LECTURE 1/2
AOSC 637 Spring 2011
Atmospheric Chemistry
Russell R. Dickerson
Copyright © R. R. Dickerson 2011
1
Topics to be covered this semester
Organic and biochemistry for physicists and meteorologists.
Laboratory techniques for detection and properties of aerosols.
Weak acids and bases.
Basic chemical thermodynamics
Experimental design.
Spectroscopy of polyatomic molecules and photochemistry.
Kinetics theory and lab techniques.
Biogeochemical cycles of Ox, NOx, SOx, HOx, CH4 and halogens.
Measurements of cloud properties.
Cloud microphysics
Dry Deposition and micrometeorology
Unanswered questions on the formation and properties of aerosols:
SOA
SO2 oxidation
Absorption and mixing
Copyright © R. R. Dickerson 2011
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AAAS Annual Meeting
17-20 February 2011
Washington, DC
When Pollution Gets Personal: Ethics of Reporting on Human
Exposures
Sunday, February 20, 2011: 8:00 AM-9:30 AM
147A (Washington Convention Center )
New technologies measure ever-lower concentrations of environmental chemicals in human blood, urine, breast milk,
and other tissues and in personal spaces such as homes. The earliest exposure measurements for emerging contaminants
precede understanding of health effects, sources, and toxicity, since chemicals are put into use without safety testing or
product labeling. As a result, exposure studies raise ethical questions about whether and how to report individual results
to participants. Medical ethics traditionally favored reporting only clinically significant results, to protect patients from
distress. However, community-based participatory research frameworks and activist biomonitoring projects focus on the
human research ethics principles of respect for autonomy and the responsibility to maximize benefits, which support
study participants’ right-to-know. This symposium will consider ethical and legal issues in reporting individual exposures
in academic, government, and activist studies. It will draw on interviews with participants from diverse socioeconomic
settings who received results for many endocrine disrupting compounds (EDCs), including compounds recently in the
news. Participants dramatically expanded their conceptions of pollution as they learned about consumer-product
chemicals in their homes and bodies. Learning personal results motivated action to reduce exposures through both
lifestyle changes and policy advocacy. Organizer:
Julia G. Brody, Silent Spring Institute
Moderator:
Julia G. Brody, Silent Spring Institute
Discussant:
Sharyle Patton, Commonweal
Speakers:
Gwen Collman, National Institute of Environmental Health Sciences
New Expectations for Individuals' Right-To-Know in Environmental Health Research
Rachel Morello-Frosch, University of California
Reporting Cord Blood Contaminants to Mothers in the California Biomonitoring Program
Shaun Goho, Harvard Law School
Is It Safe? Legal Requirements to Disclose Measurements in Homes
Aeroecology: Transcending Boundaries Among Ecology, Meteorology, and
Physics
Saturday, February 19, 2011: 3:00 PM-4:30 PM
102B (Washington Convention Center )
Birds, bats, and flying insects depend on aerial habitats for critical aspects of their life history, such as foraging, migration, and
dispersal. In marine biology, the importance of interactions between aqueous habitats and marine organisms has long been
recognized. In contrast, the ecological significance of the aerosphere for aerial organisms has received little attention. Aeroecology is
a new discipline whose unifying concept is a focus on the aerosphere and the myriad organisms that inhabit and depend on this aerial
environment for their existence. Volant animals contribute to the ecological integrity of multiple ecosystems that span geopolitical
boundaries linked by migration or dispersal through the aerosphere. Investigating behavior and movements of animals in the
aerosphere presents formidable challenges, requiring creative integration of novel technological advances for data acquisition and
analysis. In this symposium, we bring together a diverse group of physicists, meteorologists, radar scientists, and ecologists to address
major challenges and advances in aeroecological research. Through its transdisciplinary approach and emphases on biotic-abiotic
interactions at multiple spatial and temporal scales, aeroecology promises to advance understanding of the effects of climate change
and anthropogenic alteration of diverse landscapes on biodiversity, global health, and ecological integrity. Organizer:
Winifred F. Frick, University of California
Co-Organizer:
Phillip B. Chilson, University of Oklahoma
Moderator:
Thomas H. Kunz, Boston University
Discussants:
Jeffrey F. Kelly, University of Oklahoma
and Kenneth W. Howard, National Oceanic and Atmospheric Administration National Severe Storms Laboratory
Speakers:
Thomas H. Kunz, Boston University
Aeroecology as an Emerging Scientific Discipline
Phillip B. Chilson, University of Oklahoma
Enabling Aeroecological Studies Through Advancements in Radar Technology
Winifred F. Frick, University of California
Meteorological Drivers of Predator-Prey Interactions in the Aerosphere
Limiting Climate Change: Reducing Black Carbon and Tropospheric
Ozone Precursors
Sunday, February 20, 2011: 1:30 PM-4:30 PM
101 (Washington Convention Center )
This session brings together eminent scientists and pressure groups from New York to New Delhi to examine available
options that can allow society to deal with the urgent task of mitigating climate change, while continuing to improve living
conditions. Speakers will evidence that although reductions in carbon dioxide emissions and other long-lived greenhouse
gases are essential for mitigation of long-term climate change, real leverage over near-term climate comes primarily from
the tropospheric ozone precursors methane, carbon monoxide, and volatile organic compounds and, at least regionally, from
black carbon aerosols. Case studies will show how these short-lived pollutants contribute to both global warming and many
of the most alarming regional climatic changes, including the melting of Himalayan glaciers and Arctic sea ice and shifts in
regional precipitation. They also contribute to air pollution, with adverse effects on human health, agricultural yields, and
other social-economic costs. Speakers will argue for a new and more integrated approach, in which emission reductions in
specific sectors are optimized to get benefits for air quality and climate. Presentations will show how this can be more
robust and more cost-effective than policies that aim at air quality and climate targets separately. This session will report on
recent studies in support of an assessment by the United Nations Environment Program (UNEP) of the above-mentioned
options. Organizer:
Frank Raes, European Commission, JRC Institute for Environment and Sustainability
Co-Organizer:
Geraldine Barry, European Commission, JRC
Speakers:
Drew Shindell, NASA Goddard Institute for Space Studies
Limiting Near-Term Climate Change While Improving Human Well-Being
Teppei Yasunari, NASA Goddard Earth Sciences and Technology Center
Impacts of Black Carbon (BC) Pollution on Himalayan Glaciers
Markus Amann, International Institute for Applied Systems Analysis
Win-Win and Win-No-Lose Control Measures for Black Carbon and Ozone
Frank Raes, European Commission, JRC Institute for Environment and Sustainability
Benefits of BC and Tropospheric Ozone Reduction Measures for Climate, Health, and Ecosystems
Erika Rosenthal, Earthjustice
Good Practice in Reducing Black Carbon Emissions at the Local Level
Martin Williams, King's College
Global and Local Responses to the Nitrogen Challenge: Science, Practice,
and Policy
Saturday, February 19, 2011: 1:30 PM-4:30 PM
140A (Washington Convention Center )
Managing the benefits and costs associated with reactive nitrogen is one of the most dynamic challenges society faces today. On
one hand, reactive nitrogen is an essential resource. Its use in agriculture has undoubtedly increased the availability of food, fiber,
and feed and contributed to major advances in human well-being over the last half century. On the other hand, human additions of
reactive nitrogen to the environment now dominate the global nitrogen cycle. The increased quantity of reactive nitrogen in the air,
water, and soil is causing significant changes to the environment (climate change, air and water pollution, and biodiversity loss)
and threatens the health and vitality of human and natural resources worldwide. But while the challenges reactive nitrogen presents
to society are clear, the most appropriate responses are less apparent. Solutions will need to balance trade-offs among ecosystem
services and the competing objectives of key actors. This symposium brings together biologists, biogeochemists, and economists to
discuss the frontiers of knowledge and solutions on this pressing issue. Speakers will report on the results of domestic and
international assessments that link nitrogen science, practice, and policy at a variety of locations and scales. Organizer:
Todd S. Rosenstock, University of California
Co-Organizer:
Thomas P. Tomich, University of California
Moderator:
Todd S. Rosenstock, University of California
Discussant:
Eric A. Davidson, Woods Hole Research Center
Speakers:
Walter V. Reid, David and Lucile Packard Foundation
Bridging the Nitrogen Science and Policy Divide
Cheryl A. Palm, Earth Institute
Nitrogen, Development, and Sustainability: Trade-Offs Between Too Little and Too Much
Alan R. Townsend, University of Colorado
Catch 22: The Nitrogen Cycle and Human Welfare
Thomas P. Tomich, University of California
A Framework for Action: Lessons from the California Nitrogen Assessment
Cliff Snyder, International Plant Nutrition Institute
Nitrogen Stewardship: Balancing Crop Production Management and Environmental Protection
Not covered (awaiting results of
diagnostic exam.)
Basic thermodynamics of dry and wet air.
Parcel theory and stability.
General circulation and synoptic circulations.
Stefan-Boltzmann and Wien.
Basic ozone photochemistry.
Biogeochemical cycle of C.
Convection and chemistry
Unit conversions
Copyright © R. R. Dickerson 2011
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TODAY’S OUTLINE
Ia. Chemistry (Concentration Units):
1. Gas-phase
2. Aqueous-phase
Ib. Atmospheric Physics
1. Pressure
2. Atmospheric structure and circulation
A. pressure and temp. profiles
B. thermo diagram and stability
C. circulation (winds)
Copyright © R. R. Dickerson 2011
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Ia. 1. GAS-PHASE
Atoms
He Ar
monatomic
N₂ O₂
Molecules
CO₂ O₃
Radicals
H₂CO CCl₂F₂ OH HO₂
diatomic
triatomic
polyatomic
UNITS OF CONCENTRATION
Mole Fraction – for ideal gas this is the same as volume fraction. Also
called mixing ratio, or volume mixing ratio.
fraction
[O₂] = 1/5
percent
[Ar] = 1%
[H₂O] = up to 4%
parts per million (10⁶)
[CH₄] = 1.7 ppm
parts per billion (10⁹)
[O₃] = 30 ppb
parts per trillion (10¹²)
[CCl₂F₂] = 100 ppt
Copyright © R. R. Dickerson 2011
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ATMOSPHERIC CO2 INCREASE OVER PAST 1000 YEARS
Jacob: Intergovernmental Panel on Climate Change (IPCC) document, 2001
Concentration units: parts per million (ppm)
number of CO2 molecules per 106 molecules of air
CO2 CONCENTRATION IS MEASURED HERE AS MIXING RATIO
Copyright © R. R. Dickerson 2011
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For an ideal gas these concentrations are constant regardless temperature
and pressure.
Ideal Gas Law: PV = nRT
For example if T₂ = 2T₁ and if dP = 0 then V₂ = 2V₁.
Meteorologists favor the ideal gas law for a kg of air:
pα = R’T
Where R’ has units of J kg-1K-1 and α is the specific volume (volume
occupied by 1 kg of air; Mwt 29 g/mole).
If air in New York is brought to Denver (P = 83% atm) there will be no
change in the concentration of pollutants as long as the concentration is
expressed as a volume (moler) mixing ratio.
MASS PER UNIT VOLUME
Best for particles (solid or liquid)
Weigh a filter – suck 1.00 m⁻³ air through it – reweigh it
Change in weight is conc “dust” in mass/unit volume or μgm⁻³
Copyright © R. R. Dickerson 2011
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EXAMPLE
If you find 10 μg/m³ “dust” of which 2 μg/m³ are nitrate (NO₃‾), how
much gas phase HNO₃, expressed as a mixing ratio, was there in the air
assuming that all the nitrate was in the form of nitric acid? We must
convert 2.0 μg/m³ HNO₃ to ppb:
2.0 10 6 /63
1000/22.4
gm 3 / g / mole
L / m 3 / L / mole
Remember one mole of an ideal gas is 22.4 liters at STP = 0 C & 1.0 atm.
2.0 μg/m³ HNO₃ = 7.1 x 10⁻¹º = 0.71 ppb
In general: 1.0 μg/m³ HNO₃ = 0.35 ppb
Notice that the concentration in μg/m³ changes with P and T of the air.
Copyright © R. R. Dickerson 2011
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Mixing ratios area also good for writing reactions:
NO + O₃ → NO₂ + O₂
1 ppm + 1 ppm = 1 ppm + 1 ppm
Note: the [O₂] in air is not 1 ppm, rather it is 0.2 x 10⁶ + 1.0 ppm.
Above is an example of an irreversible reaction. There are also
reversible reactions.
EXAMPLE
Equilibrium of ammonium nitrate
NH₃ + HNO₃ ↔ NH4 NO3(S)
[NH 4 NO3 ]
Keq 
[NH3 ][HNO 3 ]
Ammonium nitrate is a solid, and thus has a concentration defined
as unity.
Copyright © R. R. Dickerson 2011
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Number density nX [molecules
-3
cm ]
Proper measure for
• calculation of reaction rates
• optical properties of atmosphere
# molecules of X
nX 
unit volume of air

Column concentration =  n( z )dz
0
Proper measure for absorption of radiation by
atmosphere
Column concentrations are measured in molecules cm-2 , atm*cm, and
Dobson Units, DU.
1 atm*cm = 1000 DU = 2.69x1019 cm-2.
Copyright © R. R. Dickerson 2011
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STRATOSPHERIC OZONE LAYER (Jacob’s book)
1 “Dobson Unit (DU)” = 0.01 mm ozone at STP = 2.69x1016 molecules cm-2
THICKNESS OF OZONE LAYER IS MEASURED AS A COLUMN CONCENTRATION
Copyright © R. R. Dickerson 2011
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AQUEOUS-PHASE CHEMISTRY
HENRYS LAW
The mass of a gas that dissolves in a given amount of liquid as a given
temperature is directly proportional to the partial pressure of the gas above
the liquid. This law does not apply to gases that react with the liquid or
ionized in the liquid. See Finlayson p.151 or Chameides, J. Geophys. Res.,
4739, 1984. Check out also
www.mpch-mainz.mpg.de/~sander/res/henry.html
Copyright © R. R. Dickerson 2011
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HENRY’S LAW CONSTANT
(M /atm at 298 K)
GAS
OXYGEN
O₂
OZONE
O₃
NITROGEN DIOXIDE
NO₂
CARBON DIOXIDE
CO₂
SULFUR DIOXIDE
SO₂
NITRIC ACID (effective)
HNO₃
HYDROGEN PEROXIDE
H₂O₂
HYDROPEROXY RADICAL
HO₂
ALKYL NITRATES
(RONO₂)
1.3 x 10⁻²
9.4 x 10⁻³
1.0 x 10⁻²
3.1 x 10⁻²
1.3
2.1 x 10⁺⁵
9.7 x 10⁺⁴
9.0 x 10³
1.3
Copyright © R. R. Dickerson 2011
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HENRY’S LAW EXAMPLE
What was the pH of fresh water in the preindustrial atmosphere? What
would be the pH of pure rain water in Washington, D.C. today? Assume
that the atmosphere contains only N₂, O₂, and CO₂ and that rain in
equilibrium with CO₂.
Remember:
H₂O = H⁺ + OH⁻
[H⁺][OH⁻] = 1 x 10⁻¹⁴
pH = -log[H⁺]
In pure H₂O pH = 7.00
We can measure:
[CO₂] = 280 in a preindustrial world
~ 390 ppm today
Copyright © R. R. Dickerson 2011
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Today’s barometric pressure is 993 hPa = 993/1013 atm = 0.98 atm. Thus
the partial pressure of CO₂ is
PCO 2  280 106 (0.98)  2.74 104 atm
[CO 2 ]aq  H  P(CO 2 )  3.4 10 2  2.74 10 4
 9.33 10 6 M (1.3 x10 5 M currently)
In water CO₂ reacts slightly, but [H₂CO₃] remains constant as long as the
partial pressure of CO₂ remains constant.
CO 2  H 2 O  H 2 CO3
 H 2 CO3  H   HCO 3
[H  ][HCO 3 ]
 Keq  4.3  10 7
[H 2 CO3 ]
Copyright © R. R. Dickerson 2011
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We know that:
and
THUS
[H 2 CO3 ]  9.33 10 6 M
[H  ]  [HCO 3 ]
[H  ]  K a * H 2 CO3
H+ = 2.09x10-6 → pH = -log(2.09x10-6) = 5.68
H+ = 2.5x10-6 → pH = -log(2.5x10-6) = 5.60 today.
The pH of the ocean today is ~8.1 so [H+] = 7.9x10-9.
[H+] * [HCO3-]/[H2CO3] = 7.9x10-9 * [HCO3-]/1.28x10-5 = 4.7 x10-7
[HCO3-] = 7.6x10-4 M
Most of the C in the oceans is tied up as bicarbonate.
Copyright © R. R. Dickerson 2011
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EXAMPLE 2
If fog water contains enough nitric acid (HNO₃) to have a pH of 4.7, can any
appreciable amount nitric acid vapor return to the atmosphere? Another way to ask this
question is to ask what partial pressure of HNO₃ is in equilibrium with typical “acid
rain” i.e. water at pH 4.7? We will have to assume that HNO₃ is 50% ionized.
pH  log[H  ]
[H  ]  10  4.7  2 10 5
PHNO3  [HNO 3 ]aq /H
  2 10 5 /2.1 105
  9.0  10 11 atm
This is equivalent to 90 ppt, a small amount for a polluted environment, but the actual
[HNO₃] would be even lower because nitric acid ionized in solution. In other words,
once nitric acid is in solution, it wont come back out again unless the droplet
evaporates; conversely any vapor-phase nitric acid will be quickly absorbed into the
aqueous-phase in the presence of cloud or fog water.
Which pollutants can be rained out?
Copyright © R. R. Dickerson 2011
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We want to calculate the ratio of the aqueous phase to the gas phase concentration
of a pollutant in a cloud. The units can be anything , but they must be the same.
We will assume that the gas and aqueous phases are in equilibrium. We need
the following:
Henry’s Law Coefficient: H (M/atm)
Cloud liquid water content: LWC (gm⁻³)
Total pressure: PT (atm)
Ambient temperature: T(K)
LET:
X aq be the concentration of X in the aqueous phase in moles/m³
X gas be the concentration of X in the gas phase in moles/m³
[X] aq  HPX
Where [X] aq is the aqueous concentration in M, and Px is the partial pressure
expressed in atm. We can find the partial pressure from the mixing ratio and
total pressure.
Copyright © R. R. Dickerson 2011
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PX  [X] gas PT
For the aqueous-phase concentration:
Xaq  [X]aq LWC 103
units:
moles/m³ = moles/L(water) x g(water)/m³(air) x L/g
Xaq  H[X] gas PT LWC 103
For the gaseous content:
X gas 
units:
moles/m³ =
[X] gas

T 1
 22.4 10 3

273 PT 

L(X)/L(air )
L/mole  m 3 /L

Copyright © R. R. Dickerson 2011

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X aq
X gas

1
3 T

/[X] gas
 H[X] gas PT LWC 10  22.4 10
273 PT 

3
X aq
X gas
 H  LWC 
T
 22.4  10 6
273
Notice that the ratio is independent of pressure and concentration. For a
species with a Henry’s law coefficient of 400, only about 1% will go into a
cloud with a LWC of 1 g/m³. This points out the need to consider aqueous
reactions.
Copyright © R. R. Dickerson 2011
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What is the possible pH of water in a high cloud (alt. ≃ 5km) that absorbed
sulfur while in equilibrium with 100 ppb of SO₂?
SO 2  H 2 O  SO 2  H 2 O
[SO 2 ]  100ppb
PSO 2  [SO 2 ]PTotal  [SO 2 ]P5km
In the next lecture we will show how to derive the pressure as a function of
height. At 5km the ambient pressure is 0.54 atm.
PSO 2  100 10 9 0.54  5.4 10 8 atm
[SO 2 ]aq  HPSO 2
  7 10 8 M
This SO₂ will not stay as SO₂•H₂O, but participate in a aqueous phase
reaction, that is it will dissociate.
SO2  H 2O  H   HOSO2
Copyright © R. R. Dickerson 2011
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The concentration of SO₂•H₂O, however, remains constant because more SO₂
is entrained as SO₂•H₂O dissociates. The extent of dissociation depends
on [H⁺] and thus pH, but the concentration of SO₂•H₂O will stay constant
as long as the gaseous SO₂ concentration stays constant. What’s the pH for
our mixture?


[H ][HOSO 2 ]
Ka 
[H 2O  SO 2 ]
If most of the [H⁺] comes from SO₂•H₂O dissociation, then
[H  ]  [HOSO 2 ]
[H  ]  K a [H 2 O  SO 2 ]  3  10 5
Note that there about 400 times as much S in the form of HOSO₂⁻ as in the
form H₂O•SO₂. HOSO₂⁻ is a very weak acid, ant the reaction stops here.
The pH of cloudwater in contact with 100 ppb of SO₂ will be 4.5
Copyright © R. R. Dickerson 2011
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Because SO₂ participates in aqueous-phase reactions, Eq. (I) above will
give the correct [H₂O•SO₂], but will underestimate the total sulfur in
solution. Taken together all the forms of S in this oxidation state are called
sulfur four, or S(IV).
If all the S(IV) in the cloud water turns to S(VI) (sulfate) then the hydrogen
ion concentration will approximately double because both protons come off
H₂O•SO₄, in other words HSO₄⁻ is a strong acid.
This is fairly acidic, but we started with a very high concentration of SO₂,
one that is characteristic of urban air. In more rural areas of the eastern US
an SO₂ mixing ratio of a 1-5 ppb is more common. As SO₂•H₂O is
oxidized to H₂O•SO₄, more SO₂ is drawn into the cloud water, and the
acidity continue to rise. Hydrogen peroxide is the most common oxidant
for forming sulfuric acid in solution; we will discuss H₂O₂ later.
Copyright © R. R. Dickerson 2011
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Second example - alkylnitrates
Can alkyl nitrates, R-ONO2, be removed from the atmosphere by rain (wet
deposition)? Consider the relative amount of an alkyl nitrate in the gas phase
vs. the aqueous phase in a cloud. If most of the alkyl nitrate is in the aqueous
phase, than precipitation must be important. We need the following
information:
1. Henry's Law Coef. (KH) for R-ONO2  2 M/atm at 298 K (Luke et al., 1989).
2. A thick cloud has 1.0 g liquid water per cubic meter.
3. The typical temperature of a cloud is near 0 C.
4. The typical altitude of a cloud is about 5 km thus the pressure is about 0.5 atm.
5. The most alkyl nitrate one might find in the atmosphere over a continent is
about 1.0 ppb.
First we apply Henry's law to find out what the aqueous concentration of R-ONO2
would be is the cloud is in equilibrium with the vapor phase.
[R-ONO2 ]aq = KH x [R-ONO2 ]gas x Ptotal
Where [R-ONO2 ]gas must be expressed in partial pressure, atm.
= 2.0 x 10-9 x Ptotal
= 2.0 x10-9 x 0.5
[R-ONO2 ]aq = 10-9 M
Copyright © R. R. Dickerson 2011
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How do we compare this to the gas phase concentration?
Change both values into moles/m3.
[R-ONO2 ]aq = 10-9 x 10-3 = 10-12
UNITS: moles/L(water) x L(water)/m3 (air) = moles /m3
[R-ONO2 ]gas = 10-9 x 103 x 0.5/22.4 = 2.2 x 10-8
UNITS: L (R-ONO2 )/ L (air) x L/m3 x atm/(L atm/mole) =
moles/m3
We see that the vapor phase concentration is 22,000 higher than
the aqueous phase concentration. Rainout will still matter,
however, if R-ONO2 reacts in solution and thus is removed.
This is the case for SO2 in water containing H2O2 where
H2SO4 is produced, but aqueous reactions of R-ONO2 with
species commonly found in rainwater are as yet unknown.
This implies that alkyl nitrates may have a residence time
long enough to be important in regional or global
atmospheric chemistry.
Copyright © R. R. Dickerson 2011
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For species X, a general solution to the "rain out" question is
given by an expression for the ratio of moles of gas-phase X
to moles of aqueous-phase X in a given volume of air.
Xaq /Xgas = KH x LWC x 2.24 x 10-5
Where KH is the Henry's Law coefficient in M/atm, and LWC is
the liquid water content in g/m3. This equation is valid at
273 K; to correct for temperature multiply the right side by
(T/273). The equation above shows that the ratio is
independent of pressure and concentration. For alkyl nitrates
this ratio is about 4.4 x10-5.
For a species with a Henry's law coefficient of 4 x102, about
1% will go into a cloud with a LWC of 1g/m3.
Copyright © R. R. Dickerson 2011
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AAAS Poster Session
> Location
> Sunday, 20 February 2011
> 1:00 p.m-5:00 p.m.
> Washington Convention Center
> Hall D