1. Industrial chemistry processes have enabled scientists to develop replacements for natural
products
Discuss the issues associated with shrinking world resources with regard to one identified natural
product that is not a fossil fuel, identifying the replacement materials used and/or current
research in place to find a replacement for the named material
Uses of rubber and properties that make it useful
Properties
Elastic
Waterproof
Good electrical insulator
Uses related to properties
Balls, shoes, tyres, elastic bands, elastic
bandages
Raincoats, paints
Electrical insulation
Changes in its use over time




First used by native South Americans to make shoes
During the 1900’s, it was used in Europe as erasers and raincoats
Used in tyres of military vehicles during WWII
Currently used in shoes, tyres, elastic bands, hoses, electrical insulation etc.
Sources of rubber
 Natural rubber was obtained plantations of rubber trees in tropical areas (e.g. South East
Asia, South America)
 The trees would be ‘tapped’ (an incision made into the bark of the tree and the latex sap
collected, then refined into a useable rubber)
Causes to the shrinkage of world resources of rubber
 Demand for rubber has continued to increase after WW2; traditional sources of natural
rubber could not meet the increased demand – thus leading to the shrinkage of world
resources of rubber
 Takes a long time (six years) to grow a rubber tree which can be utilised as a resource of
rubber.
Replacement material
Styrene butadiene rubber (SBR) is the most common synthetic rubber used to replace natural
rubber today.
Identify data, gather and process information to identify and discuss the issues associated with the
increased need for a natural resource that is not a fossil fuel and evaluate the progress currently
being made to solve the problems identified
Problems associated with the use of (natural) rubber
 Limited natural resource – during WWII, demand for rubber increased greatly whilst supplies
were interrupted (Japanese had control of rubber producing areas). As a result, supply of
natural rubber could not meet demands.
 Time and effort required to harvest rubber from trees – takes six years for a rubber tree to
grow to a point where it can be economically harvested; rubber trees can only produce a
certain amount of rubber each year
 In its raw state, natural rubber became sticky when hot, and brittle when cold
What is being done about the problems
 During the 1930’s, the scientist Walter Bock had discovered synthetic rubber (SBR) as a
replacement to natural rubber. During WW2, both Germany and the US began industrial
production of synthetic rubber such as SBR to replace natural rubber for use in the war
effort.
 The vulcanization process (sulfur and other substances are added to raw rubber and then
the mixture is heated) modifies rubber so that it will not become sticky when hot and brittle
when cold.
Evaluation of the progress being made to solve the problems
The development of synthetic rubber was used to solve the problems the supply of natural
rubber.
 The progress made to increase the supply of rubber has been very effective, with
synthetic rubber making up approximately 80% of the world’s total rubber production
today. The development of synthetic rubber has effectively allowed the production of
rubber to meet demand.
 Through vulcanisation, the properties of synthetic rubber have been improved, making it
more durable, more resistant to chemical attack and stronger. Furthermore, it is
produced at a relatively low cost – meaning it is economically viable.
 Synthetic rubber, however, is derived from fossil fuels, a limited resource. Recent
developments have been made to involve the use of non-petrochemical (renewable
resources) in the production of rubber.
Overall, the progress being made to solve the problems has been very successful and effective.
2. Many industrial processes involve manipulation of equilibrium reactions
Explain the effect of changing the following factors on identified equilibrium reactions: pressure,
volume, concentration, temperature
Chemical equilibria are dynamic and involve reversible reactions. Equilibrium is reached when the
rate of the forward reaction equals the rate of the reverse reaction. Pressure, volume, concentration
and temperature can alter the position of equilibrium. Le Chatelier’s principle can be used to predict
the response of an equilibrium system to a change in these conditions.
Changes in pressure
Pressure changes may affect chemical equilibria involving gaseous systems. The total gas pressure
can be increased or decreased by changing the volume of the reaction vessel, or by the injection of
more gas (increasing the partial pressure of that gas, consequently adding to the total pressure).
Note that if a gas is injected that does not react chemically with the reactants/products there is no
shift in equilibrium as the partial pressures of the gases in question have not changed.
When gaseous pressure is increased, there is a greater concentration of gas particles – increasing the
rate of collision. An increase in pressure in an equilibrium reaction will shift the system towards the
side with the least amount of gas particles; whilst a decrease in pressure will shift the system
towards the side with more gas particles.
Changes in volume
In gaseous equilibria, volume changes lead to changes in total gas pressure. However, volume
changes can also affect equilibria in non-gaseous systems
In aqueous systems, changes in the volume of the solvent (water) will change the concentration of
the dissolved particles and this can shift the equilibrium. For example, if more water is added, there
will be an immediate drop in the concentration of all species; this will cause the equilibrium to shift
to the side with more solute particles.
If we consider a 100mL solution of saturated sodium chloride in contact with excess salt crystals:
NaCl(s)
Na+(aq) + Cl-(aq)
The system is at equilibrium. If extra water is added, the concentration of dissolved ions decreases.
This causes the equilibrium to shift to the right to produce more solute particles; and eventually a
new equilibrium is reached.
Changes in concentration
The addition or removal of some reactants/products in an equilibrium system leads to sudden
changes in concentration. The system responds to counteract the change. An increase in equilibrium
reactant concentration will cause the equilibrium to shift to the products and vice versa.
Changes in temperature
Chemical equilibria are either endothermic or exothermic in the forward direction. Therefore,
changes in temperature will affect the equilibrium system. An increase in temperature in an
endothermic reaction will shift the equilibrium to the right, to consume the added heat. An increase
in temperature in an exothermic reaction will shift the equilibrium to the left to consume some of
the heat produced.
Interpret the equilibrium constant expression (no units required) from the equilibrium equation of
equilibrium reactions
Identify that temperature is the only factor that changes the value of the equilibrium constant (K)
for a given equation
For any equilibrium reaction, at its point of equilibrium, we can calculate a constant – this constant is
known as the equilibrium constant (K)
In an equation represented as:
aA + bB
cC + dD
A way to remembering this formula is PORK (Products Over Reactants equals K). Note that
calculations involving the equilibrium constant expression do not involve pure solids and liquids
(their value is just kept at 1, and so, is disregarded)
The values of the equilibrium constant give a quantitative indication of the extent of the equilibrium:
 Values of K near 1 indicate that the equilibrium does not lie strongly to the left or right
 Values of K that are very large (>104) indicate the equilibrium lies well to the right
 Values of K that are very small (<10-4) indicate the equilibrium lies well to the left
The only factor that changes the value of the equilibrium constant is temperature change. For
endothermic equilibria, K increases when temperature increases. For exothermic equilibria, K
decreases when temperature increases.
Reaction quotient (Q)
The reaction quotient is calculated the exact same way as for K; if Q=K, then the reaction is at
equilibrium, if not, then the reaction is not at equilibrium.
Choose equipment and perform a first-hand investigation to gather information and qualitatively
analyse an equilibrium reaction
2NO2(g) ↔ N2O4(g)
Brown
ΔH = -ve
Colorless
Method
1. Obtain 2 NO2/N2O4 gas tubes
2. Place one tube in a beaker containing cold ice. Leave the other tube at room temperature
3. After 5 minutes, compare the colour of the gases in each vial
4. Record observations
Result: The tube in the ice-filled beaker should be less coloured than the one at room temperature.
This is because the decreased temperature shifted the equilibrium towards the right, generating
more colourless N2O4 gas.
Cold
Room temperature
3. Sulfuric acid is one of the most important industrial chemicals
Outline three uses of sulfuric acid in industry
Sulfuric acid has a wide range of uses and is the world’s most widely used/produced industrial
chemical:
 Fertiliser – Most of the sulfuric acid produced is used to make fertiliser:
- Superphosphate fertiliser is made by reacting 70% sulfuric acid with crushed calcium
phosphate Ca3(PO4)2. This forms a mixture of calcium dihydrogen phosphate Ca(H2PO4)(s)
and calcium sulfate – this mixture is then sold as superphosphate.
- Ammonium sulfate – ammonia gas is passed through sulfuric acid, which acts as a
‘scrubber’ – forming ammonium sulfate: 2NH3 (g) + H2SO4 (l) (NH4)2SO4 (s)
 Cleaning of iron/steel – before iron or steel can be galvanised, any oxide that has formed on
the surface must be removed. Sulfuric acid is very corrosive, and therefore can remove rust.
 Dehydrating agent – concentrated sulfuric acid is a strong dehydrating agent used to
produce ethylene from ethanol:
Describe the processes used to extract sulfur from mineral deposits, identifying the properties of
sulfur which allow its extraction and analysing potential environmental issues that may be
associated with its extraction
Processes used to extract sulfur:
The Frasch process is the process used to
extract sulfur from mineral deposits. Firstly,
superheated and pressurised water (160
o
C/1.5mPa) is injected into the sulfur deposit
through the outer pipe. This melts the sulfur;
so the water and sulfur form an emulsion
(small droplets of one liquid dispersed
through another liquid). Compressed air is
injected through the central pipe, forcing the
sulfur-water emulsion to the surface through
the middle pipe. At the surface, the emulsion
is collected in large vats. As the mixture
cools, solid sulfur separates into liquid water
and solid sulfur. The sulfur obtained is 99.5%
pure.
Properties of sulfur:
 Sulfur has a very low melting point of 113oC (weak dispersion forces)
 Has a low density (2.07 g/cm3) – producing a light emulsion that can be readily transported
to the surface
 Insoluble with water; doesn’t react chemically – it can be easily recovered at the surface
 Inert, non-toxic and non-volatile – relatively safe for miners
Potential environmental issues
 Sulfur can be quite easily oxidised to SO2 which is an air pollutant that can cause respiratory
irritation and the formation of acid rain.
 The recovered water can cause thermal pollution and contaminate local ecosystems (due
to impurities dissolved into the water from the sulfur). To avoid this, it must be either
recycled or processed (purified and cooled) before releasing to the environment.
 It is very difficult to back-fill the underground caverns left by the extraction of sulfur. The
mining area is therefore prone to collapse (subsidence).
Outline the steps and conditions necessary for the industrial production of H2SO4 from its raw
materials
Describe the reaction conditions necessary for the production of SO2 and SO3
Natural gas and crude oil often contain significant amounts of hydrogen sulfide, from which
elemental sulfur can be extracted.
The main process used to produce H2SO4 industrially is the contact process.
1. Molten sulfur is firstly reacted with hot, dry air: S(l) + O2 (g)  SO2 (g)
- The dry, oxygen-rich air has been scrubbed of water vapour through dehydration by 99%
sulfuric acid. The air must be dry to avoid acid mist and corrosion of pipes. An excess of
air (which provides oxygen for the combustion) ensures the sulfur reacts completely
- Considerable heat is generated, so the gas stream must be cooled before the next stage
(from 1000oC to 400oC). This is done using a heat exchanger, which removes and recycles
the heat so it can remelt more sulfur or produce steam to power turbines to generate
electrical energy for the factory.
2. Sulfur dioxide is catalytically oxidised to form sulfur trioxide: SO2 (g) + ½O2 (g)
SO3 (g)
- The sulfur dioxide is mixed with air, and passed into a catalyst tower (called the
converter). This tower contains multiple layers of vanadium oxide (V2O5) catalyst. The
gas is passed over the catalyst and sulfur dioxide is oxidised to produce sulfur trioxide.
*Note this process is in equilibrium
- Carried out at about 400-550oC (moderate temp), 1-2 atm. pressure
- Excess oxygen is used (shifting equilibrium to right by LCP)
3. The sulfur trioxide is passed into 98% sulfuric acid to form oleum: SO3 (g)+H2SO4 (l) 
H2S2O7(l)
-
This process occurs in a second absorption tower. The sulfur trioxide enters at the
bottom of the tower and sulfuric acid is sprayed in at the top. The acid trickles down and
reacts with the gas to form an oily liquid, oleum (H2S2O7)
- Directly dissolving SO3 in water is impractical due to the high exothermic nature of the
reaction, consequently producing sulfuric acid mist, which is dangerous and hard to
utilise.
4. The oleum is diluted to form sulfuric acid H2S2O7 (l) + H2O (l)  2H2SO4 (l)
- In the diluter, water is mixed with oleum to produce 98% sulfuric acid
Apply the relationship between rates of reaction and equilibrium conditions to the production of
SO2 and SO3
Production of SO2
 Production of SO2 is a reaction that goes to completion, therefore there are no equilibrium
considerations
 However, in order to increase the rate of reaction, the sulfur is liquefied (increasing SA) and
an excess of oxygen is used
Production of SO3
SO2 (g) + ½O2 (g)
SO3 (g)
 Equilibrium considerations:
- Gas pressure: High gas pressures will favour the forward reaction
- Temperature: the reaction is exothermic, hence lower temperatures favour the forward
reaction
- Oxygen concentration: an excess of oxygen will drive the equilibrium towards the right
 Rate considerations:
- Temperature: High temperatures are favourable as the rate of reaction will increase due
to the increase in molecular collisions
-
Increase in gas pressure also helps increase the rate of reaction
Catalyst: the presence of the vanadium catalyst will decrease the activation energy and
increase the rate of the reaction.
Compromised conditions
A balance between the equilibrium yield and rate of reaction has to be made in the industrial
synthesis of sulfur trioxide. As well as this, economic practicality is a major factor in the
determination of conditions used.
 The reaction mixture is maintained at a moderate temperature of around 400-550oC
 Pressure is kept relatively low (around 1-2 atm) due to safety and economic reasons (highpressure apparatus are expensive).
 A small excess of oxygen is used (approximately twice the stoichiometry)
 A catalyst of vanadium oxide is used.
Flow chart and equations:
Describe, using examples, the reactions of sulphuric acid acting as: an oxidising agent, a
dehydrating agent
Oxidising agent
Sulfuric acid is a relatively strong oxidant (oxidising agent). The half reaction for sulfuric acid acting
as an oxidising agent is:
For example, hot concentrated sulfuric acid oxidises copper to copper ions. The copper half equation
is:
Combining the above with the general reaction forms:
Other metals such as Ag, Hg and Pb react similarly; however, more reactive metals such as Zn, Mg,
Fe and Al are oxidised by hydrogen atoms to form H2 gas (i.e. they are still redox reactions but the
oxidant is the hydrogen ion rather than the sulfuric acid itself)
Oxidation of non-metals such as carbon, sulfur and phosphorus produce carbon dioxide, sulfur
dioxide and phosphorus pentoxide respectively.
Dehydrating agent
Concentrated sulphuric acid rapidly absorbs water vapour from the air. It is classified as a
dehydrating agent as it has a high affinity for water, and can thus absorb water from mixtures. For
example, concentrated sulfuric acid can be used to dehydrate sucrose:
It can also be used to dehydrate ethanol to ethylene
*Note if the question is in the option topic, don’t write the ethanol-ethylene equation, use the
sucrose one
Describe and explain the exothermic nature of sulfuric acid ionisation
Sulfuric acid produced by the contact process is 98% H2SO4. When this is diluted with water, a large
amount of heat is released.
H2SO4 (l) + H2O (l)
HSO4− (aq) + H3O (aq) ΔH = -90 kJ/mol
This is because in concentrated sulfuric acid, there are very few ions, as most of the water is tied up
as hydrates (e.g. H2SO4.H2O). Therefore, the addition fo water to the acid provides an extra source
for the H2SO4 molecules to ionise with. Contrastingly, in concentrated hydrochloric and nitric caids,
there is sufficient water for all the acid molecules to already be ionised. Adding water therefore just
dilutes the ionic solutions and this releases much less heat.
When the H2SO4 molecules are ionised, the energy released when H3O+ is formed is much greater
than the energy absorbed when the H2SO4 is broken – therefore, the ionisation is strongly
exothermic. The acid ionises in two steps:
1.
H2SO4 (l) + H2O (l)
HSO4− (aq) + H3O+ (aq)
2.
HSO4− (aq) + H2O (l)
SO42− (aq) + H3O+ (aq)
ΔH = -90 kJ/mol
The second ionisation only occurs slightly (HSO4− is a weak acid), and is negligible. The first ionisation,
however, goes to completion and is strongly exothermic.
Identify and describe safety precautions that must be taken when using and diluting concentrated
sulfuric acid
Safety precautions
 Due to the highly exothermic nature of sulfuric acid dilution, there is a serious risk of boiling
the solution and cause splashing of the acid. In order to minimise this risk, small volumes of
the acid is added slowly to the water whilst constantly stirring to distribute the heat
generated. Never add water to the concentrated sulfuric acid!
 When pouring sulfuric acid into a beaker, pour the acid slowly down a glass rod to avoid
splashing
 Always wear safety goggles (acid splashes into eyes can cause serious damage)
 Wear protective gloves and a laboratory coat (sulfuric acid, like all concentrated strong acids,
is very corrosive to skin and clothing)
 Have a supply of sodium carbonate at hand to neutralise any spills when working with the
acid
 If sulfuric acid contacts skin, wash off rapidly with excessive amounts of running tap water. If
large amounts are spilt – it is best to wipe excess with paper towel first then wash with
water to avoid the heat generated on dissolution with water.
Use available evidence to relate the properties of sulfuric acid to safety precautions necessary for
its transport and storage
Properties
 Sulfuric acid is a strong acid that corrodes metals and other materials; it can cause extensive
damage on the structure and functioning of living organisms.
 Sulfuric acid is an oxidising agent – meaning it can attack materials normally resistant to
attack from just the hydrogen atoms of other acids
 Sulfuric acid is a dehydrating agent – meaning it can easily destroy organic materials
(including living tissue)
Transport/storage
For transport and storage of sulfuric acid, well-sealed containers must be used.
 Concentrated sulfuric acid (98%) is virtually all molecular (no hydronium ions), and does not
attack iron or steel; therefore, it can be safely stored and transported in steel
tankers/containers. Steel is much stronger than glass or plastic, and is less likely to rupture if
there is an accident
 Dilute sulfuric acid, however, contains hydronium ions, and vigorously attacks metals like
iron and steel. Hence, dilute acid must be stored in glass or plastic containers – making it
more difficult to transport.
 When storing or transporting, care must be taken to avoid contamination with water – as
this will cause significant ionisation and heat build-up will occur. Hydrogen gas will be
released, and the steel will be attacked by the acid under these conditions.
 For regular use, store sulfuric acid in smaller bottles (e.g.1 litre) for safe handling. The
bottles should be stored in a secure, cool, ventilated room, and is plastic trays in case of
breakage.
 Store sulfuric acid well away from metals, bases, and water, as sulfuric acid will react with
these substances exothermically.
Perform first hand investigations to observe the reactions of sulfuric acid acting as: an oxidising
agent, a dehydrating agent
Dehydrating agent
1.
2.
3.
4.
In a fume cupboard, place sucrose in a 150 mL beaker to a depth of 1cm
Add 5 mL of concentrated sulfuric acid
Stir gently with a glass rod
Observe
Results: In the beaker, the sugar turned black, began to smoke (water vapour) and rise to form a
large tower of black carbon. The equation was:
C12H22O11 (s)
12C (s) + 11H2O (l)
Oxidising agent
1.
2.
3.
In a fume cupboard, place a few crystals of potassium iodide
Add 5 drops of concentrated sulfuric acid.
Darkening of the mixture indicates the formation of molecular iodine
Results: The mixture darkened, indicating the formation of molecular iodine. The equation was:
Safety
 Conducted in a fume cupboard to contain the toxic sulfur dioxide gas
 Small quantities of sulfuric acid was used to prevent too much heat forming; these are all
exothermic reactions