E Environmental chemistry
E1 Primary air pollution
Primary pollutants are substances released directly into the
atmosphere by either natural processes or from synthetic sources.
Secondary pollutants are formed from primary pollutants when they
react in the atmosphere.
Carbon monoxide
Carbon monoxide (CO) is a toxic gas that can be produced by human
activities and natural processes.
CO is formed by the incomplete combustion (when the oxygen
supply is insufficient) of organic material especially in forest fires and the
incomplete combustion of fossil fuels such as coal (mostly carbon).
Learning objectives
•
Describe the sources of carbon
monoxide, nitrogen oxides,
sulfur oxides, particulates and
hydrocarbons in the atmosphere
Examiner’s tip
Throughout this topic, it is
important to be able to write
chemical equations whenever
possible.
2C(s) + O2(g) → 2CO(g)
Incomplete combustion can also occur in internal combustion (car)
engines, such as the incomplete combustion of octane in petrol (gasoline):
2C8H18(l) + 17O2(g) → 16CO(g) + 18H2O(g)
This is the largest anthropogenic source of CO and so CO levels are
usually higher in urban areas, where there are more cars.
Carbon monoxide can be removed using catalytic converters fitted to
the exhaust system. The CO is removed by the following reactions:
Examiner’s tip
Combustion of petrol in the
internal combustion engine
is an anthropogenic source of
CO – anthropogenic means
that it has been produced by
human activities.
2CO + O2 → 2CO2
2CO + 2NO → 2CO2 + N2
Ninety per cent of CO comes from natural sources, in particular the
atmospheric oxidation of methane produced from anaerobic (without
oxygen) decomposition of organic matter. The exact reactions involved
are complex but the overall reaction is:
Some cars have lean burn
engines where lower CO
emissions are obtained by using a
higher oxygen to fuel ratio.
2CH4(g) + 3O2(g) → 2CO(g) + 4H2O(l)
Nitrogen oxides
The main anthropogenic sources of nitrogen oxides are the internal
combustion engine, coal, gas, and oil-fuelled power stations and heavyindustry power generation. The combustion temperature of the fuel is
very high and oxidation of atmospheric nitrogen occurs forming NO
(nitrogen monoxide, nitric oxide or nitrogen(II) oxide).
There are several different nitrogen
oxides (e.g. NO, NO2, N2O). The
symbol NOx is usually used to
represent NO and NO2 together.
N2(g) + O2(g) → 2NO(g)
Natural sources of nitrogen oxides include soil bacterial activity and
lightning. Lightning causes the high-temperature oxidation of atmospheric
nitrogen:
N2(g) + O2(g) → 2NO(g)
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There are problems with the
operation of lean-burn engines in
cars because the catalytic converter
is not as efficient at reducing
NO under lean-burn conditions.
Additional systems, such as NOx
traps, therefore have to be used to
reduce the NOx emissions from
the exhaust.
Nitrogen oxides can be removed from engine exhaust gases using catalytic
converters (forming nitrogen gas):
2CO + 2NO → 2CO2 + N2
Heavy-duty natural gas engines can be designed to produce lower NOx
emissions by increasing the oxygen to fuel ratio (lean burn).
Nitrous oxide (nitrogen(I) oxide, dinitrogen oxide, nitrous
oxide)
Denitrifying bacteria in the soil can convert nitrate (NO−3 ) into N2O –
this is a natural source. Anthropogenic sources of N2O include the use of
artificial fertilisers in agriculture (increasing the level of nitrates in the soil)
and the manufacture of nitric acid.
Sulfur oxides
Burning petrol (gasoline) and
natural gas do not contribute SO2
because the sulfur is removed
during refining.
The two main sulfur oxides are sulfur dioxide (sulfur(IV) oxide, SO2) and
sulfur trioxide (sulfur(VI) oxide, SO3).
Anthropogenic sources of SO2 include:
• burning sulfur-containing fossil fuels, especially coal and oil
S + O2 → SO2
•
smelting of metal ores
Many metal ores contain metal sulfides and these are usually
roasted in air during the reduction process prior to reducing the
ore to release the metal. The following reactions occur in the
extraction of copper:
2Cu2S + 3O2 → 2Cu2O + 2SO2
2Cu2O + Cu2S → 6Cu + SO2
•
the manufacture of sulfuric acid.
Natural sources of SO2 include volcanic activity and decay of organic
matter. Hydrogen sulfide (H2S) produced by volcanic activity or organic
decay is oxidised in the air to form SO2.
2H2S + 3O2 → 2SO2 + 2H2O
SO2 is oxidised in the air to SO3 and H2SO4 (sulfuric(VI) acid). The exact
reactions involved are complex (see Higher Level section on page 7) but
the overall equations are:
2SO2(g) + O2(g) → 2SO3(g)
SO3(g) + H2O(l) → H2SO4(aq)
SO2 is a major pollutant and, as most of the sulfur dioxide released into
the atmosphere comes from anthropogenic sources, it is important to
develop methods to reduce emissions. The sulfur can be removed from
fuels before, during or after combustion. Here we will concentrate on the
latter two methods. The methods rely on the fact that sulfur dioxide is an
acidic gas and will react with alkalis/bases.
In fluidised bed combustion, coal is mixed with limestone (CaCO3)
and air is blown through the mixture as it combusts. The limestone
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decomposes at the high temperatures involved and the CaO produced (a
basic oxide) reacts with the SO2.
CaCO3(s) → CaO(s) + CO2(g)
2CaO(s) + 2SO2(g) + O2(g) → 2CaSO4(s)
Alkaline scrubbing can also be used to remove the sulfur dioxide from the
flue gases that are produced when the fuel is burnt. The alkali is usually
sprayed into the flue gases.Various alkalis are used and result in reactions
such as these:
Ca(OH)2 + SO2 → CaSO3 + H2O
CaCO3 + SO2 → CaSO3 + CO2
Particulates (particulate matter)
These are solid or liquid particles in the air (the term for solid or liquid
particles suspended in a gas is an aerosol). Particulate matter is usually
classified according to size, such that particles with a diameter between
10 µm and 2.5 µm are termed coarse particles and those with diameters
less than 2.5 µm are called fine particles.
Anthropogenic sources of particulate matter include:
• soot from incomplete combustion of wood, coal, petrol, diesel
• fly ash – this arises when fossil fuels are burnt in furnaces and contains
soot and metal oxides
• dust from mechanical activity, demolition, etc. (including asbestos – a
fibrous silicate mineral used in the past as a flame retardant and heat
insulator)
• metal particles (including lead and mercury) from metalworking
activities.
Natural sources of particulate matter include: pollen, dust, soot from
forest fires and sea spray.
There are many ways of removing particulates from the atmosphere
including: filtration, sedimentation, scrubbing and electrostatic
precipitation.
In electrostatic precipitation, the particles pass through a series of wires
that are negatively charged. The particles pick up a negative charge as they
pass through. The particles are then attracted to a positively charged plate.
The plates are shaken periodically to remove and collect the particles.
Electrostatic precipitators can remove more than 99% of particulate matter
(e.g. from flue gases from a coal-fired power station).
Diesel engines are a major source
of particulates resulting from
transport.
Volatile organic compounds (VOC)
This covers a wide range of organic compounds that are released into the
atmosphere. Natural sources include:
• bacterial decay of organic material forming methane CH4 (marsh gas)
• plants and trees – many plants and trees (e.g. pine trees) produce
unsaturated hydrocarbons called terpenes.
Anthropogenic sources include:
• unburnt hydrocarbons from the internal combustion engine – these can
be removed by using a catalytic converter:
2C8H18(g) + 25O2(g) → 16CO2(g) + 18H2O(l)
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Levels of VOCs are often higher
indoors than outdoors due to the
use of products containing solvents.
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•
solvents used in paints, varnishes, cleaning products, glues, marker pens,
etc.
• incomplete combustion of wood and coal produces aromatic
hydrocarbons which can be carcinogenic.
The sources, effects on health and methods for removal of primary
pollutants are summarised in Table E1.
Other ways of reducing the problems associated with these pollutants
are to combust less fossil fuels and to use alternative sources (e.g. wind,
solar) for generating electricity. Pollutants from transport can be reduced
by increasing the use of public transport and decreasing the use of private
vehicles.
Pollutant
Anthropogenic sources
Natural sources
Method of
removal/reduction
Health effects
CO
incomplete combustion of
fossil fuels
anaerobic
decomposition of
organic matter
catalytic converter;
lean-burn engines
Very poisonous. CO joins onto
haemoglobin in the blood – this prevents
it from transporting oxygen molecules
in the normal way. Leads to shortage of
breath, fatigue, coma and death.
NOx
reaction of N2 and O2 in
the internal combustion
engine
lightning;
action of soil
bacteria
catalytic converter;
lean-burn engines
Toxic. Cause irritation of eyes and nose.
Cause respiratory distress due to fluid
accumulation in the lungs which can
lead to infections and death.
SO2 (SO3)
burning coal/oil;
smelting of metal ores;
sulfuric acid production
volcanic activity;
decay of organic
matter
fluidised bed
combustion;
alkaline scrubbing
Causes respiratory irritation and
infection. Damages the mucous
membranes of the nose, throat and
lungs. This is particularly bad for those
who suffer from asthma.
Causes severe eye irritation.
particulates soot from incomplete
combustion of wood, coal,
petrol, diesel;
dust from mechanical
activity, demolition, etc.
pollen, dust, soot
from forest fires
and sea spray
electrostatic
precipitation
Particles get into lungs. Can cause
coughing, bronchitis, shortness of
breath.
VOC
bacterial decay of
organic material;
many plants and
trees produce
terpenes
catalytic converter
Irritation to eye, nose and throat,
headaches, damage to liver, kidneys
and central nervous system. Aromatic
hydrocarbons (containing benzene
rings) can cause cancer.
unburnt hydrocarbons
from the internal
combustion engine;
solvents used in paints,
varnishes, etc.
Table E1 Sources, health effects and methods for removal of primary pollutants.
Test yourself
1 Draw the Lewis structure of carbon monoxide.
2 Write a balanced equation for the incomplete combustion of
butane forming carbon monoxide.
3 Write an equation for the formation of SO2 from S and explain
why it can be described as a redox reaction.
4
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E2 Acid deposition
Learning objectives
Rain is naturally acidic because of dissolved CO2.
H2O(l) + CO2(g)
H2CO3(aq)
carbonic acid
H2CO3 is a weak acid and dissociates partially according to the equation:
H2CO3(aq)
•
•
•
−
+
H (aq) + HCO3 (aq)
Due to this process the pH of rain water is about 5.6. As this is a natural
phenomenon, rain with a pH between 5.6 and 7 is not considered to be
‘acid rain’. Acid rain is therefore considered to be rain with a pH of less
than 5.6. The average pH of rain in some areas may be as low as 4.
•
State what is meant by acid
deposition
Outline the origins of acid
deposition
Describe the free radical
HL
mechanism by which sulfuric
and nitric acids are formed in
the atmosphere
Explain the role of ammonia in
acid deposition
Acid deposition is a more general term than acid rain. It refers
to any process in which acidic substances (particles, gases and
precipitation) leave the atmosphere. It can be divided into wet
deposition (acid rain, fog and snow) and dry deposition (acidic
gases and particles).
Acidic pollutants include oxides of sulfur and nitrogen. We will first of all
consider sulfur compounds.
Sulfur dioxide can be formed by various natural and anthropogenic
processes, e.g. burning of sulfur-containing fuels:
S(s) + O2(g) → SO2(g)
The processes by which SO2 is converted into SO3 and H2SO4 in the
atmosphere are complex and do not involve simple oxidation by O2
(rather interaction with hydroxyl radicals, ozone or hydrogen peroxide).
The reactions, however, can be summarised as:
2SO2(g) + O2(g) → 2SO3(g)
sulfur(VI) oxide
SO3(g) + H2O(l) → H2SO4(aq)
sulfuric(VI) acid
SO2 can also dissolve in water to produce sulfuric(IV) acid (sulfurous acid):
SO2(g) + H2O(l) → H2SO3(aq)
sulfuric(IV) acid
NO can be oxidised in the atmosphere to NO2. Again, the exact nature
of the process is complex (see Higher Level section on page 7), but the
reaction may be summarised as:
2NO(g) + O2(g) → 2NO2(g)
The NO2 can then react with a hydroxyl radical (HO•) to form nitric(V)
acid:
NO2(g) + HO•(g) → HNO3(g)
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Again, other reactions can occur and the reactions for the formation of
nitric acid may also be shown as:
4NO2(g) + O2(g) + 2H2O(l) → 4HNO3(aq)
or
2NO2(g) + H2O(l) → HNO2(aq) + HNO3(aq)
nitric(III) acid
nitric(V) acid
Problems associated with acid deposition
•
•
•
Effect on vegetation – it is not necessarily the acid itself that causes
problems. The acid (H+ ions) can displace metal ions from the soil
which are consequently washed away (particularly calcium, magnesium
and potassium). Mg2+ ions are needed to produce chlorophyll, so
plants could be prevented from photosynthesising properly. The acid
rain also causes aluminium ions to dissolve from rocks, which damages
plant roots and limits water uptake. This can cause stunted growth, and
thinning or yellowing of leaves on trees.
Lakes and rivers – aquatic life is sensitive to the pH falling below
6. Insect larvae, fish and invertebrates, among others, cannot survive
below pH 5.2. Below pH 4.0 virtually no life will survive. Acid rain can
dissolve hazardous minerals from rocks, which can accumulate in lakes
and damage aquatic life – Al3+ ions in particular damage fish gills.
Buildings – limestone and marble is eroded by acid rain and dissolves
away exposing more fresh surface to react with more acid. A typical
reaction would be:
CaCO3(s) + H2SO4(aq) → CaSO4(s) + H2O(l) + CO2(g)
•
Human health – acid irritates mucous membranes and causes
respiratory illnesses (e.g. asthma, bronchitis).
Acidic water can dissolve heavy metals such as Cu2+ or Pb2+ which are
poisonous, and Al3+ which may be linked to Alzheimer’s disease.
Methods of dealing with acid deposition
Methods for dealing with acid deposition include:
• improving the design of vehicle engines, using catalytic converters,
removing sulfur before burning fuels
• using renewable power supplies
• greater use of public transport
• designing more efficient power stations
• liming of lakes – calcium oxide or hydroxide neutralises acidity.
CaO(s) + H2SO4(aq) → CaSO4(aq) + H2O(l)
Ca(OH)2(s) + H2SO4(aq) → CaSO4(aq) + 2H2O(l)
Test yourself
4 Explain why rain is naturally acidic.
5 Give the formulas of two acids that arise from human activities
and that are present in acid rain.
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The free radical mechanism and sulfuric and nitric
acid formation
HL
Like much of the chemistry of the atmosphere, the formation of H2SO4,
HNO3 and HNO2 (the major constituents of acid deposition) is
dominated by free radical reactions.
The formation of hydroxyl radicals (HO•) can be shown as:
UV light
O3 ⎯⎯→ O• + O2
ozone
O• + H2O → 2HO•
This can also be shown in one step as:
H2O + O3 → 2HO• + O2
The hydroxyl radicals can react with the oxides of nitrogen to form
nitric(V) acid (HNO3) and nitric(III) acid (HNO2):
HO• + •NO2 → HNO3
HO• + •NO → HNO2
These are both termination
reactions and do not result in the
formation of a free radical.
The hydroxyl radical is also involved in the oxidation of sulfur(IV) oxide
to sulfuric(VI) acid (H2SO4).
HO• + SO2 → HOSO2•
HOSO2• + O2 → HO2• + SO3
SO3 + H2O → H2SO4
The role of ammonia in acid deposition
Ammonia is produced by bacterial action (nitrogen fixation) in the soil
and in the root nodules of some plants such as peas and beans (legumes).
Ammonia is a base and neutralises a significant proportion of the acids
in the atmosphere, forming ammonium salts:
2NH3 + H2SO4 → (NH4)2SO4
A major source of ammonia in the
atmosphere is from agriculture,
especially intensive animal
husbandry (animal waste).
NH3 + HNO3 → NH4NO3
When these particles sink to the ground or are washed out of the
atmosphere or rain containing ammonium salts falls they can cause
acidification of the soil. This can happen in two ways. First, the ammonium
ion is weakly acidic and can ionise to form ammonia and H+ ions:
NH4+
NH3 + H+
Ammonium salts such as ammonium sulfate and ammonium nitrate
are salts of a weak base with a strong acid and are therefore acidic.
Secondly, nitrifying bacteria in the soil can cause oxidation of the
ammonium ion to the nitrate(V) ion and this process also produces
H+ ions:
NH4+ + 2O2 → 2H+ + NO3− + H2O
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Ammonium salts may be present
in the atmosphere in solution or as
fine particles.
How much, or how little,
ammonium is converted to nitrate
(NO3−) depends on various soil
factors including the amount of
organic matter, water content,
oxygen supply, temperature and
pH. Warm, moist soils with a good
oxygen supply provide favourable
conditions for nitrification.
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HL
Test yourself
6 Draw Lewis structures for the radicals HOSO2 and HOO.
7 Explain how ammonia can both neutralise acid in the
atmosphere and make soil more acidic.
Learning objectives
•
•
•
Describe the greenhouse effect
Describe the sources and relative
effects of the main greenhouse
gases
Explain the effects of increasing
levels of greenhouse gases
When CO2 absorbs infrared
radiation it vibrates more
energetically. As it moves back
down to a lower vibrational energy
level, the extra energy is given out
again.
E3 The greenhouse effect
The greenhouse effect is an important mechanism for maintaining the
Earth’s temperature at a reasonable level. Without some sort of greenhouse
effect the Earth would be too cold to maintain life as we know it.
Of the short wavelength solar radiation that reaches the Earth, some
is reflected back into space and the rest passes through the atmosphere
to reach the Earth’s surface. The surface absorbs some of this radiation
and heats up. The warmed surface radiates longer wavelength, infrared
radiation. Some of this radiation is absorbed by greenhouse gases such
as CO2 in the atmosphere. Of the radiation absorbed by the greenhouse
gases, some is re-radiated back to Earth. The overall effect is therefore
that the heat is ‘trapped’ by the gases in the atmosphere (Figure E1). The
natural equilibrium between incoming and outgoing radiation maintains
the Earth’s mean temperature at about 15 °C.
If the level of greenhouse gases in the atmosphere increases then more
infrared radiation will be absorbed and re-radiated back to Earth and the
global temperature should increase.
The sources and relative effects of the main
greenhouse gases
The main greenhouse gases, their sources and relative heat trapping ability
are shown in Table E2.
The contribution of a particular greenhouse gas to global warming
depends on several factors: its ability to absorb infrared radiation, its
abundance in the atmosphere, its atmospheric lifetime and the wavelength
range in which it absorbs infrared radiation.
Sun
solar radiation
reflected
outgoing infrared
radiation
atmosphere
incoming solar
radiation
(shorter wavelength)
infrared radiation absorbed
by greenhouse gases
infrared radiation emitted by
Earth’s surface (longer wavelength)
Earth
Figure E1 The greenhouse effect.
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Gas
Source
H2O
evaporation from oceans and lakes on the Earth’s surface
0.1
CO2
combustion of carbon fuels and biomass
1
CH4
anaerobic decay of organic material
agriculture: rice fields, marshes, animals
26
9.6
N2O
agricultural soil management (fertilisers), nitric acid production
216
5.4
O3
photochemical smog
CFCs refrigerants, propellants, foaming agents, solvents
Heat trapping Approximate relative amounts
effectiveness emitted in US per year
100
2 000
13 000–23 000
Table E2 The main greenhouse gases, their sources and relative heat trapping ability.
CO2 has a greater influence on global warming than some of the other
gases from anthropogenic sources because, although it does not absorb
as much infrared radiation as the other gases, it is produced in greater
amounts. The contribution of each of the anthropogenic gases to global
warming between 1980 and 1990 is shown in Table E3.
When considering factors that contribute CO2 to the atmosphere
we should also consider the effects of removing the mechanisms that
reduce the amount of CO2 in the atmosphere. For instance, if areas
of forest are cleared this can increase CO2 levels in the atmosphere
in two ways – first, CO2 is not being removed by the process
of photosynthesis and second, if the wood is burnt then CO2 is
produced.
The potential for a particular gas to cause global warming can be
described in terms of its global-warming potential. The potential for
1 kg of a particular gas to cause global warming over a particular time
period (e.g. 20 years) is compared to that of 1 kg of CO2 (see Table E4).
Methane has a much higher global-warming potential than carbon dioxide,
if we compare the same mass, but is produced in much smaller amounts.
Water vapour is another important greenhouse gas and, indeed most
scientists would regard water vapour as the most important greenhouse
gas. The amount of water vapour in the atmosphere is, however, only
directly influenced to a small extent by human activities. However, if the
Earth gets hotter, through the release of other greenhouse gases, this will
increase the evaporation of water and so further increase the amount of
water vapour in the atmosphere.
Gas
Contribution to global
warming / %
CO2
55
CH4
15
N2O
6
CFCs
24
Table E3 The contribution of anthropogenic
gases to global warming between 1980 and
1990.
Gas
Global-warming
potential (20 years)
CO2
1
CH4
72
N2O
289
CFC-11
3800
Table E4 The global-warming potential of
various gases over 20 years.
The influence of increasing amounts of
greenhouse gases on the atmosphere
In the 10 000-year period up to 1750, the CO2 concentration in the
atmosphere remained fairly constant at around 280 ppm (parts per million).
This had risen to 379 ppm by 2005. Similarly, CH4 and N2O abundances in
the atmosphere have increased since the industrial revolution.
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The effect of global warming on
global food production is difficult
to estimate but, depending on the
size of the temperature change, it
could actually result in an overall
increase in the potential for global
food production.
Learning objectives
•
•
HL
•
•
•
•
Describe the processes that result
in the formation and depletion
of ozone in the stratosphere
List the ozone-depleting
pollutants and their sources
Discuss the alternatives to CFCs
Explain how the wavelength of
light affects photolysis of O2 and
O3
Explain how CFCs and NOx
catalyse ozone depletion
Explain why greater ozone
depletion occurs in polar regions
In the 100 years up to 2005, global temperatures increased by about
0.74 °C and in the past 30 years temperatures are estimated to have
increased by about 0.2 °C per decade. Most current scientific thinking
is that the Earth’s climate is warming as a result of the increase in the
amount of greenhouse gases present in the atmosphere, although there are
a few scientists who dispute the reasons for these changes and whether it
is really an anthropogenic effect.
It is difficult to predict the effect of climate change on our planet but
some suggested consequences are given here.
• As the Earth’s temperature rises, oceanic water expands – increased
sea levels could submerge low-lying areas and many islands. Large
populations live in some of these areas. Only estimates can be made
based on complex models but these predict sea-level rises of up to
about 0.5 m during the next 100 years.
• Polar ice caps could melt (the melting of any floating ice does not cause
sea levels to rise).
• Antarctic ice, glaciers and snow/ice cover on land could melt (this does
increase sea levels).
• The occurrence of extreme weather events such as floods, droughts and
heat waves could increase.
• The amount and distribution of precipitation (rain and snow) could
change.
• A warming climate may mean that commercial crops can no longer be
produced where they grow now. This could be a massive problem in
grain-producing areas that currently produce a large amount of food.
• The distribution of pests and disease-carrying insects could change (e.g.
changes in the distribution of the mosquito population could alter the
regions where malaria is a danger).
E4 Ozone depletion
Formation and depletion of ozone in the
stratosphere
Ozone, O3, is an allotrope of oxygen. It is a toxic blue gas with a
characteristic odour.
The ozone layer is a region in the stratosphere where there is a higher
concentration of ozone. The maximum concentration occurs in the lower
regions of the stratosphere between about 15 and 35 km above the Earth’s
surface. About 90% of the ozone in the atmosphere occurs in this region.
Even in the ozone layer the concentration of ozone is very low and there
is roughly only one ozone molecule for every 100 000 air molecules.
The atmosphere can be divided into different regions: the troposphere
is the region closest to the Earth and the stratosphere is the region
between about 12 and 50 km, on average, above the Earth’s surface.
UV light in the upper atmosphere enables O2 to form molecules of O3
(ozone). Higher energy, shorter wavelength UV radiation is absorbed by
oxygen molecules in the upper layers of the atmosphere, which causes the
O=O to break producing oxygen free radicals:
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UV light
O2 ⎯⎯→ 2O•
oxygen free
radicals
The UV radiation required to do this is in the UV-C part of the spectrum
and must have a wavelength of less than 242 nm to provide sufficient
energy to break the bond. The oxygen atoms (free radicals) formed can
react with molecular dioxygen to form ozone:
O• + O2 → O3
ozone
Ozone molecules are particularly effective at absorbing lower energy
(longer wavelength) UV-B radiation (and some UV-C). Ozone can be
destroyed (depleted) in two natural processes:
Ultraviolet (UV) light from
the Sun reaching the Earth can
be divided into three regions:
UV-C (highest energy, shortest
wavelength), UV-B and UV-A
(lowest energy, longest wavelength).
The wavelength ranges are:
< 280 nm
• UV-C
UV-B
280–320 nm
•
320–400 nm
• UV-A
UV light
If the rate at which ozone is produced and destroyed balance, then a
steady state is reached and the concentration of ozone remains constant.
All of the UV-C and most of the UV-B reaching the Earth from the
Sun is absorbed by O2 in the upper parts of the atmosphere and ozone
in the stratosphere before it reaches the Earth’s surface. Most of the UV
radiation reaching the Earth’s surface is the less harmful UV-A (Figure E2).
Some molecules released by human activity react very effectively with
ozone, reducing its concentration. This has created what is known as the
ozone hole or the hole in the ozone layer. The ozone hole refers to
the fact that, at certain times of year and in certain regions (centred on
the Antarctic) the concentration of ozone in the stratosphere decreases
significantly so that it falls below a certain value. The size of the ozone
hole and the time for which it exists has on average increased over the
past 30 years. This results in more UV-B radiation reaching the surface of
the Earth.
The effects of increased exposure to UV radiation on humans include
an increased risk of skin cancer and cataracts. UV radiation can also
damage plants and phytoplankton.
Pollutants that cause ozone depletion
The main pollutants responsible for depletion of ozone are
chlorofluorocarbons (CFCs) and nitrogen oxides (NOx). The main sources
of nitrogen oxides have already been discussed on page 1.
CFCs (chlorofluorocarbons), also called Freons, were developed and
used from the first half of the 20th century as non-toxic, non-flammable
substances useful as refrigerants, aerosol propellants and foaming agents for
plastics and fire extinguishers and also as solvents for cleaning. Examples
are: CCl3F (CFC-11, trichlorofluoromethane) and CCl2F2 (CFC-12,
dichlorodifluoromethane).
CFCs are so unreactive that they can pass through the troposphere and
up into the stratosphere without reacting. In the stratosphere UV light
causes the C–Cl bond to break (it is weaker than the C–F bond) to produce
chlorine free radicals, which catalyse the depletion of ozone (see below).
One chlorine free radical can destroy thousands of ozone molecules.
CHEMISTRY FOR THE IB DIPLOMA © CAMBRIDGE UNIVERSITY PRESS 2011
Other natural processes involving,
e.g. hydroxyl (OH) radicals, can
also lead to ozone depletion.
Ozone absorbs longer wavelength
radiation than O2 as the O–O
bond is weaker and less energy is
required to break it. See Higher
Level section on page 13.
O2
2O•
O3
O• + O2
UV-A
→ 2O2
UV-B
O3 + O•
ozone
This first process is just the reverse
of the formation of ozone.
UV-C
O3 ⎯⎯→ O• + O2
Earth’s surface
Figure E2 UV-B and UV-C radiation from
the Sun is absorbed by ozone and oxygen
respectively. Only UV-A radiation gets
through to the Earth’s surface.
E ENVIRONMENTAL CHEMISTRY
11
Since 1998 the production of CFCs has decreased. No CFCs have been
produced in developed countries since 1995 and the use and supply of
CFCs was banned in the EU from 2000.
Alternatives to CFCs
Compounds containing C–H
bonds are more reactive than CFCs
and are likely to react before they
reach the stratosphere, where thay
can cause ozone depletion.
12
E ENVIRONMENTAL CHEMISTRY
Since the 1987 Montreal Protocol the use of CFCs is being phased out –
no new CFCs have been produced since 1995 in industrialised countries
such as the UK, America and Japan. Total usage of CFCs has also fallen
dramatically, particularly in aerosols. The only aerosols using CFCs in these
countries are asthma inhalers and these too are being phased out.
Alternatives to CFCs were required if they were to be phased out.
The desirable properties for a replacement are that it should be nontoxic, non-flammable, not damage the ozone layer and not contribute to
global warming as well as be able to fulfil the required role as a refrigerant
etc. Groups of compounds that have been used to replace CFCs are
hydrocarbons (HCs), fluorocarbons, hydrochlorofluorocarbons (HCFCs)
and hydrofluorocarbons (HFCs).
HCFCs (hydrochlorofluorocarbons) – e.g. CHF2Cl
(chlorodifluoromethane, HCFC-22), are more reactive than CFCs due
to the presence of the C–H bond (they can react with hydroxyl radicals
in the troposphere). They are less likely to make it through to the
stratosphere without reacting. They do, however, contain a C–Cl bond and
the molecules that survive through to the stratosphere will cause ozone
depletion.
HCFC-123 (CF3CCl2H, 2,2-dichloro-1,1,1-trifluoroethane) is used
instead of CFC-11 in air-conditioning units but, although it is less
damaging to the ozone layer, it is more toxic than CFC-11. It can cause
eye irritation and studies of its effect on animals have suggested that it
could cause liver damage. HCFCs are non-flammable.
Hydrocarbons such as butane, propane and 2-methylpropane are
used as propellants in aerosols but, unlike CFCs, these are flammable.
Hydrocarbons such as butane are non-toxic but are sometimes used
in solvent abuse. Hydrocarbons do not contain a C–Cl bond that can
be broken to produce free radicals, so have zero potential for ozone
depletion.
Hydrofluorocarbons (HFCs) such as HFC-134a (CF3CH2F,
1,1,1,2-tetrafluoroethane) are used in air conditioning and refrigeration.
HFC-134a is non-toxic, non-flammable under normal conditions, and
does not cause significant depletion of ozone (the C–F bond is much
more difficult to break than the C–Cl bond).
Fluorocarbons such as CF4 (tetrafluoromethane) are non-flammable
and non-toxic and will not deplete the ozone layer due to the strong C–F
bonds. C4F10 (1,1,1,2,2,3,3,4,4,4-decafluorobutane!) can be used as a fire
suppressant and refrigerant.
The alternatives mentioned here all have one severe disadvantage; they
are all greenhouse gases and will absorb infrared radiation and contribute
to global warming. Many of them, however, have global-warming
potentials lower than CFCs. The global-warming potential of these
substances are compared in Table E5.
CHEMISTRY FOR THE IB DIPLOMA © CAMBRIDGE UNIVERSITY PRESS 2011
Test yourself
Chemical
Global-warming
potential (over 100 years)
8 Classify each of the following compounds as a CFC, an HCFC,
an HFC or an HC.
c CH2F2
a CHCl2CHF2
b CH3CH2CH2CH2CH3
d CCl3F
CFC-11
4 600
CFC-12
10 600
9 Use the IUPAC naming system to name each of the compounds
in Question 8.
HFC-134a
1 300
C4F10
8 600
HCFC-22
1 700
HCFC-123
How the wavelength of light affects photolysis of
O2 and O3
The Lewis structures for O2 and O3 are shown in Figure E3.
a 0.121 nm
HL
120
butane
3
Table E5 Global-warming potential of
various pollutants over 100 years relative
to CO2.
b
O O
O
O
O
O
O
O
Figure E3 (a) the Lewis structure of O2; (b) two alternative Lewis structures of O3.
The arrows represent delocalisation of electrons.
The O–O bond lengths in ozone are actually equal and between that
expected for an O=O double bond and an O–O single bond. Ozone has
a delocalised structure in which the two electrons in the π component of
the double bond are shared between all three O atoms. The O–O bond in
ozone is thus longer and weaker than that in O2 (Figure E4).
This means that higher energy UV radiation is required to break the
O–O bond in O2 compared to O3. UV radiation of wavelength shorter
than 242 nm is required to break the double bond in O2:
0.128 nm
O
O
0.128 nm
O
Figure E4 The O–O bond in O3.
Ozone has a bond order of 1.5
compared to a bond order of 2
for O2.
This process is called photolysis
or photodissociation.
λ < 242 nm
O2 ⎯
⎯⎯→ 2O•
oxygen free
radicals
This is in the UV-C part of the spectrum and is absorbed in the
uppermost reaches of the atmosphere.
As the bond in ozone is weaker, lower energy (longer wavelength) UV
radiation is absorbed by ozone to break the bond.
λ < 330 nm
O3 ⎯
⎯⎯→ O• + O2
ozone
Ozone thus absorbs UV-C and UV-B radiation.
Catalysis of O3 depletion by CFCs and NOx
Examiner’s tip
There is some disagreement
about the wavelength of UV
radiation required to break
apart the ozone molecule. The
value of 330 nm is given on
the IB syllabus and should be
used in examinations.
Ozone molecules are easily destroyed by free radicals. These are present in
the stratosphere as nitrogen oxides or produced when CFCs are broken
down by UV light.
CFCs
CFC molecules such as CCl2F2, are very stable at ground level, but they
are broken down by absorbing UV radiation in the upper atmosphere:
UV light
CCl2F2 ⎯⎯→ •CClF2 + Cl•
chlorine free
radical
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E ENVIRONMENTAL CHEMISTRY
13
All reactions occur in the gaseous
state.
O atoms are available from the
dissociation of oxygen or ozone.
The chlorine free radical is a
catalyst in these reactions as it is
not used up. One Cl• free radical
can destroy many thousands of O3
molecules.
HL The C–Cl bond is weaker than the C–F bond and more readily
undergoes homolytic fission. The chlorine free radical released by this
process can then take part in a chain reaction which uses up ozone and
regenerates the Cl• free radical.
•Cl + O3 → ClO• + O2
ClO• + O• → O2 + Cl•
chlorine free
radical regenerated
The net effect of these reactions can be seen if the intermediate and the
catalyst are cancelled out:
•Cl + O3 → ClO• + O2
ClO• + O• → O2 + Cl•
O3 + O• → 2O2
Nitrogen oxides, NOx
NO can be formed at high altitudes by aircraft and from the reaction of
N2O or NO2 with atomic oxygen:
•
N2O(g) + •O(g) → 2NO(g)
•NO2(g) + •O(g) → •NO(g) + O2(g)
Similar types of reactions to those discussed above for CFCs occur with
nitrogen oxides. NO and NO2 both have unpaired electrons (their total
number of electrons is an odd number).
All reactions occur in the gaseous
state.
Extension
The bond order in NO is 2.5. A
simple approach to the bonding
does not explain this – NO has one
electron in a π* antibonding molecular
orbital.
Heterogeneous catalysts are in a
different phase to the reactants and
provide a surface on which the
reaction can occur.
The nitric acid formed in these
reactions remains in the ice
particles so that it is not available
for formation of nitrogen oxides
that can reduce the concentrations
of ClO.
14
E ENVIRONMENTAL CHEMISTRY
•NO + O3 → •NO2 + O2
•NO2 + O• → O2 + •NO
O3 + O• → 2O2
The cycles of destruction of ozone described above can stop when the
free radicals involved collide:
ClO• + •NO2 → ClONO2
Ozone depletion in polar regions
During winter a large-scale system of rotating winds develops over the
Antarctic and essentially isolates the stratosphere there from the rest of
the stratosphere. This is called the polar vortex. Normally clouds do not
form in the stratosphere but, at the poles, there is no sunlight for three
months and temperatures may drop below −90 °C. At these temperatures
stratospheric clouds containing ice particles do form.
Ice crystals formed in the polar vortex can act as heterogeneous
catalysts which allow the conversion of the relatively inactive ClONO2
to the much more active Cl2 and HOCl. These reactions occur on the
surface of the ice particles:
HCl + ClONO2 → Cl2 + HNO3
H2O + ClONO2 → HOCl+ HNO3
HCl and ClONO2 are both relatively inactive chlorine-containing species
but these reactions convert them into Cl2 and HOCl which undergo
homolytic fission much more easily. In the spring the sunlight returns
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causing particularly marked ozone depletion in the atmosphere above the HL
poles due to conversion of the Cl2 and HOCl into chlorine free radicals:
UV light
Cl2 ⎯⎯→ 2Cl•
UV light
HOCl ⎯⎯→ HO• + Cl•
As the temperature warms throughout the year the polar vortex dissipates
and the ice crystals melt allowing molecules to escape into other parts of
the atmosphere and stopping this process of reactions. This allows time for
ozone concentrations to recover before the next winter.
The Arctic is warmer than the Antarctic and ozone destruction has
not happened to the same extent. Some scientists believe, however, that
the conditions over the Arctic are becoming colder so that an ozone hole
could occur in this region as well. This could have serious consequences
for large parts of North America, northern Europe and Asia.
Test yourself
10 Explain what is meant by the terms free radical and
homolytic fission.
11 Write equations to show how ozone can be depleted by CFCs
such as CCl3F. Explain why the reaction involved can be
described as a chain reaction.
E5 Dissolved oxygen
Learning objectives
The concentration of dissolved oxygen in water at 20 °C is about
9 mg dm−3 (9 ppm) but decreases with temperature. Dissolved oxygen is
important for the maintenance of life in aquatic systems. Dissolved oxygen
comes from photosynthesis and from the atmosphere.
•
The amount of oxygen needed for fish to survive is dependent on
the type of fish and the temperature. In summer, certain fish will
not survive when the dissolved oxygen level drops below 6 mg dm−3
and if it drops below about 3 mg dm−3 few fish are able to survive.
The oxygen concentration needed for fish to survive is substantially
higher at higher temperatures. In the winter, certain fish may survive
down to levels of 0.25 mg dm−3.
•
Biochemical oxygen demand, also called biological oxygen demand or
BOD, is used as a measure of the quality of water. It is a measure of the
amount of oxygen used by microorganisms to oxidise the organic matter
in the water. Any organic pollutants in river water will be decomposed
(oxidised) by microorganisms (aerobic bacteria) in the water and this
process uses up dissolved oxygen. The higher the BOD, the more organic
waste there is in the water. If, for instance, sewage is released into a river
or lake, this will greatly increase the BOD – the water is more polluted. If
the water is fast flowing, new oxygen can be dissolved fairly quickly, but
this process is much slower in still water.
CHEMISTRY FOR THE IB DIPLOMA © CAMBRIDGE UNIVERSITY PRESS 2011
•
•
Understand what is meant by
biochemical oxygen demand
Discuss anaerobic and aerobic
decomposition of organic matter
in water
Understand what is meant by
eutrophication
Discuss thermal pollution of
water
Organic matter in water might
include leaves, animal manure, dead
plants and animals. Effluent from
water treatment plants will also
contain organic matter.
E ENVIRONMENTAL CHEMISTRY
15
BOD is defined as the amount of oxygen used by the aerobic
microorganisms in water to decompose the organic matter in the
water over a fixed period of time (usually five days) at a fixed
temperature (usually 20 °C).
Water containing a high proportion
of organic matter must be diluted
before analysing for BOD.
Good-quality river water will have a BOD of less than 1 ppm. Water
is generally regarded as unpolluted if it has a BOD of less than 5 ppm.
Untreated sewage could have a BOD of 500 ppm but treated sewage from
water treatment plants should have a BOD of less than 20 ppm.
The basic principle of measuring BOD is to compare the initial
amount of dissolved oxygen in a sample of water with the amount present
when the sample has been incubated for 5 days at 20 °C. Thus, if water has
a dissolved oxygen concentration of 9 ppm and after incubation for 5 days
this has fallen to 4 ppm the BOD is 9 − 4, i.e. 5 ppm.
A typical method for determining the amount of dissolved oxygen is to
use the Winkler titration method.
The basic chemistry behind the Winkler method is that manganese(II)
sulfate is added to the water and is oxidised under alkaline conditions to
manganese(IV) by the oxygen in the water:
2Mn(OH)2(s) + O2(aq) → 2MnO(OH)2(s)
The sample is acidified with sulfuric acid to produce manganese(IV)
sulfate:
MnO(OH)2(s) + 2H2SO4(aq) → Mn(SO4)2(aq) + 3H2O(l)
Iodide ions, which have also been added are oxidised to I2 by the
manganese(IV):
Mn4+(aq) + 2I−(aq) → Mn2+(aq) + I2(aq)
The liberated iodine can then be titrated against a standard sodium
thiosulfate solution:
I2(aq) + 2S2O32−(aq) → S4O62−(aq) + 2I− (aq)
The outcome of these equations is that the number of moles of dissolved
oxygen is 14 the number of moles of sodium thiosulfate used in the
titration or the mass of oxygen is eight times the number of moles of
sodium thiosulfate.
There is some disagreement about the reactions involved and an
alternative set of equations is:
2Mn2+(aq) + 4OH−(aq) + O2(aq) → 2MnO2(s) + 2H2O(l)
MnO2(s) + 2I−(aq) + 4H+(aq) → Mn2+(aq) + I2(aq) + 2H2O(l)
2S2O32−(aq) + I2(aq) → S4O62−(aq) + 2I−(aq)
The stoichiometry of these reactions is the same as above.
It is not even certain that the Mn2+ is oxidised to manganese(IV)
– it could also be oxidised to manganese(III)
16
E ENVIRONMENTAL CHEMISTRY
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Aerobic and anaerobic decomposition of organic
material in water
With plenty of oxygen the decay of organic material will take place
aerobically forming oxides or oxyanions.
If there is insufficient oxygen, anaerobic respiration forms different,
often smelly, toxic products. See Table E6.
Aerobic decomposition results in oxidation (increase in oxidation
number) and elements in their highest oxidation states whereas anaerobic
decomposition results in the element being in a low oxidation state.
Element in organic Aerobic decay
material
produces…
carbon
Oxidation number
of element
Oxidation number
of element
+4
methane, CH4
−4
−
+5
ammonia, NH3
−3
2−
CO2
nitrogen
Anaerobic decay
produces…
nitrate, NO3
sulfur
sulfate, SO4
+6
H2S (rotten egg smell
– very poisonous gas)
−2
phosphorus
phosphate, PO43−
+5
PH3 (phosphine gas)
−3
Table E6 Aerobic and anaerobic decomposition of organic material in water.
Redox equations for the decomposition of organic
matter
The exact equations involved when organic matter is decomposed by
bacteria are complex but some examples of reactions could be:
• Carbon compounds – under aerobic condition a carbohydrate could be
oxidised to CO2 and H2O:
(CH2O)n + nO2 → nCO2 + nH2O
Under anaerobic conditions methane is formed (oxidation number of
carbon = –4):
2(CH2O)n → nCH4 + nCO2
•
Organic nitrogen compounds are first converted to the ammonium
ion/ammonia. This will occur under aerobic or anaerobic conditions.
An example of this type of reaction is:
(NH2)2CO + H2O → 2NH3 + CO2
urea
Under aerobic conditions the ammonium ion can be oxidised to the
nitrate(V) ion, e.g.
NH4+ + 2O2 → NO3−+ H2O + 2H+
or
2NH4+ + 3O2 → 2NO2− + 2H2O + 4H+
and 2NO2− + O2 → 2NO3−
•
Organic sulfur compounds are first broken down by bacteria to
hydrogen sulfide.
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E ENVIRONMENTAL CHEMISTRY
17
Under aerobic conditions the H2S is oxidised to sulfates. The exact
nature of the reaction depends on the type of bacteria involved but
redox equations for the processes could be written as:
2H2S + O2 → 2S + 2H2O
2S + 3O2 + 2H2O → 2SO42− + 4H+
Major sources of phosphorus are
from artificial fertilisers and from
phosphorus-based detergents.
Eutrophication
Eutrophication refers to enrichment of a body of water, such as a lake,
with nutrients. These nutrients are mainly nitrates and phosphates and
come from sewage effluent or the run-off from agricultural land where
artificial fertilisers are being used.
Increased levels of nitrates and phosphates in a body of water cause
increased plant and algal growth (algal bloom) (Figure E5). When
these plants and algae die they sink to the bottom of the lake and are
decomposed by aerobic bacteria. This uses up the dissolved oxygen in the
water. If the amount of organic matter to be decomposed is excessive then
all the oxygen in the water will be used up and anaerobic bacteria will
take over the decomposition of the organic matter. This produces toxic
and foul smelling substances such as hydrogen sulfide (H2S). These toxic
products cause living species to die. This causes more organic decay and
further increases the BOD – everything dies.
An excessive amount of algae on the surface of the water will
prevent light reaching the plants underneath and will reduce the
amount of photosynthesis taking place. This will also reduce the
amount of oxygen in the water as photosynthesis produces oxygen.
Figure E5 Eutrophication.
Phosphates are usually regarded as
the limiting factor for algal growth
in fresh water and it is important
to control their entry to the body
of water. Therefore it is important
to limit the use of phosphate-based
detergents.
The two main effects of eutrophication are therefore reduction of the
dissolved oxygen in the water so that it cannot support life and the
production of toxic substances in the water, again reducing its capacity to
sustain life.
Eutrophication can also occur in the sea.
Thermal pollution in water
The effect of thermal pollution
can be reduced by having cooling
towers to reduce the temperature
of the water before it is returned to
the environment.
18
E ENVIRONMENTAL CHEMISTRY
Water is an important coolant in industry and power stations. Water (e.g.
from a lake or river) is drawn into the power station and discharged back
into the body of water at a higher temperature.
Oxygen is less soluble in hot water (Figure E6) so, at higher
temperatures, less is available for animal and plant life. The higher
temperature also increases the metabolic rate of fish and other organisms
so that they use up food sources and the oxygen in the water more
rapidly. These effects can cause sharp population declines in the aquatic
environment. The species that are able to live in the water can also be
affected by its temperature.
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16
Dissolved oxygen
concentration/ppm
14
12
10
8
6
4
2
0
0
10
20
30
Temperature /°C
40
50
Figure E6 The variation in dissolved oxygen in fresh water with temperature.
Test yourself
12 a The concentration of dissolved oxygen in
a fresh sample of water at 20 °C is 9.0 ppm.
Convert this to a concentration in mol dm−3.
b The concentration of dissolved oxygen in the
sample of water when it has been incubated
at 20 °C for 5 days is 1.8 × 10−4 mol dm−3.
Convert this to a concentration in ppm.
c Calculate the BOD of this sample of water
and suggest whether it would be regarded as
polluted.
13 The Winkler method was employed to measure
the concentration of dissolved oxygen in a
sample of water. Manganese sulfate, sulfuric acid
and potassium iodide were added to 100.0 cm3
of water. The iodine liberated was titrated against
sodium thiosulfate solution of concentration
5.00 × 10−3 mol dm−3; 16.00 cm3 of sodium
thiosulfate was required for the titration.
a Calculate the number of moles of sodium
thiosulfate used.
b Calculate the number of moles of iodine
present in the solution.
c Calculate the number of moles of
manganese(IV) that produced this number of
moles of I2.
d Calculate the concentration of dissolved
oxygen in mg dm−3 and ppm.
E6 Water treatment
Primary pollutants found in waste water and their
sources
Heavy metals
There are many ‘heavy metals’ that are regarded as pollutants (e.g. lead,
mercury, chromium, copper, nickel, cadmium). Rocks and minerals which
contain these metals can lead to local pollution as can mining and mineral
processing. Small amounts of heavy metals may get into the environment
and therefore into waste water from various industrial sources. Table E7
summarises some anthropogenic sources that may release heavy metals
into the environment.
Heavy metals often accumulate in the body and can eventually lead
to various serious health problems. For instance, higher levels of lead can
impair the mental development of children; mercury can damage the
brain, central nervous system and kidneys; cadmium can cause kidney
damage, bone disease and lung and prostate cancer; chromium compounds
can cause lung cancer.
CHEMISTRY FOR THE IB DIPLOMA © CAMBRIDGE UNIVERSITY PRESS 2011
Learning objectives
•
•
•
•
•
List the primary pollutants in
waste water and their sources
Outline primary, secondary and
tertiary water treatment
Evaluate multistage flash
distillation and reverse osmosis
as ways of obtaining fresh water
from sea water
Understand what is meant by HL
the solubility product
Solve problems relating to the
removal of heavy metal ions
from water by precipitation.
E ENVIRONMENTAL CHEMISTRY
19
The term heavy metal is a very
vague and imprecise term and
there is no clear definition of what
it means. It is usually used to refer
to a group of metals and metalloids
(such as arsenic) that have harmful
environmental effects. It also often
used to describe the compounds of
these metals.
Heavy metal
Anthropogenic source
lead
iron and steel production, lead water pipes
Lead used to be used in paints and as a petrol additive but this
is no longer permitted in most countries. There may be quite
high levels of lead in soil in inner city areas where the soil has
absorbed the lead emitted when leaded petrol was still in use.
Some older homes may also contain lead-based paint.
chromium
industrial organic chemical industries, cement production,
electroplating
mercury
waste incineration, gold mining, coal combustion, the chloralkali industry, inappropriate disposal of batteries, crematoria
Some of these sources result in
heavy metals or their compounds
entering the atmosphere.
Precipitation/dry deposition can
then cause them to enter the water
system.
copper
copper water pipes, marine paint (additives designed to
control algal growth), metal-producing industries, waste
incinerators
cadmium
burning of fossil fuels, incineration of municipal waste,
smelting of zinc, lead and copper, corrosion of galvanised
water pipes, electroplating, manufacture of batteries (NiCd)
Table E7 Heavy metals and their anthropogenic sources.
Pesticides
CCI3
C
CI
CI
H
Figure E7 DDT.
Pesticides are substances used to control or destroy pests. Pesticides
include herbicides (kill plants), fungicides (kill fungi), insecticides (kill
insects), algicides (kill algae) and rodenticides (kill rodents). Some
pesticides are added directly to water and others enter the water as run-off
from agricultural land.
One of the most famous (infamous) insecticides is DDT (Figure E7).
Although relatively non-toxic to humans, DDT has been banned in many
countries due to its adverse effects on the ecosystem – birds have been
particularly affected.
O
Dioxins
O
O
Figure E8 Compounds from which dioxin is
derived.
CI
O
CI
CI
O
CI
Figure E9 Dioxin.
a
b
CI
CI
CI
Figure E10 (a) The basic biphenyl structure;
(b) a PCB.
20
E ENVIRONMENTAL CHEMISTRY
The term dioxins is usually used to describe the polychlorinated
derivatives of the compounds in Figure E8.
Dioxins are produced as byproducts in the manufacture of some
chlorinated organic compounds. They are also produced if the
temperature is not high enough (below about 1200°C) when waste
materials containing organochlorine compounds are incinerated. The
most toxic of these derivatives is called 2,3,7,8-TCDD (or just TCDD or
2,3,7,8-tetrachlorodibenzo-p-dioxin, or just dioxin). The structure of this
is shown in Figure E9.
This is very persistent in the environment (high chemical stability and
poorly biodegradable) and very poisonous. Its effects include liver damage,
the skin disease chloracne and damage to the peripheral nervous system.
PCBs
PCBs are polychlorinated biphenyl compounds. There are 209 possible
PCBs where between 1 and 10 H atoms in the basic biphenyl structure
(Figure E10a) are replaced by Cl atoms. An example of a PCB is shown in
Figure E10b.
PCBs are chemically inert, non-flammable and stable at high
temperatures. They were used in electrical transformers and capacitors
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because of their high electrical resistance and so factories making these
would have discharged PCBs into the environment. PCBs have not been
manufactured in the USA since 1979 but because they are unreactive they
persist in the environment for a long time. They also accumulate in fatty
tissue and have been linked to low reproduction rates among some marine
animals and are thought to be carcinogenic (cancer-causing) in humans.
PCBs can be passed from mother to child in milk.
Other pollutants in water
The sources of nitrates, phosphates and organic matter in water have been
discussed on page 18 and this information is summarised in Table E8.
Pollutant
Source
Effects
nitrates
waste water effluent both from cities and industrial
plants or the run-off from agricultural land where
artificial fertilisers are being used
eutrophication
levels too high in drinking water cause
methemoglobinemia (the oxygen-carrying capacity of the
blood is reduced – this particularly affects babies (blue
baby syndrome)) and possible link with some cancers
phosphate
waste water effluent or the run-off from agricultural
land where artificial fertilisers are being used;
detergents used in the home
eutrophication
organic
matter
leaves, animal manure, dead plants and animals in
water;
effluent from water treatment plants
bacteria present in human and animal wastes affect
health
fecal coliform, E. coli and cryptosporidium can cause
diarrhoea, nausea, etc.
Table E8 The sources and environmental effects of nitrates, phosphates and organic matter in water.
Waste water treatment
Water from towns and cities is usually treated before it is returned to
rivers. Treatment involves three main stages.
Primary treatment
At a sewage works primary treatment usually involves three stages.
1 Filtration/screening: mechanical screens are used to remove large solid
objects (e.g. sticks, paper, rags). The screens are raked from time to time
and the solid material transported to a landfill site.
2 Settling/grit removal: this occurs in a grit chamber where smaller
particles (e.g. sand, grit) settle due to gravity.
3 Sedimentation: the water is transferred to a primary sedimentation
tank (primary clarifier). Grease and oil float to the surface and can be
removed by skimming. Suspended colloidal particles aggregate together
and sink to the bottom of the tank as sludge.
The process of sedimentation can be improved by adding aluminium
sulfate. Under alkaline conditions (if the water is not suitably alkaline,
extra alkali must be added) this forms insoluble aluminium hydroxide:
Al3+(aq) + 3OH−(aq) → Al(OH)3(s)
As the aluminium hydroxide forms it allows smaller particles to join
together to form larger ones which settle out. This process is called
coagulation and flocculation.
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Secondary treatment
Secondary treatment uses bacteria to remove organic waste. There are
three main methods used: the trickling filter, rotating biological contactors
and the activated sludge process. Only the activated sludge process will be
described here.
Activated sludge process
Water from the primary treatment flows into an aeration tank (Figure
E11) where air or oxygen is pumped in to ensure sufficient oxygen is
present to maintain aerobic conditions for the decomposition of the
organic matter. The water in the aeration tank is mixed with sludge
from the secondary settling tank, which contains a high concentration of
bacteria that can break down organic matter. After several hours the water
passes into the secondary settling tank where the activated sludge settles
out. Some of this sludge is returned to the aeration tank.
secondary
settling tank
water in
water out
aeration tank
sludge
settles out
air
air
air
air
activated sludge
recycled
excess sludge
removed
Figure E11 The activated sludge process.
Tertiary treatment
Tertiary treatment is used in some water treatment plants to remove
suspended solids, dissolved organic compound and dissolved inorganic
substances (nitrates, phosphates and heavy metals).
Filtration can be used to remove suspended solids (e.g. through a bed
of sand). Dissolved organic compounds can be removed by passing the
water through activated carbon (carbon with a very high surface area).
The compounds are adsorbed onto the surface of the carbon.
Removing phosphates
Phosphates can be removed by making them into insoluble precipitates
(e.g. by adding aluminium sulfate).
Al3+(aq) + PO43−(aq) → AlPO4(s)
Removing nitrogen compounds and nitrates
Waste water contains nitrogen in the form of organic nitrogen compounds
or ammonia/ammonium. The ammonium ion can be first converted to
the nitrate(V) ion by nitrifying bacteria:
NH4+(aq) + 2O2(g) → NO3−(aq) + 2H+(aq) + H2O(l)
22
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or
2NH4+ + 3O2 → 2NO2− + 2H2O + 4H+
and
2NO2− + O2 → 2NO3−
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Denitrifying bacteria can then reduce nitrate(V) to nitrogen under
anaerobic conditions. Methanol must be added as a source of carbon.
6NO3−(aq) + 5CH3OH(aq) + 6H+(aq) → 3N2(g) + 5CO2(g) + 13H2O(l)
The nitrogen compounds are thus
removed as harmless nitrogen gas.
Removing heavy metal ions
Adding calcium hydroxide (increases pH) removes some heavy metals as
insoluble precipitates, e.g.
Cr3+(aq) + 3OH−(aq) → Cr(OH)3(s)
Cu2+(aq) + 2OH−(aq) → Cu(OH)2(s)
Not all metals can be removed by increasing the pH. Mercury,
cadmium and lead are removed by bubbling hydrogen sulfide (H2S) gas
through the water to precipitate the insoluble sulfides.
Cd2+(aq) + H2S → CdS(s) + 2H+(aq)
The precipitates must then be removed from the water by sedimentation
and filtration.
Metal ions can also be removed by other techniques such as ion exchange.
Disinfection
Before waste water is released into the environment it may also be
disinfected. Disinfection involves killing microorganisms present in the
waste water. This can be accomplished using chlorine, UV light or ozone.
Obtaining fresh water from sea water
The process of obtaining fresh water from sea water is called
desalination. Desalination is an expensive process and is generally only
used when sufficient fresh water is not available from other sources (e.g.
in the Middle East). Fresh water obtained by desalination can be used for
human consumption, agriculture or in industry. The two most commonly
used processes for desalination are multistage distillation and reverse
osmosis. The energy requirement of multistage distillation is roughly five
times that of reverse osmosis per cubic metre of water produced.
Multistage distillation
Sea water is heated by passing through the condensers in the flash chamber
(Figure E12). All the chambers are maintained at low pressure so that
water has a lower boiling point (a liquid boils when its vapour pressure
equals atmospheric pressure). When sea water enters the first chamber it is
hotter than its boiling point at that pressure and flashes to steam. The steam
(fresh water) condenses (transferring its heat to the sea water inlet) and is
collected. The remaining brine then passes to the next chamber (which is
at a lower pressure). Although the brine entering this chamber is at a lower
temperature because the pressure in the chamber is lower it once again
flashes to steam. This process is repeated in up to 40 chambers.
Reverse osmosis
Osmosis is where a solvent passes through a membrane from a less
concentrated solution to a more concentrated one.
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sea water in
steam condenses
heater
fresh water out
decreasing pressure
flash chamber
semipermeable
membrane
pressure
brine out
decreasing temperature
Figure E12 Multistage distillation.
large
hydrated
ions
In reverse osmosis (Figure E13), applying high pressure (up to about
80 atm) to sea water forces water to pass through a semipermeable
membrane leaving the salt behind, i.e. the water moves from the more
concentrated salt solution. The membrane only allows small water
molecules to pass and not relatively large hydrated ions, e.g. Na+(aq) and
Cl−(aq).
Some of the advantages and disadvantages of multistage distillation and
reverse osmosis are considered in Table E9.
movement of
water
salt water
fresh water
Figure E13 Reverse osmosis.
Multistage distillation
Reverse osmosis
water must be pre-treated to remove solids and prevent microbial
growth on the membrane – risk of bacterial contamination of
membranes
water is higher purity
no membrane is perfectly impermeable to salt
land area needed for a plant is larger
capital costs of building a plant lower
more expensive energy costs
membrane needs to be replaced periodically
generally less expensive per m3 but costs can depend on local factors
thermal pollution from discharging hot waste brine
into the environment
waste brine discharged into the environment has much higher salt
concentration
higher volume of sea water required per m3 of fresh
water produced
Table E9 Advantages and disadvantages of multstage distillation and reverse osmosis.
Test yourself
14 Draw the structure of two PCBs that contain 4 Cl atoms.
A saturated solution is one in
which the maximum amount of
the substance is dissolved at that
temperature.
HL
Solubility product constant
Even substances that we regard as being insoluble in water are soluble to
a certain extent. In a saturated solution of a salt, an equilibrium will exist
between the dissolved and the undissolved salt, e.g.
BaSO4(s)
Ba2+(aq) + SO42−(aq)
This is a heterogeneous equilibrium.
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An equilibrium constant, known as the solubility product constant, can HL
be derived for this system as:
Ksp = [Ba2+(aq)][SO42−(aq)]
The concentration of the solid does not appear in the equilibrium
expression as it is essentially constant.
In general, for a salt of formula, MXn dissolving:
MXn(s)
Solubility product constants are
only applicable to sparingly soluble
salts.
Mn+(aq) + nX−(aq)
Ksp = [Mn+(aq)][X−(aq)]n
The units of Ksp can be worked out from the equilibrium
expression as (mol dm−3)x where x is the total number of ions in
the formula unit. Thus x = 2 for AgCl and 3 for Fe(OH)2 and the
units of the solubility product constant for AgCl are mol2 dm−6 and
mol3 dm−9 for Fe(OH)3. The units of the solubility product constant
are often omitted.
The solubility product constant can be worked out from the solubility.
Worked example
Given that the solubility of iron(II) sulfide is 2.5 × 10−9 mol dm−3 at 298 K, calculate the solubility product constant.
The equilibrium that is established is:
FeS(s)
Fe2+(aq) + S2−(aq)
The expression for Ksp is:
Ksp = [Fe2+(aq)][S2−(aq)]
Since the solubility of the iron(II) sulfide is 2.5 × 10−9 mol dm−3 the concentration of each ion in solution will be
2.5 × 10−9 mol dm−3 and the solubility product constant will be given by:
Ksp = (2.5 × 10−9) × (2.5 × 10−9) = 6.3 × 10−18 mol2 dm−6
If the product of the concentration of the ions (taking into account the
number of each ion present) is less than the solubility product constant,
the substance will be soluble at that temperature, but if the product is
greater than the solubility product constant some solid must precipitate
out of the solution to bring the value back down to equal the solubility
product constant. The origin of the ions does not matter.
Thus, using FeS as an example: if the product of the concentrations
of Fe2+ and S2− in a solution at 298 K is less than 6.3 × 10−18, all the FeS
will remain in solution. If sufficient Fe2+ or S2− ions are added to raise
the product of the concentrations above 6.3 × 10−18, some FeS must
precipitate out of the solution.
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For a salt AB, if [A2+(aq)][B2−(aq)]
> Ksp then AB must precipitate out
of the solution. If [A2+(aq)][B2−(aq)]
< Ksp then all AB will remain in
solution.
E ENVIRONMENTAL CHEMISTRY
25
HL
The common ion effect
A substance AB will be less soluble in an aqueous solution of A2+ or B2−
ions than in water.
If we consider the equilibrium
AB(s)
A2+(aq) + B2−(aq)
adding A or B ions will shift the position of equilibrium to the left (Le
Chatelier’s principle).
Worked example
Given that the solubility product constant of Ni(OH)2 is 6.5 × 10−18 mol3 dm−9 at 298 K, calculate the solubility of
nickel(II) hydroxide in water and in 0.10 mol dm−3 sodium hydroxide solution.
Solubility in water:
Ni(OH)2(s)
Ni2+(aq) + 2OH−(aq)
If the number of moles of Ni(OH)2 that dissolves is represented by s, the concentration of Ni2+(aq) will be s and
that of OH−(aq) 2s
Ksp = [Ni2+(aq)][OH−(aq)]2
6.5 × 10−18 = s × (2s)2
i.e. 6.5 × 10−18 = 4s3
Solving this equation we get
s = 1.2 × 10−6 mol dm−3
The solubility of nickel(II) hydroxide in water under these conditions is thus 1.2 × 10−6 mol dm−3
This calculation has been simplified by ignoring the hydroxide ions that come from the dissociation of
water. The solubility is, however, sufficiently high that this is a reasonable approximation.
Solubility in 0.10 mol dm−3 NaOH(aq)
Sodium hydroxide is fully ionised so the concentration of OH− ions in solution will be 0.10 mol dm−3. Since the
concentration of hydroxide ions is significantly higher than the solubility of Ni(OH)2 in water we will assume
the concentration of hydroxide ions remains constant at 0.10 mol dm−3. We can substitute this value into the Ksp
expression:
Ksp = [Ni2+(aq)][OH−(aq)]2
6.5 × 10−18 = [Ni2+(aq)][0.10]2
i.e. [Ni2+(aq)] = 6.5 × 10−16 mol dm−3
The concentration of nickel(II) ions is the same as the solubility of nickel(II) hydroxide since each time one
Ni(OH)2 unit dissolves, one Ni2+ ion is formed.
Therefore the solubility of nickel(II) hydroxide in 0.10 mol dm−3 sodium hydroxide is 6.5 × 10−16 mol dm−3. This
is significantly lower than its solubility in pure water. The addition of a common ion has reduced the amount of
Ni(OH)2 that can dissolve at a certain temperature.
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Removal of heavy metal ions and phosphates
from water in water treatment
HL
The common ion effect is used to precipitate heavy metal ions and
phosphates from solution.
Since many heavy metal hydroxides have extremely small Ksp values,
adding hydroxide makes even very low concentrations of these ions
become insoluble, so they precipitate out of solution, e.g.
Cu(OH)2(s)
Cu2+(aq) + 2OH−(aq)
Adding hydroxide ions causes the position of equilibrium to shift to
the left. Hydroxide ions must be added until the product [Cu2+(aq)]
[OH−(aq)]2 is greater than Ksp.
Similarly, bubbling hydrogen sulfide through the water increases the
concentration of sulfide ions and can cause heavy metal ions to precipitate
out as sulfides, e.g.:
CdS(s)
Cd2+(aq) + S2−(aq)
The position of equilibrium shifts to the left.
Phosphate ions PO43− ions can be removed from water by adding
aluminium sulfate. This precipitates aluminium phosphate (AlPO4) which
has a Ksp of 1.4 × 10−21 mol2 dm−6.
AlPO4(s)
Al3+(aq) + PO43−(aq)
In a later stage of the water treatment process, a coagulant is added to
facilitate the formation of a sludge containing the heavy metals and other
insoluble substances. The sludge settles out, is separated, dried and disposed
of in landfill sites. Since it is so insoluble it does not present significant
toxic issues to the environment.
Worked example
a A body of water contains a cadmium concentration of 1.2 × 10−15 mol dm−3. Hydrogen sulfide is bubbled into
the water to raise the concentration of sulfide ions to 5.6 × 10−15 mol dm−3. Given that the solubility product
constant for CdS is 1.6 × 10−28 mol2 dm−6 at 298 K, determine whether any cadmium sulfate will precipitate out.
Ksp = [Cd2+(aq)][S2−(aq)]
If we work out the product [Cd2+(aq)][S2−(aq)] we get:
(1.2 × 10−15) × (5.6 × 10−15) = 6.7 × 10−30 mol2 dm−6
This value is less than the solubility product constant and so all the cadmium sulfide will be soluble and none will
precipitate.
b More hydrogen sulfide is bubbled into the water so that the concentration of sulfide ions is increased to
1.0 × 10−9 mol dm−3. Determine the mass of cadmium sulfide that will be precipitated from 1.0 × 106 dm3 of water.
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The concentration of sulfide ions is significantly higher than the concentration of cadmium ions in the water so
we will assume that any changes in the concentration of sulfide ions due to CdS dissolving will be negligible so the
concentration of sulfide ions in the water will remain at 1.0 × 10−9 mol dm−3. We can then substitute this value into
the Ksp expression:
Ksp = [Cd2+(aq)][S2−(aq)]
1.6 × 10−28 = [Cd2+(aq)] × 1.0 × 10−9
[Cd2+(aq)] = 1.6 × 10−19 mol dm−3
The concentration of cadmium ions is thus reduced from 1.2 × 10−15 mol dm−3 to 1.6 × 10−19 mol dm−3. In order
for this to happen cadmium sulfide must precipitate out of the solution. The amount of cadmium sulfide that must
precipitate out is given by 1.2 × 10−15 − 1.6 × 10−19, i.e. 1.2 × 10−15 mol dm−3 to 2 significant figures.
If the volume of water is 1.0 × 106 dm3 the number of moles of cadmium sulfide that must precipitate out is given
by 1.0 × 106 × 1.2 × 10−15= 1.2 × 10−9 mol
144.46 is the Mr of CdS
The mass of CdS is given by 1.2 × 10−9 × 144.46, i.e. 1.7 × 10−7 g
Test yourself
15 Calculate Ksp for each of the following.
Compound
Solubility / mol dm−3
AgCl
1.3 × 10−5
Fe(OH)2
5.8 × 10−6
Fe(OH)3
9.3 × 10−11
16 Given the solubility product constants in the table
calculate the solubility of each substance in water
at 298 K.
Compound
Ksp
PbSO4
1.6 × 10−8
Ag2S
6.3 × 10−51
Ag3PO4
1.8 × 10−18
Bi2S3
1.0 × 10−97
18 Ksp of Co(OH)2 is 2.5 × 10−16 mol dm−3 at 25 °C;
10.0 dm3 of water is known to contain Co2+ ions
at a concentration of 1.2 × 10−7 mol dm−3. Solid
sodium hydroxide is added gradually to raise the
pH in stages from 7 to 8 then from 8 to 9, from
9 to 10 and from 10 to 11. Determine at which
stage Co(OH)2 will begin to precipitate out of
the water.
19 The solubility product constant of aluminium
phosphate is 9.8 × 10−21 mol2 dm−6 at 298 K.
Given that the concentration of phosphate ions
in 1.0 dm3 of water is 1.2 × 10−11 mol dm−3,
calculate the mass of aluminium phosphate that
precipitates when sufficient solid aluminium
sulfate is added to the water to increase
the concentration of aluminium ions to
1.0 × 10−5 mol dm−3.
17 Calculate the solubility of each of the following
in 0.10 mol dm−3 sodium hydroxide solution.
Compound
28
Ksp
Mn(OH)2
2.0 × 10−13
Cr(OH)3
1.0 × 10−33
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E7 Soil
Soil is the top layer of the Earth’s continental crust in which plants grow.
It is therefore the part of the Earth’s crust that allows the production
of most of the food required by living things. Soils are very varied and
complex but they all contain:
• inorganic material – rock fragments, clay particles, soluble nutrients,
water, air
• organic material – decaying material and living organisms.
Soils are porous and contain air spaces. Soils are usually divided into layers
called horizons, the uppermost of which is called the topsoil and contains
most of the soil organic matter (SOM).
Soil degradation
Soil degradation refers to the physical loss of topsoil (i.e. the soil is
moved from one place to another by the action of wind or water) and
to the reduction in quality of topsoil. Here we will look at some of the
factors that affect the quality of the topsoil.
Salinisation
This is the process that causes soils to contain high levels of water-soluble
salts. Ions that are present in increased concentrations are Na+, K+, Mg2+,
Ca2+ and SO42−.
One of the causes of soil salinisation is excessive irrigation. Water used
for irrigation contains dissolved salts and, on evaporation of the water,
these are left behind in the soil. If the soil is poorly drained and/or there is
low rainfall in that area, these salts accumulate.
Plants cannot grow well in salty soil because the excess salt reduces
their ability to take up water.
Soil salinisation can be treated by applying extra water, containing
very low levels of dissolved salts, to the soil, although this must be used
with caution as it can cause high levels of salts in groundwater. Another
approach to the problem is to switch the land use to growing crops, such
as barley or cotton, that have a higher tolerance to salt.
Plants will die when the concentration of salts in the soil is higher
than in the root cells as osmosis will cause the water to move from
the plant to the soil.
Learning objectives
•
•
•
•
•
•
•
•
Discuss causes of soil
degradation
Outline the functions of soil
organic matter (SOM)
Describe the importance
of SOM in preventing soil
degradation
Describe the physical and
biological functions of SOM
List common organic soil
pollutants and their sources
Explain what is meant by
HL
cation-exchange capacity
Understand the effect of soil pH
on cation-exchange capacity and
the availability of nutrients
Describe the chemical functions
of SOM
Salinisation tends to be a large
problem in arid and semiarid regions, where the rate of
evaporation of water is high, the
rainfall is low and irrigation must
be used to provide sufficient water.
The levels of dissolved salts present
in soil can be measured by mixing
the soil with water and measuring
the electrical conductivity of the
solution.
Nutrient depletion
When plants grow they remove nutrients and minerals from the soil. If
the plants die where they grow these nutrients are returned to the soil.
However, harvesting crops removes the nutrients and, if they are not
replaced, the soil becomes less fertile and the crop yields decrease. One
way of dealing with the problem is to add artificial fertilisers to the soil
to replace nutrients and minerals. Excessive use of artificial fertilisers can,
however, lead to other problems such as eutrophication, as discussed on
page 18. Crop rotation may also be used to reduce the problem.
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The soil food web is the complex
system of interdependent
organisms that live in the soil.
These range from plants and
bacteria to nematodes, arthropods
and small animals.
Soil pollution
Soil pollution (soil contamination) can arise from many sources (e.g.
overuse of fertiliser, pesticides, chemical dumping, chemical spillages,
road run-off, leaching from landfill sites). Fertilisers and pesticides (these
can kill beneficial organisms as well as harmful ones) can disrupt the soil
food web reducing the number of species present in the soil (reduced
biodiversity). The soil food web is essential for healthy, fertile soil and if
this is disrupted the soil can ultimately be ruined.
Excessive use of pesticides and fertilisers can also have other
consequences as these chemicals can run off the soil into rivers and lakes
and also move through the soil and pollute groundwater.
There are many ways to treat contaminated soil. These include digging
it all up and dumping in a landfill site, incineration, bioremediation (use of
plants, trees, microbes).
Soil organic matter (SOM)
Soil organic matter (SOM) is the organic part of soil and is vital for a
healthy, fertile soil.
The SOM includes undecayed plant and animal tissues, their
partial decomposition products and soil biomass (living organisms
in the soil, including microorganisms).
Humus constitutes about 35–55%
of the non-living part of SOM.
The decomposition products from organic matter in soil can be classified
as humic and non-humic substances. Non-humic substances are released
directly from cells and include identifiable, high-molecular-mass
molecules such as proteins and polysaccharides (e.g. starch) and simpler
substances such as amino acids and sugars. Humus or humic substances are
formed when organic matter is broken down by a series of different soil
organisms.
The substances present in humus are listed in Table E10.
Component
Solubility
Colour
fulvic acids
soluble in alkali and does not form a
precipitate when a strong acid is added
yellow, yellow-brown
humic acids
soluble in alkali but forms a precipitate
when a strong acid is added
dark brown, black
humin
insoluble
black
Table E10 Substances present in humus.
Fulvic and humic acids are not single substances but complex
mixtures of large molecules. They act as weak acids and are important in
maintaining soil pH.
The functions of SOM
The functions of SOM can be divided into physical, biological and
chemical roles. Here we will discuss the biological and physical functions.
The chemical functions will be discussed in the Higher Level section on
page 35.
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Physical
SOM improves soil water retention, alters soil thermal properties
and improves the structural stability of the soil.
The SOM allows the soil to retain more water. This protects the soil
against erosion.
Soil particles come together to form aggregates (peds) with pores
between them and it is these that give rise to the structure of soil. SOM
is essential for the formation of these aggregates. Soil structural stability
is a measure of how soil is able to maintain its structure when exposed to
various stresses such as cultivation and irrigation. SOM helps to maintain
the stability of the aggregates that are responsible for the soil structure.
Generally, the higher the level of SOM, the darker the colour of the
soil will be. Darker coloured soil absorbs more energy from the Sun. This
does not, however, always make the soil warmer as a soil with a higher
SOM content will also generally contain more water and water has a high
specific heat capacity, therefore more energy is needed to heat it up. The
thermal properties of the soil are thus influenced by the presence of SOM
but the exact nature of the effect may be difficult to predict.
Biological
SOM provides a source of energy and nutrients and contributes to
the resilience of the soil system.
An important biological function of SOM is to provide a source of energy
for the metabolic processes that occur in the soil. Plants use photosynthesis
to convert carbon dioxide and water into complex organic molecules and
when the plants die these substances are used by other soil organisms (e.g.
microorganisms) as a source of energy.
SOM is an important source of nutrients such as nitrogen, phosphorus
and sulfur that are essential for plant growth. When organic matter is
broken down in the soil by microorganisms some of the compounds are
converted into forms that can be used by plants and released into the soil.
Soil resilience is a measure of how well the soil system is able to return
to its initial state after some disturbance (e.g. adding chemicals to the soil),
i.e. how stable the soil is. Generally, soils which contain a more diverse
system of microorganisms are more resilient and resistant to disturbance.
Organic soil pollutants
Some common organic substances that can cause pollution of soil are
given in Table E11.
Anthracene (Figure E14) is an example of a polyaromatic hydrocarbon
(PAH).
Figure E14 Anthracene.
Cation-exchange capacity
Cation-exchange capacity (CEC) is a measure of the amount of
exchangeable cations (positively charged ions) in soil. The larger the CEC
the more positively charged cations that can be made available by the
soil. These cations (e.g. K+, Mg2+, Ca2+) are essential for plant health and
therefore CEC is a good measure of soil fertility. Both clay minerals in soil
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Soil pollutant
Source
petroleum hydrocarbons
spilt fuel (including from routine transport uses), leakage from
underground storage tanks, inappropriate disposal of petroleum products
agrichemicals (agrochemicals) – chemical products pesticides, fungicides, herbicides
used in agriculture
polyaromatic hydrocarbons (PAHs)
or polycyclic aromatic hydrocarbons*
or polynuclear aromatic hydrocarbons*
forest fires and volcanic eruptions, incomplete combustion of wood and
fossil fuels, petrol and diesel engines, asphalt manufacture
polychlorinated biphenyls (PCBs)
discharged from factories making electrical transformers and capacitors
organotin compounds
pesticides, anti-fouling paint, an additive to PVC
volatile organic compounds (VOCs)
unburnt hydrocarbons from the internal combustion engine, solvents used
in paints, varnishes, etc.
semi-volatile organic compounds (SVOCs) –
organic compounds that are less volatile (boiling
points 240–400 °C) than VOCs
chemical plants manufacturing plastics, pharmaceuticals and pesticides
may be formed by incomplete combustion (larger PAHs count as SVOCs)
plasticisers, flame retardants, pesticides
* Alternative names.
Table E11 Common organic substances that can cause pollution of soil.
HL and soil organic matter are able to exchange cations. Soil organic matter
has a very high CEC and contributes strongly to soil fertility.
Nutrient ions (potassium, calcium and magnesium) are taken up by
plants by exchange of the cations on the clay particles with H+ ions.
[Clay]2− M2+(s) + 2H+(aq)
Clay minerals are an important part
of most soils. Clay particles are very
small (usually less than 2 µm). Clay
minerals consist mainly of silicon,
oxygen and aluminium joined
together into a complex hydrous
aluminium silicate structure. Clays
have a layer structure where layers
are held together by hydrogen
bonding. Some of the silicon
atoms (oxidation number +4) in
clay particles can be replaced by
aluminium (oxidation number +3)
and some aluminium ions (Al3+)
may be replaced by ions such as
Mg2+ and Fe2+ that have a lower
charge. This means that the clay
particles are negatively charged and
can bind cations. These cations are
reasonably weakly bound and can
then be replaced by other cations –
cation exchange.
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H+ [Clay]2− H+(s) + M2+(aq)
Cation-exchange capacity is expressed in terms of milliequivalents
(meq) per 100 g of soil (dry weight).
1 milliequivalent is 1 millimole of unipositive cations, i.e.
1 × 10−3 mol M+.
Thus, if a soil sample has a CEC of 20 meq per 100 g, 100 g can
exchange 20 × 1 × 10−3, i.e. 0.02 mol of K+ ions or 0.01 mol Mg2+ ions.
CEC can also be expressed in terms of centimoles of charge per kg
(cmolc kg−1) – these units are equivalent to milliequivalents per 100 g.
The effect of soil pH on cation-exchange capacity
In acidic soils, with a high concentration of H+ ions, a high proportion of
the cation exchange sites are occupied by H+ ions which limits the ability
of the soil to take up other nutrient ions.
[Clay]2− M2+(s) + 2H+(aq)
H+ [Clay]2− H+(s) + M2+(aq)
At low pH the higher concentration of H+ ions means that the above
equilibrium shifts to the right and fewer metal cations are available in the
soil for take up by plants. The released metal ions can be washed out of
the soil.
The cation-exchange capacity of some clay minerals can also be
reduced at low pH by protonation of exposed surface OH groups. This
reduces the CEC as cations will be repelled.
clay–Al–OH(s) + H+(aq) → clay–Al–OH2+(s)
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At high pH some of the OH groups on the surface of the clay minerals
in the soil ionise, which increases the negative charge of the clay particles
and the ability to bind cations. The CEC thus increases.
HL
Al3+ and H+ ions do not count as
part of the CEC.
clay–Al–OH(s) + OH−(aq) → clay–Al–O−(s) + H2O(l)
In general, the CEC of soil increases with increasing pH – making
more nutrients available and the soil more fertile.
The effect of soil pH on availability of nutrients
The nutrients that will be discussed here are calcium, magnesium, iron,
phosphorus, sulfur, copper and zinc. Iron, zinc and copper are classified as
micronutrients (only required by plants in small amounts), whereas nitrogen
and phosphorus (and potassium) as well as calcium, magnesium and sulfur
are macronutrients. Aluminium will also be discussed although it is harmful
to plants.
The effect of soil pH on nutrient availability is complex. Here we will
discuss some of the factors that affect the availability of nutrients such as
whether a particular ion will be available in solution or whether it will
form part of an insoluble precipitate, which cannot be taken up by plants,
or whether the ion is adsorbed onto various mineral surfaces. pH also
affects CEC (see page 32). The ideal pH for soil is between 6.0 and 6.5 as
most plant nutrients are in their most available state in this range.
Calcium and magnesium
These form insoluble carbonates at higher pHs.
Ca2+(aq) + 2OH−(aq) + CO2(aq) → CaCO3(s) + H2O(l)
At low pH there is also less Ca2+ available as the CEC of the soil is
reduced, so the clay minerals cannot bind the Ca2+ ions as effectively and
they get leached away.
Leaching – washing of nutrients
out of the soil.
Potassium
Potassium is less available at lower pH due to the CEC being reduced. The
potassium ions are leached away and are not as available for uptake by plants.
Aluminium
When present in soil in the soluble form as Al3+(aq) ion, aluminium is
harmful to plants. Al3+(aq) is only present in soil below pH 6 and only
becomes a major problem below about pH 5. Aluminium ions can cause
direct damage to roots and also reduce the availability of other nutrients
by the formation of insoluble precipitates (e.g. aluminium phosphate).
Aluminium ions can also bind to clay mineral particles and so reduce the
CEC for desirable calcium, magnesium and potassium ions. Aluminium
ions bind more strongly to clay minerals than other ions as they have a
higher charge.
More aluminium ions are available at lower pH as some aluminium
minerals (such as kaolinite) tend to dissolve in acidic solution, e.g.
Al2Si2O5(OH)4(s) + 6H+(aq) → 2Al3+(aq) + 2H4SiO4(aq) + H2O(l)
kaolinite
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33
HL
The aluminium ion produced has a high charge density and so
causes polarisation of water. This hydrolysis reaction helps to replace
the H+ ions used up in the dissolution reaction and can cause
further breakdown of minerals.
[Al(H2O)6]3+(aq)
[Al(H2O)5(OH)]2+(aq) + H+(aq)
As the pH is raised aluminium ions will precipitate out of solution as
aluminium hydroxide and therefore aluminium is less available at higher pH.
Al3+(aq) + 3OH−(aq) → Al(OH)3(s)
Iron
Iron may be present in soil as Fe2+ and Fe3+. Iron is essential for plant
growth. Soluble iron is more available at low pH because at high pH
precipitation of the hydroxides occurs:
Fe3+(aq) + 3OH−(aq) → Fe(OH)3(s)
Fe2+(aq) + 2OH−(aq) → Fe(OH)2(s)
Nitrogen
Inorganic forms of nitrogen that are present in soil are ammonium
(NH4+) ions and nitrate(V) (NO3−) ions. At high pHs ammonium ions are
converted to ammonia, which can be lost from the soil in gaseous form:
NH4+(aq) + OH−(aq)
NH3(aq)
NH3(aq) + H2O(l)
NH3(g)
Nitrogen can be taken up by plants in the form of ammonium ions or
nitrate(V) ions. Soil pH can affect the availability of the different forms of
nitrogen by influencing microorganism activity (nitrifying bacteria) that
can convert ammonium ions to nitrate ions. At lower pHs the ability of
nitrifying bacteria to convert ammonium ions to nitrate ions is reduced
and the ratio of ammonium : nitrate will be higher. Different plants have
different preferences for taking up ammonium or nitrate ions. The relative
availability of ammonium or nitrate ions will thus affect different plants in
different ways.
The availability of soil clay minerals to bind ammonium ions will be
reduced at low pH and they may be leached from the soil.
Phosphorus
Depending on the pH the main inorganic forms of phosphorus that are
present in soils and can be used by plants are H2PO4− and HPO42−. These
two species will be in equilibrium and the position of equilibrium will be
shifted to the right as the soil becomes more acidic.
HPO4−(aq) + H+(aq)
H2PO4−(aq)
Phosphorus is most available to plants between about pH 6 and 7. Below
this pH it precipitates as aluminium or iron(III) phosphates. A simplified
version of the reaction that occurs is:
Al3+(aq) + PO43−(aq)
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AlPO4(s)
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The actual reactions that occur are more complex and species such as
Al(OH)2H2PO4 may be formed.
As discussed above, at low pH, clay minerals can develop a positive
charge on the surface and this can allow phosphate ions to bind. This can
also affect the availability of phosphate ions.
At higher pH phospate is less available because it forms a precipitate
with calcium ions. Simplified versions of the reaction that can occur are:
HL
3Ca2+(aq) + 2PO43−(aq) → Ca3(PO4)2(s)
Ca2+(aq) + HPO42−(aq) → CaHPO4(s)
As above, the actual reactions are more complex.
Sulfur
Sulfur is taken up by plants in the form of the sulfate(VI) ion, SO42−.
Sulfate is readily available in the soluble form above about pH 6. As
the pH is lowered and the surface of clay mineral particles can become
positively charged the sulfate ions can be adsorbed onto these clay
minerals and are held better in the soil and are less susceptible to leaching.
Zinc and copper
Zinc and copper are less available above pH 7. This is due to precipitation
(e.g. as the hydroxides):
The effect of soil pH on the
availability of nutrients is
summarised in Table E12.
Cu2+(aq) + 2OH−(aq) → Cu(OH)2(s)
There will be less copper and zinc available at lower pH due to the lower
CEC.
Nutrient
Low pH
Intermediate pH
High pH
calcium
less available due to reduced CEC
available
less available due to precipitation as carbonate
magnesium less available due to reduced CEC
available
less available due to precipitation as carbonate
aluminium
more available due to dissolution
of minerals
less available due to
precipitation
less available due to precipitation as hydroxide
iron
more available due to dissolution
of minerals
less available due to
precipitation
less available due to precipitation as hydroxide
potassium
less available due to reduced CEC
available
available
+
nitrogen
NH4 less available due to reduced
CEC, decrease in the activity of
nitrifying bacteria will change the
ratio of ammonium : nitrate ions
available
less available as given off as gaseous ammonia
phosphorus
less available due to precipitation
of insoluble phosphates (iron/
aluminium) and adsorption onto
mineral surfaces
available
less available due to precipitation of insoluble
phosphates (calcium)
sulfur
less available due to adsorption
onto mineral surfaces
available
available
zinc
less available due to reduced CEC
available
less available due to precipitation as hydroxide
copper
less available due to reduced CEC
available
less available due to precipitation as hydroxide
Table E12 The effect of soil pH on the availability of nutrients.
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HL
Chemical functions of soil organic matter (SOM)
• SOM increases the ability of the soil to act as an acid–base
buffer keeping soil pH fairly stable.
• It helps to bind pesticides, heavy metals and other organic and
inorganic compounds in soil, thus reducing the environmental
effects they might have.
• SOM contributes to the CEC of the soil.
A phenol group is an OH group
attached directly to a benzene ring:
OH
SOM has a high cation-exchange capacity. Thus, although a soil may
only contain a small percentage SOM, this contributes an appreciable
percentage of the total CEC. The cation-exchange capacity arises from the
presence of many acidic functional groups (carboxylic acids and phenols)
in humus. These weakly acidic groups will be partially dissociated, e.g.
RCOOH
RCOO− + H+
The carboxylate groups can then bind cations:
RCOO− + K+
RCOO−K+
The CEC of the SOM is more sensitive to pH than that of clay
minerals. This is because, at low pH the acid groups are mostly in the
undissociated form and unable to bind cations.
SOM is a very important component in buffering soil against changes
in pH. The buffering action arises mostly from the presence of carboxylic
acid and phenol groups in the SOM. The exact nature of the reactions is
complex due to the number of carboxylic acid and phenol groups present.
Both undissociated and dissociated groups will be present in the SOM
and an equilibrium will exist, that can be represented simply as:
RCOOH
RCOO− + H+
If acid is added the soil then this will react with the RCOO− group:
RCOO− + H+ → RCOOH
If base is added to the soil, this will react with the RCOOH group:
RCOOH + OH− → RCOO− + H2O
Test yourself
20 A sample of soil has a CEC
of 50 meq per 100 g. Explain
what this means.
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As long as the there a large number of dissociated and undissociated
groups present in the system the pH will remain fairly constant.
Certain functional groups (carboxylic acid and phenol groups) in the
SOM can act as ligands to metal ions and form stable complexes with
these cations. The stability of these complexes is higher with more highlycharged ions. These stable complexes mean that heavy metal ions and
aluminium are less available to be taken up by plants and are also less likely
to be leached out into ground water.
SOM binds to organic molecules such as herbicides in the soil. The
exact nature of the interactions depends on the molecules involved but
will include van der Waals’ forces and hydrogen bonding. This binding of
organic molecules is important as it holds these molecules in the soil and
does not allow them to be washed through to ground water.
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E8 Waste
Waste is basically anything that is no longer required and should be
disposed of. This could be packaging from a chocolate bar, expanded
polystyrene packaging used to protect a TV in transit, unwanted
byproducts from industrial processes or mining, spent fuel rods from a
nuclear power station.
Methods of waste disposal
Landfill
Learning objectives
•
•
•
•
Compare the various methods
of waste disposal
Describe recycling of metal,
glass, plastic and paper
Describe the different types of
radioactive waste
Compare methods for storage
and disposal of radioactive waste
Landfill sites are used for the disposal of waste by burying it in a large
hole. One of the key challenges when designing a landfill site is to contain
liquid waste (including water that has passed through the landfill due to
rain) and prevent it from contaminating the environment (ground water,
soil, rivers, etc.). This liquid waste is called the leachate.
The landfill consists of a pit with a base layer of compacted clay (low
permeability to liquids) covered with a synthetic impermeable membrane.
There is also a drainage system so that the leachate can be collected
and treated. When the pit is filled the waste is covered with another
impermeable membrane and layers of soil.
Some of the waste dumped in landfill sites is biodegradable (e.g. food
waste, paper, etc.) and some is non-biodegradable (e.g. most plastics).
Biodegradable waste can be broken down by anaerobic bacteria to
produce methane gas and carbon dioxide (both greenhouse gases). This
methane gas can be collected and used as a fuel. In the European Union
there is a move to reduce the amount of biodegradable waste that is
dumped in landfill sites so that landfill sites can be used for mostly nonbiodegradable waste.
Incineration
Incineration, or burning the waste, is an alternative to the use of landfill
sites.Various technologies are employed in incineration plants, including
moving grates, rotating kilns and fluidised bed combustion. All systems
must achieve a minimum temperature of 850 °C and conditions must
be optimised to achieve, as far as possible, complete combustion. The ash
left over after combustion (bottom ash) is buried in a landfill site (it does,
however, have substantially lower volume than the original waste) or can
be used as aggregate for making roads.
Some incinerators have systems for the recovery of recyclable materials
(such as metals) either before or after combustion.
Some of the heat produced in the incinerator can be recovered by heat
exchange between waste gases and water and this can either be used in
community heating projects or to generate electricity (the water must be
converted to steam). One disadvantage, however, of combined heating and
power schemes is that people often do not want these plants in the middle
of their communities.
Combustion of waste can produce polluting substances such as nitrogen
oxides (NOx), SO2, particulates (fly ash), PAHs, acid gases, heavy metals
and dioxins, however, more modern incinerators and stricter government
regulations have significantly reduced the output of these. The emission
of toxic substances, however, still raises health concerns for people living
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The waste gases in modern
incinerators are treated to remove,
as far as possible, pollutants,
including toxic gases and fly ash.
Alkaline scrubbing may be used
to remove acid gases such as HCl
and SO2 and activated carbon may
be used to remove dioxins. Fly ash
can be removed using filters and
electrostatic precipitation.
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37
nearby. Combustion also produces CO2, which is a greenhouse gas.
Fly ash is considered as hazardous waste and, when recovered, must be
disposed of in special landfill sites.
The discussion here focuses on modern incinerators; older plants
may burn waste at lower temperatures, which can produce higher
levels of toxic substances, and have less efficient systems for
removing hazardous substances from flue gases.
Certain European countries, such as Denmark and Switzerland,
incinerate large proportions of their waste but other countries, such as
Spain, Finland and Ireland, use predominantly landfill sites. Some of the
advantages and disadvantages of landfill and incineration are shown in
Table E13.
Method of disposal
Disadvantages
Advantages
Landfill
Unsightly, smelly and noisy
May attract vermin
Takes up large areas of land
Gives off methane and carbon dioxide
(greenhouse gases) when biodegradable waste is
decomposed anaerobically
May contaminate land
Efficient for large volumes of waste
Cheap
The only method of disposing of some waste
(e.g. ash from incineration)
Methane gas can be used as an energy source
Landfill sites may be redeveloped for use by
wildlife or for leisure activities
Incineration
Expensive (high capital costs and running costs)
Can create pollutants (especially dioxins)
Bottom ash must still be disposed of in landfill
sites
Fly ash must be disposed of as hazardous waste
Requires little space
Significantly reduces the physical volume of waste
Bottom ash may be used as construction
aggregate
Can produce energy and/or heating
Table E13 Advantages and disadvantages of landfill and incineration.
Recycling of metal, glass, plastic and paper
Many of the problems discussed in the section above can be avoided/
reduced by effective recycling programs.
Recycling:
• preserves natural resources and raw materials
• uses less energy
• releases less CO2 into the atmosphere (since less fossil fuels are
required than in the production of new materials)
• reduces waste and the need for landfill sites.
However, the recycling processes described below all use energy and
produce emissions and one way to reduce these is to reuse articles as
much as possible.
Metal recycling
Metals that are commonly recycled include iron/steel, aluminium and
copper. The advantages of recycling metals must be considered not just
in terms of the environmental impact and energy consumption of the
process to extract the metal from its ore (smelting) but also in terms of the
mining process.
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Metal recycling basically involves separating the different metals
(e.g. steel is magnetic but aluminium is not, non-ferrous metals may be
separated using eddy currents) and then melting and re-casting them.
Recycling steel can result in an energy saving of up to about 74%,
whereas recycling aluminium only uses about 5% of the energy compared
with extraction from bauxite (aluminium ore).
Metals produced by recycling may not be as pure as those produced
directly from the ores and, for example, recycled aluminium is not suitable
for some applications in the aerospace industry.
Glass recycling
Glass is separated according to colour, broken into smaller pieces then
separated from other materials such as metal, paper, plastic (which are
often present on glass bottles and jars). The glass is crushed, mixed with
sand, sodium carbonate and limestone and melted. It is then remoulded
into bottles and jars. Glass can be recycled continually without any loss of
quality.
Recycling glass uses less energy, produces less CO2 and uses less raw
materials (sand, sodium carbonate, limestone).
Every tonne of recycled steel saves:
• 1.5 tonnes of iron ore
• 0.5 tonnes of coal
• 40% of the water required in
production
• 74% of the energy needed to
make steel from the original
iron ore
• 1.28 tonnes of solid waste
• air emissions may be reduced
by 86%
• water pollution may be reduced
by 76%.
Plastic recycling
Plastics are derived from crude oil, a finite and valuable resource. There
are many different types of plastics and before recycling they must first be
separated from each other. Plastic packaging and containers have recycling
symbols on them, e.g. type 1 is PET (Figure E15) and type 2 is highdensity polythene (HDPE), to aid separation. Infrared spectroscopy may
also be used to identify different types of plastic. Each different type of
plastic is processed separately.
The plastic is shredded and washed then melted, extruded through holes
and chopped into pellets. These pellets can then be re-melted to make
them into new products
Not all plastics are equally easy to recycle. The most commonly
recycled plastics are PET and HDPE as these contain least additives. Some
plastics (e.g. PVC) containing higher proportions of additives may require
more energy to purify them than would be required to make them from
crude oil. Thermosets have cross-linking between polymer chains which
means they cannot be re-melted and reformed; they are often crushed and
used as insulation.
Pyrolysis or cracking may also be used to recycle plastics. Pyrolysis
involves heating plastics in the absence of oxygen to split them up again
into smaller molecules that can be used as a chemical feedstock to make
new plastics or as a fuel. Thermosets can also be processed in this way.
PET, the main plastic used in fizzy-drink bottles, can either be remelted and formed into new bottles or can be hydrolysed to break it
down into its monomers.
Plastic recycling is expensive (more expensive than dumping it in
landfill sites) but it can, however, reduce energy consumption, reduce
emissions of CO2, reduce the need for more landfill sites and conserve
crude oil, which is an extremely valuable natural resource.
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PET
Figure E15 PET packaging symbol.
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39
Paper recycling
The paper is sorted, and water and chemicals added to produce a pulp.
The pulp is filtered then de-inked to remove inks that can affect the
colour of the final product. The pulp can then be converted into new
paper.
The fibres in the paper are damaged by recycling so the grade of paper
gradually degrades with recycling; it cannot be infinitely recycled like glass.
The factors involved in determining whether recycled paper is better
for the environment than producing new (virgin) paper are complex
and many things must be considered. Producing recycled paper does
require quite a lot of energy including transport of waste paper and the
extra cleaning involved but overall, when the energy impact of growing
and transporting trees etc are included it probably uses less energy than
making new paper. Recycling paper does, however, consume more fossil
fuels than producing virgin paper. Overall, recycling paper releases less
greenhouse gases to the environment than burying it in landfill sites and
making new paper. Recycling paper may however be more expensive
than producing new paper.
Recycling paper does, however, have the advantage that fewer trees
must be cut down and there will be less waste in landfill sites.
Radioactive waste
Radioactive waste can be produced in nuclear power stations, research
laboratories, military establishments, industry and hospitals. Numerous
radioisotopes are used or produced in various processes and these have
half-lives that vary enormously. Some such as 131I, which is used as a
tracer in medicine has a half-life of just 8 days, whereas 239Pu, which is
produced in nuclear power stations has a half-life of thousands of years.
Radioactive waste can be divided into different categories: low-level and
high-level radioactive waste. The category ‘intermediate-level radioactive
waste’ is also sometimes used.
The classification of radioactive waste is important to determine
how it can be disposed of and the safety measures that must be used
in its transport and handling. The criteria used for the classification of
radioactive waste are quite complex and low-level waste is divided into
sub-categories (A, B, C and greater than C) depending on the activity.
Generally, low-level waste has lower activity and usually contains isotopes
with shorter half-lives.
Low-level waste
This includes items that have been contaminated with radioactive
material or have been exposed to radioactivity. Examples of items that
may be classified as low-level radioactive waste are gloves, protective
clothing, tools, soil, rubble, carcasses of animals that have been treated with
radioactive materials etc. Low-level waste may be stored on site until it has
decayed to such an extent that it can be disposed of as ordinary waste or
shipped to a central site for disposal. Low-level waste is often just buried
underground (near surface disposal) for example in individual concrete
canisters or in concrete-lined vaults. Some may even just be put into
landfill sites. Low-level waste must be contained underground for up to
500 years depending on its activity.
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High-level waste
This includes spent fuel rods or other materials from the reprocessing
of nuclear fuels. It contains fission products and transuranic (beyond
uranium) elements generated in the reactor core. These have a
high concentration of radioactive isotopes some of which (such as
plutonium-239) in spent fuel will remain hazardous to humans and
other living beings for thousands of years. Spent fuels rods may either be
considered as waste and disposed of directly or reprocessed to extract
usable material from them to be used as more nuclear fuel.
High-level waste is first stored on site at nuclear power plants in storage
pools (cooling ponds). Storage under water is usually for a minimum of
nine months but sometimes the spent fuel rods are stored in this way
for decades. After sufficient cooling the fuel rods may be transported to
a reprocessing plant, left in the pools or transferred to dry storage casks.
These dry casks have very thick walls and are made of steel and concrete.
The dry casks are then stored in concrete bunkers. Most spent fuel
rods from nuclear power stations in the US are left in cooling ponds or
transferred to dry storage awaiting a more permanent method of storage.
Permanent storage of high-level radioactive waste is a major problem and
various solutions have been suggested such as burying the waste deep
underground in stable geological areas. Over thousands of years, however,
it is difficult to predict what processes could occur to cause release of the
radioactive material.
E9 Smog
Reprocessing of nuclear fuel
produces highly radioactive liquid
waste which can be converted to
glass (vitrified) to make storage
easier.
HL
Learning objectives
Photochemical smog
•
Photochemical smog can be seen as a brown haze hanging over cities
such as Los Angeles soon after the morning rush to work. Photochemical
smog can have adverse health effects (such as eye irritation and respiratory
problems), can damage plants and damage materials such as rubber.
•
The conditions needed for the formation of photochemical smog
are the presence of volatile organic compounds, nitrogen oxides,
sunlight and a fairly stationary body of air. Ideal conditions for the
formation of this fairly stationary body of air are a city in a ‘bowlshaped’ valley, lack of wind and a temperature inversion.
•
Explain the conditions
necessary for the formation of
photochemical smog
State the source of primary
pollutants needed for
photochemical smog
Describe the formation
of secondary pollutants in
photochemical smog
Normally the temperature of the troposphere (the layer of the
atmosphere closest to the Earth) decreases with altitude. These are
favourable conditions for dispersal of pollutants formed at lower
levels as the warmer air near the surface of the Earth can rise and
cooler air can sink to take its place. In this way air can circulate and
mix and pollutants are dispersed. A temperature inversion is the
phenomenon of having a warmer layer of air above a cooler layer of
air (Figure E16). This traps the cooler layer of air close to the Earth
and prevents dispersal of pollutants.
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41
layer of warmer air
HL
cooler air trapped
pollutants trapped
Figure E16 Temperature inversion.
Primary and secondary pollutants needed for
photochemical smog
Volatile organic compounds (e.g. unburnt hydrocarbons) and nitrogen
oxides (NOx) are the primary pollutants that give rise to photochemical
smog. These come from internal combustion engines and are given out
through vehicle exhausts.
Nitrogen(II) oxide is formed when nitrogen reacts with oxygen at high
temperatures in an internal combustion engine:
N2(g) + O2(g) → 2NO(g)
NO can be oxidised to NO2 (nitrogen(IV) oxide, nitrogen dioxide) in the
atmosphere:
2NO(g) + O2(g) → 2NO2(g)
NO2 is thus a secondary pollutant.
In strong sunlight, the NO2 can break down to release an oxygen free
radical which can then react with atmospheric oxygen to form ozone.
UV light
Ozone is a secondary pollutant.
While we consider ozone
beneficial to the Earth in the upper
atmosphere, at ground level it
causes headaches, fatigue and can
aggravate respiratory problems. It
can also cause damage to plants.
•NO2 ⎯⎯⎯→ O• + •NO
O2 + O• → O3
Ozone can also react with NO to form more NO2.
O3 + •NO → •NO2 + O2
Hydroxyl radicals are formed in photochemical smog:
O• + H2O → 2HO•
These can react with NO2 to form nitric acid:
HO• + •NO2 → HNO3
Formation of peroxyacetylnitrates (PANs)
Hydroxyl radicals can also react with hydrocarbons:
RCH3 + HO• → RCH2• + H2O
The alkyl radical formed can react with oxygen to form a peroxy radical:
RCH2• + O2 → RCH2OO•
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This peroxy radical can react in a two-step reaction to form an aldehyde:
HL
RCH2OO• + •NO → RCH2O• + •NO2
alkoxy radical
O
Aldehydes are secondary pollutants
and can cause eye irritation and
respiratory problems.
RCH2O• + O2 → RC + HOO•
H
aldehyde
The aldehyde can react further with hydroxyl radicals:
O
O
RC + HO• → RC• + H2O
H
O
O
RC• + O2 → RC
O
O•
The series of reactions above all involve the using up and production of
free radical and so may be regarded as propagation steps in a free radical
chain reaction. Two free radicals can come together in a termination step:
O
RC
O
O
O• + •NO2 → RC
O
O
NO2
a PAN (peroxy acyl nitrate)
If R = CH3 the compound is called peroxyacetylnitrate (PAN). PANs
cause eye irritation and respiratory problems; they also damage plants.
The primary and secondary pollutants present in photochemical smog
are summarised in Table E14.
Substance
Type of pollutant
NO
primary
VOCs
primary
NO2
secondary
O3
secondary
PANs
secondary
aldehydes
secondary
Photochemical smog develops
during the day. The early morning
rush of traffic produces the
primary pollutants required. Then
as the Sun becomes more intense
towards midday the photochemical
reactions to produce the secondary
pollutants occur and the smog
reaches its maximum. As the Sun
goes down the photochemical
reactions stop.
Table E14 Primary and secondary pollutants in photochemical smog.
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43
Exam-style questions
1 a Identify three primary pollutants produced in car engines and explain how each one is produced.
[3]
b Write an equation to show how two of the primary pollutants in car exhaust react together when
passing through a catalytic converter.
[2]
2 Primary air pollutants include nitrogen oxide(s), sulfur dioxide and hydrocarbons.
a Explain the difference between a primary pollutant and a secondary one.
[2]
b Give an example of a secondary pollutant formed from NO.
[1]
3 Suggest one natural source for each of the following primary pollutants.
a nitrogen oxide, NO
[1]
b sulfur dioxide, SO2
[1]
c volatile organic compounds,VOCs
[1]
4 a Rainfall is naturally slightly acidic. Suggest the range of pH which would be considered to be
appropriate to describe the pH of rain as acid rain.
b Identify two polluting substances that dissolve in rain water to make it acidic and write an equation
for each one reacting with water.
[3]
c Suggest two ways in which the emissions of the polluting substances identified in part b could be
decreased.
[2]
d Explain briefly how acidic soil can damage plants and trees.
[2]
e Write an equation for the reaction of acid rain on marble statues or limestone.
[1]
f Write an equation to explain how the addition of calcium oxide to lakes neutralises the effects of
acid rain.
[1]
5 a The greenhouse effect is a natural phenomenon. It occurs as a result of the way the Earth’s atmosphere
interacts with the radiation from and to space.
i The radiation absorbed from space heats the Earth’s surface. What type of radiation is absorbed
from space?
ii The warmed Earth now radiates energy back into space. In what way is this radiation different
from that in part i?
iii Greenhouse gases reduce the effect in part ii. Explain how they can do this.
iv Why may human activity be increasing the effect of the natural greenhouse effect?
b Explain how each of the following statements can both be true?
I Carbon dioxide is the most important greenhouse gas.
II Methane is a much more important greenhouse gas than carbon dioxide.
c Many people think that global warming is at least in part caused by the increasing concentrations of
carbon dioxide in the atmosphere. Name some possible effects of global warming.
44
[1]
E ENVIRONMENTAL CHEMISTRY
[1]
[1]
[2]
[2]
[2]
[2]
CHEMISTRY FOR THE IB DIPLOMA © CAMBRIDGE UNIVERSITY PRESS 2011
6 a The presence of ozone in the upper atmosphere is important for life on Earth.
i Which type of radiation does ozone absorb in the upper atmosphere?
ii Suggest two effects that ozone depletion could have on life on Earth.
[1]
[2]
b Chlorofluorocarbons (CFCs) are one type of pollutant responsible for ozone depletion. Suggest three
properties required for hydrofluorocarbons (HFCs) to be considered as alternatives to the use of CFCs.
7 a The biological oxygen demand (BOD) is a term applied to water. Explain what this term means and
explain the difference between the BOD of pure water and water containing organic waste.
b The concentration of dissolved oxygen in river water can depend on local activities like farming and
industry. Describe and explain what happens to the concentration of dissolved oxygen in river water:
i when farmers use large quantities of fertilisers on the fields nearby.
ii when a local factory, which uses water for cooling an industrial process, releases warm water into
the river.
[2]
[3]
[2]
[2]
c The concentration of dissolved oxygen in water can be measured by the Winkler method which
involves redox reactions.
100 cm3 of water taken from a river was analysed using this method. The reactions taking place are
summarized below.
Step 1
Step 2
Step 3
2Mn2+(aq) + 4OH−(aq) + O2(aq) → 2MnO2(s) + 2H2O(l)
MnO2(s) + 2I−(aq) + 4H+(aq) → Mn2+(aq) + I2(aq) + 2H2O(l)
2S2O32−(aq) + I2(aq) → S4O62−(aq) + 2I−(aq)
i Is the O2 oxidised or reduced in step 1? Explain your answer.
ii State the change in oxidation number for manganese in step 2.
iii If 0.0002 moles of S2O32− were used in the titration in step 3, calculate the amount, in moles,
of oxygen, O2, dissolved in the original 100 cm3 of water.
8 a State two different methods by which sea water can be converted into fresh water. Explain the
essential features of each of the methods.
[2]
[1]
[2]
[3]
b The majority of the water on Earth is sea water but there is much demand for fresh water. Suggest
three main uses for fresh water.
[2]
9 a Discarded electrical equipment can pollute groundwater with substances known as PCBs which can
cause harmful medical conditions in humans including cancer and damage to the immune and
reproductive systems. What do the letters PCB stand for?
[1]
b State two possible sources of mercury compounds in water.
[2]
10 Water treatment involves three stages: primary, secondary and tertiary.
a What gets removed from the water during the primary stage?
[1]
b In the secondary stage oxygen and bacteria are used. What is removed at this stage?
[1]
c Heavy metal ions are removed in the tertiary stage. How is this done and what substances are used
in this stage?
[2]
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E ENVIRONMENTAL CHEMISTRY
45
11 a The amount of soil organic matter (SOM) is important to the structure and function of a productive soil.
It is a dark-coloured complex mixture often referred to as humus.
i What is the source of the SOM in soil?
[2]
ii List two functions of SOM in soil.
[2]
b Two problems that affect the composition of the soil and that lead to soil degradation are salinisation and
nutrient depletion.
i Explain what causes salinisation and what are its main effects.
ii Explain what causes nutrient depletion and how it can be overcome.
12 a List two advantages and two disadvantages of landfill as a means of disposal of waste materials.
[2]
[2]
[4]
b There are two types of nuclear waste: high level and low level. Explain this difference and suggest
how the problem of disposing of each type is dealt with.
[6]
c Paper, glass and steel are all recycled successfully for re-use but they are all made from readily available
raw materials.
i Why is it better to recycle these materials rather than make them from new raw materials?
ii Explain the main steps in recycling paper.
iii Why is it so much more difficult to successfully recycle plastics?
[2]
[2]
[2]
HL 13 a Identify two primary pollutants in photochemical smog.
[1]
b Ozone is a secondary pollutant in photochemical smog. Explain with the aid of equations the role
of NO2 in ozone formation.
[3]
c Explain how photochemical smog develops in some cities during the day.
[3]
14 Lead ions, Pb2+, can be precipitated from polluted water by treating the water with hydrogen sulfide, H2S.
The solubility product constant (Ksp) of lead(II) sulfide at 298 K is 1.25 × 10−28.
a Define the term solubility product constant by referring to lead(II) sulfide.
[1]
b Calculate the concentration of sulfide ions in a saturated solution of lead(II) sulfide.
[2]
c State what is meant by the common ion effect and explain how it can be used to lower the
concentration of lead ions in a saturated solution of lead(II) sulfide.
[3]
15 An important feature of any soil is its cation-exchange capacity (CEC).
46
a Explain how clay minerals contribute to the CEC of soil.
[2]
b Explain how the CEC is affected by the pH of the soil.
[4]
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