The air / environment (geosphere): Is it a reactor ?
It‘s a matter of reactions and transports and mixing !
mixing times:
• vertically lower few kilometers (boundary layer) 1h-1d, mixing with free
troposphere 2-10 days
• around the globe on the same latitude (zonal transport) within 1-4 weeks
0. Basics
0.1 Chemical conversion in the atmosphere - conceptually
It‘s a matter of reactions and transports and mixing !
mixing times:
• around the globe on the same latitude (zonal transport) within 1-4 weeks
Atmospheric transport after the Fukushima Dai-ichi accident (Marzo, 2014)
The air / environment (geosphere): Is it a reactor ?
It‘s a matter of reactions and transports and mixing !
mixing times:
• vertically lower few kilometers (boundary layer) 1h-1d, mixing with free
troposphere 2-10 days
• around the globe on the same latitude (zonal transport) within 1-4 weeks
• from mid latitudes to the pole (meridional transport)
within days to weeks, hemispheric mixing 2-6 months
• between hemispheres 1 year
• troposphere-stratosphere 1-3 years, mesosphere (>50 km) 5-8 years
Substance concentration changes result from:
• (chemical) reactions
• (meteorological) transports and mixing
Tools to understand:
• 3D transport models
• model parameters isolated in laboratory experiments
• significant ‚ingredients‘ identified in the ‚field‘
Idealized situations of atmospheric
transport and mixing
0. Basics
0.1 Chemical conversion in the atmosphere - conceptually
Pressure and temperature profiles in the atmosphere
Pressure[hPa]
Heightabovesealevel[km]
Atmospheric pressure and composition
Heightabovesealevel[km]
Pressure[hPa]
Pressure decline with altitude
ρ = 1.225 mg/cm, g = 981 cm/s² , R = 8.206 Pa cm³/mol/K
 z/p = 8 m/hPa at ground
16 m/hPa in 6km altitude
„barometric step“
p(z) = p0 e-zg/RT (T/ z = 0) → p(z,T)  p0 e-zg/RT
‚Standard‘ atmosphere(US Std 1962)
km T[°C] p[hPa] Mg[g/mol] NLu/V[molec/cm³]
0 0 1013 28.964 2.69x1019
0 15 1013 28.964 2.55x1019
0 25 1013 28.964 2.46x1019
3 - 4.4 700 28.964 1.76x1019
10 -49.9 265 28.964 6.67x1018
20 -56.5 55 28.964 1.38x1018
30 -46.6 12 28.964 0.30x1018
100 -63 0.00021 28.5 5.29x1012
Units for quantification of atmospheric trace substances
Ideal gas law: pV = nRT = mRT/Mg
Universal gas constant R = 0.082 at L/(mol K) = 8.314 J/(mol K) = NAkB
Avogadro‘s number NA= 6.023×1023 molec/mol
Boltzmann constant kB = 1.38×10-23 J/K
‚Molar‘ volume at T0=273 K and p0=101325 Pa: V = 22.414 L/mol
 ‚Molar‘ mass Mg air  28.9 g/mol
Concentration ci = mi/V [µg/m³] (for gases: = density)
1 at = 1013 hPa, 1 Pa = 1 N/m² = 1 J/m³
Mass mixing ratio µm i = ci/c [ , %, ppmm, ppbm]
Partial pressure pi = niRT/Vi [Pa]
Volume mixing ratio µV i = pi/p = Vi/V [ , %, ppmV, ppbV]
Concentration ni/V = pi/RT [mol/m³]
Number density Ni/V = niNA/V = piNA/RT [molec/cm³]
More details: Schwartz & Warneck (1995): Units for use in
atmospheric chemistry, Pure Appl. Chem. 67, 1377-1406
billion = 109 (Am., not Brit.)
TROPOSPHERE - COMPOSITION
(sum) 10-13_1!
Reaction types, kinetics
Quantifying chemical transformations: Law of mass action
Law of mass action (dt.: Massenwirkungsgesetz):
The ratio of the product of the concentrations of the products and the product of the
concentrations of the reactants (= educts) is constant for a given temperature and
pressure in a homogeneous (i.e. single-phase) reaction (Guldberg & Waage, 1867):
K1 = cC
ccD
d / (cA
acB
b) K1 = equilibrium constant (dt.: Gleichgewichtskonstante)
implications:
A reaction, A+B, will cease, when K1 is achieved
The rates of formation and decay of the products are equal when equilibrium is
established, i.e. k1 cC
ccD
d = k-1 cA
acB
b and K1 = k1/k-1
Reaction (1) aA + bB  cC + dD
Reactants Products
Back reaction (-1) cC + dD  aA + bB
Equilibrium (1) aA + bB ⇌ cC + dD
or: (1, -1)
Reaction types, kinetics
Reaction rate coefficient k
Temperature dependence:
-dci/dt = kT ci
kT = A(T) e[E/(RT)]
Universal gas constant R = kB NA= 1.38 ×10-23 J/K × 6.023×1023/mol
Providing the energy is sufficiently large, the temperature dependence
of A is negligible, and kT follows the Arrhenius expression:
kT = A e[-Ea/(RT)]
with activation energy Ea
Frequently used, too: van t‘Hoff expression:
kT = B e[- Ea/R (1/T – 1/Tref)],
The two expressions are equal via: A = B e[Ea/(RTref)]
Homogeneous gas-phase reactions
The rate law of a reaction of the general form
aA + bB  cC + dD
Is defined as
Rate (dt.: Rate) = -dcA/dt/a = -dcB/dt/b = dcC/dt/c = dcD/dt/d
example:
2 NO + O2  2 NO2
Rate = -dcNO/dt/2 = -dcO2/dt = dcNO2/dt/2
Reactions can be unimolecular, bi- or termolecular.
Second order
Usually: bimolecular
A + B  C + D
A + B  C ; A, B, C, D molecules or radicals
example:
O + O2  O3
NO + O3  NO2 + O2
Reaction rate: 2nd order (1+1=2) dcC/dt = -dcA/dt = -dcB/dt = k cA
1 cB
1
Reaction rate coefficient: k(2)
k = A exp [-(Ea/R)T]
Arrhenius expression, preexponential factor A, activation energy Ea
k(T-1)  slope –Ea/R, intercept ln A, Ea/R > 0 ↔ faster at higher T
The reaction order is given by the sum of the exponentials,
n+m+..., of the concentration terms in the rate law of the form
-dcA/dt = k cA
n cB
m (n = zero or integer or fraction*)
It is determined empirically.
* ‚overall‘ reactions only
Steil, Crutzen, et al., 1998
k(T)
Example
NO + O3  NO2 + O2
k = A e[-(Ea/R)T]
Arrhenius expression, preexponential factor A, activation energy Ea
k(T-1)  slope m = –Ea/R, intercept ln A, Ea/R > 0 ↔ faster at higher T
A = 2×10-12 cm³/molec/s
E/R = -1400 K

k298 K = 1.8 ×10-14 cm³/molec/s
k230 K = 0.45×10-14 cm³/molec/s
First order (but not uni-mo
if cB >> cA
A + B  produc
-dcB/dt  0
mechanism: Steil et al., 1998
Quasi steady state approximation
Lindemann-Hinshelwood mechanism:
(1+2) A + M  B + M
Here, M stands for any molecule or atom (i.e. N2, O2,...),
not transformed but required to absorb excess energy,
e.g. of an activated intermediate state:
(1) A + M  A* + M -dcA/dt = k1
(2)cAcM
(-1) A* + M  A + M -dcA*/dt = k-1
(2)cA*cM
(2) A*  B dcB/dt = k2
(1)cA*
A* in steady state:
dcA*/dt = k1
(2)cAcM - k-1
(2)cA*cM - k2
(1)cA* = 0
cA* = k1
(2)cAcM/(k-1
(2)cM + k2
(1))
dcB/dt = k2
(1) k1
(2)cAcM/(k-1
(2)cM + k2
(1))
If (-1) much faster than (2):
k-1
(2)cA*cM >> k2
(2)cA*, then
k-1
(2)cM + k2
(1)  k-1
(2)cM and
dcB/dt = k2
(1) k1
(2)cA/k-1
(2)
the overall process (1+2) is first order in cA.
Shifts to second order for cM  0 (i.e., low pressure):
dcB/dt = k2
(1) k1
(2)cAcM/k2
(1)
Example: thermic dissociation
HOONO2  HO2+ NO2
Reaction rate
dcC/dt = -dcA/dt = k(1)cA
Example
(1, -1) O + O2 ⇌ O3*
(2) O3* + M  O3
cO3* = k1 cO cO2 /(k-1 + k2 cM)
dcO3/dt = k2 cO3* cM
dcO3/dt = k1 k2 cO cO2 cM /(k-1 + k2 cM) = [k1 k2 cM /(k-1 + k2 cM)] cO cO2
High pressure limit, k∞
(2): k-1  0 ⇨ k∞ = k1
Low pressure limit, k0
(2): cM  0 ⇨ k0 = k1 k2 / k-1
k(2)
NO2 + HO2 + M  HO2NO2 + M
k288 K/1000 hPa = 1.9 × 10-12 cm³/molec/s
k288 K/500 hPa = 1.5 × 10-12 cm³/molec/s
k230 K/500 hPa = 2.3 × 10-12 cm³/molec/s
k230 K/1000 hPa = 3.0 × 10-12 cm³/molec/s
Reaction orders
A + B  C + D
is 2nd order or – in case dcB/dt  0 - pseudo-1st order
A + B + M  C + D + M
is 2nd order or – in case dcB/dt  0 - pseudo-1st order or
– if p << 1000 hPa - between 2nd and 3rd order
The unit of a homogeneous gas-phase rate is [molec/cm3/s].
A 1st or pseudo-1st order rate law reads:
-dcA/dt = dcC/dt = k(1) cC, with k(1) [1/s].
A 2nd order rate law reads:
-dcA/dt = dcC/dt = k(2) cC cD, with k(2) [cm3/molec/s].
A 3rd order rate law reads e.g.:
-dcA/dt = k(3) cC
2 cD, with k(3) [cm6/molec2/s].
Photochemical reactions
Absorption of radiation by molecules in the atmosphere
Gaseous molecules absorb
ultraviolet, visible and
infrared light: O3
Consequences:
• photophysical and photochemical molecular processes
• change of spectrum:
Energy ranges, correspondence between energy and wavelength
 = c/ with frequency 
E = hc/ = hc Planck relationship (wavelength , wavenumber ,
Planck‘s constant h = 6.626 x 10-34 Js)
Commonly used energy units:
Energy ranges, correspondence between energy and wavelength
= c/
E = hc/ = hc Planck relationship
(wavelength , wavenumber , Planck‘s constant h = 6.626 × 10-34 Js
The rate of photochemical reactions
Absorption
ln(I0/I) = Nd Beer-Lambert law
I/I0 = e(-Nd)
absorption cross section  (cm², default: base e)
molecule concentration N (cm-3), depth of absorptive layer d (cm)
optical depth OD = Nd
Caution: Most measurements are made to the
base 10 (log(I0/I) = 10Nd) ⇨ × 2.303 to reach base e
Unimolecular
A + h  B + C
A, B, C, D molecules or radicals
example:
O3 + h  O2 + O (3P)
Reaction rate coefficient j (photolysis rate):
dcC/dt = -dcA/dt = j cA
Ground state, A
/excitation state, A*:
A + h  A*
Example:
O + h  O*
i.e. O(3P) + h  O(1D)
Most important class of photochemical reactions:
Photodissociations
Photolysis
The photolysis rate
The photolysis rate, j (s-1), in dcA/dt = j cA is given by:
j = ∫ () () L() d
• quantum yield () ( ),
• absorption cross section  (cm²),
• actinic flux L() (cm-2 s-1)
L is the total intensity of effective light (direct + scattered +
reflected, spherically integrated).
(adopted from: Jacobson, 2005)
The actinic flux L() is a function of the solar zenith angle, cloudiness,
aerosol concentration, and surface albedo.
Spherical coordinate
system: Radial distance
r, polar angle ,
azimuthal angle φ
L is measured using a (2) radiometer or by measuring the
photolytic decay (so-called chemical actinometry).
Its value can be estimated via tabulated values of  and  for
intervals of  and estimates of L() for given conditions.
(Calvert, 1985)
Example jO3: (1) Quantum yield i() for
i = O3 + h  O (1D) + O2 (3g
-)
O3O(1D)(310-320 nm)  0.2
Data src.: Finlayson-Pitts & Pitts, 1998
Example jO3: (2) Absorption cross section (,T) of O3
Example:
O3(310-320 nm)  60 × 10-20 cm²/molec for T = 298 K
Data src.: Finlayson-Pitts & Pitts, 1998
(3) Actinic flux –
determined by radiation absorption in the atmosphere
Actinic flux L() - Example:
For z = 15 km and solar zenith angle
of 40°:
L310-320 nm = (1.69+2.08+
+2.35+2.88+2.95) ×1014 cm-2 s-1
 12 × 1014 s-1
Order of magnitude estimate of
jO3O* for a selected wavelength
interval:
(1-3) jO3O*(310-320 nm) 
 0.2 × 60 × 10-20 × 12 × 1014 s-1 
 10-4 s-1
Src: WMO: Scientific Assessment of Ozone Depletion 2006, Geneva 2007
Tropospheric chemistry
Tropospheric ozone and hydrocarbon chemistry
Tropospheric O3: Significance
Toxicology:
• concentrations > 120-150 µg/m³ are relevant, at least for sensitive persons.
No epidemiological evidences.
• Significant loss of physical performance at higher concentrations, i.e. 
400 µg/m³ (EU, 1992: 180 µg/m³ warning, 360 µg/m³ dangerous)
Climate:
• absorption in the atmospheric ‚window‘ region near  = 9.6 µm 
radiative forcing +0.35±0.15 W m-2 since 1850
Ecotoxicology:
• toxic to plants (uptake through stomatae prevails, radical formation);
sensitive crops (potato, wheat, rye, barley) and trees (larch, pine)
• for same dosis damage is highest under peak concentrations, synergistic
effects with NO2 and SO2
•
36-59°N, 1996 vs. 1970
(WMO, 1998, after Logan & Megretskaia) http://ozone.unep.org/
Trends of ozone - stratospheric and tropospheric
(%/10a)
Tropospheric ozone temporal trends
Background stations
15
20
25
30
35
40
45
50
55
Jan-87
Jan-88
Jan-89
Jan-90
Jan-91
Jan-92
Jan-93
Jan-94
Jan-95
Jan-96
Jan-97
Jan-98
Jan-99
Jan-00
Jan-01
Jan-02
Jan-03
Jan-04
OZONECONC(PPB)
Ozone Baseline monthly means 12-month moving average
Mace Head, IRL (Derwent, 2004)
(courtesy of Parrish et al., 2009)
Mauna Loa Mace Head
Ground stations
(courtesy of Barnes et al., 2011)
Tropospheric ozone temporal trends
(EDGARv4.1;
Cooper et al., 2016)
Tropospheric ozone trends
CO volume mixing ratio in the lower troposphere: 100-200 ppbv
(1) CO + OH  CO2 + H
Chemical fate of OH globally:  2/3 reacts with CO
(2) H + O2  HO2
Sum (1-2): CO + O2 + OH  HO2 + CO2
(6) O3 + NO  O2 + NO2 k(1)  10-2/s
Sum (1-6): CO + O2+ NO  CO2 + NO2
Ozone formation in the troposphere
(1) Why is CO not accumulating in urban air?
 ‚discovery‘ of the OH radical
Ozone formation in CO oxidation
(3) HO2 + NO  OH + NO2 k = 220×10-12 cm³/molec/s
(4) NO2 + h (< 420 nm)  NO + O j  5×10-3/s
(5) O + O2 + M O3 + M k(1)  105/s
Sum (1-5): CO + 2 O2  CO2 + O3
Ozone formation in the troposphere
Leighton relationship
(4) NO2 + h (< 420 nm)  NO + O j4  5×10-3/s
(5) O + O2 + M  O3 + M k5
(1)  105/s
(6) O3 + NO  O2 + NO2 k6
(2)  1×10-14 cm³/molec/s
ozone ‚titration‘ by NOx
dcO3/dt = 0 = k5 cO cO2 cM – k6 cNO cO3 equilibrium within 2 min
hence: cO3 = k5 cO cO2 cM / (k6 cNO) quasi-constant ozone level (f(jNO2))
cO available from:
dcO/dt = 0 = j4 cNO2 – k5 cO cO2 cM (hence: cO = j4 cNO2 /(k5 cO2 cM))
combined: cO3 = j4cNO2/(k6cNO) or: (cO3 cNO)/cNO2 = constant
holds as long as there are no other O3 loss reactions than (6)
Ozone depends on the background level, on NOx and on the
ratio NO2/NO upon emission:
(4) NO2 + h  NO + O
(5) O + O2 + M O3 + M
NO2 + O3 = „ Ox “
Ozone background conc.
Slope = % primary emitted NO2
NO + NO2 = „NOx“
Ozone formation in the troposphere
(Clapp & Jenkin, 2001)
Ozone formation in synthetic atmosphere
NOx + HCx + Light  partly oxygenated HCx + O3
(1a) RCH3 + OH  RCH2
. + H2O
(1b) RCH2
. + O2 + M  RCH2OO. + M
(2) RCH2OO. + NO  RCH2O. + NO2
(3) RCH2O. + O2  RCHO + HO2
(4) NO2 + h (< 430 nm)  NO + O
(5) O + O2 + M O3 + M
Sum
(1-5): RCH3 + OH + 3 O2  RCHO + H2O + HO2 + O3
= catalyzed by NO and light
Sum
(4-5): NO2 + O2 + h  NO + O3
(source: NOx)
Ozone formation from hydrocarbons, HCx, and NOx
(6) HO2 + NO  OH + NO2 = OH recycled
Sum
(1-6): RCH3 +NO + 3 O2  RCHO + H2O + NO2 + O3
= catalyzed by OH and light
2 main pathways: CO or VOC oxidation
(courtesy: Möller, 2003)
CO
VOC
CO2
VOC ´
OH
HO2
NO
2
NO
O3
hv
(1) CH4 + OH  CH3
. + H2O(slow: τCH4  8 a)
(2) CH3
. + O2 + M  CH3OO. + M
(3) CH3OO. + NO  CH3O. + NO2
(4) CH3O. + O2  HCHO + HO2
Sum:
CH4 + OH + 2 O2  HCHO + H2O + HO2
NO2 + O2 + h  NO + O3
HCx: alkanes, example methane
• Although slow, CH4 is a major chemical sink for OH (globally 
1/3 of OH reacts with CH4).
• The so formed ozone is the major contribution to the background
ozone.
• It increases with increasing methane emissions.
(1a) CH3CH2CH2CH3 + OH  CH3CH.CH2CH3 + H2O yd=85%
CH3CH.CH2CH3 + O2 + M  CH3CH(OO.)CH2CH3 + M
(2a) CH3CH(OO.)CH2CH3 + NO  CH3CH(O.)CH2CH3 + NO2
(3aa) CH3CH(O.)CH2CH3 + O2  CH3C(O)CH2CH3 + HO2yd=60%
(3ab) decomposition :  CH3C.O + .CH2CH3 + HO2yd=40%
.CH2CH3 + O2  CH3CH2OO.
(4ab) CH3CH2OO. + NO  CH3CH2O. + NO2
(5ab) CH3CH2O. + O2  CH3CHO + HO2
Sum: CH3CH2CH2CH3 + OH + 2.8 O2 
0.6 CH3C(O)CH2CH3 + 0.8 CH3CHO + H2O + HO2 + 1.4 NO2
HCx = alkanes: branching 1) C atom position attacked, 2) alkoxy radical
(1b) CH3CH2CH2CH3 + OH  .CH2CH2CH2CH3 + H2O yd=15%
CH3CH2CH2CH2
. + O2 + M  CH3CH2CH2CH2OO. + M
(2b) CH3CH2CH2CH2OO. + NO  CH3CH2CH2CH2O. + NO2
(3ba) CH3CH2CH2CH2O. + O2  CH3CH2CH2CHO + HO2 yd=?
(3bb) isomerisation:  .CH2CH2CH2CH2OH + HO yd=?
.CH2CH2CH2CH2OH + O2  .OOCH2CH2CH2CH2OH
(4bb) .OOCH2CH2CH2CH2OH + NO  .OCH2CH2CH2CH2OH + NO2
(5bb) .OCH2CH2CH2CH2OH + O2  CH(O)CH2CH2CH2OH + HO2
Sum: CH3CH2CH2CH3 + OH + 2.8 O2 
0.6 CH3CH2CH2CHO + 0.8 CH2(OH)CH2CH2CHO + H2O + HO2 + 1.4 NO2
O
OH
Alkoxy radicals: reactivity overview
k(1) (103 s-1)
RO. decomposition H-abstraction isomerization
by O2
CH3CH2CH2CH2O. 0.6 200  0
CH3CH2CH2CHO.CH3 17 40 200
CH3CH2CHO.(CH2)2CH3 34 40 200
CH3CHO.(CH2)3CH3 28 40 2000
CH3C(CH3)2CH2O. 9.8 24  80
positive p via 5- or 6-ring
dependence intermediates
Nomenclature:
Saturated and unsaturated C chains: alkanes (dt: Alkane),
alkenes and alkynes (dt: Alkene, Alkine)
Partly oxygenated hydrocarbons: ROH alcohols (dt:
Alkohole), carbonyls: RCHO aldehydes (dt: Aldehyde) and
R2CO ketones (dt: Ketone), RCOOH and R(COOH)2 monoand
dicarboxylic acids (dt: Mono- und Dicarbonsäuren)
Multifunctional partly oxygenated hydrocarbons:
RCHOHCHO -hydroxyaldehydes, RCHOHCOOH hydroxyacids,
...
CH2O
CH3
CH3
CH3
Alkenes are more reactive toward OH than alkanes: k  10-10 cm3 molec-1 s-1
The higher substituted radical is more stable, hence, formed preferentially:
(1a) CH3CH=CH2 + OH  CH3CH.CH2OH addition, yd=66%
CH3CH.CH2OH + O2 + M  CH3CH(OO.)CH2OH + M
(2a) CH3CH(OO.)CH2OH + NO  CH3CH(O.)CH2OH + NO2
(3aa) CH3CH(O.)CH2OH + O2  CH3C(O)CH2OH + HO2 yd=3%
(3ab) decomposition:  CH3C.O + .CH2OH yd=97%
(4ab) .CH2OH + O2  HCHO + HO2
HCx: alkene OH reaction, example C3, i.e. propene
(1b) CH3CH=CH2 + OH  CH3CH(OH)CH2
. addition, yd=34%
CH3CH(OH)CH2
. + O2 + M  CH3CH(OH)CH2(OO.) + M
(2b) CH3CH(OH)CH2(OO.) + NO  CH3CH(OH)CH2O. + NO2
(3ba) CH3CH(OH)CH2O. + O2  CH3CH(OH)CHO + HO2 yd=90%
(3bb) decomposition :  CH3CH(OH). + HCHO yd=10%
(4bb) CH3CH(OH). + O2  CH3CHO + HO2
Sum
(1-4) CH3CH=CH2 + OH + 2.8 O2  0.02 CH3C(O)CH2OH + 0.65 HCHO + 0.3
CH3CH(OH)CHO + 0.03 CH3CHO + H2O + HO2 + NO2
• Most alkenes react with OH addition to the double bond (positive p dependence of
kOH); only for the small alkenes the addition complex does not react further.
• H abstraction is more likely for large and branched alkenes.
• After the O2 addition step ( ROO.), decomposition is the most probable path for
 C4 while isomerisation dominates for > C4
(yields 0.04 for C4 but 0.6 for C8; Kwok et al., 1996)
HCx: alkene OH reaction
Example n-butene:
(1) CH3CH2CH=CH2 + OH  CH3CH2CH(OH)CH2
. addition
CH3CH2CH(OH)CH2
. + O2 + M  CH3CH2CH(OH)CH2OO. + M
(2) CH3CH2CH(OH)CH2OO. + NO  CH3CH2CH(OH)CH2O. + NO2
(3) isomeris.: CH3CH2CH(OH)CH2O.  .CH2CH2CH(OH)CH2OH
(4) .CH2CH2CH(OH)CH2OH + O2  .OOCH2CH2CH(OH)CH2OH
(5) .OOCH2CH2CH(OH)CH2OH + NO  .OCH2CH2CH(OH)CH2OH + NO2
(6) .OCH2CH2CH(OH)CH2OH + O2  CH(O)CH2CH(OH)CH2OH + HO2
Sum: CH3CH2CH=CH2 + OH + 2 NO + 3 O2 
CH(O)CH2CH(OH)CH2OH + H2O + HO2 + 2 NO2
 dihydroxycarbonyl
O
OH
OH
HCx: alkene OH reaction: Example isoprene (= 2-methylbutadiene)
1 or 2 addition 66%
(1a) CH2=C(CH3)CH=CH2 + OH  HOCH2C.(CH3)CH=CH2
 .CH2C(OH)(CH3)C=CH2
+ O2 + NO
  
(4aa) decomposition:  HCHO + CH3C(O)CH=CH2 yd  30%
methyl vinyl ketone
(4ab) isomerisation:  HC(O)CH(CH3)CH=CHOH yd  5%
-hydroxy-(2-methyl)butenal
OH
CH2
CH3
CH2
CH4
CH3
O
3 or 4 addition 34 %
(1b) CH2=C(CH3)CH=CH2 + OH  CH2=C(CH3)CH.CH2OH
 CH2=C(CH3)CHOHCH2
.
+ O2 + NO
  
(4ba) decomposition:  HCHO + CH2=C(CH3)CHO yd  20%
methacrolein
(4bb) 5-ring closure:  yd < 5%
3-methylfuran
(1a) RCH3 + OH  RCH2
. + H2O
(1b) RCH2
. + O2 + M  RCH2OO. + M
(2) RCH2OO. + NO  RCH2O. + NO2
(3) RCH2O. + O2  RCHO + HO2
(4) NO2 + h (< 430 nm)  NO + O
(5) O + O2 + M O3 + M
Sum (1-5):
RCH3 + OH + 3 O2  RCHO + H2O + HO2 + O3
= catalyzed by NO and light
Second option for NO:
NO + O3  NO2 + O2
= catalyst reacts with product
Ozone formation from hydrocarbons, HCx, and NOx
HCx: OH reactivity, products overview
HCx kOH oxygenated No. of NO converted
10-12cm3 intermediates initial from total
molec-1s-1 formed carbonyls
(298 K) (dep. neglected)
Alkanes:
CH4 0.006 HCHO 1 1 2
CH3CH3 0.25 CH3CHO 2 4 6
CH3CH2CH3 1.1 HCHO, CH3CHO, CH3COCH3 3 5 8
CH3CH2CH2CH3 2.4 2 CH3CHO 3 8 11
CH3CH(CH3)CH3 2.2 HCHO, CH3COCH3 3 5 8
CH3(CH2)3CH3 4.0 HCHO, CH3CHO, CH3CH2CHO,
CH3(CH2) 2CHO 3 11 14
Alkenes:
CH2=CH2 8.5 2 HCHO 2 2 4
CH2=CHCH3 26 HCHO, CH3CHO 2 5 7
CH2=CHCH2CH3 31 HCHO, CH3CH2CHO 2 8 10
cis-CH3CHCHCH3 56 2 CH3CHO 2 8 10
trans-CH3CHCHCH3 64 2 CH3CHO 2 8 10
CH2=C(CH3)CH3 51 HCHO, CH3COCH3 2 5 7
Lifetime with respect to OH reaction:
kOH
(2) = 0.01×10-12cm3molec-1s-1  kOH
(1) = 1×10-8s-1  τOH = 4 months
kOH
(2) = 1×10-12cm3molec-1s-1  kOH
(1) = 1×10-6s-1  τOH = 12 days
kOH
(2) = 50×10-12cm3molec-1s-1  kOH
(1) = 5×10-5s-1  τOH = 6 hours
Ozone formation efficiency of various hydrocarbons
Characterisation of VOCs according to their
photochemical ozone creation potential, POCP:
POCP : = mO3 / FVOCi
under defined conditions (ozone formation during several days,
NOx poor) (Carter, 1994; EK, 1994)
POCP
C2H4 100
CH4 0.7
C6H6 18.9
CH3OH 12.3
HCHO 42.1
Emission ratio of road transport source
Tropospheric ozone: Dependence on HCx(VOC) and NOx
emission reductions perspectives
in 2030 under MFR
(= maximum feasible reduction)
(Dentener et al., 2006)
Example critical levels for natural
and agroecosystems
Losses of harvested wheat > 5%, if
accumulated dose exceeding 40 ppbv
> 3000 ppbv h;
similar: SOMO35 [ppbv d] – in 2000:
in 2030 under CLE
(= current legislation)
WHO air qual. Index SOMO35:=
daily max of (8-h running average -
35ppbv ‚bckgrd‘) added over 365 d
Chemical residence time of organic substances in the atmosphere
HCx, oxygenated, halogenated HCx, and hetero atom organics
OH reactivity overview
CO volume mixing ratio in the lower troposphere: 100-200 ppbv
(1) CO + OH  CO2 + H
Chemical fate of OH globally:  2/3 reacts with CO
(2) H + O2  HO2
Sum (1-2): CO + O2 + OH  HO2 + CO2
(6) O3 + NO  O2 + NO2 k(1)  10-2/s
Sum (1-6): CO + O2+ NO  CO2 + NO2
Ozone formation in the troposphere
(1) Why is CO not accumulating in urban air?
 ‚discovery‘ of the OH radical
Ozone formation in CO oxidation
(3) HO2 + NO  OH + NO2 k = 220×10-12 cm³/molec/s
(4) NO2 + h (< 420 nm)  NO + O j  5×10-3/s
(5) O + O2 + M O3 + M k(1)  105/s
Sum (1-5): CO + 2 O2  CO2 + O3
(1a) RCH3 + OH  RCH2
. + H2O
(1b) RCH2
. + O2 + M  RCH2OO. + M
(2) RCH2OO. + NO  RCH2O. + NO2
(3) RCH2O. + O2  RCHO + HO2
(4) NO2 + h (< 430 nm)  NO + O
(5) O + O2 + M O3 + M
Sum
(1-5): RCH3 + OH + 3 O2  RCHO + H2O + HO2 + O3
= catalyzed by NO and light
Sum
(4-5): NO2 + O2 + h  NO + O3
(source: NOx)
Ozone formation from hydrocarbons, HCx, and NOx
(6) HO2 + NO  OH + NO2 = OH recycled
Sum
(1-6): RCH3 +NO + 3 O2  RCHO + H2O + NO2 + O3
= catalyzed by OH and light
Tropospheric ozone formation from CO, CH4: other products
HCHO, H2O2, HNO3 (radical sink reactions)
(courtesy: Ehhalt, 1999)
Degradation of RH in NO-poor areas
(1) RCH3 + OH.  RCH2
. + H2O
(2) RCH2
. + O2 + M  RCH2OO. + M
(3a) NO2 + h  NO + O
(4a) O + O2 + M O3 + M
(3b) RCH2OO. + NO + O3 RCH2O. + 2 O2
(4b) RCH2O. + O2  RCHO + HO2
.
Sum (1-4): RCH3 + OH. + O3  RCHO + H2O + HO2
.
 Ozone loss. The threshold NO level for formation vs. loss is 5-
10 pptv near the ground and  20 pptv near the tropopause
Sinks of tropospheric ozone
Hydrocarbon and CO chemistry in the absence of NOx
(1) CH4 + OH.  CH3
. + H2O
(2) CH3
. + O2 + M  CH3OO. + M
(3) CH3OO. + HO2
.  CH3OOH + O2
(4a) CH3OOH + h (< 330nm)  CH3O. + OH.
(5) CH3O. + O2  HCHO + HO2
.
Sum (1-5): CH4 + O2  HCHO + H2O
Degradation of methane in NO-poor areas
Much of the CH3OOH is washed out (τ  week)
 no oxidation to HCHO and CO, radical sink:
Sum (1-3): CH4 + OH. + HO2
.  CH3OOH + H2O
(4b) CH3OOH + OH.  HCHO + H2O + OH.
Sum (1-4b): CH4 + OH. + HO2  HCHO + 2 H2O
 neutral with regard to ozone
(1) CO + OH.  CO2 + H.
(2) H. + O2 + M  HO2 + M
(3a) HO2
. + O3  OH. + 2 O2
Sum (1-3a): CO + O3  CO2 + O2
Degradation of CO in NO-poor areas
(3b) HO2
. + HO2
.  H2O2 + O2
Sum (1-3b): CO + OH. + HO2
.  CO2 + H2O2
(Model results MOZART2; Horowitz et al., 2003)
Distributions of NOx (pptv) @ 510 and 970 hPa, monthly mean
NOx distribution: NOx-poor ?
(Ridley, 1991)
Ozone formation
Ozone destruction
 Most regions of the planetary boundary layer are above the NOx threshold.
ppbv
Global sources
for N oxides, NOx
1990 (Tg N/a) Natural Anthropogenic
Lightning 3.0 (2-6)
Soils and vegetation 3.2 (1.9-4.5)
Agriculture (*) 2.3 (1.4-3.2)
Biomass burning 0.3 3.3 (2.1-5.5)
Fossil fuel burning 21 (20-23)
Sum 6.5 (4.2-11) 26.6 (23-32) (*) Animal and plant production, without biomass
burning
simulation (Grewe, 2008)
(1a) O3 + h (320-800 nm)  O2 + O
(-1a) O + O2  O3 net effect: none
(1b) O3 + h (310-336 nm)  O2 + O*
quantum yield
(T,) = 5-25%:
(2ba) + M  O + M k = 26×10-12 cm³/molec/s (N2)
k = 40×10-12 cm³/molec/s (O2)
net effect: none
(2bb) O* + H2O  2 OH. k = 220×10-12 cm³/molec/s
net effect: radical formation
Radical source ozone
Radical sources
E(mT)
with element E, multiplicity m = 2s +1 (singulett, dublett,
triplett, …), term symbol T (total orbital angular momentum
number L=0  S, L=1  P, L=2  D)
s = spin angular momentum quantum number
Example E = O: Ground state O: O(3P)
Excited state O*: O(1D)
(there is also another excited state O*:
O(1S) formed from O2 with  < 133 nm)
Radical distributions - temporal
Jülich, Germany, 14.7.1987
(Dorn et al., 1988)
Berlin, Germany, 20.-21.7.98
(Barnes et al., 2007)
Radical distributions – spatial: OH
Zonally and monthly averaged data
(105 cm-3; Spivakovsky et al., 2000)
(Lelieveld et al., 2002; Krol et al., 2003)
Common acronyms for hydrogen compounds:
HOx := HO. + HO2
. + H. + CH3O. + CH3OO. + HOCH2OO.
odd-H := HOx + 2 H2O2 + 2 CH3OOH + HNO2 + HNO4
(Lammel & Cape, 1996)
another powerful OH source reaction: Photolysis of nitrous acid
Observations field: night-time accumulation (Perner & Platt, 1979)
Correlated with NO2
cHNO2 max cHNO2 max/ cNO2
(µg / m3) (% mol / mol)
Urban areas 5.2 (1.5 - 21) 4.34 (1.8 - 13)
Rural areas 1.7 (0.4 - 2.5) 3.26 (1.7 - 13)
Remote 0.092 (< 0.2 - 0.8) > 20 (7)
Day-time sink: HNO2 + h ( = 300-405 nm )  OH. + NO
24 h mean
cHNO2 / cNO2
(% mol / mol)
Urban and rural areas 2.0 (1.0 – 5.0)
(Elshorbany et al., 2012)
1 10 100 1000
NOx [ppbv]
HNO2: Significance
Efficient radical source, in particular in winter and in high latitudes
(before visible sunrise in the UV)
dcOH/dt [*106/cm³/s] =
• from HNO2 peak 34±6 (09:00h)
from ozone 23±4 (12:00h)
in urban California (Winer & Biermann, 1994)
• Integral (06:00-18:00 h) from HNO2 58±9,
HCHO 33±8, O3 38±6
in urban Italy (Acker et al., 2006)
Toxicology:
• Precursor of carcinogenic substances, i.e. nitrosamines, RNHNO
HO2: Pernitric acid HOONO2 – relevant in the upper troposphere
(1) NO2 + HO2
. + M ⇌ HNO4 + M
(2) 2 HO2
. ⇌ H2O2 + O2
K1 0 = 1.8×10-31 [M] (T/300)-3.2 cm³ molec -1 s-1
K1 ∞ = 4.7×10-12 (T/300)-1.4 cm³ molec -1 s-1 ; Fc = 0.6
K-1 0 = 4.1×10-5 [M] e-10650/T s-1
K-1 ∞ = 4.8×1015 [M] e-11170/T s-1 ; Fc = 0.6
(kinetic data taken from Atkinson et al., 2004)
3. between 2nd and 3rd order: In the range of pressures and temperatures
in the upper troposphere and stratosphere (cM small) many reactions
are in the region between 2nd and 3rd order reactions:
k0, k∞ = low, high pressure limiting rate constant (Troe, 1979)
FC = broadening factor of falloff curve between 2nd and 3rd order in the range of
atmospheric pressures
NO2 + HO2 + M  HO2NO2 + M
K0 = k0(T)cM, n = -3.2
K∞ = k∞(T), m = -1.4 
k288 K/1000 hPa = 1.9×10-12 cm³/molec/s
k288 K/500 hPa = 1.5×10-12 cm³/molec/s
k230 K/500 hPa = 2.3×10-12 cm³/molec/s
k230 K/1000 hPa = 3.0×10-12 cm³/molec/s
1.3 Combinations of thermic reactions - Quasi steady state approximationRe-call:
HO2: Pernitric acid HOONO2 – relevant in the upper troposphere
Examples T p h τ µ(NO2) µ(HO2) µ(HNO4)
(K) (hPa) (km) (HNO4) (nmol/mol)(pmol/mol)(pmol/mol)
230 300 9 6 d 0.01 1 54
0.05 1 270
270 700 3 9 min 0.02 0.4 0.08
0.2 0.4 0.8
273 1000 0 5 min 1 3 25
30 3 750
293 1000 0 25 s 1 3 1.7
30 10 165
 high solubility (acid): changes NOx phase distribution and residence time
 NOx reservoir compound
(1) NO2 + HO2
. + M ⇌ HNO4 + M
(2) 2 HO2
. ⇌ H2O2 + O2
k1 0 = 1.8×10-31 [M] (T/300)-3.2 cm³ molec -1 s-1
k1 ∞ = 4.7×10-12 (T/300)-1.4 cm³ molec -1 s-1 ; Fc = 0.6
k-1 0 = 4.1×10-5 [M] e-10650/T s-1
k-1 ∞ = 4.8×1015 [M] e-11170/T s-1 ; Fc = 0.6
(kinetic data taken from Atkinson et al., 2004)
(1) CH3C(O)CH3 + OH.  CH3C(O)CH2
. + H2O
CH3C(O)CH2
. + O2 + M  CH3C(O)CH2OO. + M
(2) CH3CH(O)CH2OO. + NO  CH3C(O)CH2O. + NO2
(3) CH3C(O)CH2O.  HCHO + CH3C.(O)
(4) CH3C.(O) + O2  CH3C(O)OO.
(5) CH3C(O)OO. + NO2  CH3C(O)OONO2
Formation of peroxyacetyl nitrate
Peroxyacetyl nitrate (PAN): thermally labile
 NOx reservoir compound ↔ changes NOx effective residence time
 delayed ozone formation
+T
CH3C(O)OONO2 ⇌ CH3C(O)OO . + NO2
-T
k = 1.95×1016 e-13543/T s-1;
τ = 2 years @ 243K
τ = 2 days @ 273K
τ = 45 min @ 298 K
PAN sink reactions:
(-6) thermal decomposition  CH3C(O)OO. + NO2
(7a) CH3C(O)OO. + NO  CH3
. +CO2 + NO2
(7b) + HO2  CH3C(O)OOH +O2
(8b) CH3C(O)OOH + h  CH3
. +CO2 + OH
Peroxyacetyl nitrate - occurrence
1) in the cold
2) in photochemical smog and in plumes of urban areas
Rural Scotland (McFadyen & Cape, 2005)
Distributions - temporal
3) ... even in plumes of whole regions:
[ppbv]
(1) NO2 + O3  NO3
. + O2
k1 = 3.2×10-17 cm³ molec -1 s-1
-T
(2, -2) NO3
. + NO2 + M ⇌ N2O5 + M
+T
k2 0 = 3.6×10-30 [M] (T/300)-4.1 cm³ molec -1 s-1
k2 ∞ = 1.9×10-12 (T/300)-0.2 cm³ molec -1 s-1 ; Fc = 0.35
k-2 0 = 1.3×10-3 [M] (T/300)-3.5 e-11000/T s-1
k-2 ∞ = 6.2×1014 (T/300)-0.2 e-11000/T s-1 ; Fc = 0.35
(kinetic data taken from Atkinson et al., 2004)
Equilibrium: cN2O5 / (cNO2 cNO3) = K = k2 / k-2
 N2O5 not measurable, but can be inferred from cNO2 and cNO3
NO2 oxidation by ozone: NO3 formation
The adduct N2O5 – a thermically labile NO3 reservoir
Nitrate radical
[h]
(Platt et al., 1980)
[pptv]
Nitrate radical - sinks
in clouds, on wet particles:
(2b) NO3
. ⇌ NO3
.
(aqu)
NO3
.
(aqu) + 2 H2O(l)  NO3
(aqu)
+ OH.
(aqu) + H3O+
(aqu)
(3) N2O5 + 2 H2O(l)  2 NO3
(aqu)
+ 2 H3O+
(aqu)
Examples r.h. µNO2 µNO3 τNO3
het τN2O5
het Yield Production
OH dcOH (aqu)/dt
(%) (nmol (nmol (s) (s) (%) (molec cm-3 s-1)
/mol) /mol)
Marit. aerosol 60 1 0.042 1960 1890 8.0 9.1x104
90 1 0.020 840 320 1.9 4.3x104
Continental aer. 60 10 0.90 5800 4200 3.3 32 x104
90 10 0.70 4500 3800 1.4 12 x104
Coastal fog > 96 1 1.8x10-5 24 6.3 0.8 1.9x104
Radiation fog > 96 10 1.1x10-5 14 3.7 0.13 2.0x104
Day-time: Rapid photolysis of NO3
(1a) CH3CH2CH2CH3 + NO3
.  CH3CH.CH2CH3 + HNO3
further similar to chemistry upon OH attack:
(2a) CH3CH.CH2CH3 + O2 + M  CH3CH(OO.)CH2CH3 + M
(3a) CH3CH(OO.)CH2CH3 + NO  CH3CH(O.)CH2CH3 + NO2
(4aa) CH3CH(O.)CH2CH3 + O2  CH3C(O)CH2CH3 + HO2
(4ab) decomposition:  CH3CO + .CH2CH3
NO3 chemistry: alkanes, example butane
Degradation of hydrocarbons during the night
(Becker, 1991)
Preferential is the addition to the higher substituted radical (like OH reaction):
(1) CH3CH=CH2 + NO3
.  CH3CH(ONO2)C.H2
k = 9.5×10-15cm³/molec/s
diferent from chemistry upon OH attack, as adducts differ:
(2a) CH3CH(ONO2)C.H2  + NO2
O
methylepoxide
(2b) CH3CH(ONO2)C.H2 + O2  CH3CH(ONO2)C(OO.)H2
NO3 chemistry: alkenes, example propene
NO3 reactivity towards hydrocarbons: overview
HCx kNO3
10-15cm3 molec-1s-1
Alkanes:
CH4 < 0.001
CH3CH3 0.0014
CH3CH2CH3 0.017
CH3CH2CH2CH3 0.046
CH3CH(CH3)2 0.11
CH3(CH2)3CH3 0.087
Alkenes:
CH2=CH2 0.21
CH2=CHCH3 9.5
CH2=CHCH2CH3 14
cis-CH3CH=CHCH3 350
trans-CH3CH=CHCH3 390
CH2=C(CH3)2 330
-pinene 590
-pinene 210