Classwork Exercise and Series (Chemistry – SS3): Metal And Their Compounds

Aluminium

Aluminium is the most common metal in the Earth’s crust, making up 7.5% by mass. Its main ore is bauxite-a clay mineral which you can think of as impure aluminium oxide. It is the most important element in group III.

Extraction of Aluminium
Aluminium is obtained largely from the ore bauxite (Al₂O₃.2H₂O). Its production is a two-step process: the purification of bauxite and extraction by electrolysis.

Purifying the bauxite (aluminium oxide) – the Bayer Process

Crushed bauxite is treated with moderately concentrated sodium hydroxide solution. The concentration, temperature and pressure used depend on the source of the bauxite and exactly what form of aluminium oxide it contains. Temperatures are typically from 140°C to 240°C; pressures can be up to about 35 atmospheres.

High pressures are necessary to keep the water in the sodium hydroxide solution liquid at temperatures above 100°C. The higher the temperature, the higher the pressure needed.

With hot concentrated sodium hydroxide solution, aluminium oxide reacts to give a solution of sodium aluminate (III) (NaAl(OH)₄).

Extraction of Aluminium by Electrolysis

After purification, aluminium oxide is mixed with cryolite (sodium aluminium fluoride) Na₃AlF₆ to lower the melting point from 2000º to 1000º, which saves money. This mixture is heated and the molten liquid used as the electrolyte. Both electrodes are made of graphite (carbon).  The anode (+ve) is graphite and the cathode (-ve) is a graphite lining to a steel case.

The anode disintegrates. The hot oxygen produced here reacts with the hot carbon anode to give carbon dioxide. Hence it must be replaced regularly.

Aluminium ions are attracted to the cathode (the negative electrode) and are reduced to aluminium by gaining electrons.

Al³⁺ (l) + 3e^– ——> Al (l)

The molten aluminium produced sinks to the bottom of the cell.

The oxide ions are attracted to the anode and lose electrons to form oxygen gas.

2O^2- (l) ——> O₂ (g) + 4e^–

Note: The extraction of aluminium is an expensive process because the large amount of electricity needed to keep the electrolytes molten is expensive. Hence using cryolite saves energy and money, as it acts as a solvent for the aluminium oxide and melts at a much lower temperature.

Physical Properties of Aluminium

1. Aluminium is a silvery white metal which is comparatively soft
2. It is a strong, malleable metal element
3. It has a low density
4. It is resistant to corrosion
5. It is a good conductor of heat and electricity.
6. It can be polished to give a highly reflective surface.

Chemical Properties of Aluminium

• Action with air: Aluminium burns in air at high temperature to form the oxide and the nitride

            4Al (s) + 3O₂ (g) ——> 2Al₂O₃ (s)

• Reaction with non-metals: On heating, aluminium combines directly with non-metals like the halogens, sulphur, nitrogen, phosphorus and carbon with evolution of heat

            2Al (s) + 3Cl₂ ——> 2AlCl₃ (s)

• Action with acids: Aluminium reacts more rapidly with the concentrated hydrochloric acid to displace hydrogen but more slowly with dilute one

            2Al (s) + 6HCl (aq) ——> 2AlCl₃ (aq) + 3H₂ (g)

• Reaction with alkalis: Aluminium reacts with both sodium and potassium hydrogen solutions giving hydrogen gas and soluble tetrahydroxoaluminate (III)

2Al (s) + 2NaOH (aq) + 6H₂O (l) ——> 2NaAl(OH)₄ (aq) + 3H₂ (g)

• Reaction with Iron (III) oxide: Aluminium reduces iron (III) oxide to molten iron. The reaction is used in thermit process and it gives out a great deal of energy

            2Al (s) + Fe₂O₃ (s) ——> Al₂O₃ (s) + 2Fe (s)

Uses

1. Low density and strength make aluminium ideal for construction of aircraft, lightweight vehicles, and ladders.
2. An alloy of aluminium called duralumin is often used instead of pure aluminium because of its improved properties.
3. Easy shaping and corrosion resistance make aluminium a good material for drink cans and roofing materials.
4. Corrosion resistance and low density leads to its use for greenhouses and window frames.
5. Good conduction of heat leads to its use for boilers, cookers and cookware.
6. Good conduction of electricity leads to its use for overhead power cables hung from pylons (low density gives it an advantage over copper).
7. High reflectivity makes aluminium ideal for mirrors, reflectors and heat resistant clothing for fire fighting.

Copper

Copper was one of the first metals discovered and used by man. It is a stable metal readily obtained from its compounds. Copper ores are widely found around the world. The main ores are copper pyrites (CuFeS2), malachite (CuCO3.Cu (OH)2), chalcocite (CuS2) and cuprite (CuO).

Extracting copper from its ores

The method used to extract copper from its ores depends on the nature of the ore. Sulphide ores such as chalcopyrite (copper pyrites) are converted to copper by a different method from silicate, carbonate or sulphate ores.

The process:

The concentrated ore is heated strongly with silicon dioxide (silica) and air or oxygen in a furnace or series of furnaces.

• The copper(II) ions in the chalcopyrite are reduced to copper(I) sulphide (which is reduced further to copper metal in the final stage).
• The iron in the chalcopyrite ends up converted into an iron(II) silicate slag which is removed.
• Most of the sulphur in the chalcopyrite turns into sulphur dioxide gas. This is used to make sulphuric acid via the Contact Process.

An overall equation for this series of steps is:

The end product of this is called blister copper – a porous brittle form of copper, about 98 – 99.5% pure.

Purification of copper

When copper is made from sulphide ores by the first method above, it is impure. The blister copper is first treated to remove any remaining sulphur (trapped as bubbles of sulphur dioxide in the copper – hence “blister copper”) and then cast into anodes for refining using electrolysis.

Electrolytic refining

The purification uses an electrolyte of copper(II) sulphate solution, impure copper anodes, and strips of high purity copper for the cathodes.

The diagram shows a very simplified view of a cell.

For every copper ion that is deposited at the cathode, in principle another one goes into solution at the anode. The concentration of the solution should stay the same.

All that happens is that there is a transfer of copper from the anode to the cathode. The cathode gets bigger as more and more pure copper is deposited; the anode gradually disappears.

In practice, it isn’t quite as simple as that because of the impurities involved.

Physical properties of copper

1. Copper is a heavy, reddish-brown metal
2. It is very malleable and ductile
3. It has a density of 8.95 g cm^-1
4. It is a good conductor of heat and electricity
5. It has a high melting point of 1083^oC
6. It has a boiling point of 2300^oC
7. It also forms alloys very readily

Chemical properties of Copper

• Reaction with Air: It is resistant to pure dry air, but over a long period of time in a moist, impure atmosphere, it becomes coated with green, basic copper(II) tetraoxosulphate (VI) CuSO₄.3Cu(OH)₂ and trioxocarbonate (IV).

On heating in air or oxygen, copper is readily oxidized to give black copper (II)oxide

2Cu(s) + O₂(g) ——–> 2CuO(s)

• Effect of Acids: Copper is lower than hydrogen in the electrochemical series, hence, it is not capable of displacing hydrogen from dilute acids. It is however attacked by oxidizing acids like trioxonitrate (V) acid and tetraoxosulphate (VI) acids

3Cu(s) + 8HNO₃(aq) ——> 3Cu(NO₃)₂(s) + 4H₂O(l) 2NO(g)

• Displacement Reaction: Because of its low position in the activity and electrochemical series, copper is easily displaced from its compound

CuSO₄(aq) + Fe(s) ——-> FeSO₄(aq) + Cu(s)

• Hydrogen gas reduces copper oxides to the metal

CuO(s) + H₂(g) ——> Cu(s) + H₂O(l)

Test For Copper(II) Ions

With sodium hydroxide: Add a few drops of sodium hydroxide solution to a solution of copper salt. The formation of a blue precipitate which is insoluble in excess sodium hydroxide confirms the presence of copper(II) ions

Cu²⁺(aq) + 2NaOH(aq) ——-> Cu(OH)₂(s) + Na⁺(aq)

Uses of copper

Amongst other things copper is used for:

1. Electrical wiring. It is a very good conductor of electricity and is easily drawn out into wires.
2. Domestic plumbing. It doesn’t react with water, and is easily bent into shape.
3. Boilers and heat exchangers. It is a good conductor of heat and doesn’t react with water.
4. Baking brass. Brass is a copper-zinc alloy. Alloying produces a metal harder than either copper or zinc individually. Bronze is another copper alloy – this time with tin.
5. Coinage. In the UK, as well as the more obvious copper-coloured coins, “silver” coins are also copper alloys – this time with nickel. These are known as cupronickel alloys. UK pound coins and the gold-coloured bits of euro coins are copper-zinc-nickel alloys.

Iron

Iron is the most important element in the industry. It is the second most abundant element on the earth crust after aluminium, but often occurs as a free metal.

The common ores are haematite found in united state, Australia and USSR. It can also occur as impure iron (III) oxide (Fe₂O₃), Magnetite or magnetic iron ore (Fe₃O₄) is found in sweden and in North America.

Siderite or spathic iro ore, (FeCO₃), found in Great Britain. Iron also occurs as iron pyrites (FeS₂) and limonite (Fe₂O₃.3H₂O).

Iron is widely present as trioxosilicate (IV) in clay soils. Iron ore is available in Itakpe, Ajaokuta, Jebba and Lokoja all in Kwara State (Nigeria).

Extraction of Iron

The extraction of iron from iron ore (haematite), using coke, limestone and air in a blast furnace

Haematite is basically iron oxide, and the oxygen must be removed to leave the iron behind. Reactions in which oxygen is removed are called reduction reactions. Since carbon is more reactive than iron, it can displace the iron from its oxide. Hence the method for extraction of iron is called ‘reduction by carbon’.

The iron ore, coke and limestone enter the blast furnace at the top. The hot waste gases at the top of the furnace are piped away and used to heat the air blast at the bottom.

Coke is impure carbon, and it burns in the hot air blast to form carbon dioxide. This is a strongly exothermic reaction which makes it an important reaction, as it helps heat up the blast furnace.

C (s) + O₂ (g) ——> CO₂ (g)

At high temperatures in the furnace, the carbon dioxide is reduced by more carbon to give carbon monoxide.

CO₂ (g) + C (s) ——> 2CO (g)

It is the carbon monoxide which is the main reducing agent in the furnace-especially in the cooler parts.

Assuming that the iron ore is haematite, Fe₂O₃:

Fe₂O₃ (s) + 3CO (g) ——> 2Fe (l) + 3CO₂ (g)

Due to the high temperatures, the iron produced melts and flows to the bottom of the furnace, where it can be tapped off.

In the hotter parts of the furnace, some of the iron oxide is also reduced by carbon itself.

Fe₂O₃ (s) + 3C (s) ——> 2Fe (l) + 3CO (g)

Notice that carbon monoxide is formed, rather than carbon dioxide, at these temperatures.

However some use this equation instead:

iron oxide + carbon    →    iron + carbon dioxide

2Fe₂O₃ + 3C    →    4Fe + 3CO₂

The limestone is added to the furnace to remove impurities in the ore which would otherwise clog the furnace with solid material.

The furnace is hot enough for the limestone (calcium carbonate) to undergo thermal decomposition. It splits up into calcium oxide and carbon dioxide. This is an endothermic reaction (it absorbs heat) and it is important not to add too much limestone to avoid cooling the furnace.

CaCO₃ (s) ——> CaO (s) + CO₂ (g)

Calcium oxide is a basic oxide, and its function is to react with acidic oxides such as silicon dioxide, SiO₂. Silicon dioxide is the main constituent of sand, and is typical of the sort of impurities that need to be removed from the furnace.

CaO (s) + SiO₂ (s) —–> CaSiO₃ (l)

The product is calcium silicate. This melts and trickles to the bottom of the furnace as a molten slag, which floats on top of the molten iron as it is less dense, and can be tapped off separately. Slag is used

Uses of iron

• Cast iron

Molten iron straight from the furnace can be cooled rapidly and solidified by running it into sand moulds. This is known as pig iron. If the pig iron is remelted and cooled under controlled conditions, cast iron is formed. This is very impure iron, containing about 4% carbon as its main impurity. Although cast iron is very hard, it is also very brittle and tends to shatter if it is hit hard. It is used for things like manhole covers, gutterings and drainpipes, and cylinder blocks in car engines.

• Mild steel

Mild steel is iron containing up to about 0.25% carbon. This small amount of carbon increases the hardness and strength of the iron. It is used for (among other things) wire, nails, car bodies, ship building, girders and bridges.

• Wrought iron

This is pure iron. It was once used to make decorative gates and railings but has now been largely replaced by mild steel. The purity of the iron makes it very easy to work because it is fairly soft, but the softness and lack of strength mean that it isn’t useful for structural purposes.

• High-carbon steel

High carbon steel is iron containing up to 1.5% carbon. Increases the carbon content makes the iron harder, but at the same time it gets more brittle. High-carbon steel is used for cutting tools and masonry nails. Masonry nails are designed to be hammered into concrete blocks or brickwork where a mild steel nail would bend. If you miss-hit a masonry nail, it tends to break into two bits because of its increased brittleness.

• Stainless steel

Stainless steel is an alloy of iron with chromium and nickel. Chromium and nickel form strong oxide layers in the same way as aluminium, and these oxide layers protect the iron as well. Stainless steel is therefore very resistance to corrosion.

Obvious uses include kitchen sinks, saucepans, knives and forks, and gardening tools. But there are also major uses for it in the brewing, dairy and chemical industries where corrosion-resistant vessels are essential.

Types of iron	Iron mixed with	Some uses
Wrought iron	(pure iron)	Decorative work such as gates and railings
Mild steel	Up to 0.25% carbon	Nails, car bodies, ship building, girders
High-carbon steel	0.25-1.5% carbon	Cutting tools, masonry nails
Cast iron	About 4% carbon	Manhole covers, guttering, engine blocks
Stainless steel	Chromium and nickel	Cutlery, cooking utensils, kitchen sinks

Physical Properties of Iron

1. Pure iron is a grey metal with density 7.8g cm-3
2. It melts at 1535^oC and boils at 2800^oC
3. It is very malleable and ductile
4. It is a good conductor of heat and electricity
5. It can easily be magnetized

Chemical Properties of Iron

• Reaction With Air: It burns in air to form reddish hydrated iron (III) oxide of variable composition

4Fe(s) + 3O₂(g) + 2xH₂O(l) ——-> 2Fe₂O₃.xH₂O

• When finely divided iron is heated in air it burns at high temperature to form magnetic iron oxide, which behaves like a compound oxide

3Fe(s) + 2O₂(g) ——> Fe₃O₄ (s)

• Reaction With Steam: When steam is passed over red-hot iron filings, tri iron tetraoxide and hydrogen are produced and the reaction is reversible

Fe(s) + 4H₂0 <——> Fe₃O₄(s) + 4H₂(g)

Tin

Tin is a common element in group (IV). Tin metal does not occur naturally, it occurs as the ore cassiterite or tin (IV) oxide in gravels and alluvial deposits, e.g. in Jos plateau (Nigeria), Malaysia, Bolivia and Indonesia.

Extraction of Tin

1. Tin exists as stanum(IV) oxide, SnO₂ in the mineral cassiterite, that is tin. Tin contains a lot of foreign matter such as sand, soil, sulphur, carbon and oil.
2.  Firstly the tin ore is made concentrated by the method of floatation. In this process, the ore is crushed and shaken in oily water. The foreign matter such as sand and soil drown while the tin ore sticks to the oil and floats on the surface of the water.

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3. The tin ore is then collected and roasted to take away foreign matter such as carbon, sulphur and oil.

4.  Lastly, the tin ore is mixed with carbon in the form of charcoal and is heated in a blast furnace at a high temperature.

5.  Stanum(IV) oxide in the ore is reduced to tin by the reducing agent carbon and carbon monoxide.

SnO₂  (s) + 2C(s) → Sn (s) + 2CO (g)

SnO₂  (s) + C(s) → Sn (s) + CO₂ (g)

SnO₂  (s) + 2CO(s) → Sn (s) + 2CO₂ (g)

6.  The melted tin that is formed collects at the base of the furnace and then is channeled out into a mould to form tin ingot.

Physical Properties of Tin

• Tin is a white lustrous metal

• It exists in three allotropic forms which have different densities

Grey tin <—–> white tin <——> rhombic tin

5.76 gcm-3           7.28gcm-3             6.6gcm-3Top of Form

• The melting point is 230^oC and boiling point of 2270^oC

• It is a good conductor of heat and electricity

• It is malleable

• It is not ductile enough to be drawn into wires

Chemical Properties of Tin

• Reaction with oxygen: Tin being unreactive only combines with oxygen above 1200^oC to form tin (IV) oxide

Sn(s) + O₂ —–> SnO₂(s)

• Reaction with acids: Tin reacts slowly with dilute acids but vigorously with hot concentrated tetraoxosulphate (VI) acid and sulphur (IV) oxide is evolved

Sn(s) + 2H₂SO₄(aq) ——> SnSO₄(aq) + SO₂(g) + 2H₂O(l)

• Reaction with Alkalis: Tin reacts with hot sodium hydroxide producing trioxostannate (IV) salts and hydrogen

Sn(s) + 2NaOH(aq) + H₂O(l) ——-> Na₂SnO₃(aq) + 2H₂(g)

• Reaction with non-metals: Tin combines with chlorine when heated to give tin (IV) chloride

Sn(s) + 2Cl₂(g) ——> SnCl₄(s)

Uses of Tin

1. It is widely used to coat iron as a protection against corrosion e.g. the cans used for food storage
2. It is also used in many alloys such as bronze (with copper), solder (Sn, Pb)Bottom of Form

Summary of Extraction and Uses of Metals

Order of reactivity	Symbol	Method of Extraction
Potassium	K	Electrolysis

The metal compound is:

1.           Melted, then
2.           Has electricity passed through it

 

These metals are very reactive and are above carbon in the reactivity series, so they cannot be reduced by it. As they are very reactive, the make very stable compounds that requires a lot of energy to separate into its elements. So electrolysis is used. SodiumNa LithiumLi CalciumCa MagnesiumMg AluminiumAl

ZincZn. Reduction by carbon

e.g. ZnO+Cà Zn + CO

Or sometimes the carbon monoxide is the reducing agent-here think of reduction as ‘taking oxygen away’ to leave pure metal. Carbon is cheap and can also be used as the source of heat. If the ore is a sulphide, it is roasted first to get the oxide. Roasting is a process where is basically heating the ore in air. IronFe

TinSn LeadPb CopperCu These metals can be found uncombined, as the metal itself because they are very unreactive. We say they are found native. (Copper and silver are often found as ores but they are easy to extract by roasting the ore.) SilverAg GoldAu PlatiniumPt

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