Chapter 2:
Protecting the Ozone Layer
1.
2.
3.
4.
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Why do we need to do to
protect the ozone layer?
Isn’t ozone hazardous to
human health?
Why is the ozone layer
getting smaller?
What can we do (if anything)
to
help stop the depletion of our
ozone layer?
Ozone Formation
Energy + 3 O2
2 O3
Energy must be absorbed
(endothermic) for this reaction to
occur.
Ozone is an allotropic form of
oxygen. The name ozone comes
from a Greek word meaning “to
smell”
An allotrope is two or more forms of the same element that differ in
their chemical structure and therefore their properties.
Element
oxygen
carbon
Allotropes
O2, O3
graphite, diamond, buckminister fullerenes
2.1
Mass number (A)
16.00
O
8
-The sum of the protons
and neutrons.
Atomic number (Z) -The number of protons.
(nuclear charge)
2.2
2.2
The electrons in the outermost energy levels are called
valence electrons.
The group number (of the representative elements) on the
periodic table tells you the number of valence electrons.
Group 1A: 1 valence electron
1A
8A
Group 3A: 3 valence electrons
2A
3A 4A 5A 6A 7A
2.2
Isotopes are two or more forms of the same element
(same number of protons) whose atoms differ in
number of neutrons, and hence in mass.
Isotopes of carbon: • Carbon 14 ( 6C)
– Used in dating old materials
C-12, C-13, C-14
– Atomic number: 6
also written as:
12C 13C
14C
– Atomic mass: 14
– # of neutrons = Atomic massAtomic #
– 6 protons, 6 electrons, 8 neutrons
*The average mass is very close to 12.000 b/c 6C is by far the most abundant
isotope.
2.2
Representing molecules with Lewis structures:
Consider water, H2O: H
O
H
1. Find sum of valence electrons:
1 O atom x 6 valence electrons per atom = 6
+ 2 H atoms x 1 valence electron per atom = +2
8 valence electrons
2. Arrange the electrons in pairs; use whatever electron pairs
needed to connect the atoms, then distribute the remaining
electron pairs so that the octet rule is satisfied:
lone pair
O
H
H
bonding pair
2.3
Representing molecules with Lewis structures:
Typical valence for selected atoms = the # of bonds an atom typically forms
Element
Typical
valence
1
Classification
O
2
divalent
N
3
trivalent
C
4
tetravalent
H,
X (X= F, Cl, Br, I)
monovalent
2.3
Representing molecules with Lewis structures:
Multiple bonds
O
H
O
Double bond
C
C
H
Triple bond
Occasionally a single Lewis structure does not adequately represent
the true structure of a molecule, so we use resonance forms:
O
O
O
O
N
N
N
O
O
O
O
O
2.3
The Nature of Light
Low E
High E
Wavelength (l) = distance traveled between successive peaks (nm).
Frequency (n) = number of waves passing a fixed point in one second
(waves/s or 1/s or s-1 or Hz).
2.4
The Electromagnetic Spectrum
The various types of radiation seem different to our senses, yet they differ only
in their respective l and n.
2.4
Visible: l = 700 - 400 nm
M MUDAS I R
Decreasing wavelength
Infrared (IR) : longest of the visible
spectrum, heat ray absorptions cause
molecules to bend and stretch.
Microwaves: cause molecules to rotate.
Short l range: includes UV (ultraviolet),
X-rays, and gamma rays.
2.4
The wavelength and frequency of electromagnetic
radiation are related by: c = ln
where c = 3 x 108 m/s (the speed of light)
The energy of a photon of electromagnetic radiation
is calculated by: E = h n
where h = 6.63 x 10-34 J.s (Plank’s constant)
Energy and frequency are directly relatedhigher frequency means higher energy.
2.5
What is the energy associated with a photon of light with a wavelength of 240 nm?
C = ln
E=hn
n =C
l
E = (6.63 x 10-34 J.s) (1.3 x 1015 s-1)
E = 8.6 x 10-19 J
n = 3 x 108 m/s
-9 m = 1.3 x 1015 s-1
10
240 nm x
nm
UV radiation has sufficient energy to cause molecular bonds to break
2.5
2.6
The Chapman Cycle
l≤
l≤
A steady state
condition
2.6
Factors influencing steady-state
concentration of O3
• Intensity of UV radiation
– Expect O3 concentration will depend on season and
sun cycle
• Concentration of O2 and other reacting species
• Temperature
• Reaction rates (some of the steps are faster
than others)
• *Many of these factors depend on altitude, so O3
concentrations also depend on altitude.
Biological Effects of Ultraviolet Radiation
The consequences depend primarily on:
1. The energy associated with the radiation.
2. The length of time of the exposure.
3. The sensitivity of the organism to that radiation.
The most deadly form of skin cancer, melanoma, is
linked with the intensity of UV radiation and the latitude
at which you live.
As l decreases, damage
to DNA increases.
Note the logarithmic
scale on the y-axis.
An Australian product uses
“smart bottle” technology;
bottle color changes from
white to blue when exposed
to UV light.
2.7
O3 concentrations vs. altitude
O3 concentrations over time
http://earthobservatory.nasa.gov/cgibin/texis/webinator/printall?/Library/Ozone/ozone_2.html
1 Dobson unit (DU) = 1 O3 molecule per 1 billion (109) molecules of air
320 Du is average ozone level over the northern hemisphere
250 DU is typical at the equator
“hole” is the area over Antarctica with < 220 DU
Natural causes of O3 destruction
1) Reactions with water vapor and its breakdown products:
H2O + UV photon  H• + •OH
These free radical products can
participate in reactions that convert O3 to O2:
O3 + H•  O2 + •OH
O3 + •OH  2O2 + H•
Free radicals very
reactive b/c they
have unpaired eOctet Rule is NOT
satisfied
Example of a chain reaction. Free radicals are very
destructive
•
2) Reactions with stratospheric nitrogen
monoxide (NO)
Formation of natural NO: N2O + O
 2NO
– N2O produced by
microorganisms
• Formation of man-made NO: N2 +
O2 + energy  NO
– Energy is from lightning or hightemp engines of supersonic
transport airplanes (like the
Concorde) that are designed to
fly at altitudes of 15-20 km (the
region of the ozone layer)
• Lewis dot structure of NO?
• Turns out that NO is also a free
radical…
•
However, these natural causes do not
explain the amount of ozone
destruction seen in recent years…
energy
Source of VOC and NO in the Upper
Atmosphere leading to Tropospheric Air
Pollution Chemistry
CFCs seem to be the culprit
Properties and uses of CFC
•
• Stable
• Nontoxic
•
• Nonflammable
• Low boiling point (makes it a
good refrigerant)
• Cheap
• Widely available
Originally developed as refrigerants to
replace the toxic ones that had previously
been used (SO2 and ammonia—NH3)
Also used as:
– Solvent for computer chips
– Propellants for aerosol cans
– Styrofoam
– Fire extinguishers
Before we knew about the ozone hole…
• Two scientists, Roland &
Molina, set out to study fate of
atmospheric CFC molecules
• The reactions they
hypothesized would occur
suggested that CFCs could
destroy significant amounts of
O3
– The destruction of O3 by
CFCs later shown by
scientific experiments.
• ~ 5 years after release in
troposphere, CFCs make
their way into the
stratosphere.
• A high-energy photon
corresponding to
wavelengths  220 nm has
enough energy to break the
chlorine-carbon bond:
CCl2F2 + photon  •CClF2
+ •Cl
Revolutionized modern life! Throughout the 1960s and 70s, cheap air
conditioning using CFCs helped spur the booming growth of Southern
cities such as Atlanta, Dallas-Fort Worth and Houston.
Reactions with •Cl
•Cl +O3  O2 + •OCl
•OCl + O  •Cl + O2
O3 + O  2O2
You can see that once •Cl is created
that it begins a chain reaction—it can
destroy many (~ 100,000!) O3
molecules.
•Cl is acting as a catalyst—it is not
used up in the reaction
The presence of •Cl will affect the
Chapman cycle balance!!!
Removal of •Cl
• Carried to lower
atmosphere by winds—
once in lower atmosphere,
there are many other
compounds for •Cl to react
with.
• If it encounters another •Cl
the following reaction may
occur: •Cl + •Cl  Cl2
• May react with other
species to form stable
compounds such as HCl
and ClONO2
Proof of relationship b/w ozone
depletion and stratospheric Cl
•
•
•
75-85% is from human activity
– CFCs
– Rocket fuel from space
shuttles (<1%)
15-20% is from methyl chloride
– Most from natural sources
and burning of biomass
~2% - Large, explosive volcanoes
– Increase amount of HCl
which can be converted to •Cl
– Only a few volcanoes have
enough explosive power to
project material into the
stratosphere.
Why is the hole over Antarctica?
• A special mechanism appears to be in effect here, related to the fact that
the lower stratosphere over the South Pole is the coldest spot on Earth.
– During the Antarctic winter (June to September), a strong circumpolar
wind develops in the middle to lower atmosphere
• keeps warmer new air (w/ new O3) from entering
– Temps can get as low as -90oC!
– At temps < 80oC, clouds of ice
crystals containing sulfates and
nitric acid can develop in the
stratosphere.
• In spring, when the sun comes out,
chemical reactions occurring on the
surface of these cloud particles convert
otherwise safe (non-ozone depleting)
molecules to more reactive species.
– ClONO2 & HCl converted to HOCl & http://earthobservatory.nasa.gov/Study/Tango/
Cl2
– HOCl
•Cl + •OH ; Cl2
2 •Cl
UV light
UV light
Summer conditions replenish hole
• In summer, sunlight
warms the atmosphere
– Ice crystals evaporate
summer
• No ice crystals means the
conversion of ClONO2 &
HCl to more reactive
species halts.
– Air from lower latitudes
flows into the polar
regions, replenishing
depleted ozone levels.
– Thus, “Ozone Hole” is
largely replenished by the
end of summer.
• Same process is developing in the Arctic
spring
The Global Response
• CFCs in spray cans banned in North America in
1978, and use as foaming agents for plastics
discontinued in 1990.
• 1987 Montreal Protocol—established a schedule to
reduce production & consumption of CFCs.
• The US and 140 other countries
agreed to a complete ban on CFC
production after Dec. 31, 1995.
• Other ozone depleting compounds to
be eliminated b/w 2002-2010:
•CFBr compounds (fire-fighting)
•CCl4 (an industrial solvent)
•CH3Br (agricultural fumigant)
Global production of CFCs, 1950-2000
Black-market CFCs
• While the protocols set dates for the halt
of production, sale of existing stockpiles
will remain legal.
– US govt. is promoting conversion to less
harmful substitute refrigerants by imposing
a tax of $5.35/lb on CFCs
• Price of CFCs has risen $1/lb in 1989 to $15/lb
(as of 2003)
• Creates a black-market for CFCs, which are 2nd
only to illegal drugs as the most lucrative illegal
import.
• “Freon busts” in 1997 led to the confiscation of
12 million lbs of illegal CFCs
Ban on CFCs is making a
difference--slowly
• Stratospheric concentration
of chlorine peaked at 4.1
ppb in the mid-90s and is
now declining.
• However, the stratospheric
chlorine is not expected to
reach 2 ppb* until 2050 or
2075.
• *Antarctic ozone hole first
appeared when chlorine
levels reached 2 ppb.
What to use
instead?
• DON’T want to go back to
previous technologies
(NH3, SO2 as refrigerants)!
• Need to balance 3
undesirable properties and
come up with the best
compromise
– Toxicity
– Flammability
– Extreme Stability
Fluorocarbons (FCs)
• Made of only carbon and fluorine
– Non-toxic
– Nonflammable
– Don’t decompose in the stratosphere
• Why? The C-F bond is stronger—the energy
requirement to break this bond is higher than
that of a UV photon.
• The negative: No destruction
pathways—the molecules will
accumulate in the atmosphere and
contribute to the greenhouse effect by
absorbing IR radiation.
Hydrochlorofluorocarbons
(HCFCs)
• Replacing one or more of the halogen atoms
with hydrogen makes the compound more
reactive
– Compounds decompose long before
reaching stratosphere
• However:
– Too many H atoms increases flammability
– Too many Cl atoms increases toxicity
Suitable HCFCs
Decompose in the
troposphere more
readily than CFCs
• HCFC-22 is the most widely-used HCFC
– Used in air-conditioners and in the production of
foamed fast food containers
– 5% ozone depleting potential of CFC-12
• HCFC-141b used to form foam insulation
– 250 million lbs produced worldwide in 1996 for
this purpose
• Montreal Protocol calls for a production ban of
HCFCs by 2030.
CFC-12
Hydroflurocarbons (HFCs)
H
• Example: HFC-134a
• Has no Cl atoms to
interact with ozone
• The 2 H atoms facilitate
decomposition in the
lower atmosphere,
without making it
flammable under normal
conditions.
C
C
H
US Emissions
http://www.eia.doe.gov/oiaf/1605/vr98rpt/chapter6.html
Timeline
•
•
•
•
•
•
•
•
•
1928 – CFCs invented
1950s-1970s – Consumption and use of CFCs rises rapidly
1971 – CFCs measured in the atmosphere
1974 – Rowland & Molina link CFCs with ozone depletion
1985 – First scientific assessment of stratospheric ozone levels
1987 – “Smoking gun” evidence linking decreasing ozone levels
with increasing stratospheric chlorine levels
– Montreal Protocol established a schedule to reduce
production & consumption of CFCs.
1991 – Multinational fund established to provide financial and
technical assistance to developing countries to enable
them to comply with control measures.
- B/w 1991 & 2004 $1.6 billion given to more than 100 countries
1995 – Production ban of CFCs in US and 140 other countries
goes into effect
– Rowland & Molina win Nobel Prize for their work
2030 – Production ban of HCFCs goes into effect