1 / 28

AME 513 Principles of Combustion

AME 513 Principles of Combustion. Lecture 6 Chemical kinetics III – real fuels. Outline. Experimental methods Types of reaction CO-O 2 Hydrocarbons NO x Flame inhibition. Experimental methods. Constant volume vessel

vance
Télécharger la présentation

AME 513 Principles of Combustion

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. AME 513Principles of Combustion Lecture 6 Chemical kinetics III – real fuels

  2. Outline • Experimental methods • Types of reaction • CO-O2 • Hydrocarbons • NOx • Flame inhibition AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  3. Experimental methods • Constant volume vessel • Heat evacuated vessel to some temperature, inject reactants quickly (shorter than reaction time scale), see if it reacts explosively or just slowly • Ideally constant T & P • Not well-defined limit, depends on injection / mixing time ~ d2/D (d = vessel dimension, D = diffusivity) being much faster that reaction, only useful for conditions with “slow” reaction • For gases, both mass D and thermal D ~ 1/P, so characteristic reaction time required for explosion varies with P! • Wall effects crucial (radical sink) • Only yields global properties (e.g. overall reaction rates) • Well-stirred reactor • Well defined limits • Must be certain that mixing of fresh reactants with products already in reactor is much shorter than residence time in reactor • Typically need very large mass flows to study flame-like conditions since reaction time scale ~ 10-3 s in flame • Only yields global properties AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  4. Experimental methods • Plug-flow reactors • Flow reactants down a preheated tube (usually constant T), measure species at varying distances along tube (~ time) • Can obtain data on observing evolution of individual species, not just global properties, thus infer elementary reaction rates • Laminar or turbulent flow • Only useful for “slow” reactions (low temperature, not flame-like conditions) (tens of milliseconds) • Need to address issue of axial dispersion of reactants • Shock tube • Pass shock wave through mixture, watch evolution of species • Step-like change in T and P, well defined, nearly homogeneous • Can obtain data on individual reactions • Amenable to quantitative laser diagnostics • Only useful for “fast” reactions (few ms) AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  5. Experimental methods • Laminar flame • “Simple” setup – measure burning velocity SL • Obviously applicable to “real” flame chemistry • Flames are very thin (< 1mm), hard to probe inside to measure species evolution • Interaction with transport via convection and diffusion – compare results with computations using detailed flame models • Not a sensitive instrument – overall reaction rate ~ SL1/2 AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  6. Types of reactions • Global reaction • Example: CH4 + 2O2 CO2+ 2 H2O • NOT an actual reaction that occurs • No relation between order of reaction (3 in this example) and actual pressure effect on reaction rate • Chain initiation • Example: H2 + M  H + H + M • Break stable molecule into radical(s) • High Ea – endothermic, must break strong bond • High Z (i.e. not very orientation sensitive) • Not needed in flames, where radical source (products) exists • Chain branching • Example: H + O2 OH + O • Use radical to create more radical • High Ea - endothermic, must break strong bond, but also make a bond, so not as high as chain initiation • Moderate Z (somewhat orientation sensitive) AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  7. Types of reactions • Chain propagation or “shuffle” reaction • Example: OH + H2 H2O + H • Use radical + stable molecule to create another pair • Moderate Ea – may be nearly thermo-neutral • Moderate Z (somewhat orientation sensitive) • Chain termination • Example: H + OH + M  H2O+ M • Recombine radicals into stable molecules (usually products) • Low or zero Ea – exothermic, no activation barrier • Need 3rd body to absorb enthalpy and conserve momentum • Moderate Z (may be orientation sensitive) • Schematic multi-step mechanism (e.g. Hautman et al, 1981) C3H8 1.5 C2H4 + H2 (Initial breakdown of fuel) C2H4 + O22 CO + 2 H2 (C2H4 = surrogate for radicals) 2H2 + O2 2 H2O (oxidation of H2) 2CO + O2 2 CO2(oxidation of CO) Reaction rates crazy (see Turns, p. 158) – sometimes order of reaction is negative – causes problems as concentration  0 AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  8. H2– O2 • Simulate explosion limits using online chemistry calculator • Use time to 50% H2 consumption as measure of explosion limit • Diffusion time ~ P so scale accordingly, e.g. at 10 atm, allow 10x more time than at 1 atm • Results for t at 1 atm = 25 s close to experiments (Fig. 5.1) • Second limit independent of time (vessel size & walls material) • Second limit where branching vs. recombination rates ≈ same H + O2 OH + O = H + O2+ M  HO2+ M when P = 103.4T0.2e-17500/RT AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  9. CO – O2 • H2-O2 already discussed – 3 explosion limits • CO MUCH different because no chain branching and no “shuffle” reaction to create product CO + O2 CO2 + O Very slow, but needed as source of O O2 + M  O+ O Very high activation energy, even slower CO + M  C + O + M No way! C=O is strongest chemical bond! CO + O + M  CO2 + M Creates product but removes radicals • As a result, pure CO – O2 oxidation is extremely slow! • Early experiments showed widely varying results because of contamination with water; with any hydrogen source CO + OH  CO2 + H Shuffle reaction to create CO2 H+ O2 OH + O Regenerate OH plus another O Neither are fast, but better than alternatives! • Stoich. CO + O2, 1 atm, 1500K, time to consume 50% of CO: • No H2O: 1.03 s; 1 ppm H2O: 0.27 s; 10 ppm: 0.038 s; • 100 ppm 0.0043 s; 1,000 ppm 0.00080 s; 10,000 ppm 0.00023 s AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  10. CO – O2 • Case shown: const. T = 1000K, P = 1 atm, CO:H2:O2 = 1:1:10 • To a good approximation, H2-O2 acts as infinitely fast (i.e. steady-state) radical source for CO • CO does not affect H2-O2 process, CO reactionjust too slow • 50% H2 consumption in ≈ 400 µs, ≈ same with or without CO AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  11. CO – O2 • Explosion limit experiments show 2nd-limit behavior, even with “dry” CO – not really dry • CSU homogeneous kinetics model (next page) does not! Dickens et al., 1964 Gordon and Knipe, 1955 AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  12. CO – O2 • Unlike classical experiments, CSU homogenous kinetics calculator predicts that dry CO has only a single limit – but CSU site does not include O3 which may be important • With H2Oaddition, behavior is very similar to H2-O2 with offset to higher T - CO is a “parasite” on the OH source AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  13. Hydrocarbon – O2 • Hydrocarbons inhibit their own oxidation, because they react with radicals more readily than O2 reacts with radicals • Once an H atom is removed from fuel, e.g. C3H8 + O2 C3H7• + HO2• then the fuel radical can lose another H to form an alkene, e.g. C3H7• + O2 C3H6+ HO2• thus alkenes are a key intermediate in alkane oxidation • Nearly all of fuel must be consumed before radical pool needed to consume CO can build up • CO oxidation is the last step • b-scission – fuel molecule breaks apart 1 C-C bond away from C missing an H atom (avoids having to move an H atom to an adjacent C atom) • This stuff only matters at “low” temperatures (<1500K) where H + O2 branching is inhibited by the hydrocarbons; at higher temperatures, H + O2 branching is sufficiently rapid AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  14. Hydrocarbon – O2 AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  15. Hydrocarbon – O2 • Hydrocarbons have no 1st or 2nd explosion limit; instead of H + O2 + M  HO2 + M being the dominant recombination reaction, it’s CH4+ H  CH3 + H2, which has same pressure dependence as the branching reaction H + O2 OH + O • C-H bonds stronger in CH4 than C3H8, thus higher explosion T AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  16. Hydrocarbon oxidation • Start with fuel molecule RH, where R is an “organic radical”, e.g. propane without an H • Abstract an H atom from RH RH + O2 R + HOO • Add an O2 to R R + O2 ROO • Produce peroxides with O-O single bond (half as strong as O=O double bond (120 kcal/mole vs. 60 kcal/mole), much easier to break) ROO + RH  R + ROOH or HOO + RH  R + HOOH • Break O-O single bond, create “chain branching” process ROOH + M RO + OH or HOOH + M HO + OH • Newly created radicals generate more organic radicals RH + OH R + HOH or RH + RO R + ROH • Note that rate of reaction will be sensitive to rates of H atom removal from fuel molecule RH AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  17. Hydrocarbon oxidation • Rate of H atom removal depends on strength of C-H bond, which in turn depends on how many other carbons are bonded to that C - stronger bond, slower reaction, less knock • Examples: n-heptane: 6 primary, 12 secondary C-H bonds 2, 2, 4 trimethy pentane: 15 primary, 2 secondary, 1 tertiary AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  18. Hydrocarbon oxidation • Does this small difference in bond strength matter? YES because activation energy is high • If we use bond strength as a measure of activation energy (dangerous in general, but ok here…) then at a typical 900K methane : primary : secondary : tertiary exp(-Emethane/T) : exp(-Eprimary/T) : exp(-Esecondary/T) : exp(-Etertiary/T) ≈ exp(-105,000 cal/mole/(1.987 cal/mole-K)(900K)) : exp(-98000/1.987*900) : exp(-95000/1.987*900) : exp(-93000/1.987*900) ≈ 1 : 50 : 268 : 820 • As a result, fuels with mostly primary C-H bonds will decompose much more slowly than isomers with more secondary & tertiary C-H bonds – higher octane number in gasoline-type fuels AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  19. Hydrocarbon oxidation • How to make CO? RO• formed by cleavage of O-O bond in peroxide then enter aldehyde route RO• + M  R’HCO + H + M, e.g. C2H5O• + M  CH3HCO + H + M • Aldehydes have weakest C-H bond (≈87 kcal/mole) thus R’HCO + M  R’CO• + H• + M, e.g. CH3HCO + M  CH3CO + H• + M R’HCO + O2 R’CO• + HO2•, e.g. CH3HCO + O2 CH3CO• + HO2• • Aldehydes also have weak C-C bond thus R’CO•+ M  R’ + CO + M, e.g. CH3CO• + M  CH3• + CO + M • Somewhat roundabout but easiest way to make CO, still takes a “long” time, see flow reactor result – first aldehydes CH2O and C2H4O form, then CO rises as aldehydes decompose AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  20. Hydrocarbon oxidation • Larger hydrocarbons also have negative temperature coefficient (NTC) behavior at low T (below H + O2 branching) • NTC especially prevalent in rich mixtures – reaction rate decreases with increasing T • R + O2 ROO is very reversible due to weak R-O bond • Equilibrium favors dissociation (ROO R + O2) at higher T, so ROO won’t stick around long enough to make ROOH • At higher temperatures HOO + RH  R + HOOH forms peroxides, lessening the need for ROO • (HOOH reaction has higher Ea than ROO + RH  R + ROOHbecause the former is more exothermic; C-O bond strength 86 vs 111 kcal/mole for H-O, thus HOOH reaction more dominant at higher temperature) • Also forms “cool flames” – exothermic propagating waves that don’t consume all reactants because of NTC shut-down – no longer “homogeneous” reaction • Negative temperature coefficient behavior not seen in calculations because no ROOH chemistry in CSU model AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  21. Hydrocarbon – O2 Hewitt and Thornes, 1937 C3H8-O2, f = 5, 0.6 liter vessel (Note T and P axes are flipped) AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  22. Chemical fire suppressants • Key to suppression is removal of H atoms H + HBr H2 + Br H + Br2 HBr + Br Br + Br + M  Br2 + M -------------------------------- H + H  H2 • Why Br and not Cl or F? HCl and HF too stable, 1st reaction too slow • HBr is a corrosive liquid, not convenient - use CF3Br (Halon 1301) - Br easily removed, remaining CF3 very stable, high CP to soak up thermal enthalpy • Problem - CF3Br very powerful ozone depleter- banned! • Alternatives not very good; best ozone-friendly chemical alternative is probably CF3CH2CF3 or CF3H • Other alternatives (e.g. water mist) AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  23. Chemical fire suppressants AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  24. Zeldovichmechanism for NO formation • Extremely high activation energy due to enormous strength of NN bond (≈ 220 kcal/mole) O + N2 NO + N N + O2 NO + O N + OH  NO + H • Reaction (1) is limiting; Z1exp(-E1/T) < Z2exp(-E2/T) except for T > 3394K • 1 NO molecule formed from (1) yields 2 NO molecules • Assuming steady-state for N, partial equilibrium for O, OH, H, with O2, Heywood (1988) shows: • T = 2200K, P = 1 atm: NO = 0.59 second • By comparison, time scale for chemical reactions in flame front ≈ 0.001 second for stoichiometric hydrocarbon-air • Thus, Zeldovich NO occurs in the burned gases downstream of the flame front, not in the flame front itself AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  25. Summary • Real fuels have very complex chemistry, not just 1 or 2 steps • Need chain branching for fast reaction • Hydrogen, CO, hydrocarbons dominated by • Initiation – H2 + M  H + H + M, RH+ M  R + H + M • Branching • Typically H+ O2 OH + O at high T / low P • Peroxide path (with HOOH or ROOH) at low T / high P where H atoms are lost due to recombination • Recombination - H + O2+ M  HO2+ M • Radical termination at walls in explosion vessel • CO oxidation • Requires CO + OH  CO2+ H • Parasitic on H2– O2 mechanism since CO + OH relatively slow • Hydrocarbons • Inhibit their own oxidation due to RH + H  R + H2 • Decomposition rate depends on C-H bond strength • First fuel decomposes, generates radical pool, generates CO then oxidizes it AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  26. Midterm exam • October 19, 9:00 – 10:30 am • Covering lectures 1 - 6 • Open books / notes / calculators • Laptop computers may be used ONLY to view .pdf versions of lecture notes – NOT .pptx versions • Note .pdf compilation of all lectures: http://ronney.usc.edu/AME513F12/AME513-F12-AllLectures.pdf • GASEQ, Excel spreadsheets, CSU website, etc. NOT ALLOWED • Followed by lecture 10:45 am – 11:50 am AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  27. Midterm exam – topics covered • Chemical thermodynamics • Stoichiometry • Heating value • Flame temperatures • Equilibrium • Degrees of reaction freedom • Equilibrium constraints • Compression/expansion • Chemical kinetics • Law of mass action, collision theory • Arrhenius form of reaction rate expression • Coupling with thermodynamics • Adiabatic constant-volume reaction • Constant-pressure Well-Stirred Reactor • Multistep reactions • Single-step irreversible & reversible • Steady-state & partial equilibrium approximations • Kinetics of real fuels • H2 – O2 • CO – O2 • Hydrocarbons – O2 AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

  28. Midterm exam – types of problems • Chemical thermodynamics • Property values will be given if needed • Stoichiometry • Heating value • Flame temperature • Equilibrium • Chemical kinetics • Exact solution (for very simple chemistry!) • Steady state or partial equilibrium • How would an explosion limit plot change if • Wall conditions changed • New species affecting certain reactions added • “Most likely” reaction steps (similar to Turns 5.7) • General - how would QR, Tad, reaction rates, homogenous explosion time, WSR blowout, etc. be affected by • Ronney Fuels, Inc. – new fuel or additive • Planet X – different atmosphere (pressure, temperature, etc.) AME 513 - Fall 2012 - Lecture 6 - Chemical Kinetics III

More Related