A Teacher’s Pack providing background knowledge for teachers, lesson plans and resources for use in class
Lisette Voûte
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Introduction
The purpose of this Teachers Pack is to provide teachers with a full resource on Green
Chemistry, covering all the specifications’ requirements for A-Level chemistry. Green
Chemistry is a new and/or enlarged topic in all the new specifications for teaching in
September 2008, following the new QCA GCE Chemistry Criteria. I reasoned that
many teachers themselves may not have had the opportunity to study environmental
chemistry as it is a relatively modern and current topic; and even if it was studied at
university, there may be some gaps in their knowledge, such as recycling and how it is
done; or they may not be up to date on the chemistry and issues as they stand today.
This resource would therefore allow them to fully understand the subject so they can
be very comfortable teaching it by knowing the background material, beyond what it
expressly required from the specifications. This would allow them, for example, to be
able to explain or answer pupils’ Green Chemistry-based questions on issues beyond
the syllabus. Furthermore, this pack would provide teachers with lesson plans and
resources, which would be useful when teaching a new and relatively unfamiliar
topic.
This resource therefore comes in two main sections: Firstly, a text-book allowing
teachers to become up-to-date on Green Chemistry today. This is based on the
required knowledge in all the A-Level Chemistry specifications (i.e. AQA, Edexcel
and OCR A and B), for both AS and A2, however in much more detail than the pupils
are required to know, in order to provide teachers with the confidence and background
knowledge required to teach these new topics well. The second main section is a
collection of lesson plans and resources on Green Chemistry-based topics. These
lesson plans can be used as just that, a plan for an entire lesson; or ideas and activities
can be used from them and attached to your other lessons where the chemistry might
link together with Green Chemistry. The plans also suggest resources to use, such as
animations, video clips, worksheets, links and PowerPoint presentations; all of which
I have attached to this resource.
The topics I have covered are, Atmospheric Chemistry, including the Ozone Hole; Air
Pollution, including smog, catalytic converters and acid rain; Global Warming,
including the notions of carbon neutral, carbon footprint, biofuels and examples of
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
them; and Recycling, including recycling, amongst other methods of disposal of
aluminium, iron, steel and polymers and issues associated with their disposal.
I hope you find this pack useful and practical; and that your pupils’ enjoyment and
understanding of Green Chemistry is increased.
February 2008
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
What is Green Chemistry?
Green Chemistry is based on and ties together a variety of strings of chemistry:
Organic, Inorganic, Physical, Environmental, Biochemistry and Analytical Chemistry.
Green Chemistry and Environmental Chemistry, while often confused are two
separate fields. Green Chemistry encourages environmentally conscious behaviour,
such as reducing and preventing pollution and the destruction of the planet. On the
other hand, Environmental Chemistry is simply the study of chemistry occurring in
the environment. i
The following page lists the Twelve Principles of Green Chemistry, reproduced from
the Royal Society of Chemistry.
i
Green Chemistry, Wikipedia, site accessed March 2008
http://en.wikipedia.org/wiki/Green_chemistry
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The twelve principles of green chemistry
•
It is better to prevent waste than to treat or clean up waste after it is formed.
•
Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
•
Wherever practicable, synthetic methodologies should be designed to use
and generate substances that possess little or no toxicity to human health and
the environment.
•
Chemical products should be designed to preserve efficacy of function while
reducing toxicity.
•
The use of auxiliary substances (solvents, separation agents, etc.) should be
made unnecessary whenever possible and innocuous when used.
•
Energy requirements should be recognized for their environmental and
economic impacts and should be minimized. Synthetic methods should be
conducted at ambient temperature and pressure.
•
A raw material or feedstock should be renewable rather than depleting
whenever technically and economically practicable.
•
Unnecessary derivatization (blocking group, protection / deprotection,
temporary modification of physical / chemical processes) should be avoided
whenever possible.
•
Catalytic reagents (as selective as possible) are superior to stoichiometric
reagents.
•
Chemical products should be designed so that at the end of their function
they do not persist in the environment, and break down into innocuous
degradation products.
•
Analytical methodologies need to be further developed to allow for real-time,
in-process monitoring and control prior to the formation of hazardous
substances.
•
Substances and the form of a substance used in a chemical process should
be chosen so as to minimize the potential for chemical accidents, including
releases, explosions, and fires.
These principles have been reprinted with permission from Paul T. Anstas and John
C. Warner Green Chemistry: Theory and Practice, New York: Oxford University
Press, 1998
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Table of Contents
1. Atmospheric Chemistry
1
1.1 Planetary Atmospheres
1
1.2 The Earth’s Atmosphere
2
1.3 Explanation for the Temperature Structure of the Atmosphere
3
1.4 Natural Catalytic Cycles: Problem with the Chapman mechanism
6
1.5 The Ozone Hole
8
1.5.1 Ozone Depletion
8
1.5.2 Why the Depletion is Dangerous
10
1.5.3 Explanation for the depletion
11
1.5.3.1 Polar Stratospheric Clouds
11
1.5.3.2 CFCs converting to Active Forms of Chlorine and Bromine
13
1.5.3.3 The Return of Sunlight: Ozone Destruction
13
1.5.3.4 Summary of Ozone Destruction
14
1.5.4 Current and Future Ozone Levels
15
1.5.5 CFC Substitutes
15
2. Air Pollution
19
2.1 Emitted Pollutants
20
2.2 Removal Processes of Compounds
21
2.3 Smog Formation
22
2.4 UK Emissions Today
25
2.5 Catalytic Converters
26
2.6 Acid Rain
27
3. Global Warming
30
3.1 Greenhouse Effect
30
3.2 Climate Change
30
3.3 Global Warming
31
3.3.1 Greenhouse Gases and How They Work
31
3.3.2 Evidence for Global Warming
31
3.3.3 Global Warming Potential
35
3.3.4 Carbon Neutral, Biofuels and Carbon Footprint
35
3.3.5 Controlling Global Warming and the Kyoto Protocol
37
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GCE Green Chemistry Teacher’s Pack
4. Recycling
40
4.1 Why Recycle
40
4.1.1 Household Waste
42
4.1.2 Aluminium and Steel
43
4.1.3 Plastics and Polymers
48
5. Lesson Plans and Resources
55
Lesson Plan: AS Module 1 – Combustion of Alkanes: Air Pollution
56
Lesson Plan: AS Module 2 – Extraction of Metals, Acid Rain and Recycling
58
Lesson Plan: AS Module 2 – Ozone Destruction
59
Lesson Plan: AS Module 2 – Global Warming
60
Lesson Plan: A2 Module 4 – Disposal and Recycling of Polymers
61
Additional Green Chemistry Resources
62
References
63
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GCE Green Chemistry Teacher’s Pack
1. Atmospheric Chemistry
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
1. Atmospheric Chemistry
1.1. Planetary Atmospheres
The Earth’s atmosphere is the only planet within the solar system which contains such
a large percentage of oxygen; it is an oxidising atmosphere; as can be seen in the table
below.
Table 1: Major atmospheric constituents of the Sun and the Planets within the Solar System and
their Surface Temperature 1
Planet/Star
Most Abundant
2nd Most
3rd Most Abundant Surface
Gas
Abundant Gas
Gas
Temperature (K)
Sun
Venus
Earth
Mars
Jupiter
Saturn
Uranus
Neptune
Titan
Planet
/Star
Mass
(kg)
•
H2
CO2
N2
CO2
H2
H2
H2
H2
N2
89%
96.5%
78.1%
95.3%
90%
96%
82%
85%
82%
He
N2
O2
N2
He
He
He
He
Ar
11%
3.5%
20.9%
2.7%
10%
4%
15%
15%
12%
H2O
SO2
Ar
Ar
CH4
CH4
CH4
CH4
CH4
0.1%
0.015%
0.93%
1.6%
0.24%
0.2%
2.3%
1-2%
3%
-732
288
223
170
130
59.4
59.3
95
Table 2: Mass of the Sun and the planets within the Solar System 2
Sun
Mercury Venus
Earth Mars Jupiter Saturn
Uranus
1.99
x1030
3.30x1023
4.87x10
24
5.97x1
024
6.42
x1023
1.90
x1027
5.68
x1026
8.68
x1025
Neptune
1.02 x1026
Mercury has a relatively low mass (see Table 2 above) hence a smaller
gravitational force, and therefore has almost no atmosphere. Its thin
atmosphere is comprised of 98% He and 2% H2.1
•
Venus has largely CO2, which causes a “runaway greenhouse effect” and
hence it’s high surface temperature.1
•
Mars’ atmosphere also consists chiefly of CO2 but as it is of a lower mass than
Venus, its atmosphere is thinner as it has a weaker gravitational force. Thus,
there is not a very strong greenhouse effect.1
•
The Outer Planets (Jupiter, Saturn, Uranus and Neptune) have a much lower
surface temperature due to their distance from the Sun. Their atmospheres are
predominately c. 90% H2 and c. 10% He, and are reducing in nature as they do
not contain oxygen. There are low levels of a range of hydrocarbons in the
atmospheres, most likely to be caused by the photochemistry of CH4.1
•
Titan is a satellite of Saturn and the only satellite to posses a massive
atmosphere; which here is of N2 and some CH4. Titan’s atmosphere contains
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
photochemical smog: most likely to be due to the oxidation of hydrocarbons.
These aerosols cause the smog to appear as coloured clouds.1
1.2. The Earth’s Atmosphere
The moon has c. 1/6 of the gravitational force than the Earth has, so it has virtually no
atmosphere. If the Earth’s atmosphere did not attenuate incoming solar radiation, the
temperature variation of the atmosphere would look like that of the moon’s 3 :
Figure 1: The temperature structure of the moon’s atmosphere and that of the Earth’s, if the
Earth’s atmosphere did not attenuate incoming solar radiation3
The light from the sun heat’s up the Moon’s surface, which radiates heat upwards, so
the surface heats the atmospheric layers directly above it. Therefore, there is a high
temperature at the surface, which falls away rapidly as the distance from the surface
increases, as heat transfer is less effective.1 However, the temperature structure of the
Earth has an S shape:
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Figure 2: Temperature variation with altitude of Earth's atmosphere 4
1.3. Explanation of the Temperature Structure of the Atmosphere
Troposphere – decrease in temperature
From the surface to the Tropopause, the temperature decreases, this is due to the same
reason for the temperature structure of the moon (the Sun heat’s the Earth’s surface,
which re-radiates heat back up, heating the layers above it, with decreasing intensity
as the altitude increases).
Stratosphere – increase in temperature
At 10-15 km the temperature begins to increase throughout the stratosphere.1 This can
be explained by the absorption of solar radiation:
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Figure 3: The solar flux at various altitudes in the Earth's atmosphere 5
As can be seen from Figure 3, the dangerous high energy radiation (wavelengths less
than 200 nm, i.e. the lower UVC region) is removed at the top of the atmosphere. This
is done through photochemistry with O2, O2+, N+, N2+, O, O+, and NO+.1 The (also
harmful – and it is vital for life that it is removed) UVB region, between ~200-300
nm, begins to be absorbed at about 50km, and is removed by the time the solar
radiation impinges on the Earth’s surface1. This is due to O2 and especially O3
(ozone) being the species that absorb UVB radiation1, acting as a ‘UV filter’.
Therefore it is clear that ozone is vital to protect humans and the ecosystem from
harmful UVB radiation.
Ozone is generated via the Chapman Mechanism (discovered in the 1930s by Sidney
Chapman)3:
O 2 + hν → O• + O•
O 2 + O• + M → O3 + M
O3 + hν → O 2 + O•
- ΔH
O3 + O• → 2O 2
Where M is a non-reactive body which can absorb excess energy
The most interesting step is the 3rd one. It is extremely exothermic (90 kJ mol-1)1 so
ozone is liberating a lot of heat in photolysis, causing the temperature to increase in
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GCE Green Chemistry Teacher’s Pack
the stratosphere. Ozone’s generation reaches a maximum in the Stratosphere because
it is a balance between number of photons and the concentration of O2 molecules1:
•
At higher altitudes: There is a high number of photons (fewer have been absorbed
by the atmosphere), but the atmosphere is thinner, so the pressure low, therefore
there is a low concentration of O2 molecules – too low to create high enough
levels of O3.
•
At lower altitudes: Despite the fact there is a higher pressure, hence higher
concentration of O2, the number of photons is too low as the layers of atmosphere
above it have attenuated the incoming radiation – this slows the rate of the first
step in the Chapman mechanism.
In the stratosphere, there is warm air sitting on top of cold air, so it is hard for the cold
air to move through it: the air remains stuck in these distinct layers, which is termed
zonally symmetric1.
Mesosphere – decrease in temperature
Again, the air begins to cool here as the altitude increases, as the pressure decreases
and hence the concentration of M and O2 molecules decrease, so there is very little to
kick-start1 the Chapman cycle, and generation of ozone is extremely low.
Thermosphere – increase in temperature
This begins at about 90 km, at this point the atmosphere is so thin and collisions
between particles are extremely infrequent. This means that most of the particles don’t
get the chance to equilibrate once they have absorbed the high energy incoming solar
radiation, and therefore their translational temperatures become very high1.
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
1.4. Natural Catalytic Cycles: Problem with the Chapman mechanism
The Chapman mechanism predicted the right ozone generation mechanism but overpredicted the production of O3 by a factor of about 5.1 This is because the last step in
the cycle:
O 3 + O • → 2O 2
is slow and there exist catalysts, X, which are species in parts of the atmosphere
which participate in the following cycle to speed up the destruction of ozone. This is
called natural catalytic cycles1:
X + O 3 → XO + O 2
XO + O → X + O 2
Net : O + O 3 → O 2 + O 2
Species of X are:
3
NO 30-40km
e.g.
NO + O3 → NO 2 + O 2
Cl and Br maximum at 45km
NO 2 + O → NO + O 2
HO above 45 km
Net : O + O3 → O 2 + O 2
H above 60 km
Sources of catalysts:
The catalysts are formed via the reactions outlined below. They involve a natural
source gas1 reacting with another molecule or undergoing photolysis (i.e. the
compound is broken down by sunlight – photons).
Despite the fact that these species catalyse the destruction of ozone, these reactions
don’t go on indefinitely, which would destroy the ozone layer. Fortunately,
termination reactions occur and this produces stable (inactive) reservoir
compounds1 from the active radicals. These reservoir compounds are often the source
gases themselves1.
•
Source of NOX:
Firstly, ozone undergoes photolysis to form oxygen and an excited oxygen
radical1:
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GCE Green Chemistry Teacher’s Pack
O3 + hν → O(1 D) + O 2 (1 Δ g )
The source gas, N2O is produced by soil, which goes on to react with the oxygen
radical to form NO1:
O(1 D) + N 2O → 2NO
Terminatio n Step:
NO 2 + HO + M → HNO 3 + M
•
Source of Cl: 1
CH 3Cl + hν → CH 3 + Cl
Termination Step:
Cl + CH 4 → HCl + CH 3
Fortunately, the amount of CH3Cl emitted into the atmosphere is very low.
However if this amount increases, it would present a problem as this is a very
efficient process (see Table 3).
•
Source of HOX: 1
O(1 D) + H 2O → OH + OH
O(1 D) + CH 4 → OH +CH 3
Termination Step:
HO + HO 2 → H 2O + O 2
HO 2 + HO 2 → H 2O 2 + O 2
Table 3: Rates of catalytic cycles
X
K220 / cm3 molecule-1 s-1
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kb/ka
H
1.7 x10
25000
HO
2.2 x10-14
32
NO
3.5 x10-15
5
Cl
8.7 x10-12
12794
O
6.8 x10-16
1
NB: kb/ka
O + O3 → O 2 + O 2
ka
X + O3 → XO + O 2 kb
Rate loss of ozone:
d[O3]/dt=- ka[O][O3]-kb[X][O3]
The rates are equal when:
ka[O][O3]/kb[X][O3]=1
∴[O]/[X] = kb / ka
Natural Chapman cycle
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
It can be seen from the above table that one only needs to add c.1 millionth of the
amount of H into the system compared to O for it to have the same effect as the
natural Chapman cycle.
1.5. The Ozone Hole
1.5.1. Ozone Depletion
In the 1970s, the British Antarctic Survey recorded an enormous decrease in ozone in
the Stratosphere over the Antarctic 6 . At first they believed their equipment to be
faulty as they were astonished by this finding 7 .This continued to decrease year on
year 8 , as shown by Figure 4.
Figure 4: Average Ozone Depletion for
October over the Antarctic6
Figure 5: Average Area of the Ozone Hole
1980-20068
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GCE Green Chemistry Teacher’s Pack
Figure 6: Graph to show the changes in Ozone Partial Pressure over the South Pole from 196720013
As can be seen from Figure 6, since about the mid-1970’s there has been a remarkable
decrease in ozone partial pressures. There has also been an observed diminution over
mid-latitudes and in the Arctic3. Every spring in the Antarctic (i.e. October), the
average ozone levels drop and ozone depletion reaches a maximum at the end of
November. The ozone hole develops seasonally over 6-8 weeks in the stratosphere,
completely destroying ozone in some places. This seasonal change can be seen from
the figure overleaf:
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Figure 7: Graph showing the variation in Ozone Partial Pressure at various altitudes in the
summer (October) and winter (July) of 2001 over the Antarctic 9
As can be seen from the July bulge around 15km, this is where the ozone is usually at
its maximum in the stratosphere. In the winter, the ozone level is almost 0 at this
altitude, showing severe depletion.
1.5.2. Why the depletion is dangerous
Low levels or lack of ozone in the stratosphere causes detrimental impacts of humans,
the ecosystem, and the economy at large. This is primarily due to the fact that reduced
levels of ozone means that less of the dangerous, high energy UV-B radiation will be
absorbed (see Figure 3), and so more will impinge upon the earth’s surface.
Approximately, a 1% decrease of ozone leads to a 2% increase of UV radiation
reaching the earth’s surface3.
•
Humans: Increased levels of dangerous high-energy UVB radiation17 impinging
upon the earth’s surface can cause “melanoma and non-melanoma skin cancer”17,
eye disorders and cataracts and suppression of the immune system in people of all
races – possibly leading to an increase in diseases and infections. According to
Tolba et al:
“The percentage increases [of skin cancer] will not be one-to-one: a sustained
ten per cent reduction in ozone would result in a 26 per cent increase in non-
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GCE Green Chemistry Teacher’s Pack
melanoma skin cancer. All other things remaining constant, this would mean
an increase in excess of 300,000 cases a year, world-wide.”17
•
Plants: Roughly half of the world’s plants are sensitive to UV-B light, and their
leaves shrink and the plant grows less when there is an increase in UV-B light
impinging on the earth’s surface. Economically, this is also problematic as it can
cause reduce food yields and plants also can change their chemical composition
with increased UV-B exposure, which can affect their quality and nutrient levels.17
•
Aquatic ecosystems: Phytoplankton experience a similar detrimental impact of
excess UV-B radiation that terrestrial plants do. This could affect species further
up the food chain and have a detrimental impact on the productivity of fisheries17,
amongst others. In addition, there could be nitrogen deficiency in rice paddies as:
“Increased exposure to UV-B radiation could lead to decreased nitrogen
assimilation by prokaryotic micro-organisms” 17
•
Air quality: Increased levels of surface UV radiation can change the levels of
reactive compounds in the troposphere, such as acids, hydrogen peroxide and
ozone where levels of NOx are high.17
•
Materials damage: Photo-oxidation3 can occur to many materials (wood, plastics
and rubber17), which is when the materials become oxidized through the action of
UV light. Thus, increased UV light can cause increased damage to these materials.
1.5.3. Explanation for the depletion
1.5.3.1. Polar Stratospheric Clouds
In the Polar Regions in the winter, the temperatures drop so low that a polar vortex
forms. This is when sunlight does not shine upon the region, and it is so dark and cold
that air descends and creates a strong downwards vortex motion1, or circumpolar
winds in the mid to low stratosphere6. This isolates the air within the vortex from the
rest of the globe, and no material can get out and material can only enter at 40 km:
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L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Figure 8: Depiction of the Polar Vortex during the winter over the South Pole 10
In these conditions: no sunlight, and very cold temperatures – below -800C – polar
stratospheric clouds (PSC) can form, these clouds remain there as the vortex isolates
the air so it remains cold. These contain high levels of nitric acid (HNO3) and icewater. 11
Figure 9: Depiction of how CFCs enter the atmosphere6
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GCE Green Chemistry Teacher’s Pack
1.5.3.2. CFCs converting to Active Forms of Chlorine and Bromine
CFCs were first created in 1928 as a non-toxic, non-flammable refrigerant, and then
many other uses for them were discovered7, due to their low reactivity and volatility.
As can be seen from the figure above, CFCs are emitted into the atmosphere by
factories and homes (such as through the use of CFCs in aerosols, refrigerants, in airconditioning and as solvents). The UV light in the upper atmosphere can easily break
the C-Cl bonds in CFCs and converts the compounds into the main reservoir species
of chlorine HCl and ClONO2, as they have a long lifetime, and move down into the
polar vortex. The PSCs provide a surface on which heterogeneous reactions can occur
to convert these two species, and their bromine equivalents, into active forms of
chlorine11:
HCl + ClONO2 → HNO3 + Cl 2
ClONO2 + H 2 O → HNO3 + HOCl
HCl + HOCl →H 2 O + Cl 2
N 2O5 + HCl → HNO3 + ClONO
N 2O5 + H 2 O → 2HNO3
These heterogeneous reactions allow for the reservoir compounds of catalysts for
ozone destruction to rapidly convert chlorine and bromine to their active forms.
1.5.3.3. The Return of Sunlight: Ozone Destruction
Remembering this occurs in the winter, as the cold temperatures that occur allow the
formation of the PSCs; when the sunlight returns in the spring (October for the South
Pole), the molecular chlorine readily undergoes photolysis:
Cl 2 + hν → Cl + Cl
This could now go on to catalyse the destruction of ozone through the cycle:
Cl + O3 → ClO + O 2
However there are not enough O atoms in the vortex to do the final step in the cycle :
ClO + O → Cl + O 2 - it is too slow at this altitude (c. 20km)
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GCE Green Chemistry Teacher’s Pack
Instead, the cold temperatures encourage the formation of dimers of ClO, which
drives the following cycle:
ClO + ClO + M → ClOOCl + M
ClOOCl + hν → Cl + Cl + O 2
2(Cl + O3 ) → 2(ClO + O 2 )
Net : 2O3 → 2O 2 - Ozone Destruction
This cycle is thought to be the predominant cycle for ozone destruction, accounting
for 70% of destruction in the South Pole.
1.5.3.4. Summary of Ozone Destruction
The greater the amount of CFCs released into the atmosphere, the greater the amount
of chlorine available as CFCs break with a high energy source:
i.e. light at 200 nm: CFCl3 + hv Æ CFCl2. + Cl.
This free chlorine then goes on to catalytically destroy ozone. The following figure
depicts how ClO is rapidly created from CFCs in the polar vortex in the winter (i.e. in
the absence of sunlight), and how this destroys ozone in the presence of sunlight (i.e.
in the spring), and ClO levels continue to increase.
Figure 10: Graphs to show a comparison in the variations of ClO and Ozone in the Antarctic
vortex in the winter (28/08/87) and the spring (16/09/87) of 1987 12
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GCE Green Chemistry Teacher’s Pack
These same factors also cause destruction over the arctic, although because of warmer
temperatures, the loss isn’t as great. In addition, the PSCs don’t occur as strongly in
the northern hemisphere because the land-ocean distribution is different (the mountain
ranges “stir the atmosphere up”11) and more favourable for their formation in the
southern hemisphere, and in the Arctic the polar vortex disperses earlier in the
spring11.
1.5.4. Current and Future Ozone Levels
Currently, the ozone hole is one and a half times the size of the United Sates (see
Figure 11), and is still getting larger. However, levels of CFCs are decreasing. This is
due to legislations controlling the use of CFCs: the Vienna Convention17 – which did
not prevent the use of bromofluoroalkanes 13 , and therefore – the Montreal Protocol 14
that were implemented in 198517 and 198914 respectively. These legislations were
supported by chemists. The Montreal Protocol aimed to reduce stratospheric halon
levels to the levels they were at before the ozone hole by 206011.
The Montreal Protocol was written so that schedules for the phasing out of
halofluorocarbons could be revised depending on the current scientific and
technological advances 15 . Thus, most recently, in September 2007 a Montreal Summit
was held whereby c.200 countries (including the US and China which had previously
been opposed to the protocol) signed a treaty to accelerate the complete ban of the use
of hydrocarbons by 2020, and developing countries were given until 2030. China
currently has CFC levels equivalent to those that were present in the 60s and 70s in
the UK3.
With the use of halofluorocarbons being phased out, all CFCs are currently banned
except for medicinal use only13. In the US, the use of CFC, HCFC or HFC gases
requires the technician to pass licensing examinations set by the Environment
Protection Agency.13
1.5.5. CFC Substitutes
Scientists have developed and are currently developing alternatives to CFCs, to meet
the legislation’s requirement. CFCs have a variety of uses, primarily as cleaning
agents, fire extinguishing agents, foam, and refrigerants15.
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GCE Green Chemistry Teacher’s Pack
The CFC substitutes developed include:
•
HFC-134a, a chlorine free compound used as a refrigerant with an Ozone
Depletion Potential (ODP – see overleaf) of zero 16 .
•
PhostrEx, the fire suppression agent used in light jets, was developed to be
free of CFCs and is now being sold to other airplane manufacturers13.
•
HCFC: the H atom increases the reactivity of the compound, so less is
required for its use. In addition, 95% of the compound is destroyed the
troposphere and never reaches the stratosphere11, for example CF3CH2F.
•
Other chlorine-free compounds have been developed as well, and their use
is in rapid growth, especially fluorinated and partially fluorinated
hydrocarbons15, such as CF2C12. This requires replacing the chlorine with
fluorine. These compounds do not destroy the ozone layer (doesn’t react
with O3) but unfortunately have a high Global Warming Potential (GWP)
and so contribute to climate change.
Therefore, compounds that are developed today not only have to be chlorine-free but
also have to abide by the Kyoto protocol by having a low GWP15.
Using models, we are able to estimate this future decrease of atmospheric chlorine
level (see Figure 12) and a complete drop by about 20701. The ozone hole itself is
hoped to level off by 2019 and eventually start to decrease by 20503.
Figure 11: The ozone hole with the area of the United States superimposed3
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Figure 12: Past and predicted levels of atmospheric chlorine3
Although this legislation appears to be taking effect, the lifetime of CFCs is very long
and particularly CF3Br is very potent at destroying ozone (see Table 4: bromine has
10 times the ODP of chlorine). In addition, unlike Cl, Br won’t react with methane1:
Cl + CH 4 → HCl + CH 3
Br + CH 4 → no reaction
Therefore it is unable to form a reservoir compound and difficult to get rid of bromine
once it has entered the atmosphere1. Therefore it could take quite some time before
the anthropogenic sources (i.e. derived from human activity) of Cl and Br are
removed from the atmosphere.
Table 4: Halocarbon abundances, lifetimes and Ozone Depletion Potentials
Abundance 17 Lifetime1
Ozone Depletion Potential1
(pptv)
(years)
(=Ozone destroyed by unit mass halocarbon/Ozone
destroyed by unit mass of CFCl3)
CFCl3
280
55
1.0
CF2Cl2
484
100
0.8
CF3Br
2
65
10.0
CH3CCl3
158
50
0.1
Halocarbon
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Finally, it can be seen that global warming did not cause the ozone hole; however
there are links between the two processes11:
•
CFCs are greenhouse gases
•
The stratosphere is actually cooling due to global warming, so more PSCs
can be formed, increasing the amount of reactions for ozone depletion.
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2. Air Pollution
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
2. Air Pollution
Figure 13: Diagram to show the range of chemicals emitted by natural and anthropogenic
activity, as well as depositions, transport and photochemistry 18
Stratosphere
Troposphere
hν
Flow of ozone from
Wet deposition stratosphere
VOCs
- Natural
emissions
NOx &
VOCs
NOx &
VOCs
DMS &
CH3X
CH4
Dry
Deposition
Industrial Activity
Termites, cows,
domestic emissions &
domestic ruminants
The industrial revolution in the 17th century instigated the development of urban
conurbations and also caused a rapid increase in anthropogenic emissions causing air
pollution 19 . This lead to heavy, stagnant combustion smog19 (so named as it is a
portmanteau of smoke and fog 20 ) over cities such as London and caused a variety of
serious heath problems including pulmonary disease and heart failure19. The smog
contained a mixture of Primary Pollutants – emitted directly from the combustion
source, especially soot particles and sulphur dioxide (SO2). Reductions in these
emissions were made leading to a reduction in combustion smog occurances19.
In the last century, increasing emissions of oxides of nitrogen and sulphur due to
industrial and domestic combustion have lead to acid rain19.
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Since the post-war era of the 1950s, both the population and use of motorized vehicles
has rapidly increased. These emissions caused a hazy photochemical smog19. This also
causes adverse health problems, such as eye irritations, sore throats, asthma and
respiratory diseases19. It was found that the component of photochemical smog (and
cause of these ailments) were Secondary Pollutants, made from Primary Pollutants.
The main components are ozone and PAN (peroxyacetylnitrate - CH3C(O)O2NO2)
and are formed from the action of primary pollutants (from car exhausts etc) such as
reactive VOCs and nitrogen oxides, with sunshine19. Every major city in the world
now experiences photochemical smog19. From the figure above, it can be seen that the
troposphere today has a complex array of emissions and processes.
2.1. Emitted Pollutants1
•
VOCs are volatile organic compounds; these contain a large variety of different
organic compounds both emitted from anthropogenic and natural sources. They
absorb IR radiation, as does methane (CH4), and so contribute to global warming
and they are therefore greenhouse gases.
o VOCs from anthropogenic sources are mainly from the burning of fuels (such
as in industry, for energy supplies and through the usage of cars). Alkanes are
used as fuels and their combustion can be complete or incomplete, so some
light, unburned hydrocarbons are emitted during combustion, such as through
car exhausts. Incomplete combustion causes emission of CO and complete
combustion emits CO2. Additionally, the light hydrocarbons in fuel evaporate
when refuelling or storing, and others during combustion. A source of light
alkanes is due natural gas leakage and a source of alkenes is biomass burning.
o Natural sources of VOCs are from vegetation (plants, trees, fungi and algae),
which they emit as pheromones (to ward off predators, to attract insects), to
regulate their temperature or even as antifreeze. The VOCs include
hydrocarbons, CO2 from respiration and CO.
•
NOx includes NO and NO2, this is formed due to high temperatures in combustion
(especially in the combustion of fossil fuels) where nitrogen and oxygen react: N2
+ O2 + heat Æ NOx. In addition, there are some natural sources of nitrogen, such
as ammonia from fertiliser and manure23. This causes the formation of smog and
acid rain.
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GCE Green Chemistry Teacher’s Pack
SO2 and particulates are also released through anthropogenic processes, as
combustion of hydrocarbons releases soot and ash and combustion of
hydrocarbons containing sulphur releases sulphur dioxide. These are emitted
through car exhausts, and as pollution from industry. In addition sulphur can be
emitted from natural sources, such as volcanoes, the action of bacteria in soils and
lightning23. These can both contribute to smog and SO2 contributes to acid rain.
2.2. Removal Processes of Compounds1
•
Wet deposition involves the incorporation of species into aqueous media such as
mist, rain, sea, and snow. This therefore has to involve soluble species such as the
following inorganic compounds: HNO3, H2SO4, HCl, HONO and SO2. VOCs
usually are hydrophobic and therefore insoluble and rarely are lost through wet
deposition.
•
Dry deposition is removal of species through their adsorption onto air mass.
When air mass comes into contact with water, earth or vegetation the emitted
species can either physisorp (through Van der Waals forces) or chemisorp
(through a chemical reaction). The rate of this determined by the flux through the
atmospheric boundary layer (1 km form the Earth’s surface3) the rate of
adsorptions. Generally the species are polar and therefore inorganic such as
HNO3, NO2, HCl, HONO, O3 and SO2. Only very polar VOCs are lost through
this process.
•
Chemical removal is the removal of VOCs through oxidation. This is mainly
done by the OH radical, but also ozone, NO3 radicals and even direct photolysis.
This is why the troposphere is said to be oxidising. Through oxidation, it reacts
with the species and eventually H2O and CO2 get out – water is then rained out (a
subcategory of wet deposition) and CO2 is a greenhouse gas.
o The OH radical is created via the mechanism:
O3 + hν → O(1 D)• + O 2
O(1 D)• + M → O(3 P) • + M
hν at λ < ~ 330nm
where M is an O 2 or N 2 molecule
O(1 D)• + H 2O → OH • + OH •
This occurs in the troposphere: a small amount of ozone is found in the
troposphere and therefore only a small amount of excited atomic oxygen
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1
(O( D)) is created by photolysis. However, as the troposphere sits below the
stratosphere, the air in it is at a higher pressure. This leads to a balance
between concentrations of water vapour and M. At these high pressures the
excited O(1D) is rapidly quenched by M, however the troposphere also has 1050,000 times higher water vapour concentrations than in the stratosphere so a
small amount of O(1D) reacts with water to form OH radicals.
o The OH radical reacts with VOCs (RH) via the reaction:
OH • + RH → R + H 2O
The rates of reaction (and hence lifetime of RH) increases with increasing
number of H’s as there is a greater chance of collision with OH • . However,
with a polar atom in the molecule, such as Cl in CH3Cl, the reaction is a lot
faster as the Cl is electronegative, inducing a dipole in the molecule and
leaving the H’s a lot more positive so they are more reactive. With alkenes, an
addition reaction occurs:
OH • + CH 2 = CH 2 + M → CH 2 (OH) − CH 2
2.3. Smog Formation1
The main ingredient of smog is O31 which is caused by the action of sunlight on NOx
and VOCs. As discussed earlier, VOCs can react very quickly with the OH radical
and in heavily polluted areas, VOCs are often very prevalent – especially aromatic
compounds, alkenes and aldehydes as these react very quickly to form smog1.
Nitrogen oxides (NOx) create ozone in the troposphere. This is unwanted and harmful
to plants and humans’ respiratory systems, as well as being a major component of
smog. This occurs due to the following mechanisms1:
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•
GCE Green Chemistry Teacher’s Pack
Without NOx:
CO + OH • → H • + CO 2
•
H • + O 2 + M → HO 2 + M
•
HO 2 + O 3 → OH • + 2O 2
Net : CO + O 3 → CO 2 + O 2
∴ O 3 is destroyed
•
With NOx:
CO + OH • → H • + CO 2
•
H • + O 2 + M → HO 2 +M
HO 2 + NO• → NO 2 + OH •
NO 2 + hν → NO • + O(3 P) • + O 2
O(3 P)• + O 2 + M → O3 + M
Net : CO + 2O 2 → CO 2 + O3
∴ O3 is produced
Figure 143 overleaf shows the complete cycle of how ozone is formed and destroyed
via the action of NOx and sunlight. In urban areas the concentration of NOx and
VOCs are always high, however one doesn’t see smog on a daily basis as intense
sunlight is required to create the OH radicals and photolyse NO2 (therefore producing
O3), so clear, still, hot days are ideal for smog formation1.
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Figure 14: Diagram to depict the cycle of formation and destruction of ozone, determined by
NOx
Image courtesy of: T. G. Harrison and D. E. Shallcross, Teacher Update 1 – Atmospheric Chemistry, University of Bristol,
Bristol, 21st June 2006.
The graph below (Figure 15) shows how production of ozone - hence smog - varies
with NOx concentration. As NOx’s are produced from anthropogenic sources (car
exhausts and industry), their concentration is higher in urban areas than rural:
Figure 15: Graph to show the variation of ozone production (hence photochemical smog) with
NOx concentration1
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•
GCE Green Chemistry Teacher’s Pack
At low NOx concentrations, there is more ozone than NOx, so ozone is destroyed
by CO. This typically occurs in a marine environment1.
•
At zero net production of ozone (the compensation point1), the amount of
destruction equals the amount of ozone production.
•
Above the compensation point there is a linear increase in the production of ozone
with increasing NOx concentrations (as explained by the mechanism above).
•
This reaches a maximum, and at very high NOx concentrations, there is
retardation of ozone production as radical termination reactions occur1:
NO 2 + OH • + M → HNO3 + M
(M is an unreactive species in the air which can absorb excess energy from the
reactants)
So we see lower levels of smog with extremely high concentrations of NOx.
Ozone only lasts on the surface for a few hours, so one can monitor the changes in
ozone levels throughout the day. For example, on a sunny day when NOx levels peak
early in the early morning rush hour, ozone levels increase for a few hours afterwards
and then drop in the evening. However, ozone rises slightly into the troposphere and
remains there for a few months which is an unwanted effect as ozone is a greenhouse
gas3.
2.4. UK Emissions Today
•
CO and SO2 levels have decreased since the 1970s3
o This is because society has made an effort in cleaning up these pollutants,
through legislations such as the Clean Air Act (introduced in 1956). This act
created zones where smokeless fuels had to be burnt, moved power stations to
rural areas (to reduce urban smog), and required industries and factories to
have tall chimneys to disperse air pollution 21 . Since then, power stations and
factories are required to have filters and catalytic converters in their chimneys.
•
NOx and VOC levels have stayed almost constant, as although cars’ emissions
have been cleaned up (through the use and improvement of catalytic converters),
the volume of cars has increased3
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GCE Green Chemistry Teacher’s Pack
Society has increased their awareness of pollution and its causes since the
industrial revolution. There are continuing efforts to find greener fuels to reduce
these emissions.
•
Infrared spectroscopy can be used to monitor air pollution through the use of
mobile Fourier transform spectrometers as the major components (CO, CO2, SO2,
NO2 and VOCs) can be detected.
2.5. Catalytic Converters
A catalytic converter treats the exhaust of a car before it enters the atmosphere to
remove sources of air pollution.
The main components of car emissions are N2 gas (due to the fact that the majority of
air in N2 gas, so this passes through the car’s engine), CO2 and H2O (both products of
combustion). As combustion is not always complete, trace amounts of CO, VOCs
(unburned fuel that has evaporated), and NOx’s. Besides carbon dioxide’s Global
Warming Potential, the trace gases are most harmful and causes of air pollution.
•
A catalytic converter is usually 3-way: it has a reduction catalyst, an
oxidation catalyst and a control system (which monitors the exhaust and
feeds back information to fuel injection system) 22 .
•
The reduction catalyst is the first stage and uses a ceramic (usually)
honeycomb structure coated with platinum and rhodium22. It reduces NOx
emissions by22:
o Adsorbing NO and NO2 at the catalyst surface
o The NOx then undergoes a chemical reaction whereby it bonds with the
nitrogen, hence breaking the N-O bond so oxygen is free and bonds with
other oxygen atoms to form O2. The N then reacts with other N atoms on
the catalyst’s surface and so forms N2. This is then desorbed from the
catalyst’s surface as free N2 gas.
•
The oxidation catalyst is the second stage and uses a platinum and palladium
catalyst coated again onto a ceramic honeycomb structure22. It oxidises
unburned hydrocarbons (VOCs) and CO22 by lowering the activation energy
for combustion of these reactants with the remaining oxygen supply in the
exhaust.
o The catalyst adsorbs CO and the VOCs onto the catalyst surface
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o Free O2 in the exhaust gas passes over the catalyst and the catalyst
therefore aids the oxidation of the CO and VOCs – causing a chemical
reaction to form CO2.
o This CO2 doesn’t adsorb as well onto the catalyst so is released into the
exhaust.
2.6. Acid Rain
Acid rain is caused by NOx and SO2 converting to nitric acid and sulphuric acid 23 .
Rain is normally slightly acidic anyway, having a pH of 5.523 due to the fact that
some acids are dissolved in it from natural sources of nitrogen and sulphur as well
as carbon dioxide. However precipitation in polluted areas or downwind from
polluted air is more acidic than usual due to increased levels of NOx and SO2.
These pollutants are emitted into the atmosphere through combustion of
hydrocarbons or hydrocarbons containing sulphur (the released H2S reacts with
oxygen during the combustion process: 2H 2S + 3O 2 → 2H 2 O + 2SO 2 ) 24 .
Sulphuric Acid Formation - Gas Phase Chemistry :
SO 2 + OH • → HOSO2
•
•
•
HOSO2 + O 2 → HO 2 + SO3
SO3 + H 2O → H 2SO 4
(Sulphuric Acid)
Sulphuric Acid Formation - Aqueous Phase Chemistry :
25
SO 2 (g) + H 2O(l) ⇔ SO 2 ⋅ H 2O Sulphur dioxide dissolves in water
SO 2 ⋅ H 2O ⇔ H + + HSO3
−
HSO3 ⇔ H + + SO3
−
Hydrolysis
2−
Nitric Acid Formation :
NO 2 + OH• → HNO3
(Nitric Acid)
Thus, both are oxidised by the reactive hydroxyl radical (see Section 2.2 for how
OH • is created). In the presence of water vapour, the sulphur trioxide (SO3) is
rapidly converted to sulphuric acid25. In clouds, liquid droplets of water react
much more rapidly with sulphur dioxide25.
These acids are both dissolved into water, precipitated down to earth, and the
pollutants may be deposited through dry deposition closer to their source onto
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vegetation and soils, where the above reactions occur to create acids23. If it is
dissolved into clouds, the acids may travel a long way, even hundreds of
kilometres before it falls23 – therefore pollution recognises no boundaries and acid
rain caused by UK emissions are known to have fallen in Norway and devastated
the forests. In fact, the UK accounts for at least 16% of Norway’s acid rain 26 .
Effects of Acid Rain
•
The increased acidity of rivers and lakes and so can destroy the ecosystems,
vegetation and organisms within them. pHs below 5 will stop fish eggs from
hatching and below that can kills even adult fish 27 . So as acidity increases, the
biodiversity of lakes decreases27. Acid rain has made certain fish and species of
insects extinct27.
•
Changes to the soil pH can harm plants and denature
enzymes and bacteria in the soil27.
•
Acid rain can slow the growth of vegetation and forests27.
•
There is an increase in the rate of people obtaining lung
diseases such as cancer, bronchitis, emphysema and
asthma23. This is thought to be caused by the inhalation of
small particulate matter with an effective diameter of
Image taken from: European
Environment Agency Website
http://reports.eea.europa.eu/2599
XXX/en/page009.html
10μm or less (PM10s); including sulphur23 – so acid rain
may be a contributing factor to these human health problems.
•
Limestone is easily dissolved by acids:
CaCO 3 + H 2SO 4 → CaSO 4 + CO 2 + H 2 O
So buildings and monuments have become
easily eroded26.
Image taken from:
http://upload.wikimedia.org/wikipedia/commons/t
humb/5/54/-_Acid_rain_damaged_gargoyle_.jpg/800px--_Acid_rain_damaged_gargoyle_-.jpg
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Reducing Acid Rain
Efforts are being made to reduce acid rain by removal of these nitrogen and sulphur
based pollutants at their source. Sulphur dioxide can be removed from flue gases
using calcium oxide:
CaO + SO 2 → CaSO3
This is then easily converted to CaSO4, known as gypsum24. Gypsum is a mineral
with many uses: it is an ingredient in plaster (of dry walls) and fertilisers 28 . Most of
the gypsum in the EU market is made from flue gas desulphurisation24. Through the
enforcement of this process, the United States (for example) has seen a 33% decrease
in SO2 emissions between 1983 and 200224.
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3. Global Warming
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
3. Global Warming
3.1. Greenhouse Effect
This is a process whereby the whereby the reflection of infrared (IR) radiation by the
Earth’s surface and incoming IR radiation from the sun is absorbed by greenhouse
gases in the troposphere. Solar radiation reaching the Earth’s surface is mainly visible
and UV light. The earth’s surface (vegetation, land and oceans) absorbs incoming
visible and UV light, which heats it up. It then re-radiates 4% of this as heat in the
form of IR radiation 29 . The reason why the earth doesn’t rapidly drop in temperature
at night is because greenhouse gases in the troposphere (i.e. the part of the atmosphere
closest to the Earth’s surface) absorb the radiated IR radiation in the ‘IR window’ (the
IR region of the spectrum where these gases show strong absorptions) and re-emit it
in all directions as heat, so some of this heat is transferred to the Earth and heats it up.
In addition, when the gases absorb IR radiation, their vibrational modes are excited,
so they vibrate more vigorously and are more likely to collide with other molecules,
transferring their energy. This increases the kinetic energy of other molecules and
raises the average temperature of the troposphere.60
Figure 16: Figure to show the proportions of radiation emitted, transmitted and reflected by the
atmosphere and Earth's surface29
3.2. Climate Change
Climate is “the characteristic weather of a region averaged over some period of
time”1. So a change in climate is not the same as a change in weather (which can vary
on a daily basis). Climate change today is referred to as the change in climate that has
been closely observed since the early 1990’s and is thought to be caused by global
warming 30 .
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3.3. Global Warming
Global warming is a controversial topic that has caused “intense and often emotional
debate” since the mid 1980s 31 . It refers to the recorded increase of the mean surface
temperature of the Earth, which is thought to be due to an increase in the
concentration greenhouse gases in the atmosphere due to human activity31. There has
almost definitely been a 1 degree temperature rise in the Earth’s surface temperature
over the past 100 years1.
3.3.1. Greenhouse Gases and How They Work
These greenhouse gases include H2O, CO2, CH4 and NOx molecules, and are sonamed because they absorb IR radiation and re-emit the heat in all directions –
thereby increasing the temperature of the atmosphere. NB, O2 and N2 are not
greenhouse gases since in order to absorb IR radiation there must be a change in the
dipole moment, and O2 and N2 are symmetric and so not stretch or bend would cause
a change in dipole moment. The IR radiation is of the correct frequency to be
absorbed by the electrons in the C=O bonds in carbon dioxide, O-H bonds in water
and C-H bonds in methane, causing them to vibrate, bend, rock, scissor and twist (for
example). The electrons have become excited, hence promoting the electrons to
higher vibrational energy levels. When the electrons return to their ground state, they
re-emit the energy with a frequency equal to the frequency of energy gap between the
two levels. 32
The “Greenhouse Effect” of a given gas – how much it heats up the Earth’s
atmosphere - is dependent upon: 48
•
Its atmospheric concentration: the greater the concentration of the gas, the more
molecules there are to absorb IR radiation.
•
Its ability to absorb IR radiation: some gases absorb and re-emit IR radiation more
strongly than others.
3.3.2. Evidence for Global Warming
•
The carbon dioxide and methane concentrations in the atmosphere are increasing
and they have been increasing rapidly since the 1800s.
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Figure 17: Graph to show the increase in carbon dioxide concentration in the troposphere from
1967-19971
•
Methane (CH4) levels have been increasing exponentially from the 1800s
onwards. This is probably due to thawing of permafrost which contains trapped
methane, increased rice cultivation (methanogenic bacteria live in rice paddies),
leaking of natural gases from pipelines and during transport, and increasing
ruminants1. However, some of this may be offset by the destruction of wetlands
(hence methanogenic bacteria1 – so it is a balance of the two and the former must
be dominating.
Figure 18: Graph to show the variation in methane concentrations since 1000 A.D.1
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GCE Green Chemistry Teacher’s Pack
The concentration of CO2 was constant until the industrial revolution, when a
sudden increase was observed. This shows it has increased due to the combustion
of fossil fuels1 – as CO2 is produced in combustion and the industrial revolution
saw a rapid increase1 in the use and combustion of fossil fuels for power.
Figure 19: Graph to show the change in atmospheric carbon dioxide concentrations 1006-2002
A.D. measured from Ice Core Data and from 1950s as a Direct Measurement1
Direct
Measurement
INDUSTRIAL
REVOLUTION
Ice Core Data
•
Our global consumption of fossil fuels is still increasing today, as are the CO2 and
CH4 concentrations. This coincides with the 3 hottest years on record have
occurred since 1998, and 19 of the 20 hottest years on record have occurred since
1980.
•
In summary, the carbon dioxide and methane levels are higher than anytime in the
past 1 million years and potentially the past 30 million years1. Although the levels
are probably in no way close to the highest levels in Earth history, the difference
here is that we know the rate of change of these levels is faster than it ever has
been by a large difference (millions of years versus 200 years). In addition, there
has been an (almost) certain global temperature rise of 1 degree over the past
century, despite of the high heat capacity of the oceans.
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Figure 20: The Combined Global Land and Ocean Surface Temperature Record from 1850 to
2007 33
This graph further depicts the marked increase in temperature over time: 2007 was the
8th hottest year on record, 1998, 2005, 2003, 2002, 2004, 2006 and 2001 respectively
were the top 733. What is clear is the difference between anthropogenic and natural
climate change. There are variations in the climate, and CO2 and CH4 levels over
hundreds of thousands of years. However we have never before seen fast a change
over such a short period of time. The Intergovernmental Panel on Climate Change
(IPCC) was set up to evaluate the risk of climate change and make suggestions to
attenuate those risks. The panel concluded:
"Most of the observed increase in globally averaged temperatures since the mid-20th
century is very likely due to the observed increase in anthropogenic greenhouse gas
concentrations." 34
"From new estimates of the combined anthropogenic forcing due to greenhouse gases,
aerosols, and land surface changes, it is extremely likely that human activities have
exerted a substantial net warming influence on climate since 1750."34
However, some scientists, such as the European Science and Environment Forum,
disagree that the observed climate change is caused by human activity as the climate
historically has been shown to fluctuate, prior to the onset of industry in modern
times.
In addition, the models (and assumptions within them) used has caused
disagreement.30
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3.3.3. Global Warming Potential
The Global Warming Potential (GWP) of a greenhouse gas is the ratio of global
warming per unit mass of the greenhouse gas to the global warming of CO2 per unit
mass 35 . It is therefore a measure of the relative strength of the greenhouse gas in
causing global warming.
Table 5: Table to show the GWP and lifetime in the atmosphere of selected greenhouse gases
Global Warming Potential 36
Lifetime
Greenhouse Gas
(years)35
Carbon Dioxide
20 Years
100 Years
500 Years
1
1
1
Methane
12
62
23
7
Nitrous Oxide
114
275
296
156
CFC-1135
55
4500
3400
1400
CFC-12
116
7900
8500
4200
3.3.4. Carbon Neutral, Biofuels and Carbon Footprint
Carbon neutral refers to “an activity that has no net annual carbon (greenhouse gas)
emissions to the atmosphere”. For example:
•
Bio-ethanol is in theory carbon neutral, as it is made from growing crops, mainly
corn and sugar cane, which takes its carbon from eh carbon dioxide in the air. This
carbon then is used to from the ethanol. When combusted the ethanol forms
carbon dioxide again – so there are no net emissions. It produces the same amount
of CO2 as it takes in from photosynthesis when growing. However, energy is
required to grow the crops (plant and harvest) and then to convert them to ethanol.
Fertiliser and pesticides are also required to grow the crops, however, they are
pollutants. In addition, the crops required for bio-ethanol could compete for land
with other food crops and grazing land for animals, as well as destroying natural
habitats and reducing biodiversity. It could even encourage deforestation. 37
“It remains unclear if the total carbon footprint of bioethanol is actually less
than that of fossil fuels”
•
The same principles above can be applied for carbon-neutral petrol and bio-diesel,
which is a mixture of methyl esters of long chain carboxylic acids.
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GCE Green Chemistry Teacher’s Pack
Bio-diesel (again, carbon neutral) is made through a transesterification reaction of
vegetable oils with methanol and a catalyst, such as potassium hydroxide 38 . The
source of the oil is usually rapeseed, palm or soybean. In the U.K. rapeseed is by
far the largest source of oil for producing biodiesel. 39
The three main methods of producing biodiesel are39:
•
Base-catalysed transesterification of the oil – Most common as most costefficient, requires the lowest temperatures and has 98% conversion
•
Acid-catalysed transesterification of the oil
•
Reaction of the oil to from its fatty acids, and then reacting on to form
biodiesel
The reaction is as follows39, with the esters formed being the biodiesel itself.
•
Hydrogen is also a carbon neutral fuel – it can be used in hydrogen fuel cells,
which uses oxygen as the oxidant.
•
Methanol can be made by reaction of carbon monoxide with hydrogen. CO and H2
are collectively called syngas and commercially produced by reacting methane
with water or a limited supply of oxygen.
CO + 2H 2
Cu, ZnO, Al2O3
→
CH 3OH
The modern-day catalyst for this reaction produces high selectivity and is a
mixture of copper, zinc oxide and aluminium oxide, at 50-100 atm of pressure at
2500C. 40 This is a carbon neutral process as any CO produced when combusting
natural gas can be reclaimed to produce methanol – so there are no net carbon
emissions into the atmosphere.
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Biofuels
A biofuel is a fuel made from a living things or the waste produced by them, and so is
renewable and potentially carbon neutral 41 . However this is debatable for the reasons
given in the argument for bio-ethanol. A second-generation of biofuels is begin
developed, which will use the cellulose found in plants and will potentially be more
efficient, requiring fewer plants and fuels can be created from a greater range of plants
and plant waste41, reducing the loss if biodiversity that production of biofuels causes.
The living things that can produce biofuels include41:
•
Wood
•
Biogas (methane) from animal excrements
•
Ethanol and diesel made form plants and plant waste
Carbon Footprint
“A Carbon Footprint is a measure of the impact our activities have on the
environment in terms of the amount of greenhouse gases we produce. It is measured
in units of carbon dioxide.” 42
3.3.5. Controlling Global Warming and the Kyoto Protocol
As can be seen from the above discussion, the emissions of greenhouse gases and
pollutants from anthropogenic sources leads to increased atmospheric concentrations
of these gases and hence causing ozone destruction, air pollution, acid rain, and
almost certainly global warming.
Therefore, it is important that we as a society control these emissions to minimise
climate change resulting from global warming. Possible solutions to reduce the CO2
levels and emissions are:
•
Carbon capture and storage: this involves the removal of waste carbon dioxide.
The CO2 can either be captured post-combustion of fossil fuels from power
stations by removal from the flue gases. The technology required for this is
already in place for other applications. It can also be removed pre-combustion by
separating natural gas (CH4) into H2 and CO (collectively called syngas) through
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partial oxidation. Then the CO is further oxidised through the water-gas shift
reaction ( CO + H 2O → CO 2 + H 2 ) into CO2, which is then captured. The H2 then
goes on to be used as a fuel in combustion and is cleaned of carbon. The CO2 can
be transported through corrosion resistant pipelines – which are already in place
for the transport of oil and natural gas. It can then either be converted into a liquid
to be injected deep into the oceans; or stored as a gas in empty oil and gas fields
and saline formations; or by reacting it with metal oxides; it will form stable solid
carbonate minerals. 43
•
Encouraging the more economical use of fuels; such as turning off lights when not
in the room and turning off computers and TVs when not in use, reducing the
amount of vehicles on roads by promoting public transport and by posing ‘green
taxes’ companies might be encouraged to combine lorry loads.
•
Through the use of alternative fuels, such as hydrogen, wind power, solar power,
wave power, HEP and tidal power; we would be burning fewer fossil fuels and
therefore reducing the amount of CO2, VOCs and NOx’s emitted into the
atmosphere.
•
By planting more vegetation there would be increased photosynthesis which
would remove carbon dioxide from the atmosphere.
Kyoto Protocol
Our progress in reducing greenhouse gas emissions can be enforced and monitored
through initiatives such as the Kyoto protocol.
•
The Kyoto protocol is an agreement signed in 1997 by at least 55 developed
countries pledging to cut greenhouse gas emissions to 5% below 1990 levels by
2008-2012. 44
•
The greenhouse gases they aim to reduce are: CO2 (carbon dioxide), CH4
(methane), HFCs (hydrofluorocarbons), PFCs (perfluorocarbons) and SF6 (sulphur
hexafluoride).44
•
Each country had its own specific target, for example EU countries pledged to cut
emissions by 8% and Japan by 5%.44
•
It does not require developing countries to cut emissions.44
•
The US has currently not agreed to the protocol, claiming it will significantly
affect their economy.44
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•
GCE Green Chemistry Teacher’s Pack
However many sceptics say the agreement is almost futile without US support as
they are the world’s largest emitter of greenhouse gases, in addition to the fact
developing countries (increasing polluters as they are currently going through
their own industrial revolution) do not have to cut emissions. Additionally, the
aim to reduce greenhouse gases by 5% is claimed to be no way near enough
according to many climate scientists, and about 60% cuts are required to avoid the
risks global warming presents.44
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4. Recycling
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
4. Recycling
4.1. Why Recycle
The world’s resources, such as metals and oil, are running out and finite (nonrenewable), so in order to continue to make use of these valuable substances waste
prevention has to be considered. There are also other social, economical and
environmental problems posed by waste such as where it is stored, inefficient use of
resources, the expense of disposal and creating new products from raw materials,
health risks and risks to the environment61. For example, landfill sites release
greenhouse gases (especially methane) and other toxic gases, waste can turn toxic and
local habitats are destroyed. By not recycling wood-based products deforestation
increases49, this leads to an acceleration of global warming and destruction of
ecosystems. To help deal with these issues, society, industry and governments
encourage people and companies to reduce, reuse and recycle their waste and products
they no longer need. There is an ongoing effort to change to renewable resources and
to reduce waste.
Recycling is defined as 45 :
“A resource recovery method involving the collection, separation, and
processing to specification of scrap materials and their use as raw materials
for manufacture into new products.”
Image taken from: http://www.unpluggedliving.com/wp-content/uploads/2007/08/recycling-image-small.jpg
There is a widely-accepted hierarchy for waste disposal61 as depicted in Figure 21. As
can be seen the most preferred solution is prevention – so not to use that material, or
reduction of its use. Disposal is the least favoured option as it doesn’t make use of any
of the materials’ properties.
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Figure 21: Waste Disposal and Prevention Hierarchy
Image reproduced Courtesy of Sligo County Council
The use of renewable resources, energy recovery and recycling can contribute to the
more sustainable use of materials:
•
Renewable resources can be replenished by natural processes and the rate of
replenishment is equal or greater than the rate of consumption 46 . They often
do not contribute to global warming or are far more environmentally
friendly. For example, in using plant-based substances such as wood to make
wood based products, the trees can be replanted – which is essentially a
carbon neutral exercise. In terms of energy, solar energy, tidal energy,
biomass, HEP and wind power are all examples of renewable energy. The
use of renewable resources leads to the more sustainable use of materials as
the resources can be used indefinitely.
•
Energy recovery also leads to the more sustainable use of materials, as it
ensures the usefulness of even waste products is exploited. For example, if
waste polymers are incinerated, the energy released can be used to drive a
turbine and contribute to the national grid.
•
Recycling allows for materials to be made into new products, therefore
making use of the substance the product is made from. There is increasingly
greater use of recycling of manufactured materials such as plastics, glass and
metals.
•
In addition, the chemical industry should endeavour to use industrial
processes that reduce waste products, hazardous chemicals (especially
pollutants or greenhouse gases) and maximise atom economy. If any waste
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products are produced, they should be able to be degraded into safe
substances in the environment, or recycled.
Atom economy is also important in waste prevention. Atom economy is how many
reactant atoms end up in the desired product as compared to waste products and other
by-products 47 . Percentage atom economy is defined as:
% Atom Economy =
Molecular Weight of Desired Product
× 100
Molecular Weight of all the Reactants Used
This is a more useful parameter than yields when considering green chemistry.
Although it is important to maximise the yield, atom economy is a better way of
measuring efficiencies between different reactions that have the aim of forming the
same product61. It can provide an extra diagnostic tool in measuring reaction
efficiencies and can sometimes compensate for low yields or poor selectivity. The
chemical industry is encouraged to use this concept when deciding on reaction type
for the production of polymers and medicines. For example, if a lot of CO2 is
produced as a by-product with a low atom economy then this is a disfavoured
reaction.
Catalysts can be used to lower the energy demand of a reaction, and reduce CO2
emissions from burning of fossil fuels. 48 Catalysts can be used to make a reaction
have a better atom economy or allow a different reaction to be used with fewer waste
products and better atom economy.
4.1.1. Household Waste
In 2003/04, UK households produced 30.5 million tonnes of rubbish, and only 17% of
that was recycled. This is low compared to some other EU countries: some recycle
50% of their household waste. 60% of rubbish that goes into household bins could be
recycled and up to 50% of waste in household bins could be composted. 49
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Figure 22: Pie Chart to Show the Average Constituents of Household Waste 50
Miscellaneous
Metal Packaging
Non3%
Combustibles
5%
Scrap
Nappies 2%
Wood 5%
Metal/White
Goods 5%
Glass 7%
Textiles 3%
Plastic Film 3%
Garden Waste
21%
Other
Combustibles
1%
Kitchen Waste
17%
Dense Plastic 4%
Paper and Board
18%
Soil and Other
Organics 3%
Fines 3%
4.1.2. Aluminium and Steel
In the UK, aluminium and steel are amongst the most common metal used. Despite
the fact that per capita consumption of steel has dropped since the 1970s, the
consumption of aluminium is still growing 51 . Global production of aluminium
averages to about 24 million tonnes per year and for steel 1.05 billion tonnes in 2004,
an increase on 2003 of 8.8% (and excluding china an increase of 4.5%)51. This
enormous production volume and requirement for aluminium and steel clearly causes
waste disposal issues when the metal products are no longer needed. Waste metal
makes up 8% of household waste, yet only roughly a third of metal waste is currently
recycled51. An advantage of recycling metals is that they never loose their properties
no matter how many times they are recycled.
Table 6: Advantages and Disadvantages of Recycling versus Extracting Metals
Pro-Recycling Scrap Metal
Pro-Extraction of Metals
Metal ore’s are non-renewable and therefore will eventually
run out – recycling will prevent this
Sometimes through the extraction of metals one can obtain a
much purer metal than one could through recycling
Less disruption to the landscape as not quarrying. When
quarrying for metals it makes large unsightly scars on the
landscape. In addition, it can pollute rivers, and produces a lot
of dust which can pollute the air; contributing to global
dimming and respiratory diseases 52
• When considering the energy required to recycle the metal,
one is not considering the energy required to collect all the
scrap metal from different recycling points (e.g. recycling
banks and kerb-side)
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Extracting metals can also damage the habitats and local
ecosystem of where the extraction is taking place
Ore’s are often located in remote mountainous regions52, so it
is expensive to:
• Transport the machinery to the region
• Transport the ore to the extraction plant
• Hire workers and house them, as not many people would
live in the area and potentially there is no
housing/infrastructure
• It would also be difficult to find workers to extract the
metals
GCE Green Chemistry Teacher’s Pack
• Therefore, the energy required to transport all the scrap
metal to the recycling plant, and to sort the metals may
require more than extraction of metals
It is difficult to organise and implement an efficient recycling
scheme as it requires households, industry and companies to
all contribute to the UK’s new metal supply. Different
councils organise this in different ways, operating recycling
banks or kerb-side pick up. However, this can be inefficient
with not all the household’s metal being recycled and difficult
to separate the waste
Reduces the waste entering landfill sites: fewer/smaller
landfill sites are required as metals are not going into it. This
also reduces disposal costs 53 . It also reduces the number of
dumped cars53
Not using up valuable resources: saving the metal ores and
saving resources and chemicals required in the manufacturing
process53
In extraction, in order to obtain the pure metal from it’s ore it
needs to be reduced. By recycling, one it not using up
expensive reducing agents, such as titanium metal. If carbon
(as coke or charcoal) is used, this is using up another finite
resource which also is polluting. If reducing the metal through
electrolysis this is also expensive
The recycling process creates jobs53
One of the major deciding factors is how much energy is
required to recycle compared to how much energy is used in
extracting the metal: in almost all cases, it requires much less
energy to recycle the material as can be seen from the
examples of steel and aluminium below
All steel cans are 100% recyclable51
All steel products can be recycled indefinitely (apart from
aerosol cans) so waste disposal problems are reduced
Recycling 1 tonne steel has the following environmental
benefits51:
•
•
•
•
•
•
•
•
Saves 1.5 kg of iron ore
Saves 0.5 kg of coal
Saves 40% of water usage
Carbon dioxide is emitted when making steel from iron
ore. As this is a greenhouse gas, this contributes to global
warming and climate change. Recycling 1 tonne of steel
scrap saves 80% of the CO2 emissions produced
Saves 1.28 tonnes of solid waste
Reduces air emissions by 86% (as not processing steel
from iron ore)
Reduces water pollution by 76%
Saves 75% of the energy needed to make steel from iron
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GCE Green Chemistry Teacher’s Pack
ore
Recycling 1 tonne aluminium
environmental benefits51:
•
•
•
•
•
•
has
the
following
Saves 6 tonnes of bauxite (aluminium’s ore)51
Saves 4kg of chemical products, as they are not required
to extract aluminium from its ore
Produces only 5% of the CO2 emissions compared with
extracting the metal
Aluminium can be recycled indefinitely, as reprocessing
does not damage its structure.
Aluminium is the most cost effective material to recycle
Saves 95% of the energy required to extract aluminium
In the extraction of copper the following damaging environmental
effects occur:
•
•
•
•
•
“The waste from crushing has to be removed
The land where the open mine is becomes devastated
The waste from froth flotation, called tailings, has to be
removed
Sulfur dioxide from smelting can cause acid rain”52
Electrolysis is required to extract the copper which is an
expensive process52
Scrap iron can also be used in a displacement reaction to extract
aqueous solutions of copper 54 . This is advantageous as:
•
•
•
A lot less energy is required than the traditional method of
high-temperature reduction of copper oxide with carbon –
saving money and fossil fuels.
CO2 is not being formed - which would be released into the
atmosphere, contributing to global warming - from the
reduction and from using fossil fuels for energy to heat up the
reaction mixture for the reduction
Scrap iron is reused and so isn’t immediately taking up space
in a landfill site
Issues Involved in the Recycling of Iron and Steel
The following steps are taken when recycling scrap metals:
•
Collection of scrap metals: Metals are either collected kerbside from
household waste or at recycling collection points56. Here, people are asked to
separate their waste into types (paper, metal and glass). All the waste metal is
then transported to a central recycling plant.
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Images taken from: http://www.highpeak.gov.uk/environment/recycle/locator.asp and http://www.recycling-guide.org.uk/
•
Cleaning by incineration: The scrap metal has to be cleaned as it will often
contain dirt, dyes, inks, coatings and other impurities on its surface. For
example in metal packaging, the label is often printed onto the metal. This is
done through incineration: where the scrap metals are heated to very high
temperatures (850-11000C) as this is when the hydrocarbons and organic
impurities on the metal is destroyed as well
as odorous gases and dioxins 55 . Aluminium
foil will oxidise in the incinerator releasing
energy51. This energy drives a turbine to
produce electricity, and in Ireland for
example the hot water from the heat
exchanger is used for district heating55.
Aluminium cans will melt and fall to the
bottom with the ash. This is easily separated once it has cooled and resolidified51. The ash can then be used as part of the materials required for road
construction55.
Figure 23: How Incineration Works55
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•
GCE Green Chemistry Teacher’s Pack
Sorting the metal by magnetic properties: Scrap metal is either ferrous or
non-ferrous. Ferrous means it contains iron (and so is magnetic) – so this
includes iron and steel. Non-ferrous scrap metal is everything else; for
example, aluminium, nickel, copper, lead, and precious metals.51 This
magnetic property (they are ferromagnetic or permanent magnets) of ferrous
metal allows the two groups two be easily separated and sorted: giant magnets
are used to attract ferrous scrap metal 56 (iron and steel) and the non-ferrous
metal (and other waste) is left behind for further separation. This allows the
iron and steel to go on and be easily cleaned by the incinerator (if it hasn’t
already) and re-melted to turn it into ingots (large blocks) of iron or steel49.
Images taken from: http://www.wasterec.co.uk/metals.html
•
Adjusting the composition of new steel: the molten iron is analysed just
before being re-added to the furnace. If the composition of the steel has to be
changed, the oxygen supply (how long and how much) is altered, and
sometimes some pure metal is added in small amounts to alter the
composition 57 , as it can oxidise the different components in steel.
•
Scrap is used to adjust the temperature of the furnace: To decrease the
temperature of the molten metal in the furnace, more scrap iron is added,
which cools it down. To increase the temperature maybe increased by
increasing the oxygen supply passing over the burners, or adjust the
hydrocarbon fuel supply to the burners or adjusting the temperature of the
“regenerated heat from the checkers”. 58
•
Thereafter, the molten
iron or steel is allowed to
cool into ingots and sent
to mills where the ingots
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are rolled out into large sheets of steel or iron. This gives the metal greater
“flexibility and strength” 49. From this, it can be remoulded into its required
shape to make new metal products.
Images taken from: http://img.alibaba.com/photo/50386658/Stainless_Steel_Ingots.jpg and
http://www.turkeyfryerexpress.com/images/BC1102.jpg
4.1.3. Plastics and Polymers
Plastic is in prevalent use globally; in the UK alone 275,000 tonnes of plastic49 are
consumed annually (this is about “15 million bottles per day”49). In addition, global
use of plastic is increasing year on year: in Europe for example the annual increase is
4%49. There is a severe need to recycle plastics:
•
Plastics
are
made
from
hydrocarbon
polymers, which are made from crude oil.
Crude oil is a finite and increasingly
expensive resource, so it will eventually run
out. It is therefore essential to recycle
plastics instead of disposing of them in a
landfill site as it will allow us to use crude oil more sparingly now; in addition
it allows us to make use of plastics even when the crude oil supplies are
exhausted.
•
Plastics can take 500 years to decompose49, so it uses up space in a landfill site
and is unsightly. Recycling plastics is the obvious solution to this.
Image taken from: http://practicalaction.org/practicalanswers/product_info.php?products_id=190
Disposal of Polymers
Polymers (including polyalkanes) have many advantages, they are chemically inert,
non toxic as solids, impenetrable to bacteria, waterproof (such as plastics), easy to
mould and process, economical, non-biodegradable so can last a long time, thermal
insulators, electrical insulators, have high strength, can be flexibility, have low
friction, rigid and are of low weights (which is useful for example in car
manufacturing as plastics are lighter than steel and other metals so the car has a lower
weight which causes a large reduction in fuel consumption – i.e. an environmental
advantage) 59 .
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However, a drawback of polymers is that as they are chemically inert, they are
therefore non-biodegradable, so disposal of waste polymers becomes a problem. They
can’t be left to degrade over time as they would just pollute the landscape, visually
and through leeching it can pollute the environment and the air. In addition, they are
made from crude oil which is a non-renewable fossil fuel. Obviously this will
eventually run out so as a society we have to ensure to recycling the raw material to
make the most use out of it as possible.
Today, there is an increased political and social desire to:
•
Reduce waste plastic
•
Recycle waste plastic to produce lower grade products or chemical feed
stocks
•
Or to use waste plastic for energy production (via incineration – see Table 7
below)
In the lifecycle of a polymer (manufactureÆ useÆ recycle/dispose) two things
important when applying green chemistry 60 :
•
“Minimising the hazardous waste during production of raw materials and
their resulting polymers to reduce any negative impact on the
environment”60 – this is done through careful selection of raw materials, the
synthetic route and ensuring any hazardous or environmentally unfriendly
by-products are safely removed.
•
“Reducing carbon emissions resulting from the ‘life cycle’ of a polymer”60 –
this is done by recycling and avoiding landfilling (which releases methane),
degradation to CO2 and H2O, or incineration (both of which release CO2).
Table 7: Table to illustrate different Methods of Disposal of Polymers and their respective
Advantages and Disadvantages
Method of Disposal
Advantages
Disadvantages
Reuse
• This is the best method as it
• Once the product has served
doesn’t require any extra energy
its purpose or is no longer
to change it into a different
usable, it needs to be disposed
polymeric
of somehow
product,
and
no
resources are used up
• The lifetime of the product is
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• Usually
indefinitely
cannot
reuse
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GCE Green Chemistry Teacher’s Pack
increased
• The product does not fill up a
landfill site or pollute the air
through incineration
Mechanical Recycling
to Lower Grade
• Doesn’t use up more crude oil
and our natural resources
• Reducing
the
processing
environmental
Products & Chemical
impact through not drilling for
Recycling to
oil
Monomers 61
• Collection, transportation and
recycling
costs
of
63
• Not all polymers can be
recycled. The main types
• Saves energy as not required to
which can are:
separate crude oil into fractions
- PTFE
and make polymers from the
- Polythene (PE) – both
high density and low
fractions
• Not using up landfill sites –
density
reduces the amount of waste
- Poly(propene) (PP)
requiring disposal
- Polystyrene (PS)
• Lower manufacturing costs for
- Polyvinyl chloride (PVC)
products made from recycled
• These different types (which
are collected altogether) have
rather than raw materials
• Creates jobs
• Chemical
to be separated. This is to
involves
ensure they remain pure and
breaking the polymer down to
the resulting material is the
its original monomers again. So
best possible quality material
there is a change in the chemical
with the required properties is
structure of the material, but the
produced 64 . This separation
new
into types is very challenging
recycling
substance
material
the
again,
original
hence
maintaining the quality e.g.
• Finding
uses
for
mixed
plastics is another difficulty64
nylon carpet recycling 62 . There
• Separation of plastics is even
is the advantage that the new
more complicated today as
“virgin” polymer has a higher
multi-layer packaging (where
value than their mechanically
different types of plastic are
61
recycled equivalents .
fused into one product) is
• Chemical recycling produces
become more common64. In
chemical feed stocks which can
addition another substance is
be used for cracking in the
often fused into the material
production of plastics and other
(a composite) to ensure it has
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chemicals
a
• Increased public awareness and
certain
property,
e.g.
strength 65
• The markets for recycled
consumer participation
materials can be unstable63
• Recycling plants can have
negative
environmental
impacts as they are unsightly,
provide noise pollution and
air
pollution
due
to
transportation to and from the
facilities63
• Energy
recovery
through
incineration is a less viable
option
as
combustible
materials (such as paper and
polymeric
materials)
from
removed
the
are
waste
63
stream
Degradation
• By
creating
biodegradable
recovering
the
raw
polymers, which can decompose
material so can not recycle it
by microbes or light; they can
back into a new product
be composted and degrade over
a relatively short period of time
• They
• The aim is that the polymer
biodegrades into CO2 and
produced
by
H2O,
resources
i.e.
contribute to global warming
are
renewable
Figure 24: Biodegradable Plant
Pots made from Corn Starch 66
• Not
but
this
would
isoprene, maize and starch – so
• Most biodegradable polymers
in essence the process is carbon
currently break into smaller
neutral
chains,
• Condensation polymers may be
fragments
photodegradable due to the
• Degraded
carbonyl
C=O
bond
monomers
and
polymer
could
64
which
harm the local ecosystem and
absorbs radiation, and therefore
environment, and is unknown
can break
as yet whether it does or is
• Condensation
(polyesters
or
polymers
toxic64, e.g. it could leech into
polyamides)
the soil and be poisonous to
may be hydrolysed at the ester
certain organisms
or amide group as part of the
• Currently the Food and Drug
degradation process, therefore
Administration prohibit the
they are biodegradable
use
• Biodegradable polymers do not
give off toxic fumes
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of
polymers
biodegradable
as
they
could
contaminate food (due to
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GCE Green Chemistry Teacher’s Pack
• Not
particularly
visually
them decomposing due to
light or the microbes in
polluting
food)64
Incineration
• Not using up landfill sites
• Destroying valuable organic
• Can use the energy evolved to
drive a turbine and create
resource
• Contributes
to
global
warming and CO2 and CO is
electricity
• Reliable, safe and efficient
• Destroys
potentially
released
unsafe
• Toxic fumes are given off and
have to be cleaned to prevent
material and bacteria
polluting the atmosphere
• Ash must be sent to a landfill
and may be toxic. Although it
can
used
be
surfacing
in
road
55
• High operating and set-up
costs
• Waste water is created
Landfill
• Inexpensive compared to other
waste disposal options
• Almost all waste is suitable for
67
landfill
pollution, air pollution and
takes up land space
• Offers
sometimes
the
only
possible method of disposal
once
• Wastes resources
• Causes visual pollution, noise
67
all
other
routes
are
exhausted: such as ash from
67
• Releases
unpleasant
and
potentially harmful odours
• Is a breeding ground for
pathogens
• The sites use up valuable
incinerators
• Nowadays sites are developed
to cover the landfill so pollution
is minimised and landfill gas
can be trapped – which is a
67
space
in
an
increasingly
populated world
• Sites
developed
before
legislation on covering and
useful low-level polluting fuel
preventing pollution are a
• Therefore new, well-designed
huge source of pollution with
landfills do not cause visual
67
unrestrained
leakages
of
67
pollution
landfill gas and leaching
• Old landfills allow wildlife to
• Open landfill sites which are
67
currently being filled are a
repopulate on the restored land
• There are still many large and
relatively empty landfill sites
67
source of air pollution
• Due to overpopulation in
areas (e.g. urban), there is no
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space
for
landfill
sites.
However this is where a large
majority
of
waste
is
generated. So there are high
transport costs (and hence
pollution)
to
transfer
the
waste to landfill sites further
afield67
• Despite collecting landfill gas,
a landfill site is the least
energy efficient method of
waste disposal. It requires a
lot of energy and the energy
stored in waste products is not
released to its full extent67
• Despite
old
landfills
providing areas of open space,
they often leach toxins hence
providing uninhabitable and
potentially dangerous land.
67
Farmers can also therefore not
use is to grow crops
• Currently
the
appeal
of
landfill sites is its versatility,
low cost and ease of use. This
deters
further
advances
innovative
in
use
other,
waste
and
more
disposal
methods. However the EU
Landfill
Landfill
Directive
Tax
will
and
shortly
make landfilling a much more
expensive option67, hopefully
leading to an increase in
recycling
Environmental damage is also minimised by removal of toxic waste products: such as
when incinerating halogenated polymers (such as PVC – polyvinyl chloride) as waste
products, HCl is formed – which has to be removed (by neutralisation) otherwise it
- 53 -
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
would contribute to acid rain and the Cl radical is a natural catalyst for ozone
destruction (see Section 1.5).
Poly(propene)
Poly(propene) is one of the most commonly used and recycled polymers. It is a
thermoplastic polymer 68 , made from the monomer propene.
Figure 25: Conversion of Propene (monomer) to Poly(propene) (polymer) via Addition
Polymerisation 69
The uses of poly(propene) are as follows: 69
Isotactic Poly(propene) – very strong material
•
Packaging e.g. crates
•
Textiles e.g. ropes
•
Stationery68
•
Plastic parts and reusable containers68
•
Laboratory equipment68
•
Loudspeakers68
•
Car parts68
Syndiotactic Poly(propene) – intermediate in softness and meting point
•
Packaging e.g. cling film
•
Medical appliances e.g. medical tubing, bags and pouches
Atactic Poly(propene) – much softer and has a lower melting point
•
Textiles e.g. roofing felt, thermal underwear and carpets
•
Road paint
•
Sealant
•
Adhesives
- 54 -
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
5. Lesson Plans and
Resources
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
5. Lesson Plans and Resources
Based upon the AQA AS and A Level Specification for teaching in 2008
These lesson plans have been created with the AQA GCE specification in mind.
However, the activities and resources suggested can be used for any of the
specifications and a lesson plan can be chopped and mixed with the others depending
on how you as a teacher see fit, and how it fits in with your teaching methods and
timetabling. Especially as the AQA specification attaches the green chemistry topic to
where it is relevant, instead of grouping it altogether as one; you may want to do the
same with the lesson plans on the following pages – attach certain activities, resources
and ideas to where it links in with other chemistry in the course.
- 55 -
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Lesson Plan 1 and
Resources:
AS Module 1 –
Combustion of Alkanes:
Air Pollution
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
AS Module 1 – Combustion of Alkanes: Air Pollution
Aim
•
•
•
•
Objective
To know that alkanes as well as fossil fuels, when combusted can
release a number of pollutants (NOx, CO2, CO and VOCs) into
the atmosphere
To understand that these pollutants can be removed using
catalytic converters
To understand that CO2, CH4 and H2O are greenhouse gases and
may contribute to global warming
To know that combustion of sulphur containing hydrocarbons
produces SO2 which causes air pollution and can be removed
from flue gases using calcium oxide
•
•
•
•
•
Be able to state which pollutants are formed during
combustion in an internal combustion engine
To be able to say NOx and SO2 cause acid rain
To able to state how a catalytic converter removes these
pollutants
To be able to write the equation of the removal of SO2
using calcium oxide and explain why this is done.
To able to state which gases are greenhouse gases and
that they therefore contribute to global warming
Links to AQA specification
AS Module 1:
Health and Safety Risk Assessment
•
Care must be taken when using a computer. Avoid straining
eyes and low light.
Overall Picture
•
•
•
•
•
•
To explain the pollutants emitted from combustion of alkanes and
fossil fuels
To explain which ones contribute to air pollution, which one
contribute to acid rain and which ones contribute to global
warming (i.e. they are greenhouse gases) – using video
Do a worksheet on air pollution
To explain how these pollutants are removed through catalytic
converters and calcium oxide.
To explain the CO2 produced during combustion is released into
the atmosphere
To plot graphs of pollutant levels against time
3.1.6. Combustion of Alkanes
Resources to be used
• Worksheets (attached) on air pollution. From:
D. Warren, Green chemistry: a resource
outlining areas for the teaching of green and
environmental chemistry and sustainable
development for 11-19 year old students,
Royal Society of Chemistry, 2001
• Video explaining photochemical smog:
http://video.aol.com/video/researchlearnglobal-warming-and-air-pollution/1399478
• www.airquality.co.uk – a website that
monitors hourly a number of pollutants in
many localities across the U.K.
Starter Activity
•
Have the pupils each answer the true/false quiz (attached – please feel free to photocopy) on the combustion of alkanes and go
through the answers together as a class
Main Activity Refer to Section 2, page 19, for notes.
•
•
•
Explain that when fossil fuels and alkanes are combusted the internal combustion engine can produce a large number of pollutants:
NOx (N2 + O2 + heat Æ NOx - i.e. NO and NO2), CO, CO2 and unburned hydrocarbons (i.e. VOCs). If the hydrocarbon contains
sulphur, combustion of it leads to SO2 which causes air pollution and can cause acid rain (as well as NOx). Can use video link
above to explain NOx forms ozone in the lower atmosphere, a secondary pollutant which is the main component of smog.
Ask the students to complete one/more of the attached worksheets on air pollution.
Removal processes:
o Explain how a catalytic converter removes NOx’s and oxidises unburned hydrocarbons and CO to CO2
o Explain how calcium oxide can remove sulphur dioxide from flue gases using the equation: CaO + SO 2 → CaSO 3
•
Explain, (ensuring the pupils make notes) that CO2, CH4 and H2O are greenhouse gases. They can all be released from
combustion, and methane also from gas pipe leakage, ruminants, cows and methanogenic bacteria (see Section 2 for more details).
Explain that greenhouse gases may contribute to global warming (no detail on global warming required).
Other possible activities (if time or for homework):
• Go to the www.airquality.co.uk website, or more specifically use http://www.airquality.co.uk/archive/data_selector.php, and have
in pupils, individually or in small groups, to select a type of pollutant, and an area where you are interested in and plot graphs of
how the pollutant levels change over the course of a day or week.
Plenary
•
•
•
Ask the pupils to name the pollutants formed when combusting alkanes
Ask which ones cause acid rain and which ones are greenhouse gases
Ask how they can be removed
- 56 -
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
True/False Quiz: Properties of Alkenes
1.
Alkanes are a type of hydrocarbon
True/False
2.
Alkanes cannot be used as fuels
True/False
3.
Combustion of alkanes can be complete or incomplete
True/False
4.
Incomplete combustion is when there is a large supply of oxygen
True/False
5.
Incomplete combustion produces carbon monoxide (CO)
True/False
______________________________________________________
True/False Quiz: Properties of Alkenes
6.
Alkanes are a type of hydrocarbon
True/False
7.
Alkanes cannot be used as fuels
True/False
8.
Combustion of alkanes can be complete or incomplete
True/False
9.
Incomplete combustion is when there is a large supply of oxygen
True/False
10. Incomplete combustion produces carbon monoxide (CO)
True/False
______________________________________________________
True/False Quiz: Properties of Alkenes
11. Alkanes are a type of hydrocarbon
True/False
12. Alkanes cannot be used as fuels
True/False
13. Combustion of alkanes can be complete or incomplete
True/False
14. Incomplete combustion is when there is a large supply of oxygen
True/False
15. Incomplete combustion produces carbon monoxide (CO)
True/False
______________________________________________________
True/False Quiz: Properties of Alkenes
16. Alkanes are a type of hydrocarbon
True/False
17. Alkanes cannot be used as fuels
True/False
18. Combustion of alkanes can be complete or incomplete
True/False
19. Incomplete combustion is when there is a large supply of oxygen
True/False
20. Incomplete combustion produces carbon monoxide (CO)
True/False
______________________________________________________
- 57 -
Pollutants and their effects on
living things
Air is a mixture of gases. Some of the gases are pollutants which are harmful to living things. Many of the pollutants
are by-products from industrial processes and car engines. Over the past 150 years the amount of pollutants has
varied. However, since about 1970 the levels have fallen as governments have enforced tighter pollution controls.
Some pollutants are more harmful than others.
Match the pollutant with the effect they have on living things.
Carbon monoxide (CO)
This is produced when fuel such as petrol does not burn
completely. Road vehicles account for 90% of emissions. CO
can survive for a month in the atmosphere.
This may be linked to
respiratory problems such as
chronic bronchitis and asthma.
Particulates (PM10)
These are very tiny particles (diameters <10 µm) of soot and
other solids that do not settle but remain in the air. Incomplete
combustion, especially from diesel engines, accounts for about
40% of UK emissions.
Some of these pollutants are
linked with cancer.
Hydrocarbons
Hydrocarbons in the air result from unburnt engine fuel,
escaping through the exhaust and evaporating before the fuel
reaches the engine. Road traffic accounts for about 35% of UK
emissions.
It interferes with the way that
red blood cells carry oxygen.
Once the cell is damaged it
cannot be repaired. In a
confined space it can be lethal.
Oxides of nitrogen (NOx)
In the high temperatures of an engine, nitrogen and oxygen
from the air combine to make a mixture of NO and NO2 gas.
NOx levels are greater in urban areas. After a day, NOx are
converted to nitric acid, which falls as acid rain.
These pose serious health
risks as they are small enough
to penetrate into the lungs,
causing an increased risk of
heart and lung disease.
Ozone
A secondary pollutant produced by the reaction of NO2,
hydrocarbons and sunlight. Ozone levels are lower in urban
areas than rural areas because NO destroys ozone as it is
formed. It tends to form downwind from urban centres.
This has been linked with
asthma and chronic bronchitis.
Acid rain damages vegetation,
building materials and pond
life.
Lead
Most airborne emissions of lead are a result of burning leaded petrol.
Levels of lead in the air have declined with the use of unleaded
petrol. Today, the battery industry is the biggest user of lead.
This can irritate the eyes and
air passages causing
breathing difficulties.
Sulfur dioxide (SO2)
The main source of SO2 is from coal-fired power stations and
other industry. It is a corrosive gas that combines with water
vapour to form acid rain.
A cumulative poison to the
central nervous system. It
affects the mental
development of children.
P
OP
Y
PHO
TO C
Pollutants and their effects on living things – page 1 of 2
Polluting gases
Complete and balance the following equations to determine
which pollutants are formed in the atmosphere.
A. N2(g)+ O2(g) → _________ (g)
B. ‘4CH’ +5O2(g) →
_________ (g) +2H2O(g)
where ‘4CH’ is the hydrocarbon fuel
C. ‘4CH’ + 3O2(g) →
_________ (g) + 2H2O(g)
D. ‘4CH’ + O2(g) →
_________(s) + 2H2O(g)
E. N2(g) + 2O2(g) →
F. S(s) + O2(g) →
_________(g)
_________(g)
G. 2NO(g) + O2(g) →
_________(g)
1. Give the letters of the reactions which have enough oxygen
for complete combustion to occur.
2. Give the letters of the reactions which have a limited supply of
oxygen.
3. Name the gases which contribute to the formation of acid
rain.
4. How could you check your answer to question 3
experimentally?
5. Do you think it is important to monitor air pollution? Give two
reasons to support your answer.
P
OP
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Pollutants and their effects on living things – page 2 of 2
PHO
TO C
Monitoring air pollution
Air pollution monitoring sites are used to collect data at regular intervals. When
evaluating the data, it is important to know where the sites are, because pollution
levels are different in cities and the countryside.
Urban Centre sites are located in town or city centre areas eg pedestrian precincts or
shopping areas. Sampling heights are typically within 2–3 m.
Rural sites are in open country locations distanced from population centres, roads and
industrial areas.
1. Predict which air pollution monitoring site, urban or rural, will record the highest
levels of the following pollutants.
Pollutant
Monitoring site
Reason
NOx
Ozone O3
Fine Particles PM10*
*PM10s are particulates which have a diameter less than 10µm in diameter.
2. Carefully study the real data taken at the start of 1999 from the Leeds Urban Centre
site and Narberth Rural site in Pembrokeshire.
160
140
Leeds NOx
Narberth NOx
NOx level /ppb
120
100
80
60
40
20
1
P
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PHO
TO C
Monitoring air pollution – page 1 of 3
3
5
7
Day
9
11
13
40
Ozone level /ppb
35
30
Narberth O3
25
Leeds O3
20
15
10
5
1
3
5
9
7
11
13
Day
40
35
PM10 level µg/m3
Leeds PM10
30
Narberth PM10
25
20
15
10
5
1
3
5
9
7
11
13
Day
Was your answer to question 1 correct? Say yes or no and explain why.
____________________________________________________________________
____________________________________________________________________
____________________________________________________________________
____________________________________________________________________
____________________________________________________________________
____________________________________________________________________
The average level of pollution recorded each day is shown in the above graphs. These
levels are compared to a standard to see if they pose a threat to our health or the
environment. The public are alerted, and advised to take precautions, if the pollutant
level is very high.
P
OP
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Monitoring air pollution – page 2 of 3
PHO
TO C
Air quality bands for three pollutants
Air Pollution
Low
Moderate
High
Very High
NOx / ppb,
hourly average
Less than 150
150–299
300–399
400+
Ozone / ppb
hourly average
Less than 50
59–89
90–179
180
PM10 / µg/m3
Less than 50
50–74
75–99
100+
3. At the start of 1999, how would you describe the levels of a) NOx, b) Ozone,
c) PM10 recorded in Leeds and Narberth?
a)
____________________________________________________________________
____________________________________________________________________
b)
____________________________________________________________________
____________________________________________________________________
c)
____________________________________________________________________
____________________________________________________________________
4. Do you find any of these results surprising? Explain your answer.
____________________________________________________________________
____________________________________________________________________
____________________________________________________________________
____________________________________________________________________
P
OP
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PHO
TO C
Monitoring air pollution – page 3 of 3
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Lesson Plan 2 and
Resources:
AS Module 2 –
Extraction of Metals, Acid
Rain and Recycling
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
AS Module 2 – Extraction of Metals, Acid Rain and Recycling
Aim
•
•
•
•
To discuss the environmental and economic
advantages and disadvantages of recycling metals
compared to extraction of metals, and why using
scrap iron to extract copper is advantageous.
To explain how sulphides are converted to oxides
and how this can contribute to acid rain
To understand the environmental problems of acid
rain
To explain the uses of polypropene and that it is
recycled.
Objective
•
•
•
•
•
To be able to state some advantages and disadvantages of recycling
scrap metals vs. extraction of metals
To be able to explain why it is more environmentally friendly to use
scrap iron to extract copper from aqueous solutions vs. carbon
reduction of copper oxide
To be able to describe which pollutants + hence acids cause acid rain
To be able to describe the environmental effects of acid rain
To be able to list the uses of polypropene and know that it is recycled.
Links to AQA specification
Overall Picture
•
•
•
•
•
To hold a ‘brainstorming’ session on the pros and
cons of recycling metals vs. extraction of metals.
To remind them of sulphur dioxide causing acid
rain, showing them the animation and PowerPoint
presentation.
To do a practical demonstrating the effects of acid
on limestone
To do a worksheet on Acid Rain
To explain the uses of polypropene, a recyclable
polymer.
Health and Safety Risk Assessment
• Wear protective lab coats and goggles
• Take car when pouring the acid, using latex gloves
•
if possible
Place the chalk into the beaker using tongs
AS Module 2:
AS Module 2:
3.2.7. Extraction of Metals
3.2.9. Polymerisation of Alkenes
Resources to be used
• Animation on acid rain:
http://www.dfes.gov.uk/psp/resources/Secondary/Science/
Year_9/Acid_rain/Acid_rain_animation.swf
• PowerPoint Presentation on Acid Rain: cause and effects
• Worksheet on acid rain:
http://chemistry.slss.ie/downloads/ch_ol_clozeacidrain.pd
f
• Britannica Online video on Acid Rain:
http://www.youtube.com/watch?v=RPsU8i2edo&feature=related
Starter Activity
•
Hold a class ‘brainstorming’ session, writing responses on a table on the board, on the environmental and economic advantages
and disadvantages of recycling metals compared to extraction of metals, and why using scrap iron to extract copper is
advantageous. Refer to Table 6 in Section 4.1.2 for a comprehensive table.
Main Activity
•
Remind the class of SO2 and NOx being pollutants that cause acid rain, asking how they are formed (the combustion of
hydrocarbons).
• Explain that metals are usually found in ores as oxides or sulphides and when they are extracted SO2 is released, another
anthropogenic way of contributing to acid rain.
• Show the animation on acid rain and the PowerPoint presentation, having the students make notes. If time, watch the Britannica
Online video – 2 minutes.
• Optional Practical activity in small groups or individually :
From: http://www.reachoutmichigan.org/funexperiments/agesubject/lessons/tnrcc/acidrainlesson.html
1. Explain that acids react chemically with limestone and that chalk is limestone
2. Fill a beaker 1/3 full with vinegar or H2SO4
3. Add a piece of chalk to the glass
4. Have the students write what they see happening
5. Discuss their observations and inferences. Discuss the slow deterioration of statues and buildings due to the weak acid rain
that falls on some statues and buildings. If the stone is limestone or has limestone in it, the deterioration is more rapid.
• Ask the students to do the worksheet on acid rain, where the students can fill in the gaps.
• To list the uses of polypropene (refer to the end of 4.1.3), and that it is a recyclable polymer.
Plenary
•
•
•
•
Ask the students to name one advantage and one disadvantage on the recycling of scrap iron
Ask the students which pollutants contribute to acid rain, which acids they form and how they are formed
Ask the class to list 3 problems caused by acid rain
Ask the class to name 3 uses of polypropene
- 58 -
Acid Rain
Lisette Voûte
Acid rain is caused by NOx and SO2
converting to nitric acid (HNO3) and
sulphuric acid (H2SO4)
Sulphuric
SO
2
Acid Formation
+ OH
•
HOSO
2
•
SO 3 + H 2 O → H 2 SO
:
•
→ HOSO
+ O 2 → HO
- Gas Phase Chemistry
2
•
+ SO 3
2
(Sulphuric
4
Acid)
Sulphuric Acid Formation - Aqueous Phase Chemistry :
SO 2 (g) + H 2 O (l) ⇔ SO 2 ⋅ H 2 O
Sulphur dioxide dissolves in water
SO 2 ⋅ H 2 O ⇔ H + + HSO
HSO
−
3
⇔ H + + SO 3
2
+ OH
•
→ HNO
Hydrolysis
2−
Nitric Acid Formation
NO
−
3
:
3
(Nitric Acid)
Rain is normally slightly acidic anyway,
having a pH of 5.5, due to:
– Natural sources of sulphur: from volcanoes,
bacteria in soils and lightning
– Natural sources of nitrogen: ammonia from
fertilisers and manure
• But…precipitation in polluted areas or downwind
from polluted air is more acidic than usual due to
increased levels of NOx and SO2
• These pollutants are emitted into the
atmosphere through:
– Combustion of hydrocarbons
– Combustion of hydrocarbons containing sulphur
– Extraction of metals from their metal sulphide form
O 2 + N 2 + heat (from combustion ) → NOx (i.e. NO 2 and NO)
2H 2 S + 3O 2 → 2H 2 O + 2SO 2
2PbS + 3 O 2 → PbO + PbS + SO 2
2
PbS + 2PbO → 3Pb + SO 2
• These acids are both:
– Dissolved into water in clouds & precipitated down
to earth
• This means the clouds could travel a long way (100’s of km!)
before it falls
• UK emissions are known to have fallen in Norway and
devastated the forests. In fact, the UK accounts for at least
16% of Norway’s acid rain
– The pollutants may be deposited through dry
deposition onto vegetation and soils, where they
create acids directly
Problems of Acid Rain
• The increased acidity of rivers and lakes and so
can destroy the ecosystems, vegetation and
organisms within them.
• pHs below 5 will stop fish eggs from hatching and
below that can kills even adult fish.
• So as acidity increases, the biodiversity of lakes
decreases.
• Acid rain has made certain fish and species of
insects extinct.
• Acid rain can slow
the growth of
vegetation and
forests.
• Changes to the soil
pH can harm plants
and denature
enzymes and
bacteria in the soil.
• There is an increase in the rate of people
obtaining lung diseases such as:
– Cancer
– Bronchitis
– Emphysema
– Asthma
• This is thought to be caused by the inhalation
of small particulate matter (PM10s); including
sulphur – so acid rain may be a contributing
factor to these human health problems.
• Limestone is easily dissolved by acids:
CaCO 3 + H 2SO 4 → CaSO 4 + CO 2 + H 2 O
• So buildings and monuments have
become easily eroded
Acid Rain
The following words are omitted from the passage below:
sulfur dioxide, volcanoes, limestone, power plants, fossil fuels, acid rain, pH,
oxidized, rain
Write down in the spaces given below the appropriate missing word corresponding to
each of the numbers 1 to 9.
of approximately 5.5 due to
Unpolluted rainwater is slightly acidic. It has a (1)
carbon dioxide in the atmosphere. Emissions of acidic gases lead to a lowering of pH,
with the formation of acid rain. The main causes of acid rain are the oxides of
(2)
damages soil, causes the poisoning of
nitrogen and the oxides of sulfur.
fish, attacks trees and erodes buildings.
Oxides of nitrogen
Oxides of nitrogen are released from car exhausts and _____ (3) ___where the high
temperatures bring about the oxidation of atmospheric nitrogen. They are also formed
in some biological processes and by lightning discharges. Nitrogen monoxide, NO, is
quickly
(4)
in air to nitrogen dioxide, NO2. Nitrogen dioxide dissolves in
water and reacts to form a mixture of HNO2 and HNO3. When this occurs in the
atmosphere the result is acid rain.
Oxides of sulfur
Oxides of sulfur are mainly formed by the combustion of fossil fuels, particularly
coal. They can also be released by
(5)
and by the decay of organic matter.
When
(6)
are burned the sulfur in them forms sulfur dioxide, SO2, which is
a dangerous pollutant. Sulfur dioxide dissolves in water to form sulfurous acid,
H2SO3. In the atmosphere, sulfur dioxide is oxidised to sulfur trioxide, SO3. Sulfur
trioxide dissolves in rainwater to form sulfuric acid, H2SO4. Thus the release of SO2
containing H2SO3 and H2SO4.
into the atmosphere is likely to result in (7)
Scrubbing waste gases
Limestone is used to reduce sulfur dioxide emissions from coal-fired power stations.
Coal is mixed with finely ground
(8)
. The high temperatures in the furnace
cause the limestone to decompose: CaCO3 → CaO + CO2
(9)
, preventing its release:
The calcium oxide reacts with much of the
CaO + SO2 → CaSO3
1________________
2________________
3 ________________
4________________
5________________
6________________
7________________
8________________
9________________
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Lesson Plan 3 and
Resources:
AS Module 2 –
Ozone Destruction
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
AS Module 2 – Ozone Destruction
Aim
•
•
•
Objective
To understand the benefit of ozone in the
stratosphere
To understand why CFCs form Cl atoms in the
stratosphere and how it causes the destruction of
ozone
To know of the legislation banning the use of CFCs
and of an example of one of the chlorine-free
alternatives developed by chemists
Overall Picture
•
•
•
•
•
•
•
•
To be able explain why ozone is beneficial in the stratosphere, listing
some adverse effects of increased UVB exposure
To be able to provide the mechanism of Cl catalysed ozone
destruction
To be able to explain why CFCs are a source of this Cl atom
To be able to state that CFCs are now banned by the Montreal
Protocol and Chlorine-free alternatives have been developed by
chemists
Links to AQA specification
Discussion of what ozone is and why it is
beneficial, making use of the PowerPoint
presentation.
An explanation of how Cl radicals catalyse the
destruction of ozone. The notes in this resource and
the Ozone Hole Tour website are useful resources
An explanation of how CFCs are a source of the Cl
radicals using the animation.
Worksheet on atmospheric chemistry
An explanation of the solution: legislation banning
CFCs and alternatives developed.
Health and Safety Risk Assessment
•
•
Care must be taken when using a computer. Avoid
straining eyes and low light.
AS Module 2:
3.2.8. Synthesis of chloroalkanes
Resources to be used
• PowerPoint Presentation: The Effects of Ozone Depletion
• Useful website that can be used to help explain ozone
destruction: The Ozone Hole Tour6
http://www.atm.ch.cam.ac.uk/tour/
• Cl. Generation animation:
http://aes.gsfc.nasa.gov/vis/a000000/a001600/a001603/mo
leculeA.mov
• Worksheet on Atmospheric Chemistry
Starter Activity Idea taken from: http://www.floridatechnet.org/GED/LessonPlans/Science/ScienceLesson5.pdf
•
Ask “what is the ozone”? Discussion should centre on how the ozone is a layer around the earth and is important in keeping out
the harmful, high energy UVB radiation from the sun. Ask “what is meant by damage to the ozone”? Answers should include such
things as holes in the layer itself or the thinning of the layer, by CFCs. You may wish to bring in a product to demonstrate the
types of materials that are viewed as damaging to the ozone such as an aerosol can or a product that contains/formerly contained
chlorofluorocarbons.
Main Activity Please refer to sections 1.2 and 1.5 for notes.
•
•
Explain the benefit of ozone in the upper atmosphere in that it attenuates harmful UVB radiation, briefly list some of the effects of
ozone depletion (see PowerPoint presentation).
Explain why chlorine atoms catalyse the decomposition of ozone: – Use the “Ozone Hole Tour” resource to help explain.
Cl • + O 3 → ClO• + O 2
ClO • + ClO • + M → ClOOCl + M
ClOOCl + hν → Cl • + Cl • + O 2
2(Cl • + O 3 ) → 2(ClO • + O 2 )
•
•
•
Net : 2O 3 → 2O 2 - Ozone Destruction
Explain that these chlorine atoms are formed in the upper atmosphere (stratosphere) when energy from UV radiation causes C-Cl
bond to break in CFCs: light at 200 nm: CFCl3 + hv Æ CFCl2. + Cl. – Use the above animation (and, if you wish, the “Ozone
Hole Tour” resource to help explain).
Ask the students to complete the attached worksheet on Atmospheric Chemistry
Explain the aims of the Montreal protocol and Vienna Convention (see section 1.5.4) to ban the use of CFCs, and that this was
supported by chemists. We have now developed alternative-chlorine free compounds, examples of which are in section 1.5.5.
Plenary
•
•
•
Ask the students to explain why ozone is beneficial
Ask the students to state which substance catalyses the destruction of ozone and each student to come to the board to write a line
for the mechanism of its destruction
Ask the student to state what the Montreal protocol does and to provide the name of an alternative to CFCs.
- 59 -
Ozone Depletion:
Why it is Dangerous
Lisette Voûte
Humans
• Increased levels of dangerous high-energy UVB
radiation impinging upon the earth’s surface can
cause:
¾Melanoma and non-melanoma skin cancer
• A 10% reduction in ozone results in a 26%
increase in non-melanoma skin cancer
¾Eye disorders and Cataracts
¾Suppression of the immune system in people of
all races
• possibly leading to an increase in diseases and
infections
Plants
• Roughly half of the world’s plants are sensitive to
UV-B light
¾Their leaves shrink and the plant grows less when
exposed to an increase in UV-B light
• Economically, this is also problematic as:
¾It can cause reduce food yields
¾Plants also can change their chemical composition
with increased UV-B exposure, which can affect
their quality and nutrient levels
Aquatic Ecosystems
• Phytoplankton experience a similar
detrimental impact of excess UV-B
radiation that terrestrial plants do.
• This could:
¾Affect species further up the food chain
¾Have a detrimental impact on the
productivity of fisheries, amongst
others
Air Quality
• Increased levels of surface UV radiation
can increase the levels of reactive
compounds in the troposphere (lower
atmosphere), such as:
¾hydrogen peroxide
¾acids
¾ozone where levels of NOx are high –
causing smog
Materials Damage
• Photo-oxidation can occur to many
materials, e.g.:
¾Wood
¾Plastics
¾Rubber
• This is which is when the materials become
oxidized & damaged through the action of UV
light
Atmospheric Chemistry
The following words are omitted from the passage below:
oxides, increase, nitrogen, oxygen, ultraviolet, ozone, reactive,
chlorofluorocarbons, chlorine
Write down in the spaces given below the appropriate missing word corresponding to
each of the numbers 1 to 9.
The atmosphere is a thin layer of gas extending about 100 km above the earth’s
surface. It becomes less dense the higher you go. The main gases in the atmosphere
are (1)
(78% by volume), oxygen (21% by volume) and argon 1%. The
remainder of the gases are present in much smaller concentrations and include carbon
dioxide, methane, water vapour and the (2)
of nitrogen. All these gases are
present in an unpolluted environment. Human activities however (3)________ the
concentrations of some gases and add other gases including chlorofluorocarbons to
the atmosphere.
In the upper atmosphere oxygen molecules break down into oxygen atoms. These
oxygen atoms react with (4)
molecules to form ozone.
Certain atmospheric gases absorb ultraviolet radiation strongly. One of the most
important of these is (5)_____, O3. Ozone is broken down on reaction with (6)_____
radiation to form an oxygen molecule and an oxygen atom.
Ozone is a very (7)
gas. It reacts very quickly with other substances and gets
destroyed in the process. Chlorofluorocarbons are believed to be the main cause of
damage to the ozone layer.(8)
break down in the atmosphere releasing chlorine
atoms. These (9)____ atoms react with and break down ozone, thus reducing the
amount of ozone in the atmosphere. There has been an increase in skin cancer and eye
cataracts as a result of the damage to the ozone layer.
1________________
2________________
3 ________________
4________________
5________________
6________________
7________________
8________________
9________________
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Lesson Plan 4 and
Resources:
AS Module 2 –
Global Warming
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
AS Module 2 – Global Warming
Aim
•
•
•
•
Explain the terms carbon neutral and biofuel
Understand how methanol and ethanol can
produced as carbon neutral biofuels
To discuss to what extent bioethanol is carbon
neutral
To explain the link between IR radiation absorption
by the bonds in greenhouse gases and global
warming
Overall Picture
•
•
•
•
•
•
To discuss what carbon natural means
To use a PowerPoint presentation to explain the
production of methanol for use as a carbon-neutral
fuel
Explain what a bio-fuel is and how bioethanol is
made
Show the video documentary about biodiesel to lead
into a discussion why bioethanol might not be
completely carbon neutral
Explain the link between IR radiation absorption by
the bonds in greenhouse gases and global warming
The students complete the global warming
worksheet
Health and Safety Risk Assessment
•
Care must be taken when using a computer. Avoid
straining eyes and low light.
Objective
•
•
•
•
To understand the terms carbon neutral and biofuels and how it is
applied to methanol and bioethanol.
To understand how methanol and bioethanol are produces
Understand to what extent bioethanol is carbon neutral
To be able to explain the link between IR radiation absorption by the
bonds in greenhouse gases and global warming
Links to AQA specification
AS Module 2:
AS Module 2:
AS Module 2:
3.2.3. Importance of equilibria in industrial processes
3.2.10. Ethanol production
3.2.11. Infra-red spectroscopy
Resources to be used
• PowerPoint presentation: Carbon Neutral Methanol
Production
• Video Documentary about biodiesel:
http://www.soton.ac.uk/chemistry/IDECAT/
Courtesy of: Dr. David Read, University of Southampton,
Southampton, U.K.
• Worksheet on Global Warming:
http://chemistry.slss.ie/downloads/ch_ol_clozegreenhouse
.pdf
Starter Activity Please refer to section 3.3.4 for notes
•
Ask the class to name 3 greenhouse gases. Ask what it means to be a ‘greenhouse gas’ (contributes to global warming). Explain
what is meant by ‘carbon neutral’ to the class. Then ask them to come up with ideas in groups or as a class processes that could be
considered carbon neutral.
Main Activity
•
•
•
•
•
•
Explain (can be done using PowerPoint presentation) how methanol can be a carbon neutral fuel by being made by reaction of CO
with H2, explain how use of carbon neutral alcohols as fuels is an important and essential application.
Use this to lead on to explain what a biofuel is (see section 3.3.4) and how ethanol, produced by fermentation of growing crops is a
carbon-neutral biofuel
Show the video about biodiesel (analogous process to bioethanol)
Set up a class discussion as to why, through this process, it is debatable whether it is carbon neutral. (see section 3.3.4)
Explain how IR radiation excites the bonds C=O, C-H and O-H in CO2, methane and water vapour respectively causing them to
vibrate more vigorously and then re-emit the energy in all directions as heat, thereby heating up the atmosphere (see section 3.3.1).
This is why they are considered to be greenhouse gases and why they contribute to global warming; which refers to the recorded
increase of the mean surface temperature of the Earth, which is thought to be due to an increase in the concentration greenhouse
gases in the atmosphere due to human activity31.
Ask the students to complete the worksheet on Global Warming.
Plenary
•
•
•
•
Ask the students what carbon neutral means
Ask them to name one carbon neutral fuel and why
Ask the students why bioethanol might not be perfectly carbon neutral
Ask the students why CO2, H21O, and CH4 are greenhouse gases in term of IR radiation.
- 60 -
Carbon Neutral Methanol
Production
Lisette Voûte
Production of Methanol (CH3OH)
• Natural gas (methane) is combusted as a fuel to
produce energy, if it combusts incompletely –
with a limited supply of oxygen – it forms
carbon monoxide, CO.
• This can then react with hydrogen using a
catalyst to form methanol
2CH 4 + O2 → 2CO + 4 H 2
CO + 2H 2
Cu, ZnO, Al 2 O 3
→
50 −100 atm, 250 C
0
CH 3OH
Carbon Neutral
• This is a carbon neutral process because
any CO produced when combusting
natural gas can be reclaimed to produce
methanol
• When the methanol is combusted as a
fuel, the CO2 released has the same
amount of carbon that was reclaimed
from CO.
¾So there are no net carbon emissions
into the atmosphere
Greenhouse Effect and Global Warming
The following words are omitted from the passage below:
greenhouse, rubbish dumps, global warming, water vapour, combustion,
warmer, forests, CFCs
Write down in the spaces given below the appropriate missing word corresponding to
each of the numbers 1 to 8.
Some gases found in the atmosphere absorb infrared radiation, thus preventing it from
escaping into space. The effect of this is to make the earth (1)
. This greenhouse
effect keeps the earth at a comfortable temperature.
Problems arise when growing amounts of certain gases in the lower atmosphere
increase the greenhouse effect. There is evidence that the earth is getting hotter, i.e.
that (2)____ _ ____ is taking place. The effects of global warming on climate
have been predicted. These include warmer weather, changes in rainfall distribution,
rising sea levels, increased flooding and loss of land.
The most important (3)____ gases are water vapour, carbon dioxide, methane, and
chlorofluorocarbons. Water vapour makes the largest contribution to the greenhouse
in
effect because of its abundance. Global warming causes the amount of (4)
the lower atmosphere to increase. The amount of carbon dioxide in the lower
atmosphere has been increasing for the past 150 years. The (5)_____ of fossil fuels
such as coal and oil is regarded as being the principal factor leading to this increase.
makes the problem worse because it reduces
The destruction of the world’s (6)
the amount of photosynthesis taking place.
The concentration of methane in the lower atmosphere is rising. It is produced in
(7)____ ___
compost heaps, marshes , slurry pits, coal mines, and in the
digestive tracts of animals. Bacteria operating in anaerobic conditions on materials
such as carbohydrates are responsible. Chlorofluorocarbons, (8)
, also absorb
infrared radiation. The fact that they have a large greenhouse factor makes their
release into the lower atmosphere undesirable.
1________________
2________________
3 ________________
4________________
5________________
6________________
7________________
8________________
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Lesson Plan 5 and
Resources:
A2 Module 4 –
Disposal and Recycling of
Polymers
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
A2 Module 4 – Disposal and Recycling of Polymers
Aim
•
•
•
To understand that polyalkenes are chemically
inert and therefore non-biodegradable
To understand that polyesters and polyamides can
be broken down by hydrolysis and are therefore
biodegradable
To appreciate the advantages and disadvantages of
different methods of disposal of polymers
including recycling
Objective
•
•
•
To be able to state why polyalkenes cannot biodegrade
To be able to state which polymers can biodegrade and how this is
done
To be able to state at least 2 advantages and disadvantages for each
method of disposal of polymers, including recycling
Resources to be used
•
•
•
•
Video of polyester hydrolysis: http://www.coolschool.ca/lor/BI12/unit2/U02L01/Hydrolysis-maltosemovie.htm
Video of polyamide hydrolysis: http://www.coolschool.ca/lor/BI12/unit2/U02L01/Hydrolysis-depeptidemovie.htm
Worksheet: Pros and Cons of Disposal of Polymers
Slide show of the process of recycling polymers:
http://www.americanchemistry.com/s_plastics/hands_on_plastics/activities/hdpe_slideshow/slide1.html
Health and Safety Risk Assessment
• If doing the practical, wear lab coat and goggles
• Take care when handling HCl and K2HPO4, use
•
•
•
latex protective gloves
Keep test tubes in a test tube rack
Have excess acid/base at hand to neutralise any
spillages
If heating with a Bunsen burner, place away from
overhead shelving and cluttered equipment and
combustible materials, do not leave open flame
unattended and shut off gas when not in use. Use
tongs to hold the test tube to the flame.
Overall Picture
•
•
•
•
To come up with the different methods of disposal of polymers
To explain that polyalkenes are non-biodegradable, but polyesters and
polyamides are degradable: show the hydrolysis videos.
To discuss the pros and cons of the different methods of polymer
disposal, creating a table (see worksheet) as a class (possible for this
to be a homework assignment)
Optional: Practical on the hydrolysis of glycogen to glucose over
time.
Links to AQA specification
A2 Module 4:
3.4.9. Biodegradability and disposal of polymers
Starter Activity Please refer to section 4.1.3 for notes
•
To brainstorm as a class, writing on the board their answers different methods of disposal of polymers
Main Activity
•
•
•
•
Explain that due to the unreactive alkane backbone of polyalkenes (e.g. polyethene), polyalkenes are chemically inert. This causes
them to be non-biodegradable, and can take up to 500 years to start to degrade.
Explain polyesters and polyamides can be broken down by hydrolysis and are therefore biodegradable, showing the videos
Discuss as a class the advantages and disadvantages as a class for the different methods of disposal of polymers, including
recycling, writing answers on the board and in the matrix on the attached worksheet (“Pros and Cons of Disposal of Polymers”).
Alternatively, it can be written on an A3 sheet which can be photocopied and distributed to the students as it is being discussed.
Table 7 is a comprehensive table of this. You can show the slide show of the process of recycling polymers (see link above).
Optional Practical Activity: Hydrolysis of glycogen to glucose monomers (Adapted From: http://staff.science.nus.edu.sg/~dbsyhh/lab3.htm)
1. Prepare a series of seven test tubes, labeled 0 through 30 minutes in 5-minute intervals.
2. Pipette 2.5 ml of 1.0 M K2HPO4 into each tube (quenches the reaction)
3. In a separate tube, add 2.5 ml of glycogen (8 mg/ml) to 2.5 ml of 4.0 M hydrochloric acid and mix.
4. Immediately withdraw 0.5 ml of the glycogen-acid mixture and transfer it to the tube marked 0 minutes.
5. Place the remainder of the glycogen-acid solution in a vigorously boiling water bath, using a foil cap to cover the top of the
tube.
6. Remove 0.5 ml samples of 5-minute intervals for 30 minutes and transfer them to the appropriately labeled tubes containing
1.0 M K2HPO4
7. Test each of the samples for increasing levels of glucose with Benedict’s Test (a solution of copper sulphate, sodium
hydroxide, and tartaric acid) and heat.
8. Write down observations of how the colour increases in depth/intensity as more glucose is formed
Plenary
•
•
•
•
Ask the students to name the different methods of disposal of polymers
Ask the students to name 2 advantages and 2 disadvantages of recycling polymers
Say which type of polymers can be biodegraded and what is the name of the reaction type for this reaction (i.e. hydrolysis)
Ask why polyalkenes cannot biodegrade
- 61 -
Environmental
Reuse
Advantages
Disadvantages
Mechanical
Recycling to
Lower Grade
Products &
Chemical
Recycling to
Monomers
Advantages
Disadvantages
Degradation
Advantages
Economic
Social
Environmental
Disadvantages
Incineration
Advantages
Disadvantages
Landfill
Advantages
Disadvantages
Economic
Social
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Additional Green
Chemistry Resources
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
Additional Green Chemistry Resources
The following pages contain extra resources on Green Chemistry, reproduced from
the Royal Society of Chemistry.
Contents:
1. The Twelve Principles of Green Chemistry
2. Plastics from renewable raw materials and bio-degradable plastics
a. Rubber & Experiment: Vulcanisation of natural rubber
b. Lactic Acid - Production of a plastic from 2-hydroxypropanoic acid
(lactic acid)
c. Experiment: production of a film using starch
Video Documentary about biodiesel can be found at:
http://www.soton.ac.uk/chemistry/IDECAT/
AQA specification link: A2 Module 4: 3.4.5. Carboxylic acids and esters and
Acylation
Courtesy of: Dr. David Read, University of Southampton, Southampton, U.K.
- 62 -
The twelve principles of green chemistry
•
It is better to prevent waste than to treat or clean up waste after it is formed.
•
Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
•
Wherever practicable, synthetic methodologies should be designed to use
and generate substances that possess little or no toxicity to human health and
the environment.
•
Chemical products should be designed to preserve efficacy of function while
reducing toxicity.
•
The use of auxiliary substances (solvents, separation agents, etc.) should be
made unnecessary whenever possible and innocuous when used.
•
Energy requirements should be recognized for their environmental and
economic impacts and should be minimized. Synthetic methods should be
conducted at ambient temperature and pressure.
•
A raw material or feedstock should be renewable rather than depleting
whenever technically and economically practicable.
•
Unnecessary derivatization (blocking group, protection / deprotection,
temporary modification of physical / chemical processes) should be avoided
whenever possible.
•
Catalytic reagents (as selective as possible) are superior to stoichiometric
reagents.
•
Chemical products should be designed so that at the end of their function
they do not persist in the environment, and break down into innocuous
degradation products.
•
Analytical methodologies need to be further developed to allow for real-time,
in-process monitoring and control prior to the formation of hazardous
substances.
•
Substances and the form of a substance used in a chemical process should
be chosen so as to minimize the potential for chemical accidents, including
releases, explosions, and fires.
These principles have been reprinted with permission from Paul T. Anstas and John
C. Warner Green Chemistry: Theory and Practice, New York: Oxford University
Press, 1998
Plastics from renewable raw materials and biodegradable plastics
Adapted from RSC material by Horn, Bader and Buchholz
Plastics are an indispensable part of modern society. They are used in many different
areas, in daily use, in technical applications and even in medicine. It is often forgotten that
many modern developments would have been impossible without the use of plastics.
Plastics may have many advantages, but they are also the subject of environmental and
political debate. The use of fossil raw materials, such as oil, and the dangers which can
occur as a result of its extraction and transport have been criticized, as well as the
disposal and recycling of plastics. The fact that some of the problems could be reduced by
using renewable raw materials for the production of plastics has been overlooked.
Particularly in recent times the interest generated by such products has increased, or
rather, increased again. They are the subject of this section.
What are renewable raw materials?
Renewable raw materials are natural substances 1 which are used by mankind for
purposes other than nutrition or foodstuffs. [i]. They can be agricultural, forestry, animal or
microbial products. Examples of renewable raw materials which are used industrially are
cellulose, starch, sugar, oils and fats. The philosophy of the renewable management of
raw materials is that only amounts which do not upset the natural biological equilibrium are
extracted, and thus no disruption of the biosphere results.
The variety of renewable raw materials and their different chemical structures leads to a
wide range of applications. Table 1 shows some examples.
Renewable raw
material
Origin
Chemical / technical Application (example)
product
Cellulose
Wood, cotton
Cellulose ethanoate
(acetate),
cellulose ether
Vegetable oils
eg rape
Lubricant
Rape seed oil
Rape seed oil methyl Fuel
ester
Sugar (saccharose)
Sugar cane, sugar
beet
Sheets, films, filtration
condensate,
construction materials
Saccharose derivates Plastics
Starch
Wheat, potatoes,
corn, starch
Physically or
chemically modified
starch (eg
hydroxypropylstarch)
alkylpolyglycoside
Pregelatinized starch,
binders, adhesives,
finishing agents, plastics
Surfactants, detergents
Flax
Fibre flax (Flax)
Flax fibre
Fibre-reinforced materials
Latex
Rubber tree
Rubber
Car tyres
Wool
Sheep, goats etc
Clothes, insulating
materials
Table 1 Examples of the application of renewable raw materials
1
These do not include mineral oil, natural gas and coal, which are natural substances, but which
were formed a long time ago.
The use of renewable raw materials to satisfy human needs has been happening for
thousands of years. The technique of boiling soap from plant oils and wood ash was used
by the Sumerians (2500 BC). The Egyptians had already discovered the method of dyeing
using henna, a powder from the leaves of the henna shrub, in the 14th century BC, as had
the people of Asia Minor in the 13th century BC, using alizarin from the madder plant.
The accumulation and decomposition of biomass are in equilibrium: 120 thousand million
tonnes of carbon a year are bound by photosynthesis, while 60 thousand million tonnes
are respired by plants in the dark and only 60 thousand million tonnes are converted to
biomass in plants. These in turn are released into the atmosphere as carbon dioxide in
winter as a result of decay processes. This is how the annual balance is worked out. This
balance can be disrupted by natural disasters and by mankind. Around 2 thousand million
tonnes of carbon dioxide are released into the atmosphere by deforestation and soil
destruction, and five thousand million tonnes by the combustion of fossil fuels.
The total biomass of the earth amounts to 1841 thousand million tonnes, or 3.6 kg m -2 of
surface area, according to estimates. Phototrophic plants, which use photosynthesis to
produce energy, make up 99% of this figure. However, mankind uses only 3% of the
biomass produced naturally (170 thousand million tonnes), by cultivation, harvesting and
processing. This amounts to 2 thousand million tonnes of wood, 1.8 thousand million
tonnes of grain and 2 thousand million tonnes of other natural substances (eg sugar cane,
turnips, oily fruits).
This can be compared with a worldwide consumption of 7 thousand million tonnes of oil
equivalent (mineral oil, natural gas, coal) for the production of energy. Only 7% is
reprocessed by the chemical industry. In 1991, the German chemical industry covered
1.8% (10%) of its raw materials demand with renewable raw materials, and the trend is
rising [ii]. In the meantime, 10% of packaging chips are made from starch, and the
proportion of biodegradable lubricants is over 30%.
Renewable raw materials have the big advantage over fossil raw materials that the
sunlight required for their growth is available in unlimited amounts, and that the carbon
dioxide released by their combustion corresponds to the amount bound by the plant during
its growth. The chemical industry makes use of this advantage, as well as the ability of
nature to synthesize new products. Thus, for example, soap can be made from vegetable
oils in one step by an alkali ester cleavage. The basic molecular structure of the surfactant
is already present and does not need to be developed via several reactions, as for
example in the oxidation of alkanes to fatty acids.
However, even the cultivation and processing of renewable raw materials can cause
environmental pollution, for example through the over-use of fertilizers and pesticides.
Furthermore, the energy expenditure required for cultivation, harvesting and processing
should not be forgotten. There are also products which are based on renewable raw
materials which are not easily degradable or even toxic. As with any chemical product, this
has to be carefully examined and taken into account.
[i]
S. Mann, Nachwachsende Rohstoffe, Ulmer Verlag, Stuttgart 1998
[ii]
M. Eggersdorfer, Perspektiven nachwachsender Rohstoffe in Energiewirtschaft und
Chemie, Spektrum der Wissenschaft 1994 Nr. 6 S. 96-102
Plastics from renewable raw materials
Plastics based on renewable raw materials are not a novelty. The first plastics which were
used in large amounts were modified natural products. Examples include rubber based on
natural rubber (1839), celluloid from cellulose (1865) and galalith which comes from casein
in milk. (1897). Natural rubber is still an important product today. In the first third of the 20th
century, polymers based on renewable raw materials were dominant. Then gradually
plastics based on fossil raw materials began to take over, due to their ready availability
and the fact that they created completely new possibilities in the world of chemistry.
Today the interest in plastics based on renewable raw materials has increased
considerably. The aim here is to move away from petroleum-based plastics towards
renewable raw materials, whilst at the same time trying to synthesize new products with
special, desirable properties. For example, sugars are used as the alcohol components in
the production of polyurethanes, and scientists are trying to better exploit raw materials,
such as cellulose, which are available in large amounts. Products which are biologically
degradable, ie which can easily be disposed of after use, are also gaining considerable
interest. Only a few examples of the many possibilities will be illustrated here. We will start
with the classic example, rubber, and extend the range from plastics based on cellulose,
starch and 2-hydroxypropanoic acid (lactic acid) to polyurethanes made using castor oil.
Rubber
Rubber can be produced either synthetically, or from natural latex (or from mixtures of the
two). Natural rubber is, today, an alternative to synthetic products, a fact which is
highlighted by its proportion of only 30% of global rubber production.
Natural rubber is extracted from latex, the sap from the rubber tree (Hevea brasiliensis),
which oozes out of the bark when the tree is damaged. Coagulation then produces solid
natural rubber which is elastic and can be stretched considerably. If natural rubber is
kneaded with sulphur and heated to around 400 K, rubber is then formed (vulcanisation)
(Figure 2). These properties depend on the one hand on the sulfur content as well as on
other additives, eg fillers such as soot and zinc oxide in the production of tyres.
Sx
Sx
+ S8
Sx
Sx
n
Sx
Natural rubber (polyisoprene)
Rubber
Figure 2 Vulcanisation of natural rubber (poly(isoprene)) with sulfur to rubber
n
Experiment: vulcanisation of natural rubber
In the following experiment we will be working with the natural latex concentrate Kagetex®
FA [i]. The rubber dispersion is stabilized with ammonia (0.7 %) and has a dryness ratio of
61.5 %. Three types of rubber with different sulfur contents are produced [ii].
Each group of students will need
•
Eye protection
•
Three aluminium dishes
•
A glass rod
•
A mortar and pestle
•
Access to an oven or drying cabinet
•
Natural latex concentrate liquid (eg Kagetex® FA)
•
Sulfur (flammable)
Safety
•
Wear eye protection
•
Sulfur is flammable and its dust may be irritating to the eyes and respiratory
system.
Method
Different mixtures with varying amounts of sulfur are set up as shown in Table 2.
Mixture
Mass of latex / Mass of sulfur / g
g
Relative
concentration of
sulfur
1
2.0
0.1
5
2
2.0
0.5
25
3
2.0
1.0
50
Table 2 Mixtures for the experiment on vulcanisation of rubber
The sulfur is ground in the mortar and mixed with the latex in an aluminium dish. The
different mixtures are then left in the drying cabinet at 140 0C.
Disposal
The rubber waste can be disposed of with the household waste.
Observation
Elastic products are formed. The lower the sulfur content, the more elastic and soft the
rubber is.
Evaluation
The degree of hardness of the rubber depends on the sulfur content. The sulfur reacts with
the double bonds in the poly(isoprene) strands. Loose bonds are formed between the
chains (vulcanisation). The more sulfur is added to the natural rubber, the higher the
degree of intermolecular cross-linking and thus the hardness of the rubber produced.
[i]
Kautschuk-Gesellschaft mbH, Frankfurt: http://www.kautschukgesellschaft.de
[ii]
H. J. Bader (Hrsg.), Kunststoffe, Recycling, Alltagschemie, Band 12 der Reihe:
W. Glöckner, W. Jansen, R. G. Weißenhorn (Hrsg.), Handbuch der experimentellen Chemie Sekundarbereich II, Aulis-Verlag, Köln 1997
Production of a plastic from 2-hydroxypropanoic acid (lactic acid)
Each group of students will need
•
Eye protection
•
Test tubes
•
Test tube rack
•
Bunsen burner
•
Acrylic glass plate (approximately 15 cm × 15 cm)
•
Boiling stone
•
Microspatula
•
Tin(II) chloride (irritant)
•
2-hydroxypropanoic acid (lactic acid)
•
Cobalt chloride paper (dried)
Safety
• Wear eye protection
• Tin(II) chloride is irritating to eyes and skin, and is dangerous with oxidising agents
such as nitrates and peroxides
Method
5 cm3 of 2-hydroxypropanoic acid (lactic acid) are added to a microspatula load of tin(II)
chloride crystals and a boiling chip in a test tube and heated for 10 minutes whilst stirring
vigorously. The vapours emitted are tested with cobalt chloride paper. If an orange-brown
colour occurs in the test tube, the hot solution is poured onto the acrylic glass plate.
Disposal
Plastic waste can be disposed of with other household waste.
Observation
The solution applied to the acrylic glass plate solidifies and forms a yellow-brown slightly
sticky mass. The cobalt chloride paper changes colour from blue to pink.
Evaluation
Cobalt chloride paper is used to test for water which, as a result of the temperature,
escapes as steam from the test tube. The blue dry cobalt(II) chloride forms a pink-coloured
hexahydrate when combined with water.
The reaction is an esterification in which the 2-hydroxypropanoic acid (lactic acid) is
converted to oligomers with loss of water. In general one molecule of alcohol reacts with
one molecule of acid with loss of water to form an ester. The 2-hydroxypropanoic acid
used is bifunctional. It consists of one hydroxyl and one carboxyl group which form an
intermolecular bond.
-nH2O
n
2-hydroxypropanoic acid
(lactic acid)
nH2O
Poly(lactic acid)
The esterification is catalyzed by acid. The proton required for the first step of the
esterification (protonation of the carboxyl group) is formed by the hydrolysis of tin(II)
chloride with precipitation of a basic salt.
SnCl2 + H2O → Sn(OH)Cl + Cl- + H+
Experiment: production of a film using starch
Part a: extraction of starch
Each group of students will need
•
Eye protection
•
Grater
•
Tea strainer
•
Mortar and pestle
•
Two 400 cm3 beakers
•
Potatoes
•
Distilled or deionized water
Safety
•
Wear eye protection
Method
100 g of peeled and washed potatoes are ground with a grater. The mush is slurried with
100 cm3 of water and poured through a tea strainer. This process is repeated twice. The
beaker is left for five minutes until the starch is precipitated on the bottom of the beaker.
The remaining water is decanted off. Another 150 cm3 of water are added, stirred briefly
and decanted again. This process is repeated again with 100 cm3 of water.
Disposal
The potato and starch waste can be added to compost or disposed of with other
household waste.
Observation
A white powder can be extracted from the potatoes.
Evaluation
The starch in potatoes can be extracted via a number of separation and purification steps.
It is known as natural starch and is treated with cold water during the extraction process
because starch does not dissolve in cold water.
Part b: production of a film
Each group of students will need
•
Eye protection
•
One 50 cm3 round flask
•
One reflux condenser
•
One 25 cm3 measuring cylinder
•
Three 3 cm3pipettes
•
Magnetic stirrer with hot plate
•
Oil bath
•
Acrylic glass plate
•
Drying cabinet
•
Thermometer
•
Spatula
•
Natural potato starch (from Part (a) or a chemicals supplier)
•
Propane-1, 2, 3-triol (glycerol) solution 50 %
•
Hydrochloric acid 0.1 mol dm-3
•
Sodium hydroxide 0.1 mol dm-3 (irritant)
•
Food dye (liquid)
Safety
•
Wear eye protection
Method
Add 25 cm3 of water to 2.5 g of potato starch, 3 cm3 hydrochloric acid and 2 cm3 proppane1, 2, 3-triol (glycerol) solution. If you use the moist starch from Part (a), 4 g of potato
starch are required. Heat for 15 minutes under reflux. The reaction is stopped by adding
sodium hydroxide. The mixture can be dyed by adding 1 to 2 cm3 of food dye. The hot,
viscous mass is poured uniformly onto the acrylic glass plate and dried for two days at
room temperature or for approximately 90 minutes at 100 °C in a drying cabinet. If the
moist starch from Part (a) is used, the amounts should be adjusted to 4 g of starch.
Disposal
Waste can be added to compost or disposed of with other household waste. Neutralized
liquid waste can be poured down the drain.
Observation
The initially highly viscous but later thin solution can be cast into a film. The clear film can
then be removed form the glass plate.
Evaluation
Starch forms a film when it is dried from an aqueous solution. The reason for this is the
intramolecular and intermolecular hydrogen bonds, in particular those between the long
chains of the amylose molecules. However, the branched amylopectin, which makes up
the larger part of the starch molecule, inhibits film formation. The reaction with diluted
hydrochloric acid partially breaks down the amylopectin. This leads to improved film
formation, but the product is brittle. For this reason propane-1, 2, 3-triol (glycerol) is added.
Propane-1, 2, 3-triol retains water as a result of its hygroscopic properties. The bound
water inhibits the formation of crystalline and brittle areas within the molecule. Thus water
acts as a softener here as does propane-1, 2, 3-triol. The later is capable of slipping
between the starch molecules and thus making the plastic softer.
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
References
L. H. Voûte
GCE Green Chemistry Teacher’s Pack
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GCE Green Chemistry Teacher’s Pack
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