Wednesday, April 28, 2021

3. Metals and Non-metals | Class 10 CBSE | Web Notes


Elements can be classified as metals or non-metals on the basis of their physical and chemical properties.

The easiest way to start grouping substances is by comparing physical properties.



Metallic lustre:

-   Take samples of iron, copper, aluminium & magnesium.

-   Clean their surfaces by rubbing with sand paper.

-   The surface becomes shining. This property is called metallic lustre.


-   Metals are generally hard and cannot be cut with a knife.

-   The hardness varies from metal to metal.

-   But sodium metal is soft and can be cut with a knife.


-   It is the ability of metals to be beaten into thin sheets.

-   Gold and silver are the most malleable metals.


-   It is the ability of metals to be drawn into thin wires.

-   Gold is the most ductile metal. One gram of gold can be drawn into a wire of about 2 km length.

-   Due to malleability & ductility, metals can be given different shapes according to our needs.

Conductor of heat:

-   Clamp an aluminium or copper wire on a stand.

-   Fix a pin to the free end of the wire using wax.

-   Heat the wire near the place where it is clamped. Heat transfers to the area of wax. It melts wax and the pin drops. But the metal wire does not melt.

-   It shows that metals are good conductors of heat and have high melting points.

-   Silver & copper are the best conductors of heat.

-   Lead & mercury are poor conductors of heat.

Conductor of electricity:

-   Set up an electric circuit as shown below.

-   Place a metal in the circuit between terminals A and B.

-   The bulb glows. It indicates that metals are good conductors of electricity.

-   The wires that carry current have a coating of PVC (polyvinylchloride) or a rubber-like material. They are insulators to prevent from electric shock.


-   It is the ability to produce sound on striking hard surface.

-   Metals are sonorous. So they are used to make bells.


-   There are very few non-metals as compared to metals.

-   Non-metals include carbon, sulphur, iodine, oxygen, hydrogen, etc.

-   Non-metals are solids or gases except bromine (a liquid).

-   Only very few non-metals have some physical properties of metals.

Element &


Type of surface




Conducts Electricity


Graphite (C)







Coal (C)







Sulphur (S)







Iodine (I)







We cannot group elements according to their physical properties alone, as there are many exceptions. E.g.

·    Metals except mercury are solids at room temperature.

·    Metals have high melting points but gallium & caesium have very low melting points. They melt if keep on palm.

·    Iodine is a non-metal but it is lustrous.

·    Carbon is a non-metal that can exist in different forms (allotropes). E.g. Diamond is the hardest natural substance and has very high melting and boiling point. Graphite is a conductor of electricity.

·    Alkali metals (lithium, sodium, potassium) are so soft and can be cut with a knife. They have low densities and low melting points.

Elements can be more clearly classified based on their chemical properties. E.g.

a) Burn magnesium ribbon. Collect the ashes and dissolve in water. Magnesium hydroxide is formed.

2Mg(s) + O2(g) → 2MgO(s)

MgO(S) + H2O(l) → Mg(OH)2 (aq)

Test this solution with red & blue litmus paper. Red litmus becomes blue. i.e., Mg(OH)2 is basic.

b) Burn sulphur powder. Collect the fumes (SO2) by placing a test tube over the burning sulphur.

Add some water to this test tube and shake. Sulphurous acid (H2SO3) is formed.

S(s) + O2(g) → SO2(g)

SO2(g) + H2O(l) → H2SO3 (aq)

Test this solution with blue and red litmus paper. Blue litmus becomes red. i.e., H2SO3 is acidic.

Most non-metals produce acidic oxides when dissolve in water. Most metals, give rise to basic oxides.


What happens when metals burn in air?

-   Try to burn various metals (aluminium, copper, iron, lead, magnesium, zinc and sodium) over a flame.

-   Sodium burns easily with yellow flame (Magnesium: white flame, Copper: Blue-green, Aluminium: White).

-   The metal surface appears silver white after burning.

-   Reactivity of metals to oxygen in decreasing order is 

Na > Mg > Al > Zn > Fe > Pb > Cu

-   Of the products (metal oxides), sodium oxide is soluble in water. Other metal oxides are insoluble.      

-   Almost all metals combine with O2 to form metal oxides.

Metal  +  Oxygen Metal oxide

E.g., when copper (Cu) is heated in air, it forms copper(II) oxide (CuO), a black oxide.

2Cu    +     O2          2CuO

Similarly, aluminium forms aluminium oxide (Al2O3).

4Al    +    3O2      2Al2O3

-   Metal oxides are basic in nature. But some metal oxides, such as aluminium oxide, zinc oxide show both acidic and basic behaviour. Such metal oxides which react with both acids & bases to produce salts and water are called amphoteric oxides. E.g. reaction of Aluminium oxide.

Al2O3 + 6HCl      2AlCl3 + 3H2O

Al2O3 + 2NaOH 2NaAlO2               +            H2O

                                  (Sodium aluminate)

-   Most metal oxides are water insoluble but some (sodium oxide, potassium oxide etc.) are soluble to form alkalis.

Na2O(s) + H2O(l) 2NaOH(aq)

K2O(s)  + H2O(l) 2KOH(aq)

-   All metals show different reactivities towards oxygen.

-   Metals such as potassium & sodium react so vigorously that they catch fire if kept in the open. Hence, they are kept in kerosene oil.

-   At ordinary temperature, the surfaces of metals such as Mg, Al, Zn, Pb etc., are covered with a thin protective layer of oxide. It prevents metal from further oxidation.

-   Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner.

-   Copper does not burn, but the hot metal is coated with a black coloured layer of copper(II) oxide.

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

Anodising: A process of forming a thick oxide layer of aluminium. This improves resistance against corrosion.  During anodising, a clean aluminium article is made the anode and is electrolysed with dilute H2SO4. The oxygen gas evolved at the anode reacts with aluminium to form a thick oxide layer. This oxide layer is dyed.

What happens when metals react with water?

-   Metals react with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to form metal hydroxide. But all metals do not react with water.

Metal + Water Metal oxide + Hydrogen

Metal oxide + Water Metal hydroxide

-   Metals like potassium & sodium react violently with cold water. The reaction is exothermic so that the evolved hydrogen catches fire.

2K(s) + 2H2O(l) 2KOH(aq)  +  H2(g) + heat energy 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) + heat energy

-   Reaction of calcium with water is less violent. The heat evolved is not sufficient for the hydrogen to catch fire.

Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)

-   Calcium starts floating because the bubbles of hydrogen gas formed stick to the metal surface.

-   Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. It also starts floating because bubbles of hydrogen gas stick to its surface.

-   Metals like aluminium, iron & zinc do not react with cold or hot water. But they react with steam.

2Al(s) + 3H2O(g) Al2O3(s) + 3H2(g)

3Fe(s) + 4H2O(g) Fe3O4(s) + 4H2(g)

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

-   Reactivity of metals with water in decreasing order:

Na > K > Ca> Mg > Al > Fe

What happens when Metals react with Acids?

-   Metals react with acids to give a salt and hydrogen gas.

Metal + Dilute acid Salt + Hydrogen

-   But all metals do not react in the same manner. E.g.

-   Put Al, Cu, Fe, Pb, Mg and Zn separately in test tubes containing dilute hydrochloric acid.

-   Suspend thermometers in the test tubes, so that their bulbs are dipped in the acid.

-   The following reactions occur.

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

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

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)

-   Here, magnesium shows fastest reaction and hydrogen bubble formation and records highest temperature (most exothermic).

-   The reactivity decreases in the order Mg > Al > Zn > Fe.

-   Copper does not react with dilute HCl. So bubbles are not formed and no change in temperature.

-   Hydrogen gas is not evolved when a metal reacts with nitric acid. It is because HNO3 is a strong oxidising agent. It oxidises the H2 to water and itself gets reduced to nitrogen oxides (N2O, NO or NO2). But magnesium and manganese react with very dilute HNO3 evolving H2.

Aqua regia, (Latin for ‘royal water’) is a mixture of conc. HCl & conc. HNO3 in 3:1 ratio. It is highly corrosive, fuming liquid. It is one of the few reagents that can dissolve gold and platinum.

How do Metals react with Solutions of other Metal Salts?

-   Put a clean copper (Cu) wire in iron sulphate (FeSO4) solution and an iron (Fe) nail in copper sulphate (CuSO4) solution taken in test tubes.

-   After 20 minutes, it is observed that blue CuSO4 solution turns green.

-   Iron is more reactive than copper. It displaces Cu from CuSO4 to form FeSO4. So CuSO4 solution turns green.

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

-   This is a displacement reaction. i.e., reactive metals displace less reactive metals from their compounds in solution or molten form.

-   Displacement reactions are better evidence for the reactivity of metals. If metal A displaces metal B from its solution, it is more reactive than B.

Metal A + Salt solution of B Salt solution of A + Metal B

-   All metals do not react with reagents like oxygen, water and acids. So we cannot put all the metal samples in decreasing order of their reactivity.

The Reactivity or Activity Series:

It is a list of metals arranged in the order of their decreasing activities.


-   Noble gases have a completely filled valence shell. So they show little chemical activity.

-   It means reactivity of elements is a tendency to attain a completely filled valence shell (noble gas configuration).

Electronic configurations of some elements:

Type of element


Atomic number

Number of electrons in shells





Noble gases

Helium (He)






Neon (Ne)






Argon (Ar)







Sodium (Na)






Magnesium (Mg)






Aluminium (Al)






Potassium (K)






Calcium (Ca)







Nitrogen (N)






Oxygen (O)






Fluorine (F)






Phosphorus (P)






Sulphur (S)






Chlorine (Cl)






-   A sodium atom has one electron in outermost shell (M).

-   If it loses the electron from M shell then its L shell becomes the outermost shell giving a stable octet. Now the number of electrons is 10 but the nucleus has 11 protons. So there is a net positive charge giving a sodium cation (Na+).

-   Chlorine has 7 electrons in outermost shell (M) and it needs one more electron to complete its octet.

-   When sodium reacts with chlorine, the electron lost by sodium is taken up by chlorine. Thus chlorine gets a stable octet and total 18 electrons in K, L and M shells. But its nucleus has only 17 protons. So, the chlorine atom gets a unit negative charge forming a chloride anion (Cl). Thus, Na and Cl have a give-and-take relation.

Na          →          Na+ + e

2,8,1                    2,8       

Cl   +     e–      →    Cl

2,8,7                    2,8,8

Formation of sodium chloride

-   Being oppositely charged, Sodium & chloride ions are held by strong electrostatic forces of attraction to exist as sodium chloride (NaCl). NaCl does not exist as molecules but aggregates of oppositely charged ions.

Formation of magnesium chloride (MgCl2):

Mg         →        Mg2+ + 2e

2, 8, 2                2, 8

                    (Magnesium cation)

Cl     +   e →        Cl (Chloride ion)

2, 8, 7                  2, 8, 8

-   The compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds or electrovalent compounds.

Properties of Ionic Compounds

-   Take salt samples such as sodium chloride, potassium iodide (KI), barium chloride (BaCl2) etc.

-   Their physical state is hard and brittle. Soluble in water, but insoluble in petrol and kerosene.

-   Heat each sample directly on the flame using a spatula. They do not melt. Flame colour changes (NaCl = orange yellow, KI = violet and BaCl2 = green).

-   Make a circuit and insert the electrodes into a solution of each salt. Salt solution conducts electricity.

Melting & boiling points of some ionic compounds

Ionic compound

Melting point (K)

Boiling point (K)




General properties for ionic compounds:

a.   Physical nature: Ionic compounds are solids and are hard due to strong force of attraction between positive & negative ions. They are generally brittle and break into pieces when pressure is applied.

b.  High Melting & Boiling points: This is because a considerable amount of energy is needed to break the strong inter-ionic attraction.

c.   Solubility: Generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.

d.  Conduction of Electricity: It occurs through a solution by the movement of charged particles. A solution of an ionic compound in water contains ions. When electricity is passed, ions move to opposite electrodes. Solid ionic compounds do not conduct electricity because ions cannot move due to rigidity. But they conduct electricity in molten state because electrostatic forces of attraction between ions are overcome due to heat. Thus, the ions move freely and conduct electricity.


-   The earth’s crust is the major source of metals.

-   Seawater also contains soluble salts such as sodium chloride, magnesium chloride, etc.

-   The elements or compounds which occur naturally in the earth’s crust are called minerals.

-   The minerals that contain a very high percentage of a particular metal are called ores. Metal can be profitably extracted from it.

Extraction of Metals

-   Metals are found in the earth’s crust in the free state or as compounds.

-   Based on the reactivity, metals are 3 types:

a)    Metals of low reactivity: The metals at the bottom of the activity series. They are found in a free state. E.g. gold, silver, platinum and copper. Cu & Ag are also found in the combined state as sulphide or oxide ores.

b)   Metals of medium reactivity: Metals in the middle of the activity series (Zn, Fe, Pb, etc.). Found in earth’s crust mainly as oxides, sulphides or carbonates.

c)    Metals of high reactivity: Metals at the top of activity series (K, Na, Ca, Mg & Al). They are never found in nature in free state.

-   Ores of many metals are oxides. This is because oxygen is very reactive and is very abundant on the earth.

-   Different techniques are used to obtain metals in each category.

-   There are several steps in the extraction of pure metal from ores. They are given below:

Enrichment of Ores

-   Ores are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue.

-   There are different separation techniques to remove the gangue based on the differences between physical or chemical properties of gangue and ore.

Extracting Metals Low in the Activity Series

-   The oxides of these metals can be reduced to metals by heating. E.g., cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted to mercuric oxide (HgO). Then it is reduced to mercury on further heating.

-   Cu2S (copper ore) on heating in air is reduced to copper.

Extracting Metals in the Middle of the Activity Series

-   These are present as sulphides or carbonates in nature.

-   It is easier to obtain a metal from its oxide, as compared to its sulphides & carbonates. So, the metal sulphides and carbonates are first converted into metal oxides.

-   The sulphide ores are converted into oxides by heating strongly in presence of excess air. This is called roasting.

-   The carbonate ores are changed into oxides by heating strongly in limited air. This is called calcination.

Roasting of Zinc ore:

Calcination of Zinc ore:

-   The metal oxides are then reduced to the metals by using suitable reducing agents such as carbon. E.g., when zinc oxide is heated with carbon, it is reduced to zinc.

ZnO(s) + C(s) Zn(s) + CO(g)

-   Obtaining metals from their compounds is also a reduction process.

-   Besides using carbon (coke) to reduce metal oxides to metals, sometimes displacement reactions can also be used. The highly reactive metals such as sodium, calcium, aluminium, etc., are used as reducing agents because they can displace metals of lower reactivity from their compounds. E.g., when manganese dioxide is heated with aluminium powder, following reaction takes place

3MnO2(s) + 4Al(s) 3Mn(l) + 2Al2O3(s) + Heat

-   Here, MnO2 is reduced and Aluminium is oxidised.

-   These displacement reactions are highly exothermic. So the metals are produced in molten state.

-   The reaction of iron(III) oxide (Fe2O3) with aluminium is used to join railway tracks or cracked machine parts. This reaction is called the thermit reaction.

Fe2O3(s) + 2Al(s) 2Fe(l) + Al2O3(s) + Heat

Extracting Metals towards the Top of the Activity Series

-   Highly reactive metals cannot be obtained from their compounds by heating with carbon. E.g., carbon cannot reduce the oxides of Na, Mg, Ca, Al, etc. This is because these metals have more affinity for oxygen than carbon.

-   These metals are obtained by electrolytic reduction.

-   E.g., Na, Mg and Ca are obtained by the electrolysis of their molten chlorides. The metals are deposited at the cathode (negatively charged electrode), whereas, chlorine is liberated at the anode (positively charged electrode). The reactions are

              At cathode:        Na+ + eNa

              At anode:           2ClCl2 + 2e

-   Similarly, aluminium is obtained by the electrolytic reduction of aluminium oxide.

Refining of Metals

-   The metals produced by reduction are not very pure.

-   The most widely used method for refining impure metals is electrolytic refining.

Electrolytic Refining:

-   Metals such as Cu, Zn, Sn, Ni, Ag, Au, etc., are refined electrolytically.

-   In this process, impure metal is made the anode and a thin strip of pure metal is made the 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 (anode mud).

Electrolytic refining of copper


-   Silver articles become black after some time when exposed to air. This is because it reacts with sulphur in the air to form a coating of silver sulphide.

-   Copper reacts with moist CO2 in the air and slowly loses its shiny brown surface and gains a green coat. This green substance is basic copper carbonate.

-   Iron when exposed to moist air for a long time acquires a coating of a brown flaky substance called rust.

Conditions under which Iron rusts

-   Take three test tubes A, B and C and place clean iron nails in each of them.

-   Pour some water in test tube A and cork it.

-   Pour boiled (to remove dissolved air) distilled water in test tube B, add 1 mL oil and cork it. The oil floats on water and prevent air from dissolving in the water. 

-   Put some anhydrous calcium chloride (drying agent to absorb moisture from the air) in test tube C and cork it.

-   After few days, the iron nails rust in test tube A, but they do not rust in test tubes B and C.

-   In test tube A, the nails are exposed to both air and water.

-   In test tube B, the nails are exposed to only water.

-   The nails in test tube C are exposed to dry air.

-   This shows that both air and moisture are needed for rusting of iron.

Prevention of Corrosion

-   Prevention of Rusting of iron: Painting, oiling, greasing, galvanising, chrome plating, anodising or alloying.

-   Galvanisation: A method of protecting steel and iron from rusting by coating them with a thin layer of zinc. The galvanised article is protected against rusting even if the zinc coating is broken because zinc is more reactive than iron and hence can be easily oxidised.

-   Alloying: A method of addition of other substances (metal or non-metal) to a metal to get new desired properties. Such a homogeneous mixture of two or more metals, or a metal and a nonmetal is called Alloy.

It is prepared by melting the primary metal, and then, dissolving the other elements in definite proportions. It is then cooled to room temperature. E.g.

·  Iron is never used in its pure state because it is very soft and stretches easily when hot. If it is mixed with about 0.05 % of carbon, it becomes hard and strong.

·  Mixing iron with nickel & chromium forms stainless steel. It is hard and does not rust.

Pure gold (24 carat gold) is very soft. So it is not suitable to make jewellery. It is alloyed with silver or copper to make it hard. In India, 22 carat gold is used to make ornaments (22 parts pure gold + 2 parts Cu or Ag).

-   The alloy containing mercury is called an amalgam.

-   Electrical conductivity & melting point of an alloy is less than that of pure metals. E.g., brass (alloy of Cu & Zn) and bronze (alloy of Cu & Sn) are not good conductors of electricity whereas copper is used for making electrical circuits. Solder (alloy of Pb & Sn) has a low melting point, it used for welding electrical wires.

The wonder of ancient Indian metallurgy

The iron pillar (built 1600 years ago) near the Qutub Minar in Delhi has high quality of rust resistance. It is 8 m high and weighs 6 tonnes.

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