What is Environmental Chemistry?
ENVIRONMENTAL CHEMISTRY
• Environmental Chemistry: "Environmental chemistry may be defined as the study of the sources, reactions, transport, effects, and fates of chemical species in water, soil, and air environments, and the effects of technology thereon." Stanley Manahan, Environmental Chemistry, 7th Ed.
APCH 211 (2013)
Introduction & Atmospheric Chemistry
Dr PG Ndungu
About Environmental Chemistry
• Environmental chemistry is an interdisciplinary science
– Air, water, soil, etc all are studied within this context
• It’s foundation is analytical chemistry
– Applied to environmental science.
• Environmental chemistry tries to establish baseline effects & how anthropogenic activities alter such
– Need to first understand is it supposed to be there?
– How much is supposed to be there?
– How does the uncontaminated environment work?
Some Perspectives …
Environmental chemistry is a mature discipline (>100 yrs.)
• Mostly studied in other disciplines (biologists noting effects of pesticides, sanitary engineers studying foaming of their water works due to detergents, etc
• Involves Flux!
• Involves chemical reactions and equations, solutions, units, sampling, and analytical techniques 1
Some Aspects of the Atmosphere and Atmospheric Note that the atmosphere Chemistry
The atmosphere to be discussed in this module is that of the Earth
of each planet is unique. Lecture Series 1 ‐ Introduction
SOME ASPECTS OF THE ATMOSPHERE AND ATMOSPHERIC CHEMISTRY
• NASA's Kepler science mission has
confirmed 135 exo‐planets & over
“3277 candidate systems”
• ~990 Exo‐planets have been
discovered to date
The Earth’s atmosphere:
Atmospheric Chemistry
• What is the Earth’s Atmosphere?
• Atmosphere is protective blanket above the Earth and protect it from hostile environment of outer space
• The atmosphere provides carbon dioxide for plants; oxygen for respiration and nitrogen required for nitrogen cycle and for ammonia manufacturing industry.
• Atmosphere is the source of many molecules:
• Atmosphere is a thin shell of gas surrounding the earth. • The major components near the surface of the Earth are:
Nitrogen 78.08%
Oxygen 20.95%
Argon 0.93%
Carbon dioxide 0.0378%
Moisture ranging 0.5 ‐3.5%
Nitrogen‐containing molecules like proteins
 Carbohydrates
Fats & oils
Note that the %concentrations are based on dry atmosphere.
7
8
2
Cont… Atmosphere
Regions of the Atmosphere
• Atmosphere as dumping ground?
•
• The atmosphere can conveniently be divided into four sections based on whether temperature decreases or increases with altitude. • Temperature fluctuations are caused by thermodynamic changes which is a measure of kinetic energy of molecules in the atmosphere.
– In the negative sense, the atmosphere has been used as the dumping ground for many pollutants.
– E.g. SOx, NOx, Atmosphere as a protective layer?
– It absorbs most of the cosmic rays from outer space and protects organisms from their effects.
– It absorbs most of the electromagnetic radiation from the sun and allows selected radiation in the regions of 300 ‐2500nm, such as near‐UV (300‐400nm), visible (400‐
800nm), and near‐IR (900‐2500nm) and radio waves in the region 1x 107‐40x 109nm.
9
Altitude above (distance above sea level)
From lower to higher altitude the regions are:
–
–
–
–
Troposphere = 0‐15 km, Temp decrease with altitude
Stratosphere = 15‐50 km, Temp increase with altitude
Mesosphere = 50‐85 km, Temp decrease with altitude
Thermosphere =85‐500 km, Temp increase with altitude
10
Stratification of the Atmosphere
• Tropopause
Troposphere
• Layer at the top of the troposphere
Lowest layer of the atmosphere
• Serves as Barrier
Extends from Sea Level to 10 – 16 km
– H2O vapor freezes
– Thus H2O can t reach altitudes were it “Tropo” Greek for turning would photo‐dissociate
~Homogeneous mixture (due to air – Thus don t loose hydrogen!
circulating) of gases other than H2O
• Decreasing Temperature with Increasing Altitude
•
•
•
•
•
– Why? – Heated by IR radiation released from the surface
– Upper limit ~ ‐ 56 C, & varies by 1 km
– Drives the weather (somewhat)!
• Contains more than 70% by mass of the entire atmosphere.
• Water content varies due to precipitation, clouds, etc
11
3
Layers of the Atmosphere Cont…
•
•
•
•
• Stratopause
Stratosphere
• Upper boundary of the Directly above troposphere
stratosphere
“Strato” Latin for layered
• Temp starts to decrease again
Temperature increases with altitude up to max of ~ ‐2° C at 50km
Exosphere
o
Thermosphere
Mesosphere
o Just above the stratosphere
o No radiation – absorbing species, thus temp drops with increasing altitude to ‐
90ºC at 85km. – Has ~ 90% of all O3
– Temperature inversion due to absorption of UV energy by O3. (up to 10 ppm by midpoint of this layer)
– Warm air on top of cold stabilizes this layer – little mixing
– ~ Years for pollutants to shift!
o Extends to ~ 500 km
o Temp can reach 1200 K due to highly energetic radiation
o But gas density so low, barely a blip on a thermometer o
o
o
Outermost region
Atoms follow ballistic trajectories & rarely collide
Begins @ ~ 500 & ends @ ~ 10, 000 km Where gas atoms escape
Noctilucent Clouds
Found in the mesosphere
Very rare, usually seen at the poles
14
Energy & Mass Transfer in The Atmosphere…Variation of Atmospheric Pressure
Atmospheric Composition & Pressure
•
Average composition of dry atmosphere, by volume (NASA)
–
–
–
–
–
–
–
–
–
–
Nitrogen, N2 78.084%
Oxygen, O2 20.946%
Argon, Ar
0.934%
Water Vapour,
0.1‐5 %
CO2
350 ppmv
Neon
18.18 ppmv
Helium
5.24 ppmv
Methane
1.7 ppmv
Krypton
1.14 ppmv
Hydrogen 0.55 ppmv
Note: • the concentration of CO2 and CH4 vary by season and location.
• *ppmv represents parts per million by volume.
• The mean molecular mass of air is 28.97 g/mol.
• Gravity Keeps it All in Place
– Thus atmospheric pressure as a function of force over area (total area of the earths surface):
• Pressure decreases with altitude
Mgh
Ph  Po e

Ph = Pressure at any height
Po = Pressure at sea level
M = Average Molar Mass of air
g = acceleration due to gravity
h = altitude
R = gas constant
RT
The atmospheric pressure at any given height is given by: Ph = Poe‐Mgh/RT
• Ph = pressure at any height.
• Po = Atmospheric pressure at Sea level = 101325 Pa
• M = average molecular mass of air molecules in troposphere = 28.97g/mol = 0.029Kg/mol
• g = acceleration due to gravity = 9.81 m.s‐1 at sea level
• R = gas constant = 8.314 J.mol‐1.K ‐1
• T = Temperature in Kelvin (K)
Pascal
(Pa)
15
1 atm =
101,325
Bar
(bar)
1.01325
Torr
(Torr)
760
Pound‐force per
square inch
(psi)
14.696
PV = nRT?!?!!!
4
Energy & Mass Transfer in The Atmosphere
λν = c
• The density of atmosphere decreases sharply with increase in altitude.
• More than 99% of total mass of atmosphere is found within approx 30km (`20miles) of the Earth’s surface. • Total mass of global atmosphere is 5.14x1015 metric tons. This is only a millionth of the Earth’s total mass. Earth’s mass ~5.14x1021 metric tons.
Energy Transport
• “Weather and climate on Earth are determined by the amount and distribution of incoming radiation from the sun.” [see ref.]
• Energy transport which is important in radiating back into space is crucial and is accompanied by three major mechanisms, namely:
• Conduction:‐ occur through interaction of adjacent atoms or molecules without the bulk movement of matter • Convection:‐ involves movement of whole masses of air which may be relatively cold or warm
• Radiation:‐ occurs thro’ electromagnetic radiation in the infra‐red region
Summary of Global Energy Flow, numbers are in W/m2: Image Copied from Ref
Ref:
Earth's Global Energy Budget
Kevin E. Trenberth, John T. Fasullo, Jeffrey Kiehl
Bulletin of the American Meteorological Society
Volume 90, Issue 3 (March 2009) pp. 311‐323
• The physical and chemical characteristics of the atmosphere and critical heat balance of the earth are determined by energy and mass transfer processes in the atmosphere.
• About half of the solar radiation entering the atmosphere reaches the earth’s surface either directly or indirectly after scattering by clouds, atmospheric gases or particles.
• The remaining half is either directed back or absorbed in the atmosphere and its energy radiated back into the space at a later time as infrared radiation.
Fate of Incoming Solar Energy
5
Anthropogenic Changes In The Atmosphere
Evolution of the Atmosphere
Ref: Scotese, C.R., 2002, http://www.scotese.com, (PALEOMAP website)
• Ever since life appeared on earth, the atmosphere has been influenced by the metabolic processes of living organisms.
• For example: carbon dioxide is consumed by photosynthesis reaction and oxygen is produced. Oxygen reacts with metal ions such as iron (II) to produce iron oxides
• Evolution of the Atmosphere
– The first atmosphere likely to have consisted of hydrogen and helium with other simple hydrogen compounds (water, ammonia, methane), eventually driven off – Planet/solar system formation!
– The Second atmosphere was formed by outgassing of gases trapped in the interior of the early Earth (~3‐4 billion years ago) & bombardment from asteroids and comets; it most likely consisted of mainly CO2, N2, liquid water in terms of oceans, etc started to condense
– The third atmosphere (~2.4 billion years ago), change from reducing to oxidizing atmosphere
Basic Processes
• Iron oxide deposits, is proof that oxygen is being consumed.
Fe2+ + O2 + 4H2O →
2 Fe2O3 + 8H+ Metal oxide formation
CO2 + H2O + h →
CH2O + O2 Oxygen production
– If plants are abundant, then oxygen is accumulated and this enables formation of an ozone shield against solar ultra‐violet radiation.
– Ozone shield protects the earth from harsh solar radiation which otherwise would destroy plants, aquatic life and animals.
Anthropogenic Activities That Destroy Ozone Layer
i) Industrial activities which emit a variety of atmosphere pollutants including – SO2
– Particulate matter
– Photochemically reactive hydrocarbons
– Chlorofluoro carbons and inorganic substances such as toxic heavy metals
ii) Burning of large quantities of fossil fuel which can introduce CO2, CO, SO2, NOx, hydrocarbons such as CH4 and particulate soot, polycyclic aromatic hydrocarbons and fly ash into the atmosphere.
iii) Alteration of land surface‐e g
deforestation; burning biomass and vegetation
iv) Agricultural practices which produce methane from digestive tracts of domestic animals and water logged anaerobic soils and nitrogen oxide from bacterial denitrification of nitrates in fertilized soils.
6
Major Effects Of Human Activities On Atmospheric Changes
From the Sun to the Thermosphere
• Thermosphere – wavelengths of 100 nm and smaller blocked by molecules present (N2 split ‐1199.998
to 2N, O2 as well!)
• Increased acidity
• Elevated levels of infrared absorbing gases, that is, greenhouse gases
• Threats to the UV‐
filtering ozone layer in the stratosphere.
• Increased corrosion of materials induced by atmospheric pollutants
N2 split ~ 945 kJ/mol
O2 split ~ 498 kJ/mol
N2+ + e‐ ~ 1500 kJ/mol
O+ + e‐ ~ 1310kJ/mol
Ionosphere?
6.6
10
100
Ozone Layer
3.0
10
10
Ozone is a light blue gas
Unique ‘electric odor’
B.p. ‐110 C
Absorbs UV in the 200 –
300 nm range
• Chapman Reaction Summarize Production & Destruction of Ozone
Ozone = O3
– Known since ancient times, due to its formation during electrical storms
– Christian Friedrich ‐ as a chemical or molecule
– Discovered in 1913 by the French physicists Charles Fabry and Henri Buisson
– Studied Extensively by British meteorologist G. M. B. Dobson
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Processes For Catalytic Decomposition Of Ozone
•
•
•
•
•
2.0
O 2  h   2O
<240nm
O  O 2  M  O3  M
O3  h  O2  O
O  O3  2O 2
*
Ozone Destruction catalysts
There are several catalytic routes that lead to removal of ozone in the stratosphere. Many of the processes share a general mechanism as follows:
X + O3
→
XO + O2
XO + O
→
X + O2
________________________ O + O3
→
2O2
Net reaction
• Note in the above general mechanism X is not consumed or generated i.e. it’s has a catalytic role
<325nm
7
Ozone Catalytic Cycles
• Three important catalytic cycles for ozone decomposition:
1. NO/NO2
2. OH/OOH (hydroxyl/hydroperoxyl radicals)
3. Cl/OCl (chloride/
• These 3 were the first catalysts determined to be the culprits!
NB: Depending on the altitude and the mixing ratio, each of these species has varying ability to destroy Ozone.
• OH/OOH at 15 km & below accounts for most of the ozone
depletion
• NO/NO2 at 30km & above is the main catalyst for removal of ozone.
• Cl/OCl catalytic cycle is said to account for ozone destruction
around between 15 ‐ 30km altitude
Ozone Depletion via OH/OOH
 This set of catalytic reaction
depends on the availability of a
source of hydrogen to combine
with oxygen.
 Water and Methane are the most
important sources.
 However the temperature in the
tropopause is about –50 °C. At this
temperature water is frozen in
crystals and in that form cannot
cross to the stratosphere; however
methane can!
 Main Sources of methane
 anoxic waters
 Soils
 farming (especially cattle)
• OH generation from methane
CH 4  O  OH  CH 3

OH/OOH Ozone destruction
OH  O3  OOH  O2
OOH  O  OH  O2
O3  O  2O2 (net  reaction)
OH is the X in the Chapman rxns
Hydrogen & Hydroxyl catalysts
• The hydrogen radical is generally produced from UV
splitting of water.

H +OH
H2O + h
• Hydrogen radical also participates in ozone destruction
catalysis as follows
OH + O2
H + O3 
OH + O 
H + O2
_________________________________

2O2
Net Reaction
O + O3
• Another large source of OH is from excited O (from ozone
rxns…see previous slides) & water
2 OH
O + H2O 
Catalysis By Nitrogen‐containing Species: NOX
• NOx (mainly NO and NO2) are found in the troposphere
• Sources of NOx
– Power plants
– Cars, planes, ships, (combustion) etc
– Rxns between nitrous oxide (N2O) and O radicals
• N2O can be produced from anorexic waters & soils
• NO and NO2 radicals – last ~ 4 days
– very water soluble in water, so easily form nitric acid and precipitate (rain, snow, etc). – Only a small fraction migrates to the stratosphere.
• Nitrous oxide, N2O is main source of NOx in the stratosphere. Mostly in the Stratosphere, but can occur in the troposphere
N 2O  O  2 NO
Also from Thermal Processes
N 2  O2  2 NO
NO in the Chapman reactions NO  O 3  NO 2  O 2
NO 2  O  NO  O 2
O 3  O  2 O 2 ( net  reaction )
Reason No Fleets of Supersonic Jets Cruising the Stratosphere
8
Catalysis By Chlorine‐containing Species: ClOX
• Chlorine and Chlorine‐containing radicals (●Cl and ClO● and their bromine analogues) are the most reactive of all the stratospheric species that catalyze ozone destruction.
• Sources of ClOX are anthropogenic as well as natural. • The most important natural source is methyl chloride, CH3Cl.
• Average [CH3Cl] ~ 550 ppt
• Sources of CH3Cl – Biological processes in the ocean
– Small amounts from burning vegetation
– Small amounts from volcanoes
• In the Troposphere:
Other Catalytic Cycles With Halogens…
Fate of CH3Cl in Troposphere CH3Cl  OH  CH2Cl  H 2O
Main Stratospheric Rxn
CH 3Cl  h  CH 3  Cl
Cl then follows normal O3 depletion route
Cl  O 3  ClO  O 2
ClO  O  Cl  O 2
HOO  ClO  HOCl  O2
HOCl  h  OH  Cl
Cl  O3  ClO  O2
Cl  O3  ClO  O2
2O3  3O2
2O3  3O2
O 3  O  2 O 2 ( net  reaction )
Together above reactions account for how ~ 60% of the Ozone depletion
Chlorofluorocarbons (CFC’s) or Freon's
• CFCs were developed in the 1930’s
• Properties include:
CFCs…
• During the 1970’s, researchers determined that CFCs were very stable in the troposphere
Low Bp
Low viscosity
Low Surface tension
Chemical & biologically inert
• When compressed at room temperature CFCs easily liquefy, & when the pressure is released they gasify taking up large amount of heat
• These are the main reason used as refrigerants, polymer foaming agents & as solvents for cleaning sensitive components (electronics)
• Before CFCs, common refrigerants were NH3 or SO2 (did kill when leaked) Br  O3  BrO  O2
OH  O3  HOO  O2
– Other natural chlorine sources (salt spray, HCl from volcanoes) are washed out by precipitation (rain, snow, etc)
–
–
–
–
BrO  ClO  Br  Cl  O2
– OH radicals had no effect
– Lifetime was determined to be > 100 years
General rule for determining the Chemical formula of a CFC:
a) Add 90 to the number
b) Then use the resulting 3 – digit number to determine number of C, H, & F
c) Add Cl to saturate the carbon/s
e.g.
CFC – 12 (Freon – 12)
12 + 90 102
1 one C atom
0 zero H atoms
2 2 F atoms
Need 2 more Cl to saturate the C
Thus CFC – 12 CCl2F2
• Mid 1970’s Rowland & Molina suggested the main environmental sink for CFCs was in the stratosphere, resulting in Cl production
– End Game Cl radical destroys ozone
CFCl3  h  CFCl2  Cl
< 290nm
Note:
This can undergo further catalytic cycles producing more Cl radicals!
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“Antarctic Ozone Hole”
• 1985 discovery of the Antarctic Ozone hole – Limited progress in limiting CFCs was done in the 1970 s & early 1980 s
– The Ozone Hole validated the concept of anthropogenic ozone depletion
– Led to the Montreal Protocol (27 countries signed in 1987) Ozone Hole…
 An Ozone hole exists when the total ozone column is less than 220 DU
During the summer months HCl is a gas & relatively stable
ClONO2( g )  HCl( S )  Cl2( g )  HNO3( s )
The formation of HCl ice crystals in PSCs catalyses formation of Cl2
• Why the Antarctic?
– Differences in the Northern & Southern hemispheres results in less ozone mixing in the Southern hemisphere so regions of depletion recover much slower
– The stratosphere over Antarctica is relatively cooler than the arctic, which allows for more frequent polar stratospheric clouds (PSCs) during the polar night in the winter season
– PSCs form at temperatures of 195 K & lower
– At these temperatures initially have nitric acid trihydrate crystals (HNO3●3H2O)
Early Spring (September & October) Sun returns & enough photons to produce Cl radicals
Cl2( g )  h  2Cl
OZONE_D2011-07-01%P1D_G^1280X720 MMERRA_LSH mp4
 Dobson Unit (DU) is the unit used to measure the total amount of ozone over a specific area
 The unit is named after Gordon Dobson (1889 – 1976)  Defined as the number of molecules of ozone needed to make a 0.01mm thick layer of pure ozone at 0°C and 1 atm
 A Column of air with 1DU will have 2 69 x 1016 O3 for every cm2 at the base of a column
 The Ozone layer has an average thickness of 300 DU, thus when compressed to a layer of pure ozone it would be 3 mm thick Greenhouse Gases And Global Warming
• Greenhouse gases are those
gases that absorb infrared
radiation. This way they allow
incoming solar radiation energy
to penetrate the earth’s surface
while re-absorbing IR
radiation.
• Examples of greenhouse gases:
carbon dioxide, methane and
chlorofluorocarbons absorb IR
radiation and thus retain heat.
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Greenhouse gases Cont…
The "greenhouse effect"
• Greenhouse gases are components of the atmosphere that contribute to the greenhouse effect.
• Without the greenhouse effect the Earth would be uninhabitable, the mean temperature of the earth would be about −19 °C (−2 °F, 254 K) rather than the present mean temperature of about 15 °C (59 °F, 288 K)]. • Greenhouse gases included in the order of relative abundance: (i) water vapour, (ii) carbon dioxide, (iii) methane, (iv) nitrous oxide, and (v) ozone. • Greenhouse gases come from natural sources and human activity.
• When sunlight reaches the surface of the Earth, some of it is absorbed and warms the surface. • Because the Earth's surface is much cooler than the sun, it radiates energy at much longer wavelengths than the sun does, peaking in the infrared at about 10µm.
• The atmosphere absorbs these longer wavelengths more effectively than it does the shorter wavelengths from the sun. Greenhouse effect…
Greenhouse effect…
•
• The absorption of this longwave radiant energy warms the atmosphere; the atmosphere also is warmed by transfer of kinetic energy and latent heat from the surface. • Greenhouse gases also emit longwave radiation both upward to space and downward to the surface. • The downward part of this longwave radiation
emitted by the atmosphere is the "greenhouse effect.".
•
•
•
•
The greenhouse effect is the process in which the emission of infrared radiation by the atmosphere warms a planet's
surface. The name comes from an incorrect analogy with the warming of air inside a greenhouse compared to the air outside the greenhouse. The greenhouse effect was discovered by Joseph Fourier in 1824 and first investigated quantitatively by Svante
Arrhenius in 1896.
The Earth's average surface temperature of 15 °C (59 °F) is about 33 °C (59 °F) warmer than it would be without the greenhouse effect. Global warming, a recent warming of the Earth's lower atmosphere, is believed to be the result of an enhanced greenhouse effect due to increased concentrations of greenhouse gases in the atmosphere. In addition to the Earth, Mars and Venus have greenhouse effects
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IR Absorption by Greenhouse Gases
• Infrared (IR) radiation is electromagnetic radiation of a wavelength longer than that of visible light, but shorter than that of microwaves. The name means "below red" (from the Latin infra, means "below"), red being the color of visible light with the longest wavelength. • Infrared radiation has wavelengths between about 750 nm and 1 mm (1000,000 nm), spanning five orders of magnitude. • Humans at normal body temperature can radiate IR radiation at a wavelength of 10 microns (10,000 nm).
IR…•
• The infrared portion of the electromagnetic spectrum is divided into three regions; the near‐, mid‐ and far‐ infrared, named for their relation to the visible spectrum. • The far‐infrared, approximately 400‐
10 cm‐1 (1000–30 μm), lying adjacent to the microwave region, has low energy and may be used for rotational spectroscopy. • The mid‐infrared, approximately 4000‐
400 cm‐1 (30–1.4 μm) may be used to study the fundamental vibrations and associated rotational‐vibrational
structure. • The higher energy near‐IR, approximately 14000‐4000 cm‐1 (1.4–
0.8 μm) can excite overtone or harmonic vibrations.
IR Absorption
• At the atomic level, infrared energy elicits vibrational
modes in a molecule
through a change in the dipole moment, making it a useful frequency range for study of these energy states. • Infrared spectroscopy
examines absorption and transmission of photons in the infrared energy range, based on their frequency and intensity.
The reciprocal centimeter is the number of wave cycles in one centimeter; whereas, frequency in cycles per second or Hz is equal to the number of wave cycles in 3*1010
cm (the distance covered by light in one second). Wavelength units can be in micrometers, microns (μ), instead of nanometers for the same reason.


E
hc
1

IR Vibrations • Infrared spectroscopy exploits the fact that molecules have specific frequencies at which they rotate or vibrate corresponding to discrete energy levels. • These resonant frequencies are determined by the shape of the molecular potential energy surfaces, the masses of the atoms and, by the associated vibronic coupling. • In order for a vibrational mode in a molecule to be IR active, it must be associated with changes in the permanent dipole of a molecule. Symmetric O‐H Stretch in H2O
10 4
Symmetric O‐H Bend in H2O
Asymmetric O‐H Stretch in H2O


3n – 6 degrees of freedom: H2O = 3(3) – 6 = 3
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IR… In Simple Molecules



The resonant frequencies of the molecules can be in a first approach related to the strength of the bond, and the mass of the atoms at either end of it. The frequency of the vibrations can be associated with a particular bond type.
Simple diatomic molecules have only one bond, which may stretch. Example‐vibration of molecules



3n – 5 degrees of freedom – Linear Molecules; CO2 = 3(3) – 5 = 4
IR radiation is relatively low energy, can only affect various vibration modes
When the frequency of IR radiation equals the frequency of a bond, the molecule absorbs the IR radiation (need a dipole!)
• For example, the atoms in a CH2 group, commonly found in organic compounds can vibrate in six different ways: symmetrical and antisymmetrical stretching, scissoring, rocking, wagging and twisting:
 Classical Theory See’s the bond between two atoms as a simple harmonic oscillator (ball on a spring) – so should only see one vibration, but see many peaks experimentally
Quantum Mechanics
 Vibrations are quantized, thus vibrations can only occur at specific frequencies
 For an energy transition, the molecule must absorb a photon with exactly the right energy
Effect of GHG on IR… IR...complex molecu
Several Atmospheric gases are IR active
Surface IR radiation emitted by the surface is absorbed by these gases “trapping” this energy • More complex molecules have many bonds, and vibrations can be conjugated, leading to infrared absorptions at characteristic frequencies that may be related to chemical groups. • The infrared spectra of a sample is collected by passing a beam of infrared light through the sample. Examination of the transmitted light reveals how much energy was absorbed at each wavelength. • This technique works almost exclusively on samples with covalent bonds.  No atmosphere Earth is a simple
black body radiator (~ 255K)
 With atmosphere we observe the
average Temp ~ 288 K
 GHG provide the added thermal
energy due to their IR activity
Iout, W / m2 = 283 8 284.4
OH
CO2 (ppm)= 390
800
http://geoflop.uchicago.edu/forecast/docs/Projects/
modtran.orig.html
13
Photochemical smog
Smog
• The term "smog" was first coined by Dr. Henry Antoine Des Voeux in his 1905 paper presented in a meeting of the Public Health Congress in London, entitled “Fog and Smoke.
• Smog is a concern in most major urban centers but, because it travels with the wind, it can affect sparsely populated areas in the countryside as well. • Two Types of Smog
– Classical or London smog is a mixture of SO2, carbon particles & water vapor. • Was very common in London for many centuries but got worse during industrial revolution
• Main cause was burning of high sulfur coal
• The SO2 in the mixture gave the smog a slightly acidic & reducing properties hence sometimes referred to as reducing smog. – Photochemical (Oxidizing) or Los Angeles Smog
•
•
•
•
Need warm air (>290 K)
Need lots of sunlight
Need lots of hydrocarbons and NOx
Stable air masses (e.g. city surrounded by mountains)
Great London Smog of 1952
• The Great Smog of 1952
darkened the streets of London and killed approximately 4,000 people in the short term (a further 8,000 died from its effects in the following weeks and months). • Reluctant to admit that coal smoke was to blame, the British government initially blamed a flu epidemic. • In 1956 the Clean Air Act
introduced smokeless zones to the capital, London. •
•
•
Smog can form in almost any
climate where industries or cities
release large amounts of air
pollution.
However, it is worse during
periods of warmer, sunnier weather
when the upper air is warm enough
to inhibit vertical circulation.
Can be observed for extended
periods of time over densely
populated cities or urban areas
–
–
–
–
–
–
–
–
London
Los Angeles
Mexico City,
Houston
Toronto
Athens
Beijing
What about Johannesburg?
Cont…. • Later in London, only smokeless fuels could be used in these areas. • Consequently, reduced sulphur
dioxide levels made the intense and persistent London smog a thing of the past. It was after this that the great clean‐up of London began and buildings recovered their original stone façades which, during two centuries, had gradually blackened.
• Smog caused by traffic pollution, however, does occur in modern London.
14
Natural Disasters as Smog triggers
• Smog can also be due to natural causes that are not anthropogenic. Examples:
• An erupting volcano can also emit high levels of sulphur dioxide creating volcanic smog, or vog. • The burning of forests in Indonesia
has on a number of occasions created prolonged smog‐like haze which have extended to parts of Malaysia, Philippines, Singapore and Thailand although a lot of the times these fires are started by farmers who want to clear away land for the start of the new planting season.
• Other examples
–
–
–
–
Russian Forest fires July 2010
California Wild Fires (2007)
Australia (2009)
RSA (Cape Town Region!)
Effects of Smog
The Chemistry of Photochemical Smog
•
•
•
•
The atmospheric smog varies with time of the day. In the morning at 6 00am, the motor traffic takes to the street and a simultaneous increase in the atmospheric concentrations of volatile hydrocarbons and nitric oxide, NO, is observed. Nitric oxide concentration rapidly reaches a maximum and then decreases while at the same time the nitrogen dioxide concentration fall off
and elevated levels of oxidizing agent and aldehydes are detected.
In the evening, a similar pattern is observed but generally the atmospheric concentrations are relatively lower. The concentrations drop to background levels and remain constant during the night. The smog http://www.ems.psu.edu/~lno/Meteo437/Smog1.jpg
consisting of a mixture of: partially oxidized hydrocarbons, ozone and other oxidants
Chemistry of Photochemical Smog… Steady State of Ozone
Troposphere’s production of ozone occurs via initial photolysis of nitrogen dioxide to produce oxygen radicals • Smog causes eye irritation, adversely affect plant growth and are implicated in serious ecotoxicological problems.
• In the evening, a similar pattern is observed but generally the atmospheric concentrations are relatively lower.The
concentrations drop to background levels and remain constant during the night. • The smog consisting of a mixture of: partially oxidized hydrocarbons, ozone and other oxidants.
• Smog causes eye irritation, adversely affect plant growth and are implicated in serious ecotoxicological problems.
NO2  h  NO  O
In the troposphere the oxygen radicals are quickly consumed by oxygen molecules
O  O2  O3
Ozone, for the most part, reacts with nitric oxide producing nitrogen dioxide
O3  NO  NO2  O2
These three reactions are fast, and the first 2 produce ozone, but the last one destroys ozone, eventually reach a balance with no net production i.e. ‘steady state concentration’
15
Chemistry of Photochemical Smog…
N 2  O2  2 NO
From Cars
This 1st rxn generates NO, which consumes ozone to generate NO2
Example Hydrocarbon (Ethane)
CH 3CH 3  OH (O2 )  CH 3CH 2O2  H 2O
CH 3CH 2O2  NO(O2 )  CH 3CHO  HOO  NO2
OH  CH 3CHO  CH 3CO  H 2O
OH  CO  (O2 )  HOO CO2
CH 3CO  O2  CH 3COOO
OH  HC  (O2 )  ROO H2O
CH 3COOO  NO2  CH 3COOONO2
These next two rxns simplify the role hydrocarbons (HC) & CO have in producing peroxy radicals HOO NO  OH  NO2
ROO NO(O2 )  R' O  NO2  HOO
• Ozone levels increase due to NO (from cars) NO2 via CO and hydrocarbon pathways
• Hydrocarbons & NO2 react to give aldehydes, ozone, and peroxyacetyl nitrate
The peroxy radicals oxidize NO, to generate NO2
The Nature of Photochemical Smog
• A number of chemicals are present at elevated atmospheric concentrations during a photochemical smog event. • Some of the chemicals are gases while others particularly aldehydes
exist as liquid droplets in form of aerosols. This is the cause of the hazy appearance present during an intense smog. • The yellowish colour is due to nitrogen dioxide.
• Peroxyacetyl nitrate (PAN) are the main cause of eye irritation. – PANs are stable and can be transported over long distances by air currents.
Typical Concentrations of some smog chemicals
• Carbon monoxide: 10,000‐30,000 ppbv in polluted area but < 200ppbv in unpolluted areas
• Nitrogen dioxides: 100‐400 ppbv in polluted areas but <20 ppbv in unpolluted areas
• Hydrocarbons: 600‐3000ppbv in polluted <300 ppbv in unpolluted areas
• Ozone: 50‐150 in polluted but <5 ppbv in unpolluted
• Peroxyacetic nitric anhydride (PAN) • 50‐250ppbv in polluted but <5 ppbv in unpolluted areas.
16
Acid Rain
• “Acid rain" – General term applied to any form of wet precipitation, usually in the troposphere, with acidic species stronger than CO2
– Includes rain, sleet, snow, fog, or dew
• Natural pH of rain water ~ 5.6 (from dissolved CO2)
– Can have deposition of dry gases and compounds – so called dry deposition
– Aqueous and dry deposition are collectively termed acid deposition
Acidic Deposition
Wet deposition • refers to acidic rain, fog and snow. • If the acidic chemicals in the air are blown into areas where the weather is wet, the acids can fall to the ground in the form of rain, snow, fog, or mist.
Dry deposition
• In areas where the weather is dry, the acid chemicals may become incorporated into dust or smoke and fall to the ground through dry deposition sticking to the ground, buildings, homes, cars, and trees. Later, when moisture content increases, acid solution is produced.
• The precursors or chemical forerunners of acid rain formation result from both natural sources, such as volcanoes and decaying vegetation and man‐made sources, primarily emissions of sulfur dioxide (SO2) and nitrogen oxides (NOx) resulting from fossil fuel combustion. • Acid rain is particularly damaging to lakes, streams, and forests and the plants and animals that live in these ecosystems. Main Culprits in Acidic Precipitation
Nitrogen species: • NOx: – Sources include burning of fossil fuels, biomass, etc
– Natural Sources; Anorexic waters, soils, etc • NH3
– Animal excreta, fertilizers and microbiological release
Sulfur Species
• SO2:
– Fossil fuels & sulfur ore smelting
• H2S, & CS2:
– Wetlands and submerged soils
• Dimethylsulfide, (CH3)2S, carbonyl sulfide, COS, methyl mercaptan, CH3SH and Dimethyl disulfide, CH3SSCH3. Atmospheric Production Of Nitric Acid
The principle reaction sequence
contributing to production of
nitric acid starts with nitric
oxide, NO from combustion
processes.
Nitric oxide Chemistry: Daytime
 NO is oxidized by O2, O3 or
ROO – e.g.
NO + O3  NO2 + O2
 This NO2 radical can then
contribute to ozone and OH
radical production i.e. plays a
role in smog formation
 Short lifetime, thus smog events
don’t last, & not that frequent
 Main removal sequence for NO
is via catalyst (M), OH rxns
NO2 + OH + M  HNO3 + M
– Ocean and soils.
17
Reactions of NO3 with hydrocarbons
Nitric acid Production at Night
 Key species is the nitrate
radical (NO3).
 Formed via O3 & NO2
 Atmospheric production
of dinitrogen pentoxide
(N2O5) occurs when NO3
reacts with NO2 is the
only way to form in the
atmosphere.
 N2O5 is a store of NO3.
 Can decompose back to
NO3 and NO2
 Can react with water to
form nitric acid
o Easily removes an H from alkanes
NO3 + RH  R + HNO3
o The R radical can then react with O2 to form peroxyl radicals
o With alkenes the NO3 radical reacts via an addition mechanism producing
nitro‐oxy substituted organic radicals which can regenerate NO2, or
relatively stable organic nitrate compounds (see ‐ Paul S. Monks. Gas‐phase
radical chemistry in the troposphere. Chem. Soc. Rev., 2005, 34, 376‐395,
for e.g. with propene)
NO3 + CnH2n  CnH2nNO3
o With aldehydes, typically form nitric acid and the corresponding radical
NO3 + RCHO  RCO + HNO3
o Overall nighttime chemistry of NO3 can recycle NOx, or form HNO3,
depending on the mix of hydrocarbons
o Can be a nighttime source of OH radicals
Easily zapped by the rising sun!
( ~ 600 ‐700nm)
Atmospheric Production of Sulfuric Acid
Removal of Nitric Acid
Oxidation of reduced sulfur species
 Removal is accomplished by either wet or dry deposition
 One of the main contributors to acid precipitation.
 Nitric acid can react with ammonia:
NH3 + HNO3  NH4NO3
 The ammonium nitrate, NH4NO3 can act as a condensation nucleus for the formation of a water droplet or it can be deposited as part of the solid aerosol
 Production of sulfuric acid is more
complex than that of nitric acid as
the starting materials cover a
wide range of reduced sulfur and
partially
oxidized
sulfur
compounds.
 These include hydrogen sulfide,
carbon disulfide, carbonyl sulfide,
methyl
mercaptan,
dimethyl
disulfide, and dimethyl sulfide. All
these compounds contain sulfur in
its oxidation state (‐2).
 Mostly from natural sources
S
H
H
S
C
S
O
C
S
18
Sequence of Reactions of Sulfur Cmpds
 Once sulfur compounds are in the
air, a sequence of reaction begin
as follows:
H2S + OH  H2O + SH
CS2 + OH  COS + SH
COS + OH  CO2 + SH
NB:



The above reactions release
thionyl radical, SH as the initial
product.
Importance of OH!
Hydrogen sulfide and carbon
disulfide unlike carbonyl sulfide,
are very reactive and therefore
are quickly consumed
Where is the OH coming from?
Further oxidation of thionyl
radical
eventually produces
sulfur dioxide:
SH + O2  SO + OH
SH + O3  SHO + O2
SHO + O2  SO + HOO
The SO radical can then react with either
O2, O3 or NO2 to give SO2 and other
products.
2SO + O2  2SO2
SO + O3  SO2 + O2
SO + NO2  SO2 + NO
NB.
SO2 is ultimately converted to sulfuric
acid, H2SO4.
 Sulfur dioxide SO2 is also released in large quantities directly into the atmosphere from sulfide ore smelting and fossil‐fuel combustion.
Understand acid deposition’s causes and effects
 understand acid deposition’s causes and effects, and to track changes in the environment.
There are several ways to reduce acid rain (i.e. acid deposition). These range from government policy to societal changes and individual action.
e.g. given by epa
(http://www.epa.gov/acidrain/reducing/)
:
The steps involved in reduction of acid deposition are:  Dimethyl sulfide is produced by phytoplankton
living in surface waters of the ocean. It is
oxidized by hydroxyl radical (OH) with a final
product being sulfuric acid.
Reducing Acid rain
Reducing Acid Rain
• Understand acid deposition s causes and effects
• Clean up smokestacks and exhaust pipes
• Use alternative energy sources
• Restore a damaged environment
• Look to the future
• Take action as individuals
More About SOx…
 Scientist  collect air, water & soil samples and measure them for various characteristics such as pH and chemical composition, and investigate the effects of acid deposition on human‐made materials.
 Scientists  understand the effects of sulfur dioxide (SO2) and nitrogen oxides (NOx), and any other acid causing species
 People to understand the process of how acid rain damages the environment
(Need to educate Policy Makers!).
 People to find out what changes could be made to the air pollution sources
that cause the problem (Need to educate Policy Makers!).
Recycle, Reuse, Reduce or the 3 R’s of waste management has evolved from the initial concepts championed in the 1970’s & now includes, amongst other things, prevention & minimization
19
Steps to solve acid deposition problem
a)
Clean up smokestacks and exhaust pipes
WHY?
Almost all of the electricity that powers modern life comes from burning fossil fuels such as coal (Over 80% in RSA), natural gas, and oil. Sulfur dioxide (SO2) and nitrogen oxides (NOx) are the main acid chemicals.


Options for reducing SO2 emissions,
include: using coal containing
less
sulfur, washing the coal, and using
devices called “scrubbers” to
chemically remove the SO2 from the
gases leaving the smokestack.
Power plants to change type of fuels
e g, burning natural gas creates much
less SO2 than burning coal.
Steps to solve acid deposition problem
a)
Clean up smokestacks and exhaust pipes
WHY?
Almost all of the electricity that powers modern life comes from burning fossil fuels such as coal (Over 80% in RSA), natural gas, and oil. Sulfur dioxide (SO2) and nitrogen oxides (NOx) are the main acid chemicals.
 Options for reducing SO2 emissions,
include: using coal containing less
sulfur, washing the coal, and using
devices called “scrubbers” to
chemically remove the SO2 from the
gases leaving the smokestack.
 Power plants to change type of fuels
e.g, burning natural gas creates
much less SO2 than burning coal.
Steps to solve acid deposition…
b)
Use alternative energy sources
 Other sources of electricity
besides fossil fuels. They include
nuclear power, hydropower, wind
energy, tidal, geothermal energy,
and solar energy.
 Alternatives to internal combustion engines  batteries, solar cells, and fuel cells
NB: All sources of energy have
environmental costs as well as
benefits.
Steps to solve acid deposition…
b)
Use alternative energy sources
 Other sources of electricity besides
fossil fuels. They include nuclear
power, hydropower, wind energy,
tidal, geothermal energy, and solar
energy.
 Alternatives to internal combustion engines  batteries, solar cells, and fuel cells
NB: All sources of energy have
environmental costs as well as
benefits.
20
Steps to solve acid deposition…
c) Restore a damaged environment
NB!
It takes many years for ecosystems to recover from
acid deposition, even after emissions are reduced and
the rain pH is restored to normal.
There are some things that people can do to bring back
lakes and streams more quickly.
 Limestone or lime (a naturally occurring basic
compound) can be added to acidic lakes to “cancel
out” the acidity. Liming, has been used extensively in
Norway and Sweden.
Steps to solve acid deposition…
e)
Take action as individuals
Yes! You too can make a difference!
 Turn off lights, computers, and other appliances when you're not
using them (Unplug chargers when not in use!).
 Use energy‐efficient appliances for lighting, air conditioners,
heaters, refrigerators, washing machines, etc.
 Use public transportation, or better yet, walk or bicycle whenever possible  Buy vehicles with low NOx emissions, and properly maintain your vehicle (In the News, Mercedes, VW, etc, complaining RSA fuel quality has too much sulfur for latest technologies).  6. Be well informed. WHAT ABOUT TALKING TO GOVERNMENT
What Policies, Programs, or commitment in general has RSA made
Steps to solve acid deposition…
d) Evaluation of the progress made on acid rain reduction process
 Monitoring Very Important!
 If the depositions are reduced, environmental protection agency (EPA) scientists must assess the reductions to make sure they are achieving the anticipated results.  If no changes, to consider additional ways to reduce emissions that cause acid deposition. Example: focus on energy efficiency and alternative energy.
The Chemistry of Urban & Indoor Atmospheres
• The chemical composition of air in places where people live & work (urban areas, homes, offices, etc) vary with modernization or industrialization of the locality.
• Urban areas are likely to be affected by atmospheric pollution due to the following major factors: – Combustion of fossil fuels (mostly cars)
– In‐space heating and cooling (wood burning?)
– Power generation and industrialization – Incineration of waste materials
21
Urban & Indoor Atmospheres…
•
•
•
•
Use of petroleum products especially in motor vehicles result in ground‐level emissions of carbon monoxides, volatile hydrocarbons, nitrogen oxides and sometimes, lead compounds. These emissions are accompanied by aldehydes and other secondary pollutants.
The combustion of biomass and coal produces substantial concentrations of solid particulate matter along with nitrogen oxides, polyaromatic
hydrocarbon (PAHs) compounds as well as sulphur dioxide.
Open burning refuse or garbage cause air pollution is a source of volatile organic carbon compounds and solid particulate matter (SPM).
Hurricanes and wind are the source of particulate matter such as dust especially in dry areas.
Quality guidelines
Pollutants In The Urban Atmosphere
• World Health Organization (WHO) Standards for Air Quality
• The WHO guidelines for air quality take into account time period over which measurements is done. This is known as human exposure.
• Potential toxicity depend on both atmospheric concentration and duration of contact with the atmosphere. That is;
• Exposure = Concentration x time
Table 1 is a summary of WHO guideline values for air
quality-values in µg.m‐3 or parts per trillion in volume (pptv).
Pollutant
• The quality guidelines must specify the acceptable concentration to be exposed to humans over a specified period. • Example: carbon monoxide at 20mg.m‐3 (20ppbv) may be acceptable if exposure time is 1 hour but not acceptable for longer period times.
• For longer exposure such as 8hrs, the allowed concentration of carbon dioxide should not exceed 10mg.m‐3, that is, 0.01ppmv or 10ppbv.
• WHO works closely with United Nations Environmental Program (UNEP) to carry out air quality monitoring.
Max.
time Average
weighted
time
(µg/ m-3)
SO2
500
10 min
CO
30,000
1 hr
NO2
400
1 hr
O3
150-200
1 hr
SPM (black smoke) 100-150
24hr
TSP
RSP; PM10
150-230
24hr
Pb
70
24hr
0 5-1
1yr
SPM=Suspended
particulate matter
TSP=Total suspended
particulate
RSP = Respirable
suspended particulate,
PM10 with particle size
< 10µm
22
About Particulate Matter
• Any type of fine solid or liquid particles that enter the atmosphere via processes near the earth's surface (Aerosol)
• Particulate matter is usually abbreviated as PM, and the size indicated after the letters.
Health Problems & SPM
Chronic:
• cardiovascular & respiratory diseases
• lung cancer
– PM10 (particles with an aerodynamic diameter smaller than 10 μm) or
– PM2.5 (aerodynamic diameter smaller than 2.5 μm).
• The major components include sulfate, nitrates, ammonia, sodium chloride, carbon, mineral dust and water. • PM are a complex mixture of solid and liquid particles of organic and inorganic substances.
Suspended Particles Matter (SPM)
• Concentration of atmospheric particulates is severe in some megacities (cities > 10million population) and average levels may range from 200 to 600 µg.m‐3 or pptv.
• Human health associated with high values depend on the nature of particulates. • Examples: those derived from coal and those in the PM10 or PM2.5 categories, have been shown to be hazardous. • NB: PM10 is particulate matter size < 10µm ; • PM2.5 < 2.5µm.
Acute:
• Respiratory problems
• Mortality in vulnerable population segments
Air quality parameters
• Carbon Monoxide (CO): depend on high traffic density – vary from city to city.
• Sulphur dioxide (SO2): Usually produced by coal. Sulfur dioxide conc. Is low in cities that use low coal fuels
• Nitrogen dioxide (NO2): Higher levels expected indoors with poor ventilation where kerosene or natural gas used for heating and cooking. • Ozone (O3): from reaction of gases in the troposphere; trace amounts may result from mass transfer from stratosphere.
• Lead (Pb): Airborne lead depends on the population of cars, the concentration of lead additives in the fuel and availability of unleaded fuel. Concentrations in leaded gasoline vary between 0.1 and 2.0 g.L‐1. • NB. Use of tetraethyl lead to augment the octane number (?) is becoming less.
23
INDOOR AIR QUALITY
Major factors that determine the quality of indoor air
• Many people spend most of their time indoors (home, office, etc.). The atmospheres encountered indoor vary a great deal.
• The nature of the ambient air, outdoor around the building plays a role. In this case, the outdoor atmosphere is influenced by air outside.
• Design and site of the building is important. This will dictate the quality of exchange of indoor atmosphere.
• Nature of materials present in the building such as polymers. The latter could be a source of formaldehydes or other partially oxidized organic compounds.
• Building materials from clays, concrete, etc., may contain traces of radioactive elements such as uranium.
• Activities that take place inside the house. These may include combustion of wood for heating, cooking gas, electric cookers, etc. Cleaning of the house may involve mechanical devices such as vacuum cleaner, that create dust. • Use of cleaning solvents and detergents, insecticide sprays, toilet sprays and air fresheners.
• The materials of house construction may vary from clay‐rich soils or other fresh or baked earth materials.
• In some cases, the homes are open and air exchange is rapid while in others heating may be done over an open fire in a room without a chimney and also a variety of fuels may be used.
• The building materials may range from bricks, stones, wood, various plastics and metals.
• Activities in the homes include: cleaning, cooking, heating over open or closed fires with varying smoke conditions.
COMMON INDOOR AIR CONTAMINANTS
Radioactivity
• Air contaminants refer to levels above the outdoor background level. • 1. Radioactive compounds: • Radioactivity is usually associated with Radon, a noble gas, Rn released by Uranium isotope 238 and also by Thorium isotope 232 with half‐lifes of 4.5 and 14 billion years, respectively. These elements are found in geological materials such as rocks and fossils
• The spontaneous emission of particles and/or energy from atomic nuclei. • The spontaneous emission of radiation from the nucleus of an atom. Radionuclides lose particles and energy through this process of radioactive decay.
24
Radioactive elements
• Radioactive elements, such as uranium (239U) thorium (234Th) and potassium (40K) break down (decay) fairly readily to form lighter atoms e.g Be, B. – The energy that is released in the process is made up of small, fast‐moving particles and high‐energy waves. – These particles and waves are, of course, invisible. (The level of radioactivity of an element varies according to how stable its atoms are). – Other elements with naturally occurring radioactive forms, (isotopes) are carbon (C13), bismuth (210Bi), radon (223R) and strontium (88Sr).
Example: Calculation of Half‐Life of Radioactive Elelemnts
• Consider strontium-90 which has a half-life of
approximately 28 years.
 Initially, at time t=0, the sample is 100% strontium-90
 After 28 years, only half the original amount of
strontium will remain: ½ x 100% = 50%
 After another 28 years, only half of this amount of
strontium-90 will remain: ½ x 50% = 25%
 After another 28 years, only half of this amount of
strontium will remain: ½ x 25% = 12.5%
 and so on.
 At any given time, the amount of strontium-90 that has
undergone decay can be calculated:
amount of strontium-90 decayed = the original amount the amount remaining.
Conti… Radioactivity process
 Radioactivity is a random process that happens naturally as the isotopes in particular elements decay. The isotopes continue to break down over time.  The length of time that is taken for half of the nuclei in an element to decay is called its 'half‐life'.  A half‐life can be very short (milliseconds to hours) or very long (hundreds of thousands of years).
 Radiation also arises from nuclear fission.  Fission can be spontaneous but is usually initiated in a nuclear reactor. Fission is a radioactive process; it releases energy as the heavy nucleus is split into two.
Calculations
• The amount of radioactive isotope remaining
can be calculated:
Nt = No x (0.5)number of half-lives
Nt = amount of radioisotope remaining
No = original amount of radioisotope
number of half-lives = time ÷ half-life
• Example
Calculate the percentage of strontium-90
remaining after 280 years.
Nt = No x (0.5)number of half-lives
Nt = ? %
No = 100%
number of half-lives = time ÷ half-life = 280
÷ 28 =10
Nt = 100 x (0.5)10 = 0.098%
25
Strontium‐90 half‐lifes
Strontium‐90 half‐lifes
%
%
Number
StrontiumTime Strontiumof Half90 that
(years)
90
lives
has
remaining
decayed
0
0
100
0
1
28
50
50
2
56
25
75
3
84
12.5
87.5
4
112
6.25
93.75
5
140
3.125
96.875
6
168
1.5625
98.4375
The Change of Concentration with Time
(Integrated Rate Equations)
Using calculus to integrate the rate law for a first‐order process gives us
ln[A]t ‐ ln[A]0 = ‐kt [A]0 is the initial concentration of A.
or
or
which is in the form
[A]t
[A]0
ln
= -kt
[A]t is the concentration of A at
some time, t, during the course of
the reaction.
Second‐Order Processes
Similarly, integrating the rate law for a process that is second‐order in reactant A, we get
1
1
[A]t = −kt + [A]0
ln[A]t = -kt - ln[A]0
y
= mx + b
therefore, if a reaction is first‐order, a plot of ln [A] versus t will yield a straight line, and the slope of the line will be ‐k.
which,
like the first order process, is also in the form:
y = mx + b
26
Indoor Pollutants Cont.…Volatile Organic Compounds (VOCs)
• Sources are: • Paints: toluene, ethylbenzene, 2‐
isopropanol and butanone.
• Cleaning agents: households solvents, detergents.
• Wood‐building materials such as plywood produc Indoor Air Pollutants Cont.
Emissions from indoor combustion. – This is combustion of fuel that contains VOCs; burning of coal, wood and biomass. – Tobacco smoking is a source of many VOCs including aldehydes, ketones, organic bases such as nicotine, organic acids.
VOC E.g. Polybrominated diphenyl ether
• Polybrominated diphenyl
ether (PBDE) is toxic. General structure is shown below. PBDE is used in commercial household Chemical structure of PBDEs
products such as plastics casings for appliances, in fabrics used for clothing, carpets, etc. • Over 209 various forms (cogeners)
• Affects the nervous system, and may have endocrine disrupting effects
Indoor particulates
– These include: solid aerosols from dust; combustion of coal & biomass material.
– Particle size is usually in the range PM10 (<10µm). Particle size <2µm, can easily enter respiratory track. – Smoking contributes to respirable particulate matter inside a building.
– Polyaromatic hydrocarbons (PAHs) are emitted from coal & biomass.
Sources
• Furnaces and ventilation systems
• Cooking
• Woodstoves / Fire places
• Mould / Mildew:
• Smoking
27
Particles in the atmosphere • Particulate is a term that has come to stand for particles in the atmosphere.
• Particulate matter makes up the most visible and obvious form of air pollution.
• Particles in the atmosphere range from 0.5 mm (size of sand) down to molecular size level (nanometer).
• Particles may consist of either solids or liquid droplets.
• Atmospheric aerosols are solid or liquid particles smaller than 100 µm in diameter.
• Pollutant particles in the 1 nm to 10 µm range are commonly suspended in the air near sources of pollution such as the urban atmosphere, industrial plants, highways and power plants.
Nature of particles
• Very small solid particles include (1 nm‐10 µm ): carbon black, silver iodide, combustion nuclei, sea‐salt nuclei‐ tend to be acidic.
• Larger particles include (100 µm ‐500 µm ) : cement dust, wind blown soil dust, foundry dust and pulverized coal‐ tend to be basic.
• Liquid particles‐mist, include raindrops, fog and sulfuric acid mixture
• Particles of biological origin: viruses, bacteria, bacterial spores, fungal spores and pollen.
• Important atmospheric contaminants‐ mainly inorganic and organic particles.
D ESC R IPTIO N
PA R TIC LES
OF
A TM O SPH ER IC
T erm s
1. A erosol
M eaning
Colloidal-sized
atm ospheric
particle
2. Condensation Form ed by condensation of
aerosol
vapors or reactions of gases.
3. D ispersion
Form ed by grinding of solids,
aerosol
atom ization of liquids or
dispersion of dusts.
4. Fog
D enotes high level of w ater
droplets
5. H aze
D enotes decreased visibility
due to presence of particles
6. M ists
Liquid particles
7. Sm oke
Particles
form ed
by
incom plete com bustion of fuel
Effects of atmospheric particles
• Effects on climate
• Damage buildings
• Reduced visibility & causes undesirable esthetic effects
• NB: Aerosols, natural and anthropogenic, can affect the climate by changing the way radiation is transmitted through the atmosphere.
28
Conti… Effects of atmospheric particles
• All aerosols both absorb and scatter solar and terrestrial radiation. • If a substance absorbs a significant amount of radiation, as well as scattering, it is called absorbing. • This is quantified as the ratio of scattering alone to scattering plus absorption (extinction) of radiation by a particle. Physical behaviour of particles
• Small colloidal particles undergo diffusion processes and coagulate together to form larger particles.
• Mechanism for removal of particles from the atmosphere is mainly through sedimentation & scavenging by rain drops and then precipitation.
• Particle size refers to diameter of the particle but in some cases radius may be used.
Process for particle formation
• Physical Process: particle formation is mainly through disintegration of larger particles > 1 µm
• Many dispersion aerosols originate from natural sources: sea‐spray, windblown dust, volcanic dust.
• Chemical process: Inorganic particles are mainly metal oxides formed by oxidation of the metal by oxygen.
• Organic particles are produced mainly through internal combustion engines. Composition of Inorganic Particles
• Aluminum oxide, iron oxide, calcium oxide and silicon dioxide are due to soil erosion, rock dust, coal combustion.
• Carbon particles‐ due to incomplete combustion
• Sodium and chlorine compounds‐ due to marine aerosols
• Antimony and selenium‐ due to combustion of oil, coal or refuse.
• Lead from combustion of leaded fuels & wastes
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Composition of Organic Particles
• A wide variety of organic compounds most of which are toxic: polycyclic aromatic hydrocarbons (PAHs) such as benzo(a)pyrene, chrysene, benzo‐fluoranthene, acridine.
• Radioactive particles
• Main source of radionuclides in atmosphere is randon: it is a noble gas produced from radium decay.
• Cosmic rays in the atmosphere produce radionuclides which are isotopes of: 7Be , 10Be, 14C, 39Cl, 3H, 22Na, 32P and 33P
•
Air Pollution Control for Particulate Emissions
• It is possible to minimize emissions of aerosol particles from point of source such as thermal electrical generating stations or industrial smelting units.
• Containment of particulate matter is achieved using devices that remove the aerosols from fast moving stack gas stream. Common collection methods include: settling chambers, cyclones, fabric filter, scrubbers, and electrostatic precipitators as shown in the slides that follow.
Control of Particulate Emissions
• Removal of particulate matter from gas streams is the most practiced means of air pollution control.
• Techniques for removal depends on particle size, loading, nature of particles and f scrubbing system.
Methods of Particle Removal
These include: • Sedimentation and inertia, i.e gravitational settling as a continuous process. • Particle filtration using fabric filters that allow gas molecules to pass through but retain the particulate matter. • Scrubbers‐ this involves use of scrubbing liquid which forms small droplets for scavenging particles from the gas stream.
• Catalytic Processes
1.
2.
3.
4.
Industrially Methods based on:
Filtration
Settling rates
Wet scrubbing
Electrostatic precipitation
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Catalytic Converters (Vehicles)
Main Devices Summarized (Industrial Scale Control)
Two‐way
A two‐way (or "oxidation") catalytic converter has two simultaneous tasks:
2CO + O2 → 2CO2
CxH2x+2 + [(3x+1)/2] O2 → xCO2 + (x+1) H2O (a combustion reaction)
Widely used on diesel engines Three‐way
2NOx → xO2 + N2
2CO + O2 → 2CO2
CxH2x+2 + [(3x+1)/2]O2 → xCO2 + (x+1)H2O.
Efficiency Comparison
Coal Combustion
Simplified Schematic
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Coal Combustion
Settling Chambers
• Are the simplest and commonly used.
• Construction includes variety of baffles and open space designed to allow the particles sufficient time to settle under the force of gravity.
• Settling rates are limited by gravity therefore method effective for large particle size >10µm.
• Sometimes Used as a First Stage (precleaner)
Ref: J.M. Beér, Combustion technology developments in power generation in response to environmental challenges, Progress in Energy and Combustion Science, Volume 26, Issues 4–6, August 2000, Pages 301‐327
Gas
Inlet
Detailed Schematic
Gas
Outlet
Dust collection hoppers
H
L
Pro’s & Cons of Settling Chambers
Advantages of Settling Chambers
•
•
•
•
•
•
•
•
•
Low capital cost; Very low energy cost; No moving parts, therefore, few maintenance requirements and low operating costs; Excellent reliability; Low pressure drop through device; Device not subject to abrasion due to low gas velocity; Provide incidental cooling of gas stream; Temperature and pressure limitations are only dependent on the materials of construction; and Dry collection and disposal.
Disadvantages of Settling Chambers
• Relatively low particulate matter collection efficiencies, particularly for particulate matter less than 50 µm in size; • Unable to handle sticky or tacky materials; • Large physical size (Take up too much space)
• Trays in multiple‐tray settling chamber may warp during high‐temperature operations
Cyclones
• Cyclones are cone‐shaped devices that cause the waste gas stream to swirl rapidly in spiral fashion causing larger particles to move towards the wall of the cone by centrifugal force.
• Once in contact with the wall, the particles slide down the inner surface of the cone to a collection container below it.
• Stoke’s law determines the extent of removal of particles but the settling rates can be greatly enhanced by the increased force due to cyclone action. In this case removal of particles <10µm can be achieved.
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Stokes Law
t 
(  p -  a ) C g d 2p
Cyclone
18
Where vt = terminal velocity of particles in m.s‐1; pp = density of particle in g.cm‐3
Pa = density of air = 1.2 x 103 g.m‐3 at Po and 25° C; C = Stokes correction factor for assuming spherical shape and discontinous of fluid interactions when the particle size is small compared with the molecular mean path in air. g = acceleration due to gravity = 9.8 m.s‐2
dp = particle diameter in meters and η = viscosity of air = 1.9 x 10‐2 g.m‐1.s‐1 at P° and 25° C
Fabric Filters: Filtration
Filtration
• Fabric filter or bags operate on a similar principle as vacuum cleaner.
• The air stream is made to pass thro a porous fabric material and is effective for particulates size in the range 0.01 ‐ 10µm range.
• Bags or fabric filters are sensitive to temperature and humidity. The fine particles clog the filters and there4 must be periodically cleaned.
Advantages
•
•
•
•
•
•
•
High Collection Efficiency (>99%)
Effective for a Wide Range of Dust Types
Modules Can be Factory Assembled
Operates Over Wide Range of Gas Flow Rates
Reasonably Low Pressure Drop
Good Efficiency for Small Particles
Dry Collection and Disposal
Disadvantages
•
•
•
•
•
Large Footprint
Temperature Limitations
Requires Dry Environment
Fire or Explosion Potential
High Maintenance Cost
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Scrubbers
Liquid Scrubber
• Scrubbers allow gas stream to be in contact with a fine mist or spray of water.
• The water droplets capture many small particles and these settle more rapidly into a collector container.
• Scrubbers come in different designs as shown below.
Absorption Method
Adsorption
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Electrostatic Precipitator
• Electrostatic precipitator causes the particles in a gas stream to become charged by electrons produced thro an electrical discharge between two electrodes.
• The negatively charged particles then migrate to the positive electrode and are collected and removed from the emission stream. Positively charged particles move to negative electrode.
Advantages
•
•
•
•
•
•
•
High Collection Efficiency
Dry Collection and Disposal
Small Pressure Drop
Capable of Handling Large Gas Flow Rates
Low Electrical Power Requirements
Low Maintenance
Disadvantages
Electrostatic Precipitator
Disadvantages
• High Capital Cost
• Particle Resistivity Limitations
• May Require Injection of SO3 or NH3 to Control Resistivity
• Relatively Large Footprint
• Special Precautions for Safe Operating at High Voltage
Minimize Emission from point Source: example SO2
EXAMPLE: Sulfur Dioxide Control
• Minimize emissions of aerosol particles from point of source such as thermal electrical generating stations or industrial smelting units.
http //www.apt.lanl.gov/projects/cctc/factsheets/puair/adflugasdemo.html
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Clean Coal Technologies
Why Clean Coal?
Impact of Coal
• Mining Coal Leads to:
• Advanced Flue Gas Desulfurization
Demonstration Project - e.g. of a
series of initiatives
• Others Include
– Carbon Capture and storage
– Underground coal gasification
Underground Coal Gasification (UCG)
– Water scarcity
– Water Pollution
• Burning Coal Leads to
– SOx, NOx, particulates, CO2, fly ash
Coal Usage
• In RSA Coal used to generate over 90% of the electricity
• Globally Coal accounts for more than 40% of Electricity produced
• Only realistic technology for next 20 –
50 years
Carbon Capture and Storage
Historically a lot of the work was done in the former USSR
Resurgence in interest (China, Australia, Europe, Americas & RSA
SASOL
Majuba
http://www.eskom.co.za/live/content.php?Item_ID=14077
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CO2 Sequestration
•
•
•
•
The Case for Renewable Energy
Clean and treat the CO2 then store it Currently used in enhanced oil recovery
Options to use saline (very salty) aquifers
Abandoned coal mines, other geologic caverns etc
Global warming over
the past millennium
• Very rapidly we have entered
uncharted territory – what some call
the anthropocene climate regime.
• Over the 20th century, human
population quadrupled and energy
consumption increased sixteenfold.
• Near the end of the last century, we
crossed a critical threshold, and
global warming from the fossil fuel
greenhouse became a major, and
increasingly dominant, factor in
climate change.
• Global mean surface temperature is
higher today than it’s been for at
least a millennium.
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Solar Energy
10 TW of Energy by 2050: But How?
•
Carbon
Neutral
Energy
(fossil
fuel
in
conjunction with carbon sequestration)
–
Need to find secure storage for 25 billion
Thermal Conversion
metric tons of CO2 produced annually (equal
Infrared Photons
to he volume of 12500 km3 or volume of lake
E=h
Photoconversion
Energetic Visible Photons
E = hc/ = 119627/ (kJ/mole)
superior!)
•
•
Nuclear Power
– Requires construction of a new 1‐GW electric nuclear fission plant everyday for the next 50 years
Renewable Energy Sources
– hydroelectric resource 0.5 TW
– from all tides & ocean currents 2 TW
– geothermal integrated over all the land area 12 TW
– globally extractable wind power 2‐4 TW
– Solar energy striking the earth 120,000 TW !!!
Heating
Electricity
Generation
Photosynthesis
Photovoltaics
Food/Fuel
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