5.5 explain the uses of aluminium and iron, in terms of their properties

Uses (due to properties) of iron
The properties of iron are all the usual properties of a metal. However, adding something to iron can change its properties (as expected), this means it an be made suitable for many different things. Fir example...
- Pure iron (wrought iron) is malleable ('bendable'), this means it is used to make things such as ornamental gates and railings (as it can be twisted into pretty shapes etc)
- Cast iron (a mixture of iron, carbon and silicon) is very hard, this means it is used to make things such as manhole covers (that undergo lots of pressure, from vehicles, daily) and also some cooking pans
- Steel (an alloy made of iron) is harder than pure (wrought) iron but is still malleable and can also be welded together. It can be easily hammered into sheets and, because of this, it is great for making things that need thin hard metal, such as car bodies and girders for construction.

NOTE: a downside to using iron is that it rusts easily. However, stainless steel (an alloy of iron and chromium) will not rust, because of this, it is used in making things such as cutlery and pans that are exposed to water often (when cleaning etc).

Uses (due to properties) of aluminium
Aluminium is slightly different to iron, it also has all the main properties of a metal however, it doesn't corrode easily. This is because it quickly reacts with oxygen in the air, producing aluminium oxide (which forms as a protective layer around the aluminium, stopping any further reaction taking place) - this stops corrosion.

Due to its non-corroding property, it is used to make products that often come into contact with water (such as coke cans etc).

Aluminium is also a lot less dense than iron, which consequently makes it lighter (less particles per certain area etc), this means it is useful for making things when the weight of a metal frame needs to be taken into consideration (for example, when producing a bicycle frame or aeroplane body)

5.4 describe and explain the main reactions involved in the extraction of iron from iron ore (hematite), using one, limestone and air in a blast furnace

In order to extract iron from hematite (iron ore) you need a blast furnace, coke (for reducing the iron oxide to iron metal) and limestone (for taking away impurities). The process is as follows...



- Hot air is blasted into the furnace (hence name, blast furnace), this makes the coke burn much faster than normal, and also raises the temperature to around 1500ºC. The coke burns to produce carbon dioxide (C + O2 ---> CO2)
- The CO2 then reacts with unburnt/leftover coke, producing carbon monoxide (CO2 + C ---> 2CO)
- The carbon monoxide will then react with the iron ore, producing iron. (3CO + Fe2O3 ---> 3CO2+ 2Fe)
- The limestone removes the silicon dioxide (SiO2) that is the main impurity. This happens as the limestone is decomposed by the heat into calcium oxide and carbon dioxide (CaCO3 ---> CaO + CO2). The calcium oxide then reacts with the silicon dioxide forming calcium silicate, aka slag (CaO + SiO2---> CaSiO3).
- The iron and slag are both molten so sink to the bottom of the furnace. However, slag is less dense than iron so the slag sits into of the iron, they are both tapped off.

NOTE: although the slag is useless in this, the process is still sustainable as the slag is not wasted. It can be used in fertilisers and road building (bit random, i know)

NOTE NOTE: It is very important to understand that this is a reduction reaction (the iron is reduced as it loses oxygen)

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5.3 write ionic half-equations for the reactions at the electrodes in aluminium extraction.

If you are looking at this post I assume you know what the electrolysis of aluminium oxide is and the implications, if you do not know, this may help... 5.2

The equations are as follows...

At the anode (positive electrode): 2O²⁻ ---> O₂ + 4e− 

At the cathode (negative electrode): Al³⁺ + 3e− ---> Al

5.2 describe and explain the extraction of aluminium from purified aluminium oxide by electrolysis including: i the use of molten cryolite as a solvent and to decrease the required operating temperature ii the need to replace the positive electrodes iii the cost of the electricity as a major factor

As aluminium is more reactive than carbon, you use electrolysis to extract aluminium from its ore. The main ore of aluminium is bauxite (which is mined, incase you were wondering). NOTE: after you purify bauxite, a white powder is left (this is aluminium oxide, Al₂O₃).

First lets start with the electrolysis process, if you are familiar with this just skip to 'problems'...

1- Aluminium oxide is melted (to form molten aluminium oxide), this contains free ions, meaning it will conduct electricity.

2- The positive Al³⁺ ions are attracted to the negative electrode (the cathode, this is lining the electrolysis cell). Here, the positive Al³⁺ ions gain electrons (3 to be precise, to balance their charge), and they turn into aluminium atoms, which sinks to the bottom as molten aluminium.

3- The negative O²⁻ ions are attractive to the positive electrodes (anodes). Are, they lose electrons. They will then react together to form O₂ (oxygen gas), which will occasionally react with the anodes forming carbon dioxide (CO₂). 

NOTE: this is a redox reaction as reduction and oxidation both take place. The equations are as follows...

at the cathode: Al³⁺ + e− ---> Al

at the anode: 2O²⁻ ---> O₂ + 4e−

The complete decomposition reaction: aluminium oxide ---> aluminium + oxygen

Problems

However, the melting point of Al₂O₃ is very high (around 2000ºC, if you were wondering). Aluminium oxide is dissolved in molten cryolite (a less common ore of aluminium), this lowers the temperature needed to melt, so less fuel is used in heating the aluminium oxide = saves energy. (NOTE: using cryolite only lowers the melting temperature to around 900ºC so, although it lowers the temperature and saves lots of energy (and also makes the extraction cheaper and easier), a very high temp is still needed.

The positive electrodes (anodes) are made of graphite (a form of carbon). During the electrolysis process, the O²⁻ ions react with the carbon anodes, forming carbon dioxide (CO₂). This means that the anodes will eventually have to be replaced, as they will start to erode where they react with oxygen.

Lots of energy is needed to heat the aluminium oxide to 900ºC and therefore electrolysis of aluminium is very expensive. Furthermore, electrolysis itself uses lots of electricity.

5.1 explain how the methods of extraction of the metals in this section are related to their positions in the reactivity series

Methods of extraction are closely linked to the position of a metal in the reactivity series. This means that, by looking at the reactivity series, you can easily see the best way to extract that metal. They are grouped as follows...

All (and only) metals below carbon in the reactivity series (zinc, iron and tin)can be extracted by a reduction reaction with carbon (method: heat the ore with carbon monoxide). This happens because the carbon is more reactive, so for example (when extracting iron from iron oxide) the carbon displaces the iron, to form carbon dioxide and iron (NOTE: this is why we react carbon monoxide, so carbon dioxide is formed, if we were to react just carbon, carbon monoxide would we be formed which is poisonous).

If an element is more reactive than carbon (higher than carbon in the reactivity series), carbon will not displace it. This means that all metals more reactive than carbon have to be displaced with electrolysis (this separates the metal from the other elements in the compound using electricity).

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