All the materials around us are made up of chemical elements, which are found in the earth crust. Earth is the source of coal, petroleum, graphite, diamond and many other minerals of metals and non-metals. We get various useful things like gasoline, kerosene, wax, coal gas and natural gas from the natural resources, which are made up of many non-metals. These elements occur as minerals and rocks in the earth’s crust. Some of these elements like oxygen, nitrogen and carbondioxide occur in atmospheric air. There are more than 115 elements known at present 80% of these elements are metals and rest are non-metals. On the basis of their properties , all the elements can be divided into two main groups: metals and non-metals.


Metals are the elements that conduct heat and electricity and are malleable and ductile. Some of the examples  of metals are : Iron, Aluminium, Copper, Silver, Gold, Platinum, Zinc. Metals are the elements which form positive ions by losing electrons (or donating electrons). Metals are known as electropositive elements because they can form positive ions by losing electrons. The most abundant metal in the earth’s crust is aluminium.


Non-metals are the elements that does notconduct heat and electricity and are neither malleable nor ductile. They are brittle. Some of the examples of non-metals are : Carbon, Sulphur, Phosphorus, Silicon, Hydrogen, Oxygen, Nitrogen. The two allotropic forms of carbon element, diamond and graphite are also non – metal. Non-metals are the elements which form negative ions by  gaining electons. Non-metals are known as electronegative elements because they can form negative ions by gaining electrons. Carbon is one of the most important non-metals, as life on this earth is based on carbon compound because the carbon compounds like proteins, fats, carbohydrates, vitamins and enzymes etc. are essential for the growth and development of living organisms. The most abundant non-metal in the earth’s crust is oxygen.


♦ Position of Metals and Non-Metals in The Periodic Table

(i) The elements which are placed on the left hand side (except hydrogen) and in the centre of the periodic table are called metals. Such as sodium, potassium, magnesium, calcium, iron, copper zinc etc.
(ii) The elements which are placed on the right hand side of the periodic table are called non-metals such as oxygen, nitrogen, chlorine, fluorine etc. These metals and non-metals are separated from each other in the periodic table by a zig-zag line. The elements placed in the zig-zag line show some properties of metals and some properties of non-metals are called metalloids. Such as boron(B), silicon(Si), germanium(Ge), arsenic(As), antimony(Sb), tellurium(Te) and polonium(Po).
(iii) The position of metals, non-metals and metalloids are shown in a simple form in figure.. Metals present at the extreme left are known as light metals, while those are present in the centre of the periodic table are called heavy metals or transition metals.
(iv) The elements at the extreme left of the periodic table are most metallic and those on the right are least metallic or non-metallic. Thus, metallic character decreases on going from left to right side in the periodic table. For example, sodium is more metallic than aluminium because sodium is on the left hand side of aluminium.
(v) However on going down in a group the metallic character increases. For example, carbon is non-metal while lead is metal because metallic character increases down in a group.



♦ Electronic View of Metal

An element is called metal, which forms positive ions (or cations) by losing electron.
Example : Sodium is a metal which forms sodium ion (Na+) by losing one electron. Similarly, magnesium metal forms Mg2+ by losing two electrons, Al metal forms Al3+ by losing three electrons.
Thus, metals are also known as electropositive elements.
The atoms of metals have 1 to 3 electrons in their outermost shell. For example, all the alkali metals have one electron in their outermost shell. (Lithium-2, 1, sodium 2, 8, 1, potassium-2, 8, 8, 1, … etc). Sodium 11(2, 8, 1) magnesium 12 (2, 8, 2)  and aluminium 13 (2, 8, 3) are metals having 1, 2 and 3 electrons respectively in their outermost shell, which lose these electron easily. The number of electrons lost by an atom of a metal is called its valency.  Thus metals have 1 to 3 electrons in their valence shell of their atoms. Exceptions : Hydrogen and Helium. Hydrogen is a non-metal having 1 electron in its outermost shell of its atom. Helium having 2 electrons in its outermost shell of its atom.


Properties of Metals

♦ Physical Properties of Metals :

1. Metals are malleable, i.e. metals can be  beaten into thin sheets with hammer (without breaking) Malleability is an important property of metals. Gold and Silver metals are some of the best malleable metals. Aluminium foils are used for packing food items like biscuits, chocolates, medicines, cigarettes, etc.

2. Metal are ductile, that is, metals can be drawn (or stretched) into thin wires.
Ductility is another important property of metals.  Gold is the most ductile metal. For example, 1 gram of gold can be drawn into a thin wire about 2 kilometer long. Copper and aluminium metals are also very ductile and can be drawn into thin copper wires and aluminium wires.

3. Metals are good conductors of heat.

Metals allow heat to pass through them easily. Take a flat aluminium  rod and stick some nails upon the rod with the help of wax. Start heating the free  end of the aluminium rod by keeping a burner  below it. We will see that the iron nails attached to aluminium rod with wax start falling one by one because heat travels from the left side to the right side along the aluminium rod. It melts the wax which holds the nails. Silver metal is the best conductor of heat. The cooking utensils and water boilers, etc., are usually made up of copper or aluminium metals because they are very good conductors of heat. Heat conductivity is an important property of metals.

4. Metals are good conductors of electricity  
Metals allow electricity (or electric current) to pass through them easily. Silver metal is the best conductor of electricity. The electric wires are made of copper and aluminium metals because they are very good conductors of electricity.

5. Metals are lustrous (or shiny) and can be polished
Metals are lustrous, they have a shining surface. For example gold, silver and copper are shiny metals and they can be polished. The property of a metal having a shining surface is called ‘metallic lustre’. The metals  lose their shine or brightness by keeping in air for a long time and acquire a dull appearance due to the formation of a thin layer of oxide, carbonate or sulphide on their surface (by the slow action of the various gases present in air).

6. Metals are generally hard (except sodium and potassium which are soft metals).
Most of the metals like iron, copper, aluminium, etc. are very hard. Some exceptions Sodium and potassium are soft metals which can be easily cut with a knife.

7. Metals are strong (except sodium and potassium metals which are not strong).
They can hold large weights without snapping (without breaking). For example iron metal (in the form of steel) is very strong. Due to this iron metal is used in the construction of bridges, buildings, railway  lines, machines, vehicles and chains etc.

8. Metals are solid at room temperature (except mercury which is a liquid metal). 

9. Metals have high melting points and boiling points (except sodium and potassium metals which have low melting and boiling points) Example, iron metal has a high melting point of 1535°C. Copper metal has also a high melting point of 1083°C. Sodium and potassium metals have low melting points (of 98°C and 64°C respectively).

10. Metals have high densities (except sodium and potassium metals which have low densitites) The density of iron is 7.8 g/cm3 which is quite high. Sodium and potassium metals have low densities (of 0.97 g/cm3 and 0.86 g/cm3 respectively)

11. Metals are sonorous. That is metals make sound when hit with an object.
The property of metals of being sonorous is called sonorousness or sonority. It is due to the property of sonorousness (or sonority) that metals are used for making bells and strings (wires) of musical instruments like sitar and violin.

12. Metals usually have a silver or grey colour (except copper and gold)



♦ Uses of some metals

(i) Many metals and their compounds are useful in our daily life. These are as follows : Aluminium is used to prepare utensils and house hold equipments like vacuum cleaner. Aluminium is extensively used in making bodies of rail, cars, automobiles, trucks and aircraft. Aluminium wires are widely used in electrical work. Aluminium foil is used to wrap chocolate cigarette and medicines and to seal milk bottles.
(ii) Major use of copper is in making electrical wires & cables. Copper is also used in making utensils, steam pipes, coin and in electroplating.
(iii) Steel is an alloy of iron which is used for making parts of machines, as building material and in the construction of refrigerator. As a matter of fact steel is said to be the back bone of industry.
(iv) Gold and silver called noble metals (or coinage metals) are used in jewellery.
(v) Mercury is used in thermometers barometers and to prepare amalgams.
(vi) Platinum is used to make crucibles and electrodes.
(vii) Zinc is used to galvanize iron, to prepare roofing material, container of dry cells and to make brass when mixed with copper.
(viii) Metal like sodium, titanium and zirconium find their applications in atomic energy, research and medical industry.
(ix) Titanium (Ti) and its alloys are used in aerospace, marine equipments, aircraft frames, chemical industries and chemical reactors. The wide application of titanium is attributed to its resistance to corrosion, high melting points and high strength.


♦ Chemical  properties of Metals:

1. Reaction of Metals with oxygen (Air)
When metals are burnt in air, they react with the oxygen of air to form metal oxides:
Metal + Oxygen     →   Metal oxide
From air                               (Basic oxide)
Metals react with oxygen to form metal oxides. Metal oxides are basic in nature. The vigour of reaction with oxygen depends on the chemical reactivity of metal.

(i) Sodium metal reacts with the oxygen at room temperature to form a basic oxide called sodium oxide: 4Na(s)   + O2(g)         →       2Na2O(s)
sodium       oxygen            sodium oxide
(Metal)      (from air)       (Basic oxide)

Potassium metal and sodium metal are stored under kerosene oil to prevent their reaction wilh the  oxygen, moisture and carbon dioxide. Some of the metal oxides dissolve in water to form alkalies.
Eg. Na2O(s)                 +             H2O(l)        →       2NaOH(aq)
sodiumoxide                              water                sodium hydroxide
(basic oxide)                               (An alkali)

(ii) Magnesium metal does notreact with oxygen at room temperature. But on heating, magnesium metal burns in air giving instense heat and light to form a basic oxide called magnesium oxide (which is a white powder)
2Mg(s)                    +         O2(g)       →          2MgO(s)
Magnesium                    Oxygen          Magnesium oxide
(Metal)                           (From air)                (Basic oxide)
Magnesium oxide dissolves in water partially  to form magnesium hydroxide solution:
MgO(s)                +            H2O(l)           →     Mg (OH)2(aq)
Magnesium oxide            water               Magnesium hydroxide

(iii) Aluminium metal burns in air on heating to form aluminium oxide:
4Al                   +          3O2          →        3Al2O3 (s)
Aluminium           Oxygen                 Aluminium oxide
(Metal)                  (From air)           (Amphoteric oxide)


Those metal oxides which show basic as well as acidic behaviour are known as amphoteric oxides. Aluminium metal and zinc metal form amphoteric oxides. Amphoteric oxides react with both, acids as well as bases to form salts and water.
Example :
(a) Al2O3(s)                +                6HCl        →            2AlCl3(aq)             +           3H2O(l)        
Aluminium oxide      Hydrochloric acid         Aluminium chloride                Water
(Base)                                                (Acid)                                    (Salt)

(b) Al2O3(s)               +                 2NaOH          →      2NaAlO2(aq)       +              H2O(l)    
Aluminium oxide         Sodium hydroxide        Sodium aluminiate               Water
(Acid)                                                 (Base)                                  (Salt)

(iv) Zinc metal burns in air only on strong heating to form zinc oxide:
2 Zn(s)                       +                  O2(g)               →            2 ZnO(aq)     
Zinc                                               Oxygen                              Zinc oxide
(Acid)                                                                                (Amphoteric oxide)

Zinc oxide reacts with hydrochloric acid to form zinc chloride (salt) and water.
ZnO(s)                   +                 2HCl(aq)            →             ZnCl2(aq)            +            H2O(I) 
Zinc oxide                          Hydrochloric acid              Zinc chloride                      Water
(Base)                                           (Acid)                                         (Salt)

(v*) Iron metal does notburn in air even on strong heating. Iron reacts with the oxygen on heating to form iron (II, III) oxide:
3Fe(s)                    +       2O2(g)              →                    Fe3O4(s)
Iron                                  Oxygen                                Iron (II, III) oxide

(vi*) Copper metal also does notburn in air even on strong heating. Copper reacts with the oxygen on prolonged heating to form a black substance copper (II) oxide:
2Cu (s)                +            O2                →                          2CuO (s)   
Copper                            Oxygen                                       Copper (II) oxide



♦ Nature of metallic oxide

Generally, metallic oxides are basic in nature except aluminium and zinc oxides which are amphoteric in nature. This means these oxides (Al2O3, ZnO) react with base as well as acid. The basic oxide of metals react with acid to give salt.

For example :
CuO                           +                 H2SO4            →        CuSO4                +              H2O
Copper(II)oxide                Sulphuric acid            Copper(II) sulphate        Water
Some oxide of metals dissolve in water and form alkalis.

Example for :
Na2O(s)                +                   H2O(l)            →             2NaOH(aq)
Sodium hydroxide
K2O(s)                   +                   H2O(l)          →                2KOH (aq)
Potassium hydroxide

Reaction showing amphoteric in nature of Al2O3 and ZnO.

Al2O3(s)             +                    6HCl(aq)       →        2AlCl3(aq)        +          3H2O(l)
Hydrochloric acid     Aluminium chloride

Al2O3(s)              +                     2NaOH(aq)      →      2NaAlO2(aq)   +       H2O(l)
Sodium hydroxide(base) Sodium meta aluminate

ZnO(s)               +                   2HCl(aq)          →       2ZnCl2(aq)            +         H2O(l)
Hydrochloric acid            Zinc-chloride

ZnO(s)              +                   2NaOH(aq)      →      Na2ZnO2(aq)       +        H2O(l)
Sodium hydroxide          Sodium Zincate



2. Reaction of Metals with water 

Metals react with water to form a metal hydroxide (or metal oxide) and hydrogen gas.

(i) Potassium react violently with  cold water to form potassium hydroxide and hydrogen gas:
2K(s)                    +           2H2O              →           2KOH(aq)                  +               H2(g)           +               Heat
Potassium             Water                                   Potassium hydroxide                  Hydrogen

(ii) Sodium reacts vigorously with cold water forming sodium hydroxide and hydrogen gas:
2Na(s)                 +             2H2O(I)           →      2NaOH(aq)                 +                H2(g)          +             Heat
sodium                                 water                      sodium hydroxide                      hydrogen

(iii) Calcium reacts with cold water to form calcium hydroxide and hydrogen gas:
Ca(s)                   +                2H2O(I)           →          Ca(OH)2(aq)         +                   H2(g)
Calcium                            water(cold)                     Calcium hydroxide                  Hydrogen

The piece of calcium metal starts floating in water because the bubbles of hydrogen gas formed during the reaction stick to its surface.

(iv) Magnesium metal does not react with cold water. Magnesium reacts with hot water to form magnesium hydroxide and hydrogen:
Mg (s)                     +                2H2O (l)              →            Mg(OH)2(aq)                   +                H2(g)
magnesium                            water(hot)                         Magnesium hydroxide                   Hydrogen

In this reaction the piece of magnesium metal starts floating in water due to the bubbles of hydrogen gas sticking to its surface.  Magnesium reacts very rapidly with steam to form magnesium oxide and hydrogen :  Mg (s)                  +                    H2O(g)             →                 MgO(s)                              +                  H2(g)
Magnesium                             Steam                               Magnesium oxide                              Hydrogen

(v) Aluminium reacts with steam to form aluminium oxide and hydrogen gas :
2Al (s)                +                  3H2O(g)             →                   AI2O3(s)                        +                     3H2(g) 
Aluminium                            Steam                                 Aluminium oxide                                Hydrogen
Aluminium metal does notreact with water under ordinary conditions because of the presence of a thin (but tough) layer of aluminium oxide on its surface.

(vi) Zinc reacts with steam to form zinc oxide and hydrogen:
Zn (s)                  +                        H2O (g)             →               ZnO (s)                           +                      H2 (g)
Zinc                                                    Steam                               Zinc oxide                                            Hydrogen

(vii) Red – hot iron reacts with steam to form iron (II, III) oxide and hydrogen :
3Fe(s)               +                      4H2O(g)              →              Fe3O4(s)                       +                      4H2(g)
Iron                                                 Steam                                 Iron (II, III) oxide                                Hydrogen
Metal like lead, copper, silver and gold does not react with water (or even steam).


3. Reaction of metals with Dilute Acids: 
Metals usually displace hydrogen from dilute acids. When a metal reacts with a dilute acid, then a metal salt and hydrogen gas are formed
Metal                       +                  Dilute acid              →            Metal salt          →                Hydrogen

(i) Sodium metal reacts violently with dilute hydrochloric acid to form sodium chloride and hydrogen:
2Na(s)                     +                    2HCl(aq)                →                 2NaCl(aq)         +                  H2(g)
Sodium                                  Hydrochloric                                Sodium chloride                        Hydrogen

(ii) Magnesium reacts quite rapidly with dilute hydrochloric acid forming magnesium chloride and hydrogen gas :
Mg(s)                   +                      2HCl(aq)                  →                    MgCl2(aq)           +                H2(g)
Magnesium                   Hydrochloric acid                              Magnesium chloride               Hydrogen

(iii) Aluminium metal at first reacts slowly with dilute hydrochloric acid due to the presence of a tough protective layer of aluminium oxide on its surface.But when the thin, outer oxide layer gets dissolved in acid. Aluminium metal reacts rapidly with dilute hydrochloric acid to form aluminium chloride and hydrogen gas :  2AI(s)                     +                    6HCI(aq)                     →                    2AlCl3(aq)                  +               3H2(g) 
Aluminium                       Hydrochloric acid                                Aluminium chloride                     Hydrogen
The reaction of aluminium with dilute hydrochloric acid is less rapid than that of magnesium, so aluminium is less reactive than magnesium.

(iv) Zinc reacts with dilute hydrochloric acid to give zinc chloride and hydrogen gas(but the reaction is less rapid than that of aluminium)
Zn(s)                     +                        2HCl(aq)                    →                         ZnCl2(aq)                 +                H2(g)
Zinc                                            Hydrochloric acid                                 Zinc chloride                               Hydrogen
This reaction shows that zinc is less  reactive than aluminium.

(v) Iron reacts slowly with cold dilute hydrochloric acid to give iron (II) chloride and hydrogen gas:
Fe(s)                      +                        2HCl(aq)                    →                            FeCl2                        +                     H2(g)
Iron                                            Hydrochloric acid                                Iron (II) chloride                              Hydrogen

(vi) Copper does notreact with dilute hydrochloric acid (or dilute sulphric acid) at all. This shows  that copper is even less reactive than iron:
Cu (s)                    +                             HCl (aq)                   →                            No reaction
Copper                                        Hydrochloric acid
Metals like copper and silver which are less reactive than hydrogen, does notdisplace hydrogen from dilute acids.


♦ Reaction of metals with solutions of other metal salts
When a more reactive metal is placed in a salt solution of less reactive metal, then the more reactive metal displaces the less reactive metal from its salt solution. This reaction is also known as displacement reaction. Let us learn it with the help of following activity.

Observation: The blue colour of copper sulphate has faded and becomes greenish. The green colour of the solution is due to the formation of iron (II) sulphate and copper is displaced. A reddish-brown coating is formed on the surface of iron nail. The reaction is represented by the chemical equation.
Fe(s)                        +                    CuSO4(aq)                     →                   FeSO4(aq)           +          Cu(s)
Iron                                     Copper sulphate solution                  Ferrous sulphate

But the greenish colour of FeSO4 does notchange. That means no reaction take place.
Conclusion : These activities shows that iron metal is more reactive than copper. Similarly,

♦ Reaction of copper with silver nitrate solution :
When a strip of copper metal is placed in a solution of AgNO3. The solution becomes gradually blue and a shining coating of silver metal gets deposited on the copper strip. The reaction may be written as:
2AgNO3(aq)                            +                       Cu(s)                    →            Cu(NO3)        +                  2Ag
Silver nitrate                                                                                                 Copper nitrate                             Silver
(colourless solution)                                                                                      (blue colour)

However, if we place silver wire in a copper sulphate solution no reaction occurs. This means copper can displace silver from its salt solution but silver cannot displace copper from its solution. i.e. copper is more reactive metal than silver.




♦ Non-metals and their general properties
Non-metals are present on the right hand side of the periodic table (exception : Hydrogen). Among the total known elements, there are only 22 non-metals, out of which 11 are gases like oxygen, nitrogen, hydrogen  one is a liquid (Bromine) and the rest 10 are solids such as sulphur, phosphorus and the allotrops of carbon (Diamond and graphite).

♦ Electronic view of non-metals
An elements is called non-metal which form ions by gaining electrons. For example, oxygen is a non-metal which form O2– ions by gaining two electrons. Similarly, nitrogen form N3– ions by gaining three electrons. Thus, non-metals also known as electronegative elements. The atoms of non-metals have usually 4 to 8 electrons in their outer most shell. For example,Carbon (At No. 6), Nitrogen (At. No. 7), Oxygen (At. No. 8), Fluorine (At. No. 9) and Neon (At. No. 10), have respectively 4, 5, 6, 7 and 8 electrons in their outermost shell. However, there are two exceptions namely hydrogen and helium which have one and two electrons in their valence shell or outer most shell, but they are non-metals.



1. Non-metals are neither malleable nor ductile. Non-metals are brittle (break easily).  Solid non-metals can neither be hammered into thin sheets nor drawn into thin wires. For example,  sulphur and phosphorus are solid non-metals which are non-malleable and non-ductile. The property of being brittle (breaking easily) is called brittleness. Brittleness is an important property of non-metals.

2. Non-metals does not conduct heat and electricity.
Non-metals does notconduct heat and electricity because unlike metals, they have no free electrons (which are necessary to conduct heat and electricity). For example, sulphur and phosphorus are non-metals which does notconduct heat and electricity. There is, however one exception, carbon (in the form of graphite) is  the only non-metal which is a good conductor of electricity because of it’s structure.

3. Non-metals are not lustrous (not shiny).
They are dull. Non-metals does nothave a shining surface. For example, sulphur and phosphorus are non-metals which have non lustre. Iodine is a non-metal having lustrous appearance.

4. Non-metals are generally soft (except diamond which is  extremely hard non-metal)

5. Non-metals are not strong. They are easily broken.

6. Non-metals may be solid, liquid or gases at the room temperature.

7. Non-metals have comparatively low melting points and boiling points (except diamond which is a non-metal having a high melting point and boiling point).
The melting point of sulphur is 115°C which is quite low. The melting point of diamond is, however more than 3500°C which is very high.

8. Non-metals have low densities,  that is, non-metals are light substances. 
The density of sulphur of 2g/cm3.

9. Non-metals are non-sonorous. They does not produce sound when hit with an object.

10. Non-metals have many different colours.
On the basis of the above discussion of the physical properties of metals and non-metals, we have concluded that elements can not be grouped according to the physical properties alone, as there are many exceptions.


For example,
(i) All metals except mercury are solids at room temperature. We know that metals have very high melting points but gallium (Ga) and caesium (Cs) have very low melting points. These two metals will melt if we keep them at our palm.
(ii) Iodine is a non-metal but it is lustrous.
(iii) Alkali metals such as Lithium, Sodium and Potassium are soft and they can be easily cut with a knife. i.e. they have very low densities and low melting points.
(iv) Carbon is a non-metal that can exist in different forms. Each form is called an allotrope of Diamond, an allotrope of carbon is the hardest natural substance. which has very high melting and boiling point. Graphite is another allotrope of carbon which is good conductor of electricity. The elements can be more clearly classified as metals and non-metals on the basis of their chemical properties.

The reactivity series

The arrangement of metals in order of decreasing reactivities is called reactivity series or activity series of metals. After performing displacement experiments the following series has been developed.


♦ Characteristics of Reactivity Series:
(i) The most reactive metal is placed at the top and the least reactive metal is placed at the bottom of the table.
(ii) Metals present above the hydrogen in reactivity series can displace hydrogen from dilute acids.
(iii) A metal can displace the metals placed below it in the reactivity series.
(iv) Metals present at the top are more elecro-positive, so they will occur in combined or compound form only in nature.
(v) Metals at the bottom are less reactive and do not react easily so they may be present in free state in nature.
Ex.1 A, B and C are three elements which undergo chmical change according to the following equations: A2O3 + 2B → B2O3 + 2A
3CSO4 + 2B → B2(SO4)3 + 3C
3CO + 2A → A2O3 + 3C
Write the anme of the most reactive and the least reactive elements.
(i) In the first reaction, B displaces A, so B is more reactive than A.
(ii) In second reaction, B displaces C, so B is more reactive than C.
(iii) In third reaction, A displaces C, so A is more reactive than C. So, B is more reactive than A and C and A is more reactive than C, So the order of their reactivities is as follows:
B > A > C


Ex.2 Explain why zinc metal can displace copper from copper sulphate solution but copper cannot displace zinc from zinc sulphate solution.
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.

When a piece of zinc is placed in a solution of copper sulphate, zinc being more reactive than copper, can displace copper from its salt solution and forms zinc sulphate and blue colour of copper sulphate fades away slowly, but when a piece of copper sulphate fades away slowly, but when a piece of copper metal is added to a solution of zinc sulphate, no change takes place as copper being less reactive than zinc, cannot displace zinc from zinc sulphate.


Electronic configuration of some elements:
Types of elementElementAtomic
Number of electrons in shells
Noble gasesHelium (He)
Neon (Ne)
Argon (Ar)
MetalsSodium (Na)
Magnesium (Mg)
Aluminium (Al)
Potassium (K)
Calcium (Ca)
Non-MetalsNitrogen (N)
Oxygen (O)
Fluorine (F)
Phosphorus (P)
Sulphur (S)
Chlorine (Cl)


It is clear from the above table that except helium, all other noble gases have 8 electrons (octet) in their outermost shell. Which represent a highly stable electronic configuration. Due to this stable configuration, the noble gases have no any tendency to lose or gain electrons. So they exist monoatomic, sodium atom has one electron in its outermost shell. If it loses the electon from its M shell the its L shell becomes the outermost shell. which has stable octet like noble gases. The nucleus of this atom still has 11 protons but the number of electrons has 10. Therefore, if becomes positively charged sodium ion or cation (Na+).

Na     \buildrel {lose1.electron} \over\longrightarrow   Na+ + e-

2, 8, 1                                      2,8
Sodium cation

On the other hand chlorine has seven electrons in its outer most shell and it require one more electron to complete its octet. The nucleus of chlorine atom has 17 protons and the number of electrons become 18. This makes chloride ion, Cl as negatively charged

CI         \buildrel {gain1.electron} \over\longrightarrow           CI
2,8,7                                                      2,8,8
                                                                            Chloride ion


So, Na+ and Cl ions being oppositely charged atoms which attract each other and are held by strong electrostatic forces of attraction to exist as NaCl. In other words, Na+ and Cl– ions are held together by electrovalent or ionic bond.

The formation of one more ionic compound magnesium chloride :The formation of one more ionic compound magnesium chloride : The electronic configuration of magnesium (Mg) and chlorine atoms are:
Mg12 : 2, 8, 7
Cl17 : 2, 8, 7

Magnesium atom has two electrons in its valence shell. It has a tendency to lose both of its electrons to attain the nearest noble gas configuration (i.e. Ne). Mg      →      Mg2+.  On the other hand, chlorine has only one electron less than the nearest noble gas (i.e. Ar) configuration. The magnesium loses its both the valence electrons to two chlorine atoms, each of which is need of one electron to form Cl ion.

The compounds formed by the transfer of electrons from a metal to a non-metal are known as ionic compound or electrovalent compounds. The structure of some common ionic compounds are given below :

Structure of some common ionic compounds:

1. Magnesium      Mg                 +                 O                    →                   Mg2+[O]2–           or       MgO
oxide                       2, 8, 2                               2, 6

2. Magnesium      Mg                +                2F                  →                    Mg2+2  [F]          or      MgF2
fluoride                  2, 8, 2                               2, 7

3. Calcium               Ca                 +               O                   →                     Ca2+[O]2–             or         CaO
oxide                 2, 8, 8, 2                               2, 6

4. Aluminium        Al                 +                O                  →                         2A3+3[O]2–        or      Al2O3
oxide                     2, 8, 3                               2, 6

5. Magnesium    Mg                +                2Cl               →                            Mg2+2[Cl]       or      MgCl2
Chloride           2, 8, 2                                 2, 8, 7

6. Aluminium       Al               +                  N                    →                             Al3+N3–            or          AlN
nitride       2, 8, 3                              2, 5



  • Following are the general properties of ionic compounds.
    (a) Physical state
    Ionic compounds are solids and relatively hard because of the strong force of attraction between the positive and negative ions. This force of attraction is also known as strong electrostatic force of attraction. These compounds are generally brittle and break into pieces when pressure is applied.
    (b) Solubility
    Electrovalent compounds are generally soluble in water (because of their polar nature) and insoluble in solvents such as kerosene, petrol, etc.
    (c) Melting and boiling points
    Ionic compounds have high melting and boiling points, due to the strong electrostatic force of attraction
    belween the oppositely charged ions. Therefore, large amount of energy is needed to break these bonds.
    (d) Conduction of electricity Ionic compounds in the solid state
    do not conduct electricity because movement of ions in the solid state is not possible due to their rigid structure. But they can conduct electricity in molten or aqueous state. (e) Colour to the flame
    Most of the salts when brought into the flame, impart characteristic colour to the flame.

Hydrogen gas is not evolved when metals such as Zn, Fe, Cu and Al reacts with nitric acid. Because HNO3 is  strong oxidising agent. It oxidises H2 gas to water and itself gets reduced to form oxides of (NO, N2O and NO2) nitrogen.

3Fe(s)                +                  8HNO3(aq)          →                3Fe(NO3)2(aq)       +       4H2O(l)             +     2NO(g)
Iron                                                Nitric acid (dil)                     Iron(II) nitrate                Water                      Nitric oxide

3Cu(s)              +                  8HNO3(aq)         →                 3Cu(NO3)2(aq)      +        4H2O(l)          +     2NO(g)
Copper                                  Nitric acid                                  Copper nitrate                  Water                       Nitric oxide

But copper reacts with hot concentrated sulphuric acid (H2SO4) to produce copper sulphate, sulphur dioxide and water.

Cu(s)               +                     2H2SO4               →                           CuSO4                    +            2SO2           +        2H2O
Mg reacts with very dilute HNO3 to evolve H2 gas.

Mg(s)             +                     2HNO3(aq)        →                           Mg(NO3)2(aq)    +                H2O(g)
Magnesium                     Nitric acid (dil)                             Magnesium nitrate

Fe react with dil H2SO4 to evolve H2

Fe(s)              +                dil H2SO4(aq)        →                           FeSO4(s)                  +                H2(g)
Iron                                  Sulphuric acid                                    Ferrous sulphate


♦ Aqua Regia (ROYAL WATER)
Aqua regia is a Latin word it means ” royal water”. It is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3 : 1. It is a highly corrosive, fuming liquid and it is used to dissolve gold and platinium.


Occurence of metals

The main source of metal is earth’s crust. Some metals also occur in sea water. The metals are found in nature in :

(1) Native state (or free state): Only a few less reactive metals like silver, gold platinum etc., are found in the free state in which they are called “native metals”.

(2) Combine state: i.e., in the from of their compounds admixed invariably with various useless impurities such as clay, sand, rocky material, etc. Usually, metals are found in the form of oxides, sulphides, carbonates, phosphates, halides silicates, etc.
(i) The naturally occurring form of metal in combined state, is known as “mineral”.
(ii) Those naturally occurring minerals, which are economically suitable for commercial extraction of metals, are known as ‘ores’. Thus, every ore is a mineral, but every mineral is not an ore.
(iii) The rocky and earthy impurities (like clay, sand) generally associated with ore, are called gangue (or matrix).

Note :
(1) Sodium a very reactive metal, and reacts readily with moisture, oxygen and carbon dioxide of air. So sodium cannot exit ‘free’ in nature. Hence, it is not found ‘native’ in nature.
(2) Sodium is highly reactive metal, and has affinity for oxygen. If it is exposed to air, a coating of the oxide is formed and sometimes, it may even  catch fire. Consequently, sodium metal should not be exposed to air. Hence, sodium is stored under kerosene.
(3) Aluminium is a reactive metal, so it is not found in free state in nature. It occurs in the form of its compounds, chief of which is bauxite (Al2O3 2H2O).
(4) Gold and silver occupr low position in the activity series. Consequently, they are least reactive elements and are not effected by most chemicals, atmospheric oxygen, moisture, carbon dioxide etc. Hence, they  often occur in free or native state in nature.


Extraction of metals: We have learnt about the reactivity series of metals, according to which, the  metals at the “bottom” of the reactivity series are the “least reactive” and these are often found in a free-state, e.g., Au, Ag, Pb and Cu. However Cu and Ag are also found in combined state as their oxides and sulphides. On the other hand, metals at the “top” of the reactivity series are so reactive, they are never found in nature as free elements, e.g., Li, K, Na, Ca, Mg etc. The metals in the “middle” of the reactivity series (e.g., Al, Zn, Fe, Pb etc.) are moderately reactive and they are found in the earth’s crust mainly as  oxides, sulphides or carbonates [e.g., Al2O3. 2H2O (bauxite), HgS (cinnaber), ZnCO3 (calamine)].

On the basis of reactivity seires, we can  have following three groups of elements:
(i) Metals of low reactivity.
(ii) Metals of medium rectivity.
(iii) Metals of high reactivity.

Metallurgy: is the process of extracting a metal in the free form from its ore and then refining it for use. Various steps involved in the extraction of metals from their ores are generally as follows:
(a) Concentration (or enrichment) of ore
(b) Conversion of concentrated ore into oxide
(c) Reduction of oxide ore into impure metal
(d) Refining of impure metal.

(a) Con centration (or enrichment) of ore: The ore is, generally, associated with useless rocky and earthy impurities (like clay, sand etc.), called ‘gangue’ or matrix. The ‘concentration’ (or enrichments) of ore means removal of gangue from the powdered ore. Thus, the percentage of the metal in the concentrated ore is higher than that in the original ore. The concentration of ore can be brought about in the following ways. depending upon the type of ore such as hydraulic washing, froth floatation method, magnetic separation etc.

(i) Levigation or gravity separation or hydraulic washing
This method is based upon the difference in the densities of the ore particles and impurities (gangue). Example: Haemetite ore of iron.

(ii) Froth floatation
This method is based on the difference in the wetting properties of the ore and gangue particles with water and oil. It is used for enrichment of sulphide ores. Example: ZnS, HgS.

(iii) Liquation
This method is based on difference in melting point of ore and gangue particles. Example: ore of tin and zinc.

(iv) Magnetic separation
This method is based on difference in the magnetic properties of the ore and gangue. Example: magnetite (Fe3O4) ore of iron.

(v) Chemical separation
When none of the physical propertry makes the difference, then we use chemical properties as the basis for enrichment. e.g. Bayer’s process for alumina enrichment.

Next steps of metallurgy depend on the type of metal to be extracted:

(a) Extracting metals low in reactivity series: Since these metals are unreactive, so the oxides of these metals can be “reduced” by heating alone. For example, cinnabar (HgS) an ore of mercury changes to mercury on heating

(b) Extracting metals in the middle of the reactivity series: Since these metals (e.g. Fe, Zn, Pb, Cu, etc.) are moderately reactive, so they are usually found in earth’s crust as sulphides or carbonates.Consequently are converted into metal oxides.

(i) The process of conversion of metal sulphide to oxide by strongly heating in the presence of excess air, is called roasting. For example:
2ZnS(g) + 3O2 (g)         \mathrel{\mathop{\kern0pt\longrightarrow}\limits_{ExcessAir}^\Delta }        2ZnO(s) + 2SO2(g)
Zinc blende

(ii) The proces of conversion of metal carbonate to oxide by heating strongly in limited air, is called calcination. For example:

Reduction of oxide to metal: The metal oxides obtained above are reduced by hearting with suitable reducing agents like carbon. For example:
ZnO(s)       +       c(S)        \buildrel \Delta \over\longrightarrow                Zn(s)       +             CO(g)

It may be pointed out here that besides using carbon (coke), to reduce metal oxides to metals, sometimes, displacement reactions are also employed. The highly reactive metals (e.g., Na, Ca, Al, etc.) are employed as reducing agents, since they displace metals of lower reactivity from their compounds, For example:                          2MnO2(s)    +     4Al(s)          →       3Mn(l)         +           2Al2O3(s)                 +                   Heat

Such a displacement reaction is highly exothermic (i.e., lot of heat is evolved), so the metal produced is in molten state [e.g., Mn(l)] Al is also used to reduce iron (III) oxide (Fe2O3) and this reaction is called thermite reaction and used to join railway trackes or machine parts.
Fe2O3(s) + 2Al(s)        →        2Fe(s) + Al2O3(s)  + Heat


(c) Extracting metals near the top of the reactivity series: Since these metals are highly reactive, so their oxides cannot be reduced by heating with carbon. For example, Na2O(s), MgO(s), CaO(s), Al2O3(s), etc. cannot be reduced by heating with carbon. This is because these metals possess more affinity for oxygen than carbon. Consequently, these metals are extracted by electrolytic reduction process. For example, when molten sodium chloride is electrolyed sodium is obtained at the cathode (the negatively charged electrode) : while chlorine is liberated at the anode (the positively charged electrode).
Thus :
At cathode : Na+ + e    →     Na(s)
At anode : 2Cl    →     Cl2 + 2e
Likewise, Al is obtained by the electrolytic reduction of Al2O3.


(d) Refining of metals: the process of purifying the crude metal to get pure metal, is called refining. The method of metal refining depends on:
(i) the nature of the metal to be purified and
(ii) the type of impurities present.

Electrolytic refining: Most of the metals are refined by this method. In this process, a large block of impure metal is made the anode in an electrolytic cell, and a thin sheet of pure metal is made the cathode. Suitable metal salt solution is made as an electrolyte. On passing electric current, pure metal deposits on the cathode sheet; while some of impurities are left in solution, and other noble metal impurities settle below the anode as ‘anode mud’. For exmple, during the electrolytic refining of a copper, a thick block of impure copper is made anode, and thin plate of pure copper is made cathode. Copper sulphate solution. is used as an electrolyte. On passing electric current, following reactions take place:

(1) Cu2+ ions (from copper sulphate solution) go to the cathode (negative electrode), where they are reduced to copper, which gets deposited on the cathode.

Cu2+(aq) +  2e-     →      Cu(s)
(From solution)               (Deposite on cathode)
(At cathode)

(2) Copper (of impure anode) forms copper ions and these go into solution of electrolyte.

Cu(s)                  →                   Cu2+(aq)
(From anode                        (Goes into solution)
(At anode)

Thus, the net result is transfer of pure copper from anode to the cathode. Impurities like zinc, iron etc., go into solution; while noble impurities like silver, gold etc., are left behind as anode mud.


Any process of deterioration (or destruction) and consequent loss of a solid metallic material, through an unwanted (or unintentional) attack by its environment, starting at its surface, is called corrosion. Thus, corrosion is a proces “reverse of extraction of metals”. The most familiar example of corrosion is rusting of iron, when exposed to the atmospheric conditions. During this, a layer of reddish scale and powder of oxide (Fe2O3 . x H3O) is formed and the iron becomes weak. Another common example is formation of green films of basic copper, when exposed to moist-air containing carbon dioxide. Similarly, silver article turns black after some time, when exposed to air. This is due to the reaction of Ag with H2S present in air to form black coloured Ag2S.
Note :
(i) It may be pointed out that noble metals such as gold and platinum do not corrode easily.
(ii) The process of corrosion is continuous and causes decrease in strength of the metal.

Prevention of rusting:
(i) By painting: The corrosion of a metal can be prevented simply by painting the metal surface by grease or varnish taht forms a protective layer on the surface of the metal which protect the metal from moisture and air.
(ii) Self prevention: Some metals form protective layers.
For example: When zinc is left exposed to the atmosphere, it combines with the oxygen of air to form a layer of zinc oxide over its surface. The oxides layer does not allow iar to go inside the metal. Thus, zinc is protected from corrosion by its own protective layer. Similarly, aluminium combines with oxygen to form a dull layer of aluminium oxide on its surface which protect the aluminium from further corrosion.

(iii) Cathodic protection: In this method the more reactive metal which is more corrosion-prone is connected to a bar of another metal which is less reactive and to be protected. In this process electron flow from the more reactive metal to the less reactive metal. The metal to be protected becomes the cathode and the mroe reactive metal becomes the anode. In this way, the two metals form an electrochemical cell and oxidation of the metal is prevented.
Example: The pipelines (iron) under the surface of the earth are protected from corrosion by connecting them to a more reactive metal (magnesium or Zn) which buried in the earth and connected to the pipelines by a wire.

(vi) Oiling and greasing: Both protect the surface of metal against moisture and chemicals etc. In addition the oil and grease prevent the surface from getting scratched.

(v) Electroplating: It is a very common and effective method to check corrosion. The surface of metal is coated with chromium, nickel or aluminium etc. by electrolysis also called electroplating. They are quite resistant to the attack by both air and water and check corrosion. If the surface of metal is electroplated by zinc, it is known as galvanisation and in case tin metal is used, then the process is called tinning.

(vi) By alloying: It is a very good method of improving the properties of a metal.
For example: Iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft and stretcheds easily when hot. But, it it is mixed with a small amount of carbon (about 0.05%) it becomes hard and strong. When iron is mixed with nickel and chromium to form stainless steel which is hard and does not rust, i.e. its properties change. In fact, the properties of any metal can be changed, if it is mixed with some other subtances.

Importance of corrosion: Sometimes corrosion of a metal prevents further corrosion of the underlying metal : For example, when Al is exposed to air a thin coating of Al2O3 on the metal article is formed. This film, quite adhering and non-porous, thereby it protect the Al metal underneath from further corrosion and damage. This is the reason why Al, being a very reactive metal, is used for making uternsils.



“An alloy is a homogeneous solid solution of one metal with one or more metals or non-metals.” such as brass, bronze, steel etc.
Purposes of alloy making : Alloys are generally, made to serve one or more of the following purposes:
(i) To modify chemical activity such as increased resistance to corrosion.
(ii) To harden a metal e.g., copper in gold ornaments.
(iii) To increase the strength and toughness.
(iv) To lower the melting point.
(v) To produce good castings.
For instance, pure iron is very soft and stretches easily, but it is mixed with some metals and non-metals, the alloys formed show considerable improvement in the qualities.

(i) Steel: When iron has carbon (0.05 to 0.5%) it is called steel. It is hard and strong. It is used for making ships, vehicles and building.

(ii) Stainless Steel : When steel is mixed with nicked and chromium, it is called stainless steel. It is hard and rust-proof. It is used for making utensils, equipments for feed and dairy industry.

Some common Alloys
(i) Brass : It is an alloy of copper and zinc (Cu-60 to 90%; Zn-10 to 40%). It is a yellow coloured alloy and used for making utensils, coins and decorative pieces.

(ii) Bronze : It is an alloy of copper and tin (Cu-88 to 96%; tin-4 to 12%). It is shining light, yellowish coloured alloy. It is used for making statures, ships and medals.

(iii) Solder: It is an alloy of lead and tin (lead 33%; tin 67%). Its melting point is low. It is used for soldering electrical wires.

(iv) Alloying of gold: The purity of gold is expressed in ‘carat’ and 24 carat gold is supposed to be 100% pure. Pure gold or 24 carat gold tis very soft and cannot be sued for making ornaments. To make is hard, it is alloyed with silver, copper or both. Mosdy 22 carat or 20 carat gold is used for making ornaments. 22 carat gold means 22 parts of pure gold mixed with 2 parts of silver or copper or both.

(v) Duralumin: It is an alloy of aluminium. It contains 95% of aluminium, 4% of copper, magnesium is 0.5% and 0.5% of manganese. It is stronger and lighter than aluminium. Duralumin is used for making bodies of air crafts, helicopters, jets, kitchen ware like pressure cooker. It is also used for making bodies of ships (due to its resistance to sea water corrosion). It is also known as duralium.

(vi) Amalgam: It is an alloy of mercury and one or more other metals is known as an amalagam. It may be solid or liquid. A solution of sodium metal in liquid mercury metal is called sodium amalgam, which is used as a reducing agent. Amalgam of silver, tin and zinc is used by dentists for filling in teeth.


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