Chapter 2: Protecting the Ozone Layer

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Chapter 2:

Protecting the Ozone Layer

1. Why do we need to do to protect the ozone layer? 2. Isn’t ozone hazardous to

human health?

3. Why is the ozone layer getting smaller?

4. What can we do (if anything) to

help stop the depletion of our ozone layer?

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Ozone Formation

Energy + 3 O₂ 2 O₃

Ozone is an allotropic form of oxygen. The name ozone comes from a Greek word meaning “to smell”

Energy must be absorbed

(endothermic) for this reaction to occur.

Element Allotropes oxygen O₂, O₃

carbon graphite, diamond, buckminister fullerenes _2.1 An allotrope is two or more forms of the same element that differ in

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O

8

16.00 Atomic number (Z)

Mass number (A)

-The number of protons.

(nuclear charge)

-The sum of the protons and neutrons.

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The electrons in the outermost

energy levels are called

valence electrons.

2.2 The group number (of the representative elements) on the

periodic table tells you the number of valence electrons.

1A

2A 3A 4A 5A 6A 7A

8A

Group 1A: 1 valence electron

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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:

C-12, C-13, C-14

also written as:

12

C

13

C

14

C

• Carbon 14 ( ₆C)

– Used in dating old materials – Atomic number: 6

– Atomic mass: 14

– # of neutrons = Atomic mass-Atomic #

– 6 protons, 6 electrons, 8 neutrons

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Representing molecules with

Lewis structures

:

2.3

Consider water, H

₂

O:

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:

O

H

lone pair

bonding pair

O

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2.3

Representing molecules with

Lewis

structures

:

Typical valence for selected atoms = the # of bonds an atom typically forms

Element Typical

valence Classification

H,

X (X= F, Cl, Br, I)

1 monovalent

O 2 divalent

N 3 trivalent

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2.3

Representing molecules with

Lewis

structures

:

Multiple bonds

O

H

C

H

Triple bond Double bond

Occasionally a single Lewis structure does not adequately represent the true structure of a molecule, so we use resonance forms:

N O

O O

N O

O O

N O

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2.4

The Nature of Light

Low E

High E

Wavelength (_) = distance traveled between successive peaks (nm). Frequency () = number of waves passing a fixed point in one second

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The Electromagnetic Spectrum

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Visible:  = 700 - 400 nm

M

UD

A

S

I

R

Infrared (IR) : longest of the visible spectrum, heat ray absorptions cause molecules to bend and stretch.

Microwaves: cause molecules to rotate.

Short range: includes UV

(ultraviolet), X-rays, and gamma rays. Decreasing wavelength

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2.5

The energy of a photon of electromagnetic radiation

is calculated by:

E = h

_



where

h

= 6.63 x 10

-34

J

.

s (Plank’s constant)

The wavelength and frequency of electromagnetic

radiation are related by

:

c =

_

_{}

where

c

= 3 x 10

8

m/s (the speed of light)

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UV radiation has sufficient energy to cause molecular bonds to break

2.5 What is the energy associated with a photon of light with a wavelength of 240 nm?

C = _

= C



x 108 m/s

E = h _

240 nm x

nm

10-9_{m =} 1.3 x 1015 s-1

E = (6.63 x 10-34 _J._{s) (1.3 x 10}15_s-1₎

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The Chapman Cycle

2.6

A

steady state

condition

 ≤

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Factors influencing steady-state

concentration of O

₃₃

• Intensity of UV radiation

– Expect O3 concentration will depend on season and

sun cycle

• Concentration of O

₂

and other reacting species

• Temperature

• Reaction rates (some of the steps are faster

than others)

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2.7

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.

An Australian product uses “smart bottle” technology; bottle color changes from white to blue when exposed to UV light.

The most deadly form of skin cancer, melanoma, is

linked with the intensity of UV radiation and the latitude at which you live.

As  decreases, damage

to DNA increases.

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O

₃

concentrations over time

http://earthobservatory.nasa.gov/cgi-bin/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

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Natural causes of O

₃₃

destruction

_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•

Example of a chain reaction. Free radicals are very destructive

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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 high-temp 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…

energy

Source of VOC and NO in the Upper Atmosphere leading to Tropospheric Air Pollution Chemistry

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CFCs seem to be the culprit

Properties and uses of CFC • Stable

• 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

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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.

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:

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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!!!

• 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

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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

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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.

http://earthobservatory.nasa.gov/Study/Tango/

• 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 & Cl₂

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Summer conditions replenish hole

spring

summer

• In summer, sunlight warms

the atmosphere

– Ice crystals evaporate

• 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.

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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)

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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.

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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.

• DON’T want to go back to previous technologies

(NH3, SO2 as refrigerants)! • Need to balance 3

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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

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Hydrochlorofluorocarbons

(HCFCs)

• Replacing one or more of the halogen atoms

with hydrogen makes the compound more

reactive

– Compounds decompose long before

reaching stratosphere

• However:

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Suitable HCFCs

• 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.

Decompose in the troposphere more readily than CFCs

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Hydroflurocarbons (HFCs)

• 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.

H H C C

http://www.eia.doe.gov/oiaf/1605/vr98rpt/chapter6.html

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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