The Harder a Metal Is to Extract the More Reactive It Is

Gold is found in the earth as gold. Iron is found combined with oxygen. Sodium and aluminium are found so tightly combined that heating with carbon will not free them at all, and only electricity can do it.

That pattern is not a coincidence — it is the reactivity series read backwards. A reactive metal gives its electrons away eagerly, so it forms very stable compounds and clings to them. An unreactive metal barely combines at all, so it lies around in its free state.

So the whole of metallurgy is decided by one question: where does this metal sit in the series?

- At the bottom — gold, silver, platinum and partly copper — found free, or freed by simple heating
- In the middle — zinc, iron, lead and copper — their oxides are reduced by heating with carbon
- At the top — potassium, sodium, calcium, magnesium and aluminium — obtained only by electrolysis of the molten compound

Three positions, three methods, and once you know a metal's place you can name its extraction route without being told.

This page covers the third part of the CBSE Class 10 Science chapter on metals and non-metals: minerals, ores and gangue, roasting and calcination, extraction by reactivity, the electrolytic refining of copper, and the prevention of rusting.

A mineral is any naturally occurring compound of a metal; an ore is a mineral rich enough to be worth extracting from; and gangue is the earthy rubbish mixed in with it.

- Mineral — the metal occurs combined with other elements in the earth's crust. Every metal has several minerals
- Ore — a mineral from which the metal can be extracted profitably, in a reasonable quantity with a reasonable effort
- Gangue — the sand, soil and rocky matter that comes up with the ore and has to be removed first

**So ore is an economic word, not a chemical one. The same mineral may be an ore where it occurs in thick deposits and not an ore where it is thinly scattered. Every ore is a mineral; not every mineral is an ore — that sentence answers the commonest one-mark question in this section.

Removing the gangue is called concentration or enrichment, and the method is chosen by a physical difference between ore and gangue:

- Washing — where the ore is denser, the lighter gangue is carried away by a stream of water
- Magnetic separation — where the ore is magnetic and the gangue is not
- Froth flotation — where the ore particles cling to oil and rise with the froth while the gangue sinks

Converting the concentrated ore to an oxide. Oxides are the easiest compounds to reduce, so sulphide and carbonate ores are first turned into oxides by one of two named processes.

Roasting — a sulphide ore is heated strongly in the presence of excess air.**

$$2\text{ZnS}+3{\text{O}}_2\stackrel{\text{heat}}{\to }2\text{ZnO}+2{\text{SO}}_2$$

Calcination — a carbonate ore is heated strongly in a limited supply of air.

$${\text{ZnCO}}_3\stackrel{\text{heat}}{\to }\text{ZnO}+{\text{CO}}_2$$
$${\text{Al}}_2{\text{O}}_3\cdot 2{\text{H}}_2\text{O}\stackrel{\text{heat}}{\to }{\text{Al}}_2{\text{O}}_3+2{\text{H}}_2\text{O}$$

The two words are swapped more often than any other pair in this chapter, so fix them by the reactant: sulphide goes with roasting and needs plenty of air because the sulphur must be burnt away as sulphur dioxide; carbonate goes with calcination and needs only enough heat to drive the carbon dioxide off. Calcination also removes water of crystallisation, as the second equation shows.

Both end at the same place. Whichever route is used, the product is the oxide, and the next step is always reduction of that oxide. That is why the two processes are taught together: they are two doors into the same corridor.

Find the metal in the reactivity series and read off the method.

Metals low in the series — simple heating is enough. Their compounds are unstable, so heat alone breaks them down.

Mercury is obtained from its sulphide ore in two steps:

$$2\text{HgS}+3{\text{O}}_2\stackrel{\text{heat}}{\to }2\text{HgO}+2{\text{SO}}_2$$
$$2\text{HgO}\stackrel{\text{heat}}{\to }2\text{Hg}+{\text{O}}_2$$

Copper is more interesting, because it reduces itself. Its sulphide ore is partly roasted, and the oxide then reacts with the remaining sulphide:

$$2{\text{Cu}}_2\text{S}+3{\text{O}}_2\stackrel{\text{heat}}{\to }2{\text{Cu}}_2\text{O}+2{\text{SO}}_2$$
$$2{\text{Cu}}_2\text{O}+{\text{Cu}}_2\text{S}\stackrel{\text{heat}}{\to }6\text{Cu}+{\text{SO}}_2$$

No carbon was needed at all — the sulphide acted as the reducing agent, which is why this is called self-reduction.

Metals in the middle — reduce the oxide with carbon. Carbon takes the oxygen away:

$$\text{ZnO}+\text{C}\stackrel{\text{heat}}{\to }\text{Zn}+\text{CO}$$
$${\text{Fe}}_2{\text{O}}_3+3\text{CO}\stackrel{\text{heat}}{\to }2\text{Fe}+3{\text{CO}}_2$$

The second is what happens in a blast furnace, and carbon monoxide does the reducing there.

A more reactive metal can also be used as the reducing agent where the heat is wanted as well:

$$3{\text{MnO}}_2+4\text{Al}\to 3\text{Mn}+2{\text{Al}}_2{\text{O}}_3$$

Metals at the top — electrolytic reduction of the molten compound. Carbon cannot pull the oxygen away from these, because they hold it more tightly than carbon does. Electricity can:

$$2{\text{Al}}_2{\text{O}}_3\stackrel{\text{electricity}}{\to }4\text{Al}+3{\text{O}}_2$$

At the electrodes, the metal ion collects electrons at the cathode and the non-metal gives them up at the anode. For molten sodium chloride:

$$\text{at the cathode: }{\text{Na}}^{+}+e^{-}\to \text{Na}\text{at the anode: }2{\text{Cl}}^{-}\to {\text{Cl}}_2+2e^{-}$$

Worked example — which method for each? Silver, iron and potassium.

- Silver is at the bottom, so heating or simple displacement is enough
- Iron is in the middle, so its oxide is reduced with carbon or carbon monoxide
- Potassium is at the very top, so only electrolysis of the molten compound will work

Notice the reason the top group needs electricity. Reduction means giving electrons to the metal ion, and a very reactive metal's ion holds onto its positive charge stubbornly. Electricity can supply electrons at any strength you like, while carbon can only supply as much reducing power as its own reactivity allows. That is why the method is set by the series and not by convenience.

Make the impure metal the anode and a thin strip of the pure metal the cathode, and pass a current through a solution of the metal's salt. The metal dissolves off one electrode and deposits on the other, leaving the impurities behind.

The three things a question always asks for:

- Anode — a thick block of the impure copper, connected to the positive terminal
- Cathode — a thin strip of pure copper, connected to the negative terminal
- Electrolyte — a solution of copper sulphate acidified with a little sulphuric acid

What happens when the current flows.

- At the anode, copper atoms lose electrons and go into solution as copper ions: the anode slowly dissolves
- At the cathode, copper ions from the solution gain electrons and deposit as pure copper: the cathode grows thicker
- Soluble impurities stay dissolved in the electrolyte
- Insoluble impurities settle below the anode as a sludge called anode mud

$$\text{at the anode: }\text{Cu}\to {\text{Cu}}^{2+}+2e^{-}\text{at the cathode: }{\text{Cu}}^{2+}+2e^{-}\to \text{Cu}$$

The two electrode reactions are exact opposites, which is why the concentration of the electrolyte hardly changes: every copper ion that leaves the anode is replaced by one arriving at the cathode. The electrolyte is a conveyor belt, not a reactant.

Why refining is needed at all. The copper that comes out of a furnace is about ninety-nine parts in a hundred pure, and that is not good enough for electrical wiring, because even small impurities raise the resistance noticeably. Electrolytic refining produces copper pure enough to carry current efficiently, which is its main industrial use.

The same method refines other metals — zinc, tin, nickel, silver, gold and aluminium are all refined this way, with the electrolyte changed to a salt of the metal being purified.

A distinction worth keeping. Extraction obtains the metal from its ore; refining purifies the metal you already have. Electrolysis appears in both, but for different reasons: in extraction it supplies the electrons that carbon cannot, and in refining it moves metal from one electrode to another. **A question about aluminium extraction and one about copper refining both say electrolysis, and they are not the same process.**

Rusting needs both air and moisture. Remove either and it stops.

The experiment that proves it. Set up three test tubes, each with an iron nail:

- The first has dry air only, with a drying agent such as anhydrous calcium chloride. The nail does not rust
- The second has boiled water with a layer of oil on top, so the water has no dissolved air. The nail does not rust
- The third has ordinary water and air together. The nail rusts

Only the tube with both rusts, and that is the complete answer to what are the conditions necessary for rusting. Rust is a hydrated iron oxide, ${\text{Fe}}_2{\text{O}}_3\cdot x{\text{H}}_2\text{O}$, which is why it is flaky and brown rather than a smooth layer.

Now the prevention methods, each judged on how it works.

Painting, oiling and greasing put a physical barrier between the iron and the air. Cheap and easy, and useless once scratched — the exposed metal rusts and the rust creeps under the coating. This is why a painted gate needs repainting and why a bicycle chain is oiled rather than painted.

Galvanisation coats the iron with a layer of zinc, used for roofing sheets, buckets, pipes and nails. It keeps working even when scratched, and the reason is the reactivity series: zinc is above iron, so the zinc is attacked in preference and the iron is left alone. That is called sacrificial protection.

Tin plating looks similar and behaves in the opposite way. Tin is below iron in the series, so a scratch in a tinned surface leaves the iron exposed and it rusts faster than it would have unprotected. Two metal coatings, opposite outcomes, decided entirely by position in the series.

Alloying changes the metal itself rather than covering it. Iron mixed with chromium and nickel gives stainless steel, which does not rust at all. Permanent and maintenance-free, but expensive, and the alloy has different properties from the iron it replaces.

Chrome plating and anodising give a bright, hard, non-reactive surface, used on taps, cycle handles and window frames.

So the evaluation asked for by the syllabus. For long outdoor exposure with rough handling, galvanisation wins, because it survives scratches. For cutlery and surgical instruments, alloying wins, because the protection can never be worn off. For a large painted structure, painting wins on cost, provided it is maintained. The best method depends on the exposure, not on the chemistry alone.

Why alloying is used for more than rust prevention. An alloy is a homogeneous mixture of a metal with one or more metals or non-metals, and alloying deliberately changes properties:

- Brass is copper with zinc, and bronze is copper with tin — both harder than copper
- Solder is lead with tin, with a low melting point, used for joining electrical wires
- Steel is iron with a small amount of carbon, far harder than pure iron
- Pure gold is too soft for jewellery, so it is alloyed with copper or silver
- An amalgam is any alloy containing mercury

The general rule: an alloy is harder, less ductile and a poorer conductor than the pure metal, and it usually has a lower melting point. Solder exists because of that last property, and pure tin or pure lead would be useless for the job.