Chemistry HSC Full Notes BEST NOTES

Chemistry HSC
Note: This is not the final revision of my notes (I’m constantly revising them as I do papers), and
there may be a few areas of error or unclear explanations. However, I’ve gone through it a number
of times, and it should be mostly very accurate and comprehensive. If you find anything wrong, it
would be nice if you could tell me on korenxii@hotmail.com so I can either discuss it or change it.
Good luck for the HSC guys
Organic Chemistry – the study of compounds containing carbon

THIS IS BACKGROUND INFO
Organic chemistry is separate because we can look at all of the included chemical groups in a
unifying way, through the bonding properties of carbon.
Study of major groups:
o Oxygen-containing compounds e.g. alcohols
o Hydrocarbons e.g. petroleum
o Carbohydrates e.g. sugars
o Nitrogen-containing compounds e.g. amino acids  proteins
Etc.
Hydrocarbons
When all bonds are single, they are called alkanes. This is a family of compounds, represented by a
general formula CnH2n+2, aka a homologous series. They have similar properties and reactions.
There are ‘straight’ chain alkanes.
e.g.
C C C
C C C
The 109o, zig-zag bonding shape is due to the tetrahedral nature of single bonds. Carbon atoms
always form 4 bonds. If they don’t you’re doing something wrong.
Branched chain (one is attached to at least 3)
C C
C C C
Methane, CH4, ethane, C2H6, Propane, C3H8, and Butane, C4H10 are all alkanes.
Physical Properties
C1 to C4 are gases at room temp, C5 to C18 are colourless liquids, others are solids.
The density of alkanes are significantly less than water (1.00g/mL), are non-conductors of
electricity and are insoluble in water. The reason for their insolubility is that C-C bonds are non-
polar, and C-H bonds are only slightly. This slight polarity is cancelled by symmetry in structure.
Weak dispersion forces, relatively low boiling/melting points. Boiling/melting points increase as
molecular weight increases, due to stronger dispersion forces (more electrons). Volatility decreases
as molecular weight increases.
Alkenes
Chem notes David Lee BHHS 2007 1
Contain a double bond between a pair of carbon atoms. Homologous series, formula CnH2n, planar
shape. There are different ways of representing structure:
(2)Full Structural Formula – shows planar geometry around double bond, and tetrahedral around
other carbon atom
(3)Intermediate type – infers tetrahedral shape
(4)Condensed structural formula – no attempt to show structure, but enough information is
provided
Isomers are different compounds with the same molecular formula but different structural
formula. The double bond can be at different positions in the compound.
e.g.
Physical Properties
Straight-chain alkenes similar to alkanes. Densities similar to corresponding alkanes, insoluble in
water.
Alkynes
Contain a triple bond between carbons. CnHn-2. As with alkenes, isomers are possible. They are non-
polar, low boiling points and insoluble in water
Naming Alkanes, Alkenes and Alkynes

o Stem telling length of carbon chain
C1 meth- C4 but- C7 hept-
C2 eth- C5 pent- C8 oct-
C3 prop- C6 hex-
o Look at the longest possible chain, then pick a prefix
o Look for branches, and use a number to denote their position, starting from the closest
end e.g. 2,3 – dimethylpentane or 2 - dimethylpentane
o If double or triple bonds present, set this as priority (start counting closest to the bond)
and first state branches then double/triple bond e.g. 2– methyl – 1 – propene (methyl on
second branch and double bond on first)
o If compound is cyclic, add a cyclo- before the name of the main branch e.g. 1,2,3 -
trimethylcyclohexane
Saturated and unsaturated compounds

Alkenes and Alkynes – unsaturated, possible to attach more hydrogen
Alkanes – saturated, max no. of H atoms that skeleton can hold
Functional Group
The functional group of carbon compounds is the most reactive area of the compound. In alkenes
and alkynes, the double/triple bonds are the functional groups. When a hydrogen atom is replaced
with a halogen atom, e.g. OH, the halogen becomes the functional group.
Molecules with a particular functional group react similarly, regardless of the attached
chains.
Alkanols are alkanes with one H replaced by an OH group. They are named with the ‘e’ replaced by
an ‘ol’, and a prefix number to denote the position of the hydroxyl group. This group is the
functional group, and provides high melting/boiling points due to polar bonds.
Primary alcohols have one carbon bound to the carbon w/ OH group, secondary have two and
tertiary have three. Extent of hydrogen bonding depends on exposure of OH group, most exposed in
primary, highest boiling/melting points etc.
• Construct word and balanced formulae equations of chemical reactions as they are
encountered
Types of Organic reactions

Substitution – replacement of one atom or group by another
Addition – adding atoms or groups of atoms to alkenes or alkynes (bond breaks, new atoms are
added on)
Elimination – a small molecule breaks off and a double bond is formed in the original (reverse of
addition)
Condensation – two molecules react, forming a new compound and a small molecule (usually
water)
Hydrolysis – the action of water on a molecule results in two new products
• Identify the industrial source of ethylene from the cracking of the fractions from the
refining of petroleum
Ethylene is produced from natural gas or crude oil (mixtures of hydrocarbons, containing mainly
alkanes and cycloalkanes and smaller amounts of unsaturated including alkenes), which is called
feedstock. The feedstock is refined by fractional distillation to obtain alkenes since alkanes are
susceptible to combustion and unreactive (not useful as starting material).
Ethylene is the most versatile, but not found in large quantities in feedstock. Produced from other
hydrocarbons in ‘cracking’ (a process where hydrocarbons of higher mol mass are converted to
lower mol mass via breaking of chemical bonds). There is greater demand for some fractions than
others (e.g. gasoline > heavier hydrocarbons), and fractions from crude oil are not in optimum
ratios, hence cracking. Note that air needs to be excluded to prevent combustion. Ethylene is simple
and can be synthesised from many different hydrocarbons. Three ways:
1. Thermal cracking – requires very high temps and generally not used. End products hard to
control since many places where bonds could break, early method. Accelerates reaction and
drives equilibrium to reactants.
2. Catalytic cracking of fractions separated from petroleum. – material is passed over a
catalyst at a temperature of about 500oC, and the particles adsorb onto the catalyst and have
their bonds weakened, resulting in decomposition. E.g. C10H22(g) -> C8H18(g) + C2H4(g). Alkane
splits further into smaller alkenes until propene/ethylene formed. Catalysts allow it to be
carried out at lower temperatures. Zeolite (by mid 1970’s) is the main catalyst, and is a
crystalline substance of Al, Si and O. Usually fine powder (higher surface area for action of
catalyst) circulated through feedstock. Zeolite gives greater control over products under
different conditions of temperature and pressure (thus increasing yields of desired products)
i.e. C18H38(g) ----(zeolite catalyst)---> 4 C2H4(g) + C10H22(g)
3. Steam cracking of ethane and propane – ethane from natural gas deposits fed into
furnaces with steam, heated between 750 – 900oC causing much ethane to be converted to
ethylene
i.e. C2H6(g) -> C2H4(g) + H2(g)
Propane can also be used:
C3H8(g) -> C2H4(g) + CH4(g)

• Identify that ethylene, because of the high reactivity of its double bond, is readily
transformed into many useful products
Ethylene’s C=C double bond is highly reactive, allowing it to react with molecules to form many
useful products
Dilute
H2So4 cat
w/ water
Br2, non
KMnO4, H+
aqueous
Acidified solvent
HCl, non
aqueous
solvent
Reaction of alkenes
Characteristic reaction of alkenes is addition reaction. Two new atoms or groups of atoms are added
across double bond, one to each carbon. The C=C is converted to a single bond and a saturated
hydrocarbon is produced. General eqn:
H2C=CH2 + X-Y => XH2C-CH2Y
1) Addition of hydrogen to ethylene (hydrogenation) - ethylene to ethane by

heating with hydrogen in presence of nickel, platinum or palladium
2) Dibromoethane - Used as a petrol additive - halogen reactions are useful for
distinguishing between saturated and unsaturated hydrocarbons. E.g. A non
aqueous solution of bromine (e.g. solvent carbon tetrachloride) when added to an
alkene causes the solution to lose its colour as bromine becomes incorporated into
the alkene:
[CH2=CH2(g) + Br2(l)  CH2Br-CH2Br(l)] (petrol additive)

Alkanes do not react with NA bromine unless exposed to UV. In aqueous
solutions, the reaction may be the same as above, but due to the presence of water
products can include:
[CH2=CH2(g) + HOBr(aq)  CH2OH-CH2Br]
Hydrogen bromide reaction:
[CH2=CH2(g) + HBr(g) -> CH3 – CH2Br] (What states?)
3) Chloroethene – Monomer for PVC
CuCl +150o C
[2CH2=CH2(g) + Cl2(g) + ½ O2(g)    → 2CH2=CH-Cl(g) + H2O(g?)]
2
4) Styrene – produces from benzene and ethylene via the intermediate ethylbenzene
5) Ethanol – Used as a fuel in automobiles and as an industrial solvent
( dilute ) H SO
[CH2=CH2(g) + H2O(l)    4 → CH3-CH2OH(l) ]
2
6) Ethylene oxide and ethanediol – fumigant (former), manufacture of polymers

(polyester fibres and PET) and antifreeze (latter)
o
Ag + 250 C
[C2H4(g) + ½ O2(g)   → C2H4O(g)]
+
H
[C2H4O(g) + H2O(l) →
 OH-CH2-CH2-OH]
• Identify the following as commercially significant monomers

o Vinyl chloride
o Styrene
By both their systematic and common names
Vinyl chloride – chloroethene CH2 = CHCl

Monomer for the production of PVC plastics which are widely used in applications such as
electrical insulation, plumbing and garden hoses with various additives to change physical
properties
Styrene – ethylbenzene C6H5CH=CH2 (also known as phenylethene)

Production of polystyrene, most stiffened of common plastics due to large
phenyl side group. Stable due to presence of C-C and C-H bonds only,
minimal chain branching means it can be formed into clear objects. Tool
handles, car battery cases, CD cases. Gas can be bubbled through to create
foam (foam drink cups), making it soft and light.
• Identify data … to compare the reactivities of appropriate alkenes with the corresponding
alkanes in bromine water
Prac – Reactions of hydrocarbons with bromine water

Risk analysis:
Hazard Risk Control
Bromine water Corrosive, and toxic, can Wear safety goggles
cause skin burns Use small amounts to
minimise vapour
Cyclohexane Highly flammable Wear safety glasses
Eye and skin irritant with Keep away from hot
severe redness and pain surfaces, flames or sparks
Toluene Highly flammable, fire Keep away from hot
hazard surfaces, flames or sparks
Eye and skin irritant with Polyvinyl gloves
severe redness and pain
Aim: To compare reactivities of an alkene (cyclohexene), alkane (cyclohexane), and an aromatic

hydrocarbon (toluene) in bromine water
Method:
1). Four semi-micro test tubes were half-filled with bromine water, cyclohexane, cyclohexene and
toluene respectively, using eye droppers
2). Bromine water was mixed with the other substances by placing a few drops of bromine water in
each micro-test tube with a dropper
3). The test tubes were tapped, and observations recorded
Results:
Cyclohexane – none
Cyclohexene – forms clear solution
Toluene – none
The functional group reacting with bromine is the double bond present in alkenes, this
decolourises bromine water. Addition reactions. These reactions are addition reactions:
Bromine w/ water
Br2( aq ) + H 2 O(l ) ←→ HOBr − ( aq ) + H + ( aq ) + Br − ( aq )
Bromine w/ cyclohexene (top)

Br2( aq ) + C 6 H 10 (l ) ←→ C 6 H 10 Br2 ( aq )
Bromine water w/ cyclohexene (bot)

HOBr( aq ) + C 6 H 10 (l ) 
→ C 6 H 10 BrOH ( aq )
Toluene did not react as aromatic molecules have

delocalised electrons which do something???
************
Bromine with cyclohexane, this is substitution:

UV
Br2 ( aq ) + C6 H 12 (l ) →
 C 6 H 11 Br( aq ) + HBr(l )
Requires UV light to break off hydrogen atom and allow reaction

Prac – Reaction of lycopene with bromine water
Aim: To determine the effect of bromine water in varying amounts on the spectrum of colours
reflected by lycopene
Method:
1). Five semi-micro test tubes were filled halfway with tomato juice
2). An eye dropper was used to place 1, 2, 3, 4 and 5 drops of bromine in each of the test tubes
respectively
3). The solutions were stirred with the stirring rod until colour appeared
4). The colours and the corresponding amounts of bromine water in each test tube were recorded
Results:
Test tube no. No. of drops of Bromine Colour
1 10 Blue
2 8 Turquoise
3 6 Green
4 4 Khaki
5 2 Orange
Varying amounts of bromine in tomato juice changes the number of delocalised electrons in
lycopene molecules, changing the spectrum of colours absorbed and resulting in different reflected
visible spectra
• Identify that ethylene serves as a monomer from which polymers are made
Polymerisation is the process of bonding monomers together to form long chains.

Polymers are macromolecules consisting of small repeating units called monomers joined by
covalent chemical bonds. Polymers can be divided into two categories:
1. Natural polymers – naturally occurring polymers used by humans since
ancient times (E.g. cellulose, silk, rubber)
2. Synthetic – more recent man-made polymers. Replacing natural since they do
not corrode, are lightweight and relatively cheap. Celluloid was first
commercially manufactured plastic, but highly flammable nature meant it
was replaced
Ethylene serves as a monomer due to the reactivity of its double bond. It has a structure that can
change to accommodate the additional bond needed to join repeating units together.
• Identify polyethylene as an addition polymer and explain the meaning of this term
Polyethylene is an addition polymer, it is created through addition polymerisation.

Def: The monomers add to the chain so that all atoms in monomer are present in polymer. It
involves unsaturated monomers (a molecule containing a double or triple bond) joining together.
One C=C is broken up and resulting molecules link up, since this provides molecules with extra
bonding capacity. E.g. for polyethylene
Addition polymerisation requires a catalyst or

initiator to start. Other polymers formed by
addition are Polyvinyl chloride (PVC), polystyrene
and Teflon.

• Outline the steps in the production of polyethylene as an example of a commercially and
industrially important polymer
General outline: Ethylene can be changed from gas to liquid under high pressure. This liquid
ethylene can be heated in the presence of a catalyst to form polyethylene.
Two forms of polyethylene can be produced, each with differing methods and varying properties:
o LDPE (reaction conditions 100 – 300oC, 1500 – 3000 atm) – Polymerisation consists
of three stages
 Initiation – organic peroxide catalyst. They produce free radicals (molecules
with unpaired electron), such as H-O. which is a hydroxy radical. This causes
the double bond in ethylene to break and form a bond with the radical.
CH2=CH2 +R●  RCH2-CH2●
 Propagation - The resulting
molecule contains an unpaired
electron. Bonds to another
ethylene molecule through the
same process etc. (Chain
propagation reactions).
Backbiting, where the chain curls onto itself and the free electrons takes a
hydrogen atom from an existing CH2 group, causes branching.
 Termination – at various times, it is possible for two free radical polymers

to react to form a covalent bond, ending propagation (chain terminating
reaction).
o HDPE (50 – 75oC, <1 atm) – polymerisation process is same as above, but Ziegler-
Natta catalyst (TiCl4, Al(C2H5)3 used. Ethylene molecules are added to chain on surface
of catalyst, reducing backbiting and branching.
Comparison of structures and properties
LDPE HDPE
Higher degree of branching, meaning less Lower degree of branching, meaning
dispersion forces between strands, more dispersion forces within strands,
making it softer and more flexible making it harder and more rigid
Less dense More dense
These products are plastics. Plastics are manufactured materials containing combinations of organic
and inorganic elements. They are solid in the finished state but fluid at some stage, and able to be
formed into shapes by application of heat and/or pressure.
Factors affecting the properties of polymers:

o Length of chain (no. of monomer units) – those with longer chains are stronger since
greater dispersion forces between chains
o Arrangement of chains relative to each other – when chains are unbranched, they are
lined up and closely packed creating crystalline areas resulting in stronger and less
flexible plastics. Amorphous regions were alignment is more random, produce weaker
and softer plastics. Polymer fibres drawn through a small hole aligns them and increases
strength
o Function groups in monomer units – polar functional groups increase intermolecular
forces between polymer molecules, increasing hardness.
o Cross-linking between polymer chains – covalent links between polymer chains makes
polymers very hard and difficult to melt.

o Additives – few polymers are used in pure form, additives improve or extend properties.
Additives can include pigments, plasticisers to soften, stabilisers to increase resistances
to decomposition etc.
The variable chain length leads to many uses, with shorter lengths for food packaging and milk
containers, to longer lengths (800,000 atoms per molecule) for artificial ice rinks.
• Describe the uses of the polymers made from the above monomers in terms of their
properties
Polymer Properties Uses

LDPE 1). Resists water and chemicals 1). Pipes for farm and
industry
2). Electrical insulator 2). Wire and cable
sheathing for telephone,
coaxial, submarine
television and radar
3). Easy and cheap to process 3). Shopping and garbage
bags
4). Waterproof 4). Milk and fruit juice
packs, food containers
5). Non-toxic
HDPE 1). Can take high pressures 1). Coating in steel pipes
in high pressure gas mains
2). High tensile strength 2). Fibres for ropes, fishing
nets
3). Chemical resistance 3). Moulded into
containers to hold petrol,
oil, detergents and acids
4). Durability and toughness 4). Children’s toys, plastic
buckets, playground
equipment
Polyvinyl 1). Soft and pliable 1). Wallpaper, clothing
Chloride OR (depending on additives) upholstery
2). Rigid 2). Water pipes, guttering
3). Coatings on materials
3). Resists burning to make flameproof
4). Flooring; tiles, roll
4). Low static electricity flooring and carpet
backing
Polystyrene 1). Rigid and electrical insulator 1). Television backing,
--- hairdryers, washing
As foam: machines
2). Chemically unreactive 2). Food containers
3). Low density 3). Marker buoys,
--- surfboards
4). Resists high impact 4). Shoe heels, toys
• Discuss the need for alternative sources of the compounds presently obtained from the
petrochemical industry
Petroleum fractions have been the most convenient and economical raw material for synthetic
polymers. However, alternatives are being sought since:
1) The current source is non-renewable, and the move to more renewable resources will allow
us to continue manufacturing petrochemical products (supplies will run out)
2) Petrochemical products are (non-biodegradable?) and contribute to the degradation of the
environment
o A solution is the use of biomass (organic material from living organisms). All living
organisms produce biopolymers, which are naturally occurring polymers made
entirely or in large part by living organisms
o They are advantageous since they are renewable, and can be used indefinitely with
careful use, and are biodegradable since the bonds within the molecule can be broken
down by bacteria and fungi, so they do not contribute to the degradation of the
environment.
• Explain what is meant by a condensation polymer
A condensation polymer is a polymer that was produced through the reaction of two different
functional groups in which a small molecule (usually water) is eliminated and the two groups
become linked together. Condensation reactions involve saturated molecules. Common groups are –
COOH (carboxylic acid), –OH (alcohol) and –NH2 (amine) group. Condensation polymers do NOT
require identical monomers
• Describe the reaction involved when a condensation polymer is formed
e.g. Condensation polymerisation of nylon
Practical
Aim: to produce nylon using interfacial polymerisation
Equipment:
2 x 100 mL beakers
Tweezers
Glass stirring rod
20mL of 1,6 – diaminohexane solution
20 mL of 10% sebacoyl chloride in hexane
Procedure:
1). 1,6 – sebacoyl chloride was added to a beaker
2). The diamino hexane was run very carefully down the side of the beaker, so that the two
solutions mix as little as possible
3). The white material formed between the two layers was clamped using the tweezers
4). The material was drawn away from the beaker and onto the glass stirring rod, being careful to
keep away from sides of beaker
5). The material was wound onto the stirring rod, dried and examined
Results/Analysis:
A variety of monomers can be used to manufacture

nylon, it is simply a generic name for a group of
polyamide polymers, the common feature being the
repeated –CONH- bond.
This experiment used interfacial polymerisation. It

is thus named since the reactants bond together and
form nylon at the contact surface between them.
PET (polyethene terephthalate) is also condensation
• Describe the structure of cellulose and identify it as an example of a condensation

polymer found as a major component of biomass
Biomass – organic material derived from living organisms

Glucose C6H12O6 is a carbohydrate of form:
The presence of five hydroxyl groups allows glucose to form polymers

such as starch, cellulose, and glycogen.
Cellulose is a biopolymer, a polymer naturally synthesised by living
organism. They are condensation polymers, since water is eliminated
from a reactive functional group when glucose units join together.
o Many glucose units linked together (polysaccharides).

o The glucose units come together, causing two hydroxyl groups to react, a hydrogen ion
dissociates from one hydroxyl and combines together with the other hydroxyl to form
water.
o The leftover oxygen atom then forms a covalent bond between the molecules. These
parts are called the functional units
o Cellulose is a linear polymer, producing a fibre-like material. The beta linkages result in
flat, ribbon-like strands which are closely packed and have strong hydrogen bonds
between them (cellulose strands). This gives cellulose its strength and rigid structure.
o It is the main component of plant cell wall and major structural component of woody
plants and natural fibres
o This makes it the most abundant polymer known on earth.
• Identify that cellulose contains the basic carbon-carbon structures needed to build
petrochemicals and discuss its potential as a raw material
Cellulose contains three-carbon and four-carbon chains with attached hydrogen and hydroxl groups.
Many polymers such as polypropylene are made from three-carbon and four-carbon monomers. If
cellulose can be broken down and these chains isolated, it can be used to produce polymers. Large
amounts occur naturally such as in plant cell walls, and large amounts left over from agriculture.
Use as raw material can be achieved by:

1. Modification of existing biopolymer chains to meet specific applications (e.g. addition of
functional groups)
2. Breakdown into smaller molecules which can then be used to build synthetic polymers e.g.
thermochemical (steam/acid) pre-treatment followed by hydrolysis using enzyme cellulase.
This produces glucose which can be dehydrated to ethene. However, this is more expensive
than using hydrocarbon sources.

Rayon is created from regenerated cellulose sourced from waste paper, straw, husks from wheat and
corn and wood pulp. The fibres are chemically treated with sodium hydroxide and carbon disulfide
to soften and break them down into smaller units.
A potential area is the use of biopolymer-based plastics is food wrappers and disposable containers,
since they are used only once. An example is a US company which has made packaging products
from corn starch.
• Use available evidence to gather and present data from secondary sources and analyse
progress in the recent development and use of a named biopolymer
o Biopol, PHA (polyhydroxyalkanoates) copolymers, a family of microbial energy reserves

accumulating as granules within the cytoplasm of cells.
o PHA’s
• polyester thermoplastics
• properties similar to oil-derived polymers (e.g. melting temp 50-180oC)
• mechanical properties can be changed to range from elastic rubber or hard crystalline
plastic.
Simplest are PHB’s

Production is carried out by the following process:
o Alcaligenes Eutrophus, a bacteria widely found in soil and water, is fed a precise
combination of glucose and propionic acid, producing PHB’s as energy storage
o The PHB’s are extracted and can be polymerised to create a plastic with properties
similar to polypropylene
• Excellent flexibility and toughness
• Stable in air, humid conditions
• Biodegrades in microbially active environments, bacterial and fungi microorganisms
can utilise PHA’s as a source of energy by breaking it down using enzymes
(depolymerases).
Potential uses
o Biocompatability - useful in several medical applications such as controlled drug release,
medical surtures, bone plates
o Flexibility and toughness - structural materials in packaging products
o Biodegradable – can be used in food packaging, natural breakdown reduces landfill
Development
Current work by Metabolix, successfully engineered bio-factories to demonstrate economic
production of a broad range of PHA’s. Demonstrated fermentation on a tonnage scale, cost to be
under a dollar a pound.
Currently working to produce PHA’s directly in non-food crop plants.
Disadvantage: Current high cost of production as opposed to crude oil sources
• Describe the dehydration of ethanol to ethylene and identify the need for a catalyst in this
process and the catalyst used
• Describe the addition of water to ethylene resulting in the production of ethanol and
identify the need for a catalyst in this process and the catalyst used
• Model the above two processes

Ethanol and is an alkanol. It is tetrahedral about each carbon and bent around oxygen atom (not
shown here). Ethylene can be made through dehydration, heating with concentrated sulfuric acid or
a porous ceramic catalyst (>350 in industry):
H 2 SO4 / H 3 PO4
CH 3 − CH 2 − O − H    → CH 2 = CH 2 + H 2 O
Reverse reaction is hydration, requires dilute aqueous sulfuric acid:

H SO dilute
CH 2 = CH 2 + H 2 O 2 4 → CH 3 − CH 2 − O − H
They are general, apply to any alkanol or alkene e.g. 1-pentanol to 1-pentene
• Describe and account for the many uses of ethanol as a solvent for polar and non-polar
substances
Risk Analysis
Hazard Risk Control
Iodine Toxic – fatal if swallowed. Safety glasses, effective
Corrosive, causes burns ventilation
and damaging to lungs if
inhaled
Oxalic acid Poisonous if swallowed, Gloves, avoid generation
inhaled or absorbed of dust
through skin
Prac – Ethanol as a solvent

Aim: To test the solubility of various materials in ethanol
Method:
1). 20 mL of ethanol was poured into each of 10 test tubes using measuring cylinders
2). A rice-grain amount of each solid was placed into successive test tubes, and a few millilitres of
each liquid placed into the remaining test tubes using an eye dropper.
3). Each test tube was agitated by tapping and gently shaking
4). Observations were recorded
Results:
Solute Solubility
Sodium chloride No
Napthalene Slightly
Cyclohexanol Yes
Glycerol Yes
Iodine Yes, dark red
Oxalic acid Yes, purple
Boric acid Yes
Glucose Yes
Wax No
Urea Yes

o Ethanol has a single hydroxyl group which is attached to an aliphatic
o Allows it to dissolve substances with polar covalent bonds, hydrogen bonds form
o Able to dissolve hydrocarbons and non-polar due to the formation of dispersion forces
with the aliphatic group.
o Widely used as alternative solvent in dissolving medicines, cosmetics, food flavourings,
alcoholic beverages, low toxicity so relatively safe
Ionic (sodium chloride): Unable to dissolve since strong ionic bonds holding atoms together, and
intermolecular formed inadequately strong to break apart lattice
Polar covalent bonds (cyclohexanol C6H11OH, glycerol C3H5(OH)3, oxalic acid C2O2(OH)2, Boric
acid B(OH)3, Glucose C6H12O6, Urea CO(NH)2) : Polar covalent bonds such as those in hydroxyl
groups allowed ethanol to bond strongly and dissolve them
Macromolecules (Wax C24H50) : Though only held together with weak dispersion forces, its large
size means a larger surface area of contact between molecules and thus more total dispersion forces.
Ethanol was unable to dissolve
Non-polar molecules (Iodine, Napthalene, heptane,
pentane) – iodine is diatomic and has very weak
dispersion forces holding together, but ethanol can
form dispersion forces with iodine molecules and pull
away. Napthalene is an aromatic hydrocarbon and
does not attract strongly, but dissolves in a similar
fashion, same for heptane and pentane which are
short-chain hydrocarbons.
With 1,2,3 – propanetriol:
• Outline the use of ethanol as a fuel and explain why it can be called a renewable resource
Ethanol is a flammable liquid, burning with the reaction:
C 2 H 5 OH (l ) + 3O2 ( g ) → 2CO2 ( g ) + 3H 2 O( g )
It is also easily transportable, and was used by hikers and campers. It has thus been proposed as an
alternative fuel source, having already been used as an ‘extender’ in world war 2. The purpose of
ethanol is to:
1. Reduce greenhouse gas emissions
2. Reduce reliability on non-renewable fossil fuels
Engines would not need any modification to run 10-20% ethanol fuel, and is renewable since
synthesised in sugar cane from carbon dioxide, water and sunlight. Burning produces carbon
dioxide and water which can then be re-used to produce ethanol, so it follows an almost indefinite
material cycle.
It has thus been promoted for motor cars to supplement and replace petrol.
• Describe conditions under which the fermentation of sugars is promoted

• Summarise the chemistry of the fermentation process
• Present information from secondary sources by writing a balanced equation for the
fermentation of glucose to ethanol
o Fermentation requires a carbohydrate as starting material, such as glucose, sucrose or

starch

o Disaccharides such as sucrose and polysaccharides such as starch need to first be
broken down into monosaccharides such as glucose/fructose by enzymes in the
mixture
o This carbohydrate is placed in the presence of yeasts, which produce enzymes that
break it down to ethanol and carbon dioxide
 C6 H 12 O 6( aq )  → 2CH 3 − CH 2 − OH ( aq ) + 2CO2( g )
yeast
o The optimum conditions are 37oC and anaerobic conditions

o Reaction is exothermic, so temperature needs to carefully controlled
o Fermentation can occur until 15% ethanol, then the yeasts cannot survive and
fermentation stops – extra ethanol can be added or distillation used to increase
concentration
• Solve problems, plan and perform a first-hand investigation to carry out the fermentation
of glucose and monitor mass changes
Prac – Fermentation of glucose

Aim: To investigate the fermentation of glucose
Method:
1). One gram of beef extract and 25 grams of glucose and 7 grams of yeast were measured out using
an electronic beam balance
2). A test tube with barium hydroxide was weighed on the triple beam balance and its weight
recorded
3). The beaker was filled with 300mL of tap water and heated over a bunsen burner until 40oC
4). The beef extract, warm water, yeast and glucose were quickly poured into the conical flask and
weighed on the electronic beam balance
5). The apparatus was set up as shown below:
6). After a week, the tubes and stopper were removed, then
the fermented solution and test tube with barium hydroxide
were weighed on the triple beam balance and analysed
7). Experiment was repeated 2 times
7). Steps 1-6 were repeated without the yeast, to act as a
control
Results:
Result Loss in mass of conical Gain in mass of test tube
flask and contents (g) contents (g)
1 8.0 5.9
2 9.8 4.8
3 9.1 5.0
Average 9.0 5.2
Control:
Result Loss in mass of conical Gain in mass of test tube
flask and contents (g) contents (g)
1 0 0
2 0 0
3 0 0
Average 0 0

The formation of carbon dioxide was evidenced by the formation of barium carbonate in the test
tube. The control showed that the loss in mass, and thus creation of carbon dioxide, was caused by
the yeast fermenting glucose.
Equation:
C6 H 12 O 6( aq )  → 2CH 3 − CH 2 − OH ( aq ) + 2CO2( g )
yeast
Reaction of Barium hydroxide with carbon dioxide:

Ba (OH ) 2 ( s ) +CO2( g ) → BaCO3( s ) + H 2 O(l )
Nine grams of carbon dioxide was given off in this experiment (loss in mass of conical flask) the
gain in the test tube was not used since some carbon dioxide escaped through holes around the
stopper.
moles of CO 2
9.0 g
=
(14.01) + (16.00) * 2
= 0.20moles
Through stoichiometry, 1 mole of glucose produced 2 moles of CO2 and ethanol each. Therefore
0.20 moles of ethanol produced
Mass of ethanol = molar mass ethanol x moles produced = 10g (3.3 % w/v)
This reaction did not go to completion, 6 grams of glucose left. This is most likely because the yeast
were saturated in ethanol and could no undergo further fermentation, or were no left for sufficient
time. Some discrepancy could have been caused by measurement error.
• Process information from secondary sources to summarise the processes involved in the
industrial production of ethanol from sugar cane
Industrial production of ethanol from sugar cane

1. Feed preparation
o Crushing – sugar cane is crushed to remove high-quality sugars and molasses, which
are used in fermentation
o Saccharification – bagasse, the other constituents of sugar cane (50% cellulose)
undergo a multi-step hydrolysis process, using an enzyme and sulfuric acid as catalyst
to produce glucose compounds
3. Fermentation – yeast and anaerobic conditions ferment glucose compounds to a certain
concentration of ethanol (max 15%). (Note: Fermentation less efficient than hydration)
4. Purification of mixture – waste products are removed and distillation is used to concentrate
ethanol
5. Addition of gasoline – varying amounts of gasoline are added to produce a commercial
product, ‘gasohol’ or E10 is 10% ethanol 90% gasoline
• Assess the potential of ethanol as an alternative fuel and discuss the advantages and
disadvantages of its use
Note: it is more expensive to dehydrate ethanol than it is to purify ethylene from crude oil (this
doesn’t really go here, I dunno where to shove it)

Ethanol is flammable liquid that is suitable as a fuel. It can be fermented from biomass such as
sugarcane and corn, requires land for agriculture and infrastructure for a fermentation industry to be
set up. If these requirements are met adequately, the advantages of ethanol as an alternative fuel are:
o Renewable – products of ethanol combustion can theoretically be completely recycled to
produce more ethanol
o Intrinsic anti-knock properties – circumvents the need for toxic anti-knock agents due to
presence of oxygen, increases octane of fuel
o Burns more cleanly due to presence of oxygen, , reducing toxic emissions (such as
hydrocarbons)
o Reduction of net emission of greenhouse gases due to reuse of carbon dioxide by
biomass used to produce ethanol
o Lower blends (<20%) do not require engine modification
o Predicted that household wastes will be recycled to produce ethanol – reduces dumping
emissions
Disadvantages are:
o Large areas of arable land needed for agriculture, associated land degradation such as
erosion and fertiliser run off
o Blends above 10% shown to be damaging to cars designed for gasoline, including
increased carbon deposits on pistons and corrosion of metallic engine components
o Greenhouse reductions hindered by use of fossil fuels and emission of toxic waste
products in transportation/production of ethanol
• Process information from secondary sources to summarise the use of ethanol as an

alternative car fuel, evaluating the success of current usage
Higher blends of ethanol require special engines suited, lower blends (<20%) do not require this. It
is extensively used in some countries, such as Brazil and the US.
o In Brazil, a large portion of cars are currently “flex-fuel”, allowing them to use both
ethanol and gasoline, 80% of cars produced in 2005 were flex-fuel
o Government subsidies and rising petroleum prices have successfully encouraged mass-
uptake of ethanol use. Pure ethanol and 25% ethanol are available at nearly all gas
stations
o Its use has had noticeable improvements on air quality due to more complete combustion
o Brazil is approaching self-sustainability in areas of ethanol use due to its large areas of
arable land and tropical climate
The Bad:
o There are situations where ethanol cannot replace fossil fuels, such as diesel fuels, and
they continue to be burnt
o The popularity of ethanol depends on government subsidies. Production of ethanol
requires a large investment of money and energy, and costs more than petroleum to
produce
o Requires destruction of rainforest which acts as carbon sink, offsets greenhouse
reductions
In Australia:
o Ethanol costs more than petrol to produce, so subsidies are provided to encourage
addition of ethanol to petrol
o Ethanol cars are in a small minority
o Public suspicion about fuel, since some independents add excessive ethanol causing
engine wear, and claims by manufacturers that blends above 10% will void warranties.
Federal government has decided to limit mixtures to max 10% ethanol
o Skepticism about mass-implementation, no reliable studies showing improvements in air
quality through ethanol industry and worry about associated environmental costs such as
land degradation
o Ethanol uptake is less successful due to lack of arable land for feedstock growth, higher
individual wealth and higher costs of labour
Evaluation:
Highly successive in some countries, but less successful in others with less land resources. Usage is
hindered economically depending on economic situation, and for some is not economically feasible
due to high energy and monetal cost.
• Define the molar heat of combustion of a compound and calculate the value for ethanol
from first-hand data
Molar heat of combustion is the amount of heat energy released when one mole of the substance
undergoes a combustion reaction.
Prac – Heat of Combustion of Alcohols

Aim: To measure the amount of energy produced when alcohols are burned and thus calculate their
molar heats of combustion
Method:
1). A copper calorimeter and spirit burner were weighed using an electronic beam balance
2). A measuring cylinder was used to pour 100 mL of water into the copper calorimeter
3). The apparatus was set up as show below:
4). The initial temperature of the

water was measured using the
thermometer
5). The spirit burner was
uncapped, lit using a matched
then allowed to burn for 30
seconds
6). The spirit burner was capped
and the temperature of the water
taken again
7). The mass of the spirit burner was weighed on the electronic beam balance
8). Steps 3-7 were repeated for the other alcohols
Results:
Molar heats of combustion
Methanol: -410 kJ/mol
Ethanol: -640 kJ/mol
Propanol: -980 kJ/mol
Butanol: -1100 kJ/mol

Butanol produced the most soot since it had the longest carbon chain, increasing tendency for
incomplete combustion:
C 4 H 9 OH (l ) + 2O2 ( g ) 
→ 4C( s ) + 5 H 2 O(l ) + energy
Shorter chain fuels release less energy per molecule and react at an overall slower rate, meaning
that the immediate availability of oxygen molecules is adequate to ensure complete combustion.
Other fuels react at faster rates due to larger release of energy per molecule, and the rate of oxygen
diffusion into the immediate environment of the fuel molecules is inadequate to prevent atoms
within the fuel reacting amongst themselves. These latter reactions occurs less readily than
complete combustion reactions, but form eventually with adequate particle energy and collision.
o Butanol produced the most heat per gram

o The larger the molecules, the more heat released per gram
o This suggests that the breaking of a CH2 group is more exothermic than breaking off an
‘OH’ group, since one gram of lighter alcohols would contain more hydroxyl groups
Ethanol would be the best fuel source since it has the best balance of achieving complete
combustion and burning with more heat per gram.
Safety considerations
Danger in transporting spirit burners – carried only while unlit so would not ignite if dropped and
shattered
Methanol – vapours are toxic in larger doses, do not open spirit burner
Ethanol – skin and eye irritant, wear safety glasses, do not open spirit b urner
Errors:
Experimental results differed from theoretical results due to:
o Conduction and radiation of heat from copper calorimeter into surrounding environment,
reducing heat of combustion values
o Inaccuracies in measuring equipment
This could be remedied by:

o Burning for shorter periods of time, less radiation/conduction of heat
• Explain the displacement of metals from solution in terms of transfer of electrons

• Identify the relationship between displacement of metal ions in solution by other metals to
the relative activity of metals
A displacement reaction is where a metal converts the ion of another metal to the neutral atom.
o Different metals have different reactivities

o Metals with higher electronegativity or lower reactivities will attract electrons more
strongly, or is the stronger oxidant
o When a solid, pure metal is in contact with a solution of another metal’s ions, the metal
with the lower electronegativity, or higher reactivity will displace the other metal in
solution, since it has weaker electron pull and loses an electron to the other metal to
become a cation
o Electrons are thus transferred from one metal atom to the other, one becoming a neutral
atom and depositing out of solution, and the other become a cation going into solution
e.g.
A granule of zinc is dropped into a blue solution of copper sulfate, zinc gets covered with reddish-
brown copper:
Zn( s ) + Cu 2+ ( aq ) 
→ Cu ( s ) + Zn 2+ ( aq )
Oxidation half-reaction:
→ Zn 2+ ( aq ) + 2e −
Zn( s ) 
Reduction half reaction:
Cu 2+ ( aq ) + 2e − 
→ Cu ( s )
The anion is a spectator ion. The activity series can be seen on the table of standard potentials
• Account for changes in the oxidation state of species in terms of their loss or gain of
electrons
1. The oxidation of an element in its stable elemental state is 0

2. The sum of oxidation states in an element or compound = 0, and for a polyatomic ion the
charge of the ion
3. An increase in oxidation state means a loss of electrons and vice versa.
Examining oxidation states is useful in determining whether a redox reaction has occurred during a
chemical reaction. Some elements display multiple oxidation states.
e.g.
In Cu2O (2Cu+, O2-)
Oxidation state of copper is +1
• Outline the construction of galvanic cells and trace the direction of electron flow
• Describe and explain galvanic cells in terms of oxidation/reduction equations
• Define the terms anode, cathode, electrode and electrolyte to describe galvanic cells
A simple galvanic cell requires:

o 2 solid metal pieces to serve as electrodes
o 2 containers (e.g. beakers)
o 2 different electrolyte solutions
o A salt bridge with electrolyte
o Connecting wire

1. Electrodes need to be matched with appropriate electrolyte, optimum metal ion same as
metal of electrode. Electrolyte must have greater or equal reactivity than electrode, else
displacement will occur
2. Salt bridge needs to be in contact with both solutions, and ions cannot form precipitate
otherwise there will be no charge neutralisation (causing opposing emf which impedes
current)
3. Conducting wire links both electrodes
The electrode are the conductors of a cell which get connected to an external circuit
The anode is the electrode where oxidation occurs
The cathode is the electrode where reduction occurs
The electrolyte is a substance which in solution or molten conducts electricity
Standard conditions are 1M solution and 25oC. Acidic conditions alter the potential difference.
Process:
1. The electrode with lower reactivity attracts electrons from the other electrode through the
conductor
2. A redox reaction results:
+
→ Cu 2+ ( aq ) + 2 Ag ( s )
e.g. Cu ( s ) + 2 Ag ( aq ) 
3. The ion formed goes into solution, and the anion dissolves. The metal ions in solution
around the cathode obtain the electron and plate onto the electrode
4. Ions flow from the salt bridge into the electrolyte solutions to neutralise charge and remove
opposing potential difference
Detailed purpose of salt bridge

If there is no salt bridge:
o As the redox reaction occurs, the electrolyte in contact with anode will have an
increasing excess of positive ions as the electrode dissolves, and the electrolyte in contact
with cathode will have an excess of negative ions as positive ions precipitate
o This imbalance of charges produces a potential difference against the direction of
electron flow, eventually stopping it
The salt bridge:
o Allows ‘migration’ of charge, and positive ions in the salt bridge move into the negative
solution, and vice versa
o This preserves electric neutrality, and eliminates negative potential difference
Cell diagrams
Type 1:
Metal/metal ion electrode
e.g.
Cu|Cu2+||Ag+|Ag
| = phase separator
|| = salt bridge
Reaction progresses from left to right
Type 2:
Inert substance such as platinum or carbon and equimolar amounts of non-metal and its ion.
e.g.
Pt(s) | I2(s)|I-(aq) || Fe2+(aq)|Fe3+(aq) | Pt(s)
Reaction progress:
2I-(aq)  I2(s) + 2e- (oxidation)

Fe3+ + e-  Fe2+(aq) (reduction)
What about gas?
• Perform a first-hand investigation to identify the conditions under which a galvanic cell
is produced
• Perform a first-hand investigation and gather first-hand information to measure the
difference in potential of different combinations of metals in an electrolyte solution
Prac – Galvanic Cells
Aim:
1. to construct a galvanic cell called a Daniell cell and investigate conditions under which it
operates (A)
2. to compare the effect of using different combinations of metals in electrodes (B)
3. to compare the effect of different volumes of electrolytes under otherwise identical
conditions (A)
Method A:
1). A zinc strip and copper strip were cleaned with sand paper
2). A half cell consisting of a copper strip resting in a 250 mL beaker half-filled with copper sulfate
solution was constructed
3). A similar cell with a zinc strip and zinc sulfate solution was constructed, and these half-cells
linked with a piece of filter paper soaked with potassium nitrate solution and folded
4). The copper strip was connected to the positive terminal of a voltmeter, and te zinc strip to the
negative terminal
5). The reading and polarity were recorded before quickly disconnecting the voltmeter
6). The beakers were then completely filled with corresponding electrolyte solutions, and the
voltage measured again
Results:
Beakers half: 0.20V Current: 0.5 mA
Beakers full: 0.25V Current: 0.8 mA
Anode: Zinc
Cathode: Copper
At anode (oxidation): Zn(s) -> Zn2+(aq) + 2e-

At cathode (reduction): Cu2+(aq) + 2e- -> Cu(s)
Overall: Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s)
The salt bridge allows migration of ions into each beaker (cations into cathode solution and vice
versa) to neutralise charge buildup and maintain cell voltage.
A larger volume of electrolyte solution means a large surface area of electrolyte in contact with
electrodes. This increases the rate at which charged particles are removed from the electrodes,
decreasing the internal resistance and thus the current. However, the voltage within the cell is
caused by the intrinsic properties of the electrodes (difference in electronegativity), and thus is not
altered significantly. It is affected slightly since a charge buildup generates a slight back emf, which
reduces the voltage, and more electrolyte action reduces this. Note: Concentration of electrolyte
does affect voltage.
The current output needs to be measured quickly since the ions in salt bridge are being used up, and
charge begins to build in half-cells opposing current flow.
Sources of error:
Electrodes not completely polished – reduces effective surface area of action for electrolyte and
thus current
Method B:
1). The copper half-cell in A was set up
2). Another half-cell consisting of a magnesium strip in magnesium sulfate solution was set up
3). The electrodes were connected to the terminals of a voltmeter, and voltage reading recorded
4). Steps 2-3 were repeated with:
a. Aluminium in 1M aluminium nitrate solution
b. Tin strip in tin(II) nitrate solution
Results:
Test Half-cell Polarity (relative to Total cell voltage
Cu/Cu2+)
Zn/Zn2+ -ve 0.50
Mg/Mg2+ -ve 1.1
2+
Sn/Sn -ve 0.35
Al/Al3+ -ve 0.90
Standard potential Cu – +0.34 V (Oxidation potential -0.34V)
Half Reaction (write in Predicted voltage E0 (V) Experimental E0 (V)
during exam)
Zn -0.76 -0.16
Mg -2.36 -0.76
Al -1.68 -0.56
Sn -0.14 -0.01
The more active the metal, the greater the potential difference.
Note: Oxidation potential is the ability of a substance to oxidise in relation to hydrogen, reductional
potential is ability to reduce in relation to hydrogen. E.g. Copper has reduction 0.34 meaning it has
higher ability to reduce, but its oxidation is -0.34, it has less ability to oxidise.
• Gather and present information on the structure and chemistry of a dry cell or lead-acid
cell and evaluate it in comparison to one of the following:
o Button cell (Silver Oxide cell)
In terms of:
o Chemistry
o Cost and practicality
o Impact on society
o Environmental impact
Criteria Dry Cell (Leclanche Cell) Button Cell (Silver Oxide Cell)

Structure
Chemistry Oxidation Oxidation

Zn(s) -> Zn2+(aq) + 2e- Amalgamated zinc
Reduction Zn(s) + 2OH-(aq) -> ZnO(s) +H2O(l) + 2e-
2MnO2(s) + 2H+(aq) + 2e- -> Reduction
Mn2O3(s) + H2O(l) Silver oxide
Hydrogen ions provided by Ag2O + H2O +2e- -> 2 Ag(s) + 2OH-
ammonium Electrolyte KOH
NH4+<-> NH3(aq) + H+(aq) Overall
Overall Zn(s) + Ag2O -> ZnO(s) + 2Ag(s)
Zn(s) + 2H+(aq) + 2MnO2(s) ->
Zn2+(aq) + Mn2O3(s) + H2O(l)
Cost and Practicality Adv: Adv:
o Inexpensive o Compact
o Robust o Provides constant voltage over long
o Easy to store and use period of time, since solid reactants
Dis: and products have fixed
o Short life concentration
o Voltage not as constant o Overall long operating life due to
as silver button solid components
(comparison) Dis:
o Cannot deliver very high o Silver is expensive
currents o Not rechargeable
o Cannot be recharged
Impact on Society o First commercial o Has allowed efficient powering of
battery, made portable miniature devices e.g. watches,
electric devices possible hearing aids
o Used widely in toys,
torches, radios etc.
Environmental o Manganese readily o Contains traces of mercury, causes
impact oxidised to stable problems with unsafe disposal
manganese (IV) dioxide, o Non-rechargeable so takes up space
becomes immobilised
and not dangerous
o Small quantities of
ammonium salts and
zinc not harmful
o Not rechargeable, large
volume in landfills so
space is an issue
• Solve problems and analyse information to calculate the potential Eo requirement of

named electrochemical processes using tables of standard potentials and half-equations
Eocell = Eocathode - Eoanode

Write down half-cell equations, then balance to get overall.
These values are under standard conditions (298K, 1M electrolytes)
A higher concentration of reactants relative to products increases spontaneity of reaction and thus
emf.
• Distinguish between stable and radioactive isotopes and describe conditions under which
a nucleus is unstable
o The spontaneous emission of radiation by certain elements is called radioactivity

o Some elements have all isotopes radioactive, some only one or some
o These particles are thus referred to as radioisotopes
Stable isotopes Unstable isotopes

No radiation emission Emission of radiation
Z ≤ 83 Z > 83
Ratio of neutrons to protons within zone Ratio of neutrons to protons out of zone
of stability of stability
Zone of stability:
A nucleus is unstable when its ratio of neutrons to protons

is outside zone of stability. For light elements (Z < 20),
1.0. Stability ratio steadily increases as atomic number
increases, up to 1.5 for Z = 83. Past this, all are unstable
due to large size of nucleus.
Mode Emission Atomic mass Atomic

number
α decay 4 -4 -2
2 He
β decay 0 0 +1
−1 e
Positron 0 0 -1
1e
emission
Electron 0 0 -1
−1 e (absorption)
capture
γ emission 0 0 0
0γ
• Describe how transuranic elements are produced
Transuranic elements are elements with atomic number above Uranium (92)
o Some isotopes undergo fission when bombarded, others undergo nuclear reactions to
form new elements
o When non-fissionable atoms such as Uranium 238 are bombarded with high speed
particles, it absorbs the particle to become an unstable atom
o It then rapidly decays to form a new element
There are two machines that are used to produce high speed positive particles to produce
transuranics.

o Linear accelerator – accelerates positive particles in straight line along axes of series of
positive and negative cylinders, accelerating it. Often more than a kilometre in length
o Cyclotrons – accelerates positive particles by passing them through alternating positive
and negative electric fields. A strong magnetic field is used to constrain particles to spiral
path, reducing size of machine.
They can also be produced in nuclear reactors, a source of neutrons. Neutrons do not experience
electric repulsion like positive nuclei, and thus speeds in nuclear reactors are adequate. These create
transuranic elements with a proton deficiency.
e.g.
238 1
→ 239
92 U + 0 n  92 U →−10 e+ 239
 93 Np
Neptunium, first discovered transuranic element obtained in by chemical separation of nuclear

fission reactor products.
This can be further bombarded to create plutonium:
239 2
→ 241
93 Np + 0 n  → 241
93 Np 
0
94 Pu + −1 e
Unstable
238 12
→250
92 U + 6 C  98 Cf → 246
 1
98 Cf + 4 0 n( )
23 transuranic elements have been created thus far.
• Process information from secondary sources to describe recent discoveries of elements
Recently discovered elements include:

o Ununoctium (118, October 10 2006) – heaviest element discovered to date. It was
indirectly detected by a team of researchers working in Russia at Dubna University’s
Joint Institute for Nuclear Research when they detected its decay products after
bombarding californium-249 atoms with calcium-48 ions. Very unstable, half-life
0.89ms. Reaction:
o Ununpentium (115, February 2 2004) – Russian scientists at Dubna “…” and

American scientists at Lawrence Livermore National Laboratory announced they
produced 4 atoms of Uup which quickly decayed into Ununtrium (113) in about 100 ms.
They bombarded Americium with calcium.
243 48 288
95 Am + 20 Ca 
→115 Uup +301 n
Note these elements are temporarily using IUPAC systematic element names, before they are
officially named.
• Identify instruments and processes that can be used to detect radiation
o Photographic film – photographic film darkens in the presence of radiation. Used in

radiation badges worn by laboratory workers handling radioactive substances to
determine radiation dosage
o Geiger-Muller tube – Radiation enters window and ionises gas particles inside Geiger
tube (inert gas such as argon), and the resulting charged particles are accelerated to the
two plates with a potential difference. They further ionise other argon atoms through

collision creating a cascade effect. This creates a signal which is amplified and converted
into an audio signal.
o Scintillation counter – ionizing radiation hits the scintillation crystal (depicted), the
electrons are excited and emit photons which can be detected and amplified by a
photomultiplier tube (depicted) to produce a reading.
o Cloud Chamber – contains supersaturated vapour of water or alcohol. Radiation (alpha

or beta particles) ionises it, forming noticeable tracks. Alpha trails are broader and
straight, whilst beta tracks are thinner and show more evidence of deflection
• Describe how commercial radioisotopes produced
Radioisotopes can be produced by bombardment of high-speed particles. Radioisotopes are

commercially produced in:
o Nuclear reactors (proton deficient, neutron enriched) – convenient source of
electrons. Target nuclei are placed in reactor core and are then bombarded by neutrons to
produce desired isotope. These are then separated chemically or physically from other
substances within reactor.
Currently operating in Australia for this purpose is HIFAR reactor, managed by ANSTO
e.g.
Creation of technetium:
98 1 99
42 Mo + 0 n 
→ 42 →9943
Mo  m
Tc + −10β
Technetium decays, releasing gamma ray inside body
o Cyclotron (proton rich, neutron deficient) – neutron deficient isotopes must be

produced in a cyclotron. They are bombarded with a small positive particle such as a
helium or carbon nucleus at high speed in order to overcome electrostatic forces of
repulsion
National Medical cyclotron

e.g.
Creation of gallium-67
68 1 67
30 Zn + +1 p 
→31 Mo + 201n
• Identify one use of a named radioisotope:

o In industry
o In medicine
• Describe the way in which the above radioisotopes are used and explain their use in terms
of their chemical properties
Technicium-99m (nuclear reactor)

Most widely used in medicine for diagnosis, such as locating brain tumours or studying other parts
of the body by being attached to red blood cells.
o Short half-life of 6 hours means patient exposure is minimised
o Versatile chemistry and can be incorporated into range of biomolecules targeting
different organs
Iodine-131 –
Produced in cyclotron or nuclear reactor, testing of thyroid function and treatment of thyroid
ailments such as overactive thyroid or thyroid cancer (beta decay destroys some thyroid cells)
o Iodine-131 is naturally absorbed by cells in the thyroid gland
o Relatively short half life to minimise exposure (8 days)
Specific problems:
o Ionising radiation of iodine-131 deals collateral damage to other cells
o Radiation penetrates the body and can damage organic tissue near to the patient
o Transport and production in nuclear reactors requires stringent safeguards, which is
problematic
Cobalt-60 – used to measure thickness of materials. With fixed geometry for source and detector,
penetration of radiation emitted from radioisotope (beta particles in this case) determines thickness
of material. Gamma ray producer
o Long half-life so does not need to be replaced frequently (5.3 years)
o Low energy emissions, so absorption is significant and can be detected
o Low energy emissions, minimises safety procedures required
Sodium-24
Used to detect leaks in water pipes or underground oil pipes. Dissolved into water source, can be
subsequently detected in soil around areas of leakage.
o Dissolves easily into water
o Relatively short half-life to minimise environmental damage (15 hours)
• Use available evidence to analyse benefits and problems associated with the use of
radioactive isotopes in identified industries and medicine
Benefits in medicine:
o Created wide range of non-invasive diagnostic procedures otherwise impossible, such as
technetium-99m used to identify brain tumours, gallium-67 for cancers, an area very
dangerous for surgery
o Allowed radiation therapy to treat many forms of cancer, e.g. iodine-131 for treatment of
thyroid cancer
Benefits in industry:

o Ability to make more sensitive, precise and reliable monitoring equipment e.g. cobalt-60
for measuring thickness of materials
o Allows otherwise difficult activities such as detecting leaks in extensive water
distribution systems, sodium-24 can be dissolved and radiation detected near leaks
Problems:
o Exposure of radiation doses to workers in medicine, industry and research can damage
tissues e.g. ionizing radiation of sodium-24 causes cancer by removing electrons from the
biological molecule DNA.
o Extra safety precautions are required for sites with radioactive materials, such as proper
storage facilities and protective clothing e.g. industries dealing with cobalt-60 and
technetium-99m need to filter out the fine radioactive dust produced, which can pose a
lung cancer risk
o Disposal of radioactive waste requires space, and can be problematic since isotopes such
as Cobalt-60 remain radioactive a long time after they are no longer useful, may leak into
environment without strict procedures
The Acidic Environment

Definitions and properties of acids/bases
An acid is a substance that produces hydronium ions (H3O+) in solution
A base is a substance containing oxide or hydroxide ions (O2-, OH-) or which in solution produces
the hydroxide ions. A soluble base is an alkali, i.e. one which dissolves or reacts in aqueous solution
to produce hydroxide ions. Note that oxygen ion-containers are insoluble or only react.
Common properties of acids:

1. sour taste
2. sting or burn the skin
3. conduct electricity in solution
4. turns blue litmus red
5. React with reactive metals to form salt + hydrogen gas
6. Reacts with carbonates to form CO2, salt and water
7. Reacts with metal oxides/hydroxides to form salt and water
Common properties of alkalis:

1. have a soapy feel
2. have a bitter taste
3. conduct electricity in solution
4. turns red litmus blue
5. React with amphoteric metals to produce hydrogen gas
Acids and bases react to form salt and water (there are exceptions). Note that the salt is in aqueous
solution, separated as ions and not precipitated. E.g. reaction of sodium hydroxide with
hydrochloric acid:
HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)
Can also be written as:
H+(aq) + NaOH(aq) - > Na+(aq) + H2O(l)

Or
H+ + Cl- + Na+ + OH- -> H2O(l) + Na+ + Cl-
The chloride and sodium ions are spectator ions. The net ionic equation is simply
H+ + OH- -> H2O(l)
A salt is an ionic compound formed when a base (alkali) reacts with an acid.
Hydrohalic acids such as HCl, HBr and HI lead to halide salts. Oxyacids (acids containing oxygen
attached to an element) e.g. sulfuric acid, nitric acid, phosphoric acid form salts that end in –ate.
Nitrous acid (HNO2) and sulphurous acid (H2SO3) create salts that end in ‘ite’. Anions formed from
oxyacids are called oxyanions. Must be familiar with acid formulas.
• Classify some common substances as acidic, basic or neutral
Acidic:
Vinegar (acetic acid) – used in cooking (~3)
Lemon juice (citric acid) (~2.5)
Vitamin C (ascorbic acid) – vitamin supplement
Hydrochloric acid – pH maintenance in swimming pools, clean bricks cement and tiles (~1)
Neutral:
Water
Salts (e.g. sodium chloride, copper sulfate)
Milk
Basic:
Baking soda (sodium bicarbonate) (NaHCO3) (~8.5)
Oven and drain cleaners (sodium hydroxide) – sodium hydroxide also used in soap, and alumina (~
13)
Lime (calcium hydroxide) – making mortar (~ 11)
Ammonia – used to make fertilisers (~ 12)
• Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol
blue can be used to determine the acidic or basic nature of a material over a range, and
that the range is identified by change in indicator colour
• Identify data and choose resources to gather information about the colour changes of a
range of indicators
An indicator is a substance that in solution changes colour depending on the pH of the solution.
There are many different indicators, and the range of pH over which these indicators change colour
varies. Litmus is the most common and is extracted from lichens. The indicator changes colour in
reaction with the pH of a substance, indicating acidity or basicity dependant on the range of the
indicator.
Universal indicator is a mixture of several indicators and works over the whole range
• Identify and describe some everyday uses of indicators including the testing of soil
acidity/basicity

o pH in soil testing – steps:
• small amount of moist soil placed in a well on a ceramic test plate
• White barium sulfate (neutral pH, insoluble) sprinkled on so colour change more
easily discerned
• 2/3 drops universal indicator added to barium sulfate
• Colour change of indicator compared to colour chart
Lime (CaO) or dolomite (CaCO3/MgCO3) added if too acidic
CaO + H2O  Ca(OH)2  Ca2+ + 2OH-
CO32- + H2O  HCO3- + OH-
Sulfur is added if too alkaline
S + O2  SO2
SO2 + H2O  H2SO3  H+ + HSO3-
o Testing home swimming pools which need to be neutral. Acidic water burns eyes,
alkaline water causes skin rashes. Operation of electrochemical cell to produce chlorine
makes pool more alkaline. A pool sample is collected into a vial and an indicator, usually
phenol red (yellow -> red) (6.6 – 8.0) is added and compared to a colour chart. HCl
added if too basic.
o Monitoring wastes from laboratories that process photographic film, as photographic
solutions are highly alkaline and discharges need to be neutral in order to not adversely
affect the environment
o pH in aquariums – fish excrete ammonia which reacts with water, making it more basic.
Universal indicator is used
• Perform a first-hand investigation to prepare and test a natural indicator
Prac – Red Cabbage indicator

Aim: to investigate the colour changes of an indicator extracted from red cabbage
Method:
1). A handful of shredded red cabbage was boiled in water over a bunsen burner for about 5 minutes
2). The resulting liquid solution was poured into a filter paper/filter funnel apparatus and collected
in a beaker
3). 20 mL of 1 M HCl solution was poured into a measuring cylinder
4). 10 mL was poured into a small test tube which was labelled ‘0’
5). The remaining 10 mL was poured into a beaker and diluted to 100 mL
6). 20 mL of the resulting solution was poured back into the measuring cylinder, and steps 4 – 5
were repeated 5 times, each test tube being labelled a successive integer higher
7). Steps 3-6 were repeated starting with 1 M NaOH solution and labelling from 14 down
8). A few drops of red cabbage indicator were added to each test tube and observations recorded
Results:
Red (0-1) Pink (2)  Purple (3-4)  Clear (5-6)  Purple (7-9) Blue (10-11)  Green (12-
13)  Yellow (14)
There are 6 discrete colour stages for this indicator. This suggests multiple molecules within the red
cabbage indicator solution acting to produce colour changes. Each molecule has a specific colour
when a proton is added or taken away. The molecules in this case are anthocyanins, and there are
about 15 different ones in red cabbage indicator.
• Solve problems by applying information about the colour changes of indicators to classify
some household substances as acidic, neutral or basic

See book
• Identify oxides of non-metals which act as acids and describe the conditions under which
they act as acids
An acidic oxide is one which either reacts with water to form an acid, or reacts with bases to form
salts (or both). E.g. carbon dioxide and diphosphorous pentoxide P2O5
CO2(g) + H2O(l)  H2CO3(aq)  2H+ (aq) + CO32-(aq) (carbonic acid)

CO2(aq) + 2NaOH(aq)  H2O(l) + 2Na+(aq) + CO32-(aq) (sodium carbonate)
Or alternatively:
H2CO3(aq) + 2NaOH(aq)  2H2O(l) + 2Na+(aq) + CO32-(aq)
The latter is more correct, as the acidic oxide would react with water to form the acid first. It would
depend on the relative concentrations of the oxide and the acid in solution, as it is an equilibrium
reaction. However, since both create the same products, this is negligible.
And similarly for P2O5
A basic oxide show basic character, and react with acids to form salts, but not with alkali solutions
e.g.
CuO + H2SO4(aq)  CuSO4(aq) + H2O(l)

CuO(s) + H2O(l)  Cu2+(aq) + 2OH-(aq)
Amphoteric oxides are those showing both acidic and basic character, and those that react with
neither acids or bases are neutral oxides e.g. NO, CO, N2O
• Analyse the positions of these metals in the periodic table and outline the relationships
between position and acidity/basicity of oxides
Acidic nature of oxides increases from left to right.
• Define Le Chatelier’s principle
If a system in equilibrium is disturbed, then the system adjusts itself so as to minimise the
disturbance
• Identify factors which can affect the equilibrium in a reversible reaction
Reversible reactions occur when products can react to generate reactants. When a reaction starts,
forward reaction generates products from reactants. Backward reaction then generates products,
which form at an increasing rate as product concentration increases. The equilibrium occurs at the
point where formation of products is equal to the rate of reactant formation, no net change in
concentration.
Factors:
o Concentration species – increasing/decreasing concentration of a species will cause
reaction equilibrium to shift so that it decreases/increases the species concentration. This
because it naturally results in more/less collisions or more/less decomposition to form
more/less of that chemical. Note that reactions involving solids and liquids experience
little effect, as concentrations remain almost unchanged (note: this does not include
dissolved substances).
o Pressure in a gaseous reaction – an increase/decrease will cause a increase/decrease in
concentration (and vice versa for volume). Depending on which side of the reaction has
more particles, the equilibrium will shift in that direction in order to reduce number of
particles and thus pressure (or vice versa). Note that increasing reaction by increasing
concentration of gas not involved in reaction e.g. argon has no effect
o Temperature – If the temperature is lowered, the amount of energy in the system
decreases and the exothermic reaction is favoured since less particles have sufficient
energy to form products with a higher potential energy. And vice versa
o Catalysts – increases speed at which equilibrium is reached, does not alter equilibrium
position as activation energy of both product and reactants formation is decreased
Notable exceptions:
o When solid or liquid is involved in reaction – the concentration of these substances stays
constant
o The addition of water to an aqueous reaction involving water – concentration of water
does not change significantly, but other substances more dilute
• Describe the solubility of carbon dioxide in water under various conditions as an

equilibrium process and explain in terms of Le Chatelier’s principle
Prac – Degassing soft drinks

Method:
1). A 300mL bottle of soft drink was taken and weighed on an electronic beam balance
2). The drink was shaken vigorously, and the top slowly and carefully unscrewed to prevent spillage
as the gas escaped
3). The bottle was reweighed
4). 6 grams of table salt was taken and added slowly to the contents of the bottle until no more
fizzing was observed
5). Bottle was reweighed
6). A new bottle was heated over a bunsen burner with cap on, then the cap unscrewed slowly to
prevent spillage before reweighing it
7).Hydrochloric acid was added until the water stopped fizzing, then it was reweighed
Discussion:
The dissolution and reaction of CO2 in water is multistep equilibria –
CO2(g)  CO2(aq) … (1)
CO2(aq) + H2O(l)  H2CO3(aq) … (2)
H2CO3(aq) + H2O(l)  HCO3-(aq) + H3O+(aq) … (3)
HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq) … (4)
Factors:

o Temperature – solubility decreases as temperature increases, opposite to liquids and
solids. Increased average kinetic energy of CO2 molecules means they have greater
overall tendency to escape from solution. Equilibrium shifts to left for all equations until
new equilibrium is reached
o Pressure – solubility increases with increased pressure, more carbon dioxide dissolves to
decrease pressure and act against change, equilibrium shifts to right.
o Dissolution of ions – dissolution of ions displaces carbonic acid ions and CO2 molecules
from hydration shells and causes equilibrium to shift to left and increase CO2 gas
concentration
o pH of water – increased pH means more hydroxide ions, which react with carbonic acid
to neutralise it and produce water, resulting in more CO2 dissolved to produce acid to
counteract change. If pH lowered, increased acidity means increased concentration of
H3O+ ions, shifting equilibrium of (3) and (4) to left to decrease its concentration. This
means increased concentration of the reactants on left, which has a cascade effect shifting
all equilibrium to left and increasing CO2 gas concentration.
• Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen
• Describe, using equations, examples of chemical reactions which release sulfur dioxide
and chemical reactions which release oxides of nitrogen
Natural sources of sulfur dioxide:

o Geothermal hot springs and volcanoes, release is unpredictable and changes with
volcanic activity. Natural levels of sulfur dioxide vary widely, but account for about ¼ of
worldwide emissions
o Bacterial decomposition of organic matter, produces H2S which oxidises to form sulfur
dioxide
2H2S(g) + 3O2(g)  2SO2(g) + 2H2O(g)
Industrial sources are mainly:

o Burning of fossil fuels – coal generally contains 0.5% to 6% sulfur as metallic sulfides in
sulfur in carbon-containing compounds, released as sulfur dioxide during combustion in
power stations. Some sulfur remains in refined petrol which is released as sulfur dioxide
in automobiles.
e.g.iron sulfide in coal
4FeS(s) + 7O2(g)  2Fe2O3(s) + 4SO2(g)
o Processing of fossil fuels - removal of sulfur from crude oil and natural gas releases some
sulfur dioxide.
o Extraction of metal from sulfide ores – first step is to roast sulfide ore in air e.g.
extraction of zinc roasting zinc sulfide
2ZnS(s) + 3O2(g) -> 2ZnO(s) + 2SO2(g)
Natural sources of nitrogen oxide/dioxide:

o Lightning – high temperatures causes atmospheric nitrogen and oxygen to combine
forming nitrous oxide:
O2(g) + N2(g)  2NO(g)
This slowly reacts with oxygen to form:
2NO(g) + O2(g)  2NO2
o Denitrifying bacteria – converts nitrates in soil into nitrous oxide (N2O), increased use of
fertiliser has increased emissions
Industrial sources:

o Combustion – includes power stations and automobiles etc. High temperatures involved
causes atmospheric oxygen and nitrogen to react, and is released into
atmosphere(equations same as above).
• Assess the evidence which indicates increases in atmospheric concentration of oxides of

sulfur and nitrogen
Nitrogen dioxide and sulfur dioxide are washed out by rain, so there is no significant buildup in
atmosphere. Nitrous oxide however, has steadily increased by about 15%, from measurements made
over the last century. There are problems associated with collecting evidence for sulfur and nitrogen
oxides, namely:
o Concentrations of both are very low, below 0.1ppm, and only recently (since about the
1950’s) are instruments accurate enough to reliably measure the levels, so trends before
this period could be invalid
o Sulfur dioxide and nitrogen dioxide form sulfate and nitrate ions which are changed
chemically as they move around the hydrosphere, so measuring traces of these
compounds is difficult
Most evidence comes from observed occurrences such as acid rain. There appears to be an increase
from data but it is inconclusive due to lack of long-term trends and inaccuracies of earlier
measurements.
• Analyse information to summarise the industrial origins of sulfur dioxide and oxides of
nitrogen and evaluate reasons for concern about their release into the environment
Concern about release into environment because it has detrimental effect on environment and can
cause harm to people since:
Sulfur dioxide:
o Sulfur dioxide irritates the respiratory tract, and can cause symptoms in people with
asthma or emphysema in concentrations as low as 1ppm
o Forms acid rain
o Dry deposition causing environmental damage
Nitrogen oxides:
o Nitrogen dioxide irritates the respiratory tract, and can cause extensive tissue damage in
concentrations 3-5 ppm
o The action of sunlight on nitrogen dioxide, hydrocarbons and oxygen increases
photochemical smog, which includes ozone – poisonous substance
o NO and NO2 participate in ozone layer depletion (NO + O3  NO2 + O2)
o Forms acid rain
o Contributes to global warming
• Explain the formation and effects of acid rain
o Sulfur dioxide and nitrogen dioxide gases released dissolve in water to form sulfuric acid
and nitric acid which is washed out of the atmosphere by rain, forming wet deposition
acid rain.
Reaction with hydroxyl radicals:
SO2(g) + 2OH  H2SO4(aq)
(OR)
2SO2(g) + O2(g)  2 SO3(g)

SO3(g) + H2O(l)  H2SO4(aq)

--
2NO2(g) + H2O(l) -> HNO2(aq) + HNO3(aq)

2HNO2(aq) + O2(g) –(catalysed by impurities) 2HNO3(aq)
(OR)
4NO2(g) + 2H2O(l) + O2(g)  4HNO3(aq)
Dry formation:
Incorporated into dust and smoke and falls to the ground
o Effects due to low pH include:

• Corrosion and tarnishing of metal and bridges, soiling and surface erosion of marble
and stone structures
CaCO3(s) + 2H+(aq) → Ca2+(aq) + CO2(g) + H2O(l)
• Crown dieback in trees

• Leeching of leaf nutrients
• Killing of leaf tissue
• Leeching of Ca2+ and Mg2+ ions from soil as they are mobilised due to decreased pH,
reducing soil fertility
• Inhibits microbial activity
• Increased acidity of lakes, killing aquatic life e.g. snails can only tolerate up to pH
6.0
• Mobilisation of Al3+ ions in soil due to reduced pH. This flows into lakes and
precipitates out, clogging fish gills and suffocating them
• Calculate volumes of gases given masses of some substances in reactions, and calculate
masses of substances given gaseous volumes, in reactions involving gases at OoC, 100 kPa
or 25oC and 100kPa
Equal numbers of molecules of different gases occupy the same volume in isothermal and isobaric
conditions. At 0oC and 100 kPa, 22.71 L per mole and 24.79 L/mol for the other.
• Gather and process information to write the ionic equations to represent the ionisation of
acids
• Define acids as proton donors and describe the ionisation of acids in water
Acids react with water in solution to form a solution containing hydronium ions and its conjugate
base.
• Identify acids including acetic acid, citric acid (2-hydroxypropane-1,2,3-tricarboxylic

acid), hydrochloric acid and sulfuric acid
o Acetic acid (CH3COOH) aka ethanoic acid – present in vinegar
o Citric acid (C6H8O7) – occurs in citrus fruit, also widely used as a

food additive for flavour or as a preservative
o Hydrochloric acid (HCl) – produced by stomach lining glands to break down food
molecules, also made commercially to clean metals, brickwork, neutralising bases etc.

o Sulfuric acid (H2SO4) – synthetic acid manufactured to make fertilisers, synthetic fibres
etc.
Phosphoric acid (H3PO4) weak
Polyprotic acids have more than one ionisable hydrogen per formula unit.
For others see book
• Describe the use of the pH scale in comparing acids and bases

• Identify pH as –log10[H+] and explain that a change in pH of 1 means a tenfold change in
[H+]
pH scale is a scale of measurement for hydrogen ion concentration.

pH = –log10[H3O+]
This obeys the significant figure rule
Water self ionises:

H2O + H2O  H3O+ + OH-
Kw = [H3O+][OH-] = 1.00 x 10-14 at 298K
• Describe acids and their solutions with the appropriate terms weak, strong, concentrated
and dilute
• Describe the difference between a strong and weak acid in terms of an equilibrium
between intact molecules and its ions
Note: In exams, define concentration and strength if used in question.
A strong acid is one in which all acid present in solution has ionised to hydrogen ions (no degrees
of strength), no equilibrium is formed. A weak acid is one in which only some of acid molecules
present in solution have ionised to form hydrogen ions, forming an equilibrium between intact
molecules and ions. The fraction of molecules ionised is called the degree of ionisation
(concentration of H+/concentration of acid originally).
• Plan and perform a first-hand investigation to measure the pH of identical concentrations

of strong and weak acids
Prac – Relative strength of acids

Aim: To compare the relative strengths of different acids using a variety of methods
Method:
1). 50mL of 0.1M hydrochloric, acetic, oxalic, citric and sulfuric acid were prepared in labelled 250
mL beakers
2). A pH meter was used in each beaker and the reading recorded
3). A pH strip was placed into each beaker for a short period and its colour compared with a chart
4). A few drops of universal indicator were dropped into each beaker and the colour compared to a
colour chart
Results:
#####
The strongest was sulfuric acid. It is a strong acid and is diprotic, meaning
that the concentration of H3O+ ions is twice the concentration of
hydrochloric acid. HCl is strong but monoprotic, meaning concentration is
identical to HCl concentration. Oxalic acid is diprotic and citric is triprotic,
but both are weak and do not completely ionise. The tendency for their conjugate bases to re-bond
with hydrogen ions limits the concentration of H3O+ in solution. Acetic is the weakest, being
monoprotic and have a low degree of ionisation.
Oxalic – C2H2O4
Citric – C6H8O7
Acetic/ethanoic – CH3COOH
Safety:
HCl - corrosive, vapour can burn mouth, throat and eyes
Oxalic acid – corrosive to tissue, corrosive to respiratory tract if inhaled
Wear safety glasses, goggles, use lower concentrations and smaller amounts
• Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric
acids and explain in terms of the degree of ionisation of their molecules
An acid is stronger than another if it has a higher degree of ionisation.
Citric is a triprotic acid, acetic is monoprotic. The degrees of ionisation are:

HCl – 0.010/0.01 = 1
Citric acid – 2.74 x 10-3/0.01 = 0.274
Acetic acid – 4.17 x 10-4 / 0.01 = 0.0417
HCl is the strongest acid and has the highest degree of ionisation.
• Use available evidence to model the molecule nature of acids and simulate the ionisation
of strong and weak acids
• Gather and process information to explain the use of acids as food additives
Acids are added to food to:

o Improve taste – e.g. carbonic acid in soft drinks, acetic acid in vinegar
o Preserve food – increases acidity to point where bacteria can no longer survive e.g.
coating freshly cut fruit with citric acid
o Increase nutritional value – e.g. adding ascorbic acid (Vitamin C, an antioxidant)
• Identify examples of naturally occurring acids and bases and their chemical composition
Acids:
o Ascorbic acid C6H8O6 – occurs widely in fruit and vegetables, essential to health
o Citric acid – found in citrus fruits
o Lactic acid CH3CH(OH)CO2H – produced by anaerobic respiration in cells, found in
muscle tissue and milk
o HCl – stomach to break down food
Bases:
o Ammonia NH3 – result of decomposition of proteins, or anaerobic decay of organic
matter found in fish urine
o Carbonates – e.g. calcium carbonate (limestone), magnesium carbonate
o Metallic oxides – e.g. Iron (III) oxide, copper oxide and titanium oxide, found in
minerals. Metals are extracted from them.
• Calculate pH of strong acids given appropriate hydrogen ion concentrations
• Outline the historical development of ideas about acids including those of:
o Lavoisier
o Davy
o Arrhenius
o Lavoisier (1780)
• acids were substances that contained oxygen
• Disproved since some oxygen-containing compounds such as metallic oxides were
basic, and distinctly acidic substances such as hydrochloric acid contained no oxygen
• Wrong but stimulated research
o Davy (1815)
• suggested that acids were substances that contained replaceable hydrogen.
• Bases were substances that reacted with acids to form salt and water.
• These definitions worked well for most of that century, but the definition made no
attempt to interpret the properties, only classify the substances
o Arrhenius (1884)
• Interpreted acidic properties in terms of ionisation to form H+, and weak/strong in
terms of degrees of ionisation
• the conductivity of acid solutions and their reaction with many metals to form
hydrogen gas evidenced that acidic solutions contained hydrogen ions
• acids were substances that ionised in solution to produce hydrogen ions
• A base is a substance that in solution produced hydroxide ions
• He defined strong acids as those that ionised completely and weak as those that
partially ionised. General equations:
HA(aq)  H+(aq) + A-(aq)
XOH(aq)  X+(aq) + OH-(aq)
Weaknesses were:
 Does not take into account role of solvent in ionisation of acid (ionisation
results from reaction of acid with solvent)
 Acid-base reactions can occur in solvents where there is no ionisation
 Not all acidic/basic substances (e.g. metallic oxides) ionised to produce
hydrogen/hydroxide ions
• Outline the Bronsted-Lowry theory of acids and bases
An acid is a proton donor, a base is a proton acceptor. Gives the broadest definition of acid/base
theory (it means that acids must have hydrogen). This definition :
o Does not restrict bases to those which ionise to produce hydroxide ions, such as in the
case of metal oxides and ammonia
o Explains how neutralisation reactions don’t require dissolution of ions into aqueous
solution e.g. NH3 + HCl in benzene (direct proton transfer)
o Exchange of proton relies on relative properties of both substances involved, accounting
for the role of the solvent
o Shows that hydrolysis of salts to change pH were acid or base reactions
o Provided basis for quantitative treatment of acid-base equilibria and pH calculations

• Trace developments in understanding and describing acid/base reactions
• Describe the relationship between an acid and its conjugate base and a base and its
conjugate acid
When an acid loses a proton, the resulting ion is called a conjugate base. If the acid is not strong,
this conjugate base can re-take a proton to reform the acid, resulting in an equilibrium reaction.
Vice versa for bases
• Identify a range of salts which form acidic, basic, or neutral solutions and explain their
nature
A salt is an ionic compound containing a cation not H+ and an anion not O2- or OH-. In aqueous
solution, salts completely dissociate into ions.
Type Acidic Basic Neutral
Salt Ammonium nitrate Sodium acetate Sodium chloride
(NH4NO3) (NaCH3COO) Potassium nitrate
Sodium hydrogen Potassium nitrite (KNO3)
sulfate (NaHSO4) (KNO2) Sodium sulfate
Anything Sodium carbonate (Na2SO4)
3+
containing Al , (Na2CO3)
Fe3+, HSO4- or Anything
-
H2PO4 containing F-, S2-
etc.
The pH of a salt solution depends on the nature of its ions, many cations/anions serve as acids or
bases. Some generalisations:
o Neutral salts have anions which are the conjugate base of strong acids, and cations the
conjugate acid of strong bases, since their reaction to accept/give protons is negligible
o Basic anions react with water to form hydroxide ions in solution. Reaction is equilibrium,
occurs to small extent since conjugate acid is stronger than water, and conjugate base is
stronger than basic anion
o Acid anions contain hydrogen atoms to react with water to form hydronium ions, derived
from polyprotic acids. The anion resulting from hydrolysis of polyprotic acids is
amphiprotic, and whether it is acidic or basic depends on the tendency for one hydrolysis
reaction (proton donation or proton accept) to occur over the other
Some examples:
(Basic anions)
S2–(aq) + H2O(l) ↔ HS–(aq) + OH–(aq)
F–(aq) + H2O(l) ↔ HF(aq) + OH–(aq)
(Acidic cations)
[Fe(H2O)6]3+(aq) + H2O(l) ↔ [Fe(OH)(H2O)5]2+(aq) + H3O+(aq)
• Perform an investigation to identify the pH of a range of salt solutions
Self explanatory. You could use NaCl + KOH for neutral, sodium bicarbonate (NaHCO3) and
sodium acetate (NaCH3COO) for basic, ammonium chloride (NH4Cl) for acidic.

Risk analysis:
Hazard Risk Control
Ammonium chloride Released vapour causes Wear safety goggles
coughing, shortness of Use small amounts to
breath minimise vapour
• Identify conjugate acid/base pairs
Acid Base Conjugate base Conjugate acid

HCl Cl-
H2SO4 HSO4-,(and SO42-?)
HNO3 NO3-
NH4+ NH3
OH- H2O
CN- HCN
CO32- HCO3-
CH3COO- CH3OOH
• Identify amphiprotic substances and construct equations to describe their behaviour in

acidic and basic solutions
Note: Hydrolysis is when a substance reacts with water
Amphiprotic substances can act as both a proton donor and proton acceptor. They react to both
accept protons and donate protons. Their behaviour changes whether in aqueous solution or
alkaline/acid solution.
e.g. HCO3- (hydrogen carbonate)

In aqueous solution:
HCO3-(aq) + H2O(l)  H3O+(aq) + CO32-(aq)

HCO3-(aq) + H2O(l)  H2CO3(aq) + OH-(aq)
In basic/acidic solution:
HCO3-(aq or s) + OH-  H2O(l) + CO32-(aq)
HCO3-(aq or s) + H+(aq)  H2CO3(aq)
Reactions go to completion since products cannot perform the reverse reaction (???). This also
applies to HSO3- (hydrogen sulfite) and HSO4-. Water is amphiprotic.
• Identify neutralisation as a proton transfer reaction which is exothermic
Neutralisation reactions are proton transfer reactions, and involve the reaction between an acid and
a base. They are exothermic and thus have a negative enthalpy change. The net ionic reaction in
Arrhenius theory is:
OH-(aq) + H+(aq)  H2O(l)
Acids and bases not fitting in Arrhenius theory do not necessarily produce salt and water during
neutralisation reactions e.g. neutralisation of ammonia. In LB theory, an acid and base react to form
conjugate base and conjugate acid. The acid gives a proton to the base. Reactions between strong
acids and bases form very weak conjugate acids/bases and go virtually to completion since the back
reaction has almost no tendency to occur. Otherwise, reactions are equilibria.
• Describe the correct technique for conducting titrations and preparation of standard
solutions
Volumetric analysis is a form of chemical analysis where the concentration of a substance is

determined.
Determining the composition of a solution require titration against another solution of known
concentration, called the standard solution. The substance dissolved is a primary standard.
Equipment:
A primary standard:
o Must be obtainable in very pure form and have known formula
o Should not alter weight unintentionally during preparation/titration e.g. absorbing
moisture from air
o Have a reasonably high formula mass to minimise weighting errors
o Purified by drying in oven and cooling in dessicator to eliminate moisture and prevent its
absorption
e.g. oxalic acid, sodium carbonate
Use of equipment:
o Pipette – solution to be used is first drawn in above mark, then solution let out until
meniscus at mark, solution let out through gravity with tip against wall of container
o Burette – first, rinse with portion of solution to be dispensed, overfilled then excess
allowed to run out
Preparation:
o Accurately measure mass of primary standard e.g. electronic beam balance
o Rinse a volumetric flask and beaker with distilled water
o Pour the primary standard into beaker and dissolve with distilled water, less than
intended volume of final solution
o Pour into volumetric flask, and repeat a few more times
o Use a pipette to add final few drops to complete solution
Titration curves:
Strong acid strong base:
Equivalence point at pH 7, steep curve. Indicator used should

have colour endpoint near equivalence point. Using indicator
changing during equivalence point is inaccurate, too difficult
to tell exact colour shade needed
Strong base/weak acid

e.g. NaOH, acetic acid

Equivalence point in basic range since salt formed is basic, the anion is a conjugate base of weak
acid and thus a weak base. Equilibrium reaction occurs, weak base reacts to reform conjugate acid,
thus decreasing acidity since less H3O+
Strong acid/weak base

Similar to above with equivalence point acid. Special case when CO2 formed during reaction e.g.
HCl + Na2CO3, CO2 forms carbonic acid.
Weak/weak
Not good since gradient around equivalence point is quite shallow, big volume difference between
indicator endpoints and equivalence points, needs to fit indicator very well or use one which
changes during equivalence, hard to distinguish.
• Perform a first-hand investigation and solve problems using titrations and including the
preparation of standard solutions, and use available evidence to quantitatively and
qualitatively describe the reaction between selected acids and bases
Prac – Titration
Aim: To standardise HCl and NaOH solutions using titration
Method:
Preparing standard:
1). A clean 250mL beaker was placed on an electronic balance, zeroed, and had 2.650g of Na2CO3
added using a spatula
2). Approx 100 mL of distilled water was added to the beaker, and solution stirred using stirring rod
3). A 250mL volumetric flask was rinsed with distilled water*
4). The Na2CO3 solution was poured into the volumetric flask, and step 2 repeated
5). The stirring rod/beaker were thoroughly washed using a wash bottle, the runoff dripping into the
vol. flask
6). Using a 25mL pipette, the volumetric flask was filled to the 250mL mark
7). A stopper was placed on the flask and contents swirled to mix
8). The pipette was rinsed with the unknown HCl*
9). 10 mL of unknown concentration HCl was poured using a pipette into a 50mL beaker washed
with distilled water*, and a few drops of methyl orange added
10). A burette was washed and filled with the Na2CO3 standard solution
11). The standard solution was quickly drained into the beaker to find an approximate end point
12). Steps 8-10 were repeated 3 times but more accurately
*
to ensure no cross-contamination
See book for calculations
• Qualitatively describe the effect of buffers with reference to a specific example in a

natural system
Buffer solutions resist changes in pH. It contains comparable amounts of a weak acid/base and its
conjugate base/acid. Take for example an acetic acid (CH3COOH) and sodium acetate
(NaCH3COO) system (acidic buffer).
CH3COOH(aq) + H2O(l)  H3O+(aq) + CH3COO- (1)

Addition of sodium acetate would increase the concentration of CH3COO – ions on the right. The
equilibrium shifts to the left, but due to the unchanged concentration of H3O+ ions, it stays enough
to the right for dissociation to cause a net increase in CH3COO – ions. This net increase enhances
buffering capacity. When hydronium ions are added to solution, the equation will shift to the left
according to Le Chatelier’s principle to reduce the concentration of H3O+ ions. When hydroxide
ions are added, the CH3COOH will react to form water and CH3COO-, reducing OH- concentration.
In both cases, pH change is reduced.
Buffer in natural system

Carbonic acid / bicarbonate ion buffer system in mammalian blood:
H2CO3(aq) + H2O(l)  HCO3-(aq) + H3O+(aq)
Maintains the blood at around 7.4 for optimum function, too high/low can result in death.
• Analyse … to assess the use of neutralisation reactions as a safety measure or to

minimise damage in accidents or chemical spills
o Many acids and bases are corrosive, can damage materials if spilt
Neutralisation reactions can:
o Reduce or nullify corrosive properties of spill, minimising damage
o Utilise common, cheap, safely handled/stored materials and produces relatively harmless
products e.g. sodium bicarbonate NaHCO3
Sodium bicarbonate is amphiprotic, so it can be used for both acidic and basic spills:
OH − ( aq ) + HCO3 − ( aq ) 
→ CO3 2− ( aq ) + H 2 O(l )
H + ( aq ) + NaHCO3 ( aq ) 
→ CO2( g ) + H 2 O(l ) + Na + ( aq )
Vinegar – commonly used in cooking, contains acetic acid:
OH − ( aq ) + CH 3COOH ( aq ) → CH 3COO − ( aq ) + H 2 O(l )

o Degree of reaction can be controlled by using different amounts of neutralising
substance, excessive amounts are wasteful and some
Overall, it is a useful, convenient and safe technique if used appropriately
• Describe the differences between the alkanol and alkanoic acid functional groups in
carbon compounds
• Explain the difference in melting point and boiling point caused by straight-chained
alkanoic acid and straight-chained primary alkanol structures
Strong intermolecular forces – higher BP/MP compared to similar mole mass (roughly similar
dispersion forces)
-OH
carboxylic acid group
Forms One hydrogen bond between Two polar bonds between molecules (C-
molecules O polar and C – O – H hydrogen bond),
higher boil/melt point
None ionised, neutral Small no. ionised, acidic
• Identify the IUPAC nomenclature for describing the esters produced by reactions of
straight-chained alkanoic acids from CI to C8 and straight-chained primary alkanols
from C1 to C8
Alkyl (e.g. methyl) alkanoate (e.g. formate, acetate, propanoate etc.)
• Identify esterification as the reaction between an acid and an alkanol and describe, using
equations, examples of esterification
Esters – carboxylic acids combined with alcohols, equilibrium reaction
Standard equation:
o H2O molecule released (H of alkanol  OH of acid)

o OR’ (R’ dummy variable) of alkanol  C of acid
H SO
e.g. CH 3COOH (l ) + CH 3OH (l ) ←2 
4 → CH 3COOCH 3 (l ) + H 2 O(l )
(acetic acid + methanol)
H 2 SO4
HCOOH (l ) + CH 3CH 2 OH (l ) ← → HCOOCH 2 CH 3 (l ) + H 2 O(l )
(formic acid + ethanol)
• Describe the purpose of using acid in esterification for catalysis
o The acid acts as a dehydration agent, removing water from the reaction (Le chatelier
argument), thus increasing yield
o Acts as a catalyst, lowering activation energy to speed reaction
o Only small amounts of acid required
o Most common is sulfuric, others include tosic, scandium (III) triflate
• Explain the need for refluxing during esterification
Reflux – the backflow of reactants into the reaction vessel

o Reactants and ester products can be volatile
o Reaction is carried out at high temperatures to speed reaction, causing evaporation of
alcohol and ester and thus loss of reactants/products
o Vapour is also dangerous as it is flammable and toxic
o To avoid loss and prevent diffusion, a cooled condenser is placed over the reaction
vessel, covering it
o The vapour condenses here and runs back into the reaction vessel, which
 Increases yield and saves on resources
 Allows the reaction to be carried at higher temperatures (faster)
 Prevents flammable gases from escaping
• Outlines some examples of the occurrence, production and uses of esters
• Identify and describe the use of esters as flavours and perfumes in processed foods and
cosmetics
Esters occur naturally, and are identified as fragrances and flavours in fruit and flowers e.g. orange
(octyl acetate). Animal fats such as butter, or oils such as linseed are also esters, as are waxes.
These fats and oils can be used to make soap
(Fat or oil) + sodium hydroxide → Salt of carboxylic acid + glycerol
The salt of carboxylic acid is the cleaning factor of soap

Artificially produced esters:
Aspirin:
o Acetylsalicylic acid active ingredient, discovered by Hoffman in 1897
o C7 H 6 O3 + C 4 H 6 O3  → C9 H 8O4 + C 2 H 4 O2
Salycilic acid + acetic anhydride
o Most widely used pain-relieving drug, e.g. headache, and prevents blood clots
Ethyl acetate:
o Solvent in industry, nail polish remover
o Acetic acid and ethanol
C2H5OH + CH3COOH → C4H8O2 + H2O
Octyl acetate:
o As a food flavouring such as in sweets, ice cream etc.
o Acetic acid and octanol
C8H17OH + CH3COOH → C10H20O2 + H2O
• Perform a first-hand investigation to prepare an ester using reflux
Prac – Esterification
1). 15 mL of acetic acid, 15 mL of butan-2-ol and 10 drops of conc.
Sulfuric acid was placed into a 100 mL round-bottom flask
2). The apparatus was set up as shown
3). The bunsen burner was lit and reflux allowed to occur for about
15 minutes until two layers clearly visible
4). A separating funnel was assembled on a retort stand and 100 mL
of distilled water poured inside
5). Contents of round-bottom flask were transferred into funnel and
the mixture shaken
6). 50mL of sodium carbonate solution was added and gently
shaken, occasionally inverted and tap opened to release gas
7). Once the layers separated, the lower layer was discarded
8). Step 6-7 was repeated
9). The remaining liquid was poured into a conical flask and a
teaspoon of CaCl2 added, shaken gently then allowed to stand for a
few minutes
CH3COOH(l) + C4H9OH(l) (H2SO4 conc.) CH3COOCH2CH2CH2CH3(l) + H2O(l)

(diagram representation here)
H2SO4 (see above)

Na2CO3: neutralise remaining acid, excess forms distinct layer so easily separated
CaCl2: anhydrous, absorbs water in mixture to leave ester
Safety:
1). Acetic acid + butan-2-ol flammable, apparatus should be secured to ensure no tipping; tipping
could pose fire risk due to contact with bunsen flame
2). Acetic acid corrosive to skin, avoid spilling
3). Condenser should have adequate water to prevent organic vapour escaping, flammable and
respiratory irritant
Chemical Monitoring and Management
• Outline the role of a chemist employed in a named industry or enterprise, identifying the
branch of chemistry undertaken by the chemist and explaining a chemical principle that
the chemist uses
Industry: Australian chemical manufacturing company

Branch: Analytical chemistry – quantitative and qualitative analysis of substances present within
materials
Roles:
o Occasional monitoring of ethylene quality, waste water (pH, suspended solids,
hydrocarbons etc.), gaseous emissions (particulates, toxic pollutants) to ensure reliability
of results by other chemists
o Check proper operation of equipment/calibrate instruments
o Train shift workers to use instruments
o Look for ways of improving processes
Principle: Adsorption in gas-solid chromatography
o Gas chromatography – a liquid or gaseous mixture is vaporised into a stream of helium
flowing over a stationary phase such as a solid
o If stationary phase is solid, the components of injected mixture adsorb (stick onto) its
surfaces to differing extents, and desorb at different rates
o This causes different substances to pass through at different rates
o A detector is able to quantitatively measure each substances as it passes out
o Can be used to determine chemical composition of substances
• Gather, process and present information from secondary sources about the work of
practising scientists identifying:
o The variety of chemical occupations
o A specific chemical occupation for a more detailed study
Many areas a chemist can work in, 13 divisions recognised by Royal Australian Chemical institute
including:
o Environmental chemistry (detailed) – determining how substances interact in the
environment, monitoring concentrations of substances especially in air, water and soil
 Environment monitoring, employed by Environmental Protection
Authority, mining companies, local government – qualification can be BSc
and postgrad qualification in fields such as scientific
communication/management. Could collect data on air/water quality, then
analyse and assess this information. Require strength in chemical analysis,
and instrumental analysis. May work in a team, providing environmental
advice to external bodies via reports.
o Physical chemistry – study and measurements of physical aspects of compounds and
reactions e.g. reaction rates, structure of substances, nature of chemical bonding

o Pharmaceutical chemistry – discovery, testing, synthesis and commercial development
of chemicals for use as medicines
o Industrial chemistry – chemistry of industrial processes such as manufacture of
ammonia, sulfuric/nitric acids and others
• Identify the need for collaboration between chemists as they collect and analyse data
o Chemistry is diverse, chemists specialise in particular branches as range of knowledge

too large
o Some real-world problems require expertise from more than one branch
o Collaboration is required for proper tackling of problem
o Also, work of one chemist may have implications in another area, requires active
communication skills
o This facilitates efficiency of scientific progress and scientific work
o Can be achieved by:
 Publishing of papers
 Collaboration between laboratories
 Direct voice communication
• Describe an example of a chemical reaction such as combustion, where reactants form

different products under different conditions and thus would need monitoring
e.g. Petrol combustion in motor vehicles (mostly octane)

Different situations/products:
o Complete combustion – excess oxygen, products carbon dioxide and water
2C8 H 18(l ) + 25O2( g ) 
→16CO2( g ) + 18 H 2 O(l )
o Incomplete combustion – insufficient oxygen, products carbon monoxide, carbon
dioxide, soot, unburnt hydrocarbons, water. Ensure adequate oxygen supply to fuel
o Nitrogen oxides – reaction of oxygen with atmospheric nitrogen due to high temps
forms nitric oxide and nitrogen dioxide (see acidic environment). Rhodium-platinum
catalyst converts more polluting gases into less harmful ones
2 NO( g ) + 2CO( g ) 
→ N 2( g ) + 2CO2 ( g )
2CO( g ) + O2( g ) 
→ 2CO2( g )
o Sulfur oxides – some sulfur compounds in fuels:
S ( 2 ) + O2 ( g ) 
→ SO2 ( g )
2 S ( s ) + 3O2( g ) 
→ SO3( g )
Monitoring can ensure minimum possible toxic chemicals released, important since:
o CO affects judgement/perception as levels as low as 10 ppm, can cause death by
asphyxiation
o Soot contributes to particulate pollution, bad for asthma sufferers
o Nitric oxide affects respiratory systems and is generally toxic, excessive production by
motor vehicles can affect health of population
o Sulfur oxides and nitric oxides contribute to acid rain
o Motor industry can use information to build more efficient engines (more complete
combustion)
• Identify and describe the industrial uses of ammonia

o Fertilisers – reacted with sulfuric acid to form ammonium sulfate or nitric acid to form
ammonium nitrate fertiliser. Application to soil provides good source of nitrates essential
for crop growth, improving yields. 2NH3 + H2SO4  (NH4)2SO4(s)
o Conversion to nitric acid – nitric acid is used in making explosives, dyes, fibres and
plastics
o Neutralisation of acid – petroleum industry uses ammonia to neutralise acid
components of crude oil and protect equipment from corrosion
o Water treatment – addition of ammonia and chlorine to water produces more stable
disinfecting residual than chlorine alone
• Identify that ammonia can be synthesised from its component gases, nitrogen and
hydrogen
• Describe that synthesis of ammonia occurs 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
• Analyse the impact of increased pressure on the system involved in the Haber process
• Explain that the use of a catalyst will lower the reaction temperature required and
identify the catalyst(s) used in the Haber process
N 2( g ) + 3H 2 ( g ) ←→ 2 NH 3( g ) ∆H = −92 Kj / mol
Nitrogen and hydrogen combine to form ammonia, which in turn decomposes to reform reactants
(reversible reaction). Equilibrium reached when rate of forward/reverse reactions the same.
At higher temperatures, the average kinetic energy of the reactants is higher. Thus:
1. Larger fraction of molecules have adequate energy to overcome activation energy and react
upon collision
2. Molecules move faster, more collisions between molecules
These increase the rate of reaction, and apply to both forward and backward reactions. However, it
affects forward reaction more.
Reaction is exothermic, higher temperatures shifts equilibrium to left to reduce temperature change.
Number of moles of gas on each side of reaction is different. According to Le Chatelier’s principle,
increasing pressure shifts equilibrium to right since there are less moles of gas, decreasing pressure
and thus minimising pressure change. Vice versa
Catalysts: iron-iron catalyst, small amounts of K2O, Al2O3

o Hydrogen and nitrogen molecules adsorbed onto surface, increasing collisions
o Allows reaction via new chemical pathways with lower Ea
o Reduces activation energy, allowing molecules with lower kinetic energy to react, thus
lowering the required temperature
o Reaction is faster, equilibrium unaffected
• Explain why the Haber process is based on a delicate balancing act involving reaction
energy, reaction rate and equilibrium
Balancing these factors to maximise yield and reaction rate is required to maintain adequate
production and use of resources:

o Higher temperatures means faster reaction rate but lower yield (equilibrium more to
reactants), and higher energy costs associated with temperature maintenance
o The reverse is true for lower temperatures, their advantages conflict, optimal temperature
currently used is 700K
o Higher pressures increases yield but places more stress on reaction vessel, current
optimum 2.5 x 104 kPa
o Catalysts can speed up reaction rate and lower reaction energy, lowering temperature
required for same production rate
• Explain why monitoring of the reaction vessel used in the Haber process is crucial and
discuss the monitoring required
o Temperature and pressure of reaction vessel – keep within range for optimum
production rate, excess temperatures can damage catalyst and lower yield
o Ratio of incoming reactants – maintain stoichiometric ratio and prevent buildup of one
reactant, slowing reaction
o Impurities in incoming gases – O2 can cause explosion, CO/CO2 can poison catalyst
and reduce its lifespan
o Rate of ammonia removal – inadequate rate of removal shifts equilibrium to reactants,
reducing yield
• Describe the conditions under which Haber developed the industrial synthesis of
ammonia and evaluate its significance in that time in world history
o Haber developed process in 1908, before WW1

o Growing population in early 20th century required large amounts of fertiliser to feed
population
o Growing militancy of Germany required more ammonia for explosives
o Haber process able to meet these demands
o Germany originally obtained nitrates as saltpere/guano from Chile, but advent of WW I
caused allies to set up a naval blockade, preventing imports, but Haber process allowed
Germany to be self-reliant
o This prolonged their resistance against the Allies, increasing the length of World War 1
and resulting in loss of many more lives
o Significant impact in world history
• Deduce ions present in a sample from a range of tests (Maybe need mixtures – go to
conquering chem for good summary)
Cations:
Cations:
Confirmation tests:
Ba2+ - apple green flame
Ca2+ - brick red flame
Cu2+ - blue-green flame, dissolves in

ammonia to form deep blue solution
Fe2+ - decolourises acidified dilute

KMnO4 solution
Fe3+ - deep red solution with SCN-

* Fe2+ can form white or green which may tun brown due to oxidisation over time (older samples)
* Copper forms CuI with iodine, not used in initial procedure
Note: Silver chloride is also a white precipitate, but silver sulfate is insoluble and colourless
Anions:
HNO3 Ba(NO3)2 Pb(NO3)2 AgNO3 Other
1
Carbonate Bubbles White - White pH between 8
CO32- precipitate precipitate, and 11
soluble in soluble in
HNO3 HNO3
*
Sulfate - (acidified) (acidified) - -
SO42- thick white *white
precipitate precipitate
Phosphate - (ammonia)* - Yellow (acidified)
3-
PO4 white precipitate (NH4)2MoO4,2
precipitate soluble in yellow
soluble in HNO3 precipitate
HNO3 (may need
warming)
Chloride Cl- - - - (acidified)  dissolves in
white ammonia,
precipitate darkens in
sunlight
1
CO32-(aq) + 2H+(aq)  CO2(g) + H2O(l)
*
Sulfate weaker lewis base than phosphate
SO42-(aq) + H3O+  HSO4-(aq) + H2O(l), equilibrium enough to left so adequate sulfate to produce
noticeable precipitation with barium/lead (CHECK – how would this effect lead?)
In basic conditions, enough phosphate for precipitation
2
12(NH4)3MoO4 + PO43- + 3H+  (NH4)3PMo12O40
• Perform first-hand investigations to identify the following ions:

o Phosphate
o Sulfate
o Carbonate
o Chloride
o Barium
o Calcium
o Lead
o Copper
o Iron
Some safety information:

o Barium compounds are toxic (see barium chloride below) (barium sulfate is mostly
harmless)
o Silver chloride – eye, skin and respiratory irritant
o Lead nitrate – poisonous if swallowed, causing spasms, nausea etc. Can be absorbed
through skin to cause irritation, redness over short periods
• Describe and explain the evidence for the need to monitor levels of one of the above ions
in substances used in society

Copper is an essential trace element for many organisms, such as humans. However, excess
amounts can be toxic and detrimental to the environment:
Sources of Copper Effects on People Effects on aquatic

ecosystems
Corrosion of copper Over consumption – Liver, Kills plants as well as
plumbing kidney damage (liver is algae (bioaccumulation)
storage point for copper)
(Chuttani et al., 1965)
Copper sulfate crystals to Over consumption - Reduces survivability of
control algal growth Vomiting, diarrhea, nausea aquatic invertebrates
(USEPA, 1980) (WSDOE, 1992)
Contact with skin – Reduces survivability of
eczema, edema of eyelids fish (Holland et al, 1960)
(Patty, 1963)
Death (NRC, 1977)
• Perform first hand investigations to measure the sulfate content of lawn fertiliser and
explain the chemistry involved
• Analyse information to evaluate the results of the above investigation and to propose
solutions to problems encountered in the procedure
Hazard Risk Control

Hydrochloric acid (if used) Corrosive. Can cause Wear safety glasses, use
permanent eye damage lower concentrations
Kills tissue
Barium Chloride Eye, skin and respiratory Wear safety glasses, use
irritant lower concentrations
Hot plate (if used) Can cause burns upon Gloves, keep body parts
contact away
Prac – Sulfate content of fertiliser

1). 1.0g of ammonium sulfate was placed into a measuring cylinder, and 2 ml of water was added;
mixture was shaken to dissolve
2). 10 mL of 1.0M barium chloride solution added to form white precipitate
3). This precipitate was allowed to settle overnight, then the clear liquid decanted
4). 10mL of water was added to precipitate to further dissolve chloride and ammonium ions,
mixture shaken
5). Steps 3-4 repeated twice
6). Two pieces of filter paper were weighed and then shaped into a cone
7). The barium sulfate slurry was evenly poured into this cone, then the cone placed on a source of
heat for an hour to evaporate remaining water
8). Barium sulfate + filter paper was weighed, results recorded and analysed
Results
Mass of ammonium sulfate: 1.000
Mass of BaSO4 precipitate: 0.655 g
96.06
Mass of SO42- = x0.655 = 0.270 g
233.36
27% sulfate ions (w/w) in fertiliser

Evaluation
The sources of error are:
1). Loss of barium sulfate due to slight solubility – using increased concentrations of reactants
would decrease its effect on results, but increases contamination by adsorption. Use between 0.005
and 0.05 mol/L of initial sulfate concentration
2). Adherence of barium sulfate to walls of the measuring cylinder
3). Incomplete drying of precipitate, contains water while weighed – repeated cycles of
drying/cooling/weighing until constant mass obtained
4). Remaining chloride and ammonium ions not completely removed by washing – more
repetitions, although this would cause greater loss of BaSO4 due to dissolution
5). Contamination by adsorption of substances in solution during precipitation – form precipitate
slowly by slowly mixing reactants, and forming in hot solutions maximises particle size and reduces
adsorption
6). Small size of barium sulfate crystals – some may have fallen through double filter paper, use of
a sintered glass funnel to trap it would result in more accurate results
OR
Weighed amount of agar added as coagulant
OR
Form in hot solution to maximise particle size + small amount of HCl
• Describe the use of AAS in detecting concentrations of metal ions in solutions and assess
its impact on scientific understanding of the effects of trace elements
• Interpret secondary data from AAS measurements and evaluate the effectiveness of this
in pollution control
Atomic absorption spectroscopy – used to measure low concentrations of elements in ppm range,
mainly metals.
Each element has unique emission spectrum, so by measuring, studying and using spectra we can
determine qualitatively and quantitatively the elements present in sample (by looking at spectra,
measuring intensity)
Practical arrangement + workings:

o Measures how much light of a specific wavelength is absorbed by sample being studied
o Sample to be analysed fed into flame, vaporises and converts molecules and ions into
atoms
o Atomic emission lamp producing a specific emission spectrum matching element to be
studied is passed through the flame
o Electrons absorb energy and are excited to higher energy levels
o Light is passed through prism to concentrate light of desired wavelength onto detector
o Wavelengths characteristically absorbed by element in question shows obvious drops,
indicating absorbance
o Absorbance linearly proportional to concentration
o This value is compared to a function (absorbance vs concentration) produced by
measuring known concentrations to find conc.
Why useful:
o Relies on absorption rather than emission, nearly 100% atoms in ground state absorb as
opposed to <0.1% excited which emit, can measure down to ppb
o Can measure concentration of only one element at a time
Uses:

o Detection of heavy metals e.g. copper, aluminium in waterways, since only found in trace
quantities
o Concentrations of micronutrients in soils
o Small amounts of contaminants in medicines and foods e.g. mercury, lead
o Elements in organisms e.g. blood and urine samples for mineral deficiencies
Trace element: Element present in concentrations < 100 ppm
Impact on scientific understanding:

o Allowed scientists to detect minute concentrations of trace elements, allowing them to
recognise its importance e.g. legumes were unable to grow in arid parts of Victoria until
AAS tests showed molybdenum deficiency
o Helped to demonstrate the importance of trace elements in living organisms, such as in
maintaining enzyme function e.g.
Evaluation on pollution control:

Useful since:
o Allows very small ppm concentrations of elements to be measured, (up to 0.01 ppm),
allowing detection of pollutants which may still be harmful in trace amounts
• Describe the composition and layered structure of the atmosphere
Layer of gas about 200 – 300 km thick surrounding the earth. Many different gases and distinct
layers which different characteristics:
Overall gas composition: 78% N2, 21% O2, 0.93% argon, trace amounts of CO2, neon, methane etc.
Layer Altitude above Most common gases Description
surface (km)
Troposphere 0 – 15 N2, O2, H2O, CO2, Ar Contains most of earth’s gases,
organisms inhabit this zone,
weather events
Stratosphere 15 – 50 N2, O2, O3, Contains ozone layer (25 km),
temperature increases with
altitude and gives stability
Mesosphere 50 – 80 Coldest layer (down to -100
celsius)
Thermosphere/ >80 Ions (O2+, NO+), O Temp rises with altitude, ionic
Ionosphere and atomic gas particles,
important in radio
communications since radio
waves reflect off
• Identify the main pollutants found in the lower atmosphere and their sources
Pollutant Source
CO Burning fossil fuels, forest fires
Airborne lead Lead smelters, leaded fuels
CFC’s Foaming agent, refrigerant-air conditioner coolant,
propellant
SO2 Combustion (fuel impurities), metal extraction from
sulfide ores, chemical manufacturing
Oxides of nitrogen (NO + Combustion (vehicles and power stations)
NO2)
Particulates Combustion, bush fires, industrial processes such as
mining
• Describe ozone as a molecule able to act both as an upper atmosphere UV radiation

shield and a lower atmosphere pollutant
Allotrope – a different physical form of the same element in the same phase
Ozone is:
o An allotrope of element oxygen
o Naturally present in atmosphere; only 0.02 ppm at ground level in clean air, 2 – 8 ppm in
stratosphere
o Detrimental in lower atmosphere – poisonous to many organisms; causes breathing
difficulties, fatigue and headache in humans
o Beneficial in upper atmosphere - filters our short wavelength UV light which can damage
living tissue
Produced in upper atmosphere through UV light:

3O2(g) -- (UV light)  2O3(g)
• Describe the formation of a coordinate covalent bond

• Demonstrate the formation of coordinate covalent bonds using Lewis electron dot
structures
Coordinate covalent bonds occur when shared the electrons come from one atom. Once formed,
identical to regular covalent bond.
• Compare the properties of the oxygen allotropes O2 and O3 and account for them on the
basis of molecule structure and bonding
Property O2 O3 Reason
Boiling point -193 -111 Polar bonds between
(oC) molecules means
intermolecular forces
stronger (O3)
Odour None Sharp, irritating
Colour None Pale blue
Density Slightly denser than 1.5 times denser than One more O atom per
air air molecule
Reactivity Highly reactive with Very highly reactive, The O-O bonds in O3
many metals and non- attacks double bonds are less strong than the
metals on alkenes double covalent bond
in oxygen
Solubility in Sparingly soluble More soluble than O2 O2 is non-polar, O3 is
water bent and thus is polar
Oxidation Lower Higher O involved only in
ability coordinate covalent
has greater electron
affinity

• Compare the properties of the gaseous forms of oxygen and the oxygen free radical
Free radical – a neutral species with an unpaired electron which can be formed by splitting a
molecule into two neutral fragments
Property O2 O
Reactivity Less reactive (full outer Very reactive (unpaired electron)
valence shell)
Oxidation Lower Higher (unpaired electron, high
ability tendency to take electrons to
complete valence shell)
• Identify the origins of CFCs and halons in the atmosphere
o CFCs (contain chlorine, fluorine and carbon) – developed as refrigerant in 1930’s to

replace ammonia, also used as propellant, solvent and foam blowing agent. Through use,
gas released into atmosphere
o Halons (carbon and halogens) – were used in fire extinguishers, recently use has been
drastically reduced
• Identify and name examples of isomers (excluding geometric and optical) of haloalkanes
up to eight carbon atoms
• Model isomers of haloalkanes using model kits
Haloalkane – hydrocarbon with one or more hydrogens replaced by halogen atoms (encompass
CFC’s)
• Discuss the problems associated with the use of CFC’s and assess the effectiveness of
steps taken to alleviate the problem
• Write the equations to show the reactions involving CFC’s and ozone to demonstrate the
removal of ozone from the atmosphere
• Identify alternative chemicals used to replace CFC’s and evaluate the effectiveness of
their use as a replacement for CFC’s
Problems:
Removal of stratospheric ozone
o CFC’s released into the troposphere are not washed out by rain (non-soluble) and not
destroyed by sunlight/oxygen at low altitudes
o Diffuse into the stratosphere and short wavelength UV breaks a chlorine off e.g.
CCl 3 F +UV 
→Cl • +CCl 2 F
o Chlorine reacts with ozone
Cl + O3 
→ •ClO + O2
o ClO reacts with free oxygen atoms
ClO + O 
→ •Cl + O2
o Net result is conversion of O3 and O to two O2
o Chlorine molecule unchanged at end, can continue to react and remove ozone
o This occurs on average a few thousand time before chlorine radical reacts with another
chemical which removes it e.g. methane
Cl + CH 4 
→ HCl + •CH 3

The problem:
o Removal of stratospheric ozone reduces filtering of short length UV radiation, meaning:
 Increased incidence of sunburn and skin cancer
 Increased damage e.g. brittleness to synthetic materials such as PVC
 Increased risk of eye cataracts
 Reduced plant growth for some species (e.g. rice) due to UV interference
with photosynthesis mechanisms
o CFC’s are greenhouse gases, and enhance global warming
Alleviation:
o Agreements to phase out use of CFC’s (e.g. Montreal Protocol, cease use in developed
by 1996)
o Agreements to phase out use of halons by 2010
o Assistance to poorer countries to phase out CFC use
o Replacement with safer alternatives
 HCFC’s – contain C-H bonds decomposable by radicals and atoms in
troposphere and are decomposed to significant extend. However, ozone-
destroying capacity is still significant (phase out by 2030 – Montreal
Protocol). Only useful as temporary substitute
 HFC’s – no C-Cl bonds, do not form Cl atoms in atmosphere, no ozone-
destroying capacity. Useful as permanent substitute, but more expensive
 Air being used to replace as foaming agent
Assessment:
o Adherence to agreements will ensure ozone layer returns to pre-CFC state since damage
is reversible
o The pace of CFC withdrawal means it will be many decades before the above happens,
meaning the effects of increased UV radiation will be felt
o Relies on co-operation of countries, will be less effective if countries withdraw
o Replacements such as HFC’s are more expensive than CFC’s, may be a burden on lower
countries until better alternatives found
Overall, it is an effective long-term solution but prolongs problems in the short term. It relies
heavily on cooperation of countries and this may be a downfall.
• Analyse the information available that indicates changes in atmospheric ozone

concentrations, describe the changes observed and explain how this information was
obtained
Evidence/information CHECK:
o Measurements of total ozone in a column of atmosphere have been conducted since 1957
o In springtime of 1980-1984, a severe depletion of ozone above Antarctica was detected
by the British Antarctic survey
o By 1985 it was approximately 30%, and in some places it had been completely destroyed
o A net decrease of 3% per decade was recorded for the period 1978-1991, factoring in
natural variations
Method of collection:
Data collected by range of ground and airborne instruments
o Dobson spectrophotometer (up to 48 km, groundbased)
 Developed in 1924, only source of long-term data

 Can be used to measure both total column ozone and profile ozone, currently
used to calibrate measurements by other methods
 Measures the intensity of four different wavelengths of UV radiation
reaching it; two are strongly absorbed by ozone, the others are not
 The ratio between the two intensities is determined and used to calculate total
ozone
Disadvantages/advantages
 Strong affected by aerosols and pollutants
 Measures only over a small area
o LIDAR (10 – 50 km, groundbased)

 Relies on absorption of laser light by ozone
 Telescope used to collect UV light scattered by two laser beams, one which is
absorbed by ozone (308 nm) and one which isn’t (351 nm)
 By comparing these values, a profile of ozone concentration vs altitude can
be measured
o Balloons (up to 40km, airborne)

 Various instruments can be mounted onto balloons
 Electrochemical concentration cells measure current produced by chemical
reactions with ozone
 Photospectroscopy utilises film or electric sensors sensitive to UV light to
measure wavelengths affected by ozone
 Can provide many days of continuous coverage
 Inexpensive
 Unpowered, flight path cannot be controlled
o TOMS (from space)

 Observes incoming solar energy and backscattered UV radiation at six
different wavelengths
 Gas molecules in the atmosphere scatter some EM radiation back, while
some is absorbed by ozone
 By comparing the intensity of backscattered radiation to incoming radiation,
amount of ozone can be obtained
 Provides global coverage
 Constant, accurate coverage
 Coverage in variety of weather and geophysical conditions
 More expensive than other methods
• Identify that water quality can be determined by considering:

o Concentrations of common ions
o Total dissolved solids
o Hardness
o Turbidity
o Acidity
o Dissolved Oxygen and Biochemical oxygen demand
Water quality is the chemical, physical and biological characteristics of water, with respect to its
intended purpose.

o Concentrations of common ions
1. Cations include Na+, Mg2+, Ca2+, K+, Al3+ and heavy metals such as Hg2+, Pb2+,
Cd2+ and arsenic. Common anions are Cl-, SO42- and HCO3-
2. Hardness (e.g. calcium, magnesium), detract from taste and appearance (iron),
Choke marine life(Al3+), toxic to humans (Hg2+, Pb2+), promote eutrophication
(nitrates, phosphates) which forms undesirable decomposition products e.g.
ammonia and detracts from taste/appearance
3. Water percolating through soil or underground estuaries, agricultural runoff
4.
a. AAS used to analyse cations
b. Gravimetric analysis used for anions e.g. precipitating Cl- as AgCl
o Total dissolved solids

1. Mass of solids dissolved in unit volume of water (ppm)
2. High TDS reduces crop and plant growth, >500 ppm is not suitable for human
consumption, >1000 ppm is unsuitable for irrigation
3. Underground aquifers, flowing through farming/grazing areas with disturbed soil
4.
a. Evaporation, but it produces very small masses of solid, easy to lose
through turbulent bubbling/spitting (use large volumes of water, 1 litre
round-bottom flask is suitable)
b. Conductivity – most solid dissolved substances are ionic and conduct
electricity; can be measured by conductivity meter to determine dissolved
solids
o Hardness
1. Concentration of Ca2+ and Mg2+ ions in water (ppm CaCO3)
2. Hard water forms a precipitate with soap, reducing cleaning power. Under high
temperatures such as in kettles, Ca2+ forms an insoluble precipitate with sulfate
and carbonate ions, reducing kettle efficiency
3. (see below)
4.
a. Titration with EDTA, which forms stable complexes with these ions.
Indicator Eriochrome Black T. In solution is blue but forms red coloured
complex with Mg2+. Endpoint when it turns blue, indicating no more Mg2+
in solution
o Turbidity
1. Measure of suspended solids in water
2. Undesirable appearance and taste, reduces sunlight penetration for plant
photosynthesis, can absorb IR and raise water temperature
3. Clay, silt, plankton, industrial wastes
4. Secchi disk, visually
o Acidity
1. pH of water
2. At extreme ranges, it can reduce survivability for aquatic organisms
3. Decomposition of organic matter, acid rain, exposure of sulfide ores in mining,
fertiliser run-off
4.
a. Universal indicator solution or paper
b. pH meter
o Dissolved Oxygen (ppm), Biochemical oxygen demand

1. Concentration of dissolved oxygen in water / capacity of organic matter to
consume oxygen
2. Regular levels of O2 (~10 ppm) indicate high quality water, since low levels
indicate high BOD and other factors such as:
a. Heat pollution which reduces O2 solubility
b. Excess of organic wastes such as sewerage which take up O2 in aerobic
bacterial decomposition
c. Eutrophication by excessive growth of aquatic plants
3. Dissolved from atmosphere, produced by photosynthesis by aquatic plants, algae /
Aerobic organisms such as bacteria, fish, worms
4.
a. (dissolved oxygen) Titration (see later)
b. (BOD) Addition of nutrient to sample and incubating at 20oC in sealed,
air-free container in dark for 5 days, then measuring residual dissolved
oxygen. Difference is BOD
• Identify factors that affect the concentrations of a range of ions in solution in natural
bodies of water such as rivers and oceans
o Pathway from rain to water body – rainwater collects ions before it runs into natural
bodies of water
• Bushland contains small amounts of nitrates and phosphates from natural nutrients on
surface
• Rainwater soaking into ground collects Ca2+, Mg2+, sulfate and chloride from soil and
rocks it flows through
• Percolation into deep underground aquifers results in collection of Fe3+, Mn2+ among
others
o Human activity
• Removal of natural vegetation or irrigation can increase salinity and thus NaCl in
rivers
• Agricultural fertilisers contribute nitrates and phosphates through runoff or dumping
• Discharge of sewage increases nitrates/phosphates, and various ions such as Cl-
• Acid rain caused by industry is better able to leech certain cations e.g. Ca2+ and Mg2+
from soil
• Motor car emission can increase lead
o Frequency of rain – more rain means more dissolved ions entering water bodies
o Bushfires – bushfires unlock nutrients and ions such as nitrates from plants, water picks
this during runoff
o Water temperature – higher water temperature increases evaporation and thus increases
concentrations of all ions in solution
• Perform first-hand investigations to use qualitative and quantitative tests to analyse and
compare the quality of water samples
Prac – Qualitative analysis

1). A sample of catchment water was taken and visually inspected for colour and hydrocarbons
(hydrocarbons produce rainbow effect on surface)
2). Some water was poured into a small conical flask and shaken for one minute, then observed for
bubbles which indicate detergent
3). Temperature of the water was recorded with a thermometer
4). A pH meter connected to a data logger was used to measure the pH of the solution
5). A turbidity tube was lowered into the sample until the bottom disappears, and reading recorded
6). Silver nitrate was added, and a white precipitate indicated the presence of chloride ions
7). A fresh sample was taken and Heavy metals (See below)
8). The above steps were repeated for tap water and distilled water
Results:
Distilled water Tap water Creek water

Colour Clear Clear Slight yellow
PO43- No No No
SO42- No No Slight
Hydrocarbons No No No
Turbidity (NTU) <10 <10 <10
Detergents No No Slight
pH 4.5 6 6
Heavy metals No No Trace
Cl- No Slight Lots
TDS 0 160 560
Bad odour mainly due to decaying vegetation in creek water.

Low pH can be due to ammonium fertiliser
High pH can be due to limestone and dolomite
Higher turbidity means lower intensity and quality of light for photosynthesis, lower DO and thus
aerobic aquatic organisms
Prac – Dissolved oxygen in water

1). A 200mL sample of water was taken in a conical flask, poured slowly to avoid aeration and
temperature measured with a thermometer
2). Using a graduated pipette, 2mL Mn(OH)2 and 2 mL alkaline NaI was added
3). The flask was stoppered and the solution inverted to mix
4). 4 ml of 1M sulfuric acid solution was added, then flask stoppered and inverted again to mix
5). 10 drops of starch indicator solution was added; this detects the presence of I2
6). A clean burette was filled with 0.02 M Na2S2O3, and the initial reading recorded
7). The water sample solution was titrated with the sodium thiosulfate, swirling constantly until the
blue colour completely disappeared
8). Steps 1-7 were repeated with tap water and another sample from same spot in creek
Results/Analysis
4 Mn2+(aq) + O2(aq) + 8OH-(aq)  2Mn2O3(s) + 4H2O (brown ppt)
Upon addition of acid
Mn2O3(s) + 6H+(aq)  2Mn3+(aq) + 3H2O(l)
2Mn3+(aq) + 2I-(aq)  I2(aq) + 2Mn2+(aq)
Redox titration:
I2(aq) + 2S2O32- -- (starch indicator)  2I-(aq) + S4O62-(aq)
Overall reaction:
O2(aq) + 4S2O32-(aq)+ 4H+(aq)  2S4O62-(aq) + 2H2O(l)
From data: 1mL 0.02M S2O3 = 0.527 mg/L DO

@ 15oC saturation = 10.1 mg/L
Creek Tap
2-
Vol. of S2O3 titrated (ml) 17.4 17.6
[DO]mg/L, % saturation 0.1698 (90.79) 9.2752 (91.83)

Very high creek saturation, probably due to recent precipitation which disturbs water and causes
aeration.
• Gather, process and present information on the range and chemistry of the tests used to:
o Identify heavy metals
o Monitor possible eutrophication of waterways
Heavy metals are transition metals plus lead which can be toxic to humans if ingested in higher
concentrations
Precipitation test (qualitative)

o Water sample acidified and Na2S added
o Precipitate formation indicates presence of one or more of Pb2+, Ag+, Hg2+, Cu2+, Cd2+¸
As3+
OR
o If no precipitate forms, water made alkaline then Na2S added
o Precipitate formation indicates Cr3+, Zn2+, Fe2+, Ni2+ …
o If precipitate forms when acidified, precipitate is filtered off and remaining solution
made alkaline to check for above ions
Chemistry:
o In solution, sulfide and hydronium react according to:
S 2− ( aq ) + 2 H 3O + ( aq ) ←→ H 2 S ( aq ) + 2 H 2 O(l )
When acidified, reaction proceeds enough to right for only minute concentrations of
sulfide left in solution
o Sparingly soluble ions such as lead can precipitate, while others such as zinc cannot
o Non-precipitation eliminates presence of these sparingly soluble ions
o In alkaline solution, enough sulfide remains to cause noticeable precipitation of ions such
as zinc and iron
o AAS
Eutrophication is the enrichment of water bodies by nutrients such as phosphate and nitrate in
excessive amounts (from agriculture, environment or discharged effluents which are decomposed
by aerobic bacteria), leading to algal blooms. This increases the BOD and decreases DO, reducing
survivability of aquatic life. Algal blooms block out sunlight for plant photosynthesis, further
reducing DO and aquatic life. Also develops unpleasant smells and unsightly appearance.
Colorimetry (quantitative) (phosphate)

o Method of measuring phosphate (a nutrient) in water, which corresponds with increases
algal growth
o Relies on absorption of light by a coloured solution
o Instrument used are colorimeters which use filters to select wavelengths of light, or
spectrophotometers which use monochromators to more precisely select wavelength
o Measured quantity of ammonium molybdate with catalyst added to sample and carefully
mixed, forming a coloured yellow complex with phosphate (phosphomolybdate)
o Measured quantity of ascorbic acid added, which reduces yellow Mo(VI) to intense blue
Mo(V)
o Colour absorbance is compared to that of standard solution to determine concentration
o Absorbance is proportional to concentration
Oxygen probe and data logger

o Aerobic oxygen decomposing organic waste into products including phosphate and
nitrate use oxygen
o Lower DO levels indicate higher microbial activity and thus higher nutrient levels in
water
o Oxygen probe lowered into sample, produces a current due to redox reaction and
datalogger converts this into a reading
Advantages of data logger:

o Can take readings over long period of time at regular intervals (reliable since larger
volume of data)
Rest is possibly useless:

o An oxygen probe consists of a cathode (platinum), and anode (usually silver) surrounded
by KCl solution
o A thin membrane permeable to O2 allows dissolved oxygen in near the cathode
o Since there is no easily reduced cation, the following reaction occurs:
O2 ( aq ) + 4 H + ( aq ) + 4e − ←→ 2 H 2O(l ) (cathode)
Ag ( s ) + Cl − ( aq ) ←→ AgCl( s ) + e −
o When electrodes are kept at constant voltage, constant separation and constant surface
area, the rate pf electrolysis is proportional to the concentration of DO
o This produces a current which is recorded by a datalogger and converted into a reading
• Describe and assess the effectiveness of methods used to purify and sanitise mass water
supplies
At the catchment:
o Preservation of natural environment – activities such as land clearing and
development cause increased TDS and turbidity; these activities prohibited around
catchments. Agriculture and industry can contribute heavy metals or nutrients from
fertilisers to water, catchments are distanced from these areas
This minimises the treatment that needs to be done at treatment plants
At treatment plant:
o Screening – large debris such as rubbish which can interfere with treatment is removed
Clarification - These processes remove turbidity and colour to give water optical clarity
o Flocculation/Sedimentation – aluminium sulphate or ferric oxide added to cause
precipitation of fine suspended particles, otherwise kept apart by surface charge. The
particles build as smaller particles adsorb onto them, coagulating to form large lumps. In
this process, dissolved particles can also become physically trapped. Resulting mixture is
settled in sedimentation tank and sludge removed
 Is able to removed dissolved and suspended particles
o Filtration – water is pushed through sand/anthracite filters which trap small particles
 Cannot remove extremely small particles
 Fast enough to produce adequate volume of water for big cities
i. Reducing dissolved organic carbon – improves taste and odour, allows water
to be safely chlorinated and makes it easier to further treat for domestic use

ii. Removes iron and manganese ions – improves taste, and eliminates stains on
laundries, fixtures and coatings on pipes
iii. Reducing phosphate concentrations – reduces algal blooms
iv. Reduces suspended particles – improves aesthetic appeal and makes it easier
to treat
v. Adsorbing organic matter – anthracite is able to remove odour and taste from
water
Sanitation – this removes anything harmful to human beings

o Disinfection – addition of a disinfecting agent to kill microorganisms, enough added so
residual amount remains to prevent further infection during distribution
 Addition of chlorine gas or liquid sodium hypochlorite, strong oxidising
agents – cheap and effective, but can fail e.g. Sydney
Giardia/Cryptosprodium scare in Sydney. Not effective against viruses
 Addition of ozone – stronger oxidant than chlorine and more effective against
bacteria and viruses but more expensive
(include membrane filters in a question)
Assessment:
1). Sydney water uses some of these methods – water out of tap is potable, clear and free of smell,
parasites such as Giardia and Cryptosporidium are kept at safe level
2). Water analysis has shown membrane filters produce water with less TDS and other
contaminants than traditional methods – method trialled in London with success
• Describe the design and composition of microscopic membrane filters and explain how
they purify contaminated water
o Thin film of synthetic polymer (e.g. polypropylene) or ceramic with microscopic holes of
approximately uniform size. Holes made by beam of ions through sheet of polymer, then
washing in alkaline solution. Changing ion changes pore size. Various forms include a
folded sheet which can be placed into pipes and used, or a bundle of hollow tubes – all of
which are very thin, so that water can flow through more quickly
Diagram:
o Reverse osmosis filters consist of cellulose
acetate or polyamid attached to another
polymer
o Water forced through the tiny holes by gravity,
vacuum or pressure pumps, which traps larger
particles and allows smaller particles through
o These trapped larger particles are water
contaminants such as bacteria and suspended
matter, whilst water molecules are small
enough to pass freely through
o This results in purer water
o The particles filtered depends on the size of the holes:
 Microfiltration – inorganic and biological particles included supsnded solids,
protozoans and bacteria
 Ultrafiltration – can remove fine suspended particles, viruses and water
borne parasites such as Giardia
 Reverse osmosis – can filter out viruses, bacteria, antibiotics and other
chemicals
• Present information on the features of the local town water supply in terms of:
o Catchment area
o Possible sources of contamination in catchment
o Chemical tests available to determine levels and types of contaminants
o Physical and chemical processes used to purify water
o Chemical additives in the water and the reasons for the presence of these additives
Warragamba dam – managed by SCA

Catchment area
Located 65 km SW from Sydney, narrow gorge on Warragamba river. Covers 9050 square km and
collects water from Coxs and Wollondily water systems, part of Sydney basic catchment area.
Possible sources of contamination

o Agriculture (fertilisers), industry (ions including heavy metals), commercial and
residential allotments (rubbish)
o Soil erosion which increases dissolved ions and turbidity
o Particulates from fire outbreaks
Chemical tests
o AAS for heavy metals
o Winkler method for DO/BOD
o pH testing using data logger
o Chloride ion titration
Physical and chemical processes to purify water

Main treatment occurs at Sydney Water treatment plants:
o Screening – sieve-like devices removes solid objects such as fish
o Coagulation – iron (III) chloride added to cause flocculation
o Filtration – water pushed through sand/anthracite filters to remove particulate matter,
cleaned by backwashing with water and air
o Microbial Treatment – Chlorine in the form of chlorine gas, liquid sodium
hypochlorite, and calcium hypochlorite tablets to kill microorganisms. Enough added so
residual remains throughout distributions. Sometimes, ammonia is added after chlorine in
fixed ratio to form monochloramine, which is less reactive and lasts longer
o Other chemical treatment
 Sodium silicofluoride and hydrofluosilicic acid are added to the water as
mandatory medication. 1ppm which reduces tooth decay and avoids
fluorosis.
 Lime and carbon dioxide, which react to form calcium bicarbonate, are added
at some treatment centres with soft water to increase resistance to pH change,
increase hardness and reduce corrosivity.
 KMnO4 added to oxidise manganese, converting to insoluble form and
filtered out
• Gather and process information from secondary sources to identify and analyse the
chemical composition of an identified range of pigments
• Analyse the relationship between the chemical composition of the metallic component(s)
of each pigment in the periodic table
Pigment Colour Composition Metal

Cadmium Yellow Yellow CdS Cadmium
Haematite Red Fe2O3 Iron
Azurite Blue Cu3(CO3)2(OH)2 Copper
Cobalt green Green CoO Cobalt

Zinc white White ZnO Zinc
Cinnabar Red HgS Mercury
Pigments all contain transition metals or cadmium and mercury - elements with valence electrons in
d-orbitals. These elements have non-degenerate d-orbitals, which have energy differences
corresponding to the visible spectrum of light. They absorb certain frequencies of incoming white
light, causing the reflected or transmitted light to be coloured as certain frequencies are missing.
• Explain that colour can be obtained through pigments spread on a surface layer (e.g.
paints) or mixed with the bulk of material (e.g. glass colours)
Paint – reflection of light

Glass – selective transmission of light
• Describe paints as consisting of the pigment and a liquid to carry it

• Explain why pigments used needed to be insoluble in most substances
Paints consist of a pigment dispersed in a liquid called a vehicle

Pigment:
o Solid material composed of very fine particles insoluble in vehicle
o Provides:
 Colour
 Hiding power (opacity)
 Strength and adhesion
 Gloss
o Must remain suspended in paint and not settle to form hard sediment
o May be organic, inorganic or metallic
o Lakes are organic pigments sourced from plants and animals e.g. cochineal (crimson)
from insects, tyrolean purple from marine snail
o They are soluble, don’t get depth and intensity of inorganic pigments, and are broken
down by UV radiation and fade
Vehicle:
o Can consist of:
 A binder which controls flow properties of coating and hardens on exposure
to air and gives bulk, gloss and toughness to paint film (natural oils, synthetic
organic compound, gum). Drying oils are polyunsaturated, polymerised due
to opening of C=C causing chain cross-linking. They physically lock in
pigments and are effective binders e.g. linseed/walnut oil
 Thinner which is a solvent to facilitate application
 Drier – a metal soap which speeds up drying of alkyl-based paint
o Affects drying time and thickness
o Needs to:
 Wet particles of pigment
 Sufficient viscosity to hold particles in suspension when drying
 Capacity to form tough adherent film on surface
Pigments need to be insoluble since they are not easily removed or affected by rain or perspiration
(makes them moisture resistant). Pigment needs to be opaque and reflective so colour is vibrant (in
many cases)

• Outline the processes used and the chemistry involved to prepare and attach pigments to
surfaces in a named example of medieval or earlier work
Madonna and Child with saints (By Sano di Pietro, medieval)
Support: Oak and pine

Ground (preparatory layer):
Linen glued to poplar with animal glue to protect from environment
Gypsum (CaSO4) or Gesso coats the linen to provide smooth white base
Coarse Gesso – coarse and thick, first layer
Smooth Gesso – smooth hard second layer made by soaking/slaking plaster in water
Underdrawing was made to outline the painting
Paint:
Pigments included: Red Lead (Pb3O4), Azurite (hydrated CuCO3, found in oxidised zones of copper
ore deposits), Yellow orpiment (As2S3), soot (carbon)
Binder was. Egg tempera method used, finely ground pigment mixed with separated egg yolk,
linseed oil and water. Layers applied thinly to prevent shrinking and cracking. Egg tempera holds
painting together and gives bulk
Resin (amber and natural tree resins) added as varnish
Other:
Attachment of gold leaf – graffito, gold leaf painted over and then design scratched to reveal it –
gilding.
Mordant gilding – thin lines of adhesive painted on surface, gold leaf cut and adheres to adhesive,
any excess is brushed away
• Identify the sources of the pigments used in early history as readily available minerals
• Identify minerals that have been used as pigments and describe their chemical
composition with particular reference to pigments available and used in art by Aborigines
Early pigments derived directly from naturally occurring coloured earth and soft rocks e.g.
Mineral Colour Chemical formula

Cerussite White, grey PbCO3
Stibnite Lead-grey Sb2S3
Cinnabar Red, brownish red HgS
Orpiment Lemon yellow As2S3
Ochres (natural earth of silica and clay), and kaolin were used extensively by the aboriginals:
Mineral Colour Chemical formula

Red Ochre (haematite) Red Fe2O3 anhydrous
Yellow ochre (limonite) Yellow Fe2O3 hydrate
Brown ochre Brown Fe2O3, partially hydrated
Kaolin white (Al2O3.SiO2.H2O)
Umbers (mixtures of manganese oxide and iron oxides) were natural brown clay pigments – also
used.
• Outline the early use of pigments for,

o Cave drawings
o Self-decoration including cosmetics
o Preparation of dead for burial

Cave Drawings:
Naturally obtainable pigments were used in the art of early cave drawings. Lascaux cave paintings
in France dating to 15,000 BC, and Aboriginal rock art such as Mimi and Bradshaw showed that
humans used:
o Yellow and red ochres (naturally occurring tinted clays)
o Black manganese dioxide (MnO2) (ore pyrolusite), or carbon from fires, or charred bones
o Kaolin (clay mineral) or chalk (CaCO3) (sedimentary rock) for white
They were ground into powder then mixed with blood, wax, saliva or cave water to create paint.
They were applied using the fingers or brushes made of kangaroo fur/a feather.
Self-decoration including cosmetics:

Aboriginals, such as Koori, used the above pigments as body paint for decoration, dance and rituals.
In PNG, Mount Hagen men painted their faces with black charcoal and white river clays.
The Egyptian, Roman and Greek cultures used pigments as cosmetics, mostly mixed with water or
saliva. Egyptians used kohl [made up variously of stibnite (Sb2S3), black manganese oxide (MnO2),
lead and CuO] as mascara after being wetted with saliva. Orpiment was used as yellow eye shadow.
All three cultures used white lead (PbCO3) for face paint. Greeks used vermillion (chemically same
as cinnabar): produced heating mercury and sulfur together in flask – taken out and ground. Henna
was also used in Egypt to dye fingernails, palms and soles, and hair.
Preparation of dead for burial:

Cro magnon man used red ochre in burial sites, place on chest and head, probably associating it
with life-giving blood.
Egyptians painted tombs with green, obtained from ground malachite (isolated from mineral by
grinding, washing and sifting). It symbolised rebirth and was significant to the afterlife.
Sarcophaguses were painted with Lapiz Lazuli, which were believed to give the dead person god-
like powers, helping them in the afterlife. Yellow and red ochre were used for paintings illustrating
the deceased’s family and slaves.
• Identify the chemical composition of identified cosmetics used in an ancient culture and
use available evidence to assess the potential health risk associated with their use
In Egypt:
Pigment Composition Use Hazard
Cinnabar HgS Rouge, lipstick Mercury is not harmful in pure form, but
Vermillion harmful if allowed to combine with oxygen
and hydrogen in water and air. Toxic by
ingestion and inhalation in large doses.
Results in numbness, staggered walk, tunnel
vision and brain damage
Galena Contains (PbS) Eyeshadow Harmful in small amounts. Most is removed
in urine, but there is risk of buildup.
Damages nervous system, causes mental
retardation and death
Yellow As2S3 Eyeshadow Highly toxic, can cause: vomiting; diarrhea;
Orpiment nausea; numbness
Malachite CuCO3.Cu(OH)2 Eye paint Can cause anemia, liver and kidney damage,
and stomach/intestinal irritation in high
doses
• Describe an historical example to illustrate the relationship between the discovery of new
mineral deposits and the increasing range of pigments

Titanium oxide (TiO2):
Historical example: titanium dioxide.
Before it was available, the principal white pigments used were calcium carbonate from shells,
bones, limestone and marble. These then moved to lead white, used by Ancient Greeks and
Egyptians, and European Ease paintings in C19th.
▪ 1791: discovered in ilmenite in England
▪ 1795: discovered in rutile – named “titanium” in Germany
▪ These Mineral sands were first mined in Australia, by 1960 mined in NSW, Qld, Vic
▪ Ilmenite is changed to synthetic rutile (more TiO2) by Becher process or newer chlorination process
to produce white pigment/very fine crystalline rutile.
▪ Mixed with other white pigments to produce different grades of whites.
▪ Now used widely in paint pigments, sunscreens, and cosmetics
These different pigments provided new shades of colours or new colours altogether.
• Describe the development of the Bohr model of the atom from the hydrogen spectra and
related energy levels to electron shells
• Explain why excited atoms only emit certain frequencies of radiation
• Explain what is meant by ‘n’, the principle quantum number
• Identify that, as electrons return to lower energy levels, they emit quanta of energy which
humans may detect as a specific colour
In 1901, Planck formulated the relationship between

frequency and energy. He proposed EM radiation was
transmitted in discrete units or quanta, called photons. He
showed that energy was proportional to frequency.
If a hydrogen lamp is subjected to high voltage, the
excited electrons will emit a violet light which can be
separated into component wavelengths using a prism.
Lymann (ultraviolet), Balmer (visible) and Paschen (IR).
Rutherford’s model did not explain this, it implied an
electron would move smoothly towards nucleus, releasing
continuous spectrum
Bohr applied Planck’s concept of energy quantisation to

explain the hydrogen spectrum. He proposed that
electrons move around the nucleus in a circular orbit
attracted by electrostatic forces without radiating energy, and that an atom could only have a
restricted set of discrete energy values (only orbits of certain radii/energies)
Bohr calculated a set of allowed energies using the hydrogen spectrum, and the principle quantum
number ‘n (a cardinal)’ denotes the energy level of a particular orbit. Energy is emitted or absorbed
by an atom when an electron moves from one stationary state to another. The difference in energy
between initial and final states is equal to the difference in energy between the initial and final
energy levels.
The lines in the hydrogen spectrum represent a drop from higher energy to lower energy (return to
ground state). Energy of photon is exactly equal to energy difference between levels. Some of this
radiation corresponds with radiation in visible spectrum.
• Solve problems and use available evidence to discuss the merits and limitations of the
Bohr model of the atom
Merits:
o Explained the observation that excited atoms generated discrete spectra
o Predicted with reasonable accuracy the emission spectrum of hydrogen
o Incorporated the idea of ‘quanta’ into a model of the atom
Limitations:
o When applied to other atoms, predictions failed to agree with experimental results
o Could not explain closely spaced emission lines, or the further energy splitting by
magnetic fields
o Did not explain why certain radii were permitted, or why moving electrons did not lose
energy
o Did not explain the different intensities of these lines
• Gather and process information from secondary sources to analyse the emission spectra
of sodium and present information by drawing energy level diagrams to represent these
spectral lines
Simplified sodium emission spectrum shows primarily

two yellow spectral lines at 589 and 589.6 nm
wavelengths, a ‘doublet’. This causes problems for
Bohr model since:
o Closely spaced doublet could not be

predicted by Bohr model, the energy
transitions between principal energy levels
could not account for this energy transition
o One yellow line is more intense than others,
and both these lines are far more intense than
other emission lines, could not be explained
by Bohr
o Magnetic field caused further splitting, could
not be explained by Bohr
o The quantum model of the atom, can account

for the spectrum through orbitals and orbital
splitting
• Identify Na+, K+, Ca2+, Ba2+, Sr2+ and Cu2+ by their flame colour
• Perform first-hand investigations to observe the flame colours above
Na+ - yellow
K+ - violet
Ca2+ - orange-red
Ba2+ - apple green
Sr2+ - red
Cu2+ - blue-green
Flame test procedure:

1). Piece of nichrome wire is cleaned by repeatedly dipping into HCl and heating it to red heat in
flame, this is to eliminate sodium which can give intense yellow colour masking other colours

2). Wire was dipped into aqueous solutions of samples and placed back into flame
3). Flame colour observed
Risk analysis:
Hazard Risk Control
HCl Corrosive. Can cause Wear safety glasses, have
permanent eye damage sodium bicarbonate to
Kills tissue neutralise spills
Bunsen flame Wire heated in hottest part Protective gloves
of flame, can cause burns
• Explain the flame colour in terms of electrons releasing energy as they move to a lower
energy level
The flame is a source of energy that can be absorbed by electrons in the atoms to be excited and
move to higher energy states. When they return to ground state, they release energy in the form of
EM radiation, which corresponds with frequencies in visible spectrum, producing colour.
• Distinguish between the terms spectral line, emission spectrum, absorption spectrum and
reflectance spectrum
o Spectral line – a discrete wavelength of light emitted by a radiant source, light emitted
by an atom is quantised, not continuous
o Emission spectrum – a set of discrete spectral lines corresponding to the energies
emitted by an excited atom, represented by bright lines against dark background
o Absorption spectrum – the specific wavelengths of light absorbed by an atom,
appearing as dark lines across a continuous spectrum. Complementary to emission
spectrum
o Reflectance spectrum – the wavelengths of light reflected by an element
These spectra can also be represented as intensity vs wavelength plots (quantitative)

Each spectrum is unique to an element
• Outline the use of infra-red and ultraviolet light in the analysis and identification of
pigments and their chemical composition
IR and UV are utilised in spectroscopy techniques

Spectroscopy: the production, measurement and interpretation of electromagnetic spectra
interacting with substances (emission, reflectance, absorbance)
Quantitative and qualitative analysis in a wide range of applications
Many techniques for analysis, includes:

Infrared (700nm to 1 mm):
Source of IR is commonly a heated ceramic e.g. silicon carbide rod.

Qualitative methods:
o IR reflectography (used for paintings, quantitative, USE THIS ONE SPARINGLY)
 Radiation penetrates through most outer pigments and reflects off white
background
 Carbon from deposited graphite, charcoal or black ink which the artist has
used to make a preliminary outline absorbs the radiation, forming a black
image
 IR also absorbed by copper-containing green pigments, can provide
information about their use
 Emission is detected by thermocouple
 Method is non-destructive (near infrared doesn’t damage it) and can be used
on whole artwork
Quantitative methods:
o IR absorption spectroscopy
 A double-beam spectrometer is used. One beam passes through the sample
and another through the reference
 A mono-chromator is used to select particular frequencies of IR to pass
through the sample and reference
 The molecules in a compound ‘vibrate’, and absorb radiation of same
frequency of their natural vibrational frequency, causing less radiation to be
transmitted
 The pattern of absorption is unique for each compound due to mass of atom,
and length/strength of bonds
 A detector compares the energy that is transmitted by the sample with that of
the reference – an absorption spectrum is plotted from the difference
 This absorption spectrum can be compared with standard absorption spectra
for certain functional groups and compounds, allowing identification of
pigments used
o ATR (Attenuated total reflectance) (Reflectance type)

 Used to examine samples in liquid or solid state without need for further
preparation
 For a solid sample, a crystal (e.g. KBr) is clamped tightly to the solid, and an
IR beam is directed through the crystal to the sample
 Some is absorbed and some reflected
 The beam is reflected internally within
the crystal and is collected by a
detector upon exit
 The resulting reflectance spectra is
compared with standard sample graphs
for compound identification, allowing
determination of pigments used
 Technique similar for liquids, but
liquid is poured onto crystal instead of clamped
 Used for analysing coatings, pastes and paints
Ultraviolet (300 nm to 10 nm):

UV spectrophotometer, with radiation source a tungsten lamp or deuterium discharge tube
(and halogen lamp for the visible region in UV/visible).
o Ultraviolet visible spectroscopy (absorption + reflectance)
 UV light is shined onto sample, and the reflected beam is collected by a
detector (photomultiplier tube)
 Absorbance is directly proportional to concentration, thus this can be used
quantitatively
 Compared with reflectance spectrum of non-absorbing substance such as
silica (SiO2) (this is the reference)
OR for absorption spectra of pigments in solution:
 Solution is placed into a silica cell
 The absorption spectrum of a reference cell containing only solvent is
determined
 The UV light is shined onto sample, and a detector compares the radiation
passing through the reference and sample
 Absorbance spectrum of sample is determined, and this is used to determine

the pigment, its concentration and chemical composition by comparing it
with absorbance spectra for different compounds
o UV fluorescence (qualitative, USE SPARINGLY)

 Different materials can fluoresce when exposed to UV light, the colour being
dependant on chemical composition and age
 ZnO (zinc white) fluoresces a pale yellow, while green malachite is a dirty
mauve
Can be used for paintings, pigments dissolved in solution etc.
• Explain the relationship between the absorption and reflectance spectra and the effect of
infra-red and ultraviolet light on pigments including zinc oxide and those containing
copper
o The reflectance and absorption spectra are complements of each other when sample is
opaque
o Represented as plots of wavelength against intensity of reflection/absorbance
o A material can either reflect or absorb a particular wavelength
o Thus, the non-reflected wavelengths must be absorbed and vice versa
o Putting the two spectra together gives you the original radiation spectrum
IR:
o Far IR radiation can change the colour of zinc oxide from white to yellow in the presence
of oxygen due to increased temperature, this is reversible by decreasing temp
o Red copper (I) oxide and malachite permanently change to black copper (II) oxide since
the heat causes breakdown
UV:
o ZnO (zinc white) fluoresces a pale yellow, while green malachite is a dirty mauve
(CuCO3.Cu(OH)2), since they absorb the radiation but the electrons don’t return
immediately to ground state, going down transitional stages and thus emitted different
wavelengths to the originally absorbed one
• Gather, process and present information about a current analytical technology to:
o Describe the methodology involved
o Assess the importance of the technology in assisting identification of elements in
samples and in compounds, and
o Provide examples of the technology’s use
Laser microspectral analysis
o A laser beam is directed at a sample of the pigment, vaporising a tiny amount

o This vapour passes between two electrodes which spark and excite the vapour particles
o An emission spectrum is obtained by a spectrophotometer when electrons return to
ground state, emitting radiation

o This can be compared with the unique emission spectrum of different elements to obtain
the chemical composition and thus the pigments present
o The intensity of the spectra is proportional to concentration
Importance:
o Laser beam is very intense, pure and focused, allowing the technique to detect trace
amounts of elements in samples
o It is useful for determining authenticity, and in restoration
o Requires minimal preparation, so can be used for a variety of solid samples
Examples of use:
o Analysis of paintings for validity (time period etc.)
o Used to analyse chemical compositions of paintings so correct restorative chemicals are
used
• Define the Pauli Exclusion Principle to identify the position of electrons around an atom
No two electrons in an atom may have identical sets of four quantum numbers
• Identify that each orbital can only include two electrons
Only two electrons which have opposite spins (anticlockwise, clockwise). These produce magnetic
fields, and result in slightly different energy levels for each electron. Note: This arrangement
produces a lower potential energy than if they were paired, since the electron-electron repulsion is
minimised
• Define the term sub-shell
A sub-shell is an energy sublevel within a principle energy level (or shell). S, P, D, F. Number of
subshells = principle energy level. S – spherically symmetrical around nucleus, P – dumbbell
shaped with 3 orientated in perpendicular planes through nucleus
• Outline the order of filling of sub-shells
• Identify that electrons in their ground-state electron configurations occupy the lowest
energy shells, sub-shells and orbitals available to them and explain why they are able to
jump to higher energy levels when excited
When an electron absorbs a quanta of energy corresponding to the difference between its current
energy level and another energy level, it moves to the higher energy level to expend this energy.
• Explain the relationship between the elements with outermost electrons assigned to s, p, d,
f blocks and the organisation of the periodic table
o Elements with similar outer shell electron configurations occur in the same group
o Periods correspond to principle energy levels
o The periodic table is divided into blocks depending on the element’s outer shell
• Explain the relationship between the number of electrons in the outer shell of an element
and its electronegativity
Electronegativity – a measure of the ability for an atom to attract electrons to itself in a chemical
bond
Electronegativity increases across a period – more protons in nucleus meaning increased nuclear
charge. Electrons added do not fully shield each other from effect of increased protons, meaning
electrostatic force on electrons around the nucleus is stronger
Electronegativity decreases down a group – the size of an atom increases, and the attractive force
of the nucleus on valence electrons is diminished
• Analyse information about the relationship between ionisation energies and the orbitals
of electrons
• Describe how trends in successive ionisation energies can be used to predict the number
of electrons in the outermost shell and the sub-shells occupied by those electrons
The ionisation energy is the energy required to

remove an electron completely from the nucleus’
electric field.
In the s-subshell, ionisation energy increases as the

number of electrons increases. The addition of an
electron a similar distance from the nucleus
corresponds with the addition of protons to the
nucleus, increasing nuclear charge. The electrons do
not shield each other enough to cancel this increased
nuclear charge, causing a net increase of
electrostatic force towards the nucleus, increasing
the ionisation energy.
In the p,d,f-subshells, the ionisation energy increase as the electrons are added into different orbitals
(See above), but decreases when an electron is first paired in an orbital. This is because the
electrons repulse each other enough to overcome the effect of increased nuclear charge, raising each
other to higher potential energies and decreasing ionisation energy.
It decreases down a group as valence electrons are further from the nucleus, and thus at higher
potential energies.
• Use Hund’s rule to predict the electron configuration of an element according to its
position on the periodic table
Hund’s rule – every orbital in a subshell must be singly occupied by one electron with identical
spin before any one of these orbitals in the same subshell are doubly occupied

• Identify the block occupied by the transition metals in the periodic table
• Define the term transition element
Transition element - those that form at least one ion with a partially filled sub-shell of d electrons
Zn is not a transition element since its ion Zn2+ has a completely filled d-subshell. Scandium ion
(Sc3+) has no electrons in d-subshell.
They have similar properties since their outer subshell is only s1 or s2
The partially filled d-subshell explains:

o Colour of metal ion complexes – d-d transitions
o Magnetic properties of metal ions – unpaired electrons in d-orbitals causes
paramagnetism, fully paired causes diamagnetism
o More than one stable oxidation state – due to the similarity of the s and d-subshell energy
levels, the electrons in both levels can be lost without high energy cost, meaning the
transition metal forms multiple oxidation states
*
In other elements, the energy levels are too separated to have commonly observable multiple
oxidation states.
• Explain why transition metals may have more than one oxidation state
• Account for colour changes in transition metal ions in terms of changing oxidation states
• Explain, using the complex ions of a transition metal as an example, why species
containing transition metals in a high oxidation state will be strong oxidising agents
Oxidation number is the charge an atom would have in a chemical bond if the bonded electrons
belonged to the more electronegative element.
The number of oxidation states for the first period of transition elements progressively increases to
a peak at VIIB then decreases, since:
o As more electrons are added to a d-orbital, the d and s subshells progressively separate in
terms of energy due to electron repulsion
o However, more electrons also means more possible oxidation states
o The “sweet spot” is Mn, with the number decreasing to the left due to lack of electrons,
and decreasing to the right due to separation of energy levels
CHECK – why do transition metals need to be in oxidised states for this to occur?
CHECK – what about return to ground state, wouldn’t it release the light again?
Transition metals ions in compounds have d orbitals with slightly different energy levels and are
incompletely filled. These energy differences correspond with the energy of visible light, and
electrons in the d subshell can absorb photons to become excited, meaning the complementary light
spectrum is able to pass through and be observed. E.g. transition metals which absorb the red end of
the spectrum will appear blue.
In different oxidation states, the transition metal has a different arrangement of filled and unfilled
3d orbitals, causing differences between d energy levels and thus causing different wavelengths of
visible light being absorbed.
Transition metals in a high oxidation state have a high deficit of electrons, which reduces the orbital
radii and decreases electron shielding. This results in a high oxidation potential and makes them
strong oxidising agents. E.g. Cr2O72- and MnO4-
• Write electron configurations of the first transition series in terms of subshells
As per Hund’s Rule. Note: The d-orbitals are in their most ‘stable’ configuration with a complete
set of unpaired or paired electrons. When the set is almost complete, an electron from the 4s orbital
will transit to make it complete (Chromium and Copper).
• Perform a first-hand investigation to observe the colour changes of a named transition

element as it changes in oxidation state
Prac – Oxidation states of Vanadium

Equipment:
100 mL of 1M NaOH
3 grams of ammonium vanadate
75 mLs of 2M H2SO4
Granulated zinc
4 large test tubes and text tube rack
250 mL conical flask and rubber stopper
Method:
1. 3 grams of ammonium vanadate was weighed on an electronic beam balance
2. 100 mL of 1M NaOH and the 3 grams ammonium vanadate were added to the conical flask
3. The mixture was swirled to dissolve
4. 75 mLs of 2M H2SO4 was added to the solution to acidify it
5. 20 mL of the resulting solution was poured into a test tube using a 50 mL measuring cylinder
6. 8 granules of zinc were dropped into the conical flask which was then stoppered
7. The solution was swirled gently until it became blue, then step 5 performed
8. The solution was swirled again until it became green, then step 5 performed
9. The flask was swirled vigorously until the solution became violet, then step 5 performed
Results and analysis:

(NH4)VO3(s)  NH4+(aq) + VO31-(aq)
(+5 YELLOW to +4 BLUE)

Oxidation: Zn(s)  Zn2+(aq) + 2e-
Reduction: VO31-(aq) + 4H+(aq) + e-  VO2+(aq) + 2H2O(l)
Redox: Zn(s) +2VO31-(aq) + 8H+(aq)  2VO2+(aq) + 4H2O(l) + Zn2+(aq)
(+4 BLUE to +3 GREEN)

Reduction: VO2+(aq) + 2H+(aq) + e-  V3+(aq) + H2O(l)
Redox: Zn(s) +2VO2+(aq) + 4H+(aq)  2V3+(aq) + 2H2O(l) + Zn2+(aq)
(+3 GREEN to +2 VIOLET)

Reduction: V3+(aq) + e-  V2+(aq) (CHECK IT)
Redox: Zn(s) +2V3+(aq)  2V2+(aq) + Zn2+(aq)
Risk analysis:
Hazard Risk Control
Sulfuric acid Corrosive. Can cause Wear safety glasses, use
Kills tissue
Sodium Hydroxide Corrosive. Can cause Wear safety glasses, use
Kills tissue
Ammonium vanadate Eye contact causes redness Safety goggles, gloves, do
and swelling not make airborne
Skin contact causes itching
and pain
Rules for balancing redox half equations

In acidic solution:
1. Write down the reactant and product [Cr2O72-  Cr3+]
2. Balance the number of atoms [Cr2O72-  2Cr3+]
3. Balance the number of atoms of oxygen by adding water to the side with least oxygen
[Cr2O72-  2Cr3+ + 7H2O]]
4. Balance the hydrogen by adding H+ ions [Cr2O72- + 14H+ 2Cr3+ + 7H2O]
5. Balance the charge by adding electrons then add states [Cr2O72- (aq) + 14H+(aq) + 6e- 2Cr3+(aq) +
7H2O(l)]
CHECK – what about 2HI  I2 ?
In alkaline solution:
1. Balance as though in acid solution
2. H+ ions are removed to form H2O by adding same number of OH- ions to both sides
3. Simplify
e.g.
[MnO4-(aq) + 4H+(aq) + 3e-  MnO2(s) + 2H2O(l)]
[MnO4-(aq) + 4H+(aq) + 4OH-(aq) + 3e-  MnO2(s) + 2H2O(l) + 4OH-]
[MnO4-(aq) + 2H2O(l) + 3e-  MnO2(s) + 4OH-]
• choose equipment, perform a first-hand investigation to demonstrate and gather first

hand information about the oxidising strength of KMnO4
Prac – Oxidising strength

1). 20 mL of 0.01 M KMnO4 and 20 mL 1M H2SO4 were added to the same 50 mL measuring
cylinder
2). 8 medium test tubes were set up on a test tube rack
3). About 5 mL of 0.5 M solutions of KI, KBr and KCl droppered into the first 3 using an eye
dropper, and zinc powder, a short magnesium strip, copper pieces, tin pieces and iron nails placed
into the five remaining test tubes corresondingly
3). The eye dropper was rinsed, then used to drop a few drops of the acidified KMnO4 solution into
each of the test tubes
4). A centimetre of toluene was poured into these three test tubes not containing metals, to help
indicate redox reactions as halides discolour it
5). Any changes were observed
5). The contents of the test tubes containing toluene were disposed of in an organic waste bin
Results:
Constant Reactant: KMnO4 solution (MnO4- is pink)
MnO4- + 8H+ + 5e-  Mn2+ + 4H2O (E0 = 1.51 V)
E0 (V) (for Variable Colour KMnO4 Product state Reaction E0
halide/metal) Reactant solution (V)
0.62 KI Colourless I2 (toluene -> pink) 0.89
1.09 KBr Colourless Br2 (toluene -> 0.42
yellow/brown)
1.40 KCl No change* 0.11
-0.76 Zn Colourless Mn2+ 2.27
-2.36 Mg Brown MnO2 3.87
2+
0.35 Cu Colourless Mn 1.16
0.14 Sn No change
0.77 Fe Colourless Mn2+ 0.74
• Some change would be expected to occur (under standard conditions), but the already low
reaction E0 + the non-standard state of the reactants meant there was too small a tendency of
reaction for observable change
Relative oxidation strengths:
Mg2+ < Zn2+ < Sn2+ < Cu2+ < I2 < Fe2+ < Br2 < Cl2 < MnO4-
Risk analysis:
Hazard Risk Control
Potassium permanganate Stains skin and clothing, Wear safety glasses, use
concentrated solutions are lower concentrations
corrosive
Toluene Very flammable Keep away from hot
Skin irritant surfaces, flames or sparks
Polyvinyl gloves
Sulfuric acid Corrosive. Can cause Wear safety glasses, use
Kills tissue
The permanganate ion is a strong oxidant due to the presence of manganese in a high (+4) oxidation
state. This gives it a high electronegativity, thus making it a strong oxidant and in turn a weak
reductant.
Some example reaction formulae,

(Iodide and permanganate)
Reduction:
MnO4- + 8H+ + 5e-  Mn2+ + 4H2O (E0 = 1.51 V)
Oxidation:
5I-  (5/2) I2(aq) + 5e- (E0 = -0.62 V)
Overall:
MnO4- + 5I- + 8H+  Mn2+ + 4H2O + (5/2) I2(aq)
• explain what is meant by a hydrated ion in solution
When an ionic solid dissolve in water, the ions dissociate and are surrounded by water molecules in
a process called hydration. The charge of the ion attracts the polar water molecule. These ions are

hydrated. For many metal cations, a covalent bond is formed between the oxygen atom and the
cation due to vacant orbitals in the metal.
The general formula is [M(OH2)n]m+ where ‘m’ is the charge of the ion and ‘n’ is the no. of water
molecules surrounding the ion. Note that it is OH2 since the “O” is donating the electrons.
Note that many other ligands are possible.
• describe hydrated ions as examples of a coordination complex or a complex ion and

identify examples
• describe molecules or ions attached to a metal ion in a complex ion as ligands
• explain that ligands have at least one atom with a lone pair of electrons
A complex ion (a lewis acid) is where a central metal ion is surrounded by ligands. Ligands (a
lewis base) are atoms, ions or molecules which donate one or more electrons in a coordinate
covalent bond with the central metal ion. Ligands therefore need a lone pair of electrons which it
can donate to an empty orbital of the central ion, in order to bond.
In a hydrated ion, the ligands are water molecules.
• identify examples of chelated ligands
Ligands can bond using different numbers of electron pairs. Monodentate ligands bond using the
electron pair of a single donor atom. Others have multiple atoms with unpaired electrons and can
bind simultaneously, and are polydentate or chelated ligands. Note: Memorise ‘en’ and ‘EDTA’
Example of chelated ligand bonding:

• discuss the importance of models in developing understanding of the nature of ligands
and chelated ligands using specific examples
Models are significant in understanding the nature of ligands. They allow us to specify mechanisms
for the formation and bonding of ligands, and other phenomena where the mechanism cannot be
directly observed but only inferred (e.g. colours). They are not necessarily true, but are adequate
explanations for current use.
1). Valence Bond Theory

Assumes:
o Ligands bond to central ion in completely covalent coordinate bonds
Bonding in the formation of complexes depends on:

o Orbitals available for coordinate covalent bond formation
o Tendency of ions or groups to share a pair of electrons
o Number of ligands that can be placed around the central ion
o The geometry assumed by the ligands
1. The electron pairs from the ligands are placed in empty orbitals of the central ion
2. These orbitals are hybridized, that is, they are mixed to form new orbitals
3. Some ligands cause unpaired electrons in d-orbitals to pair with other unpaired electrons, and
then use the newly empty d-orbitals. Others cannot do this, and use completely vacant orbitals.
4. Those ligands utilising d-orbitals previously occupied have inner spin complexes, the ones not
have outer spin complexes
5. Inner spin complex are paramagnetic, and move in the direction of a magnetic field. Outer spins
are diamagnetic, and move opposite to the direction of a magnetic field
6. The structural geometry can be determined using VSEPR rules
The advantages are:

o Accounts for magnetism of coordination complexes
o Accounts for shape of complexes through hybridisation
Disadvantages:
o Cannot show details such as the energy changes involved, rate at which reactions occur
and the mobility/flexibility of bonds involved
o Cannot account for colours of transition metal compounds
2). Crystal field theory

Assumes:
o Bonds between ligand and central ion are completely ionic
o Ligand and central ion are infinitesimally small, non-polarisable point charges
The ligand is treated as a point negative charge, and when it approaches the central metal ion, the
electron clouds of both get disturbed, resulting in changed energy states for orbitals such as the d-
orbitals. Different ligands and different metal complex geometries result in different degrees of
energy separation. The resulting separation of the d-orbitals is in the energy range of visible light,
and different energies absorb different wavelengths, resulting in different observed colours.
Advantages:
o Explains the colours of transition metal complexes
Disadvantages:
o Neglects any covalent contribution
• use available evidence and process information from secondary sources to draw or model
Lewis structures and analyse this information to indicate the bonding in selected complex
ions involving the first transition series
Some examples:
• process information from secondary sources to give an example of the range of colors
that can be obtained from one metal such as Cr in different ion complexes
The colour depends on the central metal ion (element and oxidation state) and the surrounding
ligands.
Prac – Colours of Chromium complexes

1). Place 5 mL of chromium (III) nitrate solution into five large test tubes
2). Record the colour of the solution in the first test tube0
3). 4 rice grains of sodium sulfate was added to the next test tube, and gently
warmed over a bunsen burner1
4). The colour of the solution was recorded
5). 4 rice grains of sodium chloride was added to the next test tube, and gently
warmed over a bunsen burner2
6). Colour was recorded
7). A dropping pipette was used to drop 4 drops of 1 M sodium hydroxide
solution into the next test tube, then colour recorded3
8). More sodium hydroxide was added to the test tube until the precipitate
disappeared, then colour recorded4
9). Ammonia solution was added to the last test tube so a precipitate formed,
then more added until the precipitate disappeared
10). Colour recorded5
0. [Cr(H2O)6]3+ (blue)
1. [Cr(H2O)5SO4]+ (green)
2. [Cr(H2O)4Cl2]+ (green)
3. Cr(H2O)3(OH)3 (grey-blue ppt)
4. [Cr(OH)6]3- (green)
5. [Cr(NH3)6]3+ (mauve)
Errata:
1). Added an actual evaluation in the summary of ethanol usage, changed it
2). The diagram for dry cell is incorrect (page 23-24)
3). Some content changes made to reflux (page 45)
4). Defined allotrope (page 54)
5). Added risk analysis (page 69)
6). Big error!! Explanation of why transition metals compounds are coloured was
incorrect before. See last sentence of paragraph below table (page 64)
7). Should be bromine water not iodine in the lycopene prac (page 7)
8). Ester structure clarified. The way I wrote them as word equations before was
misleading (page 44)
9). Completed the IR Absorption spectroscopy description (page 71)
10). Made risk analysis more specific for bromine water prac (page 6)
11). Scintillation counter description before was INCORRECT (page 27)
12). Modified the explanation explaining multiple oxidation states of transition
metals
13). Ester page updated with more information (page 44)
14). Made the uses of ethylene page more readable
15). Error! Iron sulfide in coal (production of sulfur dioxide) – equation is incorrect,
see correction page (34)
16). Expanded description of Arrhenius acids, including limitations, and also
Bronsted-Lowry theory (page 39)
17). Clarified conjugate base/acid explanation (page 39)
18). Changes to history of pigment use – I got kohl composition wrong
19). Amphoteric is NOT amphiprotic! Amphoteric means it has both acidic and basic
PROPERTIES. Amphiprotic means it can both accept and donate protons which ARE
properties of most acids/bases but do not encompass all of them

Chemistry HSC Full Notes BEST NOTES

Uploaded by

Copyright:

Available Formats

Chemistry HSC Full Notes BEST NOTES

Uploaded by

Document Information

Original Title

Copyright

Available Formats

Share this document

Share or Embed Document

Sharing Options

Did you find this document useful?

Is this content inappropriate?

Copyright:

Available Formats

Chemistry HSC Full Notes BEST NOTES

Uploaded by

Copyright:

Available Formats

Chemistry HSC

Organic Chemistry – the study of compounds containing carbon

Branched chain (one is attached to at least 3)

Naming Alkanes, Alkenes and Alkynes

Saturated and unsaturated compounds

Types of Organic reactions

Chem notes David Lee BHHS 2007 3

H2C=CH2 + X-Y => XH2C-CH2Y

1) Addition of hydrogen to ethylene (hydrogenation) - ethylene to ethane by

[CH2=CH2(g) + Br2(l)  CH2Br-CH2Br(l)] (petrol additive)

Chem notes David Lee BHHS 2007 4

[CH2=CH2(g) + HOBr(aq)  CH2OH-CH2Br]

Hydrogen bromide reaction:

[CH2=CH2(g) + HBr(g) -> CH3 – CH2Br] (What states?)

3) Chloroethene – Monomer for PVC

5) Ethanol – Used as a fuel in automobiles and as an industrial solvent

6) Ethylene oxide and ethanediol – fumigant (former), manufacture of polymers

• Identify the following as commercially significant monomers

Vinyl chloride – chloroethene CH2 = CHCl

Styrene – ethylbenzene C6H5CH=CH2 (also known as phenylethene)

Prac – Reactions of hydrocarbons with bromine water

Aim: To compare reactivities of an alkene (cyclohexene), alkane (cyclohexane), and an aromatic

Bromine w/ cyclohexene (top)

Bromine water w/ cyclohexene (bot)

Toluene did not react as aromatic molecules have

Bromine with cyclohexane, this is substitution:

Chem notes David Lee BHHS 2007 6

Polymerisation is the process of bonding monomers together to form long chains.

Polyethylene is an addition polymer, it is created through addition polymerisation.

Addition polymerisation requires a catalyst or

Chem notes David Lee BHHS 2007 7

 Termination – at various times, it is possible for two free radical polymers

Comparison of structures and properties

Factors affecting the properties of polymers:

Chem notes David Lee BHHS 2007 8

Polymer Properties Uses

• Explain what is meant by a condensation polymer

• Describe the reaction involved when a condensation polymer is formed

e.g. Condensation polymerisation of nylon

A variety of monomers can be used to manufacture

This experiment used interfacial polymerisation. It

• Describe the structure of cellulose and identify it as an example of a condensation

Biomass – organic material derived from living organisms

The presence of five hydroxyl groups allows glucose to form polymers

o Many glucose units linked together (polysaccharides).

Use as raw material can be achieved by:

Chem notes David Lee BHHS 2007 11

o Biopol, PHA (polyhydroxyalkanoates) copolymers, a family of microbial energy reserves

Simplest are PHB’s

Currently working to produce PHA’s directly in non-food crop plants.

Disadvantage: Current high cost of production as opposed to crude oil sources

Chem notes David Lee BHHS 2007 12

Reverse reaction is hydration, requires dilute aqueous sulfuric acid:

Prac – Ethanol as a solvent

Chem notes David Lee BHHS 2007 13

With 1,2,3 – propanetriol:

• Describe conditions under which the fermentation of sugars is promoted

o Fermentation requires a carbohydrate as starting material, such as glucose, sucrose or

Chem notes David Lee BHHS 2007 14