Chapter 2: Protecting the Ozone Layer

Isn’t ozone hazardous to human health?

What can we do (if anything) to

help stop the depletion of our ozone layer?

Why do we need to protect the ozone layer?

Why is the ozone layer getting smaller?

Ozone = O3

Ozone is an allotrope of oxygen.Allotropes have the same elemental composition, but different formulas and structures.Allotropes can have very different chemical and physical properties.

Graphite and diamond are both allotropes of carbon

Ozone can be produced by man-made activities

From high voltage dischargesAs a secondary pollutant from car exhaust

Ozone is a very reactive chemical, and a hazardous pollutant at ground level

Bleaches fabric/paperKills bacteriaCan cause lung damageHas a short lifetime due to it's high reactivity

“Ozone Layer” Not actually a layer, but a region ~10-18 miles above the earth's surface, in the stratosphere, where ozone is created by a natural process.

The Regions of the Lower Atmosphere

Atmospheric pressure changes with altitude2.1

2.1The ozone layer is a region in the stratosphere with maximum ozone concentration.

Stratospheric ozone is produced via interaction of oxygen and ultraviolet light from the sun:

step 1: O2 + energy → 2 O

step 2: O2 + O → O

3

overall reaction: 3 O2 + energy → 2

O3

(O3 can be lost by: O

3 + O → 2 O

2)

The stratospheric ozone concentration is very low. All of the ozone would form a pure layer over the earth less than 1 cm thick.

Stratospheric ozone blocks harmful ultraviolet (UV) light from the sun.

Most atmospheric gases are not able to block these rays.

To understand why some gases absorb UV, and how UV light damages the environment, we need to examine the atomic and electronic structures of matter.

Particle electrical charge mass (g) mass

(amu) Location

Proton +1 1.66E-24 1 nucleus

Neutron 0 1.66E-24 1 nucleus

Electron -1 9.03E-28 1/1839electron

cloud(orbitals)

amu = atomic mass unit

Most of the mass of an atom is in the small, dense nucleus.

The radius of an atom is about 100,000 times larger than the radius of the nucleus.

Electrons are located around the nucleus in orbitals.

Orbitals – not distinct like planetary orbits, but 3-D regions where electrons can probably be found (“electron clouds”).

The position of the outermost electrons determines the radius of the atom.

Most of an atom is empty space

An element is defined by the number of protons in its nucleus:

This is called the atomic number (Z)

The periodic table is arranged by atomic number, starting with #1 in the upper left and increasing across each period.

Most of the mass of an atom is in the nucleusmass number (A) = protons(p) + neutrons(n)

Atoms of elements are electrically neutral, so:# of electrons (e-)= # of protons (p)

All atoms of an element have the same atomic number.

All atoms of an element will therefore have the same number of electrons.

All atoms of an element may not have the same number of neutrons, so may also have different mass numbers.

Atoms of an element with different numbers of neutrons are called isotopes.

Isotope Notation

• Z = atomic # = # of protons = p

• A = mass # = p + n

• Isotopes are denoted using the chemical symbol, X, or the element name:

ZAX = AX = X-A = element name-A

So a carbon isotope with 6 neutrons could be written as:

612C = 12C = C-12 = Carbon-12

Valence electrons (ve) Atoms bond to attain more stable

configurations of electrons. Electrons occupy “shells” or “energy levels”

around an atom. Each shell can only hold a certain # of

electrons. The inner shells are filled first. Only electrons in the outer shell can

participate in bonding.

Valence electrons (ve)

The outer electrons are called “valence electrons” or “bonding electrons”

Most atoms want to get 8 ve (octet rule), which they do by gaining, losing, or sharing electrons.– Main exception: Hydrogen wants 2 ve

Noble gases already have 8 ve, which is why they almost never react.

Group Valence Electrons (v.e.)

1 (I A) 1

2 (II A) 2

13 (III A) 3

14 (IV A) 4

15 (V A) 5

16 (VI A) 6

17 (VII A) 7

18 (VIII A) 8

Lewis Structures

Use chemical symbols and valence electrons to show bonding Used to predict chemical structures Used to predict molecular geometries This information is used to predict

chemical and physical properties

To draw Lewis Structures for molecules:Add up valence electrons for all atomsDraw backbone ('skeletal structure') for the

molecule Link atoms with single bonds (two

shared electrons) Usually elements down and to left on

PT are central (carbon is usually central)

H and Group 17 elements almost always only form single bonds

Add electrons to satisfy the octet rule for all atoms: Add electrons to outer atoms first Electrons usually placed in pairs above,

below, right, or left of symbol Electrons between two atoms are

bonding electrons, and count toward valence electrons for both atoms

2 shared e- - one line – single bond 4 shared e- - two lines – double bond 6 shared e- - three line – triple bond

Use/move unshared pairs of electrons to form double and triple bonds, if needed, to satisfy octets

Double check: Final structure must have the exact

number of total valence electrons Check that all atoms obey the octet rule [(sum of desired v.e. over all atoms) –

(total valence electrons)]/2 = total # of bonds

Representing molecules with Lewis structures:

2.3

Consider water, H2O:

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:

OH H

lone pair

bonding pair

OH H

In many cases, several valid Lewis Structures may be drawn for a molecule.

More advanced chemistry can eliminate some of these structures.

However, not all molecules can be accurately represented by Lewis Structures The actual structure may be an

average of several Lewis Structures These are called Resonance Structures

2.3

Representing molecules with Lewis structures:

Multiple bonds

O O H C C H

Triple bondDouble bond

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

N

O

O ON

O

O ON

O

O O

2.3

Try these; draw valid Lewis structures for:

HNO3

CO2

H2S

H2SO4

C OO

S

H H

Sulfur is under oxygen; think of H2O

S

OO

HH

O O

S

OO

HH

O O

S

OO

HH

O O

N

O

O O

H

Sulfur is in the 3rd period: it can have an expanded octet

Can you draw other valid Lewis structures for HNO3?

Molecular GeometryThe geometric shape of a molecule may be

determined from the Lewis StructureGroups of valence electrons on an atom

repel each other, and move to maximize their angles of separation

Groups of electrons are: Unshared pairs of electrons single bonds double bonds triple bonds

With two groups – 180° between groups Linear

With three groups – 120° between groups Trigonal

With four groups – 109.5° between groups Tetrahedral

The actual shape of the molecule is dependent on the actual bonds.

Electronic Geometry Use all groups of electrons to

determine the angles between groups

Molecular Geometry Look only at bonds to determine the

shape of the molecule Angles between bonds determined by

electronic geometry For central atoms with no unshared

electrons, el. geom. = mol. geom.

Depending on it's geometry and elemental composition, a molecule may bend, twist, stretch, and compress.The bonds act like springs, and these movements vibrate at specific frequencies, like a tuning fork.Light also vibrates at specific frequencies, and if these frequencies match, the molecule can absorb that light.

Visible light is part of the electromagnetic spectrum.

Radio waves, microwaves, infrared, ultraviolet, x-rays, and gamma rays are also types of EM radiation, often generically called light.

Light acts like a particle and a wave: The smallest particle of light is a photon Photons have a wavelength and

frequency, as do sound or ocean waves.

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

(waves/s or 1/s or s–1 or Hz).

Wavelength (λ) – the distance at which a wave starts repeating itself, often measured from peak to peak.

Frequency () – the number of waves that pass a fixed point per second.

Velocity – how fast a wave moves

For almost any wave, it's velocity is constant in a given medium.

The speed of light (c) is constant in a vacuum and in air.

The Electromagnetic Spectrum

2.4

The various types of radiation seem different to our senses, yet they differ only in their respective and

Wavelength, frequency, and the speed of light are related by:

c = λ* So as the frequency increases, the wavelength drops, and vice versa.

The energy (E) of a photon increases with it's frequency:

E = h* (h is a Plancks constant)

When a molecule absorbs a photon, this energy can cause:

stretching bending twisting heating ejecting electrons breaking of bonds

Visible: = 700–400 nm

R O Y G B I V

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

2.4

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 1015 s–1)

E = 8.6 x 10–19 J

When bonds are broken in living organisms, mutations or death may occur, depending on which bonds (DNA or RNA) and how many bonds are affected.Gamma rays and x-rays can penetrate the body and break bonds internally.Ultraviolet light is not able to penetrate the body, but can cause mutations on the surface of plants and animals.Photons of visible light are not energetic enough to break most bonds

There are three classes of UV light:

1) UV-A (320-400 nm) least energetic UV Much UV-A reaches the earth's surface

2) UV-B (280-320 nm) more energetic than UV-A blocked by ozone in the stratosphere

3) UV-C (200-280 nm) most energetic UV

blocked by O2 and ozone

2.6

The Chapman Cycle

2.6

A steady state condition

≤

≤

+

+

The ozone concentration in the stratosphere is reduced by some man made pollutants:

Chlorofluorocarbons (CFCs) – contain Cl, F and C – produce chlorine free radicals in the stratosphere

Free radicals are very reactive because they need another e- to get an octet of ve.

Cl٠ + 2O3 → 3 O

2 + Cl٠

Cl٠ is a catalyst – it speeds up this reaction, but is not consumed. One of these chlorine radicals can destroy many ozone molecules

2.9

Experimental analyses show that as ClO. concentrations increase, ozone concentration decreases.

First, UV radiation breaks a carbon-halogen bond:

Photon < 220 nm) + CCl2F2 .CClF2 + Cl. (free radicals)

2.9

How CFCs Interact with Ozone

CFC-11 CFC-12Freon 11 Freon 12

trichlorofluoromethane dichlorodifluoromethane

C

F

Cl

Cl

Cl C

F

F

Cl

Cl

CCl3F CCl2F2

2.9

2 Cl. + 2O3 2 ClO. + 2 O2

2 ClO. ClOOCl

The chlorine radical attacks an O3 molecule:

Then two chlorine monoxide radicals combine:

UV photon + ClOOCl ClOO. + Cl.

ClOO. Cl. + O2

The ClOOCl molecule then decomposes:

The net reaction is: 2 O3 3 O2

The Cl. radicals are free to attack more O3

The Cl. radicals are both consumed and generated; they act as catalysts

The most common uses for CFCs have been: air conditioners aerosols plastic foams

The CFCs have been banned in most countries, and replaced with HCFCs, which do not pose the same threat to the ozone.

But the CFCs are very inert, so have a long lifetime in the environment, and will still affect the ozone layer for many years to come.

2.12

HCFCs are alternatives to CFCs: they decompose more readily in the troposphere so they will not accumulate to the same extent in the stratosphere.

HCFC-22chlorodifluoromethane

CHClF2

HCFC-141bdichlorofluoroethane

C2H3Cl2F

C

F

F

Cl

H C

Cl

F

Cl

C

H

H

H

Some effects of increased UV light due to reduction of O

3 concentration in the ozone

layer: Skin cancer Skin damage Eye damage Harms land and aquatic animals Harms plants – suppressed growth

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.