1 / 43

Gaseous Pollutants

Chap 2.3. Gaseous Pollutants. Carbon oxides Sulfur compounds Nitrogen compounds Hydrocarbon compounds Photochemical oxidants. Carbon Oxides. Two major carbon oxides Carbon dioxide (CO 2 ) Carbon monoxide (CO). CO 2. Natural atmospheric constituent Sources: Natural

Jims
Download Presentation

Gaseous Pollutants

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. Chap 2.3 Gaseous Pollutants • Carbon oxides • Sulfur compounds • Nitrogen compounds • Hydrocarbon compounds • Photochemical oxidants

  2. Carbon Oxides • Two major carbon oxides • Carbon dioxide (CO2) • Carbon monoxide (CO) CO2 • Natural atmospheric constituent • Sources: • Natural • Aerobic biological processes, combustion and weathering of carbonates in rock and soil • Anthropogenic: • Combustion of fossil fuels • Land use conversion

  3. What’s the impact if there is no CO2 in the atmosphere? Is CO2 emission regulated? Should it be? Figure 2.2 CO2 • Essential atmospheric gas • Present in variable concentrations • Not considered to be toxic • Environmental concerns are relatively new • Changes in atmospheric concentrations • Geological time • The modern period 1.5-1.7 ppmv/yr • Long atmospheric lifetime (~100 years)

  4. Figure 2.3 CO2 • Major sink processes • Oceans • Forests • Pre-industrial revolution: 98% of exchangeable CO2 were in the oceans and 2% in the atmosphere; for anthropogenic CO2, only 42% dissolves in oceans More discussion in Atmospheric Effects

  5. CO • Colorless, odorless, tasteless gas • Produced as a result of incomplete combustion

  6. Adverse effects on the consumption of OH·? • Formation of O3 Overall CO • Sink processes • Photochemistry with OH· (hydroxyl radical) • Soil uptake • Atmospheric lifetime (1 month in the tropics and 4 months in mid-latitudes) • Increase CH4 concentration thus enhancing global warming M: an energy absorbing molecule, e.g. N2 or O2 OH·: hydroperoxyl radical O(3P): ground-state atomic oxygen h: a photon of light energy

  7. Why higher in higher latitudes and altitudes? CO • Background level concentration • Vary with latitude, lower in the tropics and higher in the northern middle latitudes • Average 110 ppbv • Increasing 1%/yr, mostly in the northern middle latitudes • Urban/suburban levels • Vary from few ppmv to 60 ppmv: mainly associated with transportation emissions • Average highs (10-20 ppmv) • Higher concentrations in higher altitude cities

  8. Sulfur Compounds • Sulfur Oxides: Sulfur trioxide (SO3), Sulfur dioxide (SO2) • Reduced sulfur compounds (COS, CS2, H2S) Sulfur Oxides • Anthropogenic sources • Combustion of S-containing fuels • Smelting of metal ores • Natural sources • Volcanoes • Oxidation of reduced S compounds SO3 SO2 • Produced from SO2 oxidation • Rapidly reacts with water • Very short atmospheric lifetime • Colorless, sulfurous odor gas • Major sulfur oxide in the atmosphere • Produced on S oxidation • May be converted to SO3

  9. What is the overall picture? Data from http://www.uea.ac.uk/~e490/su/sulfur.htm

  10. Sink processes: SO2 oxidized in gas & liquid phase reactions; can be direct, photochemical or catalytic • Gas phase • Reaction with OH· (major), O3, HO2·, RO2·, O(3P) • Liquid phase • It can be further oxidized to H2SO4 by reaction with HNO2, O3, H2O2, RO2· and catalysis by Fe and Mn H2SO4: sulfuric acid H2SO3: sulfurous acid HNO2: nitrous acid H2O2: hydrogen peroxide

  11. What is the consequence of the deposition? Removal processes • Aerosol formation by nucleation/condensation • Sulfuric acid reacts with ammonia: forms sulfate salts • SO2 + aerosols removed by wet & dry deposition processes • SO2 atmospheric lifetime (1-7 days) SO2 concentration • Background levels: ~20 pptv over marine surface to 16- pptv over clean areas of US • Historical urban 1-hour highs: 1-500 ppbv • Highest 1 hr near non-ferrous metal smelters: 1.5-2.3 ppmv More discussion in Welfare Effects

  12. Reduced S compounds • (CH3)2S (Dimethyl sulfide) • Released from oceans in large quantities • Short atmospheric lifetime (0.6 days) by rapid conversion to SO2 • COS (Carbonyl sulfide) • Most abundant S species in atmosphere • Produced biogenically • Background levels (0.5 ppbv) • Limited reactivity • Atmospheric lifetime ( 44 years) • Mercaptans • Source of malodors: “Rotting cabbage” • CS2 (Carbon disulfide) • Produced biogenically • Photochemically reactive • Global concentrations range (15-190 ppbv) • Atmospheric lifetime (12 days)

  13. H2S • Major environmental and health concern (toxic): characteristic malodor (rotten egg odor, threshold of 500 pptv) • Sources: • Natural: primarily by biological decomposition • Anthropogenic sources: Oil & gas extraction, Petroleum refining, Coke ovens, Kraft paper mills • Short atmospheric lifetime (4.4 days): Oxidized to SO2 • Background concentrations( 30-100 pptv); concentrations in industrial and surrounding ambient environments can be above the odor threshold

  14. Nitrogen Compounds • Gas/Liquid phase • Nitrous acid (HNO2) • Nitric acid (HNO3) • Nitrite (NO2-) • Nitrate (NO3-) • Ammonium (NH4+) • NOx: NO and NO2 • NOy: NOx and their atmospheric oxidation products • Gas phase • Nitrogen (N2) • Nitrous oxide (N2O) • Nitric oxide (NO) • Nitrogen dioxide (NO2) • Nitrate radical (NO3) • Dinitrogen pentoxide (N2O5) • Peroxyacyl nitrate (CH3COO2NO2; PAN) • Ammonia (NH3) • Hydrogen cyanide (HCN)

  15. So, why do we care about its increase in the atmosphere? Nitrous Oxide (N2O) • Colorless, slightly sweet non-toxic gas • Also called “laughing gas” because human exposure to elevated concentrations produces a kind of hysteria • Atmospheric concentration increasing: (0.8 ppbv/yr) • Sources: • Natural: by nitrification and denitrification processes biogenically • Anthropogenic sources: Soil disturbance, Agricultural fertilizers • No known sink in the troposphere: atmospheric lifetime of 150 years • Stratosphere is only sink: photolysis and subsequent oxidation by singlet oxygen (O(1D))

  16. So, why do we care about NO emission? Nitric oxide (NO) • Colorless, odorless, relatively non-toxic gas • Natural sources: • Anaerobic biological processes • Biomass burning processes, lightning • Oxidation of NH3 • Photochemical reactions in stratosphere and transport from there into the troposphere • Anthropogenic sources • Fuel combustion (transportation, coal-fired power plants, boilers, incinerators, home space heating) • Product of high temperature combustion; concentration depends on temperature and cooling rate More details about NO formation in Reaction/Kinetics

  17. Nitrogen Dioxide (NO2) • Brown colored, relatively toxic gas with a pungent and irritating odor • Absorbs light and promotes atmospheric photochemistry • Peak levels occur in mid morning • Production by chemical reactions • Direct oxidation • Photochemical reactions

  18. Weekly pattern? Seasonal pattern? NOx concentrations • Remote locations: 20-80 pptv • Rural locations: 20 pptv -10ppbv • Urban/suburban areas: 10 ppbv - 1 ppmv • Diurnal variation

  19. (Reverse reaction under sunlight) (removed by dry & wet deposition) NOx Sink Processes • Chemical reactions convert NO to NO2 to HNO3 • Major sink process reaction with OH· • Nighttime reactions involving O3 • Reactions with organic compounds • Neutralized by ammonia to form salts • HNO3 serves as a reservoir and carrier for NOx

  20. Other N Compounds Example? • HCN (Hydrogen cyanide) • Organic nitrate compounds: Peroxyacyl nitrate (PAN), Peroxyproprionyl nitrate (PPN), Peroxybutyl nitrate (PBN) – potent eye irritants Reduced N Compounds • NH3 (Ammonia) • Sources: anaerobic decomposition of organic matter, animals and their wastes, biomass burning, soil humus formation, fertilizer application, coal combustion, industrial emissions • Background levels (0.1-10 ppbv) • Sink processes: reaction with acids, absorption by water and soil surface • Atmospheric lifetime (10 days) • Very important neutralizer for strong acids

  21. Hydrocarbons • Comprise a large number of chemical substances • Basic structure includes only carbon & hydrogen covalently bonded • Serves as a base for a number of derivative compounds • May be straight, chained, branched or cyclic • May be • Saturated (single bonds, C-C) • Unsaturated (double/triple bonds, C = C) • Unsaturated HCs more reactive • May be gas, liquid or solid phase, depending on the number of carbons: gases 1-4 C; volatile liquids 5-12 C; semivolatile liquids or solids > 12 C

  22. http://en.wikipedia.org/wiki/Toluene http://en.wikipedia.org/wiki/Xylene Example? http://en.wikipedia.org/wiki/Benzene Hydrocarbons • Types • Aliphatic • Paraffins/Alkanes - single bond • Olefins/Alkenes - have 1 double bond • Alkynes – have 1 triple bond • Aromatic • Have at least one benzene ring • Benzene • Toluene • Xylene • Lifetime • Paraffins – days • Olefins – hours • Alkyenes – weeks • Benzene (12 days), toluene (2 days), m-xylene (7 hr)

  23. Hydrocarbons • Polycyclic aromatic HCs (PAHs) • Multiple benzene rings • Solids under ambient conditions • Produced in combustion processes • Components of atmospheric aerosol • Potent carcinogens • Classification by volatility • VVOC (Very Volatile Organic Compounds): BP up to 50-100 oC • VOC (Volatile Organic Compounds): BP 50-100 to 240-260 oC • SVOC (Semi-Volatile Organic Compounds): BP 240-260 to 380-400 oC • SOC (Solid Organic Compounds): above 400 oC • NMHCs: Non-Methane HydroCarbons; Methane is excluded because of its low reactivity in the atmosphere

  24. Hydrocarbon Derivatives • Formed from reactions with O2, N2, S or halogens • Derivatives of major atmospheric concern include: • Oxyhydrocarbons • Halogenated hydrocarbons Oxyhydrocarbons • Direct emissions from industrial/commercial use: adhesives, solvents • By-products of combustion • Produced from photochemical reactions • Include • Aldehydes (C=O) • Acids (-COOH) • Alcohols (-OH) • Ketones (CO) • Ethers (C-O-C) • Esters (R-CO-OR’)

  25. Why? Nonmethane Hydrocarbons • Primary focus of air quality regulation • Biogenic sources • Trees (isoterpenes, monoterpenes) • Grasslands (light paraffins; higher HCs) • Soils (ethane) • Ocean water (light paraffins, olefins, C9-C28 paraffins) • Order of magnitude higher than anthropogenic • Question of their significance • Anthropogenic emission estimates • 40% transportation • 32% solvent use • 38% industrial manufacturing/fuel combustion • Identification is challenging; concentration of individual NMHC is not commonly measured

  26. NMHC Sink Processes • Oxidation by OH· or O3 • Produce alkylperoxyradicals (ROO·) • ROO· is converted to alkoxy radical (RO·) by reacting with NO • RO· reacts with O2 to produce aldehyde • Longer chained NMHCs result in ketones • Ethane reaction

  27. Oxidation of HCHO • Acetaldehyde more reactive than ethane • Acetaldehyde oxidized to HCHO through a series of reactions with OH· • HCHO can decompose by ultraviolet (UV) light in the range of 330-350 nm and produce CO 2nd pathway 1st pathway produces OH· for oxidizing other NMHC

  28. Photochemical Precursors • CO (above) can be eventually converted to CO2 • Aldehydes/ketones removed by wet/dry deposition • Longer chained HCs may produce condensible products • These oxidation products (e.g. ROO·, RO·, HO2· and CO) serve as major reactants in forming smog; they also serve to produce elevated tropospheric O3

  29. So, why do we care about CH4? Figure 2.5 Methane (CH4) • Most abundant HC in atmosphere • Low reactivity with OH • Little significance in urban/suburban photochemistry; hence, levels subtracted from total HC concentration • Can affect downwind of urban sources • Thermal absorber - global warming concern • Concentrations average ~ 1.75 ppmv • Significant increases over time since industrial revolution

  30. Methane • Natural Sources • Anaerobic decomposition in swamps, lakes and sewage wastes • Rice paddies • Ruminant/termite digestion • Anthropogenic Sources • Coal/lignite mining • Oil/gas extraction • Petroleum refining • Transmission line leakage • Automobile exhaust

  31. Why? Methane • Sink processes • In the troposphere, reaction with OH· • Produces HCHO, CO & ultimately CO2 • Competes with CO for OH· • Photodecomposition in stratosphere • Produces H2O • Major source of water in stratosphere • Levels in atmosphere increase with increasing CO • Atmospheric lifetime (~10 years)

  32. Halogenated Hydrocarbons • Contain one or more atoms of halogen (Cl, Br, or F); include a variety of compounds • Chlorinated HCs • Brominated HCs • Chlorofluoro HCs • Remarkable persistence (i.e. low reactivity) • Include both natural/anthropogenic sources; both volatile and semi-volatile compounds

  33. Volatile Halogenated HCs • Methyl Chloride (CH3Cl) • Methyl Bromide (CH3Br) • Methyl Chloroform (CH3CCl3) • Trichloroethylene(CH2CCl3) • Perchloroethylene(C2Cl4) • Carbon tetrachloride (CCl4) Semi-volatile Halogenated HCs • Chlorinated pesticides (DDT, Dieldrin, Aldrin) • Polychlorinated biphenyls (PCBs) • Polybrominated biphenyls (PBBs)

  34. So, why do we care about them? Chlorofluoro HCs (CFCs) • Trichlorofluoromethane (CFCl3): CFC-11 • Dichlorodifluoromethane (CF2Cl2): CFC-12 • Trichlorotrifluoroethane (C2Cl3F3): CFC-13 • Characterized by • Low reactivity • Low mammalian toxicity • Strong thermal absorption properties • Good solvent properties

  35. Halogenated HCs • Most halogenated HCs have tropospheric sinks • CFCs have no tropospheric sinks. • Atmospheric Lifetimes CH3Cl, CH3Br ~ 1 year CH3CCl3 ~ 6.3 years CCl4 ~ 40 years CFCl3 ~ 75 years CF2Cl2 ~ 111-170 years • Concentrations vary spatially, with highest in source regions over the northern hemisphere. • Concentrations in both the troposphere and stratosphere have been increasing until the early 1990s.

  36. Photochemical Oxidants • Produced in chemical reactions involving: • Sunlight • Nitrogen oxides • Oxygen • Hydrocarbons • Include • Ozone • Nitrogen dioxide • Peroxyacyl nitrate • Odd hydrogen compounds (OH·, HO2·, H2O2)

  37. Is O3 level high or low at a highway tollbooth? This doesn’t explain the high level O3 in smog! What’s wrong? Figure 2.6 Photochemical oxidants: O3 • Ozone the major photochemical oxidant of concern • Atmospheric O3 formation • Requires source of O(3P): through photolysis of NO2 at wavelengths of 280-430 nm • Nitric oxide quickly destroys O3 • Steady-state concentration of 20 ppb under solar noon conditions in mid-latitudes

  38. Tropospheric O3 Formation • Elevated O3 levels occur as a result of reactions that convert NO to NO2 without consuming O3! • Role of peroxy compounds (ROO·) derived from photochemical oxidation of HCs

  39. In summary, what are the important parameters in determining O3 level? Tropospheric O3 formation • Rate of O3 formation depends on ROO· availability • ROO· produced when OH· and HOx react with HCs • OH· is formed by photo-dissociation of O3, aldehydes and HNO2

  40. Tropospheric O3 Concentrations • Remote Locations (20-50 ppbv, summer months) • Photochemical processes • Stratospheric intrusion • Populated locations • Peak concentrations (50 ppbv - 600 ppbv) • In urban areas concentrations decline at night • In rural areas peak concentrations occur at night • Elevated rural levels associated with long-range transport (Yosemite NP,http://www2.nature.nps.gov/air/webcams/parks/yosecam/yosecam.cfm) • Transport of O3 aloft • Transport of low reactivity paraffins

  41. Tropospheric O3 levels

  42. Ozone Sink Mechanisms • Photo-dissociation • Reaction with NO in polluted area • Reaction with NO2 at night time • Surface destruction: reaction with plants, bare land, ice/snow and man-made structures

  43. Quick Reflection

More Related