Atmospheric Chemistry I

Climate Physics and Chemistry Atmospheric Chemistry I: Ozone, Chlorofluorocarbons, Nitrogen Oxides, and Dimethyl Sulfide A brief history of the atmospheric chemistry of O3, nitrogen oxides, and CFC's • • • • • • • • • • • 1970: Paul Crutzen shows that destruction by NOx can account for discrepancy between factor of three excess of calculated ozone layer relative to observed. 1971: Crutzen shows that dominant source of NOx in stratosphere is transport of N2O released from biosphere. Harold Johnston suggests that nitrogen oxides released by supersonic transports might affect ozone layer (but US SST rejected on grounds of economics rather than environmental concerns). 1973: James Lovelock, inventor of "electron capture" detector, shows that CFC concentrations in the atmosphere were increasing. 1974: Molina and Rowland suggest that CFCs would deplete ozone layer [at about same time, Ralph Cicerone and Johnston also suggest that Cl might deplete ozone, but focus on volcanism and the space shuttle as a source of Cl rather than CFC's]. 1976: US bans use of aerosol spray CFCs after 1978. 1980: "Ozone hole" opens up in Antarctic, unbeknownst to the human race. 1985: Josephan Farman found that mean Antarctic ozone (column) had decreased by nearly half of what it had been in the early 1970's. Ozone hole intensity continues to increase, with Antarctic ozone virtually disappearing at 15-21 km altitude every spring. This event was not predicted by Molina-Rowland model. Late 1980's: Antarctic ozone hole linked to heterogeneous catalysis by polar stratospheric clouds. 1987: Signing of multinational "Montreal Protocol" regulating CFC production. Gas phase chemical kinetics I • Elementary bimolecular reaction A + B -> AB • probability of reaction depends on frequency of collision, orientation (random probability), and ability to overcome activation energy threshold d[ AB] = k[ A][ B] dt Gas phase chemical kinetics II Energy Activation energy Reactants Products d ln k Ea = dT RT 2 If Ea is constant: k = Ae−Ea / RT (Arrhenius equation) Example: temperature dependence of rate constant for formation of hydrogen iodide H2 + I2 -> 2 HI d HI / dt = [H2][I2] A e-Ea/RT 3.5 log k 2.5 1.5 0.5 -0.5 -1.5 1.2 1.3 1.4 1.5 103/T 1.6 1.7 1.8 But things can get complicated: the example of hydrogen bromide H2 + Br2 -> 2 HBr Reaction kinetics, as determined by Bedenstein (~1900) : d[ HBr] k[ H 2 ] [ Br2 ] = [ HBr] dt M+ [ Br2 ] The reaction actually proceeds by 5 separate reactions: Br2 -> 2 Br Br + H2 -> HBr + H H + Br2 -> HBr +Br H + HBr -> H2 + Br 2 Br + M -> Br2 + M Statospheric Ozone - key equations I A. Ozone production • 1. O2 + hν -> O + O • 2. O + O2 + M -> O3 + M (M is some other molecule that acts as a catalyst) Statospheric Ozone - key equations II B. Ozone destruction • • • • • • • • • • • • 1. O3 + hν -> O2 + O 2. O + O3 -> 2 O2 3a. O3 + OH -> HO2 + O2 3b. HO2 + O3 -> OH + 2O2 (A1+A2+B1+B2 are sometimes referred to as the Chapman Cycle) (3a and 3b involving OH are part of a more general class of reactions involving “OX” where “OX” = HOx, NOx, ClOx) 4a. N2O + hν -> N2 + O (1D) (most N2O destruction occurs this way; note that this reaction converts N2O into forms which do not deplete ozone; O (1D) is an energetically excited oxygen atom) 4b. N2O + O (1D) -> 2 NO (only 5% of N2O is destroyed this way - but it leads to products that deplete ozone) (note: NO, NO2 and NO3 are collectively referred to as “NOx”: 4c. NO + O3 -> NO2 + O2 these plus all oxidized nitrogen species of NOx on relatively short time scales (e.g. N2O5, HNO3) are referred to as “NOy” 4d. O3 -> O + O2 4e. NO2 + O -> NO + O2 5a. Cl + O3 -> ClO + O2 5b. O3 + hν -> O + O2 5c. ClO + O -> Cl + O2 Statospheric Ozone - key equations III B. Ozone destruction (cont’d) In the presence of ice clouds: • • 6a. ClONO2 + HCl -> Cl2 + HNO3 6b. Cl2 + hν -> 2 Cl Source of stratospheric HCl: • 7. Cl + CH4 -> HCl + CH3 (these are the reactions responsible for the ozone hole) Ozone chemistry (simplified) N2O (nitrous oxide): Greenhouse gas (and ozone sink) (no laughing matter) • N2O is produced from dissolved nitrate by microbes in the earth surface environment: 2 "CH2O" + 2 NO3- + 2 H+ -> N2O + 2 CO2 + 3 H2O • Some microbes in the same environment will further reduce N2O to N2. However, gaseous N2O can diffuse out of the environments where this sink reaction can occur, and once it enters the troposphere, it is relatively unreactive. • The primary sink for tropospheric N2O is the stratospheric photochemistry. It takes about a century for the troposphere to cycle through the stratosphere, so the residence time of N2O in the atmosphere is about a century. N2O + hν -> N2 + O (1D) (most N2O destruction occurs this way; note that this reaction converts N2O into forms which do not deplete ozone; O (1D) is an energetically excited oxygen atom) (only 5% of N2O is destroyed this N2O + O (1D) -> 2 NO way but it leads to products which deplete ozone) N2O is increasing with time CFC12 N2 O CFC11 N2O is increasing with time… Mixing Ratio (ppbv) 100 N 2O 10 FEB. 1. 1979, 41oN (Golden el al, 1980) JAN. 30. 1979, 40oN (Calculated, CR) 1 SAMS. JAN., FEB., 1979, 40oN (Jones and Pyle, 1984) 40 35 30 25 20 15 Altitude (km) 10 5 N2O is a greenhouse gas and contributes to global warming N2O Concentration (ppbv) 310 300 290 280 270 1700 1750 1800 1850 1900 1950 2000 Year AD Atmospheric N2O concentrations during the last 250 years obtained by analyzing the H15 ice core collected in Antarctica. Measured values are shown by filled squares. Annual mean values of the N2O concentration at the South Pole for the period 1977-1991 AD are also plotted by circles for reference. Solid curve represents the result obtained by fitting the Reinsch-type cubic spline with a cutoff period of 50 years to the data. (millions of tons of nitrogen) Fertilizer Production or Use Fertilizer production 90 60 30 0 1920 1940 1960 1980 Year 2000 2020 Nitrate in a Greenland ice core Concentrations of Nitrate Measured in the GISP2 Ice Core Purple Line represents the 2000 year average 200 Parts/Billion (ppb) 150 100 50 1972 2000 1950 1900 1850 Calender Years 1800 1750 ESTIMATED SOURCES AND SINKS OF N2O TYPICAL OF THE LAST DECADE (Tg(N)/yr) Tropospheric N2O budget Range of Estimates Most Likely Value 3.1 - 4.7 3.9# 9 - 16 ? 12.3 Total Sinks 9 - 16 12.3 Implied Total Sources (atmospheric increase + total sinks) 13 - 20 16.2 Range of estimates Most Likely Value Increase in Atmosphere Sinks Stratosphere (N2O + h Soils N 2 + O) Known Sources Natural Oceans 1-5 3 Tropical Soils Wet Forests Dry Savannas 2.2 - 3.7 0.5 - 2.0 3 1 Temperate Soils Forests Grasslands 0.1 - 2.0 0.5 - 2.0 1 1 6 - 12 9 1.8 - 5.3 0.2 - 1.0 0.7 - 1.8 0.2 - 0.5 3.5 0.5 1.3 0.4 Total Known Anthropogenic 3.7 - 7.7 5.7 KNOWN SOURCE TOTALS 10 - 17 14.7 Total Known Natural Sources Anthropogenic Cultivated Soils Biomass Burning Industrial Sources Cattle and Feed Lots # The observed atmospheric increase implies that sources exceed sinks by 3.9 Tg(N)/yr. Dimethyl sulfide and climate • A. Marine organisms (e.g. coccolithophorids, single-celled plants) produce DMSP (dimethylsulfoniopropionate): CH3 \ S+-CH2-CH2-COO/ CH3 • • • • DMS and climate DMSP This compound is converted to DMS [(CH3)2S] when the cell is destroyed and the cell contents are degraded. B. DMS is volatile and moves into the atmosphere by gas exchange. C. DMS is oxidized in the atmosphere to two byproducts with a "branch ratio": sulfuric acid (H2SO4) and methanesulfonic acid (MSA: CH3SO3H). Cycle is complex with many intermediates; branch ratio appears to depend mainly on temperature (low MSA:nssSO4= at warmer temperatures) D. The products are transported to ice cores and recorded there as non-sea-salt sulfate (nss SO4=) and MSA. E. The non-sea-salt sulfate, (initially sulfuric acid but later reacted with sea salt particles and mineral aerosols) forms particles that act as cloud condensation nuclei (CCN). In areas that are far from land, these CCN can be an important factor regulating cloud formation and rainfall. Some people have hypothesized a feedback between DMS and climate.