Transcript Fundamentals of air Pollution – Acid Precipitation

Fundamentals of air Pollution
– Acid Precipitation
Yaacov Mamane
Visiting Scientist
NCR, Rome
Dec 2006 - May 2007
CNR, Monterotondo, Italy
photographed in 1969
Sandstone portal
Figure on Herten
Castle in Ruhr district
of Germany,
Sculpted 1702.
photographed in 1908
F U M I F U G I U M: or The Inconveniencie of the
AER 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
pH Levels in the USA, 1999
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)
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 colloidalNH
molecules.
The metal ions are then dissolved and
4
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
fish’s 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.
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.
2H  Ca 2  CO32  CO2  Ca 2
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.
CaCO3  H2SO 4  CaSO 4  CO2 
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.
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).
NiS  O2  Ni  SO 2 
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)
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
Global Sulfur Budget
(Flux Terms In Tg S Yr-1)
cloud
42
SO2
4
NO3
18
t = 3.9d
OH
t = 1.3d 8
SO42-
H2SO4(g)
OH
(CH3)2S
(DMS)
t = 1.0d
10
64
dep
27 dry
20 wet
22
Phytoplankton
Volcanoes
Combustion
Smelters
dep
6 dry
44 wet
Global Sulfur Emission To The Atmosphere
1990 annual mean
Chin et al. [2000]
Trends In Sulfate And Nitrate Wet Deposition
Rain Chemistry in the East Mediterranean
Trends of H+ and SO4= , meq/l
Line Plot (Sheet1 in Imported acid rain.stw 21v*182c)
100
1000
80
800
60
600
40
400
20
200
0
0
-20
-200
H+(L)
SO4=(exc.)(R)
18-1
17-1
23-2
16-1
30-3
05-0
15-1
23-2
04-0
14-1
06-0
04-0
10-1
17-1
Rain Chemistry in the East Mediterranean
Trends of H+ and SO4= , meq/l
Scatterplot (Sheet1 in Imported acid rain.stw 21v*182c)
H+ = 8.6346-0.0203*x
70
60
50
40
30
H+
20
10
0
-10
-20
-200
0
200
400
SO4=(exc.)
600
800
1000
Rain Chemistry in the East Mediterranean
Anions and Cations , meq/l
Scatterplot (Air Poll Course\acid rain Israel.xls 21v*182c)
Cations = 44.92+0.94*x, meq/l
2200
2000
1800
1600
1400
1200
Cations
1000
800
600
400
200
0
0
200
400
600
800
1000
1200
Anions
1400
1600
1800
2000
2200
AQUEOUS-PHASE CHEMISTRY
HENRY’S 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
CO₂
SO₂
HNO₃
H₂O₂
HENRY’S LAW CONSTANT
(M / atm at 298 K)
3.1 x 10⁻²
1.3
2.1 x 10⁺⁵
9.7 x 10⁺⁴
Use of Henry’s 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
CO 2( g ) + H 2O 
 H 2CO 3
PCO 2 =
350x10-6
K HC =

H 2CO 3 
 H + HCO 3
 
=
HCO 
3  H + CO 3
+
H 2O 
H
+
OH

  
K w = H+
OH -
pH  - log 10 H + 
= 10 -14
 H 2CO 3 
PCO 2
H +   HCO 3- 

K 1C =
H 2CO 3 




K 2C =
HCO
 3
H+
CO =
3
 H  =  OH-  +  HCO-3 + 2 CO=3 
   
K
K
K
PCO
K
K
K
PCO
2  2 HC 1C 2C
2
H  = W  HC 1C
2
H
H+
+
H
O = H
 
  K
 3
W
 
 K HC K 1C PC O  H +  
2
 2 K HC K 1C K 2C PC O
2
K1C = 4.310-7 mole/l
K2C = 4.710-11 mole/l
KHC = 0.034 M/atm
 
2

H
 K W + K HC K 1C PCO 2
= 10 -14 M 2 + 4.3  10 -7 M  0.034
~ 5.1  10 -12 M 2
 
H + = 2.3  10 -6 M
pH = 6 - 0.36  5.6
pH = 5.6
M
 350  10  6 Atm
Atm
(CO2) Total = (H2CO3) + (HCO3-) + (CO3=)
CO 
2 Total

= K HC PC O +
2
K 1C K HC PC O
H 
+
2

K 2C K 1C K HC PC O
H 
2
+ 2
 CO 2  Total = PCO 2

A + B  10pH + C  100pH

SO 2    H 2O
  H 2SO 3  = K HS PSO 2
 H 2SO 3 

HSO 3

   

=

SO

H

3 

 HSO   H 

3
K 1S =
K 2S =
K1S = 1.310-2 M
K2S = 6.610-8 M
KHS = 1.23 M/atm
H +  HSO-3 
 H 2SO 3 
H +  SO =3 
HSO-3 
 
dSO 4= 
SO =3
+
 
1
O 2  SO 4=
2
 
= t + C
=
SO
k
=
SO
 4 t 3  3 
dt
 SO 2  T
= k 3 SO =3

   + SO4= 
=  H 2SO 3  + HSO -3 + SO =
3
= K HS PSO 2 +
K 1S K HS PSO 2
H 
+
+
+
K 2S K 1S K HS PSO 2
H 
 2
+
t K 2S K 1S K HS PSO
2
+ K3 
dt
2
to
H
 
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
Today’s barometric pressure is 993 hPa = 993/1013 atm = 0.98 atm. Thus
the partial pressure of CO₂ is
PCO2 = 370106 (0.98) = 3.63104 atm
[CO2 ]aq = H  P(CO2 ) = 3.4  102  3.62 104
= 1.23 105 M
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 2CO3
 H 2CO3 = H   HCO3
[H ][HCO3 ]
= Keq = 4.3 107
[H2CO3 ]
We know that:
[H2CO3 ] = 1.23 105 M
and
[H ] = [HCO3 ]
Thus
[H ] = Ka * H2CO3
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.
pH = log[H ]
[H ] = 10 4.7 = 2 105
PHNO 3 = [HNO3 ]aq /H
 = 2 105 /2.1105
 = 9.01011 at m
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?
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
The pressure decreases as a function of height. At 5km the ambient
pressure is around half the atmospheric pressure: 0.54 atm.
PSO 2 = 100109 0.54 = 5.4108 at m
[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.
SO 2  H2O  H  HOSO2
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 =
[H2O  SO 2 ]
If most of the [H⁺] comes from SO₂•H₂O dissociation, then
[H ] = [HOSO2 ]
[H ] = K a [H2 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
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.