Key Idea 6: Aluminium (obtained from bauxite) is a strong, lightweight, rust resistant metal, used in items such as saucepans, bicycles, window frames and aeroplanes. Understand and write equations for the purification of alumina from the bauxite ore (Bayer Process). Describe and write equations for the redox process that occurs in the Hall-Heroult Cell. Aluminium: From Ore to Metal Australia is the worlds largest producer of bauxite
with substantial resources in WA, QLD and NT. Aluminium is important for manufacturing as it is lightweight, strong, durable and can be recycled. Aluminium: From Ore to Metal The steps for obtaining aluminium are: 1. Bauxite ore is mined 2. Bauxite is processed where it is treated to remove impurities main impurities are iron oxide (Fe2O3) and silicon dioxide (SiO2) BAYER PROCESS
The result is white aluminium oxide or alumina (Al2O3) 3. Further processing by electrolysis in the HALL-HEROULT CELL Bayer Process Hall-Heroult Cell The Bayer Process Bauxite is a naturally occurring material from which alumina (Al2O3) and aluminium metal are produced. The principal minerals in bauxite are gibbsite (Al2O3.3H2O) boehmite
(Al2O3.H2O) Impurities in bauxite Clay Iron (III) oxide, Fe2O3 Silicon dioxide, SiO2 The Bayer Process Step 1. The bauxite is washed and crushed, reducing the particle size and increasing the available surface area for mixture with hot, concentrated sodium hydroxide solution. Alumina and silica react with hydroxide: Alumina: Al2O3(s) + 2OH-(aq) 2AlO2-(aq) + H2O(l) Silica: SiO2(s) + 2OH-(aq) SiO32-(aq) + H2O(l)
The Bayer Process Iron oxide (Fe2O3) does not react with hydroxide (because it is a basic oxide) and is filtered off to produce red mud The Bayer Process Step 2: Water and seed crystals of alumina are added to the solution. Hydrated alumina crystallises leaving the silicate ion in solution 2AlO2-(aq) + 4H2O(l) Al2O3.3H2O(s) +2OH-(aq) The Bayer Process Producing Alumina Step 3: The crystals are then separated off and dehydrated
by roasting (heated). Al2O3.3H2O Al2O3(s) + 3H2O(g) Alumina (Al2O3) is a white powder, which is the final product of the Bayer Process. This is then ready for shipment to aluminium smelters in the Hall-Heroult Process. The Bayer Process Purification of alumina from bauxite ore Metal Reactivity Series - Aluminium
Hall-Heroult Process is the Electrolytic reduction of aluminium oxide to aluminium The commercial operation for reducing alumina to aluminium is an electrolytic process invented independently but simultaneously by Paul Heroult in France and Charles Hall in the US in 1886. Huge quantities of electricity are required (10-13 kWh of electricity per 1 kg of aluminium) so aluminium smelters need to be located close to sources of cheap electric power.
Molten electrolysis is used in the reduction of: Potassium Sodium Calcium Magnesium Aluminium Molten electrolysis is an expensive process as a considerable amount of energy is needed.
Aluminium reduction- Molten Electrolysis The reduction of aluminium ions to aluminium metal does not occur easily Molten alumina is required for the last stage of aluminium extraction. Alumina = Aluminium Oxide, Al2O3 Alumina has a very high melting point of approximately 2030C Electrolysis of aluminium https://www.youtube.com/watch?v=CglZBajYrME Carbon anodes Carbon cathode lining
Steel lining Al(l) tap Al(l) tap Molten mixture of alumina and cryolyte Molten Al How it works Direct Current Al2O3
Anode + Al3+ O O 2- 2- O2- Al3+
Al3+ Cathode - O2- O2Al 3+ O2- 2Al3+ + 3O2- At the anode O2e- e2-OO 2
anode + positive - e 2-e O CO2 O2O e-2-eO2 O2- e- e-
2O2- O2 + 4eO2 + C CO2 At the cathode Al Al3+ e- eeCathode Al Al3+ Al3+ + 3e- Al l iq uid al e- ee-
negative At the electrodes At the negative cathode Positive Aluminium ions drift to the cathode. They gain electrons and become aluminium atoms Liquid aluminium metal is
deposited the- cathode Al3+ +at3e Al At the positive anode Negative oxide ions drift to the positive anode. They lose electrons and become oxygen gas molecules. 2O2- O2 + 4eO2 + C CO2 Oxygen combines with carbon from the anode to produce carbon dioxide
Reactions Summary Hall Heroult Cell ANODE reaction: Oxide ions are oxidised to O2 and at the high temp. of the cell the carbon anodes are burnt and must be replaced periodically. 2O2- O2(g) + 4e- AND C(s) + O2(g) CO2(g) CATHODE reaction: Aluminium ions are reduced to aluminium metal. Al3+ + 3e-
Al(l) OVERALL reaction: 4Al3+(l) + 6O2- 4Al(l) + 3O2(g) OR 4Al3+(l) + 6O2-(l) + 3C(s) 4Al(l) + 3CO2(g) Hall-Heroult Cell The O2 produced at
the anode reacts with the graphite. C + O2 CO2 Aluminium reduction- Molten Electrolysis Temperature of reaction Alumina is dissolved in molten cryolite (Na3AlF6) which decreases its melting point from 2030C to just under 1000C and also increases conductivity of the solution. This is A LOT more economical but still requires a large amount of energy. Alumina electrolyte is added continuously to replace the alumina that is consumed by the reaction. Summary of Hall-Heroult cell
Cyrolite (Na3AlF6): FLUX Lowers the temperature at which Al2O3 melts from 2020oC to 900oC Saves electricity for heating and therefore money. Cathode (reduction) At 900oC and above Al2O3(s) 2Al3+(l) + 3O2-(l) Al3+(l) + 3e- Al(l) Liquid aluminium forms and settles at the bottom of the Hall-Heroult Cell. Anode (oxidation) 2O2-(l) O2(g) + 4e Problem: the anode is made up of graphite/carbon and at 900C, the oxygen formed at the anode reacts with the anode. C(s) + O2(g) CO2 (g) Therefore the anodes have to be periodically replaced. Final Stages of Hall-Heroult Process
Molten aluminium is more dense than cryolite and collects on the cell bottom where it forms the operating cathode. It is normal to operate a cell with a pool of aluminium 10 cm deep, the metal being siphoned off daily to maintain this level. The metal is generally cast as ingots, which are at least 99 % pure with small amounts of iron and silicon being the main impurities. Alternatively, the metal is kept molten in a furnace to which other elements are added to make alloys, prior to cooling. Metal of up to 99.999 % purity is produced by further refining. Suggest 3 examples of uses for aluminium and which properties of the metal make it suitable.
Give 3 reasons why to recycle aluminium Uses of Aluminium Aluminium is the most widely used metal after iron. Its properties include: low density, malleable and easily worked, corrosion resistant and good conductor of both heat and electricity. With the exception of corrosion resistance, all these properties can be improved or augmented by alloying aluminium with small amounts of other metals. It is therefore not surprising that
aluminium has a very wide range of applications. Uses of Aluminium 1) Low density and strength make aluminium ideal for construction of aircraft, lightweight vehicles, and ladders. An alloy of aluminium called duralumin is often used instead of pure aluminium because of its improved properties. 2) Easy shaping and corrosion resistance make aluminium a good material for drink cans and roofing materials. 3) Corrosion resistance and low density leads to its use for greenhouses and window frames. 4) Good conduction of heat leads to its use for boilers, cookers and cookware. 5) Good conduction of electricity leads to its use for overhead power cables hung from pylons
(low density gives it an advantage over copper). 6) High reflectivity makes aluminium ideal for mirrors, reflectors and heat resistant clothing for fire fighting. Why recycle aluminium? Aluminium ores could run out Extraction of aluminium uses lots of electricity and so is very costly Production of the greenhouse gas carbon dioxide during the process Loss of landscape due to mining, processing & transporting bauxite ore. Loss of landscape for electrolysis plant Noise pollution Pollution involved in energy generation Avoids need to dump old aluminium if it wasnt recycled.
Full process https://www.youtube.com/watch?v=WaSwimvC GA8 Ant hill castings https://www.youtube.com/watch?v=EWVx9K7K wBA Key Idea 7: Precipitation reactions are when an insoluble compound (solid) is formed from two solutions mixing. These are sometimes used in metal extractions. Observe some precipitation reactions and write fully
balanced equations for the reactions that occur. Precipitation reactions Precipitation reactions occur when cations and anions in aqueous solution combine to form an insoluble ionic solid called a precipitate. Whether or not such a reaction occurs can be determined by using the solubility table for common ionic solids.
Precipitation reactions The solids produced in precipitate reactions are crystalline solids, and can be suspended throughout the liquid or fall to the bottom of the solution. The remaining fluid is called supernatant liquid. The two components of the mixture (precipitate and supernate) can be separated by various methods, such as filtration, centrifuging, or decanting. Double displacement reactions AB(aq)+ CD(aq) AD(aq)+ CB(s)
AB and CD are usually aqueous ionic compounds consisting of aqueous ions (A+ and B-, C+ and D-). When a double displacement reaction occurs, the cations and anions switch partners, resulting in the formation of two new ionic compounds AD and CB, one of which is in the solid state. This solid product is an insoluble ionic compound called a precipitate. To determine whether a product ionic compound will be soluble or insoluble, consult a Solubility table . If both of the predicted products are soluble, a precipitation reaction will not occur. Example: 2NaOH(aq) + MgCl2(aq) Double displacement reactions AB(aq)+ CD(aq) AD(aq)+ CB(s) Example:
2NaOH(aq) + MgCl2(aq) 2NaCl(?) +Mg(OH)2(?) To determine if a precipitate occurs, use the solubility table to predict the states of the products. 2NaOH(aq) + MgCl2(aq) 2NaCl(?) + Mg(OH)2(?) Double displacement reactions Answer: 2NaOH(aq) + MgCl2(aq) 2NaCl(aq) +Mg(OH)2(s) Now you try: barium chloride + sodium carbonate
BaCl2(aq) + Na2CO3(aq) ? Double displacement reactions BaCl2(aq) + Na2CO3(aq) barium chloride sodium carbonate BaCO3(s) + 2NaCl(aq) barium bicarbonate sodium chloride
Precipitation of minerals Precipitation, like crystallisation is the reverse of dissolving. If a solid comes out of solution slowly, a regular, solid lattice has time to form and large crystals are produced. If the solid is formed quickly, there is no time for a large, solid lattice to grow and instead many particles form in the liquid (precipitation) Precipitation of minerals Mineral precipitate. A mineral deposited from a water solution in pores or other openings in rocks. Chemical reaction with
the surrounding rock, changes in pressure or temperature, or just drying up (evaporation) can cause a mineral to precipitate out of solution. Quartz veins are common products of mineral precipitation. Precipitation reactions practical BaCl2 AgNO3 Pb(NO3)2 CuSO4
FeSO4 NaCl Na2CO3 Kl NaOH - + -
- + + + + + + +
+ - + + + - + +
+ BaCl2 NaCl Na2CO3 BaCl2 + Na2CO3 Kl NaOH -
- - BaCO3+ 2NaCl AgNO3 Pb(NO3)2 CuSO4 FeSO4 AgNO3 + NaCl
AgNO3 + Na2CO3 AgNO3 + Kl AgNO3 + NaOH AgCl + NaNO3
AgCO3 + Na2NO3 AgI + KNO3 AgOH + NaNO3 Pb(NO3)2+ 2NaCl Pb(NO3)2 + 2Na2CO3 Pb(NO3)2 +2Kl Pb(NO3)2 + 2NaOH
PbCl2 + 2NaNO3 Pb(CO3)2 + 2Na2NO3 PbI2 + 2KNO3 Pb(OH)2 + 2NaNO3 CuSO4 + 2Na2CO3
CuSO4 + 2Kl CuSO4 + 2NaOH CuCO3 + Na2SO4 CuI2+K2SO4 Cu(OH)2 + Na2SO4
FeSO4 + Na2CO3 FeSO4 + 2Kl FeSO4 + NaOH FeCO3 + Na2SO4 FeI2 + K2SO4
Fe(OH)2 + Na2SO4 - - Precipitation Reactions 1) Silver nitrate + sodium chloride 2) Silver nitrate + sodium bromide 3) Silver nitrate + potassium iodide 4) Lead nitrate + sodium chloride 5) Lead nitrate + sodium bromide Precipitation Reactions 6) Lead nitrate + potassium iodide
7) Lead nitrate + sulfuric acid 8) Barium chloride + sulfuric acid 9) Copper sulfate + sodium carbonate 10) Magnesium sulfate + sodium carbonate
