel evaporation problems, implying they emit less carbon monoxide and carbon dioxide. Other advantages include:

They can be used in diesel engines, which operate at higher efficiencies than gasoline engines

They contain virtually no sulphur Ignite more readily in engines but less readily in the atmosphere.

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However, in terms of transportation, a key performance factor of a fuel is energy (heat) content. The heating content may be defined as the amount of energy released from combustion of a unit quantity of fuel.

The heating content of various fuels are depicted below:

Figure 1.3(a): Heating content of various fuels [Gossen, R. (2009)]

As noticed from [Figure 1.3(a)] above, hydrogen has the highest heating content, implying it contains more energy potential than any other fuel source.

1.4. Hydrogen

Hydrogen is a fuel which has a potential for future use in transportation. It holds significant promise as a supplement fuel to improve performance and emissions of spark ignited and compression ignited engines.

Transportation applications for hydrogen include buses, trucks, passenger vehicles, aircrafts, and trains, with technologies being developed to use hydrogen in both fuel cells and internal combustion engines, including methanol systems. Almost all major carmakers have a hydrogen-fuelled vehicle demonstration programme.

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Hydrogen-fuelled internal-combustion engine vehicles are viewed by some as a near-term, lower-cost option that could assist in the development of hydrogen infrastructure and hydrogen storage technology [Turner, J.A. (1999)]. A key advantage of this option is that hydrogen-fuelled internal-combustion engines vehicles can be made in larger numbers

The table below shows the properties of hydrogen in comparison to traditional fossil fuels methane and gasoline, which are currently used in transport engines.

Figure 1.4(a): Comparison of properties of hydrogen and conventional fossil fuels [Skottene M., Rian K.E. (2007)]

Hydrogen has several important properties that have an impact on its applicability as a fuel.

It combines with oxygen (O2) to form water. It has high energy content per weight and the flame velocity of

hydrogen is much faster than other fuels allowing oxidation with less heat transfer to the surroundings. This improves thermal efficiencies.

Efficiencies are also improved because hydrogen has a very small gap quenching distance, allowing fuel to burn more completely.

Hydrogen has the ability to burn at extremely lean equivalence ratios. Hydrogen will burn at mixtures seven times leaner than gasoline and five times leaner than methane [R. Choudhuri Ahsan & S.R. Gollahalli].

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Hydrogen is highly flammable and only requires a small amount of energy to ignite it and make it burn. It also has a wide flammability range hence can burn when it makes up 4% to 74% of the air by volume [S. Gauthier, E. Lebas and D. Baillis.].

The combustion of hydrogen does not produce CO2, particulate or sulphur dioxide (SO2) emissions. A graph showing the CO2 release of gasoline compared to hydrogen is shown below as justification:

Figure 1.4(b): Carbon dioxide release: Gas vs. Hydrogen [Douglas, K. (2008)]

1.5. Main Aims of Study

Hydrogen may step into the car transport in a few years time, and hydrogen and conventional transport fuels will co-exist in the energy market in the near future.

This project will carry out a study on the chemical kinetics of hydrogen addition to methane, and will examine the effects of mixing on the laminar burning velocity of the mixture.

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One of the important physical parameters in flame propagation is the laminar burning velocity. Knowledge of the laminar burning velocity can be used to determine a mixtures reactivity, diffusivity and exothermicity. The laminar burning velocity is also an important prerequisite in understanding turbulent combustion.

Below displays the laminar burning velocities of methane, hydrogen and ethylene in air:

Figure 1.5 (a): Laminar burning velocity for methane-, ethylene- and hydrogen-air. [Choudhuri, R. & Gollahalli,, S.R (2000)]

As noticed from [Figure 1.5 (a)], the laminar burning velocity of hydrogen is significantly higher than the rest of the other fuels. However, there are few studies which depict what happens when the fuels are mixed.

From this, the program of work is as follows:

1. A literature review will be carried out based on the detailed chemical kinetic reactions involved in the combustion of hydrogen.

2. A literature review will be carried out based on the chemical kinetics of hydrogen combustion in air.

3. A literature review will be carried out to study the laminar burning velocities of hydrogen-methane mixtures in air. The effects of

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hydrogen addition on the mixtures laminar burning velocities will also be discussed.

2. Chemical Kinetics and Combustion Modelling

Combustion affects all aspects of society, from transportation requirements to heating requirements. A current concern in modern society is to reduce energy dependence as well as explore techniques into mitigating global warming. For these reasons, the study of combustion mechanics is fundamental in improving efficiencies of all these processes.

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At present, emphasis circulates around the transport sector, as mentioned previously. The development of clean energy, in terms of alternate renewable fuels, highlights the need to study chemical kinetics, so as to understand the chemical origins of emissions. This is a vital approach in optimisation to reduce such emissions.

Experimental chemical kinetics encompasses various levels of detail, from basic kinetic experiments to complex systems which look into combustion behaviour within engines.

Focusing attention on many intermediate species, which are evolved during the combustion process, zero dimensional studies may be employed in the chemical kinetic modelling, paying key attention to ignition reliance on temperature and pressure, as well as how the concentration of species within the system changes over time.

The chemical kinetic mechanisms or models within such complicated systems can be quite multifaceted, with thousands of reactions and species. Simplification, however, is sometimes required as physical conditions get more complex. Recent developments in computational modelling aid in this understanding.

Modelling occurs at various levels of detail, with the number of molecules ranging from tenths for simpler molecules and mechanisms or hundreds even thousands, for larger, more complex chemical systems such as longer chain hydrocarbon fuels.

Of major importance is the interaction between the molecules, particularly the reactions. Reactions essentially depict molecular interaction, which can also display molecular structure. These details are usually accompanied with the thermodynamics of each molecule. The primary objective in these detailed mechanisms is inevitably the enormous network of reactions.

.

2.1. Kinetic Reactions: Chemical and Thermal Events

As a major product from many chemical processes is energy, a primary concern faced when modelling entails energy production from reacting species within a system.

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For global reactions involving fuels, a major concern might involve the formation of products, which will take into account thermal properties like the energy content of reacting species in the flame as well as the adiabatic temperature.

The concept of chemical and thermal events of kinetic reactions is evident in combustion literature, [Westbrook, CK & Dryer, FL (1984)], [Jachimowski, C.J. (1993)], etc.

Combustion processes or systems may be viewed as thermal events. In this instance, the production of thermal energy is of major importance. The chemistry that takes place to evolve this thermal energy is less important. In numerous combustion scenarios, the chemical process proceeds rapidly in comparison to other processes, for instance, nuclei diffusion. In this instance, two analogies of the chemistry can be derived regarding the fuel and oxidiser: before they react; and after the reaction is complete and products are evolved. The on goings in between these processes occurs too quickly to be able to tell apart.

However, in certain situations, knowledge about just the thermal properties can be insufficient. With raising alarm for establishing more efficient combustion processes which would consequently result in fewer emissions, the actual chemistry of combustion processes must be considered.

Understanding the true origins of harmful emissions (pollutants) involves detailed chemical kinetic models of the harmful species emitted as well as any intermediate reactions and precursors. Conventional methods utilising straightforward chemistry often does not give an ample amount chemical detail to adequately model the chemical sources and consequences.

For these reasons, combustion processes may also be viewed as chemical events. This takes into account individual reactions and intermediate species.

Understanding the complete chemistry gives a better comprehension of the rate of reaction progression, the source of thermal energy and the origins of harmful emissions.

Monitoring the behaviour of individual species under a variety of conditions leads to further understanding of the combustion process, and can give insight on how to improve the efficiency and determine under what conditions various additives can be harmful.

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2.2. Chemical Modelling Principles

The foundation of all chemical kinetic modelling, points to the actual models of molecules. Due to different levels of complexity, the question arises as to what level of modelling is suitable. This is usually based on need and results, as well as what is computationally available and affordable.

[Tomlin et al (1997)] conducted studies on the comprehensive chemical kinetics on low temperature combustion via logical programming. Based on this mathematical investigation, it was resolved that the models yielded a high degree of detail of the electronic structure, giving essential information about bonding inside the molecule, together with the three dimensional structure of the molecule. On the other hand, a computation for larger molecules, as in some combustion models, is computationally unaffordable due to the high degree of detail. In light of this, recent developments in computer aided design accommodates for larger numbers of molecules, although it may be time consuming.

At a simpler level, the molecule can be viewed as atoms connected together. Embedded in this description, is some distilled information about how the visual representation of molecules and how they react. This type of model incorporates the valence structure of the molecule into a computationally simpler form and is used to produce combustion mechanisms.

Even simpler, a model can be treated as a label. Knowledge about how it reacts unreservedly gives information about structure. This type of model is used in detailed kinetic computations in combustion.

An essential concept in modelling is the state of a system. This stems from the need to translate real world objects into mathematical models, that eventually through programming, could be used to represent the dynamics of the system. It encompasses the objects/species within the model and their descriptions.

2.3. Levels of Modelling

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Usual descriptions of the combustion process as a whole includes the Macroscopic and the Microscopic view.

Individual species and their descriptions are considered in the microscopic view. Interactions of these molecules are envisioned as collisions, resulting in the making and breaking of bonds to form different sets of molecules.

Taking into consideration an enormous number of these molecules however, results in a conglomerate of properties. In this instance, the macroscopic view might be a simpler way of examination. At this level, sets of molecules are investigated as opposed to individual molecules.

Combustion modelling makes use of both levels of modelling, as they are not exclusive but to a certain extent coincide. Each individual level helps elevate understanding of processes occurring in combustion.

2.3.1. Macroscopic Level

More global quantities are considered at this level, like temperature and pressure changes, corresponding to heat release from the process. At an engineering level, modelling is done at the macroscopic level because of the simplified modelling involved.

2.3.2. Microscopic Level

If the origins of macroscopic properties are to be understood more thoroughly, then the microscopic level has to be considered. As a result of bond making and bond breaking, alterations in individual molecules can be examined at the molecular level by conducting detailed kinetic studies. These bond making and breaking methods will usually result in energy being released or absorbed, and consequently affect the global temperature of a system. Another consequence to bond making and bond breaking is pressure, as the number of molecules within the system changes.

More importantly, the emphasis in the combustion and fuel community is on reducing emissions, understanding kinetic systems at microscopic levels are imperative.

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The origins of pollutants such as soot and NOx can only be investigated at this level, as the individual source molecules involved can be identified. Understanding the source of pollutants can lead to investigations on conditions, under which fewer pollutants are formed.

2.4. Kinetic Mechanisms

A kinetic mechanism has the complete information based on numerical calculations about how species in a reactive system combust.

It consists of:

Rate constants, giving information about molecular interaction Labels, which are used to represent molecules. Usually associated

with labels are sets of information regarding thermodynamics A list of reactions, with numbers increasing as the hydrocarbon gets

larger and the rate constants for each reaction

The structure of the molecule can be detailed from the set of reactions as well as thermodynamic information. The origins of this information are either directly from experiment or equivalence with similar molecules, where no direct information is available or from thorough calculations.

[Baulch, DL (2007)] elaborates the data needs for combustion modelling. It was emphasised that in modelling applications, whilst the principal focus is on the mechanisms and the rate of reactions, the thermodynamic and transport data of species are just as as important or in some cases, more important, and must not be neglected.

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3. Chemical Kinetics of Hydrogen-Air Combustion

Hydrogen is a simple molecule. However, it has a complex combustion mechanism with many intermediates, reactions and sub-mechanisms. Hydrogen combustion with air is considered in this section. For simplification, air will be assumed to be oxygen. This simple reaction is shown below:

2H2 + O2 → 2H2O [Reaction 3(a)]

The hydrogen-oxygen reaction mechanism has been extensively studied over the years and more recently, is receiving increasing application primarily because of the non-polluting characteristics of this combustion process. Though this may at first sight appear to be a very simple reaction, the mechanism is in fact very complex and still not fully understood.

In this context, it is essential to discuss and evaluate the hydrogen-oxygen reaction mechanism.

3.1. Overview of Mechanism

Tests conducted by [S. Gauthier, E et al (2007)] on natural gas/hydrogen mixture combustion in a porous radiant mixture suggest that the hydrogen molecule is very light and therefore strongly exposed to molecular diffusion processes. Furthermore, hydrogen is a very reactive species with a large flammability range.

A large number of detailed mechanisms of H2/O2 kinetics found in literature are dedicated to the combustion of hydrocarbons including sub-mechanisms for H2/O2 chemistry. However, the accuracy of the H2/O2

subset is also essential for the overall performance of a hydrocarbon mechanism. For this reason, the hydrogen mechanism essentially forms the base of all hydrocarbon mechanisms [Yetter R, Dryer F, Rabitz H. (1991)].

It has been observed that under ambient conditions of temperature, hydrogen and oxygen do not enter into any direct reaction between them in the absence of a catalyst [Ströhle, J. & Myhrvold, T. (2007)]. It is further seen that, if a mixture of hydrogen and oxygen get exposed to light, oxygen is activated usually by way of dissociation.

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According to [Goswami, M. et al. (2008)], in the presence of sensitizers (Cl, N2O, NH3), a set of secondary reactions take place and form H atoms. The H atoms enter into a reaction with activated oxygen thus forming H2O.

The oxidation of hydrogen can be divided into two sub-mechanisms:

1. The first sub-mechanism involves combinations of simple hydrogen and oxygen species, such as H. H2, O, H2O, O2, OH; and this includes some pressure sensitive reactions.

2. The second sub-mechanism involves the oxidation of H2O2 and evolves additional intermediate species HO2.

Accepted standardised values are used to predict the rate of control of these reactions. For initial understanding, a simplified version of the mechanism is shown:

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Figure 3.1(a): Simplified version of H2/O2 Mechanism. Rate constants k1, k2...

Each of the branching reactions (i.e. 2 & 3), produce two radical chain carriers for each chain carrier consumed. The two steps combine to give an overall branching coefficient of 3, corresponding to the hypothetical reaction:

H∙+O2+H2 → H∙+∙OH+∙OH [Reaction 3.1(b) - REVERSIBLE]

3.2. Bonding Structures

The first sub-mechanism of the hydrogen mechanism involves fundamentally all combinations of simple oxygen and hydrogen radicals. To understand them, the valence structures of both hydrogen and oxygen have to be considered. The valency is basically describes the combining tendency of an element.

3.2.1. Bonding of Hydrogen

Hydrogen has a valency of one. This implies it has the tendency to lose an electron while combining with an element, giving it a positive valency. The hydrogen radical is formed when the valence electron is not bonded (i.e. lone). The hydrogen bond is formed when this valence electron is linked with another atom.

H∙ (Hydrogen radical) H- (Single bonded hydrogen atom)

3.2.2. Bonding of Oxygen

Oxygen has six electrons in its outer shell. For it to have a complete outer shell (i.e. eight electrons in its outer shell), it can either loose six electrons or gain two electrons. It is obviously easier to gain only two electrons and for this reason, the valency of oxygen is two. The bonding of oxygen is

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much more diverse as it can result in radicals, single or even double bonds.

∙O∙ (Oxygen atom – single or triplet) : comprising of two electron pairs along with two unpaired electrons

∙O- (Single bonded oxygen): comprising of a single bond and a single electron in addition to two electron pairs

O= (Double bonded oxygen): comprising of two bonds as well as two electron pairs.

3.3. Mechanism Structure

The hydrogen mechanism is not just a list of reactions but rather a well thought-out set of sub-mechanisms and regimes that are significant under different conditions.

3.3.1. H2 O2 H OH Reactions

The analogy of the hydrogen mechanism being a chain branching process has been mentioned and validated by various researches. The different types of reactions can be represented based on the number of radicals created and destroyed between the reactants and products.

[Reaction 3.31 (a)]: Initiation Reactions: Fuel and Oxidiser

H2+M → H∙+H∙+M

The initiation reactions are those where stable species such as H2 and O2 form radicals. The dissociation energy of H2 is much lower than O2 and so the initiation step must be related to the former.

Shock tube experiments have also been carried out at low pressures and different compositions by [Wang et al (2003)]. In these studies, a different reaction is captured, which depicts the auto ignition delay more accurately and might be the most probable initiation step.

H2+O2 → ∙HO2+H∙

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Because of the fact that the hydroperoxy radical, ∙HO2 is a relatively stable radical, it does not readily participate in the rest of the necessary chain reactions.

[Reaction 3.31 (b)]: Chain Reactions: Increase of radicals

H∙+O2 → ∙O∙+∙OH∙O∙+H2 → H∙+∙OH H2+∙OH → H2O+H∙∙O∙+H2O → ∙OH+∙OH

The chain reaction steps are those where one radical species in the reactants produces one or two radicals in the products.

If these are prevalent, then the hydrogen mechanism moves faster towards ignition [Ó Conaire, M. et al (2004)].

[Reaction 3.31 (c)]: Chain Termination: Decrease of radicals

H∙+H∙+M → H2+M∙O∙+∙O∙+M → O2+MH∙+∙O∙+M → ∙OH+MH∙+∙OH → H2O

The chain termination steps are when two radicals in the reactant come together to form at least one stable species. When these reactants are dominant, then the reaction process is slowed or even stopped.

3.3.2. HO2 H2O2 Reactions

The first sub-mechanism of simple hydrogen and oxygen radicals has a transition, through the recombination to the mechanism of more complex species such as HO2, H2O2.

Since HO2 is a radical, formation of HO2 are chain branching reactions. On the other hand, the formation of H2O2 can be both terminating and branching.

Since radicals are highly reactive, their concentrations are minute and the backward reactions of all the branching and recombination chain steps are usually neglected. The rate of chain branching reactions increases with

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temperature. On the other hand, since the radical recombination steps require a third body, the rate of these reactions decreases with increasing temperature.

[Reaction 3.32 (a)]: Transition to HO2

H+O2+M → HO2+M

Active hydrogen atoms can be removed by oxygen molecules in the three-body recombination reaction as shown above. Studies conducted by [Markides, N. (2006)] at the University of Cambridge suggest that [Reaction 3.32 (a)] is kinetically favoured at lower temperatures and higher pressures.

[Reaction 3.32 (b)]: HO2 Reactions

HO2+HO2 → H2O2+OHHO2+H2 → H2O2+H

[Reaction 3.32 (c)]: H2O2 Reactions

H2O2+OH → H2O+HO2 H2O2+H → H2O+OHH2O2+H → HO2+H2

H2O2+M → OH+OH+M

[Reaction 3.32 (d)]: Surface Reactions

2HO2 → H2O2+O2

H2O2 → H2O+1/2 O2

H, O, OH → Destruction

3.4. Complex Behaviour

Since combustion is the prior concern, it is relevant to clarify the ignition behaviour of hydrogen. A well thought of approach is to connect the already stated reactions and intermediates with further catalytic reactions and their interactions with the walls, assuming containment within a

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system. All these reactions have a propensity to slow the evolution towards ignition.

A typical experiment was carried out by [Cohen, N. (1992)]. The study was based on flammability limits of H2/O2 mixtures. It was noticed that under different regimes, different reactions dominate. If the terminating reactions are in control, then ignition is inhibited. If the branching reactions dominate, then ignition is promoted.

The hydrogen combustion reaction results in a net increase of two hydrogen atoms in the reactive system, as conclusively established by various research institutes.

Unless these two hydrogen atoms are used up in the same length of time, the reaction rate increases exponentially with time and eventually leads to an explosion. Such circumstances can be evaded if the termination process is rapid.

Of particular interest in combustion systems, typically considering fuel integration, it is necessary to understand the explosion limits of the H2/O2 mechanism. This is displayed below:

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Figure 3.4(a): Explosion Limits of H2/O2 mechanism [Lu, T. & Law, C.K. (2009)]

The graph above [Figure 3.4(a)] shows the different limits of hydrogen ignition with respect to temperature and increasing pressure.

A general trend noticed is that increasing pressure of the system results in the mixture transitioning from explosive to non- explosive to explosive, as it pass through the first, second and third explosion limits.

3.4.1. First Limit

The first limit is at very low pressures and the mean free path in the gas is rather large.

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Experimental evidence clearly depicts the explosive limits dependence on the size of the vessel. These points the inhibitors are the destruction of radicals H, O, OH at the walls.

The reactions with the walls are more prominent at low pressure. This implies that the number of collisions to produce the chain reactions is relatively low. A reaction with the wall acts as a termination reaction thus lowering the explosive tendency.

If termination balances initiation, a steady reaction occurs and explosion is avoided

3.4.2. Second Limit

At higher pressures, the size of the vessel is less significant. The chain carriers react and combine prior to reaching the walls,

As pressure increases, three body termination reactions are more significant because of the increased collision frequency between molecules.

Within the hydrogen mechanism, as pressures go above the second explosion limit, HO2 is somewhat un-reactive, thus its formation can be thought of as a terminating reaction, as it is able to move towards the walls of the vessel.

The inhibition of the HO2 reaction dominating above the second explosion limit implies it requires a higher temperature to result in an explosion.

The peak pressure, which requires the highest temperature, is about half an atmosphere.

3.4.3. Third Limit

Taking into account simple density considerations, the initial density of the reactants increases as pressure goes up. For this reason, the temperature requirement for the reaction to reach a critical state to result in an explosion is low. This means the inhibiting effect of the second limit, namely the HO2 inhibition, is no longer relevant in this regime.

As pressure increases, even the inhibiting effect of HO2 formation is countered with the formation of H2O2, which in turn is a noteworthy source of hydroxyl radicals; in fact two for every HO2 radical:

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HO2+H2 → H2O2+H

H2O2+M → OH+OH+M

[Reaction 3.43 (a) – ALL REVERSIBLE]

In essence, from the above reaction, as pressure increases, HO2 reacts with H2. In addition, H2O, the major final product is an excellent third body.

3.5. Validation of Mechanisms

The number of H2/O2 mechanisms found in literature is too long to be applied to a more detailed context. Hence, only the most relevant mechanisms are selected, based on degree of validation and on frequency of being referred to.

For this section, the performance of the detailed H2/O2 mechanism is analysed. Various research groups and institutions have set out reaction kinetics outlining this mechanism.

3.5.1. Gas Research Institute

The Gas Research Institute developed the well known GRI-Mech 3.0, which contains 26 reversible reactions for H2/O2 kinetics. It is also optimised for natural gas combustion.

It was also validated against experimental data for ignition delay. Conversely, the Gas Research Institute discontinued support for the GRI-Mech project due to lack of funding, in the year 2000.

In view of this fact, significant work has been implemented regarding the H2/O2 chemical kinetic mechanism by various research institutes.

3.5.2. Yetter et al

[Yetter, R.A et al (1994)] presented a detailed CO/H2/O2 mechanism consisting of 19 reversible reactions for the H2/O2 system that was validated against experimental data at different temperatures and equivalence ratios.

29

The measurements, which were carried out assuming plug flow and adiabatic conditions, have been extensively validated for the hydrogen mechanism.The mechanism, however, was extensively studied at flow reactor conditions, but was not tested against or modified as a result of comparisons with experimental data derived in other types of experiments and in other parameter ranges.

3.5.3. Mueller et al

Based on an improvement on the chemical kinetic mechanism presented by [Yetter et al (1994)], an updated version was handed by [Mueller et al (1999)].

This revised mechanism is validated against a wide range of experimental conditions including those found in shock tubes, flow reactors and laminar premixed flame models. Excellent agreement of the model predictions with experimental observations demonstrates that the mechanism is comprehensive and has good predictive capabilities for different experimental systems.

3.5.4. Li et al

The study of [Li et al (2004)] provides an updated version of the H2/O2 mechanism, based on previous works of [Yetter, R.A & Mueller, M.A (1999)]. This improvement is based upon new thermodynamic data and rate coefficients, and has been compared over a wide range of experimental data including VPFR data, shock tube ignition delay data and new flame speeds.

This updated comprehensive detailed H2/O2 mechanism consists of 19 reversible elementary reactions, and is shown in [FIGURE 3.5 (a)]. All reactions are reversible.

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Figure 3.5 (a): Detailed H2/O2 mechanism [Li et al (2004)]

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4. Chemical Kinetics of Methane Oxidation

As clearly depicted in various literatures, combustion mechanisms have a clear hierarchical structure.

Figure 4(a): Hierarchical structure of hydrocarbon oxidation

As mentioned before, the hydrogen mechanism forms the base where many of the critical radicals driving ignition are formed, such as OH.

A critical sub-mechanism for hydrocarbons including the methane mechanism, as conclusively established by many research institutes, is the carbon monoxide (CO) mechanism. All hydrocarbon mechanisms, including methane produce intermediates that further react with these sub-mechanisms.

4.1. Carbon Bonding

In order to understand the multitude of hydrocarbon reactions, it is useful to know some of the basic Lewis bonding types of carbon.

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Hydrocarbon mechanisms are essentially many combinations of transformations between these states.

Carbon has four electrons available for bonding in its valence shell. The main stable carbon bonded states have either 4 single bonds; a double bond with 2 single bonds or a triple bond with a single bond.

These are illustrated below:

Single Bonded Carbon

Double Bonded Carbon

Triple Bonded Carbon

Simple radicals of carbon are usually of the same configurations; just one of the bonds is replaced, by a non-bonded electron.

Typical configurations are shown below:

Methylene: Radical Carbon (Singlet or Triplet)

Double Bonded with Radical

Triple Bonded with Radical

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4.2. Hydrocarbon Chemistry – Explosion Limits

As mentioned in the previous chapter, the hydrogen-oxygen mechanism is apparent in the ignition of hydrocarbons.

The carbon monoxide sub-mechanism plays an important part in the comprehension of the hydrocarbon auto-ignition chemistry. This is majorly due to the fact that the conversion of carbon monoxide to carbon dioxide is highly exothermic.

Methane, being the simplest hydrocarbon, is amongst the best understood in terms of combustion mechanism. Analysis of other hydrocarbon mechanisms usually begins with the methane mechanism.

Through extensive research, it is found in literature that hydrocarbons heavier than propane oxidise much slowly than hydrogen []. These molecules permit for the qualitative explanation of the unique explosion limits of these hydrocarbons.

This is depicted below:

Figure 4.2 (a): Explosion limits of Methane, Ethane and Propane. [Markides, N. (2006)]

34

The shape of the graph justifies how different the reactions involving hydrocarbons are from hydrogen.

As expected and noticed from [Figure 4.2 (a)], the larger the molecule, the greater the explosion shifts towards lower temperatures and pressures.

During the studies of [Miller JA & Kee RJ (1990)] on chemical kinetics and combustion modelling of hydrocarbons, it was noticed that intermolecular collisions are more prominent in heavier hydrocarbons and intermediates formed via these collisions break down forming radical pools, which initiate faster reactions.

Also, key experimental characteristics from the explosion graph include:

The very rapid reaction rates which proceed after induction levels The formation of formaldehyde, which tends to increase reaction

rates and shorten ignition times

Excluding methane and ethane, it was also noticed that cool flames are exhibited which are associated with multiple ignitions and negative temperature coefficients of reaction rates.

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4.3. Methane High Temperature Oxidation

The methane oxidation in detail can be complex with hundreds of reactions. An extensive literature survey did not yield a unique kinetic mechanism for the homogeneous combustion of methane.

The chemical reactions themselves may be discussed using various levels of detail.

The numbers quoted in this table lead to a question about the most realistic mechanism. Highly detailed and complex combustion mechanisms obtained by several authors using numerical simulations may differ widely.

Some of the detailed mechanisms based on numerical simulations are presented in the table below:

Table 4.2 (a): Different mechanisms for methane oxidation from various sources. Compiled from [Gosiewski, K (2009)]

All the mechanisms presented in literature, however varying in number of reactions, have the same basic features.

The illustration below shows the backbone of the methane oxidation under high temperature condition.

36

Figure 4.3 (a): Backbone of methane oxidation under high temperature.

H is abstracted initially, and with the addition of oxygen along with the loss of H, an aldehyde is formed. This aldehyde decomposes to CO.

Through the CO sub-mechanism, the final equilibrium products are made.

37

4.4. Side Path with CH2 Radicals

A noteworthy side reaction to the main reaction is free abstraction through singlet and triplet CH2 radicals. Following the addition of oxygen, through an OH or H2O, an aldehyde is formed within the backbone mechanism.

This side path reaction is shown below:

Figure 4.4 (a): Side path reaction with CH2 radicals.

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4.5. CH2OH Formation

From CH3, a methyl radical, the above set of side reactions in [Figure 4.4 (a)] go through a C2OH, which plays the role as the intermediate to the aldehyde. The CH2OH is formed from the addition of OH radicals with the simultaneous loss of hydrogen.

A schematic for this reaction is shown below:

Figure 4.5 (a): Aldehyde formation.

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4.6. CH4 Low Temperature Oxidation

[Simmie, J.M. (2003)] carried out the detailed chemical kinetic models for hydrocarbon fuels and found out that the low temperature oxidation of methane usually occurs around temperatures less than 1500 K. It comprises of two significant paths:

The recombination reaction; to form C2H6. This is why the propane sub-mechanism, although a larger product, is important t the methane oxidation.

The direct addition; of OH to form CH3OH, and its subsequent oxidation to form an aldehyde.

This is illustrated below:

Figure 4.6 (a): CH4 low temperature oxidation.

40

4.7. Aldehyde Decomposition

Aldehydes are an important intermediate play part in mechanisms including hydrocarbons, carbon monoxide and carbon dioxide.

Below is an illustration of the electron structure of the conversion of an aldehyde to a radical, through abstraction and loss of hydrogen.

Figure 4.7 (a): Aldehyde decomposition to radical.

The aldehyde starts as a double bond, with two single bonds, one of which is hydrogen. The hydrogen is abstracted or lost, and an aldehyde radical is produced with a double bonded carbon, with a single electron radical and a single bond, as can be seen in [Figure 4.6 (a)] above.

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4.8. CO Oxidation

CO2, the ultimate product, is produced via the two propagating steps of the CO mechanism.

The significance to these steps is that at the same time, the number of important radicals such as H and OH in the system is not diminished, hence it is propagation.

Important reactions found in the hydrogen mechanism involve further propagation and branching. Reactions with the original oxidiser, O2 produce radicals important to the two CO oxidation reactions, namely OH and HO2, one of which is a branching reaction evolving more radicals into the system.

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

Below shows the major steps in CO oxidation, made up of branching, propagating and termination steps.

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

[Reaction 4.8 (a) – ALL REVERSIBLE]: CO oxidation reactions

42

5. Laminar Burning Velocity

Hydrogen has significant prospects, both as a fuel in cells and combustions devices and mitigating the problem of global warming. However, with these advantages come stringent problems which further complicate its full implementation.

Ongoing research by various combustion science institutions have been carried out, which aim to bypass these difficulties. The opportunity of adding hydrogen to methane in order to improve performance, to extend operability ranges and to reduce pollutant emissions is one that seems rather promising.

Changes to fuel composition, particularly the addition of hydrogen to hydrocarbon fuels, affect both the chemical and physical processes occurring in flames. For hydrogen-methane mixtures to be used safely and effectively, physiochemical properties of this hybrid have to be determined. Of this, an important intrinsic property related is the laminar burning velocity.

Experimental values of the laminar burning velocity are mainly required to:

Help kineticists and modellers validate their schemes and models To provide input models for flashback, minimum ignition energy

and turbulent combustion

The laminar burning velocities of hydrogen-methane with air premixed flames have been experimentally and numerically measured at different values of equivalent ratios and fuel compositions by various research institutes.

Majority of these investigations on burning velocity conducted can be classified under two categories:

1. The effect of hydrogen addition to methane2. The effect of substitution of hydrogen by methane

Of particular interest are

The works of [Yu et al (1986)], whose results of the hydrogen-methane hybrid mixture showed a linear increase of the laminar burning velocity with the increase of hydrogen fraction in the fuel blends.

43

The effects of initial pressure and hydrogen fraction of hydrogen-methane flame was investigated by [Halter et al (2005)]. The results for this investigation showed that the laminar burning velocity increased with the increase of hydrogen fraction. A counter effect was noticed with the increase in initial pressure.

[Scholte and Vaags (1959)] first carried out measurement of the laminar burning velocity by means of a tube burner method. Together with more recent work by [Lui et al (2007)], this study is known to be extensive since a wide array of equivalence ratios and fuel compositions were investigated, varying form pure methane to pure hydrogen.

Also efforts regarding the laminar burning velocities of hydrogen-methane flames have been made on the computational side. Simulations of the premixed hybrid flames have been carried out extensively with CHEMKIN or COSILAB laminar premixed flames codes.

There are still many discrepancies on reactions and species which are dominant in hydrogen enriched methane flames. Following prior investigations in this project, the chemical kinetics of both methane and hydrogen have been reasonable well understood and outlined.

For this study, remaining uncertainties of the laminar burning velocity of hydrogen-methane flames will be considered. The chemical kinetic effects of hydrogen addition to methane will be analysed by investigating the reaction paths which would most likely occur in the premixed flames.

5.1. Case Study

The objectives of this study are to carry out simulations on the hybrid mixture of hydrogen-methane and air in order to investigate:

Individual molar concentrations of the species How the mass fractions of the species vary with distance (curved

length)

Ultimately, these findings will be related to the laminar burning velocities of other experiments found in literature and the dominant reactions and species will be identified from chemical kinetic mechanisms.

By means of two jet inlets, hydrogen and methane will be issued into a cylindrical geometry with a pressure outlet. This is illustrated below:

44

Figure 5.1 (a): Geometry for computation

45

In this simulation, swirl was also induced. The imposition of swirl was necessary to increase jet spread, jet width, jet growth, jet decay and entrainment, turbulence level and mixing with the surrounding fluid. The Swirl angle was set to 45 degrees.

In order to resolve an appropriate mesh size which will consequently lead to obtaining more accurate results, contours of y+ were plotted over the wall surfaces.

Y+ is a mesh-dependent dimensionless distance that quantifies to what degree the wall layer is resolved. Plotting contours of y+ determines values of turbulent viscosity and thermal conductivity at the centres of cells having a face on the wall.

The logarithmic wall treatment was used for these computations. This model is appropriate when these near wall centres are placed in the range of 20 <y+<200.

Y+ values for a mesh size of 0.03m are shown in [Figure 5.1 (b)].

After careful analysis of contour plots for this size mesh, it was realised that y+ values lie within a range of 38 and 24. Hence the logarithmic wall treatment is appropriate as the y+ value lies between the range of 20 <y+<200. For this reason, the chosen mesh size for computations was 0.03m, as this resolves solutions adequately for turbulent regimes.

Figure 5.1 (b): Geometry in x-y plane: 0.03m mesh size

NB: The flow rates of hydrogen and methane were kept equal and the study was based on cold flow stoichiometric simulations.

5.2. Results

Graph 5.2 (a): Mass Fraction of H2 vs. Curved Length

46

0 0.05 0.1 0.15 0.2 0.25 0.3 0.35 0.4 0.45 0.50

0.05

0.1

0.15

0.2

0.25

0.3

0.35

Mass Fraction of CH4 vs Curved Length

Curved Length (m)

Mas

s Fra

ction

of C

H4

Graph 5.2 (b): Mass Fraction of CH4 vs. Curved Length

5.3. Discussion

47

As can be seen from [Graph 5.2 (b)], the mass fraction of CH4 along the curved length decreases. This decrease is steady till about a distance of around 0.25m, when the mass fraction falls suddenly.

On the other hand, from [Graph 5.2 (a)], the mass fraction of H2 long the curved length increases.

This opposing trend can be ascribed to an increase in CH4 consumption by the reaction, and H2 switching from an intermediate species to a reactant.

A study conducted by [Salari, V. Di & Benedetto, A. Di (2007)] further validates this trend in relation to burning velocity. Their results for an experiment based on the laminar burning velocity of hydrogen-methane and air premixed flames are shown below:

Figure 5.3 (a): Laminar burning velocity of hydrogen–methane/air mixtures as function of the hydrogen content at three values of the equivalence ratio. [Salari, V. Di & Benedetto, A. Di (2007)]

As seen from the results above, as the mass fraction of hydrogen in the mixture increases, so does the burning velocity. The following elementary reactions contribute to the kinetic control of the hybrid mixture:

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From these results in [Figure 5.3 (a)], three regimes can be identified at all equivalence ratios, depending on hydrogen mole fraction in the hybrid mixture:

The first regime is a region of methane dominated combustion, where the hydrogen content is relatively low. The methane laminar burning velocity increases linearly and slightly. In this region, it was identified that methane conversion is dominated by:

H + O2 → O + OH [Reaction 5.3 (a) - REVERSIBLE]

The second regime is essentially a transition regime, where hydrogen content is at moderate levels. In this regime, the sensitivity factors of the dominating reaction in the first regime are reduced hence does not play the main role in controlling methane combustion.

The third regime is characterised by high hydrogen contents and methane-inhibited hydrogen combustion. This corresponds to hydrogen switching roles, as being an intermediate to being a reactant. For lean combustion the dominating reaction is:

OH + CH4 → CH3 + H2O [Reaction 5.3 (b) - REVERSIBE]

And for stoichiometric and rich flames, the dominating reaction is:

H + CH4 → CH3 + H2 [Reaction 5.3 (c) - REVERSIBLE]

Further investigations based on numerical studies of the effect of hydrogen addition on methane air flames have also been conducted.

Of relevance, is a recent publication in the International Journal of Hydrogen Energy by [Wang, J. Et al (2009)], based on a numerical study of the effect of hydrogen addition on methane–air mixtures combustion. The results of the reactant mole fraction of the calculated flames are shown below:

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Table 5.3 (a): Reactant mole fraction of calculated flames. [Wang, J. Et al (2009)]

It can be that the mole fraction of methane in the flame is reduced as the mole fraction of hydrogen increases. This effect will in turn affect the mole fraction of carbon related species in the flame. The reason for reduction of methane in the flame can be attributed to the enhancement of the chemical reaction as hydrogen is added.

From these findings, the dominant reactions contributing to the consumption of CH4 have been identified. These reactions are shown below:

OH + CH4 → CH3 + H2O (1)

H + CH4 → CH3 + H2 (2)

O + CH4 → OH + CH3 (3)

H + CH3 + M → CH4 + M (4)

[Reaction 5.3 (d) – ALL REVERSIBLE]

It is noticed from these reactions that the main consumption reactions of methane are the abstraction reactions attacked by H, O and OH, which are free radicals. From sensitivity analyses conducted by [Wang, J. Et al (2009)], methane mole fractions have the highest sensitivity to reactions:

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H + O2 → OH + O (1)

O + CH4 → CH3 + H2O (2)

H + CH4 → CH3 + H2 (3)

[Reaction 5.3 (e) – ALL REVERSIBLE]

Free radicals (H, O, and OH) are extremely active due because they poses an unpaired electron and are short lived during the combustion reaction. The chain branching and chain propagating reactions play an important role in this chemical reaction.

The normalised laminar burning velocity and maximum H concentration in the reaction zone as a function of hydrogen content, constructed by [Salari, V. Di & Benedetto, A. Di (2007)] is shown in the figure below:

Figure 5.3 (b): The normalised laminar burning velocity and maximum H concentration in the reaction zone as a function of hydrogen content. [Salari, V. Di & Benedetto, A. Di (2007)]

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Mole fractions and rate of production of free radicals increase as hydrogen is added. This also accounts to the enhancement of the combustion of methane and air.

From the [Figure 5.3 (b)] above, it is noticed that significant enhancement to the laminar burning velocity occurs only starting at hydrogen mole fractions higher than about 0.5. From this, it can be deduced that only at these values of fuel composition, the methane request for H radicals can be compensated by the hydrogen presence in the blend. It can also be deduced that the H radical concentration as well as the normalised laminar burning velocity are higher in rich mixtures.

The main reactions forming these radicals are:

OH + H2 → H + H2O (1)

H + O2 → O + OH (2)

[Reaction 5.3 (f) – ALL REVERSIBLE]

The first reaction listed in [Reaction 5.3 (F)] is promoted with the increase in hydrogen mole fraction and thus forms more hydrogen atoms which promote the reaction rate of the second reaction.

As formerly discussed in the previous chapter, carbon monoxide (CO) plays an important role in the oxidation of methane.

As the hydrogen fraction is increased, the CO mole fraction is reduced. This is due to a decrease in methane mole fraction as this reactant is used up quicker.

The main CO consumption reaction has been identified to be:

OH + CO → H + CO2

[Reaction 5.3 (g) – REVERSIBLE]

This reaction has also been identified as the most important reaction for hydrocarbon combustion and nearly all of the heat release associated with the overall combustion process will occur in this step.

C1 to C3 species associated with the combustion of methane also experience a similar trend to CO.

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6. Conclusions

Concerns about dwindling fossil fuel supplies have raised the need to investigate into alternate fuels. Considering the enormous ecological and economic importance of the transport sector, the question of ‘what fuel will be used in vehicles tomorrow’ touches upon core elements of sustainable development.

This research project was carried out on two fuels, natural gas (methane) and hydrogen, as both posses’ qualities to help mitigate the problem of global emissions. The study was mainly conducted to investigate the potential of hydrogen improving fuel properties. As convalescing vehicular combustion was of particular interest, the effect of hydrogen addition on improving the laminar burning velocity was delved into, basing the findings on chemical kinetics.

By comparing computational fluid dynamic simulations on the hybrid mixture of hydrogen and methane, to experimental evidence found in literature, it was discovered that hydrogen is an excellent additive to improve the combustion of natural gas. Simulation results showed that hydrogen in the hybrid mixture changes roles from being an intermediate species to being a reactant in the system. This is mainly because hydrogen has a low ignition energy and high reactivity and diffusivity.

Relating this feature to chemical kinetics, the promotion of the chemical reaction was discovered to be majorly attributed to the increase in radicals (H, OH, O) in the flame as hydrogen is added.

Substantiating simulation results to experiments on laminar burning velocity found in literature, it can be concluded that the addition of hydrogen to natural gas gives a good alternative fuel to conventional hydrocarbon fuels as it significantly enhances the chemical reaction (i.e. combustion) and also improves flame stability and burning velocity.

This project further validates the potential hydrogen has on sustainable development predominantly in the transport sector. Further research is necessary however in order to incorporate it fully in vehicles. A major focus can be on combustion chamber technology and investigating necessary modifications required to successfully incorporate hydrogenated fuels in spark ignition engines.

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