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KINETICS AND ATMOSPHERIC CHEMISTRY NC A&T Lecture February 1, 2011 John Orlando orlando@ucar

KINETICS AND ATMOSPHERIC CHEMISTRY NC A&T Lecture February 1, 2011 John Orlando orlando@ucar.edu. From Wikipedia, the free encyclopedia Boulder Kinetics “… a race from the banks of Boulder Reservoir and back by human-powered vehicles timed on speed and judged for style.”

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KINETICS AND ATMOSPHERIC CHEMISTRY NC A&T Lecture February 1, 2011 John Orlando orlando@ucar

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  1. KINETICS AND ATMOSPHERIC CHEMISTRYNC A&T LectureFebruary 1, 2011John Orlandoorlando@ucar.edu

  2. From Wikipedia, the free encyclopedia Boulder Kinetics “… a race from the banks of Boulder Reservoir and back by human-powered vehicles timed on speed and judged for style.” The Kinetics, a rock band Kinetics (physics), the study of motion and its causes What is KINETICS?

  3. Chemical kinetics - the study of chemicalreaction rates What is KINETICS? Oxford English Dictionary – “A field of study concerned with the mechanisms and rates of chemical reactions or other kinds of process; But first, let’s review a little !!

  4. What do we know so far? Macroscopic properties of gases: For atmospheric T and P, the atmosphere is an ideal gas. PV = nRT Scale Height: Lapse Rate: This is the Dry Lapse Rate, ≈10 K/km In practice, air contains humidity → 7 K/km

  5. What do we know so far? Structure of the atmosphere From Lutgens and Tarbuck, 2001

  6. What do we know so far? General Motions of Air:

  7. Atmospheric Composition Mostly ‘inert’ species – N2, O2, H2O, CO2, Ar Not much chemistry? What do we know so far?

  8. Atmospheric Composition Mostly ‘inert’ species – N2, O2, H2O, CO2, Ar LOTS OF REACTIVE TRACE GASES !! So, Actually, Lots of Chemistry!Natural sources – NO from soil, many hydrocarbons (isoprene) from plants Anthropogenic sources – hydrocarbons, NO, … What do we know so far?

  9. Atmospheric Composition (besides the basic N2, O2, H2O, etc.) Emissions – Lots of “stuff” out there. What do we know so far? Urbanski et al., Wildland Fires and Air Pollution, 2009

  10. Atmospheric Composition (besides the basic N2, O2, H2O, etc.) Emissions – Lots of “stuff” out there. What do we know so far? Urbanski et al., Wildland Fires and Air Pollution, 2009 The atmosphere needs a way to remove these species.

  11. GENERAL DESCRIPTION: The atmosphere (particularly the troposphere) acts as a low-temperature, slow-burning combustion engine. Takes all the emissions (reduced compounds) and ‘burns’ (oxidizes) them: CH4 CO2 + H2O Isoprene Other by-products, such as O3, particles, acids, DMS, NH3 nitrates, etc. (2ry POLLUTANTS)

  12. GENERAL DESCRIPTION: The atmosphere (particularly the troposphere) acts as a low-temperature, slow-burning combustion engine. Takes all the emissions (reduced compounds) and ‘burns’ (oxidizes) them: CH4 CO2 + H2O Isoprene Other by-products, such as O3, particles, acids, DMS, NH3 nitrates, etc. (2ry POLLUTANTS)

  13. GENERAL DESCRIPTION: The atmosphere (particularly the troposphere) acts as a low-temperature, slow-burning combustion engine. Takes all the emissions (reduced compounds) and ‘burns’ (oxidizes) them: OH HO2 CH4 CO2 + H2O Isoprene Other by-products, such as O3, particles, acids, DMS, NH3 nitrates, etc. (2ry POLLUTANTS) NO NO2

  14. MACROSCOPIC : (1044 molecules in the atmosphere) MICROSCOPIC : (about 25 molecules in a 10 nm cube) KINETIC THEORY OF GASES

  15. KINETIC THEORY OF GASES: (most any Physical Chemistry text, including Adamson, 1979; Atkins, 1986). 1) Molecules Move !! (they have kinetic energy): Average Velocity: For N2, can show that c is about 4 x 104 cm/sec at 298 K

  16. KINETIC THEORY OF GASES: (most any Physical Chemistry text, including Adamson, 1979; Atkins, 1986). 2) Molecules bump into walls!! (pressure on wall of a vessel) Pressure:

  17. KINETIC THEORY OF GASES: (most any Physical Chemistry text, including Adamson, 1979; Atkins, 1986). 3) Molecules collide with each other!! Frequency of collisions is about 5 x 109 sec-1 / atm at 298 K. With velocity of 4 x 104 cm/s, Mean Free Path = 7 x 10-6 cm at atmospheric P.

  18. KINETIC THEORY OF GASES: (most any Physical Chemistry text, including Adamson, 1979; Atkins, 1986). 3) Molecules collide with each other!! Frequency of collisions is about 5 x 109 sec-1 / atm at 298 K. With velocity of 4 x 104 cm/s, Mean Free Path = 7 x 10-6 cm at atmospheric P. 4) Molecules can react with each other when they collide !

  19. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) First, a couple of definitions. ELEMENTARY REACTION From Wikipedia - An elementary reaction is a chemical reaction in which one or more of chemical species react directly to form products in a single reaction step. Usually involves 1-3 molecules, with bimolecular most common: e.g., OH + CH4 CH3 + H2O

  20. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) • COMPLEX REACTION SCHEME OR MECHANISM • Made up of a bunch of elementary reactions • - e.g., the oxidation of CH4 to CH2O in the polluted troposphere leads to the following net effect: • CH4 + 4 O2 CH2O + H2O + 2 O3 • OH + CH4 CH3 + H2O • CH3 + O2  CH3O2 • CH3O2 + NO  CH3O + NO2 • CH3O + O2  CH2O + HO2 • HO2 + NO  OH + NO2 • NO2 + hn  NO + O • NO2 + hn  NO + O • O + O2  O3 • O + O2  O3

  21. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) • COMPLEX REACTION SCHEME OR MECHANISM • Explicitly describing the chemical mechanism occurring in the troposphere would probably require > millions of reactions. • - e.g., Aumont et al., ACP 2005.

  22. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) • BACK TO ELEMENTARY REACTIONS (BIMOLECULAR) • Bimolecular reactions are the most common type of elementary reaction in the atmosphere • Typically are of the form A-B + C  A + B-C • CH4 + OH  CH3 + HOH • Rate of the chemical reaction (disappearance of reactants or appearance of products): • k is the rate constant, units of (time)-1 (concentration)-1 • [AB] and [C] are concentrations • Then rate in units of (concentration) (time)-1

  23. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) ELEMENTARY REACTIONS (BIMOLECULAR) What determines the value of the rate constant??? Recall: Frequency of collisions is about 5 x 109 sec-1 / atm at 298 K.

  24. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) ELEMENTARY REACTIONS (BIMOLECULAR) Frequency of collisions is about 5 x 109 sec-1 / atm at 298 K. So, let’s say that OH and CH4 are the two reactants: OH + CH4 H2O + CH3 In the troposphere, typically daytime [OH] = 4.1 x 10-14 atm [CH4] = 1.86 x 10-6 atm So, IF REACTION OCCURRED ON EVERY COLLISION, Rate = k [OH] [CH4] = (5e9 sec-1/atm-1) * (4.14e-14 atm) * (1.86e-6 atm) = 3.8e-10 atm sec-1

  25. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) ELEMENTARY REACTIONS (BIMOLECULAR) So, IF REACTION OCCURRED ON EVERY COLLISION, Rate = (5 x 109 sec-1/atm-1) * (4.1 x 10-14 atm) * (1.86 x 10-6 atm) = 3.8x 10-10 atm sec-1 Usually, we work in molecules cm-3, rather than in atmospheres So, given that there are 2.45 x 1019 molecule cm-3 in 1 atm of gas at 298 K: Rate = (2 x 10-10 cm3 molecule-1 s-1) * (1 x 106 molecule cm-3) * (4.56 x 1013 molecule cm-3) = 9.1 x 109 molecule cm-3 s-1 BUT IN REALITY, REACTION DOES NOT OCCUR ON EVERY COLLISION!!! Why NOT? – any ideas?

  26. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) • ELEMENTARY REACTIONS (BIMOLECULAR) • BUT IN REALITY, REACTION DOES NOT OCCUR ON EVERY COLLISION!!! • Two main reasons: • There are energetic limitations. Colliding molecules must possess sufficient energy to overcome an ‘activation energy’ that typically exists. • Also, there may be ‘geometrical limitations’. Molecules must approach each other in such a way that the appropriate bonds can break / form.

  27. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) ELEMENTARY REACTIONS (BIMOLECULAR) BUT IN REALITY, REACTION DOES NOT OCCUR ON EVERY COLLISION!!! Ea k = A * exp(-Ea/RT) A is the pre-exponential factor, and accounts for the geometry limitations. Ea is activation energy. From Wikipedia

  28. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) ELEMENTARY REACTIONS (BIMOLECULAR) k = A * exp(-Ea/RT) So, Let’s go back to the OH / CH4 reaction. IF REACTION OCCURRED ON EVERY COLLISION, k = 2 x 10-10 cm3 molecule-1 s-1 Turns out that k = 2.45 x 10-12 * exp(- 3525 cal / RT) k = 6.3 x 10-15 cm3 molecule-1 s-1 at 298 K k = 5.2 x 10-16 cm3 molecule-1 s-1 at 210 K Only about 1 in 30000 OH/CH4 collisions results in reaction at 298 K.

  29. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) ELEMENTARY REACTIONS (BIMOLECULAR) Another example HO2 + NO  OH + NO2 – two radical species; no barrier to reaction (attractive forces). HO2 + NO HOO-NO OH + NO2 Reaction turns out to have a “negative activation energy”. k = 3.5 x 10-12 exp(500 cal / RT) cm3 molecule-1 s-1 (Colder molecules more likely to react – less able to overcome attraction).

  30. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) ELEMENTARY REACTIONS (BIMOLECULAR)

  31. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) EQUILIBRIUM – All elementary reactions are reversible. At equilibrium, rate of forward and reverse reactions must be equal. [ HO…H-CH3 ] Ea = 3525 calories Ea = 17450 calories OH + CH4 kf [OH] [CH4] = kr [CH3] [H2O] HOH + CH3

  32. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) EQUILIBRIUM – All elementary reactions are reversible. At equilibrium, rate of forward and reverse reactions must be equal. [ HO…H-CH3 ] Ea = 3525 calories Ea = 17450 calories kf = 6.3e-15 cm3 molec-1 s-1 kr = 1.2e-25 cm3 molec-1 s-1 OH + CH4 kf [OH] [CH4] = kr [CH3] [H2O] HOH + CH3

  33. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) “Equilibrium” and “Steady-State” are different: Equilibrium is a very precise, physical concept - established when forward and reverse rates of all reactions in a system are equal. Steady-State is more conceptual and approximate - A (short-lived) species, often an intermediate in a chemical scheme, is being produced and destroyed at roughly the same rate. Production Rate = Loss Rate O(1D) + H2O OH Reaction with CH4 HO2 + NO Reaction with CO Reaction with Isoprene

  34. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Let’s look at this in more detail. Consider a reaction scheme like this: OH + CH4 CH3 + H2O CH3 + O2  CH3O2 CH3O2 + NO  CH3O + NO2 CH3O + O2  CH2O + HO2 HO2 + NO  OH + NO2 NO2 + hn  NO + O NO2 + hn  NO + O O + O2  O3 O + O2  O3 Can assume CH3 to be in ‘steady-state’. (CH3O, CH3O2, OH, HO2, NO, NO2 too). What is the steady-state [CH3]? (What do we need to know?)

  35. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) OH + CH4 CH3 + H2O CH3 + O2  CH3O2 CH3O2 + NO  CH3O + NO2 CH3O + O2  CH2O + HO2 HO2 + NO  OH + NO2 NO2 + hn  NO + O NO2 + hn  NO + O O + O2  O3 O + O2  O3 Can assume CH3 to be in ‘steady-state’. (CH3O, CH3O2, OH, HO2, NO, NO2 too). Assume [OH] = 106, [CH4] = 4.6 x 1013, [O2] = 5 x 1018, all in molecule cm-3 Appropriate rate constants: k1 = 6.3 x 10-15 cm3 molecule-1 s-1 k2 = 1 x 10-12 cm3 molecule-1 s-1

  36. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) OH + CH4 CH3 + H2O CH3 + O2  CH3O2 CH3O2 + NO  CH3O + NO2 CH3O + O2  CH2O + HO2 HO2 + NO  OH + NO2 NO2 + hn  NO + O NO2 + hn  NO + O O + O2  O3 O + O2  O3 Can assume CH3 to be in ‘steady-state’. (CH3O, CH3O2, OH, HO2, NO, NO2 too). Assume [OH] = 106, [CH4] = 4.6 x 1013, [O2] = 5 x 1018, all in molecule cm-3 Appropriate rate constants: k1 = 6.3 x 10-15 cm3 molecule-1 s-1 k2 = 1 x 10-12 cm3 molecule-1 s-1 k1[OH][CH4] = k2[O2][CH3]ss

  37. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) OH + CH4 CH3 + H2O CH3 + O2  CH3O2 CH3O2 + NO  CH3O + NO2 CH3O + O2  CH2O + HO2 HO2 + NO  OH + NO2 NO2 + hn  NO + O NO2 + hn  NO + O O + O2  O3 O + O2  O3 Can assume CH3 to be in ‘steady-state’. (CH3O, CH3O2, OH, HO2, NO, NO2 too). Assume [OH] = 106, [CH4] = 4.6 x 1013, [O2] = 5 x 1018, all in molecule cm-3 Appropriate rate constants: k1 = 6.3 x 10-15 cm3 molecule-1 s-1 k2 = 1 x 10-12 cm3 molecule-1 s-1 k1[OH][CH4] = k2[O2][CH3]ss [CH3]ss = 0.06 molecule cm-3(Very small !!)

  38. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions). (Strangely enough, closely related to “unimolecular reactions” as we will see in a minute or two). e.g., NO3 + NO2 N2O5 N2O5  NO3 + NO2 - not as simple as they look (not elementary reactions) So, what is actually going on?

  39. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions). e.g., NO3 + NO2 N2O5 NO3 + NO2 (N2O5) * (N2O5) * has a choice – forwards or backwards N2O5 NO3 + NO2 = (N2O5) * (N2O5) * = NO3 + NO2 (N2O5) * + M = N2O5 + M where “M” = N2, or to a lesser extent, O2

  40. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions). NO3 + NO2 = (N2O5) * ka (N2O5) * = NO3 + NO2 kb (N2O5) * + M = N2O5 + M kc, where “M” = N2 or O2 Interested in the rate of formation of (stabilized) N2O5: Assume steady-state for (N2O5)* : ka [NO3] [NO2] = { kb + kc [M] } [(N2O5)*] Or [(N2O5)*] = { ka [NO3] [NO2] } / { kb + kc [M] } Substitution yields

  41. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions). NO3 + NO2 = (N2O5) * ka (N2O5) * = NO3 + NO2 kb (N2O5) * + M = N2O5 + M kc, where “M” = N2 or O2 Substitution yields: Consider low-pressure limit, [M]  0 (termolecular) And high-pressure limit [M]   (bimolecular)

  42. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions). NO3 + NO2 = (N2O5) * ka (N2O5) * = NO3 + NO2 kb (N2O5) * + M = N2O5 + M kc, where “M” = N2 or O2 Unfortunately, life is even more complicated. AHHH!!! Jurgen Troe: where Fc is the “broadening factor”.

  43. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions). NO3 + NO2 = (N2O5) * ka (N2O5) * = NO3 + NO2 kb (N2O5) * + M = N2O5 + M kc, where “M” = N2 or O2

  44. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions). (Strangely enough, closely related to “unimolecular reactions”) e.g., NO3 + NO2 N2O5 N2O5  NO3 + NO2 - not as simple as they look (not elementary reactions) N2O5 + M  (N2O5) * + M (N2O5) * + M  N2O5 + M (N2O5) *  NO2 + NO3 + MAnalogous treatment: Consider low-pressure limit, [M]  0 (bimolecular) And high-pressure limit [M]   (unimolecular)

  45. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Termolecular reactions (three-body reactions) / Unimolecular reactions: Main “Take-home” Message: Set of chemical reactions that lead to recombination of reactive species (radicals), and formation of reservoirs: NO3 + NO2 + M  N2O5 + M OH + NO2 + M  HNO3 + M HO2 + NO2 + M  HO2NO2 + M CH3C(O)OO + NO2 + M  CH3C(O)OONO2 + M ClO + NO2 + M  ClONO2 + M ClO + ClO + M  ClOOCl + M Often reversible (equilbrium). Even though the formation of the reservoir is exothermic, the reverse reaction results in a gain in entropy to compensate.

  46. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Photolysis reactions: (brief introduction to next week). Some generalities: - Sunlight provides energy across the electromagnetic spectrum, which can be absorbed by molecules. - Energies of chemical bonds typically correspond to UV photons. - Thus, absorption of UV sunlight can lead to photolytic destruction of certain molecules. e.g., NO2 + hn NO + O(3P) “j-value” – unimolecular ‘rate constant’. Vary with spectral properties of the molecule of interest, but also with solar intensity (as a fn of wavelength)

  47. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Photolysis reactions: (brief introduction to next week). e.g., NO2 + hn NO + O(3P) Assume constant solar intensity (j-value is constant), and assume no production of NO2: Then, rearrangement and integration leads to: [NO2]t = [NO2]o exp -(j*t) (exponential decay of NO2)

  48. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) Photolysis reactions: (brief introduction to next week). e.g., NO2 + hn NO + O(3P) O2 + hn  O(1D) + O(3P), upper stratosphere and above O2 + hn  O(3P) + O(3P), stratosphere and above O3 + hn  O(1D) + O2 O3 + hn  O(3P) + O2 NO3 + hn  NO2 + O(3P) NO3 + hn  NO + O2 Important at all altitudes CH2O + hn  HCO + H CH2O + hn  CO + H2 HONO + hn  OH + NO

  49. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) HOW DO WE MEASURE RATE CONSTANTS IN THE LAB??? TWO BASIC METHODS: TIME–RESOLVED METHODS (includes “Flash Photolysis” and “Flow Tube”) INDIRECT METHODS (“Relative Rate” method)

  50. REACTION KINETICS: (follows Brasseur, Orlando and Tyndall, pp. 95-114.) HOW DO WE MEASURE RATE CONSTANTS IN THE LAB??? “Flash Photolysis” – e.g., Bryukov et al., J. Phys. Chem. A., 2004, v. 108, 10464-10472. OH + CH4 H2O + CH3 Basic requirements: A method for getting CH4 in the vessel, and knowing its concentration. A method of generating reactive radicals (in this case, OH) ‘instantaneously’.

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