(METALs& NON-METALS

 

 

 

CHEMISTRY (CLASS-X)

 

Chapter-3: (METALs& NON-METALS)

 

Element are classified into two basic categories, metals and non-metals.

Before giving the definition of metal and non-metal, we must know some important terms:

(i) Malleability: Ability of an element due to which it can be beaten with hammer into thin sheets is called malleability.

(ii) Ductility: Ability of an element due to which it can be drawn into wires is called ductility.

            (iii) Sonorous: Property of an element to produce sound when it is struck with a hard substance.

            (iv) Brittleness: Property of an element to break easily into pieces on hammering or stretching.

 

            Keeping these terms in mind, now we can define metal and non-metals.

Metalsare defined as those elements which possess lusture, are malleable and ductile and good conductor of heat and electricity.

 

AmityInstitute For Competitive Examinations : Ph. : 24336143/44,  41888409

 

Or

Metals are the elements which form positive ions by losing electrons i.e. they are electropositive elements. For e.g. sodium, magnesium, potassium, aluminium, copper, silver, gold etc.

 

Non-metalsare defined as those elements which do not possess lusture and are neither good conductors of heat and electricity nor malleable. They are not ductile but are brittle.

Or

Non-metals are the elements which form negative ions by gaining electrons i.e. they are electronegative elements. For e.g. carbon, hydrogen, oxygen, nitrogen, sulphur, bromine etc.

 

œ   The metals have been placed on the left side and in the centre of the periodic table, whereas non-metals have been placed on the right side in the periodic table. The elements which show the properties of both metals and non-metals are called metalloids. The metalloids are: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te) and polonium (Po).

 

1.1       PHYSICAL PROPERTIES OF METALS

            (i)   Metals are malleable i.e. they can be beaten into thin sheets with a hammer e.g. gold, silver, aluminium and copper.

(ii)  Metals are ductile i.e. they can be drawn into thin wires when hammered. Gold and silver are among the best ductile metals.

(iii) Metals are good conductors of heat i.e. metals allow heat to pass through them easily. When a metal is heated, its atoms gain energy and vibrate more vigorously. They transfer energy to other electrons and atoms of the metal and heat is conducted. Silver is the best conductor of heat. Copper and aluminium are also good conductors of heat. Lead and mercury are poor conductors of heat.

(iv) Metals are good conductors of electricity. This is because they contain free electrons. These free electrons can move easily through the metal and conduct electric current. Silver metal is the best conductor of electricity. Copper, gold and aluminium are also good conductors.

(v)  Metals are shiny. They can be polished. On exposure to air, metals lose their shine due to formation of oxides, sulphides, carbonates etc by reaction with gases present in air.

(vi) Metals are generally hard. However, sodium and potassium are soft metals which can easily be cut with a knife.

Sodium can be easily cut with a knife

(vii)      Metals are strong and have high tensile strength i.e. heavy weight can be suspended from metallic wires without breaking the wire.

(viii) Metals have high densities except sodium and potassium.

 

œ   Osmium is the heaviest metal and lithium is the lightest metal.

 

(ix) Metals are sonorous i.e. they produce sound when hit with some solid object.

(x)  Metals are solid at room temperature except mercury which is liquid. Gallium and caesium have very low melting point. These two melt when held in hand.

Liquid Metal : Mercury

(xi) Metals have high melting and boiling points except sodium and potassium.

 

œ   The most abundant metal in earth’s crust is aluminum whereas non-metal is oxygen.

 

1.2       PHYSICAL PROPERTIES OF NON-METALS

(i)   Non-metals are non-malleable and non-ductile. Non-metals are brittle. The property of being brittle (breaking easily) when hammered is called brittleness.

(ii)  Non-metals are bad conductors of heat and electricity. Carbon (graphite) is an exception. It is conductor of electricity and is used in making electrodes.

(iii) Non-metals are non-lustrous and dull except graphite and iodine.

(iv) Non-metals are generally soft except diamond, allotrope of carbon, which is the hardest natural substance.

Diamond is lustrous and hard non-metal

(v)  Non-metals have low tensile strength and they are not strong.

(vi) Non-metals possess low densities. They are light as compared to metals.

(vii) Non-metals are non-sonorous.

(viii) Non-metals may exist in solid, liquid or gaseous state at room temperature. For example sulphur and carbon are solid at room temperature, bromine is liquid and nitrogen, oxygen are gaseous non-metals.

Liquid non-metal : Bromine

(ix) All non-metals possess low melting and boiling points except graphite.

 

1.3       CHEMICAL  PROPERTIES OF METALS

Metals are highly electropositive elements, i.e. they form positive ions by losing electrons.

M   ¾®Mⁿ⁺ + ne^–

For example:

Na  ¾®Na⁺ + e^–

Mg  ¾®  Mg²⁺ + 2e^–

The characteristic chemical properties of metal are due to their electropositive character.

 

1.3.1    Reaction with oxygen (of air)

            Metals combine with oxygen to form their respective oxides. The metal oxides are ionic compounds. Metals lose their valence shell electrons and form positive ions while oxygen atoms receive electrons to form the negative oxide ion. The two oppositely charged ions combine to produce ionic metal oxide.

Metal         +          Oxygen ¾®   Metal oxide

Metal oxides are basic in nature.

 

œ   Metal oxides, being basic, turn red litmus solution blue.

 

The reactivity of metals towards oxygen are different e.g.

(i)   Sodium and potassium react with oxygen at room temperature to form basic oxides.

4Na        +             O₂     ¾®          2Na₂O

Sodium                        Oxygen                 Sodium oxide

 

 

œ   Sodium and potassium react with oxygen even at room temperature, therefore, they are stored under kerosene oil to prevent its reaction with oxygen.

 

Most of the metal oxides are insoluble in water. But some of the metal oxides dissolve in water to form alkalis.

Sodium and potassium oxide are the two metal oxides which are soluble in water and form alkalis.

Na₂O      +           H₂O    ¾®          2NaOH

Sodium oxide                                             Sodium hydroxide

(Basic oxide)                                                     (An alkali)

 

K₂O        +       H₂O    ¾®          2KOH

Potassium oxide                                          Potassium hydroxide

(Basic oxide)                                                    (An alkali)

 

(ii)  Magnesium does not react with oxygen at room temperature. It combines with oxygen only on heating. On heating, magnesium burns in air producing intense heat and light.                     

2Mg        +         O₂     ¾®          2MgO

Magnesium                 Oxygen                  Magnesium oxide

Since reaction of magnesium with oxygen takes place at higher temperature than the reaction of sodium with oxygen, it means magnesium is less reactive than sodium.

(iii) Zinc metal burns in air only on strong heating to form zinc oxide.

2Zn         +         O₂     ¾®          2ZnO

Zinc                     Oxygen                       Zinc oxide

 

            (iv)Iron does not burn in air even on strong heating. The reaction of iron with oxygen takes place less readily than that of zinc.

2Fe         +         O₂     ¾®          2Fe₃O₄

Iron                      Oxygen                        Iron oxide

Fe₃O₄is a mixture of ferrous oxide (FeO) and ferric oxide (Fe₂O₃) and is known as iron
(II, III) oxide.

 

(v)  Copper is very less reactive towards oxygen. Copper burns in air on prolonged strong heating.

2Cu        +         O₂     ¾®          2CuO

Copper                  Oxygen                  Copper(II) oxide

So, the order of reactivity of these metals with oxygen is:

Na > Mg > Zn > Fe > Cu

Silver and gold do not react with oxygen even at high temperature.

 

1.3.2    Amphoteric oxides

Though most of the metal oxides are basic in nature but some of them show basic as well as acidic nature. Those metal oxides which show basic as well as acidic behaviour are known as amphoteric oxides. Aluminium metal and zinc metal forms amphoteric oxides. Amphoteric oxides react with both, acids as well as bases to form salt and water.

 

(i)   Aluminium oxide as basic oxide:Aluminium oxide reacts with hydrochloric acid to form aluminium chloride and water.

Al₂O₃      +           6HCl           ¾®          2AlCl₃         +          H₂O

Aluminium oxide         Hydrochloric acid                                Aluminium chloride            Water               

                                                (Acid)                     (Salt)

 

            (ii)  Aluminiumoxide as acidic oxide: Aluminium oxide reacts with sodium hydroxide to form sodium aluminate and water.

Al₂O₃      +        2NaOH          ¾®         2NaAlO₂      +          H₂O

Aluminium oxide         Sodium Hydroxide              Sodium aluminate              Water               

                                                (Base)                    (Salt)

 

Zinc oxide is an amphoteric oxide which reacts with acids as well as with bases to form salt and water.

(iii) Zinc oxide as basic oxide: Zinc oxide reacts with hydrochloric acid to form zinc chloride (salt) and water.

ZnO               +            2HCl          ¾®         ZnCl₂           +          H₂O

Zinc oxide                     Hydrochloric acid                                Zinc chloride                        Water               

                                                (Acid)                     (Salt)

 

(iv)Zinc oxide as acidic oxide: It reacts with sodium hydroxide to form sodium zincate and water.

ZnO               +            2NaOH      ¾®         Na₂ZnO₂     +          H₂O

Zinc oxide                 Sodium hydroxide               Sodium zincate                    Water               

                                                (Base)                    (Salt)

1.3.3    Reaction of metals with water

Metals react with water and produce metal oxide or metal hydroxide and hydrogen gas. However, the reactivities of metal with water are different.

Metal    +         Water  ¾®    Metal oxide or metal hydroxide    +    Hydrogen

           

2Na    +          2H₂O               ¾®                 2NaOH    +   H₂

Sodium                         Water                                           Sodium hydroxide

Ca      +          2H₂O               ¾®                 Ca(OH)₂  +   H₂

Calcium                        Water                                           Calcium hydroxide

 

            Potassium decomposes water more vigorously than sodium. Hence, potassium is more reactive than sodium. Sodium decomposes water more vigorously than calcium. Hence, sodium is more reactive than calcium.

 

œ   In case of sodium and potassium, reaction is so violent and exothermic that evolved hydrogen immediately catches fire.

 

(ii)  Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen.

Mg          +         2H₂O            ¾®     Mg(OH)₂         +          H₂

Magnesium                   Water (hot)                       Magnesium hydroxide

 

œ   Calcium and magnesium start floating because the bubbles of hydrogen gas formed stick to the surface of metal.

 

(iii) Metals like aluminium, iron and zinc do not react either with cold or hot water. They react with steam to form metal oxide and hydrogen.       

 

2Al          +         2H₂O            ¾®           Al₂O₃         +          3H₂

Aluminium        Steam                                Aluminium oxide

Zn           +          H₂O             ¾®     ZnO     +          H₂

Zinc                            Steam                                   Zinc oxide

3Fe         +         4H₂O            ¾®             Fe₃O₄      +          4H₂

Iron                           Steam                                Iron(II, III) oxide

(Red hot)

            The metals such as lead, copper, silver and gold do not react with water at all.

            Metals evolve hydrogen from water

            Water is a very weak electrolyte. It ionizes to give hydrogen ions (H⁺) and hydroxide ions (OH^–).

H₂O    ¾®     H⁺        +          OH^–

When a metal reacts with water, it loses electrons to reduce H⁺ ions to H atoms.

2H⁺     +          2e^–       ¾®     H₂

The H atom then unites to give H₂ gas.

 

1.3.4    Reaction of metals with Acids

            Metals usually displace hydrogen from dilute acids.

Metal     +       Dilute acid     ¾®       Salt      +          Hydrogen

 

All metals, however do not react with dilute acids.

 

œ   The rates of the reaction with dilute acids can be compared by observing the rate at which bubbles are formed, i.e., the rate of effervescence.

 

            (a)  Reaction of metals with hydrochloric acid

(i)   Sodium metal reacts with dilute hydrochloric acid vigorously to form sodium chloride and hydrogen gas.

2Na          +         2HCl    ¾®       2NaCl           +          H₂

Sodium chloride

This shows that sodium is a very reactive metal.

 

 

 

(iii)Aluminium metal reacts rapidly with dilute hydrochloric acid to form aluminium chloride and hydrogen gas.

 

                         2Al      +         6HCl    ¾®       2AlCl₃           +          3H₂

Aluminium chloride

 

At first, it reacts slowly due to the presence of tough protective layer of aluminium oxide on its surface. But when the outer oxide layer gets dissolved in acid, then fresh aluminium metal is exposed which reacts rapidly.

 

(iv) Zinc reacts with dilute hydrochloric acid to produce zinc chloride and hydrogen but less rapidly than that of aluminium.

 

            Zn        +         2HCl    ¾®      ZnCl₂ +          H₂

Zinc chloride

            This reaction shows that zinc is less reactive than aluminium.

 

Gold and silver metals also do not react with dilute acids.

 

 

Metals displace hydrogen from dilute acids

Those metals which are more reactive than hydrogen, i.e., those metals which lose electrons more easily than hydrogen, displace hydrogen from dilute acids to produce hydrogen gas. This is due to the fact that the more reactive metals give electrons easily and those electrons reduce hydrogen ions of acids to hydrogen gas.

The metals like copper and silver which are less reactive than hydrogen, do not displace hydrogen from dilute acids. This is because they do not give out electrons required for the reduction of hydrogen ion present in acids.

 

            (b) Reaction of a metal with dilute sulphuric acid

Metals displace hydrogen from dilute sulphuric acid. For example :

 

2Na (s)                        +          H₂SO₄ (aq)     ¾®     Na₂SO₄ (aq)   +          H₂(g)

Sodium                         Sulphuric acid               Sodium sulphate           Hydrogen

2Al (s)              +          3H₂SO₄ (aq)   ¾®     Al₂(SO₄)₃ (aq) +          3H₂(g)

Aluminium                            Sulphuric acid                      Aluminiumsulphate            Hydrogen

Zn (s)               +          H₂SO₄ (aq)     ¾®     ZnSO₄ (aq)     +          H₂(g)

Zinc                                                        Sulphuric acid                      Zinc sulphate                        Hydrogen

 

            (c)  Reaction of a metal with dilute nitric acid

When a metal reacts with dilute nitric acid, then hydrogen gas is not evolved.

Nitric acid is a strong oxidizing agent. So, as soon as hydrogen gas is formed in the reaction between a metal and dilute nitric acid, the nitric acid oxidizes this hydrogen to water. So, in the reactions of metals with dilute nitric acid, no hydrogen gas is evolved. Now, when nitric acid oxidizes hydrogen to water, then nitric acid itself is reduced to any of the nitrogen oxides (such as dinitrogen monoxide, N₂O; nitrogen monoxide, NO; or nitrogen dioxide, NO₂). The type of oxide formed depends on the nature of metal, the temperature of reaction and concentration of nitric acid.

Very dilute nitric acid, however, reacts with magnesium and manganese metals to evolve hydrogen gas. This is because the very dilute nitric acid is a weak oxidizing agent which is not able to oxidize hydrogen to water.

(i) Magnesium reacts with very dilute nitric acid to form magnesium nitrate and hydrogen gas.

      Mg(s)   +          2HNO₃ (aq)     ¾®     Mg(NO₃)₂ (aq)                        +          H₂(g)

Magnesium          Nitric acid                                              Magnesium nitrate              Hydrogen

(very dilute)

 

(ii) Manganese reacts with very dilute nitric acid to form manganese nitrate and hydrogen gas.

      Mn (s)  +          2HNO₃ (aq)     ¾®     Mn(NO₃)₂ (aq)                        +          H₂(g)

Manganese          Nitric acid                                              Manganese nitrate              Hydrogen

(very dilute)

 

œ   Aqua regia, (Latin for ‘royal water)’ is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3 : 1. It can dissolve gold, even though neither of these acids can do so alone. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that is able to dissolve gold and platinum.

Aqua Regia

 

1.3.5    Reaction of metal with solutions of other metal salts

(a) A more reactive metal displaces a less reactive metal from its salt solution. Such reactions are called displacement reactions.

 

Metal A + Salt solution of  B¾®  Salt solution of metal A + Metal B

            For example :

 

(i)

 

Thus, when zinc pieces are added to copper sulphate solution, then, zinc being more reactive metal than copper, displaces copper from its solution (CuSO₄) so that copper is set free. The blue colour of CuSO₄ solution fades due to formation of ZnSO₄ (colourless). A reddish brown deposit of copper metal is formed on the surface of zinc.

 

(ii)

(b) When iron nail is added to copper sulphate solution, then, iron being more reactive metal than copper, displaces copper from its solution (CuSO₄) so that Cu is set free. The blue colour of CuSO₄fadeds and it changes into light green due to formation of iron sulphate (FeSO₄).

 

Relative activities or reactivities of metals

Metals have been arranged in decreasing order of their activities (or reactivities) in the activity series. After performing displacement experiments, the following series known as the reactivity or activity series has been developed as follows :

 

It is clear from the series that the metals lying above the hydrogen are more reactive than the metals lying below the hydrogen. Thus any metal can displace the metals lying below it from its solution.

 

            (c) Significance of activity series

(i) The metals above hydrogen in the activity series have greater tendency than hydrogen to give up electrons in their solutions. Such metals are called electropositive metals.

 

The electropositive character of metals becomes less pronounced as we go down the series. For example, potassium (K), the first metal in the series is the most electropositive, while platinum (Pt), the last metal of the series is the least electropositive.

(ii) The metals above hydrogen in the series can liberate hydrogen when treated with an acid solution. Thus, magnesium and zinc react with dilute solutions of sulphuric acid to produce hydrogen gas.

Mg       +          H₂SO₄             ¾®                 MgSO₄            +          H₂

Zn        +          H₂SO₄             ¾®                 ZnSO₄             +          H₂

In these reactions, electrons released by metals are accepted by H⁺ or H₃O⁺ ions present in the acid solution.

Mg       ¾®     Mg²⁺                +                      2e^–

2H⁺      +          2e                    ¾®                 H₂

or         2H₃O⁺ +          2e^–                   ¾®                 2H₂O   +          H₂

Thus, H⁺ (or H₃O⁺) ions act as oxidizing agents while the metal acts as reducing agent.

(iii) A more electropositive metal can replace a less electropositive metal from the solution of a salt of the less electropositive metal. For example, when an iron rod is dipped into a solution of copper sulphate, reddish coloured copper is deposited on the iron rod.

Fe              +          CuSO₄                            ¾®                 FeSO₄             +          Cu

This is because iron is more electropositive than copper.

 

1.3.6    Reaction of metals with chlorine

Most of the metals react with chlorine to form chlorides. These chlorides are ionic in character.

Metal         +          Chlorine      ¾®         Metal chloride

During this reaction, metal loses electrons and become positively charged whereas chlorine atoms accept electrons and become negatively charged ions. Metals form ionic chlorides because they give electrons to chlorine atoms to form ions.

Metal chlorides are usually solid and conduct electricity in solution or molten state. Thus metal chlorides are electrolytes.

(i)Sodium readily reacts with chlorine to form an electrovalent chloride called sodium chloride

2Na           +          Cl₂¾®                        2NaCl

Sodium Chlorine                        Sodium chloride

 

(ii)Calcium reacts vigorously with chlorine to form an electrovalent chloride called calcium chloride.

Ca             +          Cl₂¾®                        CaCl₂

Calcium         Chlorine                        Calcium chloride

            (v)Iron combines with chlorine, when heated, to form iron (III) chloride or ferric chloride

2 Fe           +          3Cl₂   2 FeCl₃

Iron                 Chlorine                                Ferric Chloride

 

            (vi)On heating, copper reacts with chlorine to form copper (II) chloride or cupric chloride

Cu             +          Cl₂       ¾®     CuCl₂

Copper              Chlorine                                Cupric chloride

 

1.3.7    Reaction of metals with hydrogen

Most of the metals do not react with hydrogen. A few reactive metals such as sodium, potassium and calcium react with hydrogen to form ionic hydride.

For example:

(i) When hydrogen gas is passed over heated sodium, then sodium hydride is formed.

2Na            +          H₂        ¾®     2NaH

Sodium                                                             Sodium hydride

(iii) When hydrogen gas is passed over heated calcium, then calcium hydride is formed.

Ca                         +          H₂        ¾®     CaH₂

Calcium                                                         Calcium Hydride

The comparatively less reactive metals like zinc, copper and iron do not react with hydrogen.

 

1.3.8    Reaction of metal with metal oxides

A more reactive metal can displace a less reactive metal from its oxide. The reactive metals are good reducing agents because they can easily give the electrons needed for reduction. For example:

(i) Aluminium can displace iron from ferric oxide.

2Al              +          Fe₂O₃              ¾®                 2Fe                  +          Al₂O₃

Aluminium                            Ferric oxide                                           Iron                                         Aluminium oxide

(ii) Magnesium can displace copper from copper oxide.

CuO    +          Mg                   ¾®                 MgO                +          Cu

Copper (II) Oxide            Magnesium                                          Magnesium oxide                               Copper

 

1.4       CHEMICAL PROPERTIES OF NON-METALS

The non-metals are electronegative elements, i.e. they form negative ions by gaining electrons. In other words, non-metals have tendency to gain electrons and hence undergo reduction.

 

1.4.1    Reaction of non-metals with Oxygen

            Non-metals react with oxygen to form oxides. These oxides are covalent in nature and are generally gaseous. Most of these oxides when dissolved in water give acidic solution. These solutions turn blue litmus red. For example, sulphur when burnt in the presence of oxygen forms sulphur dioxide.

(a)  S    +   O₂  ¾®     SO₂

Sulphur dioxide

 

Sulphur dioxide when dissolved in water forms sulphurous acid.

 

SO₂         +          H₂O     ¾®       H₂SO₃

Sulphurdioxide                                            Sulphurousacid

 

(b) Similarly, carbon forms carbon dioxide and phosphorus forms phosphorus pentoxide. These oxides also give acidic solutions when dissolved in water.

C               +          O₂        ¾®     CO₂

Carbondioxide

CO₂           +          H₂O     ¾®     H₂CO₃

Carbonic acid

 

P₂O₅          +          3H₂O   ¾®     2H₃PO₄

Phosphoric acid

 

On the other hand, certain oxides of non-metals are neutral. Some examples are nitrous oxide (N₂O), carbon monoxide (CO), etc. These oxides do not have any effect on litmus paper.

 

1.4.2    Reaction of non-metals with water

Non-metals do not react with water or steam to evolve hydrogen gas because they do not lose electrons to reduce hydrogen ions (H⁺) into hydrogen gas.

1.4.3    Reaction of non-metals with acids

Non-metals do not displace hydrogen from acids. For the liberation of hydrogen the non-metals should be able to reduce H⁺ ions to H₂ gas by supplying electrons. However, non-metals are electronegative elements and hence have more tendency to accept electrons rather than donating. Hence, non-metals do not produce hydrogen gas on reaction with acids.

1.4.4    Reaction of non-metals with salts solution

A more reactive non-metal displaces a less reactive non-metal from its salt solution. For example :

When chlorine is passed through a solution of sodium bromide, then sodium chloride and bomine are formed :

2NaBr              +          Cl₂       ¾®     2NaCl              +          Br₂

Sodium Bromide                 Chlorine                                Sodium Chloride                 Bromine

 

1.4.5    Reaction of non-metals with chlorine

Non-metals react with chlorine to form chlorides. These chlorides are covalent in nature and are generally gaseous or liquids.

(a) Hydrogen reacts with chlorine to form covalent chloride called hydrogen chloride.

H₂              +          Cl₂       ¾®                 2HCl

Hydrogen chloride

(b) Phosphorus reacts with chlorine to form covalent chloride called phosphorus trichloride and phosphorus pentachloride.

P₄              +          6Cl₂     ¾®                 4PCl₃

Phosphorus trichloride

P₄              +          10Cl₂   ¾®                 4PCl₅

Phosphorus pentachloride

 

1.4.6    Reaction of non-metals with hydrogen

Non-metals on reaction with hydrogen form hydrides which are covalent in nature. Most of these hydrides are gases or liquids. Some examples are :

(a)  2H₂            +          O₂        ¾®                 2H₂O (Water)

(b) N₂              +          3H₂      ¾®                 2NH₃(Ammonia)

(c)  H₂              +          S          ¾®                 H₂S (Hydrogen sulphide)

 

1.4.7    Oxidising Nature

Since non-metals have great tendency to accept electrons, therefore, many non-metals act as good oxidizing agents. For example, fluorine is the strongest oxidizing agent among all the non-metals.

(a)  Zn              +          S          ¾®                 Zn²⁺ S^2–

Zinc sulphide

(b)  2Na           +          Cl₂       ¾®                 2Na⁺ Cl^–

Sodium chloride

 

œ   Although hydrogen and carbon are non-metals, they behave as reducing agents because of their low electronegativity.

 

1.6       USES OF METALS

Metals are used for a large number of purposes. Some of the uses of metals are given below:

(i)   Copper and aluminium metals are used to make wires to carry electric current. This is because copper and aluminium have very low electrical resistance and hence very good conductors of electricity.

(ii)  Iron, copper and aluminium metals are used to make house-hold utensils and factory equipments.

(iii) Iron is used as a catalyst in the preparation of ammonia gas by Haber’s process.

(iv) Zinc is used for galvanizing iron to protect it from rusting.

(v) Chromium and nickel metals are used for electroplating and in the manufacture of stainless steel.

(vi) Thealuminium foils are used in packaging of medicines, cigarettes and food materials.

Aluminium Foil

 

(vii) Silver and gold metals are used to make jewellery. The thin foils made of silver and gold are used to decorate sweets.

(viii) The liquid metal ‘mercury’ is used in thermometers.

(ix) Sodium, titanium and zirconium metals are used in atomic energy (nuclear energy) and space science projects.

(x) Zirconium metal is used in making bullet-proof alloy steels.

 

1.7       USES OF NON-METALS

The important uses of non-metals are as follows:

(i) Hydrogen is used in the hydrogenation of vegetable oils to make vegetable ghee (or vanaspati ghee).

(ii) Hydrogen is used in the manufacture of ammonia (whose compounds are used as fertilisers).

(iii) Liquid hydrogen is used as a rocket fuel.

(iv) Carbon (in the form of graphite) is used for making the electrodes of electrolytic cells and dry cells.

(v) Nitrogen is used in the manufacture of ammonia, nitric acid and fertilisers.

(vi) Due to its inertness, nitrogen is used to preserve food materials.

(vii) Compounds of nitrogen like Tri Nitro Toluene (TNT) and nitroglycerine are used as explosives.

(viii) Sulphur is used for manufacturing sulphuric acid.

(ix) Sulphur is used as a fungicide and in making gun powder.

(x) Sulphur is used to the vulcanisation of rubber.

 

When metals react with non-metals, they form ionic compounds (which contain ionic bonds). On the other hand, when non-metals react with other non-metals, they form covalent compounds (which contain covalent bonds).

The force which holds a number of atoms together within the compound or the molecule is called a chemical bond.

To understand the formation of chemical bond ‘between the atoms of metals and non-metals’ or ‘between the atoms of two non-metals’, it is necessary to know the reason for the unreactive nature (or inertness) of noble gases.

There are some elements in group 18 of the periodic table which do not combine with other elements. These elements are : Helium, Neon, Agron, Krypton, Xenon and Radon. They are known as noble gases or inert gases because they do not react with other elements to form compounds. In other words, inert gases do not form chemical bonds.

That is why they are called noble gases or inert gases and exist as monoatomic gases, i.e., as He, Ne, Ar, Kr, Xe and Rn.

 

            Modes of Chemical combination

As explained above, when atoms combined, they complete their octets (or duplets) by any one of the following methods :

(i) by losing one or more electrons (to another atom).

(ii) by gaining one or more electrons (from another atom).

(iii) by sharing one or more electrons (with another atom).

 

 

2.1       IONS

An ion is an electrically charged atom (or group of atoms). Examples of the ions are : sodium ion, Na⁺, magnesium ion, Mg²⁺, chloride ion, Cl^–, and oxide ion, O^2–. An ion is formed by the loss or gain of electrons by an atom.

 

 

2.1.1    A positively charged ion is known as cation

Sodium ion, Na⁺, and magnesium ion, Mg²⁺, are cations because they are positively charged ions. A cation is formed by the loss of one or more electrons by an atom. For example, sodium atom loses one electron to form a sodium ion, Na⁺, which is a cation.

 

Na                   –                      e–                    ¾®                 Na⁺

Sodium atom                                       Electron                                 Sodium ion

Electronic    K L M                                                                                                                                K  L

Configuration: 2, 8, 1                                                                                                                          2, 8

(unstable electron                                                                                               (Stable, neon gas electron

arrangement)                                                                                                       arrangement)

 

2.1.2    A negatively charged ion is known as anion

Chloride ion, Cl^–, and oxide ion, O^2–, are anion because they are negatively charged ion. An anions is formed by the gain of one or more electrons by an atom. For example, a chlorine atom gains (accepts) one electron to form a chloride ion, Cl^–, which is an anion

 

Cl                     +                      e–                    ¾®                 Cl^–

Chlorine atom                      Electron                                           Chloride ion

Electronic    K L M                                                                                                                                K  L M

Configuration: 2, 8, 7                                                                                                                          2, 8, 8

(unstable electron                                                                                               (Stable, neon gas electron

arrangement)                                                                                                       arrangement)

 

 

 

2.2       ELECTRONIC DOT REPRESENTATION (LEWIS SYMBOLS)

In the formation of a chemical bond between two atoms, only the electrons of the outermost shell are involved (as the inner shell electrons are well protected). These electrons present in the outermost shell are called valence electrons.

G.N. Lewis introduced a simple method of representing the valence electrons by dots or small crosses around the symbol of the atom. These symbols are known as electron dot symbols or Lewis symbols. These symbols ignore the inner shell electrons.

Electron dot (Lewis symbols) of some common elements

Element	Symbol	Atomic No.	Electronic configuration				Valence electrons	Lewis symbol
			K	L	M	N
Hydrogen	H	1	1				1
Helium	He	2	2				2
Lithium	Li	3	2	1			1
Beryllium	Be	4	2	1			1
Boron	B	5	2	3			3
Caron	C	6	2	4			4
Nitrogen	N	7	2	5			5
Oxygen	O	8	2	6			6
Fluorine	F	9	2	7			7
Neon	Ne	10	2	8			8
Sodium	Na	11	2	8	1		1
Magnesium	Mg	12	2	8	2		2
Aluminium	Al	13	2	8	3		3
Chlorine	Cl	17	2	8	7		7
Potassium	K	19	2	8	8	1	1
Calcium	Ca	20	2	8	8	2	2

 

2.3       TYPES OF CHEMICAL BONDS

There are two types of chemical bonds :

(i) Ionic bond, and

(ii) Covalent bond.

Ionic bonds are formed by the transfer of electrons from one atom to another whereas covalent bonds are formed by the sharing of electrons between two atoms. Ionic bond is also called electrovalent bond.

2.3.1    Ionic bond

The chemical bond formed by the transfer of electrons from one atom to another is known as an ionic bond. The transfer of electrons takes place in such a way that the ions formed have the stable electron arrangement of an inert gas. The ionic bond is called so because it is a chemical bond between oppositely charged ions.

When a metal reacts with a non-metal, transfer of electrons takes place from metal atoms to the non-metal atoms, and an ionic bond is formed.

The compounds containing ionic bonds are called ionic compounds. Ionic compounds are made up of ions. Lets understand with the help of examples.

 

(i)Formation of Sodium Chloride.

Atomic number of sodium (Na) = 11

\ Its electronic configuration is 2, 8, 1.

It has only one electron in the valence shell. It loses this electron to acquire the stable electronic configuration 2, 8 (similar to that of neon) and form sodium ion (Na⁺).

Na^´                  ¾®     Na⁺      +          e^–

Sodium atom                       Sodium ion

(2, 8, 1)                         (2, 8)

Atomic number of chlorine (Cl) = 17

\   Its electronic configuration is 2, 8, 7.

It has seven electrons in the valence shell. It gains one electron to acquire the stable electronic configuration 2, 8, 8 (similar to that of argon) and form chloride ion (Cl^–)

+  e^–           ¾®

Chlorine atom                               Chloride ion

(2, 8, 1)                         (2, 8)

Thus, when a sodium atom and a chlorine atom approach each other, an electron is transferred from sodium atom to chlorine atom. In other words, sodium loses one electron to form Na⁺ ion and chlorine gains that electron to form Cl^– ion. As a result, both acquire the stable nearest noble gas configuration. These oppositely charged ions are then held together by electrostatic forces of attraction forming the compound Na⁺Cl^– or simply written as NaCl. The transfer of electron may be represented in one step as follows:

 

 

Sodium atom      Chlorine atom                       Sodium chloride

(2, 8, 1) (2, 8, 7)

 

 

 

(ii)Formation of Magnesium chloride.

Atomic number of magnesium (Mg) = 12

\   Its electronic configuration = 2, 8, 2

It loses two electrons from the valence shell to acquire the nearest noble gas configuration of neon (2, 8) and form Mg²⁺ ion.

Atomic number of chlorine (Cl) = 17

Its electronic configuration = 2, 8, 7

It needs to gain only one electron in the valence shell to acquire the nearest noble gas configuration of argon (2, 8, 8) and form chloride ion (Cl^–). Now, as magnesium atom has to lose two electrons and a chorine atom can gain only one electron, therefore, two chlorine atoms will be required to accept the two electrons, one by each chlorine atom. Thus, the transference of two electrons from one Mg atom to two Cl atoms may be represented as follows:

Thus, here one magnesium atom combines with two chlorine atom. Therefore, the formula of magnesium chloride is  or Here, again the oppositely charged ions are held together by electrostatic forces of attraction.

(iii)Formation of Magnesium Oxide.

Atomic number of magnesium (Mg) = 12

\   Its electronic configuration = 2, 8, 2

It loses two electrons from the valence shell to acquire the stable configuration of neon
(2, 8) and form Mg²⁺ ion.

Atomic number of oxygen (O) = 8

\   Its electronic configuration = 2, 6

It gains two electrons in the valence shell to acquire the stable configuration of neon again (2, 8) and form oxide ion (O^2–).

Thus, in the formation of magnesium oxide, two electrons are transferred from magnesium atom to oxygen atom as represented below :

 

2.3.2    Covalent bond

The chemical bond formed by the sharing of electrons between two atoms is known as a covalent bond. The sharing of electrons takes place in such a way that each atom in the resulting molecule gets the stable electron arrangement of an inert gas.

Whenever a non-metal combines with another non-metal, sharing of electrons takes place between their atoms and a covalent bond is formed. A covalent bond can also be formed between two atoms of the same non-metal.

Covalent bonds are of three types :

(i) Single covalent bond

(ii) Double covalent bond

(iii) Triple covalent bond

 

            (i) Single Bond

Single bond is formed by the sharing of one pair of electrons between two atoms. A single covalent bond is formed by the sharing of two electrons between the atoms, each atom contributing one electron for sharing.

For example, a hydrogen molecule H₂, contains a single covalent bond and it is written as H : H, the two dots drawn between the hydrogen atoms represent a pair of shared electrons which constitutes the single bond. A single covalent bond is denoted as H–H.

            Formation of H₂ molecule

The formation of a hydrogen molecule from two hydrogen atoms can be shown by the following diagram:

 

 

 

(b)Formation of a Water Molecule, H₂O :

.

            (ii)  Double Bond

Double bond is formed by the sharing of two pairs of electrons between two atoms. A double covalent bond is formed by the sharing of four electrons between two atoms, each atom contributing two electrons for sharing.

It is represented by putting two short line (=) between the two atoms. For example, oxygen molecule, O₂, contains a double bond between two atoms and it can be written as O=O.

(a)Formation of Oxygen Molecule, O₂ : Oxygen atom is very reactive and cannot exist free because it does not have the stable, inert gas electron arrangement in its valence shell. The atomic number of oxygen is 8, so its electronic configuration is 2, 6. Thus, an oxygen atom has six electrons in its outermost shell. It requires two more electrons to achieve the stable, eight electron inert gas configuration. The oxygen atom gets these electrons by sharing its two electrons with the two electrons of another oxygen atom. So, two oxygen atoms share two electrons each and form a stable oxygen molecule.

Thus, in the oxygen molecule, the two oxygen atoms are held together by a double bond.

 

(b)Formation of Carbon Dioxide Molecule, CO₂ :

There are two double bonds in a carbon dioxide molecule.

 

            (iii) Triple Bond

A triple bond is formed by the sharing of three pairs of electrons between two atoms. A triple bond is formed by the sharing of six electrons between two atoms, each atom contributing three electrons for sharing. A triple bond is actually a combination of three single bonds, so it is represented by putting three short line (º) between the two atoms. Nitrogen molecule, N₂, contains a triple bond, so it can be written as N ºN.

(a)Formation of a Nitrogen Molecule, N₂ : A nitrogen atom is very reactive and cannot exist free because it does not have the stable electron arrangement of an inert gas. The atomic number of nitrogen is 7, so its electronic configuration is 2, 5. This means that a nitrogen atom has five electrons in its outermost shell. Since a nitrogen atom has five electrons in its outermost shell, it needs three more electrons to achieve the eight electron structure of an inert gas and become stable. So, two nitrogen atoms combine together by sharing three electrons each to form a molecule of nitrogen gas.

2.3.3    Properties of ionic compounds

(i)Physical state : Most of the ionic compounds are crystalline solids. They are relatively hard because of strong electrostatic forces of attraction between the oppositely charged ions. They are brittle and break into pieces on applying force.

Laboratory Test : Collect samples of sodium chloride, potassium chloride, barium chloride or any other salt from the laboratory. We see that they are all solids and brittle, i.e., they break into pieces if some force (e.g., hammer) is applied on them.

(ii)Solubility : They are soluble in water (which is a polar solvent) but insoluble in organic solvents like benzene, alcohol, ether, chloroform etc.

Laboratory Test : Shake the samples of the salts taken above one by one with water, alcohol, petrol and kerosene oil taken in four test tubes separately. We find that the salts are soluble in water but insoluble in other solvents.

(iii)Melting points and boiling points : They have high melting and boiling points. This is because in the ionic compounds, the oppositely charged ions are held together by strong electrostatic forces of attraction. Hence, they require a lot of heat to cut off these forces of attraction and break them into ions.

Laboratory Test : Take any of the above salts in a dry pyrex or corning glass test tube (which can withstand high temperature). Hold in a test tube holder and heat it in the Bunsen burner. We will observe that it is difficult to melt it.

(iv)Colour in the flame : Most of the salts when brought into the flame, impart characteristic colours to the flame.

Laboratory Test : Take solid sodium chloride on a metal spatula and bring it into the blue flame of the Bunsen burner. We see that it imparts golden yellow colour to the flame. Similarly, potassium salts impart violet colour and barium salts impart green colour to the flame.

Heating a salt sample on a spatula

(v)Electrical conductivity : When dissolved in water, ionic compounds dissociate to produce free ions in the solution. Similarly, on melting, the ionic compound splits to produce free ions in the melt. As ions can conduct electricity, therefore, ionic compounds conduct electricity in the aqueous solution or in the molten state.

Laboratory Test : Take water in a beaker and dissolve any one of the above salts into it. Set up the electrical circuit as shown in figure below :

Testing the conductivity of a salt solution

On switching on the current, the bulb is found to glow. This shows that the salt solution conducts electricity.

 

œ   An ionic compound does not conduct electricity in the solid state. This is because in the solid state, the ions are held together by strong electrostatic forces of attraction and hence are not free to move.

 

 

The earth’s crust is the major source of metals. Sea water also contains some soluble salts of metals like sodium chloride, magnesium chloride etc.

Some metals are found in the earth’s crust in free state where as some are found in the form of their compounds. The metals at the bottom of activity series are least reactive. Therefore, they are often found in free state. For e.g. gold, silver, platinum, copper.

The metals at the top of activity series (K, Na, Ca, Mg etc.) are so reactive, that they are never found in nature as free elements but are found in the form of their compounds. The metals in the middle of activity series (Al, Zn, Fe, Pb etc.) are moderately reactive. They are found in earth’s crust mainly as oxides, sulphides or carbonates.

Thus, the metals occur in nature in two forms :

(a)NativeState : A metal is said to exist in native state if it exists in the elementary form in the crust of the earth.

(b)Combined State : A metal is said to exist in combined state if it is found in the earth’s crust in the form of its stable compounds.

 

3.1       MINERALS AND ORES

In the combined state, the metals are found in the earth crust as oxides, carbonates, sulphides, silicates, phosphates etc.

The inorganic elements or compounds which occur naturally in the earth’s crust are known as minerals. At some places, the minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. Those minerals from which a metal can be profitably extracted are called ores.

When an ore is mined from earth, it is always found to be contaminated with sand and rocky materials. These impurities of sand and rocky materials present in the ore are known as gangue.

 

œ   Aluminium is available in earth crust in the form of two well known minerals, bauxite (Al2O3 . 2H2O) and china clay (Al2O3 . 2SiO2. 2H2O). But extraction of aluminium is cheaper and easier from bauxite. Hence ore of aluminium is bauxite. Therefore, it is concluded that all the ores are minerals but all the minerals are not ores.

Bauxite

3.2       TYPES OF ORES

Most of the ores are either oxides or sulphides. Some of the ores are also carbonates and halides. The various types of ores and their examples are given below in the table :

Various Types of Ores

S.No.	Type	Examples
1.	Oxide ores	Haematite (Fe2O3) Bauxite (Al2O3 . 2H2O) Magnetite (Fe3O4)
2	Sulphide ores	Copper pyrites (CuFeS2) Iron pyrites (FeS2) Zinc blende (ZnS) Cinnabar (HgS)
3.	Carbonate ores	Limestone (CaCO3) Siderite (FeCO3) Calamine (ZnCO3)
4.	Halide ores	Rock salt (NaCl) Horn silver (AgCl) Fluorspar (CaF2)

 

Some common ores are given in table below :

 

Some Common Ores

S.No.	Elements	Ores
1.	Iron	Haematite (Fe2O3) Magnetite (FeO . Fe2O3) Iron pyrites (FeS2)
2	Aluminium	Bauxite (Al2O3 . 2H2O)
3.	Manganese	Pyrolusite (MnO2)
4.	Calcium	Limestone (CaCO3)
5.	Limestone	Dolomite (MgCO3 . CaCO3)
6.	Copper	Copper pyrites (CuFeS2)
7.	Mercury	Cinnabar (HgS)
8.	Zinc	Zinc blende (ZnS), Calamine (ZnCO3)
9.	Lead	Lead glance (PbS)
10.	Sodium	Rock salt (NaCl)
11.	Silver	Horn silver (AgCl)

 

After the mining of the ore from the ground, it must be decomposed to give pure metal. This is known as extraction.

The extraction of metals from their ores and then refining them for use is known as metallurgy. The process of metallurgical operations consists of mainly three steps :

1. Concentration of ore or enrichment of ore
2. Reduction of concentrated ore
3. Refining

4.1       ENRICHMENT OF ORE

Ores mined from the earth usually contain large amounts of earthy impurities like soil, sand, etc., called gangue. The impurities must be removed from the ore prior to the extraction of the metal. The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore. The various methods used for the enrichment of ores are explained below :

 

4.1.1    Hydraulic washing

It is used for the enrichment of oxide ores. It is based on the difference in the densities of the gangue and the ore particles. The gangue particles are generally lighter as compared to ore particles. In this process, the crushed and finely powdered ore is washed with a stream of water. The lighter gangue particles are washed away, leaving behind the heavier ore particles.

 

Hydraulic Washing

 

 

4.1.2    Froth Floatation Process

            It is used to separate the gangue from the sulphide ores, especially those of copper, zinc and lead. In this process, the finely powdered ore is mixed with water in a large tank to form a slurry. Then some pine oil is added to it. The sulphide ores are preferentially wetted by the pine oil, whereas the gangue particles are wetted by the water. When air is blown through the mixture, the lighter oil froth carrying the metal sulphides rises to the top of the tank and floats as scum. It is then skimmed off and dried. The gangue particles being heavier, sink to the bottom of the tank.

 

4.1.3    Electromagnetic Separation

            It is used to enrich the magnetic ores (iron ore). The separation is done by using magnetic separators. A magnetic separator consists of a leather belt that moves over two rollers, out of which one is an electromagnet. The finely powdered ore is dropped over the moving belt at one end. When the ore falls down from the moving belt at the other end, the magnetic portion in the ore is attracted by the magnet and forms a heap nearer to the roller. The non-magnetic impurities fall away from the magnetic ore and form a separate heap.

 

Magnetic separation :The finely powdered ore is dropped over a moving belt which passes over a
magnetic roller. The particles attracted by the magnet form a separate pile.

4.1.4    Chemical Separation

The process of chemical separation makes use of differences between the chemical properties of the gangue and the ore.

For example, bauxite, Al₂O₃ .2H₂O, is an impure form of aluminium oxide. The impurities mainly present in it are iron (III) oxide (Fe₂O₃) and sand (SiO₂). The iron (III) oxide gives it a brown-red colour. Baeyer’s method is used to obtain pure aluminium oxide from bauxite ore. In this method, the finely powdered ore is treated with hot sodium hydroxide solution. The aluminium oxide present in bauxite ore reacts with sodium hydroxide to form water soluble sodium aluminate.

Al₂O₃(s)     +          2NaOH (aq)                ®           2NaAlO₂ (aq)           +          H₂O(l)

(From impure bauxite ore)                                                                                   Sodium aluminate

Iron (III) oxide does not dissolve in sodium hydroxide solution. It is, thus, separated by filtration. Silica reacts with sodium hydroxide to form water soluble sodium silicate.

SiO₂                +          2NaOH            ®                    Na₂ SiO₃                     +          H₂O

Sodium silicate

The filtrate containing sodium aluminate and sodium silicate is then stirred with some aluminium hydroxide to induce the precipitation of aluminium hydroxide. The impurity, silica, remains dissolved as sodium silicate in the solution.

Na AlO₂ (aq)   +          2H₂O (l)           ®        Al(OH)₃(s)                   +      NaOH(aq)

Aluminium hydroxide

Aluminium hydroxide is then filtered off, washed, dried and ignited to get pure aluminium oxide, which is called alumina.

 

Alumina

4.2       REDUCTION OR CONVERSION OF CONCENTRATED ORE INTO METAL

The method used for the extraction of the metal from the concentrated ore depends upon the nature of the metal. Based on their reactivity, the metals have been grouped into three categories :

(a)  Metals of low reactivity

(b)  Metals of medium reactivity

(c)  Metals of high reactivity

 

 

 

The activity series and related metallurgy is given below :

Metal	Method of extraction
K Na Ca Mg Al	Electrolysis of molten chloride or oxide
Zn Fe Pb Cu	Reduction of oxide with carbon
Cu Hg	Heating sulphide in air (Reduction by heat alone)
Ag Au Pt	Found in native state (as metals)

4.2.1    Extraction of metals low in the Activity Series

Metals low in the activity series are very unreactive. The oxides of these metals can be reduced to metals by heating alone. For example

(a)  Cinnabar (Hgs) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO).

2HgS(s)       +        3O₂(g)     2HgO(s)       +      2SO₂(g)

Mercury Sulphide            oxygen              Mercury oxideSulphur dioxide

 

œ   As the common ores of these metals are sulphide ores, so, to obtain metal from these ores, they are converted into metal oxide by heating strongly in presence of excess of air. This process is called roasting.

 

Mercuric oxide is then reduced to mercury on further heating.

2HgO(s)             2Hg(l)              +      O₂(g)

Mercury oxide                 Mercury metal     Oxygen

 

(b) Similarly, when copper glance (Cu₂S), an ore of copper, is subjected to roasting, it directly gives copper according to the following reactions :

2Cu₂S(s)        +      3O₂(g)                 2 Cu₂O (s)       +      2 SO₂ (g)

Copper glance                Oxygen (from air)            Cuprous oxide     Sulphur dioxide

or Copper (I) oxide

 

2Cu₂O(s)        +      Cu₂S(s)               6 Cu (s)       +      2 SO₂ (g)

Cuprous oxide                Copper glance    Copper                  Sulphur dioxide

 

 

4.2.2    Extraction of Metals in the Middle of Activity Series

The metals in the middle of the reactivity series such as iron, zinc, lead, copper etc. are moderately reactive.

These metals are found in nature in the form of their oxide, sulphide or carbonate ores. Further, as it is easier to reduce oxides than sulphides or carbonates, therefore, the sulphide and carbonate ores are first converted into the corresponding metal oxides. Thus, the different steps involved for the extraction of the metal from the concentrated ore are as follows :

(a)Conversion of the carbonate or sulphide ore into metal oxide. This is done by either of the following two methods :

(i) Calcination (For carbonate ores). It is the process of heating the ore strongly in the absence of air. The metal carbonate decomposes to form metal oxide. For example,

ZnO(s) + CO₂(g)

 

(ii)Roasting (For sulphide ores). It is the process of heating the ore strongly in the presence of excess of air. As a result, the sulphide ore is converted into metal oxide. For example,

 

 

œ   A carbonate ore is converted into oxide by calcination whereas sulphide ore is converted into oxide by roasting.

 

(b)Reduction of the metal oxide to metal. As these metals are moderately reactive, their oxides cannot be reduced by heating alone. Hence, their oxides are reduced to metals by using a suitable reducing agent such as carbon or some highly reactive metals like sodium, calcium, aluminium etc. A few examples are :

(i)Reduction by heating with carbon. Carbon in the form of coke or coal is the cheapest reducing agent. It combines with the oxygen of the metal oxide forming carbon monoxide. As a result, metal oxide is reduced to metal. For example, oxides of iron and zinc are reduced to their respective metals by heating with coke.

 

 

Carbon monoxide formed also acts as reducing agent and further reduces the metal oxide to metal.

 

 

œ   The reduction of metal oxides by heating with coke is called smelting. Before carrying out the reduction with carbon, if the calcined or roasted ore still contains infusible impurities of earthy matter (gangue), an additional substance (called flux) is added which combines with the gangue to form a fusible material (called slag). It does not mix with the molten metal, i.e., it forms a separate layer and hence can be easily separated out. Flux + Gangue ®  Slag             For example, during the manufacture of iron, to remove the impurities of silica (SiO2), limestone (CaCO3) is added as flux. At high temperature, CaCO3 decomposes to give CaO which combines with silica to form the fusible slag of calcium silicate.

 

œ   The above reactions are displacement reactions as the more active aluminium displaces the less active metal like Mn, Cr and Fe from their oxides.

 

4.2.3    Extraction of Metals towards the Top of Activity Series

The metals high up in the reactivity series are very reactive. They cannot be obtained from their compounds by heating with carbon. For example, carbon cannot reduce the oxides of sodium, magnesium, calcium, aluminium, etc., to the respective metals. This is because these metals have more affinity for oxygen than carbon. These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides.

During electrolysis, the metal ions, being positive are liberated at the cathode (the negatively charged electrode), whereas the chlorine is liberated at the anode (the positively charged electrode). So, during the electrolysis of molten salts, the metals are always produced at the cathode (negative electrode).

 

œ   Molten salt means melted salt. Salts are melted by strong heating.

 

(i)Electrolysis of molten sodium chloride. When sodium chloride melts, it splits into sodium ion (Na⁺) and chloride ions (Cl^–).

Electrolysis of molten sodium chloride

 

When electricity is passed through the melt, Na⁺ ions go to the cathode whereas  ions are liberated at the anode. ions gain electrons at the cathode and are thus reduced to sodium atoms. ions lose electrons at the anode and are thus oxidized to chlorine atoms. These chlorine atoms then combine with each other to form chlorine (Cl₂) gas. The reactions taking place are as follows :

 

At Cathode : 

At Anode :    

 

Thus, Na is obtained at the cathode whereas chlorine gas is liberated at the anode.

œ   It is important to note that if electrolysis of aqueous solution of sodium chloride is carried out, the sodium metal obtained at the cathode reacts with water to form sodium hydroxide and hydrogen gas (2 Na + 2H2O ® 2 NaOH + H2). Thus, instead of sodium metal, hydrogen gas is liberated at the cathode.

 

Similarly, calcium and magnesium are also obtained by the electrolysis of their fused chlorides.              

œ   It may be noted that melting point of alumina is very high. Hence, it is dissolved in molten cryolite (Na3AlF6) before electrolysis. The mixture has a much lower melting point and a better conductor of electricity than molten alumina.

4.3       REFINING OF METALS

The metals produced by various reduction processes described above are not very pure. They contain impurities. Therefore, they must be purified to obtain pure metals. The process of purification of impure metals is known as refining of metals. The method used for refining an impure metal depends upon (i) the nature of the element, and (ii) the nature of impurities present in it. Therefore, different methods are used for refining different metals. The impure metals are purified by any one of the following methods:

4.3.1    Liquation

This method is used for refining those metals which have low melting points such as tin, lead, etc. In this method, a sloping hearth is used. The hearth is maintained at a temperature that is slightly above the melting point of the metal. When the impure metal is placed at the top of the hearth, it melts and flows down the hearth. The solid impurities, whose melting point is higher than the melting point of the metal, are left behind on the hearth. The pure metal is collected at the bottom of the hearth.

 

4.3.2    Distillation

            The more volatile metals like mercury and zinc are purified by this method. The impure metal is heated in a retort. The pure metal distils over and is condensed in a receiver. The impurities are left behind in the retort.

4.3.3    Electrolytic Refining

            This is the most widely used method for refining impure metals. Many metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically. In this process, the impure metal is made as anode and a thin strip of pure metal is made as cathode. A solution of the metal salt is used as an electrolyte. On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode. The soluble impurities go into the solution, whereas the insoluble impurities settle down at the bottom of the anode and are known as anode mud.

 

Electrolytic refining of copper. The electrolyte is a solution of acidified copper sulphate. The anode is impure copper, whereas, the cathode is a strip of pure copper. On passing electric current, pure copper is deposited on the cathode.

 

In case of copper, the block of impure copper is made as anode. The cathode is made of a thin plate of pure copper. The electrolyte is a solution of copper sulphate, containing a small amount of dilute sulphuric acid. On passing electric current, copper dissolves from the anode into the electrolyte. An equivalent amount of copper from the electrolyte is then deposited on the cathode. Crude copper contains very small amounts of iron, silver and gold. The more reactive metals, such as iron, dissolve in the solution and remain there. The less reactive metals, such as silver and gold, fall below the anode in the electrolytic cell as anode mud. These metals are recovered in the native state. The following reactions take place :

 

            At Cathode : 

At Cathode : 

 

A summary of the methods used and the steps applied for the extraction of different types of metals is given below :

Steps involved in the extraction of metals from ores

 

One of the most destructive and annoying processes that occur in nature is the corrosion of metals. This process takes place on the surface of metal when they are exposed to air. Due to corrosion, small holes appear on the surface of the metal and the strength of the metal goes on decreasing Thus, corrosion brings about slow destruction of the metal . Corrosion is defined as :

The process of slow eating up of metals due to their conversion into oxides, carbonates, sulphides, sulphatesetc by the action of atmospheric gases and moisture.

Generally, more reactive the metal, the more easily it corrodes. Metals which lie high up in the activity series (K, Mg, Al, Zn, Fe etc.) easily undergo corrosion while noble metals (such as gold and platinum) which lie at the bottom of the activity series do not corrode readily.

5.1       CORROSION OF IRON

The corrosion of iron is known as rusting. When ordinary iron is exposed to air in presence of moisture, it gets corroded and forms reddish brown scale which is known as rust. Rust is mainly hydrated ferric oxide (Fe₂O₃ . xH₂O). Rust once formed, causes more and more rusting until the whole of the metal is eaten up.

 

 

The two conditions necessary for rusting of iron to occur are :

1. Presence of moisture, and
2. Presence of air (oxygen)

 

5.2       PREVENTION OF RUSTING

            The wasting of iron objects due to rusting causes a big loss to the country’s economy, so it must be prevented. The various common methods of preventing the rusting of iron (or corrosion of iron) are given below :

1. Rusting of iron can be prevented by painting. The most common method of preventing the rusting of iron (or corrosion of iron) is to coat its surface with a paint. When a coat of paint is applied to the surface of an iron object, then air and moisture cannot come in contact with the iron object and hence no rusting takes place. The iron articles such as window grills, railings, steel furniture, iron bridges, railway coaches, ships, and bodies of cars, buses and trucks, etc., are all painted to protect them from rusting.
2. Rusting of iron can be prevented by applying grease or oil. When some grease or oil is applied to the surface of an iron object, then air and moisture cannot come in contact with it and hence rusting is prevented. For example, the tools and machine parts made of iron and steel are smeared with grease or oil to prevent their rusting.
3. Rusting of iron can be prevented by galvanization. The process of depositing a thin layer of zinc metal on iron objects is called galvanization. Galvanisation is done by dipping an iron object in molten zinc metal. A thin layer of zinc metal is then formed all over the iron object. This thin layer of zinc metal on the surface of iron objects protects them from rusting because zinc metal does not corrode on exposure to damp air. The iron sheets used for making buckets, drums, dust-bins and sheds (roofs) are galvanized to prevent their rusting. The iron pipes used for water supply are also galvanized to prevent rusting.

œ   Zinc is more reactive metal than Iron. Therefore, it lose electrons readily and reacts with air, moisture and carbon dioxide of atmosphere to form an invisible thick layer of basic zinc carbonate, ZnCO3 . Zn (OH)3, which protects iron from further corrosion.

 

4. Rusting of iron can be prevented by tin-plating and chromium-plating. Tin and chromium metals are resistant to corrosion. So, when a thin layer of tin metal (or chromium metal) is deposited on iron and steel objects by electroplating, then the iron and steel objects are protected from rusting. For example, tiffin-boxes made of steel are nickel-plated from inside and outside to protect them from rusting. Tin is used for plating tiffin-boxes because it is non-poisonous and hence does not contaminate the food kept in them. Chromium-plating is done on bicycle handle bars and car bumpers made of iron and steel to protect them from rusting and give them a shiny appearance.

5.Rusting of iron can be prevented by alloying it to make stainless steel. When iron is alloyed with chromium and nickel, then stainless steel is obtained. Stainless steel does not rust at all. Cooking utensils, knives, scissors and surgical instruments, etc., are made of stainless steel and do not rust at all. But stainless steel is too expensive to be used in large amounts.

 

œ   There are three commercial forms of iron which differ from each other mainly in their carbon content :             (a) Cast iron or pig iron. It contains about 2 to 5% of carbon, along with impurities such as sulphur, silicon, phosphorus, manganese, etc. It is the least pure form of iron. It is brittle and is very difficult to weld. It is resistant to corrosion and is used for sewage pipes.             (b) Steel. It contains 0.1 to 1.5% carbon and other impurities.             (c) Wrought iron. It is the purest form of iron and contains carbon and other impurities less than 0.2%. It is malleable and can be easily welded. It is used for making wires, chains, electromagnets, etc.

 

An alloy is a homogeneous mixture of two or more metals or metals and non-metals.

The alloys are prepared to improve the mechanical properties like hardness and tensile strength, to resist the atmospheric and chemical corrosion, to produce sound and workable casting, and to lower the melting point.

An alloy containing mercury as one of themetals is known as an amalgam.

Some examples of common alloys are :

(a)  Brass containing copper and zinc metals.

(b)  Bronze containing copper and tin metals.

(c)  Stainless steel containing iron, carbon and chromium.

 

6.2       TYPES OF ALLOY

There are two types of alloys :

1. Ferrous alloys. These are the alloys which contain iron as one of the constituents, such as nickel steel, stainless steel, etc.
2. Non-ferrous alloys. These are the alloys which do not contain iron as one of the constituents, such as bronze, brass, etc.

 

6.4       ALLOYS OF GOLD

            Pure, gold, known as 24 carat gold is very soft. It is, therefore, not suitable for making jewellery. It is alloyed with either silver or copper to make it hard. Generally, in India, 22 carat gold is used for making ornaments. It means that 22 carat pure gold is alloyed with 2 parts of either copper or silver.

 

 

6.5       COMPOSITION, PROPERTIES AND USES OF SOME COMMON ALLOYS

The composition, properties and uses of some common alloys are given in the table below :

S.No.	Alloy	Percentage composition	Properties	Uses
1.	Steel	Iron—98.5 to 99.5% Carbon—0.1 to 1.5% Traces of other impurities	Hard, tough, strong	Construction of ships, bridges vehicles cables, machine parts, preparation of other alloy steels.
2.	Stainless steel	Iron—74% Chromium—18% Nickel—8%	Hard, Does not rust or corrode	Surgical instruments. For household utensils, shaving blades, watch cases.
3.	Alnico	Aluminium—12% Nickel—20% Cobalt—5%	Highly magnetic	For permanent magnets
4.	Magnalium	Magnesium—5% Aluminium—95%	Tough, very light does not stick to the tools	Machined articles, cheap balances
5.	Duralumin or Duralium	Magnesium—2% Aluminium—92%	Light, tough, resists corrosion	For making air crafts, pressure cooker.
6.	Brass	Copper—80% Zinc—20%	Malleable, Resistant to corrosion, good ductility, high strength	Screws, nuts and bolts, Utensils and cartridge caps
7.	Bronze	Copper—90% Tin—10%	Resistant to corrosion, hard, high tensile strength	Statues, ships, medals, coins
8.	Solder	Lead—50%, Tin—50%	Low melting point	For welding electric wires and by dentists for filling in teeth.

 

 

 

• Metals react with oxygen to form basic oxides. Aluminium oxide and zinc oxide show properties of both basic as well as acidic oxides. These oxides are know as amphoteric oxides.
• Some metals react with water and acids to liberate hydrogen gas.
• Different metals have different reactiveties with water and dilute acids.
• The arrangement of metals, in order of decreasing creativities is called reactivity or activity series of metals. The basis of reactivity of metals is the tendency to lose electrons.
• The metals above hydrogen in the activity series react with water or steam or dilute acids to liberate hydrogen gas. The metals below hydrogen in the activity series do not react with water or dilute acids.
• Most of the non-metals have properties opposite to that of metals. They are neither malleable nor ductile. They are brittle, bad conductors of heat and electricity (except for graphite) and have low melting and boiling points.
• The process of extracting metals, from their ores is called metallurgy.
• The process of removal of unwanted impurities (gangue) from the ore is called concentration or ore enrichment.
• The process of heating the concentrated ore in the absence of air is called calcination.
• The process of heating the concentrated ore in the presence of air is called roasting.
• The process of extraction of metals by electrolysis process is called electromatallurgy.
• The process of purifying the curde metal is called refining.
• The surface of some metals such as iron is corroded when they are exposed to moist air for a long time. This phenomenon is called corrosion.
• The process of covering iron with zinc to protect it from rusting is called galvanization.
• An alloy is a homogeneous mixture of two or more metals or a metal and a non-metal.

 

1. Name any one metal which reacts neither with cold water nor with hot water but reacts with heated steam to produce hydrogen gas.
2. Which one of the following metals does not react with oxygen even at high temperatures ? (i) Calcium (ii) Gold  (iii) Sodium
3. Name two metals which are both ductile as well as malleable.
4. Write the chemical equation for the reaction of hot aluminium with steam.
5. What is seen to happen when a piece of sodium metal is dropped into water ?
6. Write chemical equations to show what happens when :

(i) Manganese dioxide is heated with aluminium.

(ii) Cinnabar is roasted in a furnace.

7. Write chemical equations for reactions taking place when

(i) Cinnabar is heated in air.

(ii) Calcium metal reacts with water.

8. The oxide X₂O₃ is unaffected by water. Outline a method by which a sample of metal X can be obtained from its ore. Give one reason as to why have you chosen this method.
9. Name the alloys which are used for the following purposes :

(i) for soldering joints

(ii) for making windows and door fittings

(iii) for making aircrafts and kitchen wares

(iv) for making equipment for feed and dairy industry

10. Give reasons for the following :

(i) Metals are regarded as electropositive elements.

(ii) When a piece of Copper metal is added to a solution of zinc sulphate, no change takes place, but the blue colour of copper sulphate fades away when a piece of zinc is placed in its solution.

(iii) Articles made of aluminium do not corrode even though aluminium is an active metal.