Fundamentals of air Pollution Acid Precipitation

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Fundamentals of air Pollution Acid Precipitation

• Yaacov Mamane
• Visiting Scientist
• NCR, Rome
• Dec 2006 - May 2007
• CNR, Monterotondo, Italy

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Sandstone portal Figure on Herten Castle in Ruhr
district of Germany, Sculpted 1702.
photographed in 1969
photographed in 1908
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F U M I F U G I U M or The Inconveniencie of
theAER AND SMOAK of LONDON DISSIPATED. By John
Evelyn, 1661
It is this horrid Smoake which obscures our
Churches, and makes our Palaces look old, which
fouls our Clothes, and corrupts the waters, so as
the very Rain, and refreshing Dews which fall in
the several Seasons, precipitate this impure
vapour, which, with its black and tenacious
quality, spots and contaminates whatsoever is
expos'd to it It is this which scatters and
strews about those black and smutty Atomes upon
all things where it comes, insinuating it self
into our very secret Cabinets, and most precious
Repositories I propose therefore, that by an Act
of this present Parliament, this infernal
Nuisance be reformed enjoyning, that all those
Works be removed five or six miles distant from
London below the River of Thames I say, five or
six miles, or at the least so far as to stand
behind that Promontory jetting out, and securing
Greenwich from the pestilent Aer of Plumstead-
Marshes because, being placed at any lesser
Interval beneath the City, it would not only
prodigiously infect that his Majesties Royal Seat
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pH Levels in the USA, 1999
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• History
• As early as 1852, R. A. Smith analyzed rain that
  near the industrial city of Manchester, England
  and found that urban aerosol particles tend to be
  composed primarily of sulfuric acid, but as the
  air is transported away from sources over more
  rural areas, the acid is neutralized by
  absorption of ammonia.
• urban ? suburban ? rural
• H2SO4 NH3 ? (NH4)HSO4 (NH3) ? (NH4)2SO4
• sulfuric acid ? ammonium bisulfate ? ammonium
  sulfate
• Throughout the early part of the twentieth
  century, European scientists documented the
  sources and effects of atmospheric acids. It was
  not until 1958 that acidity of precipitation in
  the US was characterized (Junge and Werby, 1958)

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• Effects - Soils
• Soils have colloidal molecules (clay particles)
  that have a layer of negative charge. They hold
  positively charged cations such as Al³?, K?,
  Mg²?, and Ca²?.
• K?, Mg²?, Ca²? are essential plant nutrients
  while Al³? is toxic.
• Hydrogen ions from acid deposition replace these
  cations on the outer layer of colloidal
  molecules. The metal ions are then dissolved and
  leached into solution and can be washed away from
  the soil and into surface or ground water.
• Soil fertilitiy is reduced and aluminum ions can
  replace calcium in the fishs gills.
• The impact of acids on soil fertility depends on
  the structure and composition of the clays in the
  soil. The surface of the US Midwest is
  predominantly limestone (CaCO3), and lakes and
  streams have high neutralizing capacity. In the
  East granite dominates soils and surface waters
  lacking buffering capacity, are highly sensitive
  to acidification.

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• Forests can be especially sensitive to nutrient
  loss. In Europe in 1993 about a quarter of the
  trees have died or are more than 25 defoliated.
  This forest death has been attributed at, least
  in part, to environmental degradation from a
  combination of acid deposition, ground-level
  ozone, and excess nutrification, primarily
  nitrogen. In the US, loss of forests has been so
  dramatic, although several species including ash
  and oak are sensitive to acidification of soils.
• Lakes and Streams
• The sensitivity surface waters depends critically
  on their neutralizing or buffering capacity.
  Alkaline materials such as CaCO3, and MgCO3 can
  neutralize acids.

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• Materials
• The Taj Mahal, the Parthenon, the Madonna in
  Herten, Germany, and the Lincoln Memorial are
  made of marble.
• Marble, a particular crystalline form of calcite
  (CaCO3), and sandstone, are subject to attack by
  sulfuric acid.
• CaSO4 is gypsum, which is 100 times more soluble
  than CaCO3. Many priceless historic structures
  have been lost to acid deposition.
• On a more pragmatic note, the rate of corrosion
  of galvanized (zinc coated) steel is 0.62 um/yr
  in the Adirondacks, 1.01 in Washington, DC, and
  1.47 in Stubenville, OH.

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• Origins
• Primarily power generation and ore smelting.
• For example nickel is mined as nickel sulfide,
  NiS. In smelting, it is heated in air (Sudbury,
  Canada).
• The molecular weight of nickel is 57 g/mole, so
  smelting produces more than a ton of SO2 for each
  ton of nickel produced.
• Formation and Composition
• Gas Phase production of nitric acid
• OH NO2 M ? HNO3 M
• Aqueous phase production of nitric acid
• NO2 O3 ? NO3 O2
• NO3 NO2 M N2O5 M
• N2O5 H2O(l) ? 2HNO3(aq)

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• This process is important only at night, and when
  air temperatures are low because the formation of
  N2O5 is reversible, and the equilibrium
  coefficient is highly temperature dependent.
  Also, NO3 is rapidly photolyzed by visible
  radiation.
• NO3 hv ? NO2 O
• Gas Phase production of sulfuric acid
• OH SO2 M ? HOSO2 M
• Aqueous phase production of sulfuric acid
• SO2 H2O2 ? H2SO4

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Global Sulfur Budget(Flux Terms In Tg S Yr-1)
SO42- t 3.9d
SO2 t 1.3d
cloud
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OH
H2SO4(g)
8
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4
OH
NO3
(CH3)2S
64
dep 6 dry 44 wet
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dep 27 dry 20 wet
(DMS) t 1.0d
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Phytoplankton
Volcanoes
Combustion Smelters
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Global Sulfur Emission To The Atmosphere1990
annual mean
Chin et al. 2000
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Trends In Sulfate And Nitrate Wet Deposition
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Rain Chemistry in the East MediterraneanTrends
of H and SO4 , meq/l
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Rain Chemistry in the East MediterraneanTrends
of H and SO4 , meq/l
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Rain Chemistry in the East MediterraneanAnions
and Cations , meq/l
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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.
• GAS HENRYS LAW CONSTANT
• (M / atm at 298 K)
• CO2 3.1 x 10?²
• SO2 1.3
• HNO3 2.1 x 10?5
• H2O2 9.7 x 10?4

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Use of Henrys Law
Assume that the atmosphere contains only N2, O2,
and CO2 and that rain is in equilibrium with
CO2. CO2 form a weak acid H2CO3, and it is in
equilibrium with it. We should remember that H2O
H? OH? H?OH? 1 x 10?14 pH -log
H? In pure H2O, pH 7.0 We can assume that
CO2 in the atmosphere is around 350 ppm. ca.
370 ppm
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10-6x350

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K1C 4.3?10-7 mole/l
K2C 4.7?10-11 mole/l KHC
0.034 M/atm
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pH 5.6
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(CO2) Total (H2CO3) (HCO3-) (CO3)
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K1S 1.3?10-2 M K2S 6.6?10-8 M
KHS 1.23 M/atm
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What would be the pH of pure rain water in Rome
today? Assume that the atmosphere contains only
N2, O2, and CO2 and that rain is in equilibrium
with CO2. Remember H2O H? OH? H?OH? 1
x 10?14 pH -log H? In pure H2O, pH
7.0 We can assume that CO2 ca. 370 ppm
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• Todays barometric pressure is 993 hPa 993/1013
  atm 0.98 atm. Thus the partial pressure of CO2
  is
• In water CO2 reacts slightly, but H2CO3 remains
  constant as long as the partial pressure of CO2
  remains constant.

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• We know that
• and
• Thus
• H 2.3x10-6 ? pH -log(2.3x10-6) 5.6
• 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.

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• 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 will not 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?

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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?
• The pressure decreases as a function of height.
  At 5km the ambient pressure is around half the
  atmospheric pressure 0.54 atm.
• This SO2 will not stay as SO2H2O, but
  participate in a aqueous phase reaction, that is
  it will dissociate.

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• The concentration of SO2H2O, however, remains
  constant because more SO2 is entrained as SO2H2O
  dissociates. The extent of dissociation depends
  on H? and thus pH, but the concentration of
  SO2H2O will stay constant as long as the gaseous
  SO2 concentration stays constant. Whats the pH
  for our mixture?
• If most of the H? comes from SO2H2O
  dissociation, then
• Note that there about 400 times as much S in the
  form of HOSO2? as in the form H2OSO2. 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

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• Because SO2 participates in aqueous-phase
  reactions, Eq. (I) above will give the correct
  H2OSO2, 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 H2OSO4, 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 SO2H2O is oxidized to
  H2OSO4, 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.

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