chemical monitoring and management · bilal fouzi – year 12 chemistry – australian islamic...

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


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]


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


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


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



[17]


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.



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


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]


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.



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


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


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]


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


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.


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]


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


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]


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]


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


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


Bibliography

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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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body that administers it) has given a remuneration notice to Copyright Agency Limited. Full rights over

this work, can be licensed by purchase. Please call author on +61 402 013 959, for further pricing and

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