Pollution, Waste Disposal and the Environment
Geosphere
– the inorganic component of the environment
– called the lithosphere, consists of the upper mantle and the crust of
the Earth.
– made primarily of rock
– rock is made up of minerals which are naturally occurring inorganic
solids with a definite crystal structure and chemical composition
Atmosphere
– gaseous component of the environment which surrounds the planet
– gravity captured gases from space and from outgassing by the crust
– major components are N2 and O2
– minor components are noble gases, CO2, H2O
Hydrosphere
– hydrogen and oxygen combined to form water vapour which condensed as the crust cooled
creating
the oceans
Biosphere
– consists of all life forms and the components necessary for their formation
Le Chatelier’s Principle
– if a stress is placed on an equilibrium, it will shift to relieve the stress.
Two factors upset the traditional model at the beginning of the 19th century, the exploding human
population and more important, the Industrial revolution.
Anthrosphere
– technology and its effects
Geology is the study of the Geosphere.
One of the more important interactions of the spheres is the production of soil which is essential
for life on the terrestrial part of the Earth’s crust. It is produced by the weathering of rock caused
by physical, geochemical and biological processes.
Biology is study of the Biosphere
In the Biosphere, anything that is part of the environment of an organism is called biotic.
Everything else is abiotic.
Plant Photosynthesis
6CO2 + 6H2O + energy <==> C6H12O6 + 6O2
Light is packets of energy with wave properties. The wave properties are wavelength, λ, and
frequency, ν. They are related by the equation
c = λν
where c is the speed of light, ~3 x 108 m/sec.
Eαν
or
E α 1/λ
E = hν
Where h is Planck’s constant, 6.63 x 10-34 J-sec
Ecology is the study of the relationships within the Biosphere and between it and the other
spheres.
An ecosystem is a group of organisms and their environment which are interrelated and
dependent on each other.
A contaminant causes a change in the normal concentration of an environment but does no
damage.
A pollutant is a contaminant that causes damage.
- a source is where the pollutant originates
- a receptor is anything that is affected by the pollutant
- a sink is a long term repository of a pollutant
- can be due to a chemical reaction
- CaCO3 + H2SO4 <==> CaSO4 + H2O + CO2
A xenobiotic compound is one which is not normally part of an organism’s environment. If the
compound is detrimental to the organism, it is a poison or toxicant.
Discussed Section 1.6, Chapter 1 in text – Matter and Cycles of Matter Pg 9 -15
.
THE HYDROSPHERE
INTRODUCTION
– Why is water important to us?
Enters into every facet of our existence
Without water and its unique physical and chemical properties, organic life as we
know it would not likely be possible
1.
Important Properties of Water
– Is a polar molecule
=> good solvent due to hydration
=> high dielectric constant (80 vs. 1 for air; reduces forces between opposite
charges) makes water a good solvent for ionic compounds
– Forms hydrogen bonds with itself and other compounds in solution
=> higher boiling point than expected
=> water expands on freezing, max density at 4°C
=> large surface tension
– Unusual thermal properties:
- high heat capacity
- high latent heat of fusion and evaporation
=> good environmental thermal regulator
- good thermal conductor
=> good thermal regulator for organisms
2.
Hydrologic Cycle
–.Water moves around the planet in what we call the Hydrologic Cycle
– Water moves into the atmosphere by evaporation from the oceans and transpiration from the
land
– Water condenses in the atmosphere and releases heat while forming clouds, fog, etc.
– Precipitation carries water back to ocean and land
– Water dissolves and carries elements down into geosphere and groundwater.
– Important part of all geochemical cycles i.e. - important for transporting nutrient elements
essential for life
3.
Structure of Bodies of Water
– Thermal properties of water lead to stratification of bodies of water (formation of
layers)
– Upper layer is Epilimnion, lower layer is Hypolimnion
– Separated by a Thermocline
– The two layers to not mix well
– Epilimnion is heated by the sun, has relatively high levels of dissolved oxygen (aerobic
conditions); where gas exchange with atmosphere and photosynthesis occurs; species usually
in oxidized state
– Hypolimnion often depleted in oxygen due to bacterial action on organic matter; chemical
species usually in reduced state; exchange of species between water and sediment
4.
Important Chemical Processes in Water
– Gas exchange with atmosphere
– Leaching and uptake from sediment
– Reduction and oxidation
NH4+ ⇔ N03–
– Acid-base reactions
CO3– + H2O ⇔ HCO3– + OH–
– Precipitation
Ca2+ + CO32– ⇔ CaCO3(s)
– Formation of complexes and chelation
– Photosynthesis
2HCO3– + light ⇔ {CH2O} + O2 + CO32–
– Microbial action
2{CH2O} + SO42– + 2H+ ⇔ H2S(g) + 2H2O + 2CO2(g)
B.
CHEMISTRY OF NATURAL WATERS
1. Dissolved Gases and Henry's Law
X(g) ⇔ X(solv)
This is an equilibrium so we can write an equilibrium constant = [products] / [reactants]. For
gases, the concentration is expressed as a partial pressure, P(X, g)
constant = [X, solvent] / P(X, g)
– This is called Henry's Law. It governs partitioning of gases between water and atmosphere.
The constant is KH (mol/L*atm–1)
– KH = [X, solv] / P(X, g)
– Therefore when gas equilibrates with a solvent, the amount which dissolves is proportional to
the partial pressure of the gas.
– As KH increases, the solubility increases
– Assuming X(g) ⇔ X(solv), KH has units of mol/L*atm–1 but if viewed as X(solv) ⇔ X(g), can
also see units as atm*m3*mol–1
– Typical values of KH (mol L-l atm-l):
H2
N2
CO2
7.8 X 10–4
6.5 X 10–4
3.4 X 10–2
CO
O2
O3
9 X 10–4
1.3 X 10–3
1.3 X 10–2
– Gases are less soluble at higher temperatures because for the process X(g) → X(solv), ΔHo
and ΔSo are negative. Enthalpy decreases as system becomes more stabilized and greater
order decreases entropy
– Nonpolar compounds e.g. methane, are not very soluble in water because ΔSo is very
negative. Water molecules form a cage around CH4, giving a more ordered system. The process
of CH4(g) to CH4(aq) is exothermic (ΔHo is negative)
a.
Oxygen in Water
– P(O2, 1 atm) is 0.21 atm (21% by volume).
[O2, aq] = {KH = 1.3 X 10–3} * {P(O2, g) = 0.21 atm} = 2.7 X 10–4 mol L–1 * 32.00 g/mol =
8.7 mg L–1 @ 25°C
Note: this calculation ignores the vapour pressure of water P(H2O, g) = 0.03 atm @ 25°C
– Fish need 5 - 6 mg/L O2
– Thermal pollution can cause fish kills due to decrease O2 solubility. (also due to decomposition
of biomass called biological oxygen demand (BOD) and chemical oxygen demand substances
(COD)).
– O2 determined several ways:
- titration by Winkler's method
- spectrophotometry of oxidized dyes
- electrochemically by Makareth O2 electrode
– Some important terms:
- TOC: total organic carbon; determined by oxidizing all organic C to CO2 and
analyzing with gas chromatography
- COD: chemical oxygen demand; react with Na2Cr2O7 + H2SO4, backtitrate the
excess dichromate with Fe2+
- BOD: biological oxygen demand; incubate with micro-organisms, measure
O2 level before and after
Note:
- last 3 measures arbitrary because not all organics dissolve with equal ease.
- TOC measures all organic carbon
- COD measures more easily oxidizable C such as alcohols or alkenes but leaves
alkanes, carboxylic acids and aromatic rings (they react too slowly).
- BOD better but time and temperature are arbitrary.
– Slowly oxidizing organics and lower temperatures has a lesser effect on dissolved O2
than quickly oxidizing organics or higher temperatures because O2 enters from atm.
– A pulse of pollution entering flowing water will cause a local decrease in O2, but [O2]
will increase downstream as O2 enters from atmosphere
b.
Carbon Dioxide in Water
– Situation with carbon dioxide much more complex than oxygen because of acid/base
character of the molecule:
– +
H2O
–H+
H
–
CO2(g) ⇔ CO2(aq) ⇔ H2CO3(aq) ⇔ HCO3 (aq) ⇔ CO32–(aq)
– KH for dissolution of CO2 in water @ 25°C is 3.4 x 10–2 mol/L * atm–1.
[CO2(aq)] in equilibrium with CO2(g) P(CO2 = 3 X 10–4 atm) is
[CO2(aq)] = KH x P(CO2) = 3.4 x 10–2 mol/L atm–1 x 3 x 10–4 atm = 1.0 x 10–5 mol/L
as ppm, 1.0 x 10–5 mol/L * 44 g/mol = 4.4 x 10–4 g/L = 0.44 mg/L = 0.44 ppm CO2
Q: What is pH and total carbonate concentration in water in equilibrium with air?
A:
H2CO3(aq) ⇔ H+(aq) + HCO3–(aq)
1.0 x10–5
x
x
(do not have to write "1 x 10–5 – x" because H2CO3 replenished by the atmosphere)
Ka = [H+][HCO3–] I [H2CO3]
4.2 x 10–7 mol/L @ 25°C = x2 / (1.0 x 10–5)
x = 2.1 x 10–6 and pH = –log [H+] = 5.67
. Therefore 'pure' water is pH 5.67 rather than pH 7 because of dissolution of CO2. As pH of
water increases due to presence of other bases, solubility of CO2 and total carbonate increases.
. @ pH 4, log [CO32–] is about -5.0; @ pH 10, log [CO32–] ~–1.2.
Ka1= [H+][HCO3–]/[H2CO3] = 4.45 X 10–7,
Ka2 = [H+][CO32–]/[[HCO3–] = 4.69 X 10-11,
pKa1 = 6.35
pKa2 = 10.33
Henderson-Hasselbach Equation: pH = pKa + log([A–]/[HA])
– Let α(i) be the fraction of species i in solution, and
for pH < pKa1 α(CO2) is essentially 1
for pH = pKa1, α(CO2) = α(HCO3–)
for pH = 1/2(pKa1 + pKa2), α(HCO3–) at max value of 0.98
for pH = Ka2, α(HCO3–) = α(CO32–)
for pH > Ka2, α(CO32–) is essentially 1
2.
Dissolved solids in natural waters
– Amounts dissolved can vary widely
– Ground water can be very high, e.g. > 1000 ppm
– Ca2+ + HCO32– concentrations in rivers and lakes do not parallel that in the ocean because
fresh water dissolves ancient rocks containing CaCO3, whereas the oceans precipitate CaCO3
in the form of marine organism exoskeletons.
Typical Concentrations of ions in river and sea water
Ion
[ ] (river)
mol/L x 104
CI–
2.2
[ ] (ocean)
mol/L
0.55
Na+
2.7
0.46
Mg2+
1.7
0.054
1.2
0.028
0.59
0.010
3.8
0.010
9.5
0.0023
SO4
2–
K+
2+
Ca
HCO3
a.
–
Alkalinity
– Defined as concentration of all bases in solution, usually HCO3– and CO32–.
– pH is a poor measure of alkalinity because:
- weak carbonate bases more abundant than hydroxide ions.
- many natural waters not in equilibrium with the atmosphere.
– Can measure either "phenolphthalein alkalinity"
CO32– + H+ ⇔ HCO3–, endpoint pH - 8.5 using phenolphthalein
.Or can measure "total alkalinity":
CO32– + 2H+ ⇔ H2CO3, endpoint pH - 4.5 using methyl orange indicator
– Can determine carbonate composition by titration to both endpoints or measure one, then use
pH to calculate [CO32–] / [HCO3–]
– Alkalinity can be expressed either as moles H+ per litre ( how much acid used to neutralize all
bases) or in terms ofCaCO3 in mg L-1 (more common) as surrogate for bases in solution.
– Most natural waters have pH = 6 → 8.5, where HCO3– predominates
– Note that subterranean waters can have very high alkalinity due to exposure to very high [CO2]
generated by microbial oxidation of organic matter.
b.
Hardness
– primarily due to Ca2+ and Mg2+ 'alkaline earth elements'. Other elements e.g. Fe do contribute
hardness.
– Will ppt with soap
– usually measured by titration with EDTA.
– In fresh water, hardness and dissolved solids due to dissolution of rock:
limestone
gypsum
CaCO3 → Ca2+ + HCO3– + CO32– (not balanced)
CaSO4 → Ca2+ + SO42–
Note: limestone yields water that is both hard and alkaline.
gypsum yields only hardness since SO42– is very weakly basic (Kb = 8 x 10–13 mol/L):
CaSO4(s) → Ca2+(aq) + SO42–(aq)
Ksp = 3.7 x 10–5 (mol/L)2
maximum solubility is 6.1 x 10–3 mol L-1 (approx. at 25°C)
– Calcium sulphate imparts 'permanent hardness' because it cannot be removed by boiling
– CaSO4 is less prone to leave scale in hot water pipes.
– Calcium carbonate is less soluble - Ksp = 6.0 x 10–9 (mol/L)2 @ 25°C
– Becomes more soluble at lower pH because anion is basic:
CO32– + H2O ⇔ HCO3– + OH– or
CO3– + H+ ⇔ HCO3–
CaCO3(s) + H2CO3(aq) ---+ Ca(HCO3)2(aq)
– Above reaction explains why water with CO2 dissolved in it will dissolve carbonate rocks.
– why is hardness lower than usual sometimes?
- lower temp underground (~5ºC) gives higher CO2 solubility
KH (5°C) = 6.3 x 10–2
(25°C) = 3.4 x 10–2
- P(CO2) underground higher than at sea level pressure (up to 0.01 atm)
– dolomite limestone CaCO3•MgCO3(s) 1:1 composite of Ca + Mg carbonates is more soluble
than pure CaCO3
– temporary hardness' can be removed by heating, precipitating CaCO3
Ca(HCO3)2(aq) ⇔ Ca(CO3)(s) + H2O(l) + CO2(g)
– also leads to build up of scale in hot-water pipes, boilers, and kettles.
2.
Dissolved solids in natural waters
c.
Soft Water
– Soft water has low concentrations of Ca2+ and Mg2+, and low concentrations of anions,
e.g. carbonates, therefore soft water has low alkalinity (pH 7-8.5)
– Naturally found in areas with granite bedrock
– HCO3– concentrations are 10–4 to 10–3 M in soft water, > 10–3 in hard water
– Low alkalinity in soft water areas means low buffering capacity for acid rain
Artificial Water Softening
. Add lime (Ca(OH)2) or soda ash (Na2CO3) to precipitate Ca
Lime:
Ca(OH)2 + Ca(HCO3)2 ⇔ 2CaCO3(s) + 2H2O
Soda Ash:
Na2CO3 + CaSO4 ⇔ CaCO3 + 2Na+ + SO42–
– Soda ash added if hardness is mainly due to CaSO4
This method is used industrially; it has the advantage of low cost.
– Can use ion exchangers with negative ion exchange sites
Ca2+(aq) + 2Na+(A–) ⇔ 2Na+(aq) + Ca2+(A–)2
– Regenerate the resin after most of the Na+ has been removed by exchange with Ca2+
Ca2+(A–)2 + 2Na+(aq) _ Ca2+(aq) + 2Na+(A–)
– Calcium ions are washed out to drain, along with excess NaCI.
– Need a moderate equilibrium constant so that low levels of Ca2+ will displace Na+, but that high
levels of Na+ will displace Ca2+
– Common exchange resins include polystyrene with divinylbenzene, and zeolites (SiO2
and AIO4 units); the Al in the zeolites provides the exchange sites and the polymer provides the
support for the zeolites
– Deionized water produced by using 2 ion exchangers in series
cation exchanger in H+ form:
Ca2+(aq) + 2H+(A–) ⇔ 2H+(aq) + Ca2+(A–)2
regeneration step
Ca2+(A–)2 + 2H+(aq) ⇔ Ca2+(aq) + 2H+(A–)
anion exchanger using N+R3 groups in OH– form:
CI–(aq) + (C+)OH– ⇔ (OH–)(aq) + (C+)CI–
(C+) = cation exchange site
regenerate within NaOH
3.
Dissolved solids and irrigation
– Large scale irrigation can lead to increased soil salinity and aquifer depletion
– Salinity can increase to point that agriculture is impossible
– Back flushing with large amounts of water can wash salts back into the river
– A problem encountered in San Joachin valley: selenium was dissolved out of the soil, resulting
in concentrations lethal to fish
– Other problems:
1. backflush water can raise dissolved solid concentrations above 1 g/L.
2. depletion of river flow or lake volume.
3. underground aquifers can be depleted since water use can be 10 - 1000 times greater than
recharge rate e.g. Aral Sea in Russia; aquifer under Mexico City.
4.
Seawater
– Major reservoir of soluble ions carried down by rivers.
– Contains major components by mass:
CaCO3, CaSO4•2H2O, NaCI
1:5:250
Ion
Na+
Mg2+
K+
Ca2+
Cl–
SO4–
HCO3–
Input from rivers
1010 mol/y
900
550
190
1220
720
380
3200
Concentration
Mol L-1
0.046
0.054
0.010
0.010
0.55
0.028
0.0023
Resident time
yr
7 x 107
1 x 107
7 x 106
1 X 106
1 X 108
1 X 107
1 X 105
– When seas dry up, evaporates form, made of gypsum (CaSO4), sodium chloride, or
mix of Na and Ca chlorides.
– Calcium carbonate incorporated into exoskeletons of marine organisms.
– Oceans likely near equilibrium re: CaCO3(s) ⇔ Ca2+(aq) + CO32–(aq), based on observations
that very few "evaporites" form spontaneously in the oceans, and that sea shells on beachs do
not dissolve . Ksp for CaCO3 is 6.0 x 10–9 (mol/L)2.
– We calculate Qsp for calcium and carbonate:
Qsp = [Ca2+][CO32–] = (0.01)(2.7 X 10–4) = 2.7 X 10–6 M2
Since Qsp » Ksp , oceans appear supersaturated with respect to CaCO3.
– need to deal with:
1. ionic strength
2. complexation
a.
Ionic Strength
– Need to deal with activities
Activity = concentration x activity coefficient
At very low concentrations, coeff → 1, at high concentrations, coeff <1.
In sea water, Ca2+ coeff. is 0.26, and 0.20 for CO32–.
therefore Qsp = a(Ca2+) x a(CO32–) = 1.4 x 10–7 M2
– Ocean still appears supersaturated with respect to CaCO3
b.
Complexation
– Ions can form complexes with ions of opposite charges (counter ions), or with neutral species
– Ions can form complexes with specific counter ions at measurable concentrations:
(CaSO4)°, (CaHCO3)+, (MgCO3)°, (CaCO3)°, (NaCO3)–, (MgHCO3)+, (NaHCO3)°
– Ca2+(aq): 0.91 exists as free ion in presence of anions.
– CO32–(aq): 0.10 exists as free ion in presence of cations.
– Therefore can calculate Qsp:
Qsp = a(Ca2+, aq) x a(CO32–, aq)
= (0.010 x 0.26 x 0.91) x (2.7 x 10–4 x 0.20 0.10) = 1.3 x 10–8 M2
Much closer to Ksp but still greater. Ocean appears to be at least saturated with respect to
CaCO3
5.
Other Water Quality Issues
– Oxygen depletion
– industrial and sewage wastes . acidification
– non-polar toxic organics
– toxic elements e.g. mercury
C.
ACID RAIN
1.
Description
– Rain more acidic than normal e.g. < pH 5.0
– Generally associated with burning of high sulphur coal and industrial activities.
– Recall:
CO2(aq) + H2O(I) ⇔ H2CO3(aq) ⇔ H+(aq) + HCO3–(aq)
– Recorded acid rain pH <2, usual = 4 - 4.5.
– Vinegar and lemon juice pH 3.0, 2.2 respectively
2.
Sources
– SO2 and NOx in industrial areas become oxidized in atmosphere
SO2:
SO2 + H2O ⇔ H2SO3
ox. agt
SO2 ⇔ SO3 + H2O ⇔ H2SO4
NO2:
2NO2 + H2O ⇔ 1/2HNO2 + 1/2HNO3
NO2 is not very soluble in water
HX
NO2 + O3 ⇔ NO3 ⇔ HNO3
this occurs at night time: H-abstraction from a suitable HX donor
by NO3 free radical
– SO2 major contributor in acid rain areas e.g. east coast
– SO2 and HCI from volcanic eruptions can contribute
– On west coast America, NOx predominates; acid rain and photochemical smog are closely
linked
a.
Sulfur oxides
– From coal burning, and roasting metal ores.
– Coal typically 2 -3%
S + O2 ⇔ SO2(g)
– coal used mainly for thermal power stations and steel-making.
– many metals come as sulfide ores e.g. NiS, Cu2S, ZnS, PbS, HgS.
– roast to make oxides (Ni, Zn, Pb)
2MS + 3O2 ⇔ 2MO + 2SO2(g)
MO + C(s) ⇔ M + CO(g)
– for Cu2S and HgS, get direct conversion:
Cu2S(s) + O2(g) ⇔ 2Cu(s) + SO2(g)
– Ni smelter in Sudbury is largest single point source of SO2 emissions
– another source of SO2 is the oxidation of H2S, which comes from microbial decay of organic
matter, and from H2S in 'sour' natural gas
– oxidation steps:
OH•
SH•
SH•
S +
SO•
+ H2S ⇔ H2O + SH•
+ O2 ⇔ SO + OH•
+ OH• ⇔ H2O + S
O2 ⇔ SO• + O•
+ O2 ⇔ SO2 + O•
Note:
– The role of OH• here is similar to that of OH• in the oxidation of hydrocarbons
(photochemical smog formation)
– oxidation of S + O2 ⇔ SO2 does not occur directly
. removal of O atom from O2 has a lower activation energy than insertion of S into O–O.
–.Coal burning and metal smelting also produce NOx due to high heat
N2 + O2 ⇔ 2NO
ΔHº = +180 kJ/mol
heat
NO + O3 ⇔ NO2 + O2
Light
NO + HO2• ⇔ NO2 + OH•
NO2 + OH• ⇔ HNO3•
* note that there is a link between photochemical smog and acid rain. *
– On west coast of the USA, nitric acid is a major contributor to acid rain.
Heat
light, O3, HO2•
N2 + O2 ⇔ 2NO
⇔
NO2
3.
light, OH•
⇔
HNO3
Chemistry of acid rain.
– Rain normally has pH of 5.6 due to equilibrium with CO2 in the atmosphere.
– NO2 is precipitated as HNO3, SO2 as H2SO3 and H2SO4.
– H2SO4 and HNO3 are strong acids, H2SO3 has a Ka = 1.7 x 10–2 mol/L @ 25°C
– HNO3, SO2, SO3 are more soluble than CO2, so low concentrations of these gases have a
greater effect on pH than CO2
`
SO2
SO2(g) + H2O(l) ⇔ H2SO3(aq)
H2SO3(aq) ⇔ H+(aq) + HSO3–(aq)
KH = 1.2 mol/L atm–1
Ka = 1.7 X 10–2 mol/L
SO2(g) + H2O(l) ⇔ H+(aq) + HSO3–(aq)
Kc = 2.1 (mol/L)2 atm–1
CO2
CO2(g) + H2O(I) ⇔ H2CO3(aq)
H2CO3( aq) ⇔ H+(aq) + HCO3–(aq)
KH = 3.4 x 10–2 mol/L atm–1
Ka = 4.2 X 10–7 mol/L
CO2(g) + H2O(1) ⇔ H+(aq) + HCO3–(aq)
Kc = 1.4 x 10–8 (mol/L)2 atm–1
– H2SO3 is both more soluble and a stronger acid than CO2.
– 0.12 ppmv SO2(g) in rainwater results in pH = ~4.3
. The chemistry of acid rain from sulphur oxides complex because it can be deposited in various
forms:
1. can react to form H2SO3(aq)
2. can be oxidized to SO3(g) and react to form H2SO4(aq)
3. can precipitate as wet deposition or assoc. within particulate matter (deposits as sulfite
or sulphate ions.
4.
Oxidation of SO2
. Routes:
1. homogeneous oxidation in the gas phase
2. homogeneous oxidation in the aqueous phase in raindrops
3. heterogeneous oxidation on surface of particles
SO2(g) + 1/2O2(g) ⇔ SO3(g)
ΔG°(298 K) = –71 kJ/mol
– The negative Gibbs Free Energy, ΔG° indicates that at 298 K and close to the emission
sources where there is not yet any SO3 formed but a high concentration of SO2, the reaction to
form SO3 is spontaneous
– However, SO2 is known to be stable in dry air at 300 K
- SO2 will oxidize at higher temperatures in the presence of a catalyst
a.
Homogeneous gas phase oxidation
– Most important reactions in troposphere involve the OH radical:
SO2(g) + OH• ⇔ HSO3• (radical)
K = 9 x 10–13 cm3 molec–1 s–1
– For average [OH•] = 8 x 105 molec/cm3, half-life estimated to be 1O days.
– HSO3 oxidized by reaction with molecular oxygen:
HSO3 + O2 ⇔ SO3 + HO2
. Could oxidize SO2 by reactions with O (atomic O) but concentration is too low in troposphere
(1 x 105 molec/cm3) for the reaction to be important
– Oxidation by HO2 or O3 also slow due to small rate constants
b.
Homogeneous water phase reaction
– There is extensive aqueous phase chemistry going on in cloud-, fog-, and rain-water.
– OH radicals can either partition into water phase or they can be formed directly in the aqueous
phase, then oxidize SO2
– Hydrogen peroxide in water can also oxidize SO2
– H2O2 forms by disproportionation of HO2 radicals:
2HO2• → H2O2 + O2
SO2(aq) + H2O2(aq) ⇔ H2SO4(aq)
H+
HSO3–(aq) + H2O2(aq) ⇔ HSO4– + H2O
.details of mechanism of oxidation of SO2 in clouds or raindrops still uncertain:
1. HO2• could disproportionate before entering drop or dissolve then disproportionate
2. What is the pH dependence of the reactions, and effects of metal ion catalysis
e.g. Fe (II)
3. Other oxidants other than H2O2, such as O3 or NO2 (present at higher levels in polluted
air) may be reactive
HSO3–(aq) + O3(aq) ⇔ HSO4–(aq) + O2(aq)
SO2(aq) + NO2(aq) ⇔ SO3(aq) + NO(aq)
c.
Heterogeneous oxidation of SO2 on particles
. The heterogeneous oxidation of SO2, thought to be similar to industrial oxidation of SO2; not
known in detail
catalyst
SO2 + 1/2O2 ⇔ SO3 ⇔ H2SO4 or SO42–
– Could be catalyzed by salts of V, Mn, Fe
– Could be electron-transferring semi-conductor substances on particles that could act to reduce
O2 to H2O2
5.
Oxidation and Deposition of SOx
– Regional problem due to short half-life of SO2 and SO3 (about all gone in 1 week).
– Air mass is able to move - 3,500 km in only 1 week
– Oxidation rates vary greatly:
1. in dry air, 0.2% / hr (predominantly homogeneous gas phase oxidation by OH radical)
2. typical rates, 1 to 10% / hr
3. humid air up to 30% / hr (where oxidation takes place mainly in the aqueous phase).
– Rates also depend sunlight intensity, which generates the oxidants OH• and H2O2. In general,
as humidity increases, oxidation rate increases since aqueous oxidation rate>>air oxidation rate
– Can derive reaction rates that combine rates of all processes and average conditions:
SO2
K1 = 0.08 hr–1
⇔
SO3
SO2
K2 = 0.025 hr–1
⇔
dep. H2SO3
SO3
K3 = 0.03 hr–1
⇔
dep. H2SO4
Fig. 6.2 shows changes in concentrations of SO2 and SO3 with time:
a. [SO2] initially high, then it decays within time;
b. [SO3] rises, then falls again
c. rain depositing close to source will be H2SO3 (sulphite), farther away H2SO4 (sulphate)
will predominate
– In Sudbury (single site emitting about 1% of world's total SO2 emissions) the 'superstack'
penetrates the troposphere boundary layer and sends SO2 far away; therefore, northern Ontario
affected mainly from upwind sources other than smelter
– An extreme and natural example is the Smoking Hills, NWT.
-Combustable shales burning and releasing sulfurous smoke.
- small ponds have pH of 2 or lower
- pH 1.8 @ 40 m to pH 8.1 @ 4.4 km
- acidity 0.12 M @ 40 m to 6 x 10–4 M @ 670m.
- [SO4–] 16,000 ppm @ 40m to 110 ppm @ 4.4 km
6.
Effects of acidic emissions
1. effects of gaseous pollutants
2. effects of lowered pH
a.
Effects on Natural Waters
– Acidification is a problem where the rocks underlying lakes can not buffer the acidity . Buffered
lakes have chalk and limestone in soil beneath to neutralize acid
- Poorly-protected lakes set on granite
– Reaction within limestone:
CaCO3(s) + 2H+(aq) ⇔ Ca2+(aq) + CO2(g) + H2O(I)
– occurs as H+ reacts with HCO3– to produce H2CO3 ⇔ H2O + CO2, and so CaCO3(s) dissolves
to restore equilibrium.
– Northern and eastern Canada, north-eastern US and Scandanavia sit on granite rock, so are at
risk.
– normal pH range 6.5 - 7.
– many lakes now have pH < 5.0.
– Great Lakes basin receiving 400,000 tonnes S as SO2 and 200,000 tonnes N as NO2
– Some effects of lowered pH on aquatic organisms
pH
6.0
5.5
5.0
4.5
Effects
death of snails and cructaceans
death of salmon, rainbow trout
death of perch and pike
death of eel and brook trout
– Lakes in Adirondack region depleted of sport fish (Up-state NY)
- Correlation within loss of alkalinity
– Also similar problem N. Ontario
– Get a pulse of activity within melting of "acid snow" which affects hatching of fish eggs and
viability of hatched fry.
– Can try to counteract within depositing powdered limestone, but really need to control source
of acidity.
i)
Dissolution of metal ions
–Metal ions dissolved out of the bedrock
–Can be toxic metals: Cd2+, Pb2+, Hg2+
–dissolution due to reaction of H+ with a basic anion associated within metal
PbS: Ksp = 1 x 10–28 (mol L-1)2
– in pure water at equilibrium, [Pb2+, aq] = 1 x 10–14 M
– at pH = 4, [Pb2+, aq] = 1 x 10–5 M
– This could compromise drinking water quality
– Dissolution of aluminium important due to toxicity to fish
– Aluminum also has a complex aqueous chemistry
– Can exist in many dissolved forms: A13+, AIOH2+, AI(OH)2+, and poly anions depending on pH
and [AI] (the following discussion will ignore the polyanions)
– When the pH drops to less than 6.5, the solubility of Al rises steeply
– at pH > 6.5, dominant species is AI(OH)4–
– at a minimum between 5 to 6 (rain water pH 5.67 naturally)
– In Adirondack lakes, Al solubility greater than predicted.
- May be due to colloidal particles containing AI
- or due to F– in water, which forms a complex with AI, increasing solubility.
AI3+(aq) + F–(aq) ⇔ AIF2+(aq)
Kf = 2.5 X 106 L mol-1
– Why is an increased Al concentration toxic to fish?
– Lake has pH = 5, fish gill blood has a pH = 7.4
. There is a steep [H+] gradient across the gill membrane
– AI(OH)3 will precipitate as Al meets the higher pH.
– The precipitate is gelatinous and suffocates fish
b.
Environmental Toxicity
i.
Gas Toxicity
– SO2 toxic to plants at levels as low as 0.1 ppmv
– Plants can be damaged after a few hours exposure at levels 0.1 to 1 ppmv; these values can
reached in large cities
– Plants can be damaged in urban areas with SO2 and O3 at elevated levels
– Some combinations of pollutants, SO2/NO2 and SO2/O3 are synergistic
– Synergism occurs when the toxicity of the mixture is greater than sum of individual toxicities.
– The effect of nitrogen oxide is more complicated: NO3– , an ultimate sink of NOx reactions, is a
nutrient, while NO2 is a plant toxicant, esp. in combination within SO2
if.
Effects of Acidity
– Leaves are damaged by precipitation at pH <3.5
– Acidity changes soil chemistry, esp. poorly buffered soils (which are often naturally acidity)
– Fewer plants can tolerate acidity as can tolerate alkaline conditions
– Seed germination and seedling growth inhibited.
– Acid rain implicated in forest decline in Europe, but is controversial.
– acidic cloud could be a major sink for deposition of H+, NH4+, SO42, NO3– to upland forests
– Cloud droplets scavenge pollutants from the atmosphere, eventually forming a toxic mist that
can damage northern forests
c.
Effects on Health
– SOx, NOx are irritants to respiratory tract
– Upper tract affected at concentrations of 1 - 2 ppmv
– Lower tract affected at concentrations of > 25 ppmv
– combined within particulate effects
– need to set standards to protect plants since they are more susceptible to this problem.
d.
Effects on Structures
– Acidic gases will dissolve limestone in buildings in the same way as natural limestone
CaCO3(s) + H2CO3(aq) ⇔ Ca2+(aq) + 2HCO3–(aq)
K = 5.3 x 10–5 (mol/L)2
– This reaction occurs slowly in clean air
– Limestone will dissolve much more quickly in presence of acid
CaCO3(s) + H+(aq) ⇔ Ca2+(aq) + HCO3–(aq)
K = 1.3 x 102 mol/L
– Another process which damages structures is sulfation, which involves replacement of
carbonate by sulfate
– CaSO4 more soluble the CaCO3 and less structurally strong
CaCO3(s) + SO4(g) + 1/2O2(g) ⇔ CaSO4(s) + CO2(g)
Ksp (CaSO4) = 5 x 10–4 (mol/L)2
Ksp (CaCO3) = 6 x 10–9 (mol/L)2
– medieval stained glass being etched and frosted by acidic atmospheric pollutants
– increased rate of iron and steel corrosion, which is a reduction-oxidation (redox) reaction
General form:
Anode - oxidation:
Cathode - reduction:
Sum:
For Fe:
M ⇔ M2+ + 2e–
O2 + 4H+ + 4e– ⇔ 2H2O
2M + 4H+ + O2 ⇔ M2+ + 2H2O
2Fe(s) + 4H+(aq) + O2(aq) ⇔ 2Fe2+(aq) + 2H2O(I)
– Fe2+ oxidized to Fe3+ and precipitate as rust, Fe2O3, and easily falls off. Even steel protected by
zinc can be corroded in heavy polluted areas
c.
DRINKING WATER
1.
Sources of Drinking Water
–.Groundwater or surface water
– Groundwater is pumped up from aquifers
- Aquifers are recharged by rainwater percolating down into the aquifer
- The recharge rate can be very slow, and the rate of use can exceed the recharge rate,
depleting the aquifer
- Depletion of aquifers
- can cause land to subside
- contamination by infiltration of salt water
– Some of the water is several thousands of years old: "Fossil water"
– This water can be considered a nonrenewable resource on the time-scale of human life.
– Groundwater is generally less contaminated than surface water because:
1) soil bacteria can decompose organics;
2) groundwater is filtered during percolation
– Earliest water purification was filtration through sand beds, mimicking filtration of groundwater
– Sometimes surface water is re-injected into aquifers to replenish it, and permit purification
– Groundwater is threatened by industrial pollution
2.
Surface water
– comes from lakes and rivers
– More processing is required of surface waters than groundwater to make it drinkable. The
water is often contaminated by sewage (or poorly treated water) from cities and by industries
upstream
Industries contribute to metal pollution e.g. cadmium, lead and mercury, and can spill process
chemicals and wastes into rivers
– Need to treat groundwater and surface water before use
– What needs to be treated or removed:
volatiles: hydrogen sulphide, methane, odorous bacterial metabolites
solids: suspended particulate or colloidal particles
excess water hardness
excess iron and manganese
dissolved organics
pathogens: viruses, bacteria, etc.
Outline of Water Treatment
Categories of treatment:
- drinking water
- industrial use
- treatment of waste water for release back into the environment
Focus in this discussion will be on treatment for drinking water
a.
b.
c.
d.
e.
a.
Primary Settling
Aeration
Removal of excess hardness, iron and manganese
Coagulation
Disinfection
Primary Settling
– Allow the larger particulates to settle
– Can add lime at this point if the pH is < 6.5
b.
Aeration
– Aeration oxidizes the more easily oxidizable substances which could consume chlorine or
other disinfectants.
– removes more volatile compounds: H2S, CO2, CH4 and volatile organics.
– Ferrous iron, Fe2+, is oxidized to ferric iron, Fe3+, which precipitates as Fe(OH)3 at pH> 3.5
– For the homeowner, ferrous iron is removed by an oxidizing filter
c.
Removal of Excess Hardness, Iron and Manganese
– Usually remove hardness by addition of lime, Ca(OH)2 and/or soda ash, Na2CO3. Ca2+
precipitates as CaCO3, Mg2+ as Mg(OH)2
– Soda ash used to prevent pH from going too low and converting the carbonate to bicarbonate:
– Mg precipitates as Mg(OH)2 and requires a higher pH
– Can recover CaCO3 and convert it back to CaO by heating above 825°C, then reacting
with water to produce slaked lime, Ca(OH)2
– The lime-soda process can leave some CaCO3 or Mg(OH)2 in solution due to supersaturation
– These species could precipitate at a later time in pipes
– Ca and Mg are redissolved by bubbling CO2 through water to adjust the pH to 7.5 to 8.5 . When
pH is alkaline, [Ca2+] is adjusted to a point close to CaCO3 saturation and the water is said to be
"chemically stabilized" so that CaCO3 will not precipitate, nor will scale be dissolved . If the Ca2+
concentration is too low, water is termed "aggressive"
d.
Coagulation
– Want to remove colloidal minerals, bacteria, pollen and spores.
– To do this, can use filter alum at pH 6-8:
AI3+(aq) + 2HCO3–(aq) ⇔ AI(OH)3(s) + 3CO3(aq)
– AI(OH)3 is a gelatinous hydroxide that settles slowly, carrying down particulates
– Ferric hydroxide, Fe(OH)3, or activated silica (which forms gelatinous alkali metal silicates) can
serve the same purpose
d.
Disinfection
– This process makes water safe to drink by killing of pathogenic (disease-causing) organisms
– It also prevents recontamination of water, which is necessary since water can remain
in the system for 5 days before it is used
– Recontamination also accomplished by keeping the water pressure high throughout the
system to keep leaks flowing outward.
– Common disinfection agents: chlorine, ozone, chlorine dioxide, UV radiation
– Chlorine is the only agent that maintains a residual disinfection ability in the water after the
treatment at the source; other methods must be followed up by some further addition of chlorine
3.
Chemistry of Disinfectants
a.
Chlorine
– Chlorine dissolved in water forms several equilibria:
Cl2(g) ⇔ Cl2(aq)
KH = 8.0 x 10–3 mol/L atm–1
Cl2(aq) + H2O ⇔ H+(aq) + CI–(aq) + HOCI(aq)
Kc = 4.5 x 10–4 (mol/L)2
HOCI(aq) ⇔ H+(aq) + OCI–(aq)
Ka = 3.0 X 10–8 mol/L
–Hydrochloric acid is formed, reducing alkalinity
– the "active chlorine" species are: Cl2, HOCI, OCl–; - these species are oxidizers whereas Cl– is
not.
– The pH needs to be controlled to give the correct ratio of HOCI:OCI–
– HOCI is 100 times more effective than OCI– since it can penetrate cell membranes
– HOCI predominates at pH < 7.5, therefore the amount of Cl2 required increases as pH
decreases
– Some terms related to chlorination:
Chlorine dose:
amount of chlorine added
Chlorine residual:
amount of chlorine remaining
Chlorine demand:
amount of chlorine used; the difference between chlorine
dose and chlorine residual
Free available chlorine:
the amount of HOCI and CIO– in solution
– Chlorine gas is added under pressure as a liquid into a tank
– The residence time is 20 to 60 minutes, sufficient to completely disinfect the water The
concentration of Cl2 in the finished water is < 1 ppm
– The use of chlorine has some drawbacks related to taste and odour, and to toxicity. If phenols
are present in water, chlorination will form chlorinated phenols: 2-; 4-, 2,4-;
2,6-chlorophenols.The phenols come from use of some herbicides, pulp mill effluent, and
plastics manufacture. The chlorophenols have very low odour thresholds: 2-CP and 2,4-CP - 2
ppb; 2,6-CP - 3 ppb; 4-CP - 250 ppb
– Concentrations above 1 ppm render the water undrinkable
– Another major concern is the formation of trihalomethanes from the reaction of HOCI on humic
acids:
R-C(=O)CH3 + 3HOCI ⇔ R-CO2H + CHCI3 + 2H2O
– Humic acids are the breakdown products of lignin from plant material
Can remove organics such as phenols, lignin and humic acids by use of activated charcoal, but
this is expensive and not often used
– The chlorinated phenols and trihalomethanes could be toxic
- 2-CP fetotoxic in rats, 2,4,6-CP carcinogenic in rats and mice
– However, odour and taste threshold are so low that levels of these compounds are kept low for
that reason
– Chloroform is a promotor but not an initiator of cancer in rats
– Levels of chloroform in drinking water are very low
– There is some evidence that some drinking water (1 in 4 systems tested) have some
mutagenic activity (mutagenesis considered necessary for carcinogenesis)
– Cancer rates in communities which chlorinate water are very slightly higher than in those that
do not; there are possible confounding factors
– Chlorine itself is not carcinogenic, even at 100 ppm (objectionable taste at that level)
– The hazards of not disinfecting water are much greater than disinfection due to great risk of
water-borne diseases such as typhoid or cholera.
b.
Chlorine Dioxide
– CIO2 is a more effective disinfectant than HOCI, but more expensive. It can be used at lower
doses than Cl2, 0.1 to 5 ppm. Sometimes it is used where there is a taste or odour problem when
Cl2 is used. Unfortunately, CIO2 can not be stored as Cl2 can, it must be generated on site:
10NaCIO2 + 5H2SO4 ⇔ 8CIO2 + 5NaSO4 + 2HCI + 4H2O
– The reaction is a disproportionation of CIO2– to CIO2 and CI–, giving CIO2 free of Cl2 (aq)
– Sodium chlorite is a powerful oxidizer, so it must be handled carefully
– CIO2 is also a powerful oxidizer, and it disinfects through oxidation. It also oxidizes organics,
and introduces an oxygenated functional group into a molecule:
H2O
CIO2 + C6H5OH ⇔ CIO2– + [C6H5OH]+ ⇔ C6H4(OH)2
– CIO2 leaves not residual activity, so Cl2 must be added in low amounts to finished water to give
the residual activity
– Possible to create toxic chlorate ion CIO3–
OCI– + CIO2– ⇔ CIO3– + CI–
3OCI– ⇔ CIO3– + 2CI–
c.
Ozone
– Ozone has been used as a disinfectant since the early 20th century in Europe . Has been used
in North America since the 1940's; is used in Montreal
– Ozone generated by a 15 kV electrical discharge in dry air
– Ozone is absorbed into water (KH= 1.3 x 10–2 mol/L atm–1)
– Like CIO2, it is a more powerful disinfectant than chlorine, but more expensive. It decomposes
rapidly in water. Need chlorine for residual activity
–.The equipment is expensive, but economical when used on a large scale
d.
Ultra- Violet Radiation
– Radiation with a wavelength less than 300 nm will damage wildlife
– Germicidal lamps use λ = 254 nm
– The effects are due to absorption by DNA, creating photochemical defects, leading to fatal
errors of transcription
–UV can be used in large or small scale
– Some urban home systems and rural water treatment systems use UV disinfection
– Since UV need to penetrate water, need to remove suspended matter, organics, coloured
material first
– Advantages of UV: needs only a short contact time, so can disinfect water in a flow–through
system; low cost; no toxic residues; not influenced by pH or temperature
4.
Inorganic Contaminants in the Drinking Water Supply
– Iron is a nuisance because it has a 'metallic' taste and leaves stains. It precipitates out when
oxidized, so it is not present in toxic amounts
– However, lead, cadmium, mercury, aluminum and sodium can be present in amounts that may
have adverse effects
– Excess nitrate levels can also pose a problem
a.
Lead
–.Lead is toxic and has a long residence time in the body i.e. it is a cumulative poison. Fetuses
and babies are more susceptible to its neurotoxic effects than adults. Lead is a problem where
lead piping was used for in-house plumbing, also lead in solder in copper pipes
– In hard water areas, CaCO3 scale can cover the metal and prevent water from contacting the
lead
– However, soft water will dissolve lead more easily since it is a bit more acidic than hard waters
(due to the elevated carbonate content of hard waters). The hot water side of the domestic water
system generally has more lead than the cold side because: lead more soluble in hot water;
scale dissolves in hot water heater and leaves the pipes unprotected.
– Home with lead pipes should run the taps out for 1 to 2 minutes each morning to flush out any
water that could have accumulated lead overnight. In a survey of northeastern US homes, 90%
had lead above the 50 μg/L 'safe' standard. In extreme cases of Pb > 0.8 mg/L, brain damage
was seen in children (both pregnant women and developing children drinking contaminated
water)
– Standards have been strengthened: max allowable level reduced to 10 μglL
b.
Cadmium and Mercury
– These metals are also cumulative poisons
– They are usually found as contaminants in raw water
– In northeastern US homes (see Pb survey above), 2% of population exposed to Cd> 10 μg/L,
22% exposed to Hg > 2 μglL
c.
Aluminum
– While aluminum is very abundant in the earth's crust, there is no apparent biochemical function
of aluminum.
– Al can cause following health problems:
- anemia (AI interferes with iron transport in the body)
- softening of bones (calcium phosphate replaced by aluminum phosphate)
- a form of senile dementia
- implicated in Alzheimer's Disease (not yet substantiated as a cause; may be a spurious
correlation)
– Al is present naturally in water, but levels greater than the 2.7 ppm maximum solubility may be
due to use of large amounts of filter alum in water treatment.
– People can also be exposed to Al through aluminum cooking ware (acidic foods such as
tomato sauces, citrate juices) can cause Al to leach into the food.
d.
Sodium
– Sodium is present in naturally soft water
– for NaCl levels above 300 to 400 ppm, water is not palatable
– Na is present in softened water; it is recommended that unsoftened water be used for drinking
– Na pollution comes from industry and the use of NaCI as a road and sidewalk deicer. Should
keep Na below 100 ppm to reduce the chance of hypertension (high blood pressure).
e.
Nitrates
– The main source of contamination is from agriculture:
- manure seepage
- excessive use of nitrate fertilizers
– Fields are often cultivated right up to margins, so can promote run-off into waterways and
dugouts
– Nitrate contamination of rural wells and dugouts (a source of drinking water for many farms) is
a problem
– Nitrate is reduced to nitrite by intestinal bacteria, such as e. coli; this is less of a problem for
adults than for children
– Nitrite is toxic because it binds with hemoglobin and reduces absorption of oxygen; the
condition is called methemoglobinemia
– This can lead to mental impairment in infants
– The daily intake of nitrate/nitrite is mainly from food such as processed cured meats . Limits in
the U.K. are 500 ppm NaNO3
– In the stomach, low pH promotes conversion NO2– to H2NO2+, which can nitrosate secondary
amines and secondary amides
– Nitrosamines are initiators of cancer in animals, but is not a confirmed human carcinogen
– Demethylnitrosamine from industrial activity can also contaminate drinking water; another
source is microbial degradation of proteinaceous materials
5.
Organics in the Drinking Water Supply
– Natural sources:
- decay of biological materials
– Human activities:
- food and meat processing plants
- manure from feedlots
- synthetic compounds: insecticides and herbicides, sewage from municipal and industrial
waste dumps, sewage outflows, cooling water, organic air pollutants
– Concern about organics in drinking water rose in the 1960's when analytical techniques were
developed that could identify and quantify sub-ppm levels of organics
–Due to complaints and concerns about taste and odour problems, a landbreaking study of
contamination carried out in New Orleans: checked raw water and drinking water
– Finished water contained hundreds of organic contaminants which were present in raw water
and were passing through the water treatment process to enter the finished drinking water
– Among the compounds found: agricultural chemicals (insecticides and herbicides),
industrial organics (BTEX, phthalates, chlorinated benzenes and phenols, petroleum
hydrocarbons)
– There are concerns about adverse health effects of organic contaminants
– Elevated cancer levels have been found in fish from the Great Lakes
– Even if man is not significantly affected, need to reduce effects on the aquatic environment
– Some organics are of natural origin, e.g. Geosmin
– This is a monoterpene produced by Actinomyces bacteria
– It gives the smell of freshly dug earth or cooking beets
– Gives water an undesirable odour
– Is a tertiary alcohol that is resistant to oxidation, including chlorine and ozone
– It is a problematic compound in Regina because of algal blooms on Buffalo Pound Lake. The
city had to install activated charcoal filters to remove geosmin and other organics
–.The charcoal is produced by partially burning wood chips under oxygen-starved conditions;
produces a highly absorbent material.
– The charcoal is reactivated by burning; this removes the absorbed organics but also lose about
10% of the mass of the charcoal
– The public can also be exposed to VOCs by showers in addition to drinking. Such volatile
organics can be removed by aeration in towers.
– Reverse osmosis is a technique often used to remove dissolved salts and organics from water.
– Normal osmosis arises when a semi-permeable membrane separates two solutions of differing
concentration. Water will flow from the dilute to the concentrated solution, generating an osmotic
pressure.
– In the reverse case, the side containing the concentrated solution is .pressurize to counter the
natural tendency for water to migrate from dilute to concentrated, and the water flow across the
membrane is reversed.
– Clean water is generated on the low pressure side, and the concentrated solution of brine is left
on the high pressure side. Can be used to purify sea water or brackish water for drinking
D.
SEWAGE AND WASTE DISPOSAL
– In this section we deal with the treatment and disposal of municipal wastes, which are mainly
sewage wastes, and other aqueous wastes from industrial and agricultural sources
1.
Sewage Treatment
– We wish to discharge waste waters into lakes or rivers without unacceptable adverse effects
on the quality of the receiving waters
– The effluent should be free of organic and inorganic toxic compounds, pathogens, and
excessive biological oxygen demand (BOD).
– Industrial discharges imperil ecosystems and human health. The contamination of drinking
water by untreated sewage can cause cholera and typhoid
– Many cities continue to discharge raw sewage directly into receiving waters e.g. Victoria; and
some have only recently added treatment systems e.g. Montreal
– There are three steps in treatment of municipal sewage:
1. primary settling
2. secondary treatment (and disposal of generated sludge)
3. tertiary (advanced) treatment and removal of special wastes such as phosphorus,
a.
Primary Settling
– This step involves the removal of solid materials from the waste stream
– There are several stages, including a coarse screen, a grit tank, then a lagoon or clarifier to
allow smaller, lighter material to settle
– The steel screen removes large objects
– The grit tank collects grit, sand, etc.
– The effluent from the primary settling step is nearly clear but has a high biological oxygen
demand (BOD)
– The solids collected at this stage are sent to a municipal landfill
– An advanced primary treatment system has been tried as a replacement for conventional
primary and secondary treatments (see below)
– It involves the use of various coagulants (filter alum, ferric chloride, charged ionic synthetic
polymers) which could remove the tiny particles responsible for a large portion of the BOD
– The advanced primary treatment has the advantages of: lower capital costs and less sludge
production; disadvantages include: higher coliform counts, addition of more chemicals to the
waste stream
b.
Secondary Treatment
– Meant to reduce BOD about 90% (to less than 80 ppm)
– There are two approaches: trickling filter and activated sludge reactor
i)
Trickling Filter
– Composed of a bed of gravel and sand that becomes finer with depth
– It is colonized by micro-organisms which degrade organic material
–. Its purpose is to use up the BOD under controlled conditions before the effluent reaches the
receiving waters.
– The same bed can be used for several decades if cared for properly
– The micro-organisms can be poisoned by toxic substances in industrial effluents.
– The large bed requires a lot of land area
– It doesn't work well at lower temperatures, which is a problem for some Canadian locations
ii)
Activated Sludge Reactor
– This is an enclosed, heated tank where biological activity is maintained under optimum
conditions
– The tank contents are agitated and aerated, and micro-organisms that leave the tank with the
treated effluent are recycled back into the tank
– After secondary treatment, the effluent can be passed through a settling tank to capture solids;
filter alum may be added at this point to promote settling
c.
Sewage Sludge
– A sludge has a very high water content (> 95%)
– It can be de-watered by heating (digestion), causing coagulation of the solids
– Digestion is anaerobic, releasing reduced gases such as CH4and H2S
– The de-watered sludge can be air-dried to a solid
– There are various disposal options:
- land fillings
- incineration
- ocean dumping
- land spreading
– Sludges are rich in organic matter, especially nitrogen and phosphorus
– It can be used as a fertilizer and a soil conditioner
– As a fertilizer, sludge is not as good as regular fertilizers, so it is not economic to haul long
distances
– It has been used to remediate severely disturbed land, where it is spread at rates of 100's of
tonnes/hectare (organic matter helps retain moisture)
– Unfortunately, improperly treated sludge can contain toxic inorganic compounds and organic
compounds
Organics:
oxidation-resistant organo-chlorine compounds which become bound to the
organic portion of the sludge
Inorganic:
arsenic, chromium, cadmium, copper, lead, mercury, nickel, silver, zinc
(mainly industrial sources); metals can be taken up by crops, or leach from soil
and contaminate groundwater
Chromium:
Nickel:
Copper:
Zinc:
Cadmium, silver, mercury, lead:
not concentrated by plants;
toxicity to plants is only observed in highly acidic soils;
sheep are susceptible to copper poisoning-;
while a necessary micronutrient, it is toxic at higher levels;
Cd is of most concern-; heavy metals could reduce yields
of leguminous crops due to toxicity to N-fixing bacteria
– There are regulations pertaining to the amount of sludge that can be applied to land to be used
for crops and to the waiting time between application and harvest
he movement of metals through soils depends on soil type:
Sandy soils:
these soils are highly porous to water and dissolved metals, so
movement through sandy soils is rapid;
Highly organic soils:
the organics will bind metals, but not strongly enough to prevent uptake
by plants if the concentration is high enough;
Clay soils: clay soils will bind metal tightly and movement is slow; clay has a net
negative charge and cation-exchange-capacity; the negative charges
are balanced by interstitial cations; when sludge reacts with clay, the
Na+, K+, Ca2+, Mg2+ cations from the clay move into the sludge and
the metals in the sludge move into the clay
– The availability of metals to crops depends on pH
– Metals have insoluble hydroxides and carbonates that become more soluble as acidity
increases
– Liming can reduce acidity and hence availability
– A high acidity inside the clay limits the exchange of heavy metals into the clay . The availability
of metals also depends on character of the sludge
– sludge binds cations tightly, so metals are less available than they would be as simple salts in
solution
What can be done with the effluent from secondary treatment?
– can increase the oxygen content of the effluent before discharge into a river or lake so that
after the BOD of the effluent has depleted some oxygen content, there is still enough oxygen left
to support life
– can use as irrigation water for crops, but must be careful of salt content
d.
Tertiary Treatment
– Tertiary treatment is done to remove specific compounds
– The most commonly removed component is phosphorus (phosphates)
i)
Phosphorus
– Sources of anthropogenic phosphate include:
- sewage
- fertilizers and manure from agriculture (phosphate fertilizers, manure spreading,
seepage of wastes from feedlots
– Phosphates are nutrients that contribute to the eutrophication (accelerated ageing) of lakes
– Eutrophication: excessive levels of nutrients cause algal blooms, which are large floating mats
of algae, during the summer
– Dying algae greatly increase BOD, decreasing the amount of dissolved oxygen
– The decrease in oxygen causes problems for other aquatic life
– Dead algae contributes to increased rates of sedimentation
,.
– Phosphorus is usually the limiting nutrient.
– Uptake of nutrients is usually in the ratio of C:N:P = 100:12-15:1
– C as carbonate is usually replenished by atmospheric CO2.
– N is usually provided by N-fixing blue-green algae
– Major problem with phosphorus arose when phosphate-containing detergents entered the
marketplace in the 1960's.
– Phosphates comprised up to 50% of the formulations. The phosphates passed through the
treatment systems used at the time and entered natural waters in large amounts
– The large amounts of nutrient phosphates caused extensive algal blooms, especially in Lake
Erie where oxygen depletion nearly "killed" the lake
– Phosphate levels in detergent have been reduced to 5% (measured as P2O5) or so by
regulation
– The treatment of municipal waste is only one aspect of controlling phosphate problem because
as the level of phosphates in detergents drop, agricultural pollution becomes more important.
– To understand the methods used to remove phosphates, one needs to understand the basic
chemistry of phosphates in water
– The H2P04– anion can be written as HO-PO2-OH–
– Sodium tripolyphosphate (STP) Na3H2P3O10:
– Trimetaphosphate anion (P3O93–) is cyclic:
– Polyphosphates and metaphosphates are hydrolyzed to orthophosphates:
– Phosphorus removed by precipitation with lime (Ca(OH)2) to raise pH to ~9
– Low pH promotes solubility of phosphate salts
PO43–
H+, pKa = 12.3
⇔
HPO42–
H+, pKa = 7.2
⇔
H2PO43–
– Can also precipitate phosphates with Al3+ or Fe3+; this works at a lower pH as compared with
use of lime
M3+(aq) + PO43–(aq) ⇔ MPO4(s)
– Need to carefully control pH, since species actually exist as Mx(OH)yPO4)z
– Could get precipitation of the hydroxide rather than the phosphate
– Can use filter alum as a source of Al3+.
– Precipitation of the aluminum phosphates occurs in a narrow window (pH 4.5 to 5.5), if
pH >5.5, Al(OH)3 will precipitate.
– Can also remove phosphates using biological processes
– Use activated sludge process to incorporate P into microorganisms, leading to sludge that has
higher fertilizer value
– The water leaving secondary treatment still contains significant BOD, about 50 ppm O2 .
– This can be further removed by filtration of the very fine steel screens or coagulation
using filter alum, so that both coagulation and phosphate control can be combined
2.
Other Aqueous Wastes
. We can classify other wastes into 2 categories: high-strength and low-strength
a.
High Strength Aqueous Wastes
– The food industry releases organic wastes that are very high in BOD
– These wastes can put a very high strain on municipal sewage treatment systems.
– Pulp mills produce pulp by separating lignin from cellulose
– Lignin can be ~20% of the dry weight of wood, and lignin in the aqueous waste stream
generates much BOD
– Both these industries have to reduce their BOD and other pollutants to regulatory levels .
– The most common strategies involve either aerobic or anaerobic treatments
– Aerobic treatment involves use of micro-organisms convert waste to CO2, H2O and biomass
– Anaerobic treatment involves conversion of wastes to CH4, CO2, H2O and biomass.
– Aerobic treatments are more common because:
- Faster: this reduces the residence time, and can use small facility to treat a given volume
of waste per day;
- Oxidative: avoids the production of unwanted reduced N and S species
– The oxidative process is usually carried out in large aerated lagoons
– Closed bioreactors are sometimes used to reduce releases of air pollutants, to allow more
precise control of reaction conditions, and to allow use during cold winter months
– The sludge that is generated is digested, dewatered then land-filled
– When dealing with industrial wastes, one must be careful that toxic components in the wastes
do not poison the micro-organisms
– If the wastes are limited in necessary nutrients, then can add fertilizers, for example
ammonium phosphate
b.
Low Strength Aqueous Wastes
–.These wastes are defined as containing only a few ppm of organic compounds which are
resistant to degradation
Oily wastes
– Oily wastes are generated by the oil industry among many (eg. natural gas plants).
– Petroleum oils are not miscible with water and form a film on the surface of water.
– This film can impede oxygen exchange with the atmosphere, and are generally toxic to
organisms that reside in the interface (micro-organisms, birds, etc)
– The oils are mainly hydrocarbons, predominantly alkanes
– These oils need more treatment than wastes from the food industry, such as fatty acids and
triacylglycerides
– The oily aqueous waste can be treated by physical separation of the oil and water layers by
physical means, or by degradation/destruction of the oil in place.
– Can remove organics by use of an absorbent such as activated charcoal
– This method does not get rid of the problem, it just substitutes one problem for another: what to
do with the absorbent?
– The absorbent can be landfilled, incinerated, or regenerated through partial incineration and
combustion of volatilized pollutants.
Oily wastes can be landfarmed: spread over land and mixed in with soil with tillage down to a
depth of about 30 cm
– Makes use of natural biological action to degrade the hydrocarbons
– The land has to be regularly tilled, and needs addition of N and P fertilizers
– Contaminated lands, such as a beach contaminated by an oil spill, can be remediated in place
by biodegradation after addition of fertilizers
– In northern areas, landfarming is limited by short growing seasons
– Other problems include the release of volatile organics into the air, and build-up of resistant
toxic compounds such as PAHs, PCBs, other chlorinate compounds
– Land used for landfarming should not be 'used for growing of crops at a later date, so the use of
such land is limited
– As alternatives to landfarming, use of composting or closed bioreactors
– Other organic wastes include those from pesticide and explosives manufacture
(nitroaromatics, aminonitroaromatics, nitrophenols, etc.).
– Nitrophenols can also cause problems for drinking water during chlorination
– These compounds just mentioned are toxic and will kill the micro-organisms in bioreactors, so
alternatives must be found
– There are several alternative methods available for destruction of the organics:
- oxidation by UV and ozone or hydrogen peroxide
- oxidation by semiconductor catalyst
- electrochemical oxidation
UV and oxidants
– Use aerated water to which ozone or hydrogen peroxide is added . UV radiation cleaves
oxidants to OH• (aq)
H2O2(aq) + hv ⇔ 2OH•(aq)
– OH• radicals are very reactive towards organics in water:
OH• (aq) + R-H(aq) → H2O(I) + R•.
R• + O2 → ROO• → more oxidation products
– Oxidizing power actually comes from O2
– This method can be expensive due to cost of H2O2, power and Hg lamps.
– Strong UV can cleave pollutant bonds directly but it mainly works through OH• generation
Semiconductor-catalyzed Oxidation
– Generates OH•(aq) using visible light and titanium oxide (TiO2) semiconductor
– 400 nm light causes promotion of electrons into an upper conduction band, leaving a positively
charged 'hole' in the electronic structure ofTiO2
“+” hole (TiO2) + OH–(aq) → OH•(aq) + TiO2–
TiO2– + O2 → O2–(aq) + TiO2
(sample reaction, compound reacting could be
something other than O2)
– These reactions compete with the recombination of the promoted e– with the “+” hole which just
generates heat.
– OH• and O2– are highly reactive and start a chain of reactions as described above
– This approach eliminates the use of H2O2 and costly UV lamps of the technique above.
Electrochemical Oxidation
– Technique involves the generation of Ag(II) from Ag(l).
– Ag(II) is a very powerful oxidant
– OH•. is generated by a series of reactions:
Ag2+ + NO3– → Ag+ + NO3•
NO3• + H2O → NO3– + H+ + OH•
Ag2+ + H2O → Ag+ + H+ + OH•
. Use electrical energy to generate Ag2+
. Since Ag is expensive, use only catalytic amounts and immobilize the metal in a reactor
ii)
Removal of Volatile Organic and Inorganic Pollutants
– Volatile pollutants can be removed from water by air stripping
– The waste stream is sprayed down in a tower against a countercurrent of air
– Volatile components are transferred to the air phase; transfer governed by Henry's Law
constant for the substance, KH = [X(aq)]/ P X(g)
– Heating the water or using steam (steam stripping) will increase transfer to air phase.
– Is used to remove chlorinated ethylenes (eg. perchloroethylene from dry cleaning waste
streams) and volatile inorganics such as ammonia
– Problem is that it exchanges water pollution for air pollution
– Can oxidize organics in the air stream by passing the air over a catalyst bed and generating
CO2 and H2O
– A variant of this technique uses slime layer to absorb and degrade organics
iii)
Inorganic Aqueous Wastes
Ammonia
– Is toxic to aquatic life
– Since it is volatile, it can be removed from waste streams by air stripping
– Can also be removed by nitrification and breakpoint chlorination
– Breakpoint chlorination involves the use of chlorine or hypochlorous acid to convert ammonia
to elemental nitrogen N2
2NH3(aq) + 3HOCI(aq) → N2(g) + 3HCI(aq) + 3H2O(I)
– Chlorine is reduced from + 1 to -1 oxidation state
– The reaction proceeds sequentially through Chloramine (NH2CI) and Dichloramine (NHC12)
NH3(aq) + HOCI(aq) → NH2CI(aq) + H2O
NH2CI(aq) + HOCI(aq) → NHCI2 + H2O
NH2CI(aq) + NHCI2(aq) → N2(g) + 3HCI(aq)
– The technique is called "breakpoint" chlorination because of the formation of the "combined
chlorine" residual composed of chloramine and dichloramine
– Chlorine residual rises to a maximum, then falls back to zero (the "breakpoint") as the
chloramines react with each other to generate N2(g)
Cyanide
– Used in metal plating and finishing processes, and in the mining industry for ore separation.
– Cyanide as NaCN can be as high as 50 ppm in effluent from ore separation, higher as a result
of spills in metal plating
– Cyanide can be removed by reactions with chlorine or hypochlorous acid:
CN–(aq) + HOCI(aq) → OCN–(aq) + HCI(aq)
CN–(aq) + Cl2(aq) + H2O → OCN–(aq) + 2HCI(aq)
OCN–(aq) + H+(aq) + H2O → NH3(aq) + CO2(aq)
Heavy Metals
– Term usually refers to toxic metals in first transition series such as vanadium and chromium
– Can be removed from waste stream by precipitation with a suitable anion, usually sulphide (S-)
H2S(g) + M2+(aq) → MS(s) + 2H+(aq)
– Works best at high pH, to generate basic anion and to remove generated acidity. Can add lime
to effluent to raise pH
SOLID AND HAZARDOUS WASTES
A.
MUNICIPAL SOLID WASTES
1.
Introduction and Scope of the Problem
– We live in a high-throughput consumer society that take for granted planned obsolescence, a
high rate of consumer spending and almost total disregard for wastes. This is leading to
increasing costs of resources, scarcity of energy and a deteriorating environment.
– Previous statement was paraphrased from Science Council of Canada 1977 "Canada as a
conserver society; resource uncertainties and the need for new technologies”. Report No. 27,
Ottawa; reported in Environment Canada State of the Environment Report 1991, p. 25-4.
– In period 1988 to 1992 Canadians produced about 18 million tonnes/yr of municipal solid
wastes, about 637 kg per person and 1.7 kg per person /day
– About 1/2 of that from residential sources, rest is construction and residential wastes (wastes
from building, road and bridge work)
– Total production of wastes about 33 million tonnes/year.
– Canada usually cited as biggest per capita generators of wastes in the world
– National target of 50% reduction in wastes from 1988 levels by year 2000
– Decline in total solid waste generated by 16% 1988 to 1992; amount discarded declined 23%
in the same period (larger proportion recovered or recycled)
– Decline in construction and demolition and industrial and packaging wastes due in part to
economic recession 1990 to 1992.
2.
What is it?
– Municipal wastes are divided into several categories: residential; commercial, institutional
and light industrial; construction and demolition waste
– Municipal solid wastes made up of: paper and paperboard (25%), food wastes (19%), yard
wastes (13%) and plastics (12%); also metals, wood and glass.
– Paper wastes comprised of: newspapers, fine papers, card and boxboard, magazines, phone
books, wallpaper, tissues, etc.; some are contaminated with food wastes
– Generated 6 million tonnes of paper wastes in 1988
– Metals and glass can come from: spent and broken light bulbs and fluorescent bulbs, window
glass, ceramics, steel food and beverage cans, glass jars, etc.
– Yard wastes comprised of grass clippings, leaves, hedge trimmings
– Plastics wastes include beverage containers, packaging, etc.
– Small but important part of the household waste is household hazardous waste, to be
discussed later
– Composition largely unchanged 1988 to 1992
– Construction and demolition waste: asphalt (35%), concrete, rubble, wood, gypsum (from
drywall) and others.
3.
Who Generates the Wastes
– Household wastes can make up to 35 to 40% of the total quantity of garbage generated for
disposal
– Commercial wastes come from stores, offices and restaurants, for example.
– Institutional wastes come from schools, airports, hospitals.
– Industrial wastes usually divided into light industrial and heavy industrial: light industrial
wastes usually end up in municipal waste stream, while heavy industrial wastes do not
4.
Management and Disposal of Municipal Solid Wastes
Municipal solid waste are sent to landfills, incinerated, or recycled/recovered
a.
Landfills
– Many landfills are no more than rudimentary garbage dumps: dump garbage, cover with soil,
dump more garbage, cover and forget.
– Sanitary landfills: controlled to keep hazardous material out, constructed to prevent leakage,
to protect human health and the environment
– Most landfills generate thousands of litres of leachate
– Leachate comes from rainwater percolating through the soil and wastes, from groundwater
flowing through the subsurface, and from liquid wastes and liquids generated by
decomposing wastes
– The composition of the leachate depends on the materials in the landfill
– The danger comes from contamination of surface waters and groundwater
– Sanitary landfills have engineered liners (clay and''or synthetic materials) to prevent liquid
wastes from moving away from the site
– Also have a network of pipes to collect leachates which can be either treated on-site or sent
to the sewer system.
– If hazardous wastes are kept out of the landfill, levels of contaminants are generally low
– Landfills also produce landfill gases, which are generated from the anaerobic breakdown of
organic matter
– Gas moves upward through the site, carrying gaseous contaminants
– Many older sites (containing a mix of municipal and industrial wastes) have a gas
composition of: CH4, 53%; CO;, 25%; N2, 10%; O; 0.9%; other hydrocarbons, 0.1%
– The gases can be collected and burned (flared).
– These gases can pollute the atmosphere: about 50% of anthropogenic methane comes from
landfills
– Methane can also move horizontally and enter buildings up to several hundred metres away
– Methane forms explosive mixtures with air
– Generally a bad idea to build over a closed landfill site
– Landfills use up land that could be used for other purposes: agriculture, recreation, wildlife
habitat, housing
– It is becoming more expensive to site, build, operate, monitor and decommission landfills
– Most people don't want a landfill in their backyard
Incineration
– Much of the garbage that is not landfilled is burned
– Some household wastes are still burned in woodstoves (eg. papers and food scraps)
– In some places can still burn leaves and yard wastes outdoors
– While some wastes are burned at landfill sites, most are burned in incinerators
– Incinerators can be designed to recover energy as heat and converted to electricity and
steam hot water
– In 1991, 16 municipal solid waste incinerators were operating in Canada, each burning 15
tonnes day
– Some solid waste is burned in private incinerators eg. hospitals and apartment buildings
– Incinerators need to be operated at high temperatures to ensure nearly complete combustion
and conversion of combustibles to CO2 and H2O, otherwise, can generate toxic chemicals
– Incinerators emit acid gases (SO2, HCl, NOX), CO, particulate matter, and toxic substances
such as chlorinated dioxins and dibenzofurans
– Incinerators also generate ash (80% bottom ash, 20% fly ash), which is generally
contaminated with toxic metals and organics
– Ash needs to be tested (leaching of controlled substances) before disposal
– In 1988, 100,000 tonnes sent to municipal disposal sites, 100,000 sent to hazardous landfill
sites
– Like landfill sites, they are very difficult to site as there is usually much public resistance
c.
New Directions in Municipal Solid Waste Management
– Need to substantially reduce the amount of waste going to landfills
– Major initiatives include reducing the amount of packaging, and banning the disposal of
certain types of wastes, including batteries, tires, drywall gypsum
– 4R's: reduce, reuse, recycle and refuse
– Such initiatives work best when management of wastes are considered when designing a
product:
- design for recycling
- reduce toxicity of components by replacing them with more benign substances
– Source reduction: produce less waste at the source of production
– Waste reduction: reduce amount of waste going to disposal
– Reuse: repeated use of an item e.g. refillable bottles, or items with replaceable or repairable
parts
– Recycling: take apart products, use material to make a new product eg. plastics, newsprint
paper, metals
– Recycling can occur at the source: use of scraps put back into production Problem with postconsumer recycling is finding a market for material
– Composting: organic refuse decomposed by micro-organisms to produce material that can
be returned to the soil
– Done on household scale or municipal scale
– Composting could reduce quantity of waste going to landfill by 30%
B.
HAZARDOUS WASTES
1.
Introduction and Scope of the Problem
– Usually defined as: wastes that pose a risk to human health or the environment and require
special disposal techniques to make them harmless or less dangerous
2.
What are Hazardous Wastes
- based on physical/chemical characteristics
a.
Ignitability
–liquids having a flash point below 60 °C
– nonliquids liable to cause fires through friction, absorption of moisture, spontaneous
chemical change; will burn so vigorously and persistently that it is a hazard
– ignitable compressed gases
– oxidizers
– These are substances that could bum during storage, transportation or disposal
– Oxidizers will aggravate an already burning fire
– Burning materials may generate toxic fumes and create convection currents that may
transport particulate matter
– Such materials as tree bark, wood chips, sawdust and paper boxes are usually excluded
b.
Corrosivity
– aqueous wastes with pH < 2.0 or > 12.5
– aqueous wastes capable of corroding steel at a rate > 0.25"/year
– Total acidity or total alkalinity is not a criterion, but should be (higher levels indicate larger
amounts of bases or acids to neutralize corrosive material)
– Corrosion of steel a concern because of corrosion of steel waste drums
c.
Reactivity
– readily undergo violent chemical change
– react violently or form potentially, explosive mixtures with water
– generate toxic fumes in a quantity sufficient to present a danger to the environment or to
human health when in contact with water
– explode when subjected to a strong initiating force
– explode at normal temperature and pressure
– fall within regulations regarding transportation of explosive materials
d.
Toxicity
– Usually defined in terms of levels of contaminants in a leachate from hazardous waste using
a standard procedure (Toxicity Characteristic Leaching Protocol, TCLP)
– Extract with acidic water at pH 5.0
– Material is toxic if contaminants in leachate exceed l00 times drinking water guidelines or
other suitable guidelines
– Based on a worst-case scenario in which groundwater is contaminated by leachate from an
improperly constructed disposal site
– The factor of 100 is an educated guess and allows for attenuation in the soil
e.
Other Definitions of Hazardous Wastes
– Specific compounds, formulations, classes of substances
– Far to many to list here
– Include waste materials such as sludges and used cleaning solvent produced by industries,
discarded and spilled products and their containers, some chemical intermediates in
manufacture of certain products
3.
Who Generates Hazardous Wastes
– 1986, 8 million tonnes from industry'
– Consumers can generate hazardous waste: 2.5 kg/person/year of things like old paint,
batteries, pesticides, cleaners, other household products
– Total from consumers 60,000 tonnes/year; !% of industrial production
– Biomedical wastes 8,300 to 31,300 tonnes/year,'/: % of total hazardous wastes
– In Canada, the largest proportion was produced in Ontario and Quebec (59 % and 21 %
in 1986)
a.
Case Study - Petroleum Refining
– In 1990, there were 29 refineries in Canada generating a total of 240,000 m" crude oil per
day
– They generate air, water and solid waste pollution
– Air pollutants include: carbon monoxide, sulphur oxides, nitrogen oxides, particulates,
hydrocarbons, hydrogen sulphide, mercaptans, aldehydes, ammonia and phenolics
– Water pollutants include: crude oil; acid catalysts; process chemicals such as caustic soda,
sulphuric and phosphoric acids; amines; ammonia; detergents; nutrients for biotreatment;
additives for foam suppression, corrosion inhibition and water treatment; reaction products
and product additives such as corrosion inhibitors, antiknock compounds, anti-icing
compounds
– Principle contaminants in discharges are: oil and grease, suspended solids, phenols,
sulphides and ammonia nitrogen
– From 1972 to 1987 there has been a great improvement: 81% of the sites apply secondary
and tertiary treatment of water wastes
– Solid wastes include: sludges in bottom of tanks; spent catalysts, alumina, clay, and sand
from filters
– Sludges may contain volatile organics (benzene, ethyl benzene, toluene, xylenes), phenols,
PAHs and heavy metals
– 30 % of sludges are reused or recycled, 36 % are disposed of in landfills, 18% are spread on
land, 7 °o are incinerated and 1% injected into deep wells
b.
Case Study - Pulp and Paper Production
– Production of pulp and paper requires large volumes of water, on average 100 m3 per tonne
– Releases large volume of liquid wastes, which contain wood fibres, finely divided solids, and
complex mixtures of compounds derived from wood and process chemicals
– Also release large amounts of BOD and Total Suspended Solids (TSS); resins, fatty acids
and sulphur compounds which are acutely toxic to fish
– Mills using chlorine as the primary bleaching agent are a source of chlorinated dioxins and
dibenzofurans, though not as much as 10 years ago
– In 1990, 47 mills used chlorine, but now most have been modified to reduce the use of
chlorine and dioxin generation has also decreased
– In 1991, 125 out of 155 mills discharged directly into water bodies rather than into a
municipal water system
– However, there have been improvements: in period 1978 to 1985, BOD decreased from
3337 tonnes/day to 1961 tonnes/day and TSS decreased from 2106 tonnes/day to 816
tonnes/day; there have been more reductions since then
– Unfortunately air pollutant emissions have increased in the same period
c.
Case Study - Household Hazardous Wastes
– In 1990, Canadians generated 2.5 kg/person/year of hazardous wastes
– There is concern about this, because those wastes largely end up in unlined municipal
landfills, so that some of these wastes contaminate leachates and therefore groundwater and
surface water, and are emitted as toxic gases
– Some communities provide collection systems of various types to keep hazardous wastes
out of municipal landfills
– Identifying hazardous household hazardous wastes can be difficult: often use warning labels
such as "corrosive", "toxic", "reactive", "flammable"
– Some of the ingredients causing problems are: heavy metals, toxic organics and chlorine
– Benzene, methylene chloride, toluene, ethylene dichloride are found in: spot removers,
lubricating oils, latex paints, glues and nail polish
– Chlorine bleach is discarded in containers with residual levels still left
– Lead can be found in many components of batteries, inks, plastics (eg. some manufacturers
of vinyl miniblinds)
– Other products contributing hazardous compounds to municipal waste stream are: tile
cleaners, carpet shampoos, oil, batteries, turpentine, paints and stains, wood preservatives,
pesticides, drain cleaners, rat poisons, some pet care products, some cosmetics, pool
chemicals, ammonia, scrap tires
– Outdated pharmaceuticals are generally kept out of landfills by programs run by pharmacies;
these are generally incinerated
4.
Management and Disposal of Hazardous Wastes
– Until recently, such wastes were dumped without cleanup e.g. in unsecured landfills and
improperly decommissioned industrial sites
– Improperly stored toxic wastes at sites on the American side of the Niagara River are a major
source of pollution in Lake Ontario
– The major site, Love Canal, had wastes including organochlorine pesticides, PCBs and
dioxins
– Over 1000 people were forced to leave their homes
– In Canada at Port Hope, Ontario, they are faced with cleanup of low-level radioactive waste
– There are over 1000 hazardous waste sites, 3-5% of which the owners are unknown
– Cleanup costs for the "orphan" sites have to borne by the government
– Want to destroy or detoxify wastes, then dispose them securely
– Some materials that can not be neutralized, such as high-level nuclear wastes, have to be
properly disposed of in some way, such as deep in stable geological formations
e.g. Canadian Shield
– Detoxification:
- caustic alkaline material can be mixed with acid, to product salt and water
- oils can be degraded by micro-organisms
- can destroy some materials by incineration eg. PCBs, organochlorine pesticides
– In 1986, 65 % of industries disposed or treated their hazardous wastes on site (where
generated)
– Alternative is off-site treatment of wastes
– Off-site management has risks: need to store wastes safely until disposal; need to safely
transport wastes to disposal site
– Releases can occur through leaks, spills, fires eg. PCB warehouse fire in St. Basile-le-Grand
– The possibility of release is greater during transport, so reduce distance travelled or
necessity for transportation
– There is some transborder shipment of wastes; over the years 1987-1990 Canada imported
more than it exported.
– The amounts exported to the US increased in the same period, perhaps because more
treatment sites in the US.
– As controls become more stringent in the industrialized countries (US, Canada, Europe,
Japan), some companies try to ship overseas to countries with less restrictive laws
– There are conventions, such as the Basel Convention that controls the transboundary
movement of hazardous wastes
– Canada signed the convention in late 1992 and as of early 1996, 97 countries and the
European Union have signed on.
– Some waste are disposed of illegally through mismanagement, or by deliberate intent
(dumping in illegal dumps, ditches, or in other countries)
– In a 1990 study, Quebec could not account for about 1/3 of all wastes shipped off-site
– Problem is where to put permanent off-site treatment facilities
– Untreated wastes are accumulating faster than they can be treated
– Due in part to a lack of permanent off-site facilities
– The facility at Swan Hill Alberta is the only (1992) comprehensive treatment site in Canada
– While others have been planned, there has been much local opposition i.e. the NIMBY
syndrome
– While the risks may be small, there is usually vigorous opposition due to perceived risks
– There is a problem of credibility on the part of the company wanting to build the plant
– Such sites are usually sited based on technical, environmental, economic reasons
– It is better to involve communities at the earliest opportunity to make the process less
combative
– Let the communities "ask' for site and cooperate in designing the facility
– Mobile treatment facilities are also possible, such as mobile incinerators for disposal of PCBs
at over 3000 sites in Canada
– The use of mobile treatment facilities minimizes the risk of transport, and reduces the
likelihood of strong public opposition
a.
Ultimate treatment methods
– These refer to permanent disposal methods
– Distinguish between organic and biological wastes on one hand, and inorganic (metallic
wastes)
– The latter can be immobilized in a non-migrating non-leaching form but can not be destroyed
because the hazardous agent is an element
i.
Incineration
– Used for hospital and chemical hazardous waste
– Advantages of incineration:
- Reduced volume; industrial chemical wastes can be reduced to negligible volume,
municipal solid waste to about - 25 % original volume
- Efficient destruction of hazardous wastes and sterilization of hospital wastes
- Can recover energy as heat and or electricity
- May be able to use ash to make concrete, but problem with some heavy metals
– Biggest problem is the emission of particulates and toxic chemicals such as dioxins and
dibenzofurans; emissions of dioxins and dibenzofurans from incineration of hospital wastes
are usually greater than that for MSVV due to more chlorine-containing plastics and a lack of
pollution controls
– There has been limited incineration at sea for very toxic wastes, but this is controversial;
ship-board incinerators usually operate at much lower destruction efficiencies than landbased incinerators (99.99 % versus 99.9999 %)
ii.
Solidification
Process of making a liquid or sludge into a solid mass This is suitable for:
- Flue gas cleaning sludges
- Waste streams with toxic inorganics
- Waste streams with toxic organics
– Sometimes a distinction between solidification and stabilization; latter usually refers to
immobilisation through chemical reaction, or entrapment in an impermeable and inert
structure
Cement and lime based technologies
– Concrete is made from cement, sand and water
– Cement, e.g. Portland cement, is made from heating limestone and silicate minerals and
some aluminum oxide together at 1500 °C until the calcium carbonate is convened to
calcium oxide (CaO) and CO2; in a process called calcining; a clinker is formed that is
ground to produce cement
– The concrete sets when water is added to produce amorphous hydrated calcium silicate
– The concrete hardens as silicate fibres grow and CaO absorbs CO: from the air to reproduce
CaCO3
– Concretes and cements are generally alkaline
– Dry solid wastes are immobilized by replacing the sand with solid waste; or sludges can be
immobilized by replacing the sand and water with the sludge
– Suitable wastes include: flue gas cleaning sludges (CaSO3, CaSO4) or dry mixes of
CaSO3/CaSO4 flyash
– Can also include sulphide, hydroxides and phosphate precipitates of metals from waste
streams
– Concrete immobilization is good for these compounds because they will not leach under
basic conditions
– If the concrete is exposed to acidic conditions, may need to seal the surface
– A variation is the use of 'pozzolanic concrete' made of lime (Ca(OH);, water and pozzolanic
silicate minerals
– Can be used with fly ash, cement kiln dust, ground blast furnace slag
– This form of concrete is more susceptible to acid leaching than concrete based on Portland
cement
Vitrification
– Formation of a glass, which is a fused mixture of metallic oxides
– Usually contain CaO (usually added as CaCO3), Na2O, SiO2, B2O3, etc.
– Waste disposal by vitrification works because other metal oxides can be substituted for part
of the usual oxide content
– Production of glasses very energy intensive (need temperatures greater than 1300 "C), so
used only if must be absolutely no leaching into the environment
– Disposal method of choice for radioactive wastes from nuclear fission power plants
– Remove valuable radioisotopes such as uranium, plutonium, etc first
– The vitrified radioactive material is sealed in stainless steel drums, which are then sealed in
concrete
– The sealed drums would be buried deep (several km) in stable geological formations
b.
New Directions in Hazardous Haste Management
– Better to reduce the amount of wastes being generated
– Can be done by changing manufacturing processes or using alternative chemicals ("green
chemistry')
– Can re-use wastes within plant or find an industry that could use the waste as a raw material
or a feed stock i.e. exchange wastes
– The Canadian government wants to reduce the volume of hazardous waste from 1988 levels
by 50% by the year 2000
– The advantages of recycling hazardous wastes include:
- savings in raw material costs
- savings in time and energy
- lower treatment and disposal costs
- lower risk of legal liability
- less employee exposure
- better corporate image
– A suggested heirarchy (from best to worst) for disposal of hazardous wastes is:
1. On-site reuse or recycling
2. Waste exchange
3. Chemical treatment
4. Incineration
5. Disposal in a permitted landfill
6. Discarded into the environment
– Need to audit wastes to identify and quantify materials in order to "exchange" wastes with
other companies
THE ATMOSPHERE
A.
OVERVIEW
– Forms a thin layer of gas around the planet
– Total mass about 5 x 1018 kg; 99% of it is below 30 km. At 150,000 m, pressure is
1 x 10–9 atm
– Protects surface from ultraviolet radiation, cosmic rays and solar wind
– Source of O2, CO2 and N for respiration, photosynthesis and components of biological
molecules (eg. proteins)
– It stabilizes temperatures by absorbing and reradiating solar radiation, and transferring
heat from equatorial to higher latitudes
– Average temperature 16°C; only varied 7°C in last 150,000 yr
– Forms part of hydrological cycle
– Medium for transport of many substances
– Pollution problems do not just involve new compounds added to the atmosphere, but also
large and often local concentrations of naturally-occurring compounds which can change
the chemical balance of the atmosphere
– Local emissions of pollutants can become regional or global problems e.g. SO2 up large
stacks, or persistent pesticides moving to the Arctic
– In last 150,000 yr, P(CO2) has risen from 200 ppm to 350 ppm
– 10% of this rise has occurred in the last 30 yr
– Levels of methane are rising 1%/yr
– Levels of nitric oxide are rising 0.3%/yr
– Levels of chlorofluorocarbons (CFC) rising 5-6%/yr in period (1988-1990)
– Remember that everything is connected: land-use, atmospheric chemistry, heat balance all
influence climate, and vice versa.
B.
ORIGIN OF THE EARTH'S ATMOSPHERE
The Beginning
The Big Bang Theory
– Astronomers noticed two phenomenons when studying the movement of the stars and
galaxies which led to the Big Bang theory.
1. The universe seemed to be expanding uniformly in all directions
2. The outermost objects were moving the fastest.
– This could be explain by having the universe being created by an sudden expansion from a
single point (singularity, black hole)
– why the sudden expansion occurred is unclear but physicists have used present knowledge
about subatomic particles and an assumption about the mass of the universe to predict what
conditions must have been like in the first few fractions of a second after the expansion
began.
– Stephen Hawkings is famous physicist in this area.
– At time = 10–32 s, T = 1 x 1027 °C, any matter and antimatter exists as subatomic particles
– called quarks, leptons and electrons with each having an antimatter counterpart.
– matter and antimatter combine to destroy each other, converting to energy
– matter is in excess by 1 ppb so a universe created from matter results.
– as the universe expands, it cools.
– at time = 10–6 s, T = 1 x 1013 °C, quarks begin to combine to form protons and neutrons
– at time = 3 min, T = 1 x 109 °C, protons and neutrons begin to combine to form the atomic
nuclei of the hydrogen isotopes (deuterium and tritium) and helium
proton (1H+) + neutron –––> deuterium nucleus (2H+)
deuterium nucleus (2H+) + neutron –––> tritium nucleus (3H+)
2 deuterium nuclei (2H+) –––> helium nucleus (4He2+)
– at time = 300,000 yr, T = 1 x 104 °C, the nuclei created earlier begin to collect electrons to
become neutral atoms
– at time = 1,000,000,000 yr, T = –200 °C, slight irregularities in the early expansion result in
clouds of hydrogen and helium which begin to collapse due to gravity
– if the clouds are big enough, the heat generated by the collapse triggers a fusion reaction in
the core of the star where hydrogen atoms are being forced together to form helium.
– this is what has been occurring in our Sun for ~5 billion years and will continue for another
5 billion years.
– what happens when the hydrogen is used up?
– depends on the size of the star
Our Sun
5 x Our Sun
15 x Our Sun
white dwarf
neutron star
black hole
– in the star, gravity is trying to pull the material to the centre.
– this is balanced by the radiation from the fusion reaction trying to push material out
– when fusion quits, gravity succeeds. The core collapses.
– In a star like our Sun, the He atoms are forced together forming carbon atoms. The repulsion
of the electron clouds around the atoms is enough to halt the collapse. Considerable heat
and a shockwave are generated. The shockwave expels most of the other material in a event
called a supernova. What is left is called a white dwarf. It is extremely dense, 1 x 109 kg/m3.
Earth’s density = 5.4 x 103 kg/m3.
– in a star 5 times our Sun, the collapse continues. The electrons surrounding the atoms in the
core are forced into the nucleus where they combine with the protons to make neutrons. At
this point the collapse stops and the shockwave is generated. What is left is a neutron star.
An amount the size of a sugar cube would weigh 100 million tons.
– in larger stars, the collapse may generate enough heat to start another fusion reaction in the
core producing heavier elements . The cycle stops with iron. The final collapse results in a
black hole.
– the supernova explosion generates the heavier elements.
1.
General
a)
b)
c)
d)
present major components: N2, O2
noble gases: He, Ne, Ar, Xe, Kr
abundant variable components: CO2, H2O
others: H2 CH4, CO, NH3, N2O, H2S, (CH3)2S, SO2, HC1
– The atmosphere is a dynamic system and a balance of processes that delivers volatiles to
the atmosphere or removes it
– A source is a process that delivers a substance to the atmosphere.
– The source may be either a) the ultimate source, the source of the substance in the inventory
of the earth, or b) the immediate source, a process that sustains the concentration of the
substance in the present atmosphere
– A sink is a process that removes a substance from the atmosphere, either by:
- a chemical reaction,
- or by physical means, such as the loss of hydrogen to space at the edge of the
upper atmosphere
– Concentration controlled by difference of rates: Sources minus sinks
– For most substances, there is a constant turn-over
2.
Ultimate Sources Of Volatiles In The Early Atmosphere
a)
Capture of primordial gases from the solar nebula (gas cloud)
– As planets formed by accretion, gases would be captured by gravity
– Escape velocity for a gas on Earth now is 11,300 m/s
Vgas = 157(T/M)1/2 (m/s)
– Rule of thumb is that a planet can hold onto a gas permanently if the velocity of the gas is less
than one-sixth of the escape velocity.
Vgas < 1/6 Vescape
– At T <1400K, gases of mol. wt. > 10 would be captured but H2 (M = 2) and He (M = 4) would
be able to escape.
b)
Outgassing from the planet's crust
– Solid particles in the condensing solar nebula would absorb volatiles on their surfaces;
retention by physical and chemical absorption
– Volatiles collected would be those substances that are "sticky" i.e. those that are more polar
and reactive and would "stick" to the solid surfaces, rather than those unreactive substances
(noble gases, H2)
– If not heated and driven off on impact, volatiles would be trapped as particles accreted
– As planet heats up, would be differentiation of materials as denser materials would sink and
volatiles would rise and volatilize
– Predominant outgassed C and N species would be CO or CO2, and N2
c)
Importation from comet-like bodies late in planet's formation
– "Dirty snow-balls" of rock, ice, and other volatiles
– Volatiles would also be predominantly polar and reactive
– Planet would receive a thin veneer of volatiles enriched in those volatiles last to condense in
the solar nebula
3.
Sinks
a)
Photochemical reactions
– CH4 and NH3 would be rapidly degraded by the intense UV from young sun; estimate lifetime
of an CH4/NH3 atmosphere would be about 1 million yr.
– H2O would also degrade to reactive species such as OH
H2O + hυ ——> H• + OH•
– This also a source of O2:
H2O + hυ ——> H2 + O•
2O• –—> O2
b)
Escape
– Light species such as H2 and He would escape
c)
Solar Wind Stripping
– Since no magnetic field early in planetary development, solar wind would have directly
interacted with the upper atmosphere: "cosmic sandblasting"
4.
Sequence Of Events In Development Of The Atmosphere
Accretion of Planet
– Solids formed in the nebula which were rich in chemically-active volatiles but not large
enough to capture noble gases; the planet formed by accretion, incorporating volatile-rich
solids
Loss of Primordial Atmosphere
– Either the primordial atmosphere was lost or never existed
– Evidence: should be more 36Ar or Kr in present atmosphere than there is now
Formation of Secondary Atmosphere
– In comparing concentrations of selected components of solar nebula versus terrestrial
concentrations, note that most efficient collection (earth conc'n / nebula conc'n) is associated
with chemical reactivity i.e. volatiles formed involatile phases
– Secondary atmosphere would have formed from outgassing of volatiles and from importation
– Atmosphere has a concentration of 40Ar much higher than 36Ar; 40Ar outgassed from crust
after radioactive decay of 40K, while 36Ar would be derived from nebula
Early Composition of the Atmosphere
– Crust of the Earth likely much less oxidized 3 billion yr ago than now, eg. metallic iron
– Outgassing C, H, S-containing species would reaction and equilibrate with reduced crust to
form H2, H2O, CO and H2S
– N2 would outgas and remain unreactive
– H2 would be lost, C would be bound as carbonates in weathering reactions, H2O would
precipitate for form the oceans; therefore atmosphere would be predominately N2
Rise of Biological Activity and Biogeochemical Cycles
– Earliest preserved rocks dated 3.8 billion yr are sediments containing partially oxidized Fe2+,
organic-C, and carbonates
– Sedimentary rocks require atmospheric pressures and temperatures consistent with liquid
water
– Evidence for CO2 levels l00x present atmospheric level (PAL), leading to a "greenhouse
effect" that could keep surface temperature > 0°C even though sun still dim
– However, 3.8 billion yr old rocks are poorly preserved because of extended heating
(1,000,000 yrs at T>500°C)
– Rocks of age 3.5 billion yr old better preserved and they contain abundant organic-C and
micro-fossils
– The ratio of 13C/12C indicates that a biologically-mediated C-cycle was operating at this time
f)
Rise of Molecular O2
– Early on, there were abundant sinks of readily-oxidizable materials: ferrous oxide, sulfides,
organic compounds
– Any O2 formed would be readily bound
– Oxidatively-weathered crustal materials would be recycled and new material continuously
exposed through the geological cycle: sedimentation, burial, igneous activity, uplift,
weathering); about half of rock recycled every 600,000,000 yr
– The strength of the sink would prevent accumulation of O2 for about 100,000,000 yr from first
onset of biological O2 production
– Photochemical reactions of H2O in upper atmosphere can generate O2, but transport of H2O
to upper atmosphere hindered by a cold layer in the atmosphere
– Strength of this source likely much less than strength of available sinks
– Sea-level conc'n of O2 from photochemical source estimated to be 10–8 PAL, just before
onset of photosynthesis
– Evidence supports onset of photosynthesis between 3.5 billion yr to 2.8 billion yr ago
– By 2.8 billion yr ago, ratio of O2/CO2 about 1.3 (635 now), with O2 conc'n = 0.01 PAL, and
CO2 conc'n about 9x PAL (i.e. 3,000 ppm, 0.3 %)
– By 1.7 billion yr ago, Fe formation disappear from geologically record
– By 1.4 billion yr ago, eukaryotic organisms arose; they need O2 conc'n of 0.02 PAL
– From 0.9 billion yr to 0.6 billion yr ago, the ratio of 13C/I2C for organic-C in sediments
shows that much {CH2O} was being deposited and buried without oxidation, which leads to
higher O2 levels in the atmosphere
– Land plants developed about 0.4 billion yr ago and would have driven O2 up to present
levels, if not there already
C.
COMPOSITION AND STRUCTURE OF THE ATMOSPHERE
1.
Chemical Composition
– As seen from discussion on the Origin of the Atmosphere discussion, the atmosphere is
constantly changing on a geological time scale, and has evolved from about 1.4 billion yr ago
to present composition
– Some constituents of the atmosphere do not vary much in concentration over 1000 yr or
even 1,000,000 yr: O2 and N2
– Other constituents change over considerably shorter times: CO2, which has doubled in
concentration over last 20,000 yr.
– Some constituents change concentration over very short term: H2O, which changes
concentration rapidly over time and over distance; diurnal (day and night) changes, seasonal
changes
– The Index of Variability is a useful parameter for expressing these differences:
I.V. (unit of time) = (1/c)(Δc/Δt)
where
- c = concentration of constituent, and
- Δc/Δt = time rate of change of concentration of
the constituent
– Stable or fixed constituents have I.V. of order of 10–6 < (1/c)( Δc/Δt )] < 10–2 hr–1
– Variable constituents have I.V. of order of 10–1 < [(1/c) (Δc/Δt )] < 10 hr–1
– Another useful parameter is Residence Time defined as:
τ = (amount in the "reservoir") / (rate of inflow or outflow); units of time
– Residence time also called "lifetime": time for concentration to fall to 1/e (~37%) of initial
concentration
– Lifetime is different from the half-life (t1/2), which is the time for the concentration to fall to
50% of the original concentration
– Can also classify constituents on the basis of residence time:
a) Quasi-permanent gases: τ values longer than 1000 yr;
b) Variable gases: τ values in order of years
c) Highly variable gases: τ values in the order of days or weeks
– residence times are related to how well a substance is mixed in the environment:
- a long residence time => well mixed and widely distributed e.g. CFCs
- a short residence time => poorly mixed and regionally distributed e.g. acid rain
– Important: The longer the Residence Time, the smaller the Index of Variability
- Therefore, it is not possible to precisely define the chemical composition of the atmosphere
because that composition changes with pressure, time of day or year, location, etc.
Gas Composition of Unpolluted Air near the Earth's Surface
Gas
Formula
%Volume
ppm
τ
Quasi-Permanent
Nitrogen*
N2
78.08
~106yr
Oxygen*
O2
20.95
5 x l03yr
Argon*
Ar
0.93
Neon
Ne
18.18
Helium
He
5.24
107yr
Krypton
Kr
1.14
Xenon
Xe
0.087
Variable
Carbon Dioxide
CO2
~330
5-6 yr
Methane
CH4
1.3-1.6
4-7 yr
Hydrogen
H2
~0.5
6-8 yr
Nitrous Oxide
N2O
0.25 - 0.35
~25 yr
–2
Ozone
O3
(1 – 5) x l0
~2yr
Highly Variable
Water
H2O
(0.4-400) x 102
10 d
Carbon Monoxide
CO
0.05-0.25
0.2-0.5 yr
Nitrogen Dioxide
NO2
(0.1 – 5) x l0–3
8-10 d
–3
Ammonia
NH4
(0.l – l0) x l0
~5 d
Sulfur Dioxide
SO2
(0.03 - 30) x l0–3
~2 d
–3
Hydrogen Sulfide
H2S
(0.006 - 0.6) x l0
0.5 d
a)
Oxygen
– amount – 3.8 x 1019 mol or 1.2 x 1018 kg
– principal source is photosynthesis, which generates about 4.0 x 1014 kg/yr
– processes which consume oxygen (sinks) are:
- respiration
- decay
- combustion of fossil fuels
- weathering of rocks and ancient sediments e.g. oxidation of Fe2+ to Fe3+; note that this
was likely a major sink in the past, but now only ~0.1% as strong as other sinks today
– residence time of O2 is 3300 to 7600 yr, which is a time longer than necessary to provide
complete mixing
b.
Water
– about 7 x 1014 mol of H2O (g) in atmosphere, vs. 9.5 x 1019 mol of H2O (l) on surface –
evaporation rate from ocean: 2.2 x 1016 mol/yr
– evaporation rate from lakes and rivers: 5.5 x 1016 mol/yr
– precipitation over land: 5.5 x 1015 mol/yr
– precipitation over oceans: 1.9 x 1016 mol/yr
– average residence time about 10 days (3 x 10–2 yr) => water is poorly mixed and much local
variation
– water is not a permanent gas at atmospheric pressure
– the atmosphere becomes saturated, the level varies with temperature
– relative humidity is P(atm) as a fraction of the equilibrium value
– note that P(H2O) is non-zero even below O° C, i.e. snow can evaporate
c.
Nitrogen
– regulated principally by biological processes
– atmosphere has 39 x 1018 kg of N
– sinks: biological nitrogen fixation (2 x 1011 kg/yr)
– NO produced during thunderstorms and through combustion, leading to HNO3 in rainwater
(7 x 1010 kg of N/yr)
high temp
N2(g) + O2(g) ———————> 2NO (g)
– Industrial fixation by the Haber process ~ 5 x 1010 Kg/yr.
catalyst
N2 (g) + 3H2(g) ———————> 2NH3 (g)
– NH4+ + NO3– are cycled through the biosphere to make proteins and nucleic acids
– process of decay returns N as N2 and N2O by denitrifying bacteria
- denitrification:
- nitrification:
NO3– to N2O + N2
NH4+ to NO3–
(reduction)
(oxidation)
– almost all N fixed by Haber process is used as fertilizer
– about 1/2 of fertilizer denitrified before being taken up by plants
– this leads to increase in atmospheric N2O levels from 0.29 to 0.31 ppm over 20 years.
– N2O rather unreactive and has a lifetime of- 20 years
Carbon Dioxide
– 1.4 x 1016 mol in the atmosphere
– sources: respiration, combustion, decay
– major sink is photosynthesis (1.5 x 1015 mol/yr)
– CO2 is soluble in water, so must account for air:water exchange: uptake ~ 7 x 1015 mol/yr +
release - 6 x 1015 mol/yr
– lifetime ~ 2 years
– is a pronounced annual cycle in conc'n: peaks in April and through October
– due to photosynthetic activity in northern hemisphere, maximum photosynthesis occurs in
May to October; reverse in southern hemisphere
– CO2 conc'n also steadily rising: 315 ppm in 1958, 350 ppm present
– rising trend also seen in CH4 and N2O
– reasons controversial: out-of-balance between inflow and outflow
- emissions from fossil fuel burning increased
- are losses decreasing? deforestation can increase emissions due to biomass burning
and decrease of photosynthesis
e)
Trace Gases
Gas
CH4
CO
PPM
1.6
~0.12
N2O
NOX
0.3
10–6 - 10–4
HNO3
NH3
10–5 - 10–1
10–4 - 10–3
H2
H202
H2CO
CS2
SO2
0.5
10–4 - 10–2
10–4 - 10–3
10–5 - 10–4
2 x 10–4
Source
Biogenic
Anthropogenic, biogenic,
photochemical
Biogenic
Photochemical, lightening,
anthropogenic
Photochemical
Biogenic
Sink
Photochemical
Photochemical
Photochemical
Photochemical
Washed out by precipitation
Photochemical,
washed out by precipitation
Photochemical
Washed out by precipitation
Photochemical
Photochemical
Photochemical
Biogenic, photochemical
Photochemical
Photochemical
Anthropogenic, biogenic
Anthropogenic, photochemical,
volcanic
CC12F2
2.8 x 10–1
Anthropogenic
Photochemical
–––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––
2.
Units and Dimensions in Atmospheric Chemistry
– Need to become familiar with the units used in expressing concentrations of constituents in
the atmosphere
a)
Dimensioned units
Partial Pressure
1 std atm (at STP) = 101.35 kPa = 1.0135 bar = 760 torr = 760 mm Hg
1 bar = 100 kPa = 106 dynes/cm2 = 105 N/m2 (105 dynes = 1 N)
1 Pa = 1 kg/m/s2 = 1 N/m2
Mass Concentration (mass/volume)
Micrograms per cubic metre (µg/m3)
Molarity (moles/volume)
Moles per litre (moles/L)
Molecular Concentration (molecules/volume)
Molecules per cubic centimetre (cm–3)
– Particulate matter most commonly expressed in unit of mass concentration (µg/cm3) or
particle concentration (cm–3)
b)
Dimensionless units
– These units are usually expressed in volume percent (%)
- parts per million by volume (ppmv), parts per billion (ppbv), etc.
– Can express all these units in terms of the mole fraction Xi, where:
Xi = ni/nT
ni = # moles of species i, nT = total # of moles of gaseous species
– Common dimensionless units for gases:
Mole fraction
Mole per cent
Parts per million
Parts per billion (U.S.)
Parts per trillion (U.S.)
Xi
%
ppm
ppb
ppt
Shown above
100Xi
106Xi
109Xi
1012Xi
– If gases are behaving ideally, then relation PV =nRT holds; can find equivalent expressions
for mole fraction in terms of pressure and volume:
At constant volume, PiV = niRT => ni =Pi(V/RT), similiarly nT = PT(V/RT)
therefore ni / nT = Pi / PT i.e. ratio of partial pressure versus total pressure
At constant pressure, PVi = niRT => ni =Vi(P/RT), similiarly nT = VT(P/RT)
therefore ni / nT = Vi / VT i.e. ratio of partial volume versus total volume
3.
Structure of the Atmosphere
a.
General
– 99.999% of the mass of the atmosphere lies within 80 km of the surface
– The atmosphere comprises 1 x 10–4 % of the total mass of the Earth and atmosphere
– The composition of the atmosphere is constant only in the lower 80-100 km due to constant |
mixing; this is called the homosphere
– Above this altitude is the heterosphere where the composition of the air changes with
altitude i.e. the molecular weight of the air decreases with altitude
– Variations in composition are due to photo-dissociation of air molecules by far UV and soft
X-ray electromagnetic (EM) radiation from the Sun
– Reactions also ionize air, especially at 300-400 km altitude; this region is called the
ionosphere
– The homosphere is divided into the troposphere, the stratosphere and the mesosphere
– This division is made on the basis of the thermal structure of the atmosphere
– Our concern in this course will be on the troposphere and the stratosphere, where most of
the physical and chemical processes involved with weather and atmospheric pollution take
place
b.
Pressure Variations and the Scale Height
– The force of gravity compresses the gases in the atmosphere so that pressure decreases
with height above the surface.
– Can derive the variation of pressure with altitude:
dP/dz = -gρ;
where P is pressure, z is altitude, g is the acceleration due to
gravity, ρ is the density of the atmosphere
– Density is related to pressure:
PV = nRT
Ideal gas equation; V is volume, n = molar mass, R is gas constant
(8.314 J/K-mol), T is temperature K
P = (m/M)RT/V
m is the mass of gas, M is the molecular weight (average for air is
28.966 g/mol)
P = (m/V)RT/M
or
– Rearranging, we get
P = ρRT/M
ρ = PM/RT
– Substituting
dP/dz = g(PM/RT)
– Rearranging again
dP/P = (gM/RT)dz
– integrate, and use P0 as integration constant,
Pz = P0 exp(-gMz/RT) => Pz = P0 exp(-z/H)
where H = RT/gM = Exponential Scale Height
– The Exponential Scale Height is the change in height necessary to give a pressure drop of
1/e; also the thickness of the atmosphere if it were compressed into a layer of constant
density equal to the density and pressure at 0 km altitude: 8,434 m
– The change in pressure with altitude can be approximated by the relation:
log P = -0.06z; P in atm, z in km
c.
The Troposphere
– At standard temperature and pressure (STP) (T=273.15K, 101.3 Kpa), the density at sea
level is 1,293 x 10–3 g/cm3; density and pressure decrease with altitude
– This layer extends from 0 to -12 km altitude
– It is heated by thermal energy emitted from the surface
– Intense convection in both vertical and horizontal directions keeps everything well mixed
(there is no stratification based on density or molecular weight)
– Air cools with height about 6-10°C/Km [Lapse Rate]
– This cooling cause the formation of clouds i.e. involved with the hydrological cycle, which
affects the cycling, fate and transport of many constituents and pollutants in the atmosphere
d.
The Tropopause
– Extends 9-16 km altitude
– The temperature drops from about 15°C to ~-60°C at the top of the troposphere
– It is the boundary layer between the troposphere and the stratosphere
– Above the tropics, the layer lies 13-16 km high, being highest over the equator
– Above the temperate and polar regions, the layer lies 9-13 km high
–There is a gap between the tropical and temperate layers, in which lies the "jet streams"
–This gap and the jet streams are important for the transport of pollutants into the
stratosphere, especially the chlorofluorocarbons (CFCs) implicated in ozone depletion
e.
The Stratosphere
• It extends up to 52 km
– The temperature increases from the tropopause due to absorption of sunlight wavelengths in
the middle to near UV range by oxygen (O2) and ozone (O3):
O3 + hυ ——> O2 + O (1)
O2 + hυ ——> O + O (2)
– The energy of the photo has to be greater than the bond energy to break the bond
– E = hυ = hc/λ; i.e. a photon's energy is proportional to its frequency and inversely proportional
to its wavelength (Note: υλ= c; units: υ = Hertz, 1/s, λ = metres)
– For reaction (1) need λ ≤ 310-330 nm (nm = 10–9 m)
– For reaction (2) need λ = 190-200 nm ; this reaction is more common in the upper
atmosphere (in the "ozone layer")
– The reaction products recombine to recreate O2 and O3 and heat:
O + O ——> O2 + heat (3)
O2 + O ——> O3 + heat (4)
– The heat released is the cause of increase in T with altitude
– There is a minor heat contribution from excess energy of photon bond-breaking (reactions 1
and 2)
– Since both O2 and O3 absorb light below 330 nm, both species absorb UV radiation and heat
the atmosphere by increased kinetic energy of the molecules
f.
The Stratopause and Mesosphere
–The Stratopause starts about 52 km altitude
–At this level, the temperature starts to fall again.
–The maximum temperature is achieved here because the concentration of absorbing
molecules becomes too low and not enough UV is absorbed to heat the air; below this
altitude, while the concentration of UV-absorbing species in increasing, the intensity of UV light
is decreasing as much of it has been absorbed already at higher altitudes
– This also the start of the Mesosphere
– The atmospheric pressure is very low here; < 0.0005 atm
– Very short wavelength light can cause ionization and dissociate N2, among other species:
O2 + hυ ——> O2+ + e–
O2 + hυ ——> O + O
N + O + M ——> NO+ + e–
NO + hυ ——> NO+ + e–
– This region of increased concentration of ions and free electrons is called the ionosphere; it
is also the lower part of the thermosphere
– The Aurora Borealis occurs here at altitudes between 80 – 110 km
g.
Mesopause and Thermosphere
– At 90 km altitude, the temperature is -95 to -100°C, and starts to rise again due to
photo-dissociation and ionization reactions with oxygen;
134-176 nm
O2 + hυ —————— > O• + O• + ionization reactions + heat
– In the thermosphere, the "temperature" can be quite high e.g. 1,500K, but a thermometer (or
you) suspended there would feel very cold
– This is because the "mean free path" of the molecules is very high, 10° to 104 m
– Atoms and molecules do not reach thermal equilibrium with each other, so their
"temperature" is a function of their average kinetic energies (and velocities)
mean free path (m): l = 1/[(√2)πρ2n•)
where:
- ρ is the molecular diameter in metres (do not confuse with density, which
unfortunately uses the same symbol)
- average molecular diameter = 0.17 µm
- average atomic diameter 75 pm (0.000075 µm)
n• is the no. of molecules per unit volume (calculate from ideal gas equation)
average velocities (m/s): v = [3RT/M]
where:
R is the gas constant 8.314 kg m2 / (s2 mol K)
T is temperature K
M is molar mass (kg/mol)
Table: Overview of Atmospheric Layers
Layer
Troposphere
Stratosphere
Mesosphere
Thermosphere
Height (km)
0 to 12
12 to 52
52 to 90
90 +
T(°C)
15 to –60
–60 to 0
0 to –100
–100 to 300+
Main Species
N2, O2, H2O, Ar, CO2
N2, O2, O3
N2, O2, O2+, NO+
N2, O2+, O+, O, NO+, O2
D.
PHYSICAL CHEMISTRY BACKGROUND FOR ATMOSPHERIC CHEMISTRY:
KINETICS AND PHOTOCHEMISTRY
– We need to understand the basics of chemical kinetics and photochemistry to understand the
chemical processes going on in the atmosphere (and water)
a.
Chemical Kinetics
– The gases in the atmosphere, especially trace gases and pollutants, are not in a state of
equilibrium, as predicted by thermodynamics.
– The apparent stability of the composition of the atmosphere is due to its being in "steady
state" in which there is a rough balance between sources and sinks
– The chemistry of the atmosphere is governed more by gas phase reaction rates than by
thermodynamic equilibria
– Consider the hypothetical reaction:
aA + bB —— > cC + dD
vf = kf [A]a[B]b
vr = kr [C]c[D]d
– at equilibrium, vf= vr
K' = kf/kr = [C]c[D]d / [A]a[B]b;
– for example,
NO + 1/2O2 <===> NO2
Keq = P(NO2)/{P(NO)} {P(O2)}1/2 = 1.38 x 10–6 atm–1/2 at 25 °C
at P(O2) = 0.21 atm (sea level oxygen partial pressure)
P(NO2)/P(NO) = 6.3 x !05
– This result implies that the ratio of NO2/NO molecules approaches 1 million
– This is rarely the case. Why?
– Other reactions can prevent equilibrium from being reached
– Forward and reverse reaction rates are very slow, so it takes a very long time for equilibrium
to be reached
– How fast is NO converted to NO2?
2NO + O2 –––> 2NO2
rate (NO2)= Krate [NO]2 [O2]
Krate = 14.8 x 103 L2/mol2 s
– Convert O2 partial pressure to moles/litre:
Gases occupy 24.5 L/mole @ 25°C
1 mole/24.5 L x 21% O2 = 8.6 x 10–3 mol/L
– The concentration of NO2 is usually 0.1 ppm, so:
1 mole/24.5 L x 0.1 x 10–6 = 4.1 x 10–9 mol/L
–
Therefore
rate = 14.8 x 103 x (4.1 x 10–9)2 x (8.6 x 10–3) = 2.1 x l0–15mol/L-s
– Convenient to convert back to ppm:
rate = 2.1 x 10–15 mol/L-s x 24.5 L/mole x 106 ppm x 60 s/min x 60 min/hr
= 1.8 x l0–4 ppm/hour
– Note that the maximum rate of conversion of NO to NO2 is very slow, and that the conversion
rate decreases with the decrease in concentration of NO
– The slow reaction rate explains why the prediction of P(NO2)/P(NO) is not observed
– Equilibrium is seldom reached in the atmosphere due to slow chemical kinetics
– The most important reactions tend to be the fastest reactions
i)
Effect of Temperature on Reaction Rates
– Reaction rates depend on temperature:
Arrhenius equation K = A exp(–Ea/RT)
where: A is the pre-exponential factor (effects of collision frequency and steric effects),
Ea is the activation energy,
R, T have usual definitions
– If the activation energy is too high, the reaction rate may be very slow, even if it is
thermodynamically favoured
– One can also get one set of products (Ea, P) at a lower temperature, and another set (Ea', P')
at a higher temperature
Ea
R ––––> P
low T
ii)
Ea'
R ––––> P'
high T
Catalysis
– A catalyst is a substance that alters the rate of a chemical reaction without being consumed
by the reaction
– The catalyst only changes the rate of the reaction, not the equilibrium of the reaction, since it
changes the rates of the forward and reverse reactions equally
– Catalysts change the rate of a reaction by changing the mechanism of the reaction, allowing
progression via a transition state of lower energy (Ea)
Example: Destruction of ozone
O• + O3 ——> 2O2
- This reaction is slow
O• + NO2 ––––> NO• + O2
O3 + NO• ––––> NO2 + O2
O• + O3 —–—> 2O2
The sum of the reactions gives the same results as
above. Note that NO is consumed in the first reaction
but regenerated by the second reaction. The overall
rate is faster than the single step reaction.
– Atmospheric chemistry usually dominated by kinetics and kinetic reactions
example: N2 + O2 ——> 2NO
N2 + 2O2 ——> NO2
- occurs at high temperatures during combustion
- occurs at room temperature
Note that NO formed in the first reaction will slowly convert to NO2
Figure 2.1
b.
Energy changes associated with chemical reaction processes.
Photochemistry
– Very important to processes occurring in the atmosphere
– Reactions driven by energy input in the form of light, rather than by kinetic energy (heat)
i)
Principles of Photochemistry
1. Only light that is absorbed by a substance is effective in producing a photochemical
change
2. Each quantum of radiation absorbed by a molecule activates one molecule in the primary
step of a photochemical process
– Energy associated with light:
E = hν = hc/λ,
h = Planck's constant (6.626 x 10–34 J s)
ν = frequency (Hz, s–1)
λ = wavelength (m)
c = speed of light (2.997 x 108 m/s)
I/ λ = wavenumber (cm–1)
where
= (6.626 x 10–34 J s)(2.997 x 108 m/s)/ λ (m)
= 1.986 x 10–25 J m/ λ (m)
E (mole–1) = 1.986 x 10–25 J m x 6.022 x 1023 photons/mole / λ (m)
= 1.19 x 10–1J m mol–1/ λ (m)
E (mole–1) = 1.19 x 10–1 J m mol–1 x (1 kJ/103 J) x (109 nm/ m) / λ (m)
= 1.19 x 105 kJ nm mol–1 / λ (nm)
– The amount of heat (enthalpy) absorbed by a molecule or atom depends on the energy of
the photon and the work done on the surroundings (if the number of moles of particles
changes):
ΔH° = ΔUphoton + ΔnRT
(ΔnRT accounts for PV-work done on the surroundings
as the number of moles of particles changes)
ΔUphoton = ΔH° – ΔnRT
– an activated molecule may become deactivated by collision before a reaction occurs
– an activated molecule may set off a chain reaction
Quantum yield Φ = moles of substance reacted
moles of light absorbed
ii)
Example: Dissociation of O2
hν
O2 (g) ——> 2O• (g)
ΔH°298 = 494 kJ/mol
(homolytic cleavage of O2 into ground state O• atoms)
– 300 nm photons will not cause photo-dissociation
– In the stratosphere, light of λ < 242 nm will dissociate O2, leading to:
O2 + O• ——> O3
ozone formation
– Also possible to absorb energy at a level below dissociation energy:
SO2 + hν (λ = ~384 nm) —— > SO2*
* excited state
– The excess energy can accelerate reactions:
SO2* + O2 ——> SO3 + O•
ii)
Example: NO/NO2/O3 system
– One of the more important reactions in the lower atmosphere
– Photolysis of NO2 occurs at λ < 435 nm
NO2 + hν ——> NO + O•
(produces a ground state O atom)
rate = –d[NO2]/dt = ΦΙabs
– Φ falls off as λ increases
– Ιabs is intensity of absorbed light, units of moles of photons/L - s
– Ιabs also a function of [NO2]; at low concentrations, little light is absorbed and little reaction
occurs
Ιabs = Ka [NO2]
Ka is the wavelength dependent absorption constant for NO2
rate = –d[NO2]/dt = ΦKa [NO2] = J [NO2]
- J = ΦKa, is a first-order rate constant
- J also depends on intensity of incident light, so that:
J = 0 at night
J = ~20 h–1 hi bright noon sun
– After generation of O•, next reaction is usually
O• + O2 ——> O3
rate = K2 [O•][O2]
Another reaction completes the cycle:
O3 + NO ——> NO2 + O2
rate = K3 [O3][NO]
– The lower atmosphere contains a low background level of O3 (0.03 to 0.05 ppm)
– The photochemical conversion of NO to NO2 is more rapid than the thermal reaction of NO
with O2 discussed earlier
– The reactions are rapid and they establish a photostationary equilibrium state, which governs
the NO2/NO ratio in air
– NO2 is formed as rapidly as it is formed,
J [NO2] = K3 [O3][NO] –––> [NO2]/[NO] = K3 [O3]/J
– This implies that the ratio is governed by [O3] and J (light intensity)
- at night in excess O3, the ratio is large
- in bright sunshine, ratio is small
– The predictions of this model are borne out by observations and serve to demonstrate that
kinetics can have a much greater controlling effect on concentrations than thermodynamics
E.
THE GREENHOUSE EFFECT AND CLIMATE CHANGE
1.
Introduction
– The Earth intercepts about 340 W/m2 of solar radiation (see figure below)
– Of this amount, about 30% is reflected by the atmosphere and 70% is emitted as infrared
radiation
– However, the Earth's surface receives only 45%, while emitting 133% of what reaches the
surface
– Where does this extra energy come from? From the Greenhouse Effect (GHE)
– The atmosphere traps heat near the surface by water, CO2 and other trace gases that are
transparent to the incoming solar radiation but absorb infrared (IR) radiation emitted from the
surface (about 92% of IR emitted by surface is absorbed by these gases)
– The warmer atmosphere re-emits IR back to the surface
– The GHE results in the Earth appearing to radiate energy equivalent to a black body at 255K,
while receiving energy from the sun at 6000K (see figure below)
– The atmosphere is largely transparent at the wavelengths corresponding to the maximum of
the Sun's emission spectrum, corresponding to visible light (400 to 700 nm).
– Oxygen and ozone absorb at wavelengths below 300 nm while water absorbs radiation of
wavelengths greater than 20,000 nm
– Except for the window from 7,000 to 13,000 nm, H2Oand CO2 absorb most incoming
radiation from 9,000 to 20,000 nm
– Without the GHE, the average surface temperature would be about -30°C rather than +15°C
– The warming effect caused by absorption and re-radiation is called radiative forcing, and is
a measure of the amount of warming a greenhouse gas provides
– Why be concerned if GHE makes the surface of the planet livable for us?
- The concern is about predicted increase in temperature due to energy trapping by increasing
levels of these gases
– Why is there so much interest in global warming issue?
1. Some modelers claim that first signal of global warming (a statistical trend) has been
detected
2. The rapid development of the ozone hole and its link to CFCs showed that man-made
alterations to the atmosphere can have serious and unexpected consequences
3. Concentrations of other greenhouse gases are also rising; these gases will soon produce a
degree of warming equivalent to doubling the concentration of CO2
4. Possibilities of catastrophic effects in worst case scenarios
– The big questions are: How much, when, and where will warming take place
– Interest waxes and wanes depending on the most recent climate conditions in the region
1.
Greenhouse Gases
– These gases absorb IR radiation
– N2, O2, Ar do not absorb IR radiation
– The "natural" greenhouse gases (GHGs) are: CO2, H2O, CH4, N2O, O3
– Note that mankind contributes to increased concentrations of these "natural" gases
– Chlorofluorocarbons (CFCs) contribute to the GHE by absorbing radiation in the 7,000 to
13,000 nm band transparent to CO2 and H2O.
– While methane, nitrous oxide, ozone and fluorocarbons are at concentrations much less than
CO2 or H2O, they absorb IR at wavelengths not covered by CO2 or H2O, are more efficient
than H2O or CO2, and are increasing relatively more rapidly than CO2 or H2O.
Summary of Greenhouse Gas Properties and Effects
CO2
CH4
N2O
CFC-11
CFC-12
Concentration
ppb(1990)
Trend
%/yr
Lifetime
yr
353,000
1,720
310
0.28
0.48
0.4
1
0.3
4
4
100-250
8
65
120
Relative
Contribution
Strength (CO2=1) 1980-1990
1
20
270
3400
7100
55%
11%
6%
12%*
* CFC-11 and CFC-12 combined
a.
Carbon dioxide
– Accounts for about one-half of human contribution to GHE
– Current P(CO2) 350 -360ppm, has been increasing last 30 years
– 25% increase since the beginning of the Industrial Revolution
– Rate of increase ~ 0.5%/yr, since source and sink not in balance
–Sources of CO2 include respiration, fossil fuel burning and biomass (forest) burning
– Oceans are a sink for CO2; 60x more CO2 (aq) in ocean than CO2 (g) in atmosphere
– Equilibrium, reactions for CO2 (g):
CO2 (g) ——> H2CO3 (aq) ——> HCO3– (aq) ——> CO32– (aq) ——> CaCO3 (s)
– Uptake of CO2 by ocean's surface relatively slow (t1/2 - 1.3 years)
– Surface waters (0 -100 m depth) do not mix well with deeper waters (t1/2 – 35 years)
– In medium term surface can only remove a fraction of increase of CO2.
– Has been a suggestion that CO2 from fossil fuel can be injected into the deep ocean
– The effect on and by the carbon cycle is a large unknown in climate predictions
b.
Methane
– Accounts for about 18% of human contribution to GHE
– About 20x more effective GHG than CO2
– Concentration in atmosphere -1.7 ppm and increasing 1 - 2% /year
– Concentration has doubled over the last 300 yr; from ice cores, was ~ 0.7 ppm until -200
years ago
– Sources: anaerobic decay of vegetation (wetlands, rice paddies, landfill sites; 33% total),
cattle, wood burning, natural gas
– Atmospheric lifetimes -10 years; due to 10% excess of sources over sinks (?)
– Removal of methane from atmosphere:
H abstraction by hydroxyl radical:
CH4 + OH• ——> CH3• + H2O
– [OH•] may be decreasing due to increase in carbon monoxide:
CO + OH• ——> CO2 + H•
– Could get positive feedback from melting of arctic permafrost to release CH4 from clathrate
hydrates:
CH4•6H2O (c) ——> CH4(g) + 6 H2O(l)
c.
Nitrous oxide, N2O
(N=N=O)
– Accounts for 6% of human contribution to GHE
– [N2O] ~ 300 ppb, increasing 0.25%/year
– Increased 5-10% over the past 200 years
– 200x more effective than CO2
– Chief source of N2O is biological denitrification
– No known sink in troposphere; lifetime 120 years
– Diffuses into stratosphere where it photo-decomposes or reacts with excited state O atoms
d.
Chlorofluorocarbons (CFCs)
– Accounts for 14% human contribution to GHE
– First produced in 1930's
– 1000x or more effective as GHE gas as CO2
– Lifetime of few 100 years
– CFCs mainly of concern due to stratosphere's ozone destruction, but also important to global
warming
– No tropospheric sinks
– Major commercial CFCs: CFCl3, CF2Cl2, CHFCl2 which were growing ~ 6% up to 1984.
– See table pg 19 Bunce on CFCs.
e.
Ozone and Carbon Monoxide
– Usually very low concentration in an unpolluted atmosphere (troposphere)
– Ozone is one of the more active greenhouse gases
– Total Radiative Forcing: est. 2.5 W/m2
CO2: 60%; CH4: 15%-20%; O3, N2O, CFCs: 20%-25%
2.
Predicting Climate Change
–There appears to be a close correlation between rising CO2 levels and rising temperature
(energy)
– Other greenhouse gases combined roughly equal effect of CO2
e.g.: CO2-50%; CH4 - 20%; CFCs - 20%, N2O - 5%
– Predictions based on complex computer models (global circulation models, GSMs) which
are estimates only
– Based on our present understanding of relevant processes and on certain assumptions
about air and water circulation patterns
– Also based on knowledge of reservoirs, sources and sinks, effects of clouds, mixing vertically
and horizontally (e.g. latitude)
– Current estimates give rise 1° - 3° over next 50-100 years
– Depends upon p (CO2) rising to 600 -1000 ppm (amount depends on fossil fuel use)
– There are many uncertainties in the climate models:
1. Effect of clouds
2. Ocean heat capacity
3. Deep circulation in the ocean ("Atlantic Conveyor")
4. Microscopic life
5. Variability of solar energy output
6. Biogeochemical carbon cycle (CO2 and CH4)
7. Relationship between temperature and methane releases by permafrost
8. Effect of sulphate aerosols
9. Dimethylsulphide releases by phytoplankton
– Controversy about whether warming trend has started yet.
– Other long-term atmospheric cycles and events such as volcanic eruptions (Mount Pinatubo)
which inject large amounts of SO2, creating stratospheric sulphate aerosols
– Dimethylsulphide emissions from ocean plankton also a source of sulphate aerosols via
oxidation of DMS in the troposphere
– Some suggest 50% reduction in CO2 emissions are required by industrial countries
– Some studies suggest that CO2 increases may be a consequence of warming, rather than a
cause
– Correlation of CO2 and warming may be influenced by where measurements made:
– Some areas may be cooling while others are warming; some temperature monitoring stations
are near large population centres where cities slowly encroach on station, increasing
temperature.
3.
The Consequences of Climate Change
a)
Environmental Effects
– The production of biomass could double for each 10°C increase in temperature, assuming
adequate moisture and other nutrients
– The poleward limit of agriculture, permafrost and sea ice would move north as temperature
increases
– The start and end of growing seasons and freeze-up and break-up of ice in rivers and lakes
will change
– A warmer ocean will expand and ice will melt
– Warming which occurred from 1900 to 1940 is thought to have caused:
- a new cod fishery off western Greenland
- 1930's drought and crop failures in North America
– Droughts and heat waves can lead to:
- reduced volumes of surface waters
- concentrated waters salts and pollution
- increased algae blooms famine and water shortages
–Climate change can make ecosystems more vulnerable to toxins, acid rain, ozone damage
and UV radiation
– Different regions of the world have different perspectives on the threat
– North America expects less rainfall in American Mid-west and Canadian prairies grainbelts;
more rainfall in American Southwest, agricultural region in Canada would move north.
– Other: less rainfall in drier regions of North Africa, Northern India, coastal flooding in coastal
areas like Bangladesh.
–Effects predicted for Canada:
- coastal inundations
- reduced ice hazards and year-round shipping in NW Arctic Ocean
- growing season length 1 month longer in Yukon
- increased forest fires, insect pests and disease in N. BC.
- increased drought frequency and shortage of irrigation water in S. prairies
- more crop growth in S. prairies due to increase in CO2
- agricultural impacts in Ontario: large yields, greater risk, moisture stress
- negative impact on skiing industry
- losses to hydro power and shipping
- lower space heating costs
- changes in fish migrations in Atlantic
- more icebergs from north into Atlantic
b)
Human Health Effects
– Changes in the food supply:
- force relocation of people and disrupt patterns of food supply and distribution
- use of more pesticides and fertilisers to maintain food supplies may have negative
health effects
– Spread of disease:
- insect spread can spread diseases
- more tropical diseases into temperate areas
– Water shortages:
- salinization
- ground water shortages
– Rise in ocean levels
- salinization
- release of contaminants from sludge dredged to maintain navigation channels
- flooding of sanitation systems, garbage dumps and water treatment facilities
F.
STRATOSPHERIC OZONE
1.
Ozone Layer
– Ozone good absorber at 240 - 320 nm
– Ozone not a major constituent, even in the 'ozone layer'
– Max [O3] ~10 ppm at ~35K m high
– If “compressed” at STP, layer would be ~ 3 mm thick
– Dobson units: 0.01 mm of O3 at STP
2.
Formation and Destruction
– O2 converted to O3 and back in steady-state
– Ozone depletion due to pollutants that increase rate of destruction
– Chapman reaction for formation and depletion of ozone:
ΔH kJ/mol
λ<240nm
(1)
O2 ——————> 2O
495 -E (photon)
(2)
O + O2 —————> O3
–105
(3)
λ<325 nm
O3 —————> O2 + O
105 -E (photon)
(4)
O + O3 ————> 2O2
–389
– O has very short lifetime since it is very reactive
– Reactions stop at night
– Reactions absorb both at 240 and 325 nm.
– Reactions convert photon energy to heat, accounting for the higher temperature of this
region of stratosphere
– Light moves [O3]/ [O2] ratio away from chemical equilibrium i.e. light pumps O2 to O3, forming
O3 that rises in concentration until destruction rate matches formation rate.
– Chapman scheme not complete because it predicts 2 - 3x higher ozone levels than observed
– This means that there are additional sinks
– Discovered new catalytic processes:
General forms:
(A)
X + O3 —————> XO + O2
(B)
XO + O —————> X + O2
– Added together, reactions (A) and (B) give the same results as reaction (4):
4)
O + O3 —————> 2O2
– This sequence increases the rate of (4) if X is present
– 4 catalytic cycles have been discovered; X is an odd - electron species
– Example with Cl = X:
(5)
Cl + O3 —————> CIO + O2
(6)
CIO + O —————> Cl + O2
(Cl is regenerated in this series)
(7)
Cl + HO2 —————> HCl + O2
(termination step)
– Example with NO:
NO + O3 ––––––> NO2 + O2
O• + NO2 –––––––> NO + O2
– Other catalytic reactions have X= H, OH for cycles with
NO/NO2 cycle > uncatalyzed reactions = Cl/ClO cycle > OH/HO2 cycle > H/OH cycle
– Concentrations of reactive intermediates have been measured to give rates for various
cycles
– For all cycles, reaction (B) is rate-limiting.
– H/OH cycle not important at 30 Km because
(8)
H + O2 ––––––> HO2
is very fast
– is of some significance at higher altitudes where [O2] and [M] are smaller, and reaction with
O3 can compete more completely.
a)
Additional reactions
i) temporary reservoirs of active species
ii) interaction between temporary cycles
iii) null cycles
iv) initiation and termination cycles
i)
temporary reservoirs
NOX
NO2 + OH ———> HNO3
– HNO3 inactive but can be converted back to NO2 by photolysis
Cl
Cl + CH4 ———> HCl + CH3
– HCl reacts with OH to regenerate Cl
HCl + OH ———> Cl + H2O
ii)
interaction between cycles
– catalyst pairs X + XO can react cross-cycle e.g. across NO/NO2and OH/HO2 cycles:
NO + HO2 ———> NO2 + OH
iii)
null cycles
– Photon energy converted to heat with no net change in chemicals
hν, λ < 400 nm
NO2 ————————> NO + O
M
O + O2 ————> O3
O3 + NO ————> NO2 + O2
– All the reactants and products cancel: no net reaction
– This reaction only converts light into heat
iv)
initiation and termination reactions
– Initiation: generation of free radicals from non-radical precursors
- Initiation rates depend on intensity of sunlight
– Termination: conversion of free radicals to non-radical products
Initiation reactions:
(14)
hν
HNO3 ————> NO2 + OH•
(15)
hν
NO2 ————> NO + O•
(16)
hν
O3 ————> O2 + O•
(17)
hν
CH3C1 ———— > CH3• + Cl•
(18)
O• + H2O ————> 2OH•
(19)
N2O + O• ————> 2NO•
Reactions (18) and (19) are not photochemical themselves, but the O atoms are generated by
the photochemical dissociation of ozone to oxygen molecules and O• atoms
Termination reactions:
2ClO• ————> ClOOCl
2HO2• ————> H2O2 + O2
NO2 + Cl• –––––––> NO2Cl
3.
Chlorofluorocarbons
– Discovered in atmosphere in 1970's
– Tropospheric levels rose from 50 pptv in 70's to 270 pptv in 1993
– In troposphere CFCs are unreactive and are important greenhouse gases
– Migrate to stratosphere with half-life of 3 -10 years.
– Can undergo photolysis in upper atmosphere
λ<250nm
CF2Cl2 ——————> CF2Cl• + Cl•
– Cl radical participates in ozone depletion in following reactions:
Cl• + O3 ————> CIO• + O2
ClO + O• –————> Cl• + O2
– These reactions occur naturally as well, therefore, CFCs increases strength of existing sink
– CFCs important because:
- increase sink rate without increasing the source rate
- CFCs are persistent (lifetimes ~ 100 years)
– Each CFC has different ozone depletion potential (ODP). They are compared to the ODP of
CFCl3 (CFC-11)
Compound
CFCl3
CF2Cl2
CHF2Cl
CF3Cl, .CFCl2
4.
(CFC No.)
Lifetime(yr)
%Increase
ODP
(11)
(12)
(22)
(113)
70
110
25
90
6
6
>10
—
1
0.86
0.05
0.8
ΔO3(%)
2.0
2.1
0.03
0.5
Brominated CFC analogs (Halons)
– These compounds used as fire suppression chemicals
– C-Br bond weaker than C-C1 bond and are more easily photolytically cleaved => greater
ODP than CFCs.
– Methylbromide CH3Br being phased out along with CFCs
– CF3I is possible substitute for Br-Halons since will decompose in troposphere.
5.
Southern Polar Ozone Holes
– No ozone generated during long polar night
– At end of winter, crystals of H2O ice and HNO3 • 3H20 form
– Generated Cl2 and HOCl cleaved by light after polar sunrise to generate Cl
– Ozone broken down by Cl
(5)
(6)
HCl + ClONO2 ————> Cl2 + HNO3
H2O + ClONO2 ————> HOCl + HNO3
HOCl + HC1 ————> Cl2 + H2O
Cl + O3 ——————> CIO + O2
CIO + O ——————> Cl + O2
– Ozone restored by two factors:
1. greater sunlight intensity creates more ozone.
2. ice crystals sublime as temperature increases, releases HNO3 (g), which leads to
NO2; NO2 reacts with CIO to form C1ONO2, a temporary reservoir for CIO; a dimer
ClOOCl is also a temporary reservoir, this dimer cleaves easily to form O2 + 2Cl.
– Ozone depletion also seen over arctic, but infusion of air from lower latitudes helps prevent
more ozone loss.
– Some loss of troposphere ozone seen in arctic
– May be due to bromine chemistry
4.
Consequences of Ozone Depletion
– UV spectral ranges:
400 - 700 nm
320 - 400 nm
290 - 320 nm
<290nm
visible
UV-A
UV-B
UV-C
– Climate could change because the temperature of stratosphere would decrease, leading to a
lower altitude for the tropopause, leading to possible changes in atmospheric circulation
– There could be biological effects from less efficient filtering of UV-B
– Many possible changes in microorganisms, plants and animals.
– Possible increases in human skin cancer (5% decrease in O3 = ~20% more cancers)
5.
Montreal Protocol
– 1974 - CFCs identified as a threat to ozone
– 1978 - banned aerosols with CFCs
– 1987 - Montreal protocol to ban use of substances that deplete the ozone layer
– targets: to mid -1986 levels by mid -1989 to 80% of baseline by 1993 to 50% of baseline by
1998
– was strengthened to phase out CFC-11 and CFC-12 by 1996, along with CCl4 and CH3CCl3
– Were revisions of Montreal Protocol in 1990 & 1992
6.
CFC replacement compounds
– Want compounds that are volatile and have low toxicity
– Replacements should be more reactive in troposphere so that they do not migrate to
stratosphere; should contain few or no chlorine atoms.
– Chief candidates are partially fluorinated hydrocarbons with minimal or zero chlorine atoms
– H-atoms in molecules confer reactivity because C–H is a site of OH attack
CF3CH2F + OH• —————> H2O + CF3CHF•
CF3CHF• + O2 –> –> –> CO2, HF, H2O
– Foams: need blowing agents of lower flammability
– Replacements include: HCFC-22 (CHF2Cl) + HCFC-146 (CH3CFCl2, b.p. 32°C) with ODP of
0.12 relative to CFC-11.
– Refrigerants: need low boiling point liquids:
HFC-131a (CF3CH2F boiling point - 26°C),
HFC-152a (CF3CHF2 boiling point - 25°C)
– As replacements for CFC-12
– 152a is flammable but is compatible with lubricants used in pre-1993 equipment.
– 134a now used extensively in new equipment.
– After 1995, only recycled CFC used
– Cleaning solvents: low surface tension and low viscosity
- CFC-113 formally used
- tried mix HCFC-22 +methanol
– Aerosols: had used CFC-11 or CFC-12
- as replacement, can use mix of isobutane and methylene chloride, or mix of diethyl
ether and water.
Nitrogen Oxides N20
– N2O migrates to stratosphere where it degrades photochemically:
hν
N2O ————> N2 + O*
– NO can also react further with O*
N2O + O* ––––––> N2 + O2
N2O + O*————> 2NO
b.
NO and NO 2
– Remember that NO and NO2 can decompose ozone catalytically:
NO + O3 ———> NO2 + O2
NO2 + O ——— > NO + O2
– Compounds are naturally present in the atmosphere, and human activities increase
concentration by combustion, use of nitrogeneous fertilisers, and SSTs and space shuttles.
N2 + O2 ————> 2NO
ΔH° = +180 kJ/mol
– Note that atmospheric nuclear tests depleted ozone by injecting NOX directly into the
atmosphere
– Reaction above is endothermic so its equilibrium constant increases with temperature.
G.
TROPOSPHERIC CHEMISTRY
– Tropospheric chemistry is very much the chemistry of reactive radicals species, especially
hydroxy (OH•), peroxy (HO2•, RO2•) and nitrate (NO3•)
– These radicals are involved in the formation of photochemical smog, ground level ozone and
acid rain
– Discussions in section will also include organic pollutants, particles and aerosols, and acidic
precipitation (acid rain)
1.
Hydroxyl Radical
– Centre of tropospheric chemistry
– Has no charge, not same as OH–
– Highly reactive, short half-life
– Continually created and consumed, so there is a low, steady-state concentration
a.
Characteristic reactions
i) Abstraction of H- atom from substrate:
(1)
OH• (g) + CH4 (g) ———> CH3• (g) + H2O (g)
This is the preferred reaction for most substrates with hydrogen, since the O-H bond is
stronger than the C-H bond.
ii) Addition to an unsaturated centre:
(2)
OH• (g) + NO2 (g) ———> HNO3 (g)
This reaction is preferred for unsaturated organics (e.g. benzene), since Csp2-H bonds
are stronger than O-H bonds.
– OH• does not generally react with CFCs (This is why CFCs are inert in lower atmosphere).
– See the rate constants for reactions of various compounds with OH• on p. 68 of Bunce
– Note: Since OH. has an unpaired electron, it reacts with even electron substances to give
another free radical, as in reaction (1) above, which generates a methyl radical CH3•
– Radicals such as methyl are very reactive towards oxygen:
(4)
CH3
•
M
+ O2 ———> CH3OO•
– Reaction sequence (1) then (4) replaces a C-H bond with a C-O bond, which can lead to
complete oxidation of C-atoms in the atmosphere
b.
Formation of OH• radical in the Troposphere
- The major route is a multi-step reaction powered by light:
i)
(5)
hν, λ<400 nm
NO2 ————————> NO + O
(6)
O + O2 –––––> O3
(7)
hν, λ<320 nm
O3 ––––––––> O2* + O*
(8)
O* + H2O –––––> 2HO•
(excited state species)
Reaction (5)
– NO2 is natural constituent of atmosphere
– NO2 and NO referred to as NOX since they interchange
– Nitric oxide formed in heated air by the reaction:
(12)
ΔH = +180 kJ/mol
N2 + O2 –––––> 2NO
- Equilibrium [NO] increases with temperature
– Lightning and combustion form NO
– Reverse of (12) is slow at normal temperature, so NO formed gets "locked in" i.e. is limited
by kinetics
– Nitric oxide oxidized by atmospheric O2 to NO2:
2NO + O2 –––––> 2NO2
rate = K [NO]2 [O2] (3rd order)
– Slow rate of reaction due to low [NO] (<1 ppm)
– Oxidation of NO can take other routes:
(13)
NO + O3 —————> NO2 + O2
(14)
NO + HO2• (or RO2•) –––––>NO2 + OH• (OR•)
– NO2 is a brown gas, absorbs in visible parts of the spectrum
– NO bond cleavage only when λ<400 nm.
ii)
Reactions 6 and 7
– The formation of O3 (reaction 6) is faster in the troposphere than in the stratosphere since
the concentrations of O2 and M are higher in the troposphere
– The photochemical cleavage of O3 yields O2 and O either in ground state or the excited state
– The cleavage requires light of X < 300 nm (the UV-B region)
iii)
Reaction 8
– Only a small fraction of O* goes to OH* because O* is rapidly deactivated by collisions with
other air molecules, generating heat:
O* + M ———> O + M + kinetic energy
c)
Tropospheric Concentration of OH•
– Average concentration (globally, seasonally, diurnally) is 8 x 105 molecules/cm3
(3 x 10–5 ppb)
– Average not very useful if component reactive and has a short lifetime.
– Formation depends on flux of solar photons (time of day, season, latitude) and concentration
of NO2.
– Middle day [OH]: 2.5 - 25 x 105 molecules/cm3 in rural areas
– [OH] > 1 x 107 molecules/cm3 under highly polluted conditions
– Need sensitive analytical methods to monitor temporal variations
– Measurements show that [OH] in rural areas follows closely the intensity of sunlight
– Steady state concentration of OH can be obtained if OH formation controlled by Reaction 7
O3 –––––> O2* + O*
rate of formation = rate of disappearance
rate of disappearance = Σ k [OH•] [Substrates]
[OH•] = rate of photolysis
Σ k [OH•] [Substrates]
d)
Oxidation Reactions of OH•
– chief substrates are CO and CH4
i)
Carbon monoxide (CO)
– Natural concentration ~ 0. 1 ppmv; from biological processes, incomplete combustion,
- intermediate in tropospheric oxidation of substances such as methane.
– City air can have higher concentrations 2-20 ppmv (peak 100 ppmv)
– Can reduce CO emissions by burning oxygenated fuels (gasohol).
– tropospheric sinks:
- uptake by soils
- microbial oxidation to CO2
- atmospheric oxidation by OH•
– Last one is only known atmospheric sink for CO
(16)
OH• + CO –——> H + CO2
(17)
H + O2 ——––> HO2•
k = 2.7 x 10–13 cm3/molecules - s at 300°K
– Reaction 16 may actually be a 2 - step process:
(16a) OH• + CO –——> HOCO•
(16b) HOCO•——> H + CO2
– in the unpolluted troposphere, 70% of all OH• disappears through reactions with CO; most of
remainder with methane.
ii)
Methane
– reactions are very complex in getting to CO2 and H2O.
– generally: successive attack on the C-H bonds of methane, and their replacement by CO
bonds, or the elimination of H2O.
(1)
OH• + CH4 ––——> CH3• + H2O
(4)
(18)
CH3• + O2 —––—> CH3OO•
CH3OO• + NO —––—> CH3ONO2 or CH3O• + NO2
(19)
CH3O• + O2 —–—> CH2O (formaldehyde) + HO2•
(20)
CH2O + OH• —–—> H2O + HCO•
(direct photolysis of CH2O to H and HCO also occurs)
(21)
HCO• + O2 —–—> CO + HO2•
(22)
OH• + CO ——> H• + CO2
–addition at unsaturated carbons (alkenes and aromatics) also occur
–modelling tropospheric chemistry very difficult since reactive intermediates are involved in
many different reactions.
2.
Photochemical Smog
– Smog is a term first used to refer to the "smoke and fog" episodes which occurred
over London
– Smoke, sulphur dioxide and particulates from the burning of high-sulphur coal for
heating and industry mixed with fog
– 4,000 people died in an episode in December 1952
– Pollutants were trapped by a temperature inversion
– Particulate concentration reached 4,000 µg/m3 and visibility dropped to 2 m
– This was a reducing smog because it was chemically reducing
– Photochemical smogs are different in that they are oxidizing smogs, containing
oxidizing species.
– Characterized by the yellow-brown haze due to a high concentration of NO2
– oxidizing conditions, reduced visibility, substances that irritate eyes and respiratory
tract due ozone, aliphatic aldehydes, organic nitrates
– The chemistry of photochemical smog is the same as natural photochemistry i.e. the
formation of OH• radicals, ozone, and the oxidation of hydrocarbons; difference is in
levels of NOX and hydrocarbons that lead to generation of oxidizing byproducts
– Reactions leading to photochemical smog are initiated by sunlight
– The most important step is the photochemical cleavage of NO2 to NO and O.
a.
Formation of Photochemical Smogs
– Requires four conditions to be met simultaneously:
- nitrogen oxides (NOJ
- sunlight (UV and visible)
- hydrocarbons
- temperatures above 18°C
– See figure 3.1 in Bunce p. 75
– In the morning as people drive to work, the level of NO and hydrocarbons in the atmosphere
rise; both are emitted by automobiles
– As the sun rises, photochemical reactions start which converts NO to NO2
– The levels of ozone, aldehydes, and other oxidants rise, maximizing about noon
– Note: the NO and hydrocarbons emitted by auto engines are considered primary pollutants;
the results of reactions of the primary pollutants with other atmospheric species, or with light,
produce secondary pollutants such as ozone, aldehydes and other oxidants; tertiary
pollutants are the final form of the pollutant, often in a form that is removed from the
compartment, such as NO2 reacting with OH• to form HNO3 — nitric acid, which forms
nitrates in rain and fog.
– See Manahan 8th ed. pp. 364-366 to see how all of the following reactions fit together.
– Recall that auto engines form NO from the high temperature reactions of N2 and O2:
(12)
N2 + O2 <———> 2NO
ΔH = +180 kJ/mole
- Equilibrium [NO] increases with temperature
– Reverse of (12) is slow at normal temperature, so NO formed gets "locked in" i.e. is limited
by kinetics
– Nitric oxide can be oxidized by atmospheric O2 to NO2:
2NO + O2 ————> 2NO2
rate = K [NO]2 [O2] - (3rd order)
Slow rate of reaction due to low [NO] (<1 ppm) (However, if you form NO in the lab in an open
container, it will quickly convert to NO2 because of the high local concentration of NO, forming
the brown fumes of NO2)
However, NO can be converted to NO2 by reactions with other oxidants, such as O3 and HO2•
(13)
NO + O3 ————> NO2 + O2
(14)
NO + HO2• ————— > NO2 + OH•
– Recall the equations for the formation of OH• and O3 from NO2:
(5)
hν, λ<400 nm
NO2 ——————> NO + O
(6)
O + O2 ——————> O3
(7)
hν, λ<400 nm
O3 ——————> O2* + O*
(8)
O* + H2O ——————> 2HO•
(excited state species)
– NO2 is photolyzed to NO and O, the rate depending on light intensity
– The concentration of O3 builds up, leading to increased OH• concentration
– The OH• reacts with methane and non-methane hydrocarbons (NMHC) to eventually form
CO2 and H2O, but along the way they form byproducts: aldehydes and organic nitrates
– Strong relationship between NOX and ozone in smog:
– If only these species were involved, one could calculate the relative concentrations of NO,
NO2 and O3 based on the intensity of light, temperature and reaction rates.
– However, in the presence of high levels of hydrocarbons and NO, O3 is generated as a result of
reactions of hydrocarbons and OH•
16)
CO + OH• ———> CO2 + H
17)
H + O2 –––––> HOO•
18)
HOO• + NO ———> N02 + OH•
19)
NO2 + O2 + hν ——— » NO* + O3
Net reaction: CO + 2O2 + hν ———> CO2 + O3
When NO levels are low, the oxidation of CO by OH• results in decrease in O3 concentration:
16)
CO + OH• –––—> CO2 + H
17)
H + O2 ––——> HOO•
23)
HOO• + O3 ———> OH• + O2
Net reaction: CO + O3 ———> CO2 + O2
– Reactions with methane and NMHCs also result in ozone production, but form aldehydes
and organic nitrates
– When NO concentration falls and is not available to react with ozone, the ozone starts
reacting with other compounds, especially hydrocarbons to form reactive organic radicals:
O3 + NMHC ———> aldehydes, ketones, peroxides, free radicals (R•, RO•, ROO•)
Aldehydes: R(C=0)H
Ketones: R(C=O)R'
R•, RO•, ROO•: such as CH3•, CH3O•, CH3OO•
– These reactive compounds can react with NO or O2:
ROO• ⎫
⎧ RO•
⎬ + NO –––––> NO2 + ⎨
•
RO ⎭
⎩ R•
R•
2
–––––> ROO•
– The organic free radicals can react with each other and with NO2:
R•
⎫
•
RO ⎬ + NO2 ———>
ROO• ⎭
⎧ RNO2
⎨ RONO2
⎩ ROONO2
– The most important of these organic nitrates is PAN: peroxyacetylnitrate
CH3(C=O)OONO2
– The overall result of the reaction of ozone with NMHCs: [NO]↓, [NO2]↑, [O3]↑ initially, then
[oxidation products]↑, [PAN]↑, [NMHC] ↓
– The reaction scheme is actually much more complex, with over 100 separate reactions
having been identified.
b.
Effects of Photochemical Smogs
– These smogs are toxic and irritating
– The adverse effects are due to ozone and the partially oxidized and reactive intermediates
such as PAN
– Ozone reacts with membranes in the lungs and the eyes, since they contain unsaturated
fatty acids which are susceptible to attack
– The result in impairment of respiratory function and eye irritation at 0.1 ppm
– Perhaps more important, these oxidants can destroy plants, inhibit plant growth and reduce
crop yields at levels > 0.05 ppm (important to crops downwind of Los Angeles)
– Ozone will also degrade materials such as rubber, reducing tire life
– The ozone and hydrocarbons form an aerosol that reduced visibility.
3.
Particulates in the Atmosphere
a.
General
– Microscopic sized particles 0.001 µm to > 100 µm
– Scatter light (similar to Tyndall Effect), resulting in hazy appearance to distant objects
– Large releases of particulates can cool weather due to reflection of sunlight
– Have a high surface area / mass ratio, leading to high chemical and catalytic activity,
e.g. Vanadium on fine particles can promote oxidation of SO2 to SO3
– Terms pertaining to aerosols:
Aerosols
Condensation Aerosols
Dispersion Aerosols
Fog
Haze
Mists
Smoke
Colloidal-sized particles
Formed by condensation of vapours or reactions of
gases
Formed from grinding of solids, dispersion of dusts,
atomization of liquids
Water droplets
Reduced visibility from particulates
Liquid particles
Formed from incomplete combustion
– Classes of aerosols
- carbonaceous particles (containing predominantly carbon)
- metal oxides and glasses (e.g. SiO2, Fe2O3)
- water
- dissolved ionic species (electrolytes, e.g. SO32") ionic solids (salts, e.g. NaCl)
– Processes affecting particulates: diffusion
- coagulation
- sedimentation (dry deposition)
- scavenging by precipitation (wet deposition: rain-out and wash-out)
- condensation of atmospheric water and other vapours
- reaction with atmospheric gases
– the chief sink is wet scavenging and wet deposition
– dry deposition is much slower than wet deposition
– Settling rates: depend on size and density of the particles, as well as the viscosity of the air
– described by the Stokes Law:
rate (m/s) = gd2(Δp)/18η
g = acceleration of gravity (m/s2)
d = diameter of particle (m)
Δp s difference in density (kg/m3)
η = viscosity (Pa s; kg m/s)
– This equation only applies to d > 1 um
– For d < 1 um and p = 2 g/cm3, particle can remain suspended for weeks
– thus small particles can have regional or global impact
– Canadian arctic dust collectors found to contain: soot, sulphate, vanadium and lead (from
combustion, acid precipitation, coal, gasoline)
– This dust originates in western Europe and Russia
– North American particles move out over the Atlantic Ocean and are washed out there
– Large releases of particulates can cool weather due to reflection of sunlight:
- forest fires can darken skies far away from the site of the fire e.g.. fires in Alberta in
1950 significantly darkened skies in Washington DC
- volcanic eruptions inject large clouds of particles into the atmosphere
e.g.. Mount St. Helens
- 'nuclear winter' scenario proposes that large dust clouds arising from a large-scale
nuclear weapons exchange could cause agriculture to collapse due to months-long
darkness
– Fine particles (< 10 um, so called PM10) can also be respirated deep into lungs to cause
health problems
b.
Common Types of Particulates
– Solid Particles
- smoke from:
- forest fires
- industry
- domestic heating
- windborne materials:
- soil particles
- pollen
- bacterial spores
- sea salt
- finely divided rock from volcanoes
- industrially-derived particles:
- fly ash
- grinding operations (cement manufacturing)
– Liquid particles
- fog and clouds
- haze-forming aerosols from smog
– Sea Salt
- wind picks up surface sea water, water evaporates to leave salt behind as particles
- particles will grow at relative humidity > 75% since salts are deliquescent (absorb
moisture from the air)
- rain near coasts can be corrosive due to salt content
– Fly Ash
- very fine dust from coal burning and from municipal and industrial waste incinerators
- metallic and non-metallic oxides, especially SiO2, Al2O3, Fe2O3 and CaO
- contains small amounts of catalytically-active metals such as V, Mn, Cr, Co and toxic
elements such as As, Cu, Ag, Pb
- elemental carbon present in form of soot and carbon black
- toxic organics eg. PAHs, chlorinated dioxins and dibenzofurans also formed on
surfaces
– Soot and smoke
- from incomplete combustion of organics
- derived from coal, petroleum products and wood
- predominantly impure graphite
- very black, reduces visibility
– Sulphate and nitrate aerosols
- salts of sulphuric and nitric acids
- found predominantly in industrial and urban areas areas
b.
Formation
– Two main sources:
1. primary sources, particles generally formed through physical processes, also through
combustion
2. secondary sources, particles formed by chemical processes
i)
Primary Sources
– Dispersion aerosols:
- particles > 1 um such as dust, formed by disintegration of larger particles
- since it takes more energy to break down larger particles than for smaller particles to
bind together, dispersion aerosols tend to be larger particles
- larger particles not as much a health problem as smaller particles
- natural: sea spray, wind-blown dust, volcanic dust
- man-made: coal grinding, cooling tower spray, quarrying, agriculture, "off-road
vehicles"
- dust from abraded and suspended material can travel long distances e.g. yellow dust
from Asian deserts can travel to Arctic and are found in snow, reddish dust from
Sahara found in Europe
- volcanoes emit huge amounts, but most particulates are coarse (1 to 10 µm)
- bursting bubbles in ocean surface source of sea salt aerosols, due to breaking waves
and biological activity; particles vary in size from 0.8 µm to 25 µm
- meteoric dust is estimated to be from 1.6 x109 to 1.7 x 1011 g/yr
– Biological materials:
- spores, insect fragments, grains and pollen, decay of leaf litter - 5 to 25 um
- leaves can directly emit particles containing volatile metals e.g.. zinc, cadmium,
lead, mercury, arsenic, selenium
- particles of biological origin in indoor air: vegetable dusts, hair, skin, feathers, dander,
bacteria, fungi, micro-organisms
– Combustion:
- smoke and soot formed from combustion of carbonaceous fuels (wood, charcoal, coal)
forest fires can produce several tonnes per hectare (1 x 104 m2); 6 to 30 x 1012 g
smoke/yr (as elemental carbon)
- fly ash produced from burning fuel with a high ash content: coal, solid wastes
ii)
Secondary Sources
– Formed by chemical reactions of a primary pollutant
– Can form metal oxides on burning of fuel:
- roasting of sulphide ores generating sulphur dioxide
3FeS2 + 8O2 ——> Fe3O4 + 6SO2
- calcium carbonate in coal converted to calcium oxide and emitted
CaCO3 + heat ——> CaO + CO2
- Pb from leaded gasoline reacting with halogenated scavengers in the fuel:
Pb(C2H5)4 + O2 + C2H4X2 ——> CO2 + H2O + PbX2
(X = Cl or Br)
– Aerosol mists usually involves oxidation of atmospheric sulphur dioxide to sulphuric acid:
2SO2 + O2 + H2O ——> 2H2SO4
which reacts with ammonia or calcium oxide:
H2SO4(droplet) + 2NH3(g) ——> (NH4)SO4(droplet)
H2SO4(droplet) + CaO(particulate) –—> CaSO4(droplet) + H2O
– Volcanoes and industry important source of sulphate aerosols
– Plants emit volatile organic precursors, eg. α-pinene from the sap of pine trees, that undergo
oxidation by OH•, O3 etc. that lead to various ketones, aldehydes, carboxylic acids, phenols
– Overall primary sources
Forest Fires
Dust
Sea Salt
Volcanic Dust
Meteoric Dust
– Overall Secondary Sources
Sulphates
Nitrates
Organic
c.
1012 g/vr
35
300
1000
50
1
150
250
100
Elemental and Inorganic Components
– The composition of particles usually reflects composition of parent material but can be
altered by reactions with the atmosphere, e.g.
2SO2 + O2 + 2H2O ——> 2H2SO4
H2SO4 + 2NaCI ——> Na2SO4 + 2HCI(g)
– "Natural" elements: K, Na, I, Ca, Si, Fe, Cl, Ti
– "Anthropogenic" elements (emitted through industrial activity): Be, Pb, Br, Mn, Co, Cu, V, Zn,
Mg, Ba, Fe, Ti
– Atmospheric reactions contribute to the levels of these compounds in particulates: H2O, HBr,
NO2, SO2, NH3, HC1
i)
Sources:
– Al, Fe, Ca, Si: soil erosion, rock dust, coal combustion
– C: incomplete combustion
– Na, Cl: marine aerosols, chloride from incineration of organochlorine-containing wastes
– Sb, Se: volatile elements, combustion of coal, oil or refuse
– V: combustion of residual petroleum e.g. from Venezuelan crude oil
– Zn: combustion
– Pb: combustion of leaded fuels and wastes
ii)
Asbestos
– Fibrous silicate mineral - approx. composition: Mg3P(Si2O5)(OH)4
– Used as brake linings, insulation, fire barriers, pipe manufacturer, structural materials
– Being phased out due to health concerns
– Inhalable particles can lead to asbestosis (pneumonia-like symptoms), mesothelioma (tumor
of mesothelial tissue lining chest cavity), bronchogenic carcinoma (cancer in air passages of
the lungs)
iii)
Mercury (Hg)
– Mobile and volatile (bp. 375°C)
– Associated with particulate matter
– Source is coal and refuse combustion, volcanoes
– Volatile organomercury compounds such as dimethylmercury (CH3)2Hg, monomethylmercury
salts e.g. CH3HgBr
– Threshold limit value (TLV) of Hg is 0.05 mg/m3; of organo-Mg 0.01 mg/m3
– Mercury is neurotoxic, causing muscle tremors, depression
– Organo-Hg are more toxic because they are more lipophilic and bioconcentratable than
metallic Hg
– Inorganic Hg salts can cause kidney damage
iv)
Lead(Pb)
– Major source used to be tetraethyl lead from leaded gasoline
– During combustion, lead halides form in reactions with halogenated scavengers
dichloroethane and dibromoethane
– During 1970s (max emission) 200,000 tonnes/yr emitted in US
– Other sources include incinerators, Pb paint, batteries
v)
Beryllium (Be)
– Used in specially alloys in electrical equipment, space application, nuclear reactor
components
– Be was used as phosphors in fluorescent lamps
– Was eliminated from this use due to high toxicity
d.
Organic Components
i)
Organic particulates
– Can collect on filters that collect 0.3 µm particles with an efficiency of 99.9%
– Most dangerous particles are those PM10 to PM0.1 since they are respirable
– Extract particles with benzene or ether, then fractionate by polarity or acidity
– Of most concern are aromatics and PAHs, and many oxygenated neutrals: aldehydes,
ketones, epoxides, peroxides, esters, quinones and lactones
– Some are mutagenic or carcinogenic
– The acidic fraction: fatty acids, non-volatile phenols
– The basic fraction: alkaline N-heterocyclic HC’s e.g. acridine
N
– Much organic particulates are oxidized, polymerized HCs and nitrogeneous azaheterocyclic
substances
ii)
Soot
– Highly condensed product of PAHs
– Formed from crystallite, made up of tiny platelets of graphite
– The platelets are made up of several hundreds of fused rings
iii)
PAHs
– Poly cyclic aromatic hydrocarbons: fused aromatic benzene rings
– PAHs are not very volatile, and are found on soot
– Most commonly found in urban atmospheres:
- ~ 20 µg/m3
- ~ 1,000 µg/m3 in flue gas emitted from coal burning
- ~ 100 µg/m3 in cigarette smoke
– Generally products of incomplete combustion
– Formed by pyrosynthesis: at 500°C, bonds are broken to form organic free radicals, followed
by dehydrogenation and reactions of the radicals to create larger organic molecules
– These reactions can form PAHs from small molecules, such as ethane
– Pyrolysis cracks larger hydrocarbons to smaller and less stable molecules and radicals
– naphthalene, phenanthrene, biphenyl usually partitions into the gas phase
– PAHs with 5+ rings almost all are on particles
– Intermediate molecular weight PAHs are partitioned between gas and particles
– PAHs transformed in the atmosphere by the addition of OH• and reactions with NOX
– PAHs on particles are less reactive than in gas phase
– Some PAHs are mutagenic and carcinogenic, mainly after reactions to oxygenated species
by the body's detoxification mechanisms
– They enter through the respiratory tract, can lodge in nasal passages, pharynx and lungs
– Respiratory system can be affects by particulates that enter blood or lymph system
– PAHs can be sorbed in lungs, move into blood, and be transported to other organs
– Higher mortality correlates with high pollution
iv)
Health Effects of Selected Organic Compounds
Type of Compound
Possible Effect
Polar Organics
Extracted by aqueous body fluids, enabling
penetration of tissue in a soluble form
Aldehydes
Suspected of being the primary causal agents in
smoke poisonong; postulated denaturation of amino
acids and RNA; inflamation and necrosis result
Other oxygenated organics (acids,
neutral oxygen compounds)
Some have suspected carcinogenic effects
Polynuclear azaheterocyclic compounds Some of these, such as benz(a,h)acridine
(benzacridines, dibenzacridines,
and dibenza(a,j)acridine are known carcinogens
benzoquinolines)
e.
Reactions on Particulate Surfaces
– Reactions on particles
- uptake of gases and liquids e.g. H2O
- oxidation reactions on surface e.g. SO2 adsorbed and in presence of H2O, oxidized to
H2So4; on graphite soot, but better on metallic oxides
- Photochemical reactions, e.g.
hν, TiO2/ZnO
2CO + O2 ————> 2CO2
hν, TiO2
2N2 + 6H2O ——————> 4NH3 + 3O2
f.
Control of particulates
– emission controls in industry
– bag filters or electrostatic precipitators
– electrostatic precipitators:
- charged tube with central rod
- potential difference 50 to 100 kV
- particles move to oppositely charged surface and discharge
- particles coagulate and eventually fall off
– sedimentation and inertia methods for particulate control
- cyclone separators (dry centrifugal separators)
– Note: must also control dust in grain elevators to prevent explosion hazard
H.
INDOOR AIR POLLUTION
– This discussion will concentration mainly on offices and homes, not workplaces
– Indoor air pollution has been a problem ever since we were burned wood and other fuel
indoors
– Most outdoor pollutants are usually at a lower concentration indoors
– Thought to be due to adsorption of gases on surfaces
– Sulphur dioxide and ozone concentrations will decrease rapidly in a closed room with no
source renewing concentration
– For SO2, t1/2 = 1 hr, for O3, t1/2 = 10 min.
– But can often measure higher concentration of hazardous pollutants indoors than outdoors
– A contributing factor is energy conservation in which homes better sealed against leaks of
warm air, leading to need of additional outside air and use of heat exchangers
– Commercial buildings may recirculate a portion of the Air to save money
– Gases can build up:
- CO2: respiration and combustion
- CO: combustion e.g. smoking, heaters
- formaldehydes and organics: outgassing
- radon: ground vapour infiltration particles: smoking, wood burning
– Why do they build-up?
R = k[ci]V
R = rate of emission (mg/h)
k = # of air exchanges per hour (1/h)
[ci] = concentration (mg/m3)
V = volume of building (room) (m3)
– Rearrange:
[ci] = R/(kV) or
[ci] = [c0] + R/(kV) if outside concentration of species [co]not zero
– See figure 4.2 p. 116 Bunce
– The concentration rises quickly as the number of air exchanges (k) falls below 1/hr, which is
common in most energy efficient homes and buildings
– If the outdoors is the source of the contamination, then increased ventilation will increase
concentration e.g. ozone, nitric acid
– If indoors is the source, then increased ventilation will help
– In extreme cases get "Sick Building Syndrome"
– Can alleviate symptoms by increasing ventilation rate or increase proportion of fresh air
1.
Formaldehyde
– Is used to make resins e.g. urea-formaldehyde [urea (NH2)2CO] or phenol-formaldehyde
(Bakelite) resins
– These materials (esp. urea-formaldehyde) can release excess formaldehyde or can
hydrolyze polymers
– They are used in: plywood and particle board, fiberglass insulation, foam insulation
– Urea-formaldehyde (UFFI) insulation was a major problem in 1970s and 1980s but
apparently now outgassing of foam almost gone now.
– Formaldehyde is an irritant gas, can impair respiratory function in test animals at levels 0.3 to
0.5 ppmv
– In homes with UFFI, people exhibit symptoms of drowsiness, headaches, nausea (also
allergy sensitizer)
– Not yet proven to be human carcinogen
– Emission rates (table 4.2)
plywood and particle board
paneling
glass fibre insulation
clothing and drapery
[µg/g/day
<0.1 to 9
0.8 to 2
0.3 to 2.3
0 to 5
– usually best to increase ventilation
– can also remove chemically by reaction with an oxidant e.g. Al2O3/KMnO4 or a polymeric
amine
– max concentration allowable 0.1 ppmv
2.
Carbon dioxide (CO2)
– Not poisonous but is an asphyxiant at very high concentrations
– Can be 2,000 ppmv in sealed office buildings (355 ppmv in normal in the atmosphere)
– Causes fatigue and inability to concentrate
3.
Carbon Monoxide (CO)
– Much more serious than CO2
– Binds with hemoglobin to form carboxyhemoglobin
- at 10 ppmv, 2% of hemoglobin is inactivated
- at 100 ppmv, 15% of hemoglobin is inactivated
– Generated by kerosene space heaters, wood stoves, tobacco smoking, fumes from attached
garages
– Smoldering cigarette can release 100 mg CO into a room
– Malfunctioning chimneys in homes with gas or oil heat a potential hazard, which is why CO
monitors are recommended for homes
4.
NOX
– Hot flames (gas cooking and kerosene heaters)
– Max acceptable limit in Canada 110 ppbv, the max desirable level is 32 pptv
– Can cause increased incidence of respiratory diseases in children at levels > 30 u.g/L (ppbv)
– NO2 is converted to nitrous acid in a scheme similar to night-time outdoor oxidation:
O3 + NO2 ———>NO3 + O2
NO3 + RH ———> HNO3 + R•
NO3 + NO2 ———> N2O5
N2O5 + H2O ———> 2HNO3
5.
Particulate matter
– Tobacco smoke and wood stove smoke contains PAHs, especially benzo[a]pyrene
– Outdoor concentration of BaP is area of high wood burning for heat is 10 to 1,000 ng/m3,
whereas the average US outdoor concentration is 0.7 ng/m3
– "Average percent excess risk of lung cancer per each 1 ng BAP per m3" over a 70 year
average lifetime
– For workers and urban dwellers in US and UK, 0.2% to 5% excess
– Second hand cigarette smoke declared a "proven human carcinogen"
– Cigarette smoke emits 5 x 1012 particles and 5 µg BaP per cigarette
6.
Radon
– Is a natural substance from spontaneous decay of radium
– Density of 9.7 g/L, bp. -62°C
– Responsible for higher rates of lung cancer in miners
– Radon daughters are actually the big problem (p. 118 Bunce):
222
Rn ———>
218
Po + α (alpha particle – 4He)
– Radon has a half-life of 3.8 days (226Ra parent has a half-life of 1,600 yr)
– It eventually decays to 2I0Pb, which is stable
– Radon is not reactive and non-polar so will not lodge in lungs if inhaled
– But if Radon decays while in the lung, non-volatile metallic daughters are formed
– The daughters readily form oxides that are scavenged by particles or lung tissue
– Damage is done by the alpha-particles that ionize matter in cells, which can lead to cancer
– Radon usually enters through basement of homes, especially those that are air tight
and have negative pressure
– Come from soil gases and ground water
– Hot spots in western Canada include houses in Regina, Uranium City, Winnipeg
– Radon is responsible for -50% of natural background of radioactivity
– In Canada, the 'action level' is 20 pCi/L (800 Bq/m3) (1 Bq is 1 disintegration/s; 1 Ci is
number of Bq of activity from 1 g Radium (3.7 x 1010 Bq)
– 10 to 20 pCi/L (400 to 800 Bq/L) thought to be safer