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Chemistry Notes
HSC Core Topic 3
Chemical Monitoring and Management
1. Much of the work of chemists involves monitoring the
reactants and products of reactions and managing reaction
conditions
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Outline the role of a chemist in a named industry or enterprise, identifying the
branch of chemistry undertaken by the chemist and explaining a chemical
principle that the chemist uses
Case study example (from CCHSC)
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A plant chemist in chemical manufacturing company that makes ethylene
(ethene) from ethane, and then polymerises it or sells it for further processing
Branch of chemistry: analytical chemistry
Monitors:
o Output (ethylene, propylene) quality (impurities)
o Waste water for environmental standards (pH, suspensions, sulfate,
grease, hydrocarbons)
o Working with cracking (ethane -> ethene) to adjust operating conditions
for optimal product yields
o Correct operation of equipment (periodic recalibration)
Mostly uses gas chromatography:
o Mixture vaporised into stream of helium flowing through tube over
stationary phase (solid or liquid)
o Gas-solid chromatography: different components of the mixture
accumulate on the solid, and thus pass through the column at different
rates
o Gas-liquid chromatography: components dissolve into liquid at different
extents. The more soluble a substance is, the slower it moves through
the column.
o A device at the end of the column detects each substance as it exits and
thus can measure its concentration quantitatively when compared to
standards
Describe an example of a chemical reaction such as combustion, where
reactants form different products under different conditions and thus would
need monitoring
Reaction condition effects
Incomplete combustion occurs if there is not enough oxygen gas available for reaction. Rather
than producing CO2, it produces poisonous CO and solid C (soot). It also does not release as
much energy as wanted from the complete combustion reaction. Therefore it is necessary for
the reaction to be monitored in order to ensure that there is an excess of oxygen so that
complete combustion occurs:
Complete combustion:
C8H16(l) + 14O2(g) --> 8H2O(l) + 8CO2(g)
Incomplete combustion:
C8H16(l) + 6O2(g) --> 8H2O(l) + 4CO(g) + 4C(s)
Another example is the reaction of ethene with oxygen. In a plentiful supply of oxygen at a
high temperature, ethene reacts to form carbon dioxide (normal combustion).
C2H4(g) + 3O2(g) - 2CO2(g) + 2H2O(g)
However at lower temperatures and with lesser amounts of oxygen and with suitable catalysts,
quite different products are formed. It is important to monitor reaction conditions to ensure
that the yield of the desired product is maximised.

Gather, process and present information from secondary sources about the
work of practising scientists identifying: the variety of chemical occupations,
and a specific chemical occupation for more detailed study
Environmental chemist: are employed by the Environmental Protection Authority (EPA) as
well as by mining companies, industries and local government to collect, analyse, and assess
environmental data. Water resource authorities and air quality management authorities
employ environmental chemists who monitor water and air samples for pollutants.
Metallurgical chemist: are scientists who specialise in the properties, applications and
development of metals and alloys in out technological society, They give advice on the
extraction of metals from ores and ways in which they could be combined with other materials
such as polymers or ceramics.
Biochemist: studies the chemical structure and functions of molecules (eg. carbohydrates,
fats, proteins, nucleic acids) in living things. Through their research new medical, industrial
and agricultural products are developed.
Polymer chemist: investigate the properties of large polymeric molecules. They manipulate
their structure to alter their properties in order to produce new and useful plastic products and
materials.
Industrial chemist: study the structure and chemical reactions of materials that can be used
in industry. Their research and development programs lead to the production of a wide variety
of commercial products ranging from petrochemicals, detergents, and plastics to
semiconductors. Industrial chemists may be specialists in analytical chemistry, organic
chemistry or inorganic chemistry. An industrial chemist works in a team. Research ideas are
developed into experimental procedures on a small scale. Consideration in normally given to
factors such as reaction speed, use of catalysts and the position of the reaction equilibrium.
With the aid of chemical engineers, the procedures are scaled up to the industrial production
levels. Chemical monitoring procedures, quality control testing and environmental monitoring
procedures are designed. An industrial chemist works closely with environmental officers and
marketing and management personnel. Planning and organisational skills as well as analytical
and laboratory skills are vital for in industrial chemist.
2.
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Chemical processes in industry require monitoring and
management to maximise production and ensure quality
control
Identify and describe the industrial uses of ammonia
Ammonia ranks second to sulfuric acid in terms of quantity provided worldwide per year. It is
used to make:
 Fertilisers (sulfate of ammonia, ammonium nitrate, urea)
 Fibres and plastics (rayon, acrylics, nylon)
 Nitric acid, which in turn is used to make fertiliser (ammonium nitrate), dyes, fibres and
plastics, and explosives such as ammonium nitrate, TNT (trinitrotoluene) and nitroglycerine (in dynamite)
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Household cleaners
Detergents (non-ionic ones).
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Gather and process information from secondary sources to describe the
conditions under which Haber developed the industrial synthesis of ammonia
and evaluate the significance at that time in world history
By the beginning of the twentieth century there was a growing need for an industrial method
of synthesising ammonia. Increasing demands for nitrogenous fertiliser to grow food for
increasing world populations were placing strains on the supply of naturally-occurring Chile
saltpetre (sodium nitrate), the main ‘artificial’ fertiliser at that time. In addition the growing
militancy of Germany was promoting calls for more explosives (generally made from nitric acid
which in turn was mainly produced from saltpetre).
The German, Fritz Haber, in 1908 first developed a method of synthesising ammonia from its
elements, through it was not until 1914 that Carl Bosch successfully converted it into an
industrial process. This synthesis contributed significantly to the German war effort in World
War I in that it insulated German agriculture from any harm caused by interrupted supplies of
saltpetre from South America and it facilitated the production of nitric acid and hence of
explosives.
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Identify that ammonia can be synthesised from its component gases, nitrogen
and hydrogen
Describe the synthesis of ammonia as a reversible reaction that will reach
equilibrium
Identify the reaction of hydrogen with nitrogen as exothermic
Explain why the rate of reaction is increased by higher temperatures
Explain why the yield of product in the Haber process is reduced at higher
temperatures using Le Chatelier’s principle
Explain why the Haber process is based on a delicate balancing act involving
reaction energy, reaction rate and equilibrium
Explain that the use of a catalyst will lower the reaction temperature required
and identify the catalyst(s) used in the Haber process
Analyse the impact of increased pressure on the system involved in the Haber
process
The synthesis of ammonia uses the simple exothermic reaction:
N2(g) + 3H2(g) -> 2NH3(g)
H = -92 kJ/mol
This is an equilibrium reaction which at ordinary pressures and temperatures lies well to the
left.
Equilibrium Considerations
Le Chatelier’s principle shows us how to maximise the conversion of nitrogen and hydrogen to
ammonia.
1. If the pressure on a reaction system is increased, the equilibrium moves in the direction
which tends to reduce pressure; that is, the direction which corresponds to a decrease
in the number of moles of gas, since a decrease in the number of moles of gas in a
container of fixed volume leads to a decrease in pressure.
If the pressure on equilibrium mixture of N2 and H2 is increased, the reaction moves to
the right: some N2 and H2 react to form NH3. This happens because that direction
corresponds to a decrease in the number of moles of gas:
4 moles of gas - 2 moles of gas
2. If the temperature is lowered the equilibrium will move in the direction which tends to
increase temperature (release heat). It is exothermic, so if temperature is lowered, it
will move towards the right (form more ammonia).
The percentage conversion of nitrogen to ammonia varies with pressure (at constant
temperature) and temperature (at constant pressure). On equilibrium considerations
alone the reaction should be conducted at high pressure and low temperature.
Rate Considerations
However another consideration is how long will it take for the reaction to reach equilibrium:
that is we need to consider the rate of the reaction. As for most reactions, the rate of the
reaction decreases as temperature decreases. If we lower the temperature in order to move
the equilibrium towards the right, we make the reaction very slow and so it takes a very long
time to reach equilibrium.
One way to increase the rate of reaction is to find a suitable catalyst. Iron is a good catalyst.
While this catalyst does speed up the reaction, the rate is still too slow at room temperature to
be practical. Remember that while a catalyst speeds up a reaction, it does not affect the
position of equilibrium. This is because it speeds up both the forward and reverse reactions.
Hence we have these situations: a low temperature produces a high yield (say 90% conversion
of hydrogen to ammonia), but a very long time (weeks to months) is required to reach
equilibrium, even with a catalyst. A high temperature causes equilibrium to be reach more
quickly (in a few minutes) but the equilibrium yield is extremely low (say 0.1%).
Compromise
Compromise conditions are therefore used. A moderate temperature produces a moderate
yield moderately quickly. Typical conditions for the industrial process, called the Haber
process are:
 A temperature of about 700 K (or about 400C) and
 A total pressure of about 250 atmospheres.
With a reactant mixture having H2 and N2 in the ratio of 3 : 1, these conditions give an
equilibrium conversion to ammonia of about 40%. The catalyst is magnetite, Fe 3O4, with its
surface layer reduced to free iron.
(Re: pic CCHC 193). Reactants pass through the catalyst reactor, then the mixture is cooled to
condense out the ammonia formed: this can be drained off as required. Unreacted gases are
fed back into the catalyst chamber along with incoming fresh reactants. None of the reactant is
wasted. We essentially drive the reaction to a completion by condensing out the reaction
product.
A stoichiometric mixture of hydrogen and nitrogen is used because, as ammonia is formed and
condensed out, left-over reactants can be recycled through the process (with some fresh
reactant mixture added) without any build-up of one reactant over the other.
An important factor in designing an industrial process is energy management. In the Haber
process we would like to use the head released to heat up, at least partially, the incoming
reactants and so minimise our energy costs: this has the added advantage that it stops the
catalyst overheating and so losing activity. This is the reason for the incoming reactants
flowing over the outside of the catalyst chamber entering it.
The Source of Reactants
Nitrogen can be obtained from the atmosphere, so hydrogen is the ‘difficult’ or expensive
reactant to obtain.
In the laboratory we make hydrogen gas by reacting zinc of magnesium with hydrochloric acid.
This process is too expensive for industrial use. Alternatively hydrogen can be made by the
electrolysis of water, but except in special situations this is also very expensive. Industrially,
hydrogen is generally produced by reacting methane (natural gas) or some other hydrocarbon
with steam in the presence of a nickel catalyst at the temperature of about 750C:
CH4(g) + H2O(g) -> CO(g) + 3H2(g)
H = +206 k/J mol
Carbon monoxide poisons the iron catalyst in the Haber process and so must be removed. This
is done with other catalytic reaction:
CO(g) + H2O(g) -> CO2(g) + H2(g)
Which has the added advantage of producing more hydrogen. The catalyst used is either Fe 3O4
at 500C or Cu at 250C.
A complication for the Haber process is that we want a mixture of hydrogen and nitrogen that
does not contain any oxygen (which can react explosively with hydrogen under the conditions
used). Methane us also used to remove oxygen from air (in effect just normal combustion to
CO2 and steam). By adjusting the quantities of methane, steam and air used, this combination
of reactions can be made to product a 3 : 1 mixture of hydrogen and nitrogen. The only
unwanted gas in the mixture is carbon dioxide. This is removed by reaction with a base, if
oxygen has to be excluded and if the reactant mixture has to be recycled through the reactor,
the use of a 3 : 1 mixture is the most efficient way of making ammonia.
Re. picture CCHSC p195
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Explain why monitoring of the reaction vessel used in the Haber process is
crucial and discuss the monitoring required
Because many different conditions must be maintained for efficient and safe operation of the
Haber process, monitoring is essential. First temperature and total pressure must be
monitored to keep them in the range for optimum conversion of reactants to products: in
addition excessive temperature can damage the catalyst. Then it is essential to monitor the
composition of the incoming gas stream: we need to ensure that the ratio of H 2 to N2 is kept at
3 : 1 (to avoid a build-up of one reactant), that oxygen is absent (to avoid risk of explosion)
and that concentrations of carbon monoxide and sulfur-containing species are sufficiently low
to prevent poisoning of catalyst. With a well-maintained plant the catalyst can last up to eight
years. Any build-up of unreactive gases such as argon and methane needs to be watched also
because that can lower the efficiency of the conversion.
In Summary:
Ammonia can be synthesised from nitrogen and hydrogen (equilibrium).
N2(g) + 3H2(g) <--> 2NH3(g) ΔH = –92 kJ/mol
The rate of reaction would be increased by higher temperature, since particles move faster,
causing more collisions and reactions.
But since the reaction is exothermic, an increase in temperature will favour the back reaction
and thus reduce yield of NH3.
Since some changes that would usually be made to the set-up to increase reaction rate and
decrease activation energy shift equilibrium left, there needs to be a delicate balancing of
these factors to compromise between yield and reaction rate.
A catalyst in the reaction lowers activation energy and thus less heat is required for the
reaction to occur. Iron (as surface on magnetite, Fe 3O4) is the heterogenous catalyst used (N 2
& H2 break apart and react on its surface).
Increased pressure shifts equilibrium right, in order to gain fewer overall moles of gas,
therefore increasing yield in the Haber process.
The Haber process is thus performed with:
 approximate temperature 700K, pressure 250atm
 catalyst Fe3O4 with Fe surface
 the product constantly being condensed and removed (shifting equilibrium
right)
 left-over reactants are recycled
 incoming reactants heated up by previous reactions by flowing over reaction
chamber (it also stops overheating the catalyst)
Monitoring is required:
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Temperature and pressure monitored to keep in optimum range (and not to
damage the catalyst)
H2:N2 :: 3:1 (to avoid build up of one reactant)
Oxygen not present (explosion)
Other contaminants not present (CO or species with S can poison catalyst)
Avoid build-up of unreactive gases (reduces efficiency of reaction)
Gather and process information from secondary sources to perform
calculations to demonstrate the effect of volume, temperature and
concentration changes on product formation in the Haber process
Quantitative analysis of equilibrium:
for reaction aA + bB <--> cC + dD, K 
Cc Dd  products
Aa Bb reactants
This is known as the equilibrium constant, for a particular reaction at particular temperature.
When concentrations of involved species change, the other concentrations adjust in order for K
to remain constant.
A higher number means more to the right in equilibrium.
Hence subbing in numbers for volume, temperature and concentration on the Ammonia
equilibrium will give us different K values.
3. Manufactured products including food, drugs and household
chemicals, are analysed to determine or ensure their chemical
composition and/or energy content
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Describe chemical tests to identify the following: anions: phosphate, sulfate,
carbonate, chloride, cations: barium, calcium, lead, copper and iron
Perform a first-hand investigation to carry out a range of tests, including flame
tests, or use information from secondary sources to identify the following ions:
phosphate, sulfate, carbonate, chloride, barium, calcium, lead, copper and iron
Tests for cations
Cation
Chemical tests
Flame test colour
2+
–
Pb
Cl precipitate (white)
I– precipitate (yellow)
2+
Ba
SO42– precipitate (white)
Pale green
2+
Ca
SO42– precipitate (white)
Brick-red
F– precipitate (white)
Cu2+
OH– precipitate (blue)
Blue-green
Blue
Precipitate with NH3 dissolves
Fe2+
OH– precipitate
Pale green
MnO4– decolourises
3+
Fe
OH– precipitate (brown)
Yellow
SCN– deep red
2+
2+
Pb and Ca in small concentrations may not precipitate; if no precipitate is found add KI, to
distinguish.
Tests for anions
Anion
Tests (performed with nitrates for multiple)
CO3–
PH test 8 – 11
Add H+ Bubbles form
SO42–
Ba2+ precipitates (in acid)
–
Cl
Ag+ precipitate
PO43–
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Add NH3 and Ba2+ precipitates
Deduce the ions present in a sample from the results of these tests
From the above tests, we can deduce the ions present in a sample.
For multiple ions, tests need to be performed in an order, making sure that salts that will not
react are used (ie. NO3– tests with anions). When precipitates form, the reaction needs to be
driven to completion and then the precipitate filtered (or preferably centrifuged off) and the
filtrate tested.
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Describe the use of atomic absorption spectroscopy (AAS), in detecting
concentrations of metal ions in solutions and assess its impact on scientific
understanding of the effects of trace elements
Atomic absorption spectroscopy (AAS) is a method to determine concentrations of cations
to very fine accuracy (below 1ppm). When light of an element’s emission spectrum is shone on
it, it is absorbed. Each element has a unique set of wavelengths whose energy it will absorb.
The substance is placed in a flame, and lamps shine possible emission spectra onto the
substance (now melted into atoms). The resulting intensity of light is measured and calibrated
against the intensity without the substance (and with known substances), and thus the
concentrations of the ions in the substance can be determined. The major disadvantage is the
cost of equipment and need for different light sources for each spectrum tested. The electrons
of the exposed atom absorb the relevant energy. Since nearly 100% of the atoms absorb, AAS
is extremely sensitive. It is used to find metals in the environment (eg. Mercury in fish),
micronutrients in soils, contaminants in foods, trace elements in organisms. By using AAS, it
was discovered that organisms require some elements in very minute amounts (1-100ppm),
and thus allowed research into their essentiality & effects within organisms.
While titration is accurate to 10-4 molL-1, AAS is accurate to 10-6 molL-1.
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Gather, process and present information to describe and explain evidence for
the need to monitor levels of one of the above ions in substances used in
society
Need to monitor ion levels
Lead ion concentration in air needs to be monitored in order to ensure that people are not
exposed to excessive concentrations. This is because the lead (which is released from some
petrol) is poisonous and retards intellectual development.
Phosphate contents need to be monitored, since excessive amounts in natural waterways can
cause algal blooms. This rapid growth can completely cover lake surfaces and kill natural life.
The algae’s death after phosphate has depleted is also detrimental to the water body. Thus the
amounts of phosphate in consumer goods that ultimately end up in waterways need to be
monitored.
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Identify data, plan, select equipment and perform first-hand investigations to
analyse the: phosphorous content in detergent, nitrogen content of fertiliser,
sulfate content of lawn food or fertiliser, ethanol in wine or beer, and citric
acid in orange juice, and explain the chemistry involved in each case
Analysis of substance content in products
(1) Phosphorus in detergent
 To digest sample (polyphosphates -> phosphates, boil in sulfuric acid)
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Weigh 1g into 250mL conical flask
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Add 25ml 2M H2SO4
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Suspend flask (with loose lid) in beaker of water (with boiling chips) above
Bunsen flame
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Boil for 30 mins
Prepare accurate standards
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From 0 g/L to about 0.40 g/L PO43–
(Make overall standard and add in different proportions to make other
standards)
 Prepare unknown for colour
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Take unknown, rinse the conical flask a few times into a 250mL volumetric
flask
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Fill to graduation mark with distilled water
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Take 5mL by bulb pipette and add to another 250mL volumetric flask
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Fill with H2O to graduation
 Prepare molybdate mixture
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Make up a common solution of ammonium molybdate - ?1M
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Add to it concentrated (5/6M) sulfuric acid
 Add colour reagents – (NH4)6Mo7O24, H2SO4 and C6H8O6
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Accurately add 10mL molybdate mixture, and a few crystals to each sample
of 100mL (conical flask via measuring cylinder)
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Heat flask on tripod and gauze until 95°C (or it starts to boil)
 Compare colours
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Colorimeter with 415 nm filter, measure absorbance (lowest to highest
calibration) and then find unknown on calibration curve
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Line up standards in same-diameter beakers, in order
(2) Nitrogen in fertiliser
 Assume nitrogen is present as NH4+
 Weigh 1g of fertiliser into conical flask and dissolve in warm water
 Add [accurate] 25mL of 0.7M NaOH solution
 Boil away NH3 over Bunsen until fumes cannot be smelt
 This solution now contains excess OH– from NH4+ + OH– -> H2O + NH3
 Allow the solution to cool
 Titrate against 0.1M HCl, with methyl red (this performs back titration)
 Knowing remaining OH– used to find reacted OH– and thus initial NH4+
(3) Sulfate in lawn food / fertiliser
 Prepare BaCl2 Add 5g BaCl2 to 100mL of deionised water (5% w/v)
 Weigh 1g fertiliser into beaker, with deionised water to dissolve
 Add 10 drops concentrated HCl
 Add 50mL barium chloride solution
 Heat beaker (covered with watch glass) at 70°C for 60 mins
 Filter through a weighed Buchner funnel lined with filter paper (collect in conical
flask with side-arm attached to vacuum pump)
 Wash precipitate with warm water, ethanol and ether
 Allow to dry (10 mins) and then desiccate
 Weigh filter when white solid is dry
(4) Ethanol in wine or beer
 Distil
–
Add 100mL of drink to volumetric flask (rinsed with drink)
–
Pour alcohol into distillation flask, with boiling chips
–
Add half a teaspoonful of Ca(OH)2 to react out any volatile acids
–
Set up condenser (water travelling up) in distillation set-up (thermometer
bulb at height of condenser opening)
–
Heat the distillation flask on low flame, distilling into conical flask
–
It should stay at 78°C and stop heating when it nears 100°C
–
Take conical flask with distillate (ethanol) and cool it
 Measure density
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Add distillate to rinsed 100mL volumetric flask and fill
–
Mix flask contents well (shake)
–
Measure temperature, and cool/heat to 20°C
–
Pipette 25mL into weighed beaker
–
Calculate mass of ethanol and water, and therefore density
–
Repeat this a couple of times and average for accuracy
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 Compare to calibrated densities vs %ethanol at 20°C
(5) Citric acid in orange juice
 Find ascorbic acid percentage on product packaging (though it is rather
negligible since citric acid is the major acid present)
 Note that ascorbic acid is monoprotic for our use, and citric acid is triprotic (2hydroxypropane-1,2,3-tricarboxylic acid, H3C6H5O7)
 Measure out 100mL orange juice and filter out solids with Buchner filter lined
with filter paper
 Pipette out 25mL of filtered juice into conical flask
 Add 50mL of deionised water to dilute the colour
 Titrate (with phenolphthalein) against 0.1M NaOH
4. Human activity has caused changes in the composition and the
structure of the atmosphere. Chemists monitor these changes
so that further damage can be limited

Describe the composition and layered structure of the atmosphere
Atmospheric layers defined according to weather trends:
 Troposphere decreases temp from approx. 15°C at 0km above surface to –50°C
at 15km (tropopause)
 Stratosphere from 15km to 0°C at 50km (stratopause)
 Temp decreases from here (mesosphere)
 Air pressure decreases with gravity and thus radius
 Since hot air rises and cool air falls (convection), air mixes very easily
throughout the troposphere, and does not very much in stratosphere. Therefore
troposphere is relatively unstable and contains weather phenomena;
stratosphere does not.
 Because of temperature barrier, very little leaks down from stratosphere into
troposphere, and pollutants very slowly leak up / diffuse into stratosphere
Composition:
 By volume (& mol): nitrogen 78.11%, oxygen 20.96%, argon 0.93%
 Many trace gases and pollutants (ppm concentration), some of which vary from
place to place depending on natural and man-made environments
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Identify the main pollutants found in the atmosphere and their sources
Main air pollutants and
Carbon monoxide (CO)
Ozone (O3)
Haloalkanes
NOx
Hydrocarbons
Lead
SO2
Radioactivity
Carcinogens
sources
Cars, cigarettes, fires
Photochemical smog
Refrigeration, air conditioning, plastics, fire extinguishers,
electronics cleaning
Combustion
Vehicles and factories (solvents)
Leaded petrol, smelters
Combustion, metal extraction, chemical manufacture
Combustion, uranium mining, medical and scientific
radioisotope use, nuclear testing & power plants
unleaded petrol [benzene, toluene], plastics manufacture
[vinyl chloride], cigarette smoke
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Describe ozone as a molecule able to act both as an upper atmosphere UV
radiation shield and a lower atmosphere pollutant
Ozone, O3, is a pollutant in the lower atmosphere, is poisonous to most life forms. Causes
respiratory problems, headaches and premature fatigue. It can be creating by passing an
electric current through O2 gas (3O2 -> 2O3).
It is the most harmful component of photochemical smog. Sunlight splits an O atom off
relatively highly concentrations of NO2 from car exhausts, which combines with O2 to form O3
in the air.
In the upper atmosphere (stratosphere), ozone acts as a UV radiation shield. Oxygen gas
shields out short wavelength UV (< 240nm): O2 + UV -> 2O2
These oxygen atoms react with oxygen gas to form ozone, or react with O3 to form O2,
creating a balance in O2 and O3:
O2 + O -> O3
O3 + O -> 2O2
ΔH = -106 kJ/mol
ΔH = -390 kJ/mol
These exothermic reactions explain the heat increase with altitude in the stratosphere.
Ozone absorbs the longer wavelength UV radiation (200-300nm): O3 -> O + O2, thus
protecting life on earth from the harmful UV.

Describe the formation of a coordinate covalent bond
A coordinate covalent bond is one where both of a pair of shared electrons come from the
one atom. It is formed when an atom in a bond does not have a full outer shell and the other
does and has an unshared electron pair. It therefore shares this pair with the atom and creates
a coordinate covalent bond.
To form a coordinate covalent bond, 1 atom has a pair of non-bonding electrons, and the other
has to have 2 holes.

Demonstrate the formation of coordinate covalent bonds using Lewis electron
dot structures
Ozone

Compare the properties of the oxygen allotropes O2 and O3 and account for
them on the basis of molecular structure and bonding
Allotropes of oxygen
Oxygen gas, O2
 Quite reactive
 Moderate oxidant (quite
explosive)
 To split requires double-bond to
be broken
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ΔH=+498
O 2 ( g )  2O( g )
kJ/mol
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Ozone, O3
Much more reactive (and
unstable)
Powerful oxidant (highly
explosive)
To split requires coordinate
covalent bond to split
O3 ( g )  O 2 ( g )  O( g ) ΔH=+106
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Non-polar
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Colourless gas, pale blue liquid
Odourless
Essential for all living matter
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Boils at –183°C
Sparingly soluble in water

Compare the properties of the gaseous forms of oxygen and the oxygen free
radical
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kJ/mol
Polar
Colourless gas, distinct blue liquid
Strong, distinctive odour
Poisonous (very reactive free
radical O joins onto molecules in
organisms)
Boils at –111°C (150% mass)
More soluble in water
A free radical is an uncharged atom or group of atoms with one or more unpaired electrons.
The oxygen free radical is very reactive, since it has 2 unpaired electrons in 2 outer orbitals.
Because of this high and rapid reactivity, it does not usually exist alone in nature.
Oxygen: 1s2 2s2 2p4
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Explain the effects of ultraviolet radiation on many polymers including those in
living tissue
Gather, process and present information from secondary sources such as
visual models to identify the effect of UV radiation on biological models
Short wavelength UV radiation contains enough energy to break chemical bonds, particularly
double bonds. Many bonds in polymers (biological or not) can be broken by the UV radiation.
In the case of biological molecules (proteins and nucleic acids), their chemical changes caused
by UV creates molecules unable to perform their original function, killing cells in the organism
or causing severe genetic mutation in offspring. To humans they cause increased cancer risk,
damage to the immune system, severe sunburn, and eye cataracts.

Identify the origins of chlorofluorocarbons (CFCs) and halogens in the
atmosphere
CFCs are compounds containing only chlorine, fluorine and carbon. Their (and other
halogenated hydrocarbons, like halons) origins in the atmosphere are from use as refrigerants,
propellants, solvents, and in plastic manufacture.


Identify and systematically name straight chain haloalkanes and alkenes with
up to eight carbon atoms
Identify and name examples of isomers (excluding geometrical and optical) of
haloalkanes up to eight carbon atoms
Systematic naming of haloalkanes & alkenes
(1) Bromo-, chloro-, fluoro- and iodo- prefixes for halogens in hydrocarbons
(2) Number according to position in carbon chain (prioritise by position of double bond
or alphabetical order of halogen names if necessary)
(3) If more than one of a halogen is present, prefix the halogen name by di-, tri-, tetra,… give numbers of positions separated by numbers

Discuss the problems associated with the use of CFCs and assess the
effectiveness of steps taken to alleviate these problems
Problems due to CFC usage
 Do not get destroyed at lower altitudes by sunlight and O2



Insoluble in H2O – do not wash out of atmosphere
Thus remain in troposphere and diffuse into stratosphere
Short wavelength UV in stratosphere can remove chlorine from CFCs
radiation
Eg. CCl 3F u.v.

 Cl  CCl 2 F





Chlorine free radical reacts with ozone:
Cl + O3 -> ClO + O2
Chlorine is released again in chain reaction 2ClO -> Cl2O2 -> 2Cl + O2
Thus one CFC destroys many O3 in stratosphere
Causes decrease in stratospheric ozone concentration
More harmful UV radiation passes through stratosphere, harming life on earth
The only way to stop ozone destruction by CFCs is to stop releasing them to the atmosphere,
because once they are released it is impossible to remove them, and they do not decompose
or get washed out by rain.
International Agreements
Since 1987 there have been several international agreements in which nations have
undertaken to phase out the use of CFCs and other ozone-destroying compounds. The original
agreement was called the Montreal Protocol. The most recent agreement (1992) was to:
 Stop using halons by the end of 1994
 Cease the manufacture and use of CFCs and 1,1,1-tricholoroethane by 1996
 Phase out HCFCs by early in the 21st century
 Allow less developed countries some period of grace and provide some financial
assistance to help them phase out CFCs.

Analyse the information available that indicates changes in atmospheric ozone
concentrations, describe the changes observed and explain how this
information was obtained
Stratospheric ozone is monitored from ground-based instruments, from instruments in a
satellite and from instruments in balloons.
The ground based instruments are u.v. spectrophotometers pointing vertically upwards
through the atmosphere. They measure the intensity of the light received at a wavelength at
which ozone absorbs and then at wavelengths either side of this at which ozone doesn’t
absorb. This gives an indication of the total ozone in the atmosphere in a particular area.
The total ozone mapping spectrophotometers (TOMS) that have been aboard several US
satellites over the past 20 years work similarly, and have been able to scan through the
atmosphere and measure the ozone concentration as a function of altitude and of geographic
position from orbit. They can produce contour maps.
Total ozone per unit are above a location is generally measure in Dobson Units, DU. 1 DU =
0.01 mm thickness of pure ozone at atmospheric pressure. Normal ozone is 300 DU = 3 mm.
Huge helium-filled balloons have been used to carry instruments into the stratosphere to
measure ozone the same way.

Present information from secondary sources to identify alternative chemicals
used to replace CFCs and evaluate the effectiveness of their use as a
replacement for CFCs
Alternatives to CFCs
International agreements (initially the Montreal Protocol) involved many nations in phasing out
the use of CFCs. HCFCs (hydrochlorofluorocarbons) were used as temporary replacements,
whose C-H bond can be broken down in the troposphere, thus causing less to reach the
stratosphere and become destructive to ozone. HFCs (hydrofluorocarbons) lack a C-Cl bond
and thus are not destructive to ozone, but are more expensive and not as efficient as CFCs.
Since CFCs are very slow at diffusing into the stratosphere, they will continue to destroy ozone
for 50-100 years, though restrictions put in place are reducing our emissions and thus will
eventually recover normal atmospheric ozone concentrations.

Gather, process and present information from secondary sources to compare
amounts of components of the atmosphere by moles, mass, by volume and by
part per million

Present information from secondary sources to write the equations to show
the reactions involving CFCs and ozone to demonstrate the removal of ozone
from the atmosphere
In the stratosphere ozone come into contact with short wavelength u.v. radiation (before the
ozone has filtered it out) and this short wavelength u.v. breaks a chlorine atom off the CFC
molecule:
CCl3F + u.v. light - Cl + CCl2F
CCl2F2 + u.v. light - Cl + CClF2
The chlorine atom formed then abstracts an oxygen atom from an ozone molecule, forming a
ClO ‘molecule’:
Cl + O3 - ClO + O2
ClO is a free radical and is very reactive with an unpaired electron. ClO reacts with a free
oxygen atom to form a molecule of oxygen, O2 and this regenerates the chlorine atom:
ClO + O - O2 + Cl
This process continues in a cycle, and is a chain reaction. One chlorine atom can destroy
thousands of ozone molecules.
5. Human activity also impacts on waterways. Chemical
monitoring and management assists in providing safe water
for human use and to protect the habitats of other organisms

Identify and describe tests that can be used to determine the quality of water
samples as: concentrations of common ions, total dissolved solids, hardness,
turbidity, acidity, dissolved oxygen and biological oxygen demand, and
nitrogen-to-phosphorus ratio
Good quality drinking water:
 Is completely colourless and clear
 Is odourless
 Has a pleasant taste
 Has a relative low salt content
 Contains no pathogens (disease causing agents)
 Contains no poisonous chemicals.
The concentration of common ions
AAS is the most common technique used to identify the concentration of metal ions. In testing
water quality the concentration of the following cations are usually determined: sodium,
magnesium, calcium and potassium.
Gravimetric analysis can be used to determine the quantities of both cations and anions, e.g.
for chloride ions precipitated and weighed as silver chloride. Common anions concentrations
measured include chloride, sulfate, bicarbonate and fluoride ions.
A method for the precipitation of barium sulfate from ammonium sulfate fertilizer is described
previously.
Total dissolved solids (TDS)
Total dissolved solids (TDS) are determined by evaporation to dryness of a known volume of a
filtered sample. The value is converted to parts per million (ppm) and expressed in mass per
volume units, ppm (m/v).
Since most of the dissolved solids are ionic, their presence can be determined by data loggers,
which have the electrical conductivity probe attached. They can be set up to record
continuously if needed.
Hardness
Hardness is due to the presence of calcium and magnesium ions in the water. These form
insoluble compounds with soap ions, resulting in a scum on the water surface and around sink
basins. The removal of soap ions from solution reduces the ability of soap to lather.
The test for hardness involves precipitating the calcium and magnesium ions from a known
volume of the water sample with a solution of sodium carbonate (of known concentration),
followed by filtering and drying of the precipitate. Most of the insoluble salt is assumed to be
calcium and the concentration of calcium ions is calculated and reported in parts per million
(ppm).
Turbidity
This can be one measure of the ability of the water to support life. Turbidity results from the
presence of suspended solids in the water. Water with a high turbidity reduces penetration of
light and decreases photosynthesis, which in turn reduces the oxygen concentration.
The test for turbidity is conducted using a turbidity tube standing on a white tile. The tube has
a black cross marked on the base. The water sample is poured into the tube until the cross just
disappears when looking from above.
The use of a turbidity tube is less reliable than a transmittance of light test done with a
colorimeter. The lower reliability is due to variations in human eyesight and the intensity of
background light during testing.
Acidity (pH)
A pH reading below 7 would be expected where there are acid sulfate soils or where there is
acid produced by decomposition of organic matter in stagnant situations. The test can be
conducted with a data logger and pH probe, universal indicator solution or paper, or a pH
meter. If using the universal indicator, comparison with a coloured pH scale provides the pH
value. If it is less than 7, the solution is acidic.
Dissolved oxygen (DO)
There are several tests for determining the DO in a water sample. The Winkler method fixes
the amount of dissolved oxygen, which is later determined by titration.
The amount of manganese dioxide produced by adding manganese(II) ions and hydroxide ions
is a measure of the DO. Acidified iodide ions are added to cause the manganese dioxide to
produce a yellow iodine solution. This is then titrated against a standard sodium thiosulfate
solution using starch as the indicator. The indicator turns a blue colour with the iodine and the
blue disappears at the endpoint. To conduct the test, no air is to be trapped with the sample
and it is to be kept in the dark to reduce algae photosynthesis increasing the DO.
Reaction 1:
Reaction 2:
Reaction 3:
The overall reaction can be represented by the equation:
This shows that, for each mole of thiosulfate (S2O3
moles of dissolved oxygen in the original sample.
) used in the titration, there was 0.25
2–
Biochemical oxygen demand (BOD)
BOD measures the amount of oxygen used by bacteria and other microorganisms during a
five-day period. The sample bottles are held below the water surface and away from the bank.
One sample is measured for DO as soon as possible while the other sample is kept in a dark
place for 5 days and then tested for DO. The BOD is calculated by subtracting the DO value
after 5 days from the initial DO value. The reading is given as milligrams per litre (mg L -1).
Nitrogen-to-phosphorus ratio
The nitrogen:phosphorus ratio refers to the total nitrogen level (inorganic and organic nitrogen
in dissolved soluble form and insoluble solid form but not nitrogen gas) and total phosphorus
level (mostly phosphates, hydrogen phosphate, dihydrogen phosphate and polyphosphate
ions). The test is either via volumetric analysis, gravimetric analysis or instrumentally, by
visible light absorption spectroscopy.
The volumetric technique requires the nitrogen to be converted to ammonium sulfate using
sulfuric acid with the addition of excess alkali to produce ammonia gas, which is collected. This
is neutralised by a known volume and concentration of acid in excess. The excess is titrated
with a base to determine the amount of acid that reacted with the ammonia.
The phosphate determination can be done using a gravimetric technique. Sulfuric and nitric
acid mixture converts the phosphorus to phosphate. Addition of a soluble magnesium salt and
ammonia solution precipitates the phosphate as insoluble magnesium-ammonium phosphate6-water. This is then collected, dried and calculations are done to determine the amount of
phosphorus. Then the ratio of N:P is calculated and compared with accepted standards for
different types of waterways.

Identify factors that affect the concentrations of a range of ions in solution in
natural bodies of water such as rivers and oceans
The ions present in natural water bodies (such as rivers, dams and oceans) and their
concentrations depend upon the following factors:
 The pathway from rain to water body
Rain contains very few ions. When rain runs off bushland into streams, it picks up small
amounts of nitrates and phosphates from natural nutrients on the surface. If, however, rain
water soaks into the ground and flow through underground aquifers and then into a stream, it
will contain increased amounts of Ca2+, Mg 2+, SO42-, Cl- and CO32- by dissolving them from the
soils and rocks that it flows through.
 The pH of the rain
Water from acid rain is better able to leach certain cations such as Ca 2+, Mg2+ and Fe3+ from
the soil it passes over or though and so concentrations of these ions increase.
 The nature and amount of human activity in the catchment
Land clearing generally leads to more water rapidly running across the land and into streams.
This increases sediment loads. Agricultural pursuits (growing crops and improving pastures)
often lead to fertiliser run-off, and increases nitrate and phosphate concentrations in the
water.
 Effluents discharged into the water bodies
Discharges of raw and/or treated sewage into rivers or oceans increase the concentrations of
many ions, particularly phosphate and nitrate. Even good sewage treatment which removes
nearly all suspended matter, BOD and pathogens can increase the TDS of the water by 200
ppm or more. Storm water run-off in urban areas can similarly increase the concentration of a
variety of ions in water bodies. Industrial effluents, if not carefully monitored and controlled,
can discharge heavy metal ions into water bodies.
 Leaching from rubbish dumps
When rain and storm water flows over and seeps through poorly designed rubbish tips, it
dissolves many harmful substances and carries them into waterways. This is known as
leaching. Dangerous cations from batteries and other items such as phosphates, zinc and
nitrates are carried into streams.
Because of their smaller volumes, rivers and dams are more vulnerable than oceans to
increases in ion concentrations. Nevertheless discharge of poorly treated sewage from big
cities (such as Sydney) and large run-offs from flooded rivers that flow through agricultural
land can affect the ion concentrations in coastal ocean water quite significantly.

Describe and assess techniques for accurate determination of masses of solids
in water samples
A common technique to determine insoluble solids in water is to filter a known volume
through a weighed filter paper. The filter paper and trapped insoluble solids are dried and
weighed. The drying and reweighing are repeated until a constant weight is obtained.
An assessment of the accuracy of this technique would involve comparison of results with a
range of solid concentrations and with a range of filter papers, from coarse papers to papers
with a fine pore size. School laboratory filter paper typically removes 98% of the solids.
Another method is to use a turbidity tube, which relies on a measure against the amount of
light that can be transmitted through a sample of water.
The methods described above are not effective for very small amounts of suspended solids in a
water sample. More accurate measurements are made using electronic instruments that can
detect and record the amount of light that is transmitted through a sample or the amount of
light that is scattered by the suspended solids when a light beam is directed through the
sample.
Total dissolved solids (TDS) can be determined by evaporation to dryness of a known
volume of a filtered sample and weighing of the solid left behind. This technique requires a
large volume of water, great care towards the end of the heating so that the solid does not
splatter out of the heated container, and careful collection of and weighing of the solid.
A simpler and more accurate way of measuring TDS is with a electrical conductivity meter.
Practically all dissolved solids are ionic. Measurement of the conductivity of filtered water can
give a good estimate of the amount of dissolved solids. Automated spectroscopy techniques
such as Atomic Absorption Spectroscopy can be used to give accurate measurements of very
low concentrations (ppm or even ppb) of many dissolved solids.

Describe and assess the effectiveness of methods used to purify and sanitise
mass water samples
There are several methods used to purify mass water supplies. Most are variations on the
following process:

Water is collected in dams and pumped to a treatment site where the larger solids such
as pieces of plastic are removed by screens. Fine particles suspended in water have
electric charges on their surface that keep the particles from joining together and
becoming large enough to settle as sediment. Separation of fine particulate matter
suspended in water involves the addition of coagulants, such as iron (III) chloride,
FeCl3, to the water. The added FeCl3 overcomes these surface charges so the particles
come together. Iron (III) hydroxide is formed by reaction of FeCl 3 and precipitates out
as a floc (flocculant). The floc collects the neutralised particles into large masses that
are more easily filtered. The chloride ions added with the FeCl 3 remain in the water.
They do not have a harmful impact on the water quality. In some instances, after
flocculation the particle size of suspended solids is too fine for filtration through sand
beds and so membrane filters are used. They are more effective than sand filters as
they can have a very small pore size. However membrane filtration is more costly than
sand filtration.

Sanitising mass water supplies involves disinfecting with chlorine gas, Cl2, liquid sodium
hypochlorite solution, NaOCl (aq), or solid calcium hypochlorite, Ca(OCl) 2.
Sedimentation and filtration removes some harmful organisms, such as bacteria,
viruses, cryptosporidium and giardia, but disinfection is needed to ensure
concentrations are acceptably low. In some parts of the water distribution system,
ammonia, NH3, is added to form monochloramine, ClNH2, by reaction with chlorine. This
is a less reactive disinfectant, but lasts much longer in the distribution system.

Assessment of the effectiveness of sanitisation methods should involve microbiological
testing of water samples throughout the distribution system, in particular before and
after sanitisation processes public health surveys and reporting by medical doctors of
incidences of illnesses that are possibly transmitted through water supplies.

A few years ago, some water supplies in NSW were contaminated by the
microorganisms, cryptosporidium and giardia. As a result, water supplies are monitored
daily at water treatment plants and throughout each catchment during storms or other
events that cause a rise in stream water levels and could influence water quality. This is
considered effective, as the cost of treatment of all water supplies with membrane
filters would make treatment very expensive.

Describe the design and composition of microscopic membrane filters and
explain how they purify contaminated water
Microscopic membrane filters have microscopic pores and the use of appropriate sized
filters can avoid the need to chemically treat the water. The filters can be classified as
microfiltration, ultrafiltration, nanofiltration or reverse osmosis membranes depending on the
size of the pore.
The membrane is made from synthetic polymers dissolved in a mixture of solvents and watersoluble powders of a particular size. The mixture is spread out over a plate and left for the
solvent to dry. The polymer membrane containing particles of water-soluble powder is then
placed in water. The remaining solvent and the powder particles dissolve, leaving a very thin
polymer sheet with definite sized microscopic pores where the water-soluble particles were.
Semi-permeable membranes used in reverse osmosis are either made of cellulose acetate or a
layer of polyamide attached to another polymer. Under pressure these polymers allow the
passage of water molecules but not that of most atoms, ions or other molecules.
Water is made to flow across the membrane not through it. This reduces the blockage factor.
Microfiltration removes protozoans, bacteria, colloids, some colouration and some viruses. The
size of the pore determines which sized particle or organism may pass through the membrane.
The finer the pore size the smaller the particle trapped and the more expensive the
membrane.
Re. Pictures CCHSC p. 281 – 283
Advantages
 They can filter out much smaller particles than paper or sand filters can
 They filter out virtually all particles larger than their specified pore sizes
 They are quite thin and so liquids flow through them fairly rapidly
 They are reasonably strong and so can withstand pressure differences across them of
typically 2 to 5 atmospheres
 They can be cleaned (back-flushed) and reused
Membrane filters are widely used for filtering both drinking water and treated sewage. No
filters can remove substances that are actually dissolved in the water such as phosphates and
nitrates and heavy metal ions.

Identify the need for collaboration between chemists as they collect and
analyse data
The analysis of local waterways requires the collaboration of a wide variety of chemists and
other scientists. This allows for effective monitoring and management of our waterways.
AAS testing is carried out in a laboratory while simple chemical and physical measurements,
such as pH and temperature, are carried out in the field. A chemist using complex equipment
in the laboratory on water samples needs to have confidence in the sampling techniques used
by field workers collecting the samples.
A fine example of collaboration is the Streamwatch program in NSW. Streamwatch is a school
and community water quality monitoring and action program. Each year more than 30,000
people in NSW participate in Streamwatch, including students and staff in schools, people in
community groups, thousands of individuals and scientists and project officers in Sydney
Water, the Department of Land and Water Conservation and a number of catchment
management committees and trusts. Data from water quality tests are collected locally and
shared via the Internet. The collaboration between these people results in an effective early
warning program for water quality in NSW.

Gather, process and present information on the range and chemistry of the
tests used to: identify heavy metal pollution of water, and monitor possible
eutrophication of waterways
Heavy metal ions such as Pb2+, Hg2+, Cd2+, Cr2+, Cu2+ and Zn2+ can be identified using the
tests mentioned above.
Eutrophication is the increase in nitrates (NO3-) and phosphates (PO43-) from fertiliser
overuse and sewage. These can be identified using the tests mentioned before.
This leads to:
 Increase of algae in waterways which need these two nutrients
 Algae grow at the surface reducing light for normal water-plants below – these die
 Increased number of bacteria of decay which consume oxygen
 Water becomes oxygen depleted
 End with a smelly pond covered with green scum.
Other Syllabus Points

Analyse information to evaluate the reliability of the results of one of the
above investigations and to propose solutions to problems encountered in the
procedure

Solve problems by devising a procedure that could be used to identify the
water content, alcohol content and aspirin content of a medicine

Gather, process and present information to interpret secondary data from AAS
measurements and evaluate the effectiveness of this in pollution control

Gather, process and present information from secondary sources including
simulations, molecular model kits or pictorial representations to model
isomers of haloalkanes

Perform a first-hand investigations to use qualitative and quantitative tests to
analyse and compare the quality of water samples

Perform first-hand investigations to test for solids in water samples from the
local environment

Gather, process and present information on the features of the local town
water supply in terms of: catchment area, possible sources of contamination in
this catchment, chemical tests available to determine levels and types of
contaminants, physical and chemical processes used to purify water, and
chemical additives in the water and the reasons for the presence of the
additives