Experimental study of the combustion properties of methane/hydrogen mixtures Gersen, Sander

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1 University of Groningen Experimental study of the combustion properties of methane/hydrogen mixtures Gersen, Sander IMPORTANT NOTE: You are advised to consult the publisher's version (publisher's PDF) if you wish to cite from it. Please check the document version below. Document Version Publisher's PDF, also known as Version of record Publication date: 2007 Link to publication in University of Groningen/UMCG research database Citation for published version (APA): Gersen, S. (2007). Experimental study of the combustion properties of methane/hydrogen mixtures s.n. Copyright Other than for strictly personal use, it is not permitted to download or to forward/distribute the text or part of it without the consent of the author(s) and/or copyright holder(s), unless the work is under an open content license (like Creative Commons). Take-down policy If you believe that this document breaches copyright please contact us providing details, and we will remove access to the work immediately and investigate your claim. Downloaded from the University of Groningen/UMCG research database (Pure): For technical reasons the number of authors shown on this cover page is limited to 10 maximum. Download date:

2 CHAPTER 1 Combustion properties of homogeneous reacting gas mixtures 7

3 1.1. Motivation to study the combustion properties of CH 4 /H 2 gas mixtures Consider a homogeneous (premixed) fuel/oxidizer gas mixture. A premixed mixture is characterized by the equivalence ratio, ϕ, which expresses the ratio of fuel and oxidizer the unburned mixture. This is given by, [ Fuel] 1 ϕ =., (1.1) [ Oxidizer] f st where the amounts [fuel] and [oxidizer] can be expressed in molar, volume or mass units, and f st is the ratio of fuel to oxidizer under stoichiometric conditions using the same units. Here we will generally use moles or mole fraction as units. A mixture is said to be stoichiometric (ϕ=1) when fuel and oxidizer are present in the ratios prescribed by the balanced chemical reaction for combustion: C H n n + ( x + ) O 2 xco 2 + H O, (R1.1) 4 2 x n 2 If the oxidizer in the unburned mixture is in excess, the mixture is said to be fuel-lean (ϕ<1), while the mixture is called fuel-rich (ϕ>1) when an excess of fuel in the unburned mixture is present. The gas mixture can remain unreacted, such as in fuel-air mixtures at room temperature in the absence of an ignition source, or the fuel and oxidizer react (combust) to form products. Combustion can take place either in a flame (a reaction front propagates subsonically through the mixture) or in a non-flame mode ( homogeneous combustion, reaction occurs simultaneously everywhere in the mixture). To understand which mode of combustion takes place under a given set of conditions (temperature, pressure, equivalence ratio), it is necessary to study the chemical processes in the system in detail. The overall combustion process can be described by reaction (R1.1). However, it is unrealistic to think that combustion proceeds via this single reaction because it would require breaking high-energy bonds, which at room temperature makes this reaction extremely slow. Instead, combustion occurs in a sequential process involving many reactive intermediate species. To illustrate this process, we first consider a H 2 -O 2 gas mixture. The conversion of hydrogen and oxygen to water starts with the formation of reactive species (radicals) 8

4 to initiate a chain of reactions [1]. The reactions in which radicals are formed from stable species are called chain-initiation reactions, and an example of a chaininitiation reaction is the (endothermic) dissociation reaction: H + M = 2H + M 436kJ / mole, (R1.2) 2 The H radicals formed in reaction (R1.2) can react further with oxygen molecules, forming two new radicals, OH and O, H + O2 = OH + O 70.6kJ / mole. (R1.3) This reaction (R1.3), in which two radicals are created for each radical consumed is called a chain-branching reaction and is crucially important in combustion processes. The formation of the radicals OH and O can lead to further chain branching via O + H 2 = OH + H 8.2kJ / mole. (R1.4) In addition, there are reactions in which the number of radicals does not change, such as OH + H = H O + H 63.21kJ / mole, (R1.5) which are called chain-propagating reactions. The reactions in which radicals react to stable species without forming another radical are called chain-terminating reactions, as in H + OH + M = H 2 O + M kJ / mole. (R1.6) Although not strictly chain terminating, since HO 2 is a radical, the reaction H + O + M = HO + M kJ / mole (R1.7)

5 is often considered chain terminating because, compared with the flame radicals, H, O and OH, the HO 2 radical is relatively unreactive. At low pressures, reaction (R1.7) is an important chain-terminating reaction because the mildly reactive HO 2 radicals diffuse to the wall, where they react at the surface. Summing the chain-branching and chain-propagating reactions (R1.3+R1.4+R1.5+R1.5) results in H + H + O 3H + 2H O 47.62kJ / mole, (R1.8) from which we can see that starting with one radical, three radicals are formed from the reactants in this simplified mechanism. For quantitative description of chemical processes the rate of change of the species concentrations (formation and consumption) should be determined. For species A in an arbitrary bimolecular elementary reaction, k f aa + bb cc + dd, (R1.9) kr this is expressed as: da a b c d = k f [ A] [ B] + kr[ C] [ D], (1.2) dt where A,B,C,D denote the different species in the reaction, a,b,c,d are the stoichiometric coefficients of species A,B,C,D, respectively, and k f and k r represent the forward and reverse rate coefficient of the reaction. For example, the rate of change of the oxygen radical in reaction (R1.3) is expressed as, d[ O] R1.3 R1. 3 = k [ H ][ O2 ] k r [ OH ][ O] dt f +. (1.3) The rate coefficients k f and k r are connected through the equilibrium constant K w. The reaction rate constant k of a reaction is generally assumed to have a modified- Arrhenius temperature dependence, b E k AT exp A = RT, (1.4) 10

6 where A is the pre-exponential factor, b the temperature exponent, T the temperature, R the universal gas constant and E a the activation energy, which corresponds to the energy barrier that has to be overcome during reaction. Here should be mentioned that the activation energy is always higher than the heat of the formation. Generally, the more exothermic a reaction is, the smaller the activation energy. As can be seen from reactions (R1.3-R1.5), the rate of formation of the important free radicals ( H + OH + O) is proportional to the concentration [n] of the radicals with some coefficient α. The rate of consumption of radicals (chain termination) is proportional to the concentration [n] as well (R1.7), with some rate β; the rate of chain initiation, denoted as γ, is independent of the concentration [n] (R1.2). Thus, in generalized form, the rate of change of the concentration of free radicals can be expressed as, d[ n] = γ + ( α β )[ n], (1.5) dt From equation (1.5) three different scenarios can be derived, which are presented schematically in figure 1.1 a [2]. Figure 1.1a) Schematic of the growth of the concentration of free radicals [n] in time. b) Schematic of the growth of free radicals [n] in time for the cases with and without heat release. 11

7 For the condition α>β the concentration [n] increases exponentially in time, and ignition takes place. When α<β, d[n]/dt becomes zero, and no exponential growth of free radicals occurs (no ignition). The condition α=β results in a linear growth of the concentration of free radicals, and defines the ignition limit. Equation (1.5) shows that the growth of the concentration of free radicals is determined by the competition between the chain branching (R1.3) and chain terminating reactions (R1.7). A very important parameter in this competition is the temperature, since the chain branching reaction (R1.3) has large activation energy E a [3] while that of the chain terminating reaction (R1.7) is small [4]. Thus, the value of α (chain branching) is strongly dependent upon the temperature, and β (chain terminating) is more or less independent of the temperature. At low temperatures the endothermic chain branching reaction will not proceed rapidly, so the value of α is much smaller than β (α<β), and ignition does not occur. Increasing the temperature results in an increase in α, while β remains unchanged; at sufficiently high temperatures α will be larger than β and ignition occurs. The temperature during the early period of the ignition process remains more or less constant because the heat release from the branching and propagating reactions (R1.8) is small, but as the radical concentration grows, exothermic reactions such as (R1.6) and (R1.7) will produce substantial quantities of heat. If the heat produced by the exothermic reactions in system exceeds the rate of heat loss to the surrounding, the temperature in the system will rise. Since the rate of reaction, and thus the rate of heat release, grows exponentially with temperature, the overall reaction will auto-accelerate, that is, the system will explode. The time before explosion takes place is called autoignition delay time (figure 1.1b). In this example, if no heat accumulates in the system (T=constant), no auto-acceleration of the reaction rate takes place; in this case we speak of non-explosive reactions, both situations are shown in figure 1.1b. As described above, the dominance of chain branching reaction (R1.3) characterizes the high temperature regime, while at low and intermediate temperatures the in essential chain terminating reaction (R1.7) competes effectively with reaction (R1.3) [5]. Since the rate of a three-body reaction increases with pressure much faster than the rate of a two-body reaction, there exists a pressure above which reaction (R1.7) exceeds the rate of the competing reaction (R1.3). Reaction (R1.7) is only a chain terminating reaction when the produced HO 2 radicals will diffuse to the wall 12

8 and recombine to stable molecules without having undergone a reaction. At high pressures, species collide much more frequently. Therefore, at sufficiently high pressures the HO 2 radicals will be frequently interrupted in its path to the walls by reacting with H 2 molecules to produce H and H 2 O 2 (R1.10) [5], HO + H = H O + H 72.45kJ / mole. (R1.10) The H radicals so produced can contribute to chain branching via (R1.3) or will generate more HO 2 via reaction (R1.7) and H 2 O 2 itself can contribute to branching via [5], H O + M = 2OH + M 214.6kJ / mole. (R1.11) 2 2 Thus at high pressure and moderate temperatures reaction (R1.7) will dominate over (R1.3), and the number of active centers will grow (α>β) via the sequence of reactions (R1.7), (R1.10), and (R1.11). Instead of heating the cold gas mixture by the heat release of exothermic reactions as in this example of a closed homogeneous system, in flames, the mixture is heated by conduction from the hot flame gases. Furthermore, radicals needed to decompose the fuel are also transported from the high temperature region of the flame. As in the closed system described above, the rate of formation of radicals controls the overall rate of reaction in flames. For flames, however, the chain initiation and chain-terminating reactions are less important in the formation/destruction of radicals (α>>β), and the reactions (R1.3)-(R1.5), responsible for the growth of the radical pool, dominate the overall reaction rate in H 2 -O 2 flames [6]. The combustion chemistry of hydrocarbon fuels is much more complicated than that of hydrogen. As an example, we consider a CH 4 -O 2 mixture. Under flame conditions (α>>β), the most important chain branching reaction in the oxidation of methane is also reaction (R1.3) [6]. After radicals are transported into the unburned gas mixture, methane is attacked by the radicals, as in CH + H CH + H 13kJ / mole. (R1.12)

9 As can be seen from figure 1.2, the rate coefficient of reaction (R1.12) [7] is much larger than that for reaction (R1.3) [3]. Thus reaction (R1.12) competes effectively with reaction (R1.3) for H atoms and reduces the chain branching rate. Figure 1.2. Reaction rate expressions for reaction R1.3 [3] and reaction R1.12 [7]. Furthermore, the very reactive H radical is replaced in reaction (R1.12) by the unreactive CH 3 radical. These two processes result in a slow conversion of methane and contributes to the low burning velocity of methane (40 cm/s) as compared to hydrogen (340 cm/s) [8]. If the CH 4 -O 2 mixture under consideration is at moderate or low temperature (T below roughly 1100K), chain branching reaction (R1.3) is too slow to provide a sufficient branching rate for autoignition, and a different reaction path dominates [9]. These paths are extremely complex and strongly dependent upon temperature and pressure [10]. At sufficiently high pressures, reactions involving the radical HO 2 become important in the low temperature regime, for example [9], CH + HO CH + H O 85.5kJ / mole. (R1.13) The oxidation of the CH 3 formed, and the development of the radical pool, is complicated and slow. This process dominates most of the ignition delay period and is 14

10 characterized by the accumulation of significant amounts intermediate species such as H 2, C 2 H 6, CH 2 O and H 2 O 2. The decomposition of H 2 O 2 via reaction (R1.11) and the oxidation of CH 2 O via a sequence of reactions lead to a sharp increase in the concentration of free radicals [10] and ultimately ignition occurs (α>β) [10]. The H (R1.12) and HO 2 (R1.13) scavenging reactions compete effectively with R1.3 and R1.10 respectively, and thus effectively reduce the chain branching rate in the CH 4 -O 2 system. This, together with the formation of the relatively unreactive CH 3 radical, contributes to the fact that CH 4 -O 2 mixtures tend to auto ignite slower than H 2 -O 2 mixtures [6], illustrated in figure 1.3. Figure 1.3. Computed autoignition delay times for stoichiometric H 2 /air and CH 4 /air mixtures at P=30 bar. Calculations were made using the GRI-Mech 3.0 mechanism [11]. Besides their effects on the combustion properties, such as burning velocity and ignition, the differences in the combustion chemistry of methane and hydrogen have significant consequences for pollutant formation. One of the main consequences is that during the combustion of methane (and other hydrocarbons) carbon containing pollutant species like soot, HCN and CO are formed, while the only pollutant from hydrogen combustion is NO x. Moreover, the formation of NO in methane combustion is different than that from H 2 combustion; in methane flames an additional mechanism exist that produces NO via the hydrocarbon intermediate CH [12]: 15

11 CH + N 2 products NO. (R1.14) ( NCN?, HCN?) This mechanism is particularly important under fuel rich conditions. A challenging task is to understand the possible changes in the combustion chemistry caused by addition of hydrogen to methane and how this affects combustion properties like pollutant formation and ignition delay. Since there is a clear distinction between the chemistry in flames and ignition, it is necessary to study both. To gain understanding in the underlying chemical kinetics and physical processes involved in these two kinds of combustion process, it is necessary to analyze the combustion processes quantitatively by solving the governing equations with detailed chemical mechanisms. 1.2 Governing equations for a homogeneous reacting gas mixture in a closed system The time-dependent behavior of a closed system containing a reacting gas mixture is described by the system of the conservation equations for mass and energy. Since no mass can be formed or destroyed by chemical reaction, the total mass of the closed systems remains constant over time: d( ρv) d K = ρvy k = 0, (1.6) dt dt k = 1 where ρ is the overall mass density, V is the system volume and Y k is the mass fraction of k-th component in the gas mixture. The mass fractions of individual species change in time according to dy ρ k = ωk Wk, k = 1.K (1.7) dt 16

12 where ω k and W k are the molar chemical production rate per unit of volume and molecular weight of the k-th species, respectively. The density is related to temperature T, pressure p and composition through the ideal gas equation of state: p ρ = W, (1.8) RT K Y k where W = 1/ is the average molecular weight of the mixture. The system k = 1 M k (1.6) (1.8) consists of (K+1) linearly independent equations and contains (K+3) unknown parameters: ρ, p, T, V and Y k. Since the system contains more unknowns than equations, it can be solved only when two unknown parameters (for example, the measured temperature and pressure) are used as input. The number of input parameters can be decreased to one if the energy conservation equation is added to the system. The energy conservation equation can be derived from the first law of thermodynamics, which states that heat δq added to the system is equal to the sum of the change of its internal energy du and the work PdV of the system against an external force: du + PdV = δq. (1.9) Equation (1.9) can be rewritten as du dt dv dq + P = = Q dt dt loss. (1.10) The internal energy of the mixture is given by U K = ρvykuk, (1.11) k= 1 where u k is a specific energy of the k-th component. After differentiating expression (1.11) and substituting in (1.10) one receives 17

13 1 d dt ρ 1 K. Cv + p + ukωkwk = qloss, (1.12) dt dt ρ k= 1 where q loss = Q /(ρv) is the heat loss per unit of mass and C v is the specific heat of loss system at constant volume. For an adiabatic mixture of inert gases (ω k = 0 and q loss = 0), equation (1.12) can be easily integrated, resulting in the following expression T CvW R T0 dt T ρ = ln. (1.13) ρ 0 It is common to use the ratio of molar heat capacities at constant pressure and constant volume, γ, and a specific volume v, instead of C v and ρ. In this case, equation (1.13) can be rewritten as T 1 ( γ 1) T 0 dt T V = ln 0. (1.14) V Several simulation programs have been developed to solve the set of governing equations. The program used in this study is SENKIN [13], and runs in conjunction with pre-processors from the CHEMKIN library [14], which incorporate the chemical mechanism and thermodynamic properties. 1.3 Laminar premixed flames Flat laminar premixed flames are ideally suited for combustion research, since the one-dimensional character offers great advantages for modeling and unambiguous model-experiment comparison. Moreover, the structure of these flames is representative for many practical flames. The structure of laminar premixed flames 18

14 can be divided in three zones (figure 1.4), the preheat zone, the flame front (reaction zone) and the post-flame zone (burned-gas zone). In the preheat zone the unburned gas mixture is heated by conduction and diffusion of species from the flame front; this zone can be considered as chemically inert. The flame front, located downstream of the preheat zone is a thin zone (in the order of 1 mm at atmospheric pressure) in which the fuel is rapidly oxidized by radicals from the post-flame zone as described above, leading to a steep gradients in temperature and species concentrations. The flame front is rich in radicals and intermediate species. Although the temperature and major species in the post-flame zone are close to their equilibrium value, the concentrations of minor species can differ substantially from their equilibrium value. In the post-flame zone, the system goes to equilibrium predominantly via radicalrecombination reactions such as (R1.6). Figure 1.4. Schematic illustration of the structure of a premix one-dimensional flame. Premixed flat flames can be characterized by the free-flame laminar burning velocity, v L. In the laboratory system, where the cold gas moves with velocity v u, the flame front propagates with velocity v u -v L. We can consider three situations regarding the stability of a idealized one-dimensional flame. If the cold gas velocity is larger than the laminar burning velocity, v u >v L, the flame front propagates upstream. When 19

15 the burning velocity, v L is equal to the velocity of the unburned gasses, (v=v L ) the flame front is stationary in space, and if v u <v L the flame front will be convected downstream. The laminar burning velocity and the temperature of the burned gas are completely determined by the properties of the unburned mixture, such as the equivalence ratio, temperature and the identity of the fuel [2]. Figure 1.5 shows the interaction of the idealized 1-D flame with a porous-plug burner. Figure 1.5a illustrates a flat flame stabilized on a burner where the unburned gas velocity is set equal to the free-flame laminar burning velocity (v u =v L ). In this situation, all heat generated during combustion is transferred completely into the gas mixture and the flame is essentially adiabatic (neglecting flame radiation). Lowering the unburned gas velocity v u causes propagation of the flame front towards the burner surface. Since the porous plug is too dense to allow propagation of the flame into the burner, the flame is stopped in its upstream propagation. In this case, the flame transfers heat to the burner by conduction, lowering the flame temperature and thus lowering the actual burning velocity of the flame v L'. The flame transfers enough heat to the burner to reach a stationary situation (v u = v L' ), illustrated in figure 1.5b. This type of flame is called a burner-stabilized flame. Figure 1.5. a) Adiabatic flat flame (freely propagating flame) b) Burner-stabilized flat flame Further decrease in the unburned gas velocity results in increasing heat loss and drop in temperature; ultimately the temperature drops to such a level that α<β and the flame extinguishes. 20

16 Governing equations for a one-dimensional laminar flame The description of one-dimensional laminar flames is based on the conservation equation for mass, species mass fraction and energy. Using the assumptions that onedimensional laminar flames are: (1) stationary (all flame parameters are independent of time), (2) the system is at constant pressure and (3) effects due to viscosity, radiation and external forces are negligible [2,15], the conservation equations governing the behavior of these flames can be summarized as follows: overall conservation of mass d( ρv) dx dm = dx = 0, (1.15) where v is the mass averaged flow velocity, x is the distance along the line normal to the burner surface and M is called the mass flux. conservation of species d( ρyk( Vk + v)) = ω k W k, k=1.k (1.16) dx where Y k is the mass fraction and V k is the diffusion velocity, which accounts for the effect of molecular transport due to concentration gradients of the k th species [2,16]. Since mass can neither be destroyed nor formed in chemical reactions it follows from (1.15) and (1.16) that, K d( ρy ( )) K k Vk + v d( ρv) = ω k W k = = 0 dx dx k= 1 k= 1 k=1 K. (1.17) Addition of the ideal gas equation of state (1.8) to the system of equations (1.15, 1.16) results in a system containing (K+1) linear independent equations. Assuming that the diffusion velocity V k is a known function of temperature and species concentrations, 21

17 the system contains (K+2) unknown parameters (T, ρ, v and Y k ). Therefore an additional equation should be introduced to solve the system of equations: conservation of energy d dt ρyk( v+ Vk) Hk λ = 0, (1.18) dx dx k where H k the specific enthalpy of species k and λ the thermal conductivity coefficient. With the proper choice of the boundary conditions for one-dimensional flames, it is possible to solve the governing equations [2,16]. Various software packages have been developed, which are able to calculate the one-dimensional flame structure in only a few minutes by solving the set of governing equations. The simulation program used in this study is the PREMIX code [17]. This code is included in the CHEMKIN II simulation package [13]. This package operates using a reaction mechanism data file as input, along with thermal and transport properties of the species involved in the mechanism. The program is able to calculate temperature- and mole fraction profiles in both burner-stabilized and free flames. 1.4 Chemical mechanisms In the last decades chemical kinetic mechanisms have been developed to model combustion of hydrogen (for example, see [18]) and hydrocarbon mixtures ([11,19], among many others). These mechanisms, used to describe the transformation of reactants into products, may contain hundreds of species and thousands of elementary reactions. Improvement of the mechanisms currently in use is necessary, since none of them can be regarded as comprehensive [20], i.e. accounting for all combustion phenomena and the predictive power is only accurate for a small range of parameters. In order to improve the existing chemical mechanisms, they should be validated against experimental data, where parameters are varied in a well-defined manner. Sensitivity and rate-of-production analyses are used to design and optimize models. Using these methods rate-limiting steps and characteristic reaction paths can be identified [2]. 22

18 The experimental data obtained in this study have been modeled using different mechanisms. One of these mechanisms is GRI-Mech 3.0 [11], which is widely used and has arguably become the industry standard for methane in the research community. This mechanism is optimized to model natural gas combustion and contains 325 reactions and 53 species, including reactions that describe NO formation and reburn chemistry. 23

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20 M. Caracotsios, W. E. Stewart, P. Glarborg, C. Wang, and O. Adigun, CHEMKIN Collection, Release 3.6, Reaction Design, Inc., San Diego, CA, (2000). 17. R. J. Kee, J. F. Grcar, M. D. Smooke, J. A. Miller, Fortran program for modelling steady one-dimensional premixed flames. Sandia Report SAND Sandia National Laboratories, (1985). 18. O. M. Conaire, H. J. Curran, J. M. Simmy, W. J. Pitz, C. K. Westbrook, Int. J. Chem. Kin. 36 (2004) J. M. Simmie, Prog. Energy Combust. Sci. 29 (2003)