Introduction to Photochemical Smog Chemistry

1. Introduction to Photochemical Smog Chemistry • Basic Reactions that form O3 • Distinguish between O3 formation in the troposphere and stratosphere • How hydrocarbons and aldehydes participate in the formation of smog ozone • Formation of free radicals • Nitrogen loss mechanisms • Running simple simulation models

2. Why are we interested in the Smog Chemistry???

3. Ozone • ozone is a form of oxygen; it has three atoms of oxygen per molecule • It is formed in the lower troposphere (the atmosphere we live up to 6 km) from the photolysis of NO2 • NO2 + light --> NO + O. • O. + O2 -----> O3 (ozone) • its concentration near the earth’s surface ranges from 0.01 to 0.5 ppm

4. Ozone • background ranges from 0.02 to 0.06 ppm • What is a ppm?? • A ppm in the gas phase is one molecule per 106 molecules air or • 1x10-6 m3 O3 per 1 m3 air or • 1x10-6 atmospheres per 1 atmosphere of air • A ppm in water is 1x10-3grams /L water

5. Ozone • let’s convert 1 ppm ozone to grams/m3 • start with: 1x10-6 m3 per 1 m3 air • we need to convert the volume 1x10-6 m3 of O3 to grams • let’s 1st convert gas volume to moles and from the molecular weight convert to grams • at 25oC or 298K one mole of a gas= 24.45liters or 24.45x10-3 m3

6. Ozone • we have 1x10-6 m3 of ozone in one ppm • so: 1x10-6 m3 --------------------- = #moles O3 24.45x10-3 m3/mol • O3 has a MW of 48 g/mole • so # g O3 in 1ppm = #moles Ox 48g/mole per m3 • = 4.1x10-5 g/m3

7. Ozone Health Effects • Ozone causes dryness in the throat, irritates the eyes, and can predispose the lungs to bacterial infection. • It has been shown to reduce the volume or the capacity of the lungs for air • School athletes perform worse under high ambient O3 concentrations, and asthmatics have difficulty breathing • The current US standard has been just reduced from 0.12 ppm for one hour to 0.075 ppm for 8 hours.

8. Lung function before exposure to O.32 ppm O3

9. Lung function after exposure to O.32 ppm O3

10. Athletic performance

11. How do we measure Ozone • 40 years ago chemists borrowed techniques that were developed for water sampling and applied them to air sampling • for oxidants, of which O3 is the highest portion, a technique called “neutral buffered KI was used. • a neutral buffered solution of potassium iodide was placed in a bubbler

12. How do we measure Ozone • a neutral buffered solution of potassium iodide is placed in a bubbler • KI + O3 --> I2 • measure I2

13. How do we measure Ozone • a neutral buffered solution of potassium iodide is placed in a bubbler • KI + O3 --> I2 • measure I2 KI solution

14. How do we measure Ozone • A top is added to the bubbler so that air can enter the KI solution • KI + O3 --> I2 • measure I2 KI solution

15. How do we measure Ozone • a pump is attached to the bubbler pump KI solution

16. How do we measure Ozone • Air goes in through the top of the bubbler and oxidants are trapped in the KI liquid and form I2 Air goes in KI solution + I2

17. How do we measure Ozone • The absorbance of the I2 in the KI solution is then measured with a spectrophotometer KI solution + I2

18. How do we measure Ozone • The absorbance of the I2 in the KI solution is then measured with a spectrophotometer KI solution + I2

19. How do we measure Ozone • The absorbance of the I2 in the KI solution is then measured with a spectrophotometer Spectrophotometer KI solution + I2

20. A calibration curve • A standard curve is constructed from known serial dilutions of I2 in KI solution • to do this I2 is weighed out on a 4 place balance and diluted with KI solution to a known volume

21. A calibration curve • A standard curve is constructed from known serial dilutions of I2 in KI solution • to do this I2 is weighed out on a 4 place balance and diluted with KI solution to a known volume I2

22. Serial dilutions from stock solution I2

23. Serial dilutions from stock solution I2 5 mg/Liter

24. Serial dilutions from stock solution I2 5 3 mg/Liter

25. Serial dilutions from stock solution I2 5 3 2 mg/Liter

26. Serial dilutions from stock solution I2 5 3 2 1 mg/Liter

27. Spectrophotometer absorbance absorbances are measured for each of the serially diluted standards

28. Making a plot I2 adsorbances are plotted vs. concentration

29. We then compare our sample absorbance to the standard curve I2 absorbances are plotted vs. concentration absorbance air sample 1 2 3 4 5 concentration (mg/liter)

30. Problems • anything that will oxidize KI to I2 will give a false positive response • NO2, PAN, CH3-(C=O)-OO-NO2, give positive responses • SO2 gives a negative response

31. Instrumental techniques of measuring Ozone • Chemilumenescene became popular in the early 1970s • For ozone, it is reacted with ethylene • ethylene forms a high energy state of formaldehyde, [H2C=O]* • [H2C=O]*--> light + H2C=O • A photomultiplyer tube measures the light • The amount of light is proportional O3

32. Chemilumenescence measurement of Ozone pump sample air with O3 O3 PM tube {H2C=O}* waste ethylene ethylene catalytic converter CO2 + H2O

33. Using UV photometry to measure Ozone • This is the most modern technique for measuring ozone • sample air with O3 enters a long cell and a 254 nm UV beam is directed down the cell. • at the end of the cell is a UV photometer which is looking at 254 nm light • we know that: light Intensityout= light intensityin e- aLC

34. Photochemical Reactions • Oxygen (O2) by itself does not react very fast in the atmosphere. • Oxygen can be converted photochemically to small amounts of ozone (O3). O3 is a very reactive gas and can initiate other processes. • In the stratosphere O3 is good, because it filters uv light. At the earth's surface, because it is so reactive, it is harmful to living things

35. In the stratosphere O3 mainly forms from the photolysis of molecular oxygen (O2) • O2 + uv light -> O. • O. + O2 +M --> O3 + M • In the troposphere nitrogen dioxide from combustion sources photolyzes • NO2 + uv or visible light -> NO + O. • O. + O2 +M --> O3 (M removes excess energy and stabilizes the reaction)

36. O3 can also react with nitric oxide (NO) • O3 + NO -> NO2 + O2 • both oxygen and O3 photolyzes to give O.O2 + hn-> O. +O. (stratosphere) O3 + hn -> O. + O2 • O. can react with H2O to form OH. radicalsO. + H2O -> 2OH.

37. OH. (hydroxyl radicals) react very quickly with organics and help “clean” the atmosphere; for example: • OH. + H2C=CH2-> products ;very very fast • If we know the average OH. radical concentration, we can calculate the half-life or life time of many organics [org] in the atmosphere.

38. from simple kinetics we can show that: d[org]/dt = -krate [org] [OH] If [OH.] is constant • ln [org]t = ln [org]t=o -krate[OH.]x time1/2 • Let’s say we want to know the time it takes for the organic to go to 1/2 its original [conc].

39. ln [org]t = ln [org]t=o -krate[OH.]x time1/2 • rearranging ln {[org]t / [org]t=o }= -krate[OH.]x t1/2 • The time that it takes for the conc to go to half means [org]t will be 1/2 of its starting conc. [org]t=o . • This means [org]t / [org]t=o = 1/2 • and ln (1/2) = -0.693= -krate[OH.]x t1/2

40. if we use CO as an example, it has a known rate constant for reaction with OH. • CO + OH. -> CO2 krate= 230 ppm-1 min-1 • If the average OH. conc. is 3 x10-8 ppm • for t1/2 we have: ln(1/2) = -krate[OH.] x t1/2 • -0.693= -230 ppm-1 min-1 x 3 x10-8ppm x t1/2 • t1/2 = 100456 min or69.7days

41. What this means is that if we emit CO from a car, 69.7 days later its conc. will be 1/2 of the starting amount. In another 69.7 days it will be reduced by 1/2 again. • For the same average OH. conc. that we used above, what would be the t1/2 in years for methane and ethylene, if their rate constants with OH. radicals are 12.4 and 3840 ppm-1 min-1 respectively? CH4 H2C=CH2

42. Why is the reaction of OH. with ethylene so much faster than with methane? H H1. H-C-H....OH . -> H-C. + .HOH . H H 2. H2C=CH2attack by OH.is at the double bond, which is rich in electrons

43. What happens in urban air?? • In urban air, we have the same reactions as we discussed before • NO2 + uv light -> NO + O. • O. + O2 +M --> O3 + M • O3 + NO -> NO2 + O2 • This is a do nothing cycle

44. What is the key reaction that generates ozone at the surface of the earth? • What is the main reaction that generates it in the stratosphere? • How would you control O3 formation?

45. In the urban setting there are a lot of ground base combustion sources Exhaust hydrocarbonsNO & NO2 CO aldehydes

46. If organics are present they can photolyze or generate radicals • H2C=O + hn -> .HC=O + H. • H. + O2 -> .HO2 • if we go back to the cycleNO2 + uv light -> NO + O. O. + O2 +M --> O3 + M O3 + NO -> NO2 + O2 • .HO2 can quickly oxidize NO to NO2 • NO + .HO2 -> NO2 + OH.

47. OH. + can now attack hydrocarbons which makes formaldehyde and other radical products • for ethylene CH2=CH2 + OH.-> OHCH2CH2.OHCH2CH2. + O2-> OHCH2CH2O2. • OHCH2CH2O2. + NO ->NO2+ OHCH2CH2O. • OHCH2CH2O. + O2 -> H2C=O + .CH2OH • O2 + .CH2OH -> H2C=O + .HO2

48. These reactions produce a host of radicals which “fuel” the smog reaction process First OH radicals attack the electron rich double bond of an alkene Oxygen then add on the hydroxy radical forming a peroxy-hydroxy radical the peroxy-hydroxy radical radical can oxidize NO to NO2,just like HO2 can

49. Further reaction takes place resulting in carbonyls and HO2 which now undergo further reaction; the process then proceeds…

50. What if NO is very low; the primary reaction is with HO2 Jenkin et al, 1997 Atmos. Envir., Pitts and Pitts , 186 RO2.+ HO2. ROOH + O2 ROOH carbonyl +H2O +O2 andROOH alcohol + O3 (dark O3) does the R group matter??? rate constants range: 5.2x10-12 R is CH3 7.7x10-12R is C2H5 1.5x10-11larger