Chapter 7: Metallurgy
In December 1884, workmen in Washington finished the tallest stone structure in the world and set a small pointed cap on its very top. The cap was a pyramid of metal, less than three kilograms of it, no taller than your hand. It was made of aluminium, and at that moment it was one of the largest pieces of aluminium anyone had ever cast.
That sounds like an odd choice today. Aluminium is the metal of window frames, cooking vessels, foil and aircraft bodies. It is the most plentiful metal in the earth’s crust — there is more of it under our feet than iron. Yet in 1884 an ounce of aluminium cost about as much as an ounce of silver, and a lump of it was a fitting ornament for a national monument.
The puzzle is this. The aluminium was always there, in ordinary clay, in the red soil of the Deccan, in the rock we call bauxite. What was missing was not the element but a way of tearing it away from the oxygen it clings to. Within two years of that capstone being fitted, two young chemists working independently found that way. The price of aluminium collapsed, and it has never risen again.
Diagram
Fig. 7.0-1 — The Aluminium Cap Of A Stone Monument Tall, tapering stone obelisk drawn from a low angle against a plain sky, with a magnified inset of its apex; label the small pyramid-shaped aluminium tip, its height marked as about 22 cm, and the surrounding marble capstone. Keep the inset simple and diagrammatic rather than photographic.
7.1 Metals, Non-Metals And Metalloids
One hundred and eighteen elements are known to us today. They are spread through the earth’s crust, some as the free element and most of them locked up in compounds.
They are not spread evenly. Just two elements, oxygen and silicon, make up nearly three-quarters of the crust between them, and everything else shares what is left. Table 7.1 gives the approximate share of some common elements.
Table 7.1 Approximate relative abundance of some elements in the earth’s crust
| Element | Abundance | Element | Abundance |
|---|---|---|---|
| Oxygen | 46.6% | Silicon | 27.7% |
| Aluminium | 8% | Iron | 5% |
| Calcium | 3.6% | Sodium | 2.8% |
| Potassium | 2.8% | Magnesium | 2.1% |
| Hydrogen | 0.14% | Phosphorus | 0.12% |
Read the table once more and notice which elements are metals. Aluminium heads that list, with iron next, and both are far commoner than the metals we think of as precious.
On the basis of their properties, the elements are sorted into metals and non-metals. Of the known elements only 17 are non-metals, so the great majority are metals.
7.1.1 What Makes An Element A Metal
Take any familiar metal — a copper wire, an aluminium sheet, an iron rod — and a list of shared properties builds itself. A metal can be hammered into sheets, which is to say it is malleable, and it can be pulled out into wire, which is to say it is ductile. Metals have high densities, high melting and boiling points and high tensile strength, and they conduct heat and electricity well.
That description tells you how a metal behaves in the hand. It does not tell you why. The chemical answer is shorter and far more useful, because it comes from what the atom does with its outermost electrons. A metal atom holds its valence electrons loosely and gives them up, and in giving them up it becomes a positively charged ion.
\[\mathrm{Na} \rightarrow \mathrm{Na^{+}} + e^{-}\]
This is the definition to carry through the rest of the chapter, because everything that follows — why some metals are found free, why others need electricity to release them — comes back to how firmly an atom keeps its outer electrons.
Metals In Use And In Living Things
People have used metals for a very long time. Copper was the first metal put to work by human beings, for utensils, weapons and tools, and gold, silver, copper, lead, iron, tin, antimony and mercury — mercury was once called quick silver — have all been known since ancient times.
The uses today are everywhere. Metals go into jewellery and coins, into utensils and furniture, and into heavy machinery, tools, transport vehicles and manufacturing equipment. Many act as catalysts in industry — in the hydrogenation of vegetable oils and in the manufacture of ammonia, sulphuric acid, nitric acid, dyes and drugs.
Living bodies need metals too, in tiny amounts. Iron sits at the centre of haemoglobin, the pigment that carries oxygen in your blood; zinc is present in insulin; cobalt is part of vitamin B-12; and magnesium is the metal at the heart of chlorophyll, the green pigment of plants.
7.1.2 Non-Metals And Their Importance
Non-metals are few in number, but life would be impossible without them. Their chemical behaviour is the mirror image of a metal’s: a non-metal atom takes in electrons rather than giving them up, and so it forms a negatively charged ion.
Hydrogen, the lightest element known, is used to harden vegetable oils into ghee, as a fuel, and in the manufacture of many compounds. It is also part of almost every organic compound. Proteins, fats, carbohydrates, enzymes and vitamins are all compounds of carbon, and all of them are essential for the growth of living things.
Oxygen gas in the air keeps respiration and combustion going. Nitrogen is the most abundant element in the atmosphere, and because it is unreactive it slows the rate of combustion and is used to keep packed food from spoiling.
The physical states divide up neatly as well. Every metallic element except mercury is a solid. Among the non-metals, 11 are gases, one — bromine — is a liquid, and the rest are solids. Astatine lies at the foot of the halogen group, and it is the rarest of all naturally occurring elements; so little of it exists that nobody has ever seen enough to say what it looks like.
Table 7.2 sets the two families side by side.
7.1.3 Metalloids
A few elements — the metalloids — refuse to sit in either family. They look metallic but behave chemically like non-metals in some reactions, and they conduct electricity, though far less freely than a metal does.
There are seven of them: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te) and polonium (Po). Silicon is the one you meet most often, since its middling conduction is exactly what makes it useful in the chips inside a phone or a computer.
7.1.4 Metals And Non-Metals In The Periodic Table
The two families are not scattered at random through the periodic table. Metals occupy the left-hand side and the whole middle block, while non-metals are gathered at the right-hand side, as Fig. 7.1-1 shows.
Group 13 is metallic apart from boron, and the metallic character of its elements increases on moving down the group. The elements of Groups 14, 15, 16 and 17 are mostly non-metallic, although a few members at the bottom of each of these groups begin to show some metallic nature. One non-metal sits out of place: hydrogen [H] is placed on the left-hand side of the table, above the alkali metals, because it too forms a positive ion.
Between the two regions runs a zig-zag staircase, and the metalloids lie along it. That line is the boundary between the metals and the non-metals.
Diagram
Fig. 7.1-1 — The Metal, Non-Metal And Metalloid Regions Of The Periodic Table Outline grid of the long form of the periodic table, groups 1 to 18 across and periods 1 to 7 down, with the f-block shown separately below; shade the left side and the middle block as metals and the top right corner as non-metals, and draw a bold zig-zag staircase line between them. Label the staircase as the metalloid boundary, mark the seven metalloid boxes B, Si, Ge, As, Sb, Te and Po along it, mark hydrogen at the top left with a note that it is a non-metal placed among the metals, and label Group 13 and the noble gas column.
7.2 Occurrence Of Metals In Nature
The earth’s crust is the biggest source of metals we have. A metal may be found there in either of two conditions: as the free element, or joined to other elements in a compound.
A metal found as the element itself, uncombined, is said to occur in the native state, which is also called the free state. A metal found joined to other elements is said to occur in the combined state.
Reactivity Decides The State
Which of the two states a metal is found in is not a matter of luck. It follows from how readily the metal gives up its electrons, and that is what the activity series records — a list of the metals arranged in decreasing order of reactivity, which you will meet in full in Section 7.8.
A metal low in that series is unreactive. It does not react under ordinary conditions with oxygen, water, carbon dioxide or the other reagents it meets in the ground, so it survives as the metal. Gold and platinum are the clearest cases, and both are found native.
Most metals, though, are reactive. Over millions of years they have combined with whatever was around them, and so they are found in the combined state — as oxides, carbonates, halides, sulphides and sulphates, mixed in with mud, clay, sand and stone.
The Form The Compound Takes
Which compound a metal is found as also follows a pattern. The activity series arranges the metals in decreasing order of electropositivity — the readiness of an atom to give away electrons and become a positive ion. You would therefore expect the metals near the top to be found joined to the most electronegative elements, the ones most eager to take those electrons. That is roughly what we find, as Table 7.3 shows.
Table 7.3 The usual combined form of some common metals
| Metal | Usually found as |
|---|---|
| Sodium, potassium | Chlorides |
| Calcium, magnesium | Carbonates |
| Aluminium | Oxide |
| Zinc, lead, copper, mercury, silver | Sulphides |
There is one further point about where these compounds end up. Insoluble compounds stay put and are found as rock. Soluble ones, such as the chlorides of the alkali metals, are dissolved out by rain over long periods and carried away, which is why so much of the world’s sodium chloride is found in sea water rather than in the ground.
7.3 Minerals And Ores
Since almost every metal is found combined, the material dug out of the ground is never the metal itself. It is a rock, and the metal’s compound is one ingredient of that rock.
A naturally occurring compound of a metal is called a mineral. In the ground, minerals come mixed with soil, mud, sand, silica (\(\mathrm{SiO_2}\)), limestone and rock. These earthy impurities have their own name — they are called the gangue, or the matrix.
Not Every Mineral Is Worth Working
A metal can occur in many different minerals, but a factory cannot use all of them. Extracting the metal must be cheap enough and easy enough to be worth doing. The minerals that pass that test are called ores.
Aluminium makes the point neatly. There is aluminium in ordinary clay, and clay is certainly a mineral of aluminium, but no one extracts the metal from clay — the process would cost far more than the aluminium is worth. Clay is a mineral of aluminium and not an ore of it.
So the two words are not interchangeable, and the relation between them runs one way only, as Fig. 7.3-1 shows. All minerals are not ores, but all ores are minerals.
Diagram
Fig. 7.3-1 — The Relation Between Minerals And Ores Two concentric ovals: the large outer oval labelled ‘Minerals — all naturally occurring compounds of metals’, and a smaller oval drawn wholly inside it labelled ‘Ores — minerals worth extracting the metal from’. Place the word ‘clay’ in the outer region only and the word ‘bauxite’ inside the inner oval, with a caption arrow noting that every ore is a mineral but every mineral is not an ore.
7.4 Common Ores Of Aluminium, Iron And Zinc
Every metal has its own short list of ores, and the list is worth knowing, because the first decision in any extraction — what to do with the rock once it is out of the ground — depends entirely on what the metal is combined with. Table 7.4 gives the main ores of the metals you will meet most often.
Table 7.4 Some common metals and their main ores
| Metal | Main ores |
|---|---|
| K | Carnallite (\(\mathrm{KCl\!\cdot\!MgCl_2\!\cdot\!6H_2O}\)) |
| Na | Rock salt (\(\mathrm{NaCl}\), also present in sea water) |
| Ca | Limestone and marble (\(\mathrm{CaCO_3}\)), gypsum (\(\mathrm{CaSO_4\!\cdot\!2H_2O}\)) |
| Mg | Carnallite (\(\mathrm{KCl\!\cdot\!MgCl_2\!\cdot\!6H_2O}\)), magnesite (\(\mathrm{MgCO_3}\)) |
| Al | Bauxite (\(\mathrm{Al_2O_3\!\cdot\!2H_2O}\)), cryolite (\(\mathrm{Na_3AlF_6}\)), feldspar (\(\mathrm{KAlSi_3O_8}\)), mica (\(\mathrm{KAl_3Si_3O_{10}(OH)_2}\)) |
| Zn | Zinc blende (\(\mathrm{ZnS}\)), calamine (\(\mathrm{ZnCO_3}\)) |
| Fe | Red haematite (\(\mathrm{Fe_2O_3}\)), brown haematite (\(\mathrm{2Fe_2O_3\!\cdot\!3H_2O}\)) |
| Pb | Galena (\(\mathrm{PbS}\)) |
| Cu | Copper pyrites (\(\mathrm{CuFeS_2}\)) |
| Hg | Cinnabar (\(\mathrm{HgS}\)) |
| Ag | Argentite (\(\mathrm{Ag_2S}\)), horn silver (\(\mathrm{AgCl}\)) |
Three of these metals — aluminium, iron and zinc — are worth looking at more closely, because they are the three whose extraction this chapter follows.
The Ores Of Aluminium
The ore from which aluminium is actually extracted is bauxite, hydrated aluminium oxide. The other aluminium minerals in Table 7.5 matter for a different reason: cryolite has a part to play later in the extraction itself, and corundum is the pure anhydrous oxide.
Table 7.5 Ores of aluminium
| Ore of Al | Chemical name | Formula |
|---|---|---|
| Bauxite | Hydrated aluminium oxide | \(\mathrm{Al_2O_3\!\cdot\!2H_2O}\) |
| Cryolite | Sodium aluminium fluoride | \(\mathrm{Na_3AlF_6}\) |
| Corundum | Anhydrous aluminium oxide | \(\mathrm{Al_2O_3}\) |
The Ores Of Iron
Iron is mainly extracted from its chief ore haematite, \(\mathrm{Fe_2O_3}\). Notice in Table 7.6 how varied the rest of the list is — an oxide, a hydrated oxide, a mixed oxide, a sulphide and a carbonate — because iron combines readily with a great many things.
Table 7.6 Ores of iron
| Ore of Fe | Chemical name | Formula |
|---|---|---|
| Red haematite | Anhydrous ferric oxide | \(\mathrm{Fe_2O_3}\) |
| Brown haematite | Hydrated ferric oxide | \(\mathrm{2Fe_2O_3\!\cdot\!3H_2O}\) |
| Magnetite | Triferric tetraoxide | \(\mathrm{Fe_3O_4}\) |
| Iron pyrites | Iron disulphide | \(\mathrm{FeS_2}\) |
| Siderite | Ferrous carbonate | \(\mathrm{FeCO_3}\) |
The Ores Of Zinc
The main ore of zinc is zinc blende, \(\mathrm{ZnS}\), from which the metal is extracted. Table 7.7 gives it with the other two zinc ores, which are the oxide and the carbonate.
Table 7.7 Ores of zinc
| Ore of Zn | Chemical name | Formula |
|---|---|---|
| Zinc blende | Zinc sulphide | \(\mathrm{ZnS}\) |
| Zincite | Zinc oxide | \(\mathrm{ZnO}\) |
| Calamine | Zinc carbonate | \(\mathrm{ZnCO_3}\) |
Look across all three tables and one thing stands out. Aluminium comes to us as an oxide, iron chiefly as an oxide, and zinc chiefly as a sulphide. That single difference decides the treatment each ore gets in the sections that follow.
7.5 The Stages In The Extraction Of A Metal
Getting a metal out of its ore takes several steps, some physical and some chemical. Which steps are needed, and in what form, depends on two things: the nature of the ore and the properties of the metal itself. There is no single method that works for every metal.
The whole business of winning metals from their ores has its own name — metallurgy.
The Five Stages
Whatever the metal, the route from rock to usable metal passes through the same five stages, set out in order in FC 7.1.
- Pulverisation — crushing and grinding the ore.
- Concentration, also called dressing or benefaction of the ore — removing the earthy impurities.
- Conversion of the concentrated ore to its oxide, by roasting or calcination.
- Reduction of the metallic oxide to the metal. When this is done by heating with a reducing agent in a furnace, along with a substance that carries off the remaining gangue, it is called smelting.
- Refining, or purification, of the metal obtained.
Flowchart
FC 7.1 — The Stages In The Extraction Of A Metal Ore → pulverisation → concentration (branches: gravity separation / froth floatation / magnetic separation / chemical method or leaching) → roasting and calcination → reduction (branches: chemical reduction / electrolytic reduction) → refining → pure metal
Why The Ore Is Crushed First
The ore comes out of the mine as lumps of rock, and nothing useful can be done with a lump. It is fed into big jaw crushers and then into ball mills, which grind it down to a fine powder. This first stage is pulverisation.
Grinding does two jobs at once. It breaks apart the mineral and the gangue, which were bound together in the rock, so that the next stage can separate them. It also exposes far more surface, and every chemical step that follows — the action of air, of a reagent, of heat — works on the surface of the powder. The finer the powder, the faster and the more complete the reaction.
7.6 Dressing Or Concentration Of The Ore
The powdered ore leaving the ball mill is still mostly rock. Mixed in with the mineral is the gangue — sand, silica, clay and stone — and none of it is wanted. Removing it is the second stage, called concentration of the ore.
There is no single way of doing this, because what separates the mineral from the gangue is some physical or chemical difference between them, and that difference is not the same for every ore. Four methods are in common use, compared in Table 7.8 and taken one at a time below.
7.6.1 Hydraulic Washing
Hydraulic washing is also called gravity separation or levigation, and it is the oldest trick in the book — the same one used by anyone panning for gold in a stream.
Principle: the difference in the densities of the ore and the gangue is the main criterion.
Method: the powdered ore is poured over a sloping, vibrating corrugated table cut with grooves, and a jet of water is allowed to flow over it. The denser ore particles are too heavy to be carried along, so they settle down into the grooves. The lighter gangue particles are washed down the slope by the water and run off the lower end. The arrangement is shown in Fig. 7.6-1.
Oxide ores of iron and tin are concentrated by this method.
Diagram
Fig. 7.6-1 — Concentration By Hydraulic Washing Side view of a sloping corrugated table mounted on a frame, with a curved pipe above the higher end delivering a spray of water over the powdered ore. Show the grooves running across the table with dense dark ore particles collecting in them, and lighter gangue particles being carried down the slope into a collecting trough below. Label: spray of water, denser ore particles settling in the grooves, sloping vibrating corrugated table, water and gangue running off.
7.6.2 Magnetic Separation
Some ores, or some of their impurities, are attracted to a magnet, and magnetic separation puts that to use. Where that is so, a magnet will do the sorting.
Principle: the magnetic properties of the ore, or of the impurity present with it.
Method: the crushed ore is fed onto a conveyor belt that runs around two metal wheels, one of which is a magnet. As the belt turns over the magnetic wheel, the non-magnetic particles fly off and fall straight down, while the magnetic particles cling to the belt a little longer and drop in a separate heap further along. Two clean piles are formed, as Fig. 7.6-2 shows.
This method is used to concentrate tin stone, \(\mathrm{SnO_2}\), which contains the magnetic oxide of iron, \(\mathrm{Fe_3O_4}\), as an impurity. Here it is the impurity that is magnetic and the ore that is not — the method works either way round.
Diagram
Fig. 7.6-2 — Electromagnetic Separation Side view of a hopper feeding powdered ore onto a horizontal conveyor belt that runs around two wheels; label the left-hand wheel as the magnetic wheel and the right-hand one as the ordinary wheel. Show arrows on the belt, non-magnetic particles falling in a heap directly below the end of the belt, and magnetic particles carried round and dropping in a separate heap closer to the magnetic wheel. Label: ore, magnetic wheel, non-magnetic wheel, magnetic gangue, non-magnetic ore.
7.6.3 Froth Floatation
Froth floatation is the method used for the sulphide ores, and it depends on a difference not in density but in the way two surfaces behave towards oil and water.
Principle: the process depends on the preferential wettability of the ore with oil, and of the gangue particles with water.
Method: the crushed ore is taken in a large tank containing pine oil and water. Potassium ethylxanthate is added to it, and a current of compressed air is blown in while the whole mass is agitated. The oil clings to the surface of the ore particles and will not wet the gangue.
The rising air bubbles, coated with oil, pick up the oil-wetted ore particles and carry them to the top, where they collect as a froth that can be skimmed off. The gangue particles, wetted by water and not by oil, are left behind and settle to the bottom of the tank. Fig. 7.6-3 shows the tank in section.
Sulphide ores such as zinc blende, \(\mathrm{ZnS}\), and galena, \(\mathrm{PbS}\), are concentrated by the froth floatation process.
Diagram
Fig. 7.6-3 — Concentration Of Ore By Froth Floatation Sectional view of a rectangular tank holding a mixture of powdered ore, water and pine oil, with a pipe entering from the top delivering compressed air and a second inlet on the left for the powdered ore, water and pine oil mixture. Show rising air bubbles carrying oil-coated ore particles upward, a layer of froth overflowing at the right into a collecting channel, and heavier gangue settled at the bottom. Label: compressed air, powdered ore with water and pine oil, air bubbles with oil-wetted ore, froth with sulphide ore, gangue or matrix.
7.6.4 Chemical Method Or Leaching
The three methods so far have all been physical — nothing has changed chemically. Leaching is different, because here the mineral is actually dissolved.
The ore is treated with a suitable reagent, an acid, a base or some other solution, chosen so that the ore dissolves in it while the impurities do not. The undissolved impurities are then removed by filtration. The filtrate, which is the solution of the ore, is treated with a second reagent that precipitates the ore back out in a pure form.
The leaching method is used for the concentration of the ores of silver, gold and aluminium. You will see it worked through in full detail in Section 7.11, where sodium hydroxide is the reagent that dissolves the ore of aluminium.
7.6.5 Fluxes And Slag
Even after concentration, the ore is not perfectly clean. A little gangue always survives into the next stages, and it must be dealt with when the ore is finally heated.
To remove it, a substance called a flux is mixed with the concentrated ore. The flux is chosen to react with whatever the remaining gangue is, and the product of that reaction is a fusible mass — one with a low enough melting point to run freely — called the slag.
If the gangue is acidic, such as silica, a basic flux such as limestone is added. Silica and limestone react to give calcium silicate, which is the slag:
silica + calcium carbonate → calcium silicate + carbon dioxide
\[\mathrm{SiO_2} + \mathrm{CaCO_3} \rightarrow \mathrm{CaSiO_3} + \mathrm{CO_2}\]
If instead the gangue is basic, such as ferrous oxide, an acidic flux is used. The principle is the same either way: acid needs base and base needs acid.
Slag is usually lighter than the molten metal, so it floats on top as a separate layer and can be run off easily, leaving the metal below.
7.7 Conversion Of The Concentrated Ore Into Its Oxide
The concentrated ore may be a sulphide, a carbonate or an oxide. Before the metal can be released from it, whatever it is, it is first turned into the oxide.
There is a plain reason for that. The reducing agents used in the next stage — carbon, carbon monoxide and hydrogen — all work by taking oxygen away, and they take it away readily. An oxide is therefore far easier to reduce to the metal than a sulphide or a carbonate is. So it is worth spending a stage converting the ore into the one compound that reduces most easily.
Two methods are used, and which one is chosen depends on what the ore is made of. Sulphide ores are roasted; carbonate ores and hydrated oxides are calcined.
7.7.1 Roasting
Roasting is the treatment given to a sulphide ore. The ore is heated strongly in a plentiful supply of air, but never so strongly that it melts — a molten mass would stop the air reaching the inside of it.
Four things happen while the ore is roasting.
- Moisture present in the ore is driven off.
- Organic matter is oxidised and removed.
- Impurities of sulphur, phosphorus and arsenic are oxidised to their volatile oxides, which escape as gases:
\[\mathrm{S} + \mathrm{O_2} \rightarrow \mathrm{SO_2}\]
\[\mathrm{P_4} + 5\,\mathrm{O_2} \rightarrow \mathrm{P_4O_{10}}\]
\[4\,\mathrm{As} + 3\,\mathrm{O_2} \rightarrow 2\,\mathrm{As_2O_3}\]
- Metal sulphides are oxidised to metal oxides. This is the main purpose of the whole operation:
zinc sulphide + oxygen → zinc oxide + sulphur dioxide
\[2\,\mathrm{ZnS} + 3\,\mathrm{O_2} \xrightarrow{\ 800^\circ\text{--}900^\circ\text{C}\ } 2\,\mathrm{ZnO} + 2\,\mathrm{SO_2}\]
\[2\,\mathrm{PbS} + 3\,\mathrm{O_2} \rightarrow 2\,\mathrm{PbO} + 2\,\mathrm{SO_2}\]
\[4\,\mathrm{FeS} + 7\,\mathrm{O_2} \rightarrow 2\,\mathrm{Fe_2O_3} + 4\,\mathrm{SO_2}\]
There is a fifth effect, easy to overlook but important. As the gases stream out of the solid, they leave holes behind, and roasting makes the ore porous. A porous ore heats evenly right through and reacts far better in the reduction stage that follows.
7.7.2 Calcination
Calcination is the treatment given instead to a carbonate ore. Such an ore has no sulphur in it, so there is nothing to burn off, and blowing air through it would achieve nothing. What such an ore needs is simply to be broken down by heat.
On calcination, metal carbonates and metal hydroxides decompose to give the metal oxide.
zinc carbonate → zinc oxide + carbon dioxide
\[\mathrm{ZnCO_3} \xrightarrow{\ \Delta\ } \mathrm{ZnO} + \mathrm{CO_2}\]
\[\mathrm{CaCO_3} \xrightarrow{\ \Delta\ } \mathrm{CaO} + \mathrm{CO_2}\]
Along with the carbon dioxide go the moisture, the water of hydration and the other volatile impurities. Here too the escaping gases leave the ore porous, so that it heats uniformly afterwards.
7.7.3 Roasting Compared With Calcination
The two processes are easy to confuse, because both involve heating an ore and both end with the oxide. Table 7.9 sets out where they differ.
7.8 Reduction Of Metal Oxides To Metals
After roasting or calcination we have the metal oxide. The metal is still combined, and the oxygen has to be taken away from it. That is what reduction means, and this is the stage at which the metal is finally born.
The usual way of doing it is to heat the oxide with a reducing agent — carbon, carbon monoxide or hydrogen — which takes the oxygen for itself and leaves the crude metal behind. But this does not work for every metal, and the choice of reducing agent depends on where the metal stands in the reactivity series.
7.8.1 The Activity Series Of Metals
The series is set out in Fig. 7.8-1. Potassium is kept at the top because it is the most reactive metal: it is the most electropositive, giving up its valence electron more readily than any other metal in the list. Gold, the least active, is placed at the bottom. Hydrogen, though a non-metal, is included in the series because it too can form a positive ion,
\[\mathrm{H} - e^{-} \rightarrow \mathrm{H^{+}}\]
and it is placed between lead and copper, the position its reactivity earns it when compared with the metals around it.
Diagram
Fig. 7.8-1 — The Activity Series Of Metals Vertical ladder of element symbols, most reactive at the top and least reactive at the bottom, in this order: K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au. Draw hydrogen in square brackets to show it is a non-metal. Add a downward arrow beside the ladder labelled ‘reactivity and electropositive character decrease’. On the right, bracket the ladder into three bands and label them: K to Al — extracted by electrolysis of the fused compound; Zn to Cu — extracted by reduction of the oxide with carbon; Hg, Ag and Au — obtained by heating the compound alone.
What The Series Tells You
Six statements follow from the order of the list, and between them they explain most of the chemistry of the rest of this chapter.
- Any metal higher up in the series displaces any metal below it from a solution of the latter’s salt.
- The larger the difference in the positions of the two metals in the series, the more rapidly does the displacement occur.
- Metals above hydrogen may displace hydrogen from water and from dilute acids; the metals below hydrogen cannot.
- The oxides of potassium, sodium, calcium, magnesium and aluminium cannot be reduced by hydrogen, carbon monoxide or carbon.
- The oxides and nitrates of mercury, silver and gold decompose to give the metal on being strongly heated.
- Metals below copper do not rust easily.
One caution about the series. It is drawn up from the reactivity of metals towards acid, air and water, so the position of one or two elements in it may differ from their position in the electrochemical series, which is based only on the tendency to lose valence electrons.
7.8.2 Metals High In The Series — Electrolytic Reduction
The oxides of the highly active metals — potassium, sodium, calcium, magnesium and aluminium — have a great affinity for oxygen. Their grip on it is stronger than the grip carbon, carbon monoxide or hydrogen can exert, so none of these common reducing agents can take the oxygen away. Heating such an oxide with coke simply leaves you with the oxide.
These metals are therefore obtained by electrolytic reduction: an electric current is passed through the fused salt of the metal, a halide or an oxide, using inert electrodes. The cathode is usually made of iron and the anode of graphite. Electricity does what no chemical reducing agent can — it hands the electrons straight to the metal ion.
Three examples show the pattern.
Potassium
Electrolyte: fused potassium bromide.
\[\mathrm{KBr} \rightleftharpoons \mathrm{K^{+}} + \mathrm{Br^{-}}\]
At the cathode: \(\ \mathrm{K^{+}} + e^{-} \rightarrow \mathrm{K}\)
At the anode: \(\ \mathrm{Br^{-}} - e^{-} \rightarrow \mathrm{Br}\), and then \(\ \mathrm{Br} + \mathrm{Br} \rightarrow \mathrm{Br_2}\)
Sodium
Electrolyte: fused sodium chloride.
\[\mathrm{NaCl} \rightleftharpoons \mathrm{Na^{+}} + \mathrm{Cl^{-}}\]
At the cathode: \(\ \mathrm{Na^{+}} + e^{-} \rightarrow \mathrm{Na}\)
At the anode: \(\ \mathrm{Cl^{-}} - e^{-} \rightarrow \mathrm{Cl}\), and then \(\ \mathrm{Cl} + \mathrm{Cl} \rightarrow \mathrm{Cl_2}\)
Calcium
Electrolyte: fused calcium chloride.
\[\mathrm{CaCl_2} \rightleftharpoons \mathrm{Ca^{2+}} + 2\,\mathrm{Cl^{-}}\]
At the cathode: \(\ \mathrm{Ca^{2+}} + 2e^{-} \rightarrow \mathrm{Ca}\)
At the anode: \(\ \mathrm{Cl^{-}} - e^{-} \rightarrow \mathrm{Cl}\), and then \(\ \mathrm{Cl} + \mathrm{Cl} \rightarrow \mathrm{Cl_2}\)
Notice that in every case the electrolyte is the fused, that is molten, salt, and never a solution.
Aluminium is also obtained from its oxide by electrolysis, and Sections 7.11 and 7.12 follow that extraction through from the ore.
7.8.3 Metals In The Middle — Reduction By Carbon
The metals in the middle of the activity series — zinc, iron, lead, copper and the like — are only moderately reactive. In nature they are usually found as sulphides or as carbonates, so the first job is the one described in Section 7.7: the sulphide ores are roasted and the carbonate ores calcined, and either way the oxide is obtained.
The oxide is then reduced with a suitable reducing agent. When the reducing agent is carbon, the operation is called the carbon reduction process.
For zinc, the roasted ore and coke are heated together in a retort, the heat being supplied by burning producer gas, a mixture of carbon monoxide and nitrogen.
zinc oxide + carbon → zinc + carbon monoxide
\[\mathrm{ZnO}\text{(s)} + \mathrm{C}\text{(s)} \xrightarrow{\ 1400^\circ\text{C}\ } \mathrm{Zn}\text{(s)} + \mathrm{CO}\text{(g)}\]
The zinc obtained in this way is known as zinc spelter.
Zinc oxide is a special case worth pausing on. It cannot be reduced by carbon monoxide or by hydrogen, because neither of them can overcome the affinity of zinc for oxygen. Only carbon, which is the stronger reducing agent of the three, is able to do it — which is why zinc oxide is reduced by heating with carbon and by nothing else.
Iron, lead and copper are less demanding. Their oxides are reduced by carbon, by carbon monoxide and by hydrogen alike.
Iron(II) Oxide
\[\mathrm{FeO} + \mathrm{C} \xrightarrow{\ \Delta\ } \mathrm{Fe} + \mathrm{CO}\]
\[\mathrm{FeO} + \mathrm{CO} \xrightarrow{\ \Delta\ } \mathrm{Fe} + \mathrm{CO_2}\]
\[\mathrm{FeO} + \mathrm{H_2} \xrightarrow{\ \Delta\ } \mathrm{Fe} + \mathrm{H_2O}\]
Iron(III) Oxide
\[\mathrm{Fe_2O_3} + 3\,\mathrm{C} \xrightarrow{\ \Delta\ } 2\,\mathrm{Fe} + 3\,\mathrm{CO}\]
\[\mathrm{Fe_2O_3} + 3\,\mathrm{CO} \xrightarrow{\ \Delta\ } 2\,\mathrm{Fe} + 3\,\mathrm{CO_2}\]
\[\mathrm{Fe_2O_3} + 3\,\mathrm{H_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{Fe} + 3\,\mathrm{H_2O}\]
Lead(II) Oxide
\[\mathrm{PbO} + \mathrm{C} \xrightarrow{\ \Delta\ } \mathrm{Pb} + \mathrm{CO}\]
\[\mathrm{PbO} + \mathrm{CO} \xrightarrow{\ \Delta\ } \mathrm{Pb} + \mathrm{CO_2}\]
\[\mathrm{PbO} + \mathrm{H_2} \xrightarrow{\ \Delta\ } \mathrm{Pb} + \mathrm{H_2O}\]
Lead(II) oxide is yellow, and the lead that forms is greyish white, so the change can be followed by eye.
Copper(II) Oxide
\[\mathrm{CuO} + \mathrm{C} \xrightarrow{\ \Delta\ } \mathrm{Cu} + \mathrm{CO}\]
\[\mathrm{CuO} + \mathrm{CO} \xrightarrow{\ \Delta\ } \mathrm{Cu} + \mathrm{CO_2}\]
\[\mathrm{CuO} + \mathrm{H_2} \xrightarrow{\ \Delta\ } \mathrm{Cu} + \mathrm{H_2O}\]
Copper(II) oxide is black, and the copper left behind is pink-brown. Pass dry hydrogen over heated copper(II) oxide in a tube and you see the black powder turn pink-brown while droplets of water collect at the cool end of the tube. Both observations together tell you that the oxygen has gone from the copper to the hydrogen.
7.8.4 Metals Low In The Series — Reduction By Heat Alone
The metals at the foot of the series are so unreactive that their oxides fall apart of their own accord when heated hard enough. No reducing agent is needed.
Mercury shows the sequence in full. Its ore is cinnabar, mercury(II) sulphide, which is first roasted in air to the oxide:
\[2\,\mathrm{HgS} + 3\,\mathrm{O_2} \rightarrow 2\,\mathrm{HgO} + 2\,\mathrm{SO_2}\]
and the oxide, on further heating, gives up its oxygen and leaves the metal:
\[2\,\mathrm{HgO} \xrightarrow{\ \text{above}\ 300^\circ\text{C}\ } 2\,\mathrm{Hg} + \mathrm{O_2}\]
Silver oxide behaves in the same way:
\[2\,\mathrm{Ag_2O} \xrightarrow{\ \text{above}\ 300^\circ\text{C}\ } 4\,\mathrm{Ag} + \mathrm{O_2}\]
Copper found in nature as \(\mathrm{Cu_2S}\) can be obtained by heating it in air alone. Part of the sulphide is first oxidised to the oxide, and the oxide then reacts with the sulphide that is left:
\[2\,\mathrm{Cu_2S} + 3\,\mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{Cu_2O} + 2\,\mathrm{SO_2}\]
\[2\,\mathrm{Cu_2O} + \mathrm{Cu_2S} \xrightarrow{\ \Delta\ } 6\,\mathrm{Cu} + \mathrm{SO_2}\]
Because the ore supplies both the oxidising step and the reducing step, this is called the auto reduction process.
Silver can also be obtained by displacement. A more electropositive metal such as zinc, added to a solution containing silver ions, pushes the silver out as the free metal:
\[\mathrm{Zn} + 2\,\mathrm{Ag^{+}}\text{(in solution)} \rightarrow \mathrm{Zn^{2+}} + 2\,\mathrm{Ag}\downarrow\]
This is the first of the six statements about the activity series being put to work: zinc is higher than silver, so it displaces silver from the salt solution.
7.8.5 Reactive Metals Used As Reducing Agents
There is one more route to a metal, and it turns the activity series to a different use. A highly reactive metal can itself act as a reducing agent, displacing a metal of lower reactivity from its compound.
To obtain manganese, its oxide is heated with aluminium powder:
\[3\,\mathrm{MnO_2}\text{(s)} + 4\,\mathrm{Al}\text{(s)} \rightarrow 3\,\mathrm{Mn}\text{(l)} + 2\,\mathrm{Al_2O_3}\text{(s)} + \text{heat}\]
Aluminium is a powerful reducing agent. When a mixture of aluminium powder and iron oxide is ignited, the iron oxide is reduced to the metal, and this process is called aluminothermy:
\[\mathrm{Fe_2O_3} + 2\,\mathrm{Al} \rightarrow 2\,\mathrm{Fe} + \mathrm{Al_2O_3} + \text{heat}\]
These displacement reactions are strongly exothermic. So much heat is given out that the metal produced is not merely hot but molten, and it runs out as a liquid — which is what makes the reaction useful.
7.9 Refining Of The Metal
The metal that comes out of the reduction stage is not pure. It is called the crude metal, and cleaning it up is called refining. Mixed in with it are two kinds of impurity: other metals and non-metals such as silicon and phosphorus, and unreduced oxides and sulphides of the metal itself.
There is no single refining method either. Which one is used depends on three things: the nature of the metal, the nature of the impurities present in it, and the purpose for which the metal is going to be used. A girder does not need the purity that an electrical conductor needs.
7.9.1 Electro-Refining
This method uses electrical energy to drive a chemical reaction, in an electrolytic cell. The cell has three parts: an electrolyte, and two conducting electrodes — a cathode and an anode — held in a vessel that does not itself conduct.
The arrangement is the opposite of what a student usually expects, so it is worth stating plainly. The impure metal is made the anode, and a thin sheet of the pure metal is made the cathode. The electrolyte is a salt solution of the very metal being refined, to which a little dilute sulphuric acid is added to increase the conductance. In preparing that solution the acid is always added slowly to the water, never water to the acid. The two electrodes are then suspended in the electrolyte, as Fig. 7.9-1 shows.
When the current is passed, two things happen at once. At the anode the impure metal dissolves, its atoms losing electrons and entering the solution as ions. At the cathode, metal ions from the solution gain electrons and are deposited as pure metal.
At the anode: \(\ \mathrm{M} - ne^{-} \rightarrow \mathrm{M^{n+}}\)
At the cathode: \(\ \mathrm{M^{n+}} + ne^{-} \rightarrow \mathrm{M}\) (M is a metal)
So the anode wastes away and the cathode grows, and an equivalent mass of metal is transferred from one to the other. The impurities cannot dissolve in this way. They fall from the anode and collect at the bottom of the vessel as a sludge called the anode mud.
Metals such as copper, silver, gold, nickel, aluminium and zinc are refined by this process. In a factory the single cell of Fig. 7.9-1 is simply scaled up, with many anode plates and cathode sheets hanging alternately in one large tank.
Diagram
Fig. 7.9-1 — Electro-Refining Of A Metal Sectional view of a rectangular tank containing the electrolyte, drawn with a stippled fill and labelled ‘solution of a salt of the metal’. A thick block on the left is labelled ‘impure metal (anode)’ and connected to the positive terminal of a battery; a thin plate on the right is labelled ‘pure metal (cathode)’ and connected to the negative terminal. Show a dark deposit on the floor of the tank beneath the anode, labelled ‘anode mud’. Mark the direction of the current in the external circuit.
7.9.2 Distillation And Liquation
Two simpler methods work where the metal has a convenient physical property.
Distillation is used for volatile metals — those with low boiling points, such as mercury and zinc. The impure metal is heated until the metal itself boils and distils over in the pure form, while the non-volatile impurities are left behind in the vessel. Zinc spelter, the crude zinc of Section 7.8, gives pure zinc on distillation.
Liquation is used for metals with low melting points, such as lead and tin. The impure metal is placed on the sloping hearth of a furnace and heated gently. The metal melts and flows down the slope into a collecting vessel, while the impurities, which have higher melting points, do not melt and stay behind on the hearth. Fig. 7.9-2 shows the furnace.
Diagram
Fig. 7.9-2 — Refining By Liquation Sectional view of a brick furnace with a sloping hearth. Show lumps of impure metal fed in at the upper left of the hearth, a fire burning in the chamber below, molten metal running down the slope and dripping into a collecting pot at the lower right, and unmelted material remaining on the upper part of the hearth. Label: impure metal, infusible impurity, pure metal.
7.10 Aluminium And Its Ore
Everything in this chapter so far — ores, dressing, roasting and calcination, reduction, refining — now comes together on a single metal. Aluminium is worth following in full, because it is the metal whose extraction shows the whole of metallurgy at work.
Table 7.10 The element aluminium
| Property | Value |
|---|---|
| Symbol | Al |
| Atomic number | 13 |
| Relative atomic mass | 27 |
| Electronic configuration | 2, 8, 3 |
| Valency | 3 |
| Colour | Silvery white metal |
| Specific gravity | 2.7 |
| Position in the periodic table | Period 3, Group IIIA (13) |
The valency in that table is not a fact to be memorised on its own; it follows from the line above it. An aluminium atom has three electrons in its outermost shell, and by losing all three it reaches the stable arrangement 2, 8 of the nearest noble gas. Losing three electrons leaves it with three units of positive charge, so aluminium forms the ion \(\mathrm{Al^{3+}}\) and its valency is 3, as Fig. 7.10-1 shows.
\[\mathrm{Al} \rightarrow \mathrm{Al^{3+}} + 3e^{-}\]
Diagram
Fig. 7.10-1 — The Aluminium Atom And The Aluminium Ion Two shell diagrams side by side. On the left, a nucleus marked 13p, 14n with three concentric shells holding 2, 8 and 3 electrons, labelled ‘aluminium atom, Al’. On the right, the same nucleus with only two shells holding 2 and 8 electrons, labelled ‘aluminium ion, Al³⁺’. Between them draw an arrow labelled ‘loses 3 electrons’, with the three departing electrons shown leaving the outer shell.
Abundant, But Never Free
Aluminium is the most abundant metal in the earth’s crust — you saw its 8% share in Table 7.1, ahead of every other metal including iron. Yet not one gram of it exists there as the metal.
The reason is in the activity series. Aluminium is a reactive metal, high in the list, and it holds oxygen so firmly that any free aluminium would have combined long ago. So it does not occur free in nature, and every atom of it must be prised out of a compound.
The Ore
Aluminium is extracted from its main ore bauxite, hydrated aluminium oxide, \(\mathrm{Al_2O_3\!\cdot\!2H_2O}\).
Bauxite as it is mined contains about 60% \(\mathrm{Al_2O_3}\). The remaining 40% is sand, ferric oxide and titanium oxide, and all of it has to go. Notice that this ore cannot be dressed by any of the physical methods of Section 7.6 — the impurities are too finely mixed through it for washing or a magnet to separate. Bauxite is concentrated by a chemical method, which is the leaching process of Section 7.6.4, and Section 7.11 follows it step by step.
Once the pure oxide has been obtained, there is only one way to get the metal out of it. Aluminium oxide, because of the great affinity of aluminium for oxygen, is a very stable compound, and it is not reduced by the common reducing agents — carbon, carbon monoxide or hydrogen. Electrolytic reduction was therefore chosen as the method for reducing alumina, and that is the subject of Section 7.12.
7.11 Purification Of Bauxite — Bayer’s Process
The problem is this. Bauxite is aluminium oxide mixed with sand, ferric oxide and titanium oxide, and the mixture is too fine to sort out by density or by a magnet. It has to be sorted out chemically, by finding a reagent that dissolves the aluminium oxide and leaves everything else untouched.
Such a reagent exists, and the reason it works lies in a property of aluminium oxide.
Aluminium oxide is amphoteric. Ferric oxide is basic, and sand is acidic, so neither of them dissolves in an alkali. Aluminium oxide does. Treat bauxite with a strong alkali and only the aluminium goes into solution — which is exactly the separation we need. The purification is carried out in three steps, set out in FC 7.2.
Flowchart
FC 7.2 — Bayer’s Process For The Purification Of Bauxite Finely ground bauxite → heated under pressure with concentrated caustic soda at 140–150 °C for 2 to 8 hours → sodium meta aluminate in solution, with insoluble red mud filtered off → solution diluted with water and cooled to 50 °C → aluminium hydroxide precipitated, soluble impurities left in solution → precipitate filtered, washed and dried → ignited at 1000 °C → pure alumina; the sodium hydroxide filtrate returns to the start of the process
Step 1 — Dissolving The Aluminium Out
The bauxite is first ground finely and heated to drive off volatile impurities. It is then heated under pressure with concentrated caustic soda solution for 2 to 8 hours at 140 °C to 150 °C. This is the leaching step of Section 7.6.4, applied to an ore.
bauxite + sodium hydroxide → sodium meta aluminate + water
\[\mathrm{Al_2O_3\!\cdot\!2H_2O} + 2\,\mathrm{NaOH} \xrightarrow{\ 140^\circ\text{--}150^\circ\text{C}\ } 2\,\mathrm{NaAlO_2} + 3\,\mathrm{H_2O}\]
The bauxite dissolves and forms sodium meta aluminate, because of the amphoteric nature of aluminium. The insoluble impurities do not dissolve, and they are left behind as a reddish sludge called red mud — ferric oxide, sand and the rest — which is removed by filtration. The red colour of the mud, and of the ore, comes from the iron.
Step 2 — Bringing The Aluminium Back Out
The clear solution now contains the aluminium, but so does it contain the other soluble impurities. The next step separates them, by turning the aluminium back into a solid while leaving the impurities in solution.
The sodium meta aluminate solution is diluted with water and cooled to 50 °C. It is then hydrolysed, and aluminium hydroxide comes down as a precipitate.
sodium meta aluminate + water → sodium hydroxide + aluminium hydroxide
\[\mathrm{NaAlO_2} + 2\,\mathrm{H_2O} \xrightarrow{\ 50^\circ\text{C}\ } \mathrm{NaOH} + \mathrm{Al(OH)_3}\downarrow\]
The other soluble impurities remain dissolved in the sodium hydroxide, so filtering at this point separates them from the aluminium for good.
Step 3 — Making The Oxide
The precipitate of aluminium hydroxide is filtered, washed and dried, and then ignited at 1000 °C. The hydroxide decomposes, and what is left is pure alumina.
aluminium hydroxide → aluminium oxide + water
\[2\,\mathrm{Al(OH)_3} \xrightarrow{\ \text{heat},\ 1000^\circ\text{C}\ } \mathrm{Al_2O_3} + 3\,\mathrm{H_2O}\]
Nothing in this process is wasted. The filtrate from step 2 contains the sodium hydroxide released by the hydrolysis, and it is returned to step 1 to dissolve the next batch of bauxite.
7.12 Electrolytic Reduction Of Alumina — Hall–Héroult Process
Pure alumina is now in hand, and only one step remains: tearing the aluminium away from the oxygen. Section 7.10 established that no chemical reducing agent can do it, so the current must do it instead. But electrolysis brings a difficulty of its own.
An electrolyte has to be either molten or in solution, and a reactive metal cannot be got from an aqueous solution. So the alumina must be molten — and alumina melts at about 2050 °C. That temperature is not merely awkward to reach. At 2050 °C aluminium itself vaporises, so the metal would boil away as fast as it formed and could not be collected.
7.12.1 The Electrolytic Cell And The Electrolyte
The cell. It is a rectangular iron tank with a sloping bottom, lined on the inside with gas carbon. The sloping floor lets the molten aluminium, which collects there, be run off cleanly through a tap hole. The carbon lining is not merely a lining: it is the cathode.
The electrolyte. It is a mixture of molten alumina 20%, cryolite 60% and fluorspar 20% — that is, alumina, cryolite and fluorspar in the ratio 1 : 3 : 1. Powdered coke is sprinkled over the surface of the electrolyte.
The temperature. The molten electrolyte is kept at 950 °C by electrical heating.
The voltage. 5 to 6 volts. A low voltage is used deliberately, to avoid decomposing the molten cryolite along with the alumina.
The whole arrangement is shown in Fig. 7.12-1.
Diagram
Fig. 7.12-1 — The Cell For The Extraction Of Aluminium Sectional view of a rectangular iron tank with a sloping floor, the inside faces lined with a thick layer labelled ‘carbon lining (cathode)’ and connected to the negative terminal. From above, four thick vertical rods labelled ‘carbon (graphite) anodes’ hang down into the tank from a horizontal bus bar connected to the positive terminal. Fill the tank with a stippled layer labelled ‘electrolyte — molten alumina, cryolite and fluorspar at 950 °C’, with a thin dotted layer on top labelled ‘powdered coke’. Show a pool of molten aluminium collected on the sloping floor and a tap hole at the lower end labelled ‘tap hole for aluminium’.
7.12.2 The Electrode Reactions
In the molten bath, all three components ionise.
\[\mathrm{Na_3AlF_6} \rightleftharpoons 3\,\mathrm{Na^{+}} + \mathrm{Al^{3+}} + 6\,\mathrm{F^{-}}\]
\[\mathrm{CaF_2} \rightleftharpoons \mathrm{Ca^{2+}} + 2\,\mathrm{F^{-}}\]
\[\mathrm{Al_2O_3} \rightleftharpoons 2\,\mathrm{Al^{3+}} + 3\,\mathrm{O^{2-}}\]
So the positive ions present are \(\mathrm{Na^{+}}\), \(\mathrm{Ca^{2+}}\) and \(\mathrm{Al^{3+}}\), and the negative ions are \(\mathrm{F^{-}}\) and \(\mathrm{O^{2-}}\). Three kinds of positive ion travel to the cathode, but only one of them is discharged. Aluminium lies lower in the electrochemical series than sodium and calcium, and the lower ion is discharged in preference, so it is aluminium that is deposited. At the anode, oxygen is discharged in preference to fluorine.
At the cathode (the inner carbon lining of the cell):
\[4\,\mathrm{Al^{3+}} + 12e^{-} \rightarrow 4\,\mathrm{Al}\]
The molten aluminium is denser than the electrolyte, so it sinks to the sloping floor and gathers there.
At the anode (thick rods of graphite suspended in the fused electrolyte):
\[6\,\mathrm{O^{2-}} - 12e^{-} \rightarrow 6[\mathrm{O}]\]
\[6[\mathrm{O}] \rightarrow 3\,\mathrm{O_2}\]
The oxygen set free does not simply bubble away. It meets the white-hot graphite of the anode and burns it:
\[2\,\mathrm{C} + \mathrm{O_2} \rightarrow 2\,\mathrm{CO}\]
\[2\,\mathrm{CO} + \mathrm{O_2} \rightarrow 2\,\mathrm{CO_2}\]
The anode is therefore eaten away as the cell runs, and it has to be replaced from time to time — a working cost that no amount of clever design removes.
Knowing When To Feed The Cell
As the alumina in the bath is used up, the cell tells the operator. The voltage suddenly jumps far above its working value, to something of the order of 40 to 60 volts, and a bulb connected across the cell glows brightly. Fresh alumina is then added and the voltage falls back.
The aluminium obtained from the cell is 99.8% pure. For most purposes that is enough; where it is not, further purification can be done by electrolysis, and Section 7.13 describes it.
7.12.3 The Purpose Of Cryolite, Fluorspar And Coke
Three substances in the cell are not there to supply aluminium, and it is worth being clear about what each one is for.
- Cryolite lowers the fusion temperature of the charge from about 2050 °C to about 950 °C, and it increases the conductivity of the melt.
- Fluorspar and cryolite together act as the solvent for the electrolytic mixture, and both increase its conductivity — necessary because pure alumina is almost a non-conductor of electricity.
- Powdered coke, sprinkled over the surface of the electrolytic mixture, does two things: it reduces heat loss by radiation from the top of the bath, and it prevents the burning of the anode by keeping air away from it.
7.13 Refining Of Aluminium — Hoope’s Process
The aluminium tapped from the Hall–Héroult cell is 99.8% pure. For a cooking vessel or a window frame that is entirely good enough. For electrical wiring and circuits it is not, because even small amounts of impurity raise the resistance of the metal, and a conductor is bought for its low resistance. Aluminium for such uses is refined further, by electrolysis, in Hoope’s process.
Three Liquid Layers
The cell is a single vessel in which three molten substances float one above the other. They stay apart because their specific gravities differ, the densest at the bottom and the least dense at the top, in the same way that oil floats on water. Fig. 7.13-1 shows the arrangement.
The bottom layer is molten impure aluminium. The floor and sides of this layer are carbon-lined, and this layer serves as the anode.
The middle layer is a mixture of the fluorides of sodium, barium and aluminium. It is the electrolyte.
The top layer is pure molten aluminium, into which carbon electrodes dip from above. This layer serves as the cathode.
Diagram
Fig. 7.13-1 — Refining Of Aluminium By Hoope’s Process Sectional view of a rectangular tank with three horizontal molten layers drawn in different fills. Label the lowest layer ‘molten impure aluminium (anode)’ with a carbon lining along the floor and lower walls, connected to the positive terminal; the middle layer ‘molten fluorides of sodium, barium and aluminium (electrolyte)’; and the top layer ‘pure molten aluminium (cathode)’, with several graphite rods dipping into it from above and connected to the negative terminal. Add a tap at the top right for drawing off the pure metal.
What The Current Does
At the anode, which is the impure metal at the bottom, aluminium atoms lose electrons and enter the middle layer as ions:
\[\mathrm{Al}\ \text{(impure)} - 3e^{-} \rightarrow \mathrm{Al^{3+}}\]
At the cathode, which is the pure metal at the top, aluminium ions from the middle layer gain electrons and join the pure metal:
\[\mathrm{Al^{3+}} + 3e^{-} \rightarrow \mathrm{Al}\ \text{(pure)}\]
So aluminium passes from the middle layer up into the top layer, and an equivalent amount shifts from the lower layer into the middle one to replace it. The impurities cannot make that journey; they stay in the bottom layer. Pure aluminium is tapped from the top as the top layer grows, and the metal so obtained is 99.9% pure.
7.14 Alloys
A pure metal is rarely the most useful form of that metal. Pure iron is soft, pure gold bends under a thumbnail, and pure aluminium is too weak to build with. What engineers actually use, almost always, is a mixture called an alloy.
The word homogeneous matters: the components are mixed so evenly that no part of the alloy differs from any other part. And the properties of the mixture are often greatly different from those of the metals that went into it.
Steel shows how far a small addition can go. Carbon present up to 1.5% profoundly affects the properties of iron. Adding 15 to 18% of chromium, often with a few per cent of nickel, markedly increases the corrosion and oxidation resistance of steel and gives stainless steel. A low percentage of molybdenum improves its toughness and wear resistance. Hundreds of special alloys exist, each one tailor-made to meet a particular set of operating conditions.
Two named groups are worth knowing. An alloy steel is a steel containing up to 10% of elements such as chromium, molybdenum or nickel, usually with a low percentage of carbon; these added elements improve hardness, wear resistance and toughness. A fusible alloy is one melting in the range of about 51 °C to 260 °C, and such alloys usually contain bismuth, lead and tin. There are also amorphous alloys, used in transformer coils, which are made by quenching the molten metals so quickly that they have no chance to crystallise.
7.14.1 Why Alloys Are Made
Alloys are made to change the property of their major constituent in order to achieve some specific objective. Six such objectives cover most of the alloys you will meet.
- To modify appearance and colour. Aluminium bronze, an alloy of aluminium and copper, is bright yellow and resembles gold.
- To modify chemical activity. Sodium amalgam, an alloy of sodium with mercury, is far less reactive than sodium itself and is much easier to handle.
- To modify casting ability. Type metal, an alloy of lead, tin and antimony, expands slightly on solidifying instead of shrinking, so it fills every corner of a mould and is easily cast. It was the metal of printing type.
- To lower the melting point. Solder, an alloy of lead and tin, melts at 180 °C, which is lower than the melting point of either lead or tin. That is precisely what makes it useful for joining electrical connections: the joint can be made without melting anything else.
- To increase hardness and tensile strength. Brass, an alloy of copper and zinc, is harder than copper and stronger than either of its components. Duralumin, an alloy of aluminium with copper, has a strength up to six times greater than that of pure aluminium. Carboloy, an alloy of tungsten, carbon and cobalt, is as hard as diamond.
- To increase resistance to electricity. Nichrome, an alloy of nickel, iron and chromium, offers far more resistance than copper does, and a wire that resists a current heats up strongly — which is the whole point of a heating element.
Other properties can be adjusted too. Bell metal, 78% copper and 22% tin, is more sonorous than either copper or tin, and rings for longer. Alnico, an alloy containing aluminium, nickel and cobalt, makes such a strong magnet that a piece of it can lift sixty times its own mass.
One caution runs through the whole subject. The purpose of an alloy is to improve the specific usefulness of the primary component, not to adulterate it or degrade it.
7.14.2 Methods Of Making Alloys
There are two ways of making an alloy.
- By fusing the metals together. The metals are melted, mixed in the liquid state and allowed to solidify. Brass is made by melting zinc and copper together.
- By compressing finely divided metals. The metals are powdered and pressed together. Wood metal, an alloy of lead, tin, bismuth and cadmium powder, is made this way. It is a fusible alloy, and it is used in the automatic sprinklers that release a spray of water to stop a fire from spreading — the heat of the fire melts a small plug of the alloy and opens the valve.
7.14.3 Alloys As Solid Solutions
An alloy can be thought of as a solid solution. The metal present in the higher concentration is the solvent, and the other metals dissolved in it are the solutes. Brass, on this view, is a solid solution of zinc, the solute, in copper, the solvent.
The picture is a useful one, because it explains why an alloy is homogeneous and why its composition can be varied continuously — brass may be 60% or 70% copper, just as a solution may be more or less concentrated.
7.14.4 Common Alloys, Their Compositions And Uses
Table 7.11 collects the alloys that matter most, with what each is made of, what it is like, and what it is used for.
Table 7.11 Common alloys, their compositions, properties and uses
| Principal metal | Alloy | Composition | Properties | Used for making |
|---|---|---|---|---|
| Aluminium | Duralumin | 95% Al, 4% Cu, 0.5% Mg, 0.5% Mn | Light, hard and resistant to corrosion; highly ductile. Al imparts lightness, Mg imparts strength | Bodies of aircraft, buses and tube trains; light tools; pressure cookers |
| Aluminium | Magnalium | 90–95% Al, 10–5% Mg | Resists corrosion; light; strong | Aircraft; scientific instruments; metal mirrors; light tools; beams of balance; household appliances |
| Iron | Stainless steel | 73% Fe, 18% Cr, 8% Ni, 1% C | Resists corrosion; lustrous and hard; resistant to acids and alkalis. Ni and Cr impart lustre, C imparts hardness | Utensils; cutlery; ornamental pieces; surgical instruments |
| Iron | Manganese steel | 85% Fe, 14% Mn, 1% C | Durable, tough and hard | Safes; rock drills; armour plates |
| Iron | Tungsten steel | 84% Fe, 5% W, 1% C | Very hard | Cutting tools for high-speed lathes |
| Zinc and copper | Brass | 60–70% Cu, 40–30% Zn | Malleable and ductile; easily cast; resists corrosion; yellow to silvery in colour | Decorative hardware and utensils; screws and handles; cartridges and containers; parts of watches and taps; musical instruments; electrical goods; marine engines |
| Zinc and copper | Bronze | 80% Cu, 18% Sn, 2% Zn | Hard and easily cast; takes a polish; resists corrosion | Medals; statues; utensils; bearings; coins |
| Zinc and copper | German silver | 50% Cu, 30% Zn, 20% Ni | White and light, like silver; malleable and ductile; high electrical resistance | Decorative articles; electric heaters and rheostats; resistors |
| Lead | Solder or fuse metal | 50% Pb, 50% Sn | Low melting point, 180 °C; high tensile strength; tin lowers the melting point | Soldering purposes; fuses |
| Lead | Type metal | 75% Pb, 10% Sn, 15% Sb | Tin and antimony improve the casting property and the capacity to expand on solidification | Printing blocks |
7.15 Amalgams
One metal makes alloys in a way no other metal can. Mercury is liquid at room temperature, so it does not have to be melted before another metal will dissolve in it — the other metal simply dissolves, as sugar dissolves in water. An alloy made in this way has its own name — an amalgam.
Mercury does not amalgamate with every metal. It does not form an amalgam with iron, and that single fact has a practical consequence: mercury is commonly stored and transported in iron bottles, since the container cannot be attacked by what is inside it.
Two amalgams are worth knowing by name.
- Dental amalgam is a mixture of mercury with a silver–tin alloy. It is soft enough to be packed into a prepared cavity in a tooth and hardens in place.
- Sodium amalgam, which you met in Section 7.14.1, is much less reactive than sodium itself, and when the proportion of sodium in it is small it is a liquid alloy at room temperature.
7.16 Summary
- Of the 118 known elements, only 17 are non-metals. Oxygen (46.6%) and silicon (27.7%) are the commonest elements in the earth’s crust, and aluminium (8%) is the commonest metal, ahead of iron (5%).
- Metals are malleable and ductile, dense, high-melting and good conductors of heat and electricity. Chemically, metals are elements which form positive ions by the loss of electrons.
- Non-metals, except hydrogen, are elements which form negative ions by the gain of electrons. Hydrogen is a non-metal that forms a positive ion, which is why it is placed on the left of the periodic table and included in the activity series.
- Elements showing the properties of both families are metalloids: boron, silicon, germanium, arsenic, antimony, tellurium and polonium. They lie along the zig-zag line that separates the metals on the left and centre of the periodic table from the non-metals on the right.
- A metal that is unreactive occurs in the native or free state; gold and platinum are examples. Reactive metals occur in the combined state, as oxides, carbonates, halides, sulphides and sulphates.
- A mineral is a naturally occurring compound of a metal. An ore is a mineral from which the metal can be extracted commercially, at low cost and with least effort. All ores are minerals, but all minerals are not ores. The earthy impurity present with a mineral is the gangue or matrix.
- The chief ores are: bauxite \(\mathrm{Al_2O_3\!\cdot\!2H_2O}\) for aluminium, haematite \(\mathrm{Fe_2O_3}\) for iron, and zinc blende \(\mathrm{ZnS}\) for zinc.
- Metallurgy is the extraction of metals in the pure form from their ores, together with their purification and the manufacture of alloys. It has five stages: pulverisation, concentration, conversion to the oxide, reduction, and refining.
- Concentration removes the gangue. Hydraulic washing uses the difference in density; magnetic separation uses magnetic properties; froth floatation uses the wetting of the ore by oil and of the gangue by water, and is used for sulphide ores; leaching dissolves the ore in a chosen reagent.
- A flux is added to combine with the gangue that survives into the furnace, forming a fusible slag which floats on the molten metal and is run off, for example \(\mathrm{SiO_2} + \mathrm{CaCO_3} \rightarrow \mathrm{CaSiO_3} + \mathrm{CO_2}\).
- Roasting is the heating of a concentrated ore to a high temperature below its melting point in excess of air. It is used for sulphide ores: \(2\,\mathrm{ZnS} + 3\,\mathrm{O_2} \rightarrow 2\,\mathrm{ZnO} + 2\,\mathrm{SO_2}\).
- Calcination is the heating of a carbonate ore or hydrated oxide in the absence of air, or in a limited supply, to a temperature insufficient to melt it: \(\mathrm{ZnCO_3} \rightarrow \mathrm{ZnO} + \mathrm{CO_2}\). Both processes make the ore porous and leave it as the oxide, which is easier to reduce.
- The activity series arranges the metals in decreasing order of reactivity: K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au. A metal higher in the series displaces one below it from its salt solution; metals above hydrogen displace hydrogen from dilute acids.
- The oxides of K, Na, Ca, Mg and Al cannot be reduced by carbon, carbon monoxide or hydrogen. These metals are obtained by electrolytic reduction of the fused salt, for example \(\mathrm{Na^{+}} + e^{-} \rightarrow \mathrm{Na}\) at the cathode of fused sodium chloride. They cannot be obtained from aqueous solutions, since they react with water.
- Metals in the middle of the series are obtained by reducing their oxides with carbon, carbon monoxide or hydrogen: \(\mathrm{ZnO} + \mathrm{C} \rightarrow \mathrm{Zn} + \mathrm{CO}\) at 1400 °C gives zinc spelter; \(\mathrm{PbO}\), \(\mathrm{CuO}\), \(\mathrm{FeO}\) and \(\mathrm{Fe_2O_3}\) are reduced similarly. Zinc oxide is reduced only by carbon, which is the stronger reducing agent.
- Metals low in the series need no reducing agent; their oxides decompose on heating, as in \(2\,\mathrm{HgO} \rightarrow 2\,\mathrm{Hg} + \mathrm{O_2}\) above 300 °C.
- A reactive metal can act as a reducing agent for a less reactive one. In aluminothermy, \(\mathrm{Fe_2O_3} + 2\,\mathrm{Al} \rightarrow 2\,\mathrm{Fe} + \mathrm{Al_2O_3}\), and the heat released leaves the iron molten.
- Refining purifies the crude metal. In electro-refining the impure metal is the anode, a thin sheet of pure metal the cathode, and a salt solution of the metal the electrolyte; pure metal deposits at the cathode and the impurities collect as anode mud. Volatile metals are refined by distillation and low-melting ones by liquation.
- Aluminium (Z = 13, configuration 2, 8, 3, valency 3) is the most abundant metal in the crust but never occurs free, because it is reactive and its oxide is very stable.
- Bauxite is purified by Bayer’s process, which works because aluminium oxide is amphoteric while the impurities are not: \[\mathrm{Al_2O_3\!\cdot\!2H_2O} + 2\,\mathrm{NaOH} \xrightarrow{\ 140^\circ\text{--}150^\circ\text{C}\ } 2\,\mathrm{NaAlO_2} + 3\,\mathrm{H_2O}\] \[\mathrm{NaAlO_2} + 2\,\mathrm{H_2O} \xrightarrow{\ 50^\circ\text{C}\ } \mathrm{NaOH} + \mathrm{Al(OH)_3}\downarrow\] \[2\,\mathrm{Al(OH)_3} \xrightarrow{\ 1000^\circ\text{C}\ } \mathrm{Al_2O_3} + 3\,\mathrm{H_2O}\] The insoluble residue removed by filtration is red mud.
- In the Hall–Héroult process the electrolyte is molten alumina 20%, cryolite 60% and fluorspar 20%, at 950 °C and 5 to 6 volts. Cryolite dissolves the alumina, lowers the working temperature from about 2050 °C, and makes the melt conduct. Powdered coke on the surface reduces heat loss and protects the anode.
- The electrode reactions are \(4\,\mathrm{Al^{3+}} + 12e^{-} \rightarrow 4\,\mathrm{Al}\) at the carbon-lined cathode and \(6\,\mathrm{O^{2-}} - 12e^{-} \rightarrow 3\,\mathrm{O_2}\) at the graphite anode. The oxygen burns the anode away, so the anode must be replaced from time to time. The aluminium obtained is 99.8% pure.
- Hoope’s process refines it further, in a cell of three molten layers held apart by their densities: impure aluminium below as anode, fluorides in the middle, pure aluminium on top as cathode. The metal obtained is 99.9% pure.
- An alloy is a homogeneous mixture of two or more metals, or of metals with certain non-metals. Alloys are made to change colour, chemical activity, casting ability, melting point, hardness and tensile strength, or electrical resistance. An alloy may be regarded as a solid solution.
- Compositions worth remembering: duralumin 95% Al, 4% Cu, 0.5% Mg, 0.5% Mn; magnalium 90–95% Al with 10–5% Mg; stainless steel 73% Fe, 18% Cr, 8% Ni, 1% C; brass 60–70% Cu with 40–30% Zn; bronze 80% Cu, 18% Sn, 2% Zn; solder or fuse metal 50% Pb, 50% Sn.
- An amalgam is an alloy of mercury with another metal. Mercury forms no amalgam with iron, so it is stored in iron bottles. Dental amalgam is mercury with a silver–tin alloy.
- The extraction of aluminium, from ore to refined metal, is condensed in FC 7.3.
Flowchart
FC 7.3 — The Extraction Of Aluminium From Bauxite Bauxite (Al₂O₃·2H₂O) → Bayer’s process: leaching with concentrated NaOH at 140–150 °C, red mud filtered off → sodium meta aluminate solution → diluted and cooled to 50 °C → Al(OH)₃ precipitated and filtered → ignited at 1000 °C → pure alumina (Al₂O₃) → dissolved in molten cryolite with fluorspar, 1 : 3 : 1, at 950 °C → electrolysis at 5–6 V in an iron tank with a carbon lining → aluminium at the cathode, tapped from the sloping floor, 99.8% pure; oxygen at the graphite anode, which burns away to CO and CO₂ and is replaced → Hoope’s process, three molten layers → aluminium 99.9% pure
7.17 Key Terms
Activity series — the arrangement of the metals in decreasing order of their reactivity.
Alloy — a homogeneous mixture of two or more metals, or of one or more metals with certain non-metallic elements.
Alloy steel — a steel containing up to 10% of elements such as chromium, molybdenum or nickel, with a low percentage of carbon.
Alumina — pure anhydrous aluminium oxide, \(\mathrm{Al_2O_3}\), the compound electrolysed to obtain aluminium.
Aluminothermy — the reduction of a metal oxide by igniting it with aluminium powder, which releases enough heat to leave the metal molten.
Amalgam — an alloy of mercury with another metal or alloy.
Amphoteric — describing an oxide or hydroxide that reacts with acids as a base does and with alkalis as an acid does.
Anode mud — the sludge of impurities that settles beneath the anode during electro-refining.
Bauxite — hydrated aluminium oxide, \(\mathrm{Al_2O_3\!\cdot\!2H_2O}\), the chief ore of aluminium.
Calcination — heating a carbonate ore or hydrated oxide in the absence, or a limited supply, of air to convert it into the oxide.
Carbon reduction process — the extraction of a metal by heating its oxide with carbon.
Combined state — the condition of a metal found in nature joined to other elements in a compound.
Concentration — the removal of the gangue from a powdered ore; also called ore dressing.
Cryolite — sodium aluminium fluoride, \(\mathrm{Na_3AlF_6}\), which dissolves alumina and lowers the working temperature of the extraction cell.
Distillation — refining a volatile metal by boiling it off and condensing it, leaving non-volatile impurities behind.
Dressing — another name for the concentration of an ore.
Ductile — able to be drawn out into a wire.
Electrolytic reduction — obtaining a metal by passing a direct current through its fused compound.
Electropositivity — the readiness of an atom to lose electrons and form a positive ion.
Flux — a substance added to the furnace to combine with the remaining gangue and form a fusible slag.
Froth floatation — concentrating a sulphide ore by wetting it with oil so that it is carried up on air bubbles as a froth.
Fusible alloy — an alloy melting between about 51 °C and 260 °C, usually containing bismuth, lead and tin.
Gangue — the earthy impurity, such as sand, clay or rock, present with a mineral in the ground; also called the matrix.
Haematite — ferric oxide, \(\mathrm{Fe_2O_3}\), the chief ore of iron.
Hydraulic washing — concentrating an ore by washing it down a sloping table, using the difference in density between ore and gangue.
Leaching — concentrating an ore by dissolving it in a reagent in which the impurities are insoluble.
Liquation — refining a low-melting metal by melting it on a sloping hearth so that it runs away from its infusible impurities.
Magnetic separation — concentrating an ore by using a magnetic wheel to draw off whichever of the ore and the gangue is magnetic.
Malleable — able to be beaten into sheets.
Metallurgy — the extraction of metals in pure form from their ores, together with their purification and the manufacture of alloys.
Metalloids — elements such as silicon and germanium that show the properties of both metals and non-metals.
Metals — elements which form positive ions by the loss of electrons.
Mineral — a naturally occurring compound of a metal.
Native state — the condition of a metal found in nature as the free, uncombined element.
Non-metals — elements which, except for hydrogen, form negative ions by the gain of electrons.
Ore — a mineral from which a metal can be extracted commercially, at low cost and with least effort.
Pulverisation — crushing and grinding the ore into a fine powder.
Red mud — the insoluble residue, chiefly ferric oxide and sand, filtered off during the purification of bauxite.
Refining — the process by which crude metal is purified.
Roasting — heating a concentrated ore, usually a sulphide, in excess of air at a temperature below its melting point.
Slag — the fusible compound formed when a flux reacts with the gangue, which floats on the molten metal and is run off.
Smelting — reducing a roasted or calcined ore in a furnace, with a flux, to obtain the crude metal.
Zinc blende — zinc sulphide, \(\mathrm{ZnS}\), the chief ore of zinc.
7.18 Points To Ponder
Whether a mineral counts as an ore depends on cost, and cost changes with time, with technology and with where the deposit happens to lie. Does that make ‘ore’ a chemical term at all? What would be lost, and what gained, if chemists defined an ore purely by the percentage of metal in it?
Aluminium stands high in the activity series, well above iron, and yet an aluminium window frame survives decades of rain while an iron gate rusts through. What could a reactive metal do that protects it, and why does the same thing not save iron?
Every metal above zinc in the activity series is obtained by electrolysis, which is expensive because electricity is expensive. Suppose electricity cost almost nothing. Which parts of this chapter would be rewritten, and would carbon reduction survive at all?
Zinc displaces silver from a solution of its salt because zinc is higher in the activity series. Aluminium is higher still. Why, then, is aluminium not used industrially to displace other metals from their solutions, in place of electrolysis? Think about what aluminium would meet first in an aqueous solution.
Brass is nothing but copper and zinc mixed, yet it is harder than either, and steel is iron with a trace of carbon, yet it is far stronger than iron. If nothing new has been created chemically, where do the new properties come from? What does this suggest about how the properties of a pure metal arise in the first place?
Roasting a sulphide ore releases sulphur dioxide into the air; extracting aluminium consumes enormous quantities of electrical energy. Both are costs, but of quite different kinds. Is there any honest way to compare them, and what information would you need before you could?
7.19 Objective Questions
Choose the correct option. [1 each]
- The ore of zinc that is concentrated by froth floatation is:
- calamine
- zincite
- zinc blende
- carnallite
- The substance sprinkled over the surface of the electrolyte in the Hall–Héroult cell is:
- fluorspar
- powdered coke
- cryolite
- red mud
- Which of the following oxides can be reduced to the metal by carbon monoxide?
- \(\mathrm{Al_2O_3}\)
- \(\mathrm{MgO}\)
- \(\mathrm{PbO}\)
- \(\mathrm{ZnO}\)
- In the electro-refining of copper, the impurities collect:
- on the cathode
- in the electrolyte as ions
- below the anode as a mud
- on the surface of the electrolyte
- An alloy that contains no copper is:
- bronze
- magnalium
- German silver
- duralumin
- The property of aluminium oxide that makes Bayer’s process possible is that it is:
- basic
- acidic
- neutral
- amphoteric
- The ore of zinc that is concentrated by froth floatation is:
Assertion–reason. Each question below has an assertion (A) and a reason (R). Choose from the codes: (1) both A and R are true and R is the correct explanation of A; (2) both A and R are true but R is not the correct explanation of A; (3) A is true but R is false; (4) A is false but R is true. [1 each]
Assertion (A): Sulphide ores are roasted and not calcined. Reason (R): Roasting is carried out in a plentiful supply of air.
Assertion (A): Zinc oxide is reduced by heating with carbon rather than with carbon monoxide. Reason (R): Carbon is a stronger reducing agent than carbon monoxide.
Assertion (A): Sodium is obtained by the electrolysis of fused sodium chloride and not of its aqueous solution. Reason (R): Sodium chloride does not dissolve in water.
Assertion (A): The graphite anodes of an aluminium cell have to be replaced from time to time. Reason (R): The oxygen liberated at the anode burns the carbon away.
Assertion (A): Mercury is stored in iron bottles. Reason (R): Mercury does not form an amalgam with iron.
Match each item in Column A with the one that belongs to it in Column B. [5]
| Column A | Column B |
|---|---|
| a. Pine oil | i. Lowers the fusion temperature of alumina |
| b. Cryolite | ii. Combines with the gangue to form slag |
| c. Limestone | iii. Wets the sulphide ore in the concentration tank |
| d. Caustic soda | iv. Reduces zinc oxide at 1400 °C |
| e. Coke | v. Dissolves aluminium oxide out of bauxite |
Fill in the blanks. [1 each]
- The earthy impurity present with a mineral is called the …………… .
- Heating an ore in the absence of air to a temperature below its melting point is called …………… .
- The residue filtered off during the purification of bauxite is called …………… mud.
- In Hoope’s process, pure aluminium collects in the …………… layer of the cell.
- An alloy of mercury with another metal is called an …………… .
State whether each of the following is true or false. [1 each]
- All minerals are ores.
- Metals below hydrogen in the activity series displace hydrogen from dilute acids.
- Roasting makes an ore porous.
- The cathode of the Hall–Héroult cell is made of graphite rods.
- An alloy has a fixed composition expressed by a formula.
Name the following. [1 each]
- The most abundant metal in the earth’s crust.
- A metal that occurs in the native state.
- The gas evolved when a sulphide ore is roasted.
- The process used to purify bauxite.
- An alloy used for making the heating element of an electric iron.
- The method of refining used for a metal with a low boiling point.
Give one word or phrase for each of the following. [1 each]
- The process of extracting metals in the pure form from their ores.
- The fusible substance formed when a flux reacts with the gangue.
- The crushing and grinding of an ore to a fine powder.
- The reduction of a metal oxide by ignited aluminium powder.
- An alloy that melts between about 51 °C and 260 °C.
In each list below, pick the odd one out and give a reason for your choice. [2 each]
- Bauxite, corundum, cryolite, calamine.
- Gravity separation, magnetic separation, calcination, froth floatation.
- Brass, bronze, German silver, magnalium.
- Potassium, calcium, copper, magnesium.
7.20 Check Your Concepts
Define an ore, and state how it differs from a mineral. [2]
Give a reason for each of the following. [2 each]
- Gold is found in the free state while zinc is not.
- A finely powdered ore is used at every stage after pulverisation.
- A low voltage of 5 to 6 volts is used in the aluminium cell.
State what is meant by the gangue of an ore, and name the substance added to remove it in the furnace. [2]
Name the four methods of concentrating an ore, and state in one line the difference between the ore and the gangue that each of them makes use of. [4]
Differentiate between roasting and calcination, stating clearly the basis of the difference in each case. Give one balanced equation for each process. [4]
Explain, with reference to the activity series, why the oxides of some metals can be reduced by carbon while the oxides of others cannot. Support your answer with one example of each kind. [3]
Why does the extraction of a metal such as calcium use the fused salt and never an aqueous solution of it? [2]
With reference to Bayer’s process: [5]
- State the property of aluminium oxide on which the process depends.
- Describe what happens to the ferric oxide present in the ore.
- Explain why the sodium meta aluminate solution is diluted and cooled.
- State what is done with the filtrate obtained after the aluminium hydroxide has been precipitated.
Explain the purpose of each of the following in the electrolytic extraction of aluminium: cryolite, fluorspar, powdered coke and the sloping bottom of the tank. [4]
Differentiate between electrolytic reduction and electro-refining, stating the basis of the difference. [3]
What would happen if, during the electrolysis of alumina, the supply of fresh alumina to the cell were not maintained? State what the operator would observe. [2]
What is an alloy? Explain, with two examples, why the properties of an alloy differ from those of the metals it is made from. [3]
Give the composition of stainless steel and of duralumin, and state one use of each that depends directly on the composition you have given. [4]
Aluminium is the commonest metal in the earth’s crust, yet it was one of the last common metals to come into everyday use. Account for this, referring to the stability of its oxide and to the method by which it must be extracted. [5]
7.21 Equations, Conversions & Identification
Write the balanced chemical equation for each of the following. [2 each]
- The roasting of zinc blende.
- The calcination of calamine.
- The reduction of lead(II) oxide by carbon monoxide.
- The reduction of iron(III) oxide by aluminium powder.
- The action of heat on aluminium hydroxide.
Copy and complete each of the following, and balance it. [2 each]
- \(\mathrm{CuO} + \mathrm{H_2} \rightarrow \ldots\ldots + \ldots\ldots\)
- \(\mathrm{Al_2O_3\!\cdot\!2H_2O} + \mathrm{NaOH} \rightarrow \ldots\ldots + \ldots\ldots\)
- \(\mathrm{SiO_2} + \mathrm{CaCO_3} \rightarrow \ldots\ldots + \ldots\ldots\)
- \(\mathrm{ZnO} + \mathrm{C} \rightarrow \ldots\ldots + \ldots\ldots\)
Write the equation for the reaction taking place at each electrode: [4]
- at the cathode and at the anode during the electrolysis of fused sodium chloride,
- at the cathode and at the anode during the electrolytic reduction of alumina.
A sulphide ore is strongly heated in a plentiful supply of air. [3]
- Name the process.
- Name the gas evolved and describe one test by which you would identify it.
- State what change would be seen in the ore itself.
Dry hydrogen is passed over black copper(II) oxide contained in a heated tube. State two things you would observe, and explain what each observation tells you. [3]
How would you obtain the following? Give the equation for each step. [3 each]
- Pure alumina from bauxite.
- Zinc from calamine.
- Mercury from cinnabar.
Identify the substance in each case. [2 each]
- A metal that is liquid at room temperature and forms no amalgam with iron.
- A reddish insoluble residue left when bauxite is treated with hot concentrated caustic soda.
- An ore that is attracted by a magnet and is an oxide of iron.
- A compound added to alumina to bring the working temperature of the cell down to about 950 °C.
Study Fig. 7.12-1, which shows the cell used for the extraction of aluminium. [5]
- Name the material of which the lining of the tank is made, and state which electrode it serves as.
- Name the material of the rods suspended from above, and state why they must be replaced periodically.
- State the purpose of the sloping floor of the tank.
- Give the three substances present in the electrolyte and the proportion in which they are mixed.
Study Fig. 7.9-1, which shows a cell used for electro-refining. [3]
- State what the thick block on the left and the thin plate on the right are made of.
- Name the material collecting at the bottom of the tank and state what it consists of.
- Explain why a little dilute sulphuric acid is added to the electrolyte.