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Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
CHEMICAL
MONITORING AND
MANAGEMENT Year 12 HSC
ABSTRACT This report consists of research gathered from
secondary sources relating to the outcomes of Core
Module 3: Chemical Monitoring and Management.
Questions regarding our research will also be
addressed in this report
Bilal Fouzi Due Date: Week 2, Thursday, 15th May
Pg. 1 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Describe the composition and the layered structure of the earth’s atmosphere
Contrary to common belief the earth’s atmosphere is not simply just its
composition. The atmosphere is conclusive of all available gaseous matter in and
even around our planet. The Earth's atmosphere is divided into five main layers:
the exosphere, the thermosphere, the mesosphere, the stratosphere and the
troposphere. The atmosphere thins out at each level higher until the gases scatter
out into space.
The troposphere is the layer closest to Earth's surface and where all life inhabits. It
is 4 to 12 miles (7 to 20 km) thick, constituting half of Earth's atmosphere. Over
90% of Earth's gases are in the troposphere. Air is warmer near the ground and
gets colder higher up. Nearly all of the water vapor and dust in the atmosphere are
in this layer and that is why clouds are found here. [1]
The stratosphere, starts above the troposphere and ends about 31 miles (50 km)
above ground. The famous protective layer of ozone is abundant here and it heats
the atmosphere while also absorbing harmful radiation from the sun. The air here
is very dry, and it is about a thousand times thinner here than it is at sea level.
Because of that, this is where jet aircraft and weather balloons fly. [2]
The mesosphere is marked approximately at 31 miles (50 km) and extends to 53
miles (85 km) high. The coldest part of Earth's atmosphere is the mesopause-the
top of the mesosphere-with temperatures averaging at about minus 130°F (minus
90°C). Since jets and balloons don't go high enough to this layer, and satellites and
space shuttles orbit above, this layer is perhaps the hardest to study. Scientists
however, do know that meteors burn up in this layer. [3]
The thermosphere extends from about 56 miles (90 km) to anywhere between 310
and 620 miles (500 and 1,000 km). Temperatures can get up to 2,700 °F (1,500°C)
at this altitude. The thermosphere is considered part of Earth's atmosphere, but air
density is so low that most of this layer is what is normally thought of as outer
space. In fact, this is where space shuttles fly and where the International Space
Station orbits Earth. The auroras also occur at this layer. [4]
The exosphere, the highest atmospheric layer, is extremely dilute and is where the
atmosphere fuses into outer space. It is composed of very widely dispersed
particles of hydrogen and helium. [5]
The Aurora Lights! Charged particles from space collide with atoms and molecules in the thermosphere, exciting them into higher states of energy. The atoms shed this excess energy by emitting photons of light, which we see as the colorful Aurora Borealis and Aurora Australis.
“NASA – NASA and World Book". Nasa.gov. 5 May 2014
Pg. 2 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
As for the composition of the atmosphere, the earth's atmosphere near the surface is composed
primarily of Nitrogen and Oxygen. It is clear that the main gas is nitrogen. Oxygen - the gas that allows
animals and plants to respire, and fuels to burn - is the next most abundant gas. These two gases are
both elements and account for about 99% of the gases in the atmosphere. Together, the two comprise
about 99% of the gas in the atmosphere. The remaining gases, such as carbon dioxide, water vapour and
noble gases such as argon, are found in much smaller proportions. The exact proportions in which the
atmosphere is composed of gases is as follows: [6]
Nitrogen - 78.084%
Oxygen - 20.95%
Argon - 0.934%
Carbon Dioxide - 0.036%
Neon - 0.0018%
Helium - 0.0005%
Methane - 0.00017%
Hydrogen - 0.00005%
Nitrous Oxide - 0.00003%
Ozone - 0.000004%
Identify the main pollutants found in the lower atmosphere and their sources
The main pollutants in the lower or the tropospheric atmosphere are carbon dioxide, carbon monoxide,
nitrogen oxides, ozone, volatile organic compounds, sulfur oxide, lead and particulates. Below is a table
that best identifies the main sources from which these pollutants enter the troposphere.
Main pollutants Main sources
Carbon monoxide Incomplete combustion in stoves, cars, fires and cigarettes. Road traffic emissions account for up to 90% of carbon monoxide emissions. [8]
Carbon dioxide The burning of fossil fuels such as coal and oil products.
The transportation of goods and people is the second largest source of anthropogenic carbon dioxide emissions.
Nitrogen oxides These gases form when fuel is burned at high temperatures, and come principally
[9]
[7]
Pg. 3 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
from motor vehicle exhaust and stationary sources such as electric utilities and industrial boilers. Examples of the nitrogen oxide family (NOx) include of nitrogen monoxide and nitrogen dioxide.
Naturally occurring bi-products of some organic reactions in soil bacteria
Exhibited when used in agricultural processes as a form of soil fertilizer [10]
Ozone The two major sources of natural ground-level ozone are hydrocarbons, which are released by plants and soil, and small amounts of stratospheric ozone, which occasionally migrate down to the earth's surface. However, emissions from neither of these two sources are in harmful amounts.
Harmful amounts of tropospheric ozone are formed when sunlight, particularly ultraviolet light, reacts with hydrocarbons and nitrogen oxides, which are emitted by automobiles, gasoline vapors, fossil fuel power plants, refineries, and certain other industries.
Volatile organic compounds Volatile organic compounds are released in the exhaust gases of vehicles when hydrocarbon compound fuels are burnt.
Sulfur dioxide Irritating, poisonous gas produced by the
combustion of fuels (such as coal and oil)
that contain sulfur minerals and from
metal extraction processes-in ores that
contain traces of sulfide, which oxidize
with oxygen at high temperatures in
smelters. Naturally generated by volcanoes and
bacterial action Lead Leaded fuels, metal extraction, renovating
old houses containing leaded paints and electrical wire coverings.
Run-offs from lead-acid battery manufacturing and recycling plants
Particulates Incomplete combustion, earthmoving dust pollution, dust storms and some agricultural and industrial practices
Pg. 4 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Describe ozone as a molecule able to act both as an upper atmosphere UV radiation shield
and lower atmosphere pollutant
Senior Principle Research Fellow (Honorary), Peter Gies, at the James Cook University states in his article
‘How does the ozone layer protect Earth from radiation?’ that ozone is found in the highest levels in the
stratosphere, in a region also known as the ozone layer, stationed between 10 to 50 km above the
surface (or 6 to 31 miles). [11] This layer screens the shorter wavelength and highly hazardous ultraviolet
radiation (UVR) from the sun, protecting life on Earth from its potentially harmful effects. However,
even in this layer, ozone concentrations are only 2-8 parts per million
Ozone in the stratosphere is mostly produced from short-wave ultraviolet rays (in the UVC band) but it
can be also produced from x-rays (mainly α, β radiation) reacting with oxygen:
O2 + photon (radiation, wavelength < 240 nm) → 2O
O + O2 → O3 + M (excess energy from reaction)
α + β− + O2 → He + O3 (where the alpha and beta denote nuclear x-ray
radiation) [12]
The charged ozone molecules react with other roaming molecules such as nitrogen or oxygen or
photons may be released (for example, the latter equation could be → He + O3 + M, where M denotes
the energy released as photons), so that the energized ozone
molecules may stabilize momentarily.
Production and constant levels of ozone in the stratosphere are vital
to a sustainable future of our planet earth. This is because, as
aforementioned, due to the ability of ozone molecules to absorb and
convert dangerous ultraviolet radiation into heat (200 to 310 nm
range). [13]
However, these same ozone molecules can also act as harmful
pollutants in the troposphere, due to the properties of ozone. Ozone is a very reactive molecule capable
of oxidizing many substances. Ozone is very poisonous at levels above 20 ppm. It readily oxidizes organic
tissue and thus disrupts normal biochemical reactions in the body. It irritates the eyes and causes
breathing difficulties. Ozone is also toxic to plants, including agricultural crops. It is a much stronger
oxidizing agent than oxygen especially in acidic environments. It readily attacks rubber and plastics. [15]
Describe, using the Lewis dot structure, the formation of a co-ordinate covalent bond and
relate this to the structure of ozone.
A coordinate covalent bond, more commonly known as a dipolar bond, is a description of covalent
bonding between two atoms in which both electrons shared in the bond come from the same atom.
In the formation of a simple covalent bond, each atom supplies one electron to the bond - but that is
not the case with a coordinate covalent bond. A co-ordinate bond (also called a dative covalent bond) is
a covalent bond in which both electrons come from the same atom. Examples of coordinate covalent
bonding are present in many of the industrially used ions, such as ammonium ions, hydronium ions,
carbon monoxide and so forth.
[14]
Note: Species in bold
are energized molecules
Pg. 5 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Ammonium ions
Ammonium ions are an example of a
coordinate covalent bond where the pair of
shared electrons originates from the nitrogen
atom and the hydrogen ions share the pair. The
hydrogen ion having lost its only one electron,
has now a valence of 2+ and a positive charge.
The pair of electrons on nitrogen’s end are thus
shared to form a coordinate covalent bond. The
arrow between the NH3 molecule and the H+
ion, is to show that Nitrogen is the atom donating its pair to the hydrogen ion. [16]
Carbon Monoxide
Certain molecules can also have both general covalent bonding as well as coordinate covalent bonding
present simultaneously, such as in the example of CO. A simple demonstration of the covalent bonding
between carbon and oxygen in CO is given by:
Here, denoted by x, carbon has 4 electrons in its outermost shell
and oxygen’s electrons in the outermost shell, represented by
the dot, are 6. To satisfy its outermost shell, oxygen forms a two
covalent bond pairs with the carbon atom, having an orbit of 8
complete electrons. This is an example of normal covalent
bonding.
However carbon only manages to gain two
extra electrons by forming 2 pairs with
oxygen with two of its existing electrons,
resulting in an orbital of 6 electrons. The
carbon atom’s outermost shell is still not
complete. We can see another pair of
electrons from oxygen on the reactive site.
The extra pair is shared by oxygen, allowing
carbon to have a complete outermost shell.
This bond is a form of a coordinate covalent bonding, since both electrons shared in the bond come
from the same atom.
Hydronium ions
The oxygen atom in water has two non-bonding
pairs of electrons. Hydronium ions form when one
of these non-bonding pairs are donated to a
hydrogen ion. This can be easily understood by the
following Lewis dot structural formula.
G. Thickett, Chemistry 2, Milton, Queensland, 2006. Pg. 298
G. Thickett, Chemistry 2, Milton, Queensland, 2006. Pg. 298
G. Thickett, Chemistry 2, Milton, Queensland, 2006. Pg. 298
[17]
Pg. 6 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Ozone
Ozone, an allotrope of oxygen, is just like carbon monoxide in the sense that it also contains both:
normal covalent bonding as well as coordinate covalent boding. The structure of ozone is bent. The
bond angle is 117° and each bond is equidistant from the central atom in the molecule. Unlike
conventional O2, ozone (O3) has three atoms of oxygen bonded to each other. The Lewis dot electron
structure of ozone best describes its nature physical makeup of the covalent bonds it contains.
In both the structures, we can see that one of the
oxygen atoms (Far left on the left structure and far
right on the right structure), has 3 paired electrons
at rest, yet there is still a bond with the central
atom. The central part of the molecule tends to
have only one pair of electrons at rest, while the
other two pairs are involved in the covalent
bonding.
One pair is shared in normal covalent bonding,
whereas the other pair is donated to the atom with
three electron pairs at rest. Therefore, the ozone molecule consists of conventional covalent bonding as
well as cordinate covalent bonding.
The reason for the two structures shown in the above diagram, is because ozone is a resonance
structure. In simple terms, it means that the cordinate and the conventional covalent bonding can
alternate on either of the two protuding bonds, since we cannot know exactly where the bonds are
present. [18]
Compare the properties of the oxygen allotropes O2 and O3 and account for them on the
basis of molecular structure and bonding.
The properties of the two oxygen allotropes, O2 and O3 differ physically and chemically due to their
different make up.
Properties Oxygen (O2) Ozone (03) Explanation
Molecular Formula O2 O3 -
Appearance Transparent Bluish colored gas -
Odor Odorless Strong. Human noses can identify ozone gases at around 10ppm
-
Melting point -218.79 oC -192.5 oC The melting point of
diatomic oxygen is
lower than that of the
ozone as diatomic
oxygen has less
molecular bonds
requiring less energy
in the melting process to break the bonds.
G. Thickett, Chemistry 2, Milton, Queensland, 2006. Pg. 298
Pg. 7 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Boiling point -182.95 oC -111.9 oC The boiling point of diatomic oxygen is lower than that of the ozone as diatomic oxygen has a lower molecular mass requiring less energy in the boiling process.
Density (0 °C, 101.325 kPa) 1.429 g/L
2.144 g/L (0 °C), gas Ozone is denser because it is a tri-atomic gas, while Oxygen is diatomic. The molecular weight of Ozone is 48 & of Oxygen is 32. 1unit volume of Ozone weighs more than 1 unit volume of Oxygen.
Solubility in water sparingly soluble more soluble than oxygen
O2 is non-polar, therefore it does not form strong intermolecular forces in the polar water. Ozone has a bent structure, which provides some polarity of the molecule (its dipole nature), in its interaction with water. [19]
Chemical stability Relatively stable, much more stable than ozone gas
Much less stable than oxygen gas
Ozone is readily decomposed into oxygen gas molecules. This is mainly because upon reacting with atomic oxygen, ozone decays, since that leads to a more stable form oxygen. O3 + O → 2 O2
Oxidation ability Gentle, less powerful oxidant than ozone
Readily oxidizes organic matter, as well as inorganic matter. Much more powerful than oxygen
To look at metallic oxidization for example, oxygen only forms one metallic oxide as a product, whereas ozone forms a
Pg. 8 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
metallic oxide as well as an oxygen molecule, which further oxidizes the metallic substance. Furthermore the electronegative nature of the dipole in the ozone molecules, readily stimulate the molecule to react with extrinsic matter
Uses in industry Diatomic oxygen has many uses in industry. These include its involvement during oxyacetylene welding and, in the steel industry, in the conversion of iron to steel. In medicine, compressed oxygen bottles are vital for sustaining patients during operations, as well as those suffering from respiratory diseases. In the space industry liquid oxygen is used as an oxidizer for the liquid hydrogen fuel in rockets and the space shuttle [20]
Ozone is an excellent bleaching agent and can be used to bleach wood pulp in the preparation of paper. Because it kills micro-organisms, ozone can be used to disinfect water.
The physical and chemical properties of the two allotropes allow for them these uses in the industrial and commercial world
Diatomic oxygen (O2)
Diatomic oxygen (O2) consists of two oxygen atoms linked by a double
covalent bond. This double bond is very strong and has a high bond
energy (498 kJ/mol). Such bond stability makes it less reactive than
ozone, which has a lower bond energy (445 kJ/mol). [21] Diatomic
oxygen supports combustion and rekindles a glowing splint of wood.
This test is commonly used as a test for oxygen. Hot metals such as
magnesium burn in oxygen to form metallic oxides
2Mg(s) + O2 (g) 2MgO(s)
The molecular structure of O2 plays a major role in its functionality.
The oxygen molecule contains one double covalent bond O=O. The
picture aright shows the Lewis dot structure of O2 and its molecular makeup.
[22]
Pg. 9 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Triatomic Oxygen (O3)
Ozone consists of three oxygen atoms arranged to form a bent molecule with bonds of equal length.
Simple Lewis electron dot diagrams do not adequately explain its structure and properties. Each bond in
ozone can be considered as being intermediate between a single bond and a double bond. Ozone is a
pungent and a poisonous gas. It is very powerful oxidizing agent and as a consequence is able to oxidize
living tissue. In its reactions it is able to split off a reactive oxygen atom (or radical), because its bond
energy is lower than that of diatomic oxygen. This free oxygen atom then rapidly combines with
material that is being oxidized. Ozone very rapidly oxidizes metals to form metal oxides with the release
of oxygen gas.
Zn(s) + O3(g) ZnO(s) +O2(g)
Ozone readily oxidizes sulfides to form sulfates.
CuS(s) + 4O3(g) CuS04(s) + 4O2(g)
The structure of ozone is different to oxygen’s. It has two double covalent bonds as well as a singular
coordinate covalent bond. The coordinate covalent bond is represented by an arrow.
The bonds between the oxygen atoms in an ozone molecule are of equal length (128pm) and strength,
thus the resonant ozone can be represented by: [23] [24]
There are two identical oxygen to oxygen bonds in ozone, which consist of a
single bond and a partial bond. The presence of a partial bond results in lower
stability of the ozone molecule, compared with the diatomic oxygen
molecule.
Compare the properties of the gaseous forms of oxygen and the oxygen free radical
In their ground state oxygen atoms have 8 electrons. The electron configuration of a ground state
oxygen atom is 2, 6. Where, in the second shell the 6 electrons are paired in groups of two, to form
three pairs. When oxygen atoms combine to form oxygen
molecules (O2), a pair of valence electrons from each atom
separate and form a double covalent bond with the
opposing pair of electrons.
When these same, double covalently bonded oxygen
molecules split into separate oxygen atoms, for example,
by the absorption of UV light in the stratosphere, the
atoms of oxygen that are formed are called oxygen free
radicals.
G. Thickett, Chemistry 2, Milton, Queensland, 2006. Pg. 301
Pg. 10 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
These radicals are different from oxygen atoms in their ground state, because
once these atoms split, the third electron pair in the outermost shell has
separated and thus two free roaming electrons cause the oxygen free radical to
be very reactive. The unpaired electrons exist in higher energy (or excited states)
than the ground state.
These oxygen free radicals exist only briefly in the lower layers of the
atmosphere before they react with other radicals or molecules. In the
thermosphere, oxygen free radicals are formed when UV photons cause
photodissociation of oxygen molecules. The great separation of particles in the
thermosphere allows radicals to predominate. Oxygen free radicals are even
more reactive than ozone.
Natural formation of ozone
Stratospheric oxygen absorbs the ultraviolet radiation (UV wavelengths 200-310
nm range) and photodissociation occurs leading to the formation of oxygen free
radicals. [13] These reactive radicals combine with oxygen molecules to form an energized ozone
molecule. As mentioned earlier, the excess kinetic energy of the energized ozone molecules is exhibited
out to other molecules such as nitrogen or oxygen, or photons may be released. The process prevents
energized ozone molecules decomposing.
O2(g) UV radiation 2O. (g)
O2(g) + O.(g) O3(g)
O3(g) + N2(g) O3(g) + N2(g)
Identify and name examples of isomers of haloalkanes which impact upon the ozone
concentration
When alkanes react with halogens (members of group VII of the periodic table) they form new
compounds that are collectively called haloalkanes. Haloalkanes often exist in isomeric forms. The
variable location of the halogen functional groups within the molecule leads to the formation of
isomers. One example of an isomeric haloalkane (a halon isomer C3H5BrCl2) is:
G. Thickett, Chemistry 2, Milton, Queensland, 2006. Pg 304
O Free Rad
Oxy
gen
allo
tro
pe
reac
tivi
ty s
erie
s
Note: Species in bold
are energized molecules
Pg. 11 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Chlorofluorocarbons and halons
Chloroalkanes can be used as solvents. Two types in particular, known as chlorofluorocarbons (CFCs),
which are widely used in aerosols and fridges and halons, which are used in extinguishing fire, are
studied in atmospheric chemistry.
Chlorofluorocarbons are haloalkanes
containing chlorine and fluorine atoms but not
hydrogen atoms, e.g. CCl2F2, CClF3 or CH3Cl.
Halons are bromofluorocarbons and typically
contain at least one ‘bromo’ group as well as
‘fluoro’ functional groups, e.g. CBr3F or
CBrClF2. These small chloroalkanes are gases
and can escape into the atmosphere. Ozone
(O3) is a naturally occurring substance found in
the upper atmosphere. [25]
Chlorinated haloalkanes and other halogenated
hydrocarbons are the reason for the thinning of the ozone layer. Some chlroinated compounds such as
CH3Cl and HCl are naturally occuring, but rarely reach the stratosphere in significant amounts as they are
quickly oxidised in the troposphere. A vast species of chlorinated compounds is synthethic. These
synthetic halogenated hydrocarbons-usually referred to as chlorofluorocarbons (CFCs) and
bromofluorocarbons (halons)-slowly diffuse from the lower atmosphere into the stratosphere. Once in
the stratosphere they break down under UV light (photodissociation) to produce reactive chlorine and
bromine radicals that readily attack and deplete ozone molecules.
The following reactions show the photodissociation reactions incolving CFCs and halons. CFC-11 has an
atmospheric lifetime of 70 years.
a) trichlorofluoromethane (CFC-11) (CFCl3)
CFCl3(g) + UV CFCl2. + Cl.(g)
b) bromotrifluoromethane (Halon 1301) (CF3Br)
CF3Br(g) + UV CF3.(g) + Br.(g)
Bromine radicals, produced by halons, cause more ozone reduction than Cl. radicals. Halon 1301 has an
atmosperic lifeftime of 110 years. It also comprises ten times the ozone depleting potential (ODP) than
that of the CFC-11. CFC-11 is used as a standard haloalkane and is given an ozone depleting potential of
1.0. CFCs generally carry ODP values of between 0.01 to 1.0. Whereas, halons have ODP values of up to
10. Below are examples of some of these haloalkanes, both CFCs and halons, with their ODP values. [26]
Compound IUPAC name ODP
CFC-11 (standard) trichlorofluoromethane 1.0
CFC-113 1,1,2-trichloro-1,2,2-trifluorothane 0.8
CFC-115 Chlroropentalfluoroethane 0.6
Halon 1211 Bromochlorodifluoromethane 3.0
Halon 1301 Bromotrifluoromethane 10.0
CCl4 Terrachloromethane 1.1
CH3Br Bromoethane 0.6
G. Thickett, Chemistry 2, Milton, Queensland, 2006. Pg 304
Pg. 12 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Identify the origins of chlorofluorocarbons (CFCs) and halons in the atmosphere
Chlorofluorocarbons (CFCs) and halons are both examples of haloalkanes. These haloalkane gases are
extremely small and can very easily escape into the atmosphere. The difference between CFCs and
halons being, that CFCs are haloalkanes that contains both fluorine and chlorine atoms, and no
hydrogen atoms. Whereas halons are also referred to as a brominated CFC, a haloalkane that contains
bromine, chlorine and/or fluorine atoms, and no hydrogen atoms.
CFCs
During the late 1800’s and early 1900’s ammonia (NH3), methyl chloride (CH3Cl), and
sulfur dioxide (S02), were used as refrigerant. The problem with these compounds was
that they often led to fatal accidents. Consequently three American corporations-
Frigidaire, General Motors, and Du Pont developed refrigerants to substitute these
noxious compounds. [27]
CFCs were first manufactured in 1928 by Thomas
Midgley, Jr. of General Motors. At the time, their properties were
found to be ‘safer’ than the ammonia and the other refrigerants and
were used in large commercial applications. The non-toxicity of CFC’s
and their safety led to CFCs being the preferred coolant in large air-
conditioning systems. Many American cities revised their public
health codes to designate CFCs as the only species of coolants that
could be used in public buildings.
Over time, CFCs were being used as propellants for bug
sprays, paints, hair conditioners, and other health care
products, specifically after WWII. Their low boiling points
(near room temperature), inertness and ability to phase
change at low pressure made them ideal working fluids.
During the late 1950s and early 1960s air conditioning
became the mainstream in many automobiles, homes,
and office buildings because of CFCs; which accentually
led to more than one million metric tons of CFCs being
produced. Overtime CFC emissions increased and then
came to a halt, as shown by the graph. [30] The reasons
for the plateau will be explained in the following sections.
Halons
Halons can be referred to as CFCs containing bromine. Halons were introduced into the
commercial and the industrial world concurrently with the progression in the study of
CFCs. Their main use was in fire extinguishers for electrical fires or to protect computer
systems. Fortunately they were never used as extensively as CFCs were since their uses
were limited. Halon use has been drastically reduced since studies show that bromine
atoms are much more effective than chlorine atoms in the chain reactions that lead to
ozone depletion.
[28]
Concentrations of CFC-11 overtime [31]
[29]
[32]
Pg. 13 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Write equations to show the reactions involving CFCs and ozone to demonstrate the removal
of ozone from the atmosphere
The cause of ozone depletion
Ozone plays an important role in absorbing ultra-violet radiation from the sun and preventing it from
getting to the earth’s surface. The first piece of evidence that would lead to an explanation of the
thinning of the ozone layer were provided by the measurements of the levels of chlorine oxide radicals
(ClO.) in the stratosphere. Chlorine oxide is formed when atmospheric chlorine compounds, such as
chloromethane (methyl chloride), undergo photodissociation from UV radiation to form methyl radicals
and reactive chlorine radicals. Chloromethane is actually an industrially manufactured chloroalkane.
These chlorine radicals are viable in the stratosphere by the emissions of CFC gases and halon gases via
industrial and commercial use. A small quantity of chloromethane reaches the stratosphere and
undergoes photodissociation. The chlorine radicals rapidly attack ozone molecules and produce ClO.
CH3Cl(g) UV Radiation CH3.(g) + Cl.(g)
Cl.(g) + O3(g) ClO.(g) + O2(g)
The product of chlorine oxide radicals can further react with many other species in the stratosphere to
produce more chlorine radicals. The chlorine oxide radicals may react with oxygen free radicals to
produce dioxygen and additional chlorine radicals.
ClO.(g) + O.(g) O2(g) + Cl.(g)
Below is another alternative chemical reaction pathway. The ClO. Radical forms a dimer (Cl2O2), which
eventually produces, via photodissociation, more Cl. Radicals to catalyze further ozone decomposition.
2ClO.(g) Cl2O2(g)
Cl2O2(g) + UV ClO2.(g) + Cl.(g)
ClO2.(g) Cl.(g) + O2(g)
This process can repeat itself almost indefinitely, meaning that even small quantities of chlorine radicals
can significantly destroy the ozone layer. It is estimated that one chlorine radical can destroy up to tens
of thousands of ozone molecules before it is removed from the stratosphere by other processes. [33]
There are only certain narrow pathways by which chlorine oxide radicals may be deactivated, for
example by reaction with nitrogen dioxide gas. Chlorine nitrate is the product of the reaction.
ClO.(g) + NO2(g) ClONO2(g)
This graph shows the correlation
between ozone and ClO levels.
G. Thickett, Chemistry 2,
Milton, Queensland, 2006.
Pg 308
Pg. 14 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Discuss the problem associated with the use of CFCs and assess the effectiveness of steps
taken to alleviate these problems
Problems with CFCs
As previously discussed CFCs and halons are very unreactive, inert hydrocarbons. They are also insoluble
in water, meaning that they remain in the atmosphere even after it has rained. Furthermore, they don’t
decompose in the lower atmosphere and have the tendency to stay the same. CFCs and Halons migrate
slowly into the stratosphere, diffusing out of the lower atmosphere and causing a major environmental
problem. Once exposed to ultra-violet light wavelengths (U.V.) in the stratosphere, they consume
enough energy to break the bonds and dissociate the CFC molecule. Thus a chlorine atom, i.e. a chlorine
radical is released from the molecule. This is given by the equation
CF2Cl2(g) UV Radiation CF2(g) + 2Cl.
As previously mentioned, the equation process further causes the chlorine radicals (Cl.) to react with
ozone to deplete it.
Cl.(g) + O3(g) ClO.(g) + O2(g)
The Chlorine oxide radical is then regenerated with the reaction with free Oxygen atoms which forms
the Chlorine radical (Cl) and oxygen.
ClO.(g) + O.(g) O2(g) + Cl.(g)
This is a cyclic process which has the potential for a single chlorine atom to destroy up to a hundred
thousand ozone molecules, over its two year life-span.
Ozone, is poisonous to humans if inhaled, but in the stratosphere it is responsible for efficient human
living on our planet. It filters out and absorbs the short dangerous wavelengths of ultraviolet radiation
and converts them into heat up in the stratosphere. The range of wavelengths it can absorb is an
amazingly small range of 200-310nms. And it is the CFCs which have caused the depletion of the ozone
layer. Halons, are even more easily dissociated by UV, and the bromine atoms are far more dangerous at
depleting ozone.
Steps taken to alleviate these problems: International Treaties and their impact
The harmful effects caused by CFCs have received high levels of alertness over the past few decades.
These detrimental effects were first discovered in 1974, by the two chemists, Professor F. Sherwood
Rowland and Dr. Mario Molina, of the University of California. Their findings showed that CFCs could be
a major source of inorganic chlorine in the stratosphere following their photolytic decomposition by UV
radiation. They established that CFCs were the main reason behind the depletion of the ozone in the
stratosphere. [33]
Ever since, growing concern over the depletion of the ozone layer has led to a ban being imposed on the
use of CFCs in aerosol-spray dispensers in the late 1970s by many governments such as the United
States, Cana, and by the Scandinavian countries. The measurements of ozone showed that the depletion
was worsening every year. This point was described by British researcher Joe Farman and his colleagues,
in 1985. It was then realized that the only way to stop ozone depletion would be to terminate the
emissions of CFCs into the atmosphere and thus an international covenant was made.
Pg. 15 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Two years after the study of the British scientist Joe Farman, in 1987, 27 countries across the world
signed a global environmental pact, known as the Montreal Protocol, to stop the use of ozone depleting
substances and products that had a provision to reduce 1986 production levels of these compounds by
50% before the year 2000. [34] This international treaty also imposed restrictions on many of the CFCs
such as CFC-11, -12, -113, -114, -115, and halons (brominated CFCs used in fire extinguishers). The
effectiveness of the Montreal Protocol is best described by the above graph, taken from the UNEP
report “HFCs: A Critical Link in Protecting Climate and the Ozone Layer”. [34]
An amendment to the Montreal agreement, approved in London in 1990, was more aggressive and
called for stopping the production of CFCs by the year 2000. 93 nations agreed to the London
amendment and by 1992 majority of those same countries agreed to bring this target closer to 1996. [35]
The chlorinated solvents, methyl chloroform (CH3CCl3), and carbon tetrachloride (CCl4) were also added
to the London Amendment as banned chemicals. [36]
The science that became the basis for the Montreal Protocol resulted in the 1995 Nobel Prize for
Chemistry.
During the winter of 1992, significant amounts of reactive stratospheric chlorine, in the form of chlorine
oxide radical, (ClO) were observed by instruments onboard the NASA ER-2 aircraft and the UARS (Upper
atmospheric Research Satellite) over regions in North America. [37] The environmental concern for CFCs
follows from their long atmospheric lifetime, for example the lifetime of 55 years of CFC-11 and 140
[34]
Pg. 16 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
years for CFC-12 and CCL2F2, which limits our ability to reduce their abundance in the atmosphere and
associated future ozone loss. [38]
These readings resulted in the formation of the Copenhagen Amendment, further limiting production of
CFCs and was approved later in 1992. By January 1, 1996, the production of these chemicals ended for
the most part. The only exceptions to their production were for production on minute scales within
developing countries and for some excused applications in medicine (i.e., asthma inhalers) and research.
[39] More revisions of the Montreal Protocol included applying economic and trade penalties should a
participant country yield or trade these banned chemicals. As of now a total of 148 countries have
signed to the Montreal Protocol. [39]
By measuring the atmospheric content of
CHC-11 and CFC-12, readings showed that
their growth rates were reducing as result of
both voluntary and authorized reductions on
productions and emissions. The main and the
most crucial effectiveness of these treaties
being that many CFCs and selected
chlorinated solvents have either leveled off or
decreased in concentration by 1994, as
shown by the graph above. It is estimated
that these will only be on the decline
henceforth.
Other steps
Globally, there have been a few more steps taken. Organizations have been funded by governments to
produce alternative chemicals such as hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons
(HFCs). These substitute compounds will be discussed in the following section. Via the Montreal Protocol
and the voluntary need for action against the depletion of stratospheric ozone, advanced countries are
now providing assistance to less developed countries to eventually phase out the use of CFCs.
Identify alternative chemicals used to replace CFCs and evaluate the effectiveness of their use
as a replacement for CFCs
The demand for CFCs has been further reduced by the use of substitutes. Some applications, for
example cleaning solvents for circuit boards, which once used CFCs now use halocarbon-free fluids. The
industry developed two classes of halocarbon substitutes for other uses.
HCFCs (Hydrochlorofluorocarbons)
At first the CFCs were replaced by HCFCs (Hydrochlorofluorocarbons). HCFCs are compounds that
comprise of at least one Hydrogen (H) atom. Examples of these are shown in the table on the next page.
[40] The presence of a C-H bond makes these compounds more reactive in the lower or the tropospheric
atmosphere. This means that they are more easily decomposed, prior to them reaching and diffusing
into the stratosphere, however this is a slow reaction and there is still a chance that the HCFCs can reach
and destroy stratospheric ozone. The environmental impact HCFCs had was still noticeable and under
the Copenhagen amendment, the production of all HCFCs will be finished by the year 2030. The use of
Pg. 17 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
HCFCs was not as effective as at first suggested since they did continue to add on to potential ozone
depletion, but they did initially provide a much better substitute for CFCs.
Name Formula Uses
Chlorodifluoromethane CHClF2 Air conditioning, refrigeration, foams
1-Chloro-1,1-Difluoroethane
CClF2CH3 Aerosols
1,1-Difluoroethane CHF2CH3 Aerosols, refrigeration
HFCs (Hydrofluorocarbons) The other substitute, HFCs (Hydrofluorocarbons), are considered one of the best substitutes for CFCs
while helping reduce stratospheric ozone loss. This is because of their short lifetime and lack of chlorine
and bromine. They are however efficient greenhouse gases and were targeted for emission reductions
in another global treaty, which was aimed at dealing with the effects greenhouse gases, the Kyoto
Protocol. HFCs are now the most commonly used substitutes, consisting of Hydrogen (H), Fluorine (F)
and Carbon (C) atoms. There are two types of HFCs: high-GWP HFCs and low-GWP HFCs, the latter being
safer for the environment, since GWP is a relative index that enables comparison of the climate effect of
the emissions of various greenhouse gases
HFCs are more expensive to synthesize and less efficient in performance than CFCs but they do have the
advantage of a null ozone depleting potential. Their C-H bond victimizes them as subjects to
decomposition in the troposphere. An example of the effectiveness of their use is given by the
[42]
Pg. 18 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
demographic fact that in the United States, HFC-134a (CF3CH2F) has been used in all new domestic
automobile air conditioners since 1993. [41]
The bar graph on the previous page further seconds the effectiveness of the use of HFCs by showing that
they are now being used been
used in: air conditioning,
refrigeration, fire suppression,
solvents, foam blowing agents,
and aerosols.
The graph aright, taken from the
UNEP report “HFCs: A Critical Link
in Protecting Climate and the
Ozone Layer” shows how reliant
industry has become on the two
substitutes overtime. [42] CFCs
have practically dropped to
nothing, HCFCs have been on a
constant rise and lastly HFCs have
rapidly increased.
Sequential Progression of alternative substitutes
Although HFCs have been the latest substitute used to replace CFCs, after HCFCs, scientists aim to make
future technologies which have no influence on the climate because HFCs have a greenhouse effect on
the climate. Currently Low-GWP HFCs are the most viable and favored substitute for the original CFCs
[43]
Pg. 19 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Analyze the information available that indicates changes in atmospheric ozone
concentrations, describe the changes observed and explain how this information was
obtained.
Stratospheric ozone depletion
In 1976, the British Antarctic Survey at Halley Bay noted a 10%
drop in ozone levels in the stratosphere over Antarctica in the
southern spring (August to October). This was unusual as level
had remained constant since measurements had begun in 1957.
These scientist made their measurements using ground based
Dobson UV spectrometers as well as on air samples collected by
high altitude balloons and aircraft.
Initially, they considered that either their instruments were
malfunctioning or that the apparent seasonal losses were due to
natural events such as sunspot activity or volcanic action. They
become very concerned in 1983, however, when they observed
record losses of ozone that spring
Measuring the ozone
By 1985, atmospheric measurements over Antarctica showed
a 50% reduction in ozone concentrations in the stratosphere
over the precious decade. This result correlated with
independent data recorded by the total ozone mapping
spectrometer (TOMS) and a solar backscatter ultra-violet
detector orbiting the Earth in the Nimbus-7 satellite. Since
then other satellites (including some that use infra-red
radiometers) have been used to scan the upper atmosphere to
determine ozone levels. [44]
Another technique used to measure ozone involves the use of
UV lasers. Pulses of different wavelength UV laser light are
fired from several lasers at ground levels into the stratosphere. The degree of absorption of this light at
various levels is measured using UV spectroscopes attached to telescopes. From this data the ozone
concentration can be calculated. The use of many different methods, including new electrochemical and
chemiluminescence techniques, to measure ozone levels has improved the reliability of the collected
data and has demonstrated that ozone loss is a real phenomenon.
Finding the ‘ozone hole’ The thinning of the ozone layer results in what is often called the ‘ozone hole’ in the 1980s the ozone
loss worsened, and the area over which the ozone loss occurred became wider. During 1987, the ozone
hole broke up. And ozone-depleted air spread over large areas including southern Australia and New
Zealand. Over several days in mid-December 1987 the ozone layer had thinned by 12% over Melbourne.
Further significant ozone depletion were recorded (by later launches of TOMS instrument) in the
southern springs of the 1990s. The largest ozone holes measured so far occurred in 2003, 2000 and
Pg. 20 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
1998. The CSIRO report for September 2005, however, has
revealed that the 2005 ozone hole was the fourth largest since
1979. The hole covered an area of 26.4 million square
kilometers. The ozone hole, however was not confined to the
Antarctic. Small decreases in ozone levels were noted above
the Arctic in the winters of 1994 and 1995. By 1996 the
thinning of the ozone layer over the Arctic had reached 40%
[45]
Scientists became alarmed over the decreasing concentrations
of ozone, as this would lead to more UV radiation reaching
ground level. The vast numbers of phytoplankton and
zooplankton in the surface waters of the ocean would be
affected by an increase in UV.
As these organisms are vital components of the ocean food chain, a reduction in ozone levels could
produce a significant decline in marine organisms. Studies of skin cancer and suburban rates in Punta
Arenas, Chile (the southern city in the world), in the 14-year period 1987-2000 have shown that there
has been an increase of 66% in skin cancers in the second half of this time interval compared with the
first half. These results correlate well with the 56% reduction in peak stratospheric ozone recorded over
the same period. [46]
End of Report
Pg. 21 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
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Pg. 23 Bilal Fouzi – Year 12 Chemistry – Australian Islamic College of Sydney
Books used
G. Thickett, Chemistry 2
R. Smith, Conquering Chemistry: HSC course
David E. Alexander and Rhodes W. Fairbridge, The Chapman & Hall Encyclopedia of
Environmental Science
Matsumi, Yutaka; Kawasaki, Masahiro (2003). "Photolysis of Atmospheric Ozone in the
Ultraviolet Region
Special Acknowledgments from author
One of the most influential person in my life, my mother, my teacher, my mentor and the one who
believed in me when all had given up on me, Mrs. James, deserves to be credited for the work of this
report. The work and the areas of extensive knowledge I have uncovered and elaborated on in this
report, is all the fruits of her work. You reap what you sow. And this report is a mere metaphorical
reflection of what perfection she has gained via me, my most respected, most honored teacher. It is as if
I am, the natural process of ozone depletion and she is the Cl radical in my life. Catalyzing me, and
reducing the initiation energy required in my natural processes of producing quality work. I hope I can
truly pay her back, by becoming the state topper for chemistry and show her that yes I actually am a
positive ion.
I would also like to acknowledge, my stage 5-6 science teacher, Mr. Sallahuddin Ahmed, due to whom
my interest in science developed. My strong basic knowledge and curious interest in the fields of
science, particularly in biology and chemistry, is due to influence of Mr. Ahmad’s caliber as a physicist,
but moreover as a fatherly figure, always looking out for me. I can never pay him back in all aspects of
life.
Moreover, I would like to thank God Almighty for enabling me to be able to do this assignment and get
24/24 in this part of my assignment. For all the favors he has bestowed upon me, in particular, giving me
teachers and life guides like Mrs. James and Mr. Ahmad.
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