Ncert class 10 - science - chapter 3 - metals and non-metals

1. Metals and Non-
metals
Class X – Science Chapter 3 Chemistry

2. Metals vs Non-metals – Physical properties
Metals Non-Metals
Lose electrons to become positive ions
 Cations
e.g. Na+, Mg 2+
Gain electrons to become negative
ions  Anions
e.g. Cl- , O2-
Solid at room temperature except
Mercury
May be solid (Carbon, Sulphur), liquid
(Bromine) or gas (Chlorine, Oxygen)
Lustrous (Shining) Non-lustrous except Iodine
Strong except (Sodium Na or
Potassium K)
Brittle
Hard except Na & K which can be cut
with Knife
Soft except diamond
Malleable – drawn into sheets Non-malleable
Ductile – draw into wires Non-ductile
High melting and boiling point
(except Na and K)
Low Melting and Boiling point except
diamond

3. Metals vs Non-metals – Physical properties
Metals Non-Metals
High densities (Heavy elements) Low densities (Light elements)
Good conductor of electricity and
heat. Lead is poor conductor of heat
Poor conductor of electricity (except
graphite) and heat
Sonorous Non-sonorous
Lies on the left hand side of the
periodic table.
Lies on the right hand of the periodic
table.

4. • Activity to show that Metals are lustrous
Inference: Metal surfaces are lustrous in nature.
Activity 3.1 – Lustrous property
Take a dull metal strip such
as Iron or Copper
Rub the metal surface with
sand paper
Metal shows Shining
surface

5. • Activity to show that Metals are hard
Inference: Metals are hard by nature except few metals such as
Sodium, Magnesium and Potassium.
Activity 3.2 – Hardness of metal
Take different metals
such as Iron, Copper,
Aluminium, Magnesium
and Sodium
Only Magnesium
and Sodium can
be cut with knife.

6. • Activity to show that Metals are malleable
Inference: Metals can be beaten into thin sheets. This property is
called Malleability.
Activity 3.3 – Malleability of Metals
Take metals such
as Iron, Zinc,
lead and Copper.
Hammer
them
Metals change
their shape
and beaten
into thin sheets

7. • Activity to show that Metals are ductile
Inference: Metals can be pulled into wires. This property is
called Ductility.
Activity 3.4 – Ductility of Metals
Observe the materials
common wires are made
of
One could note that
wires are made of
metals such as
Copper, Iron,
Aluminium and Lead

8. Activity 3.5 – Metals conduct heat
• Activity to show that Metals
are good conductors of heat
Inference: Metals conduct
heat which melts the wax at
the other end.
Fix a copper wire
on the Clamp
stand as shown in
figure.
Fix a pin on the
free end of the
wire using wax
Heat the wire
with burner at
other end.
After some
time, pin
drops.

9. Activity 3.6 – Metals conduct Electricity
• Activity to show that Metals
are good conductors of
electricity
Inference: Metals conduct
electricity which is indicated
by the Glowing Bulb.
Setup an electric
circuit as shown in
the figure
Place a metal
between terminals
A and B

10. Activity 3.7 – Physical properties of Non-Metals
• Collect samples of Carbon, Sulphur and Iodine and note down the
physical properties.
Element Symbol Types of
Surface
Hardness Malleab
ility
Ductility Conduction Sonorit
y
Heat Electrici
ty
Carbon C Non-
lustrous
Hard --- --- No Graphit
e – Yes
--
Sulphur S Non-
lustrous
Soft --- --- No No --
Iodine I Jon-
lustrous
Soft --- --- No No ---

11. Metals vs Non-metals – Chemical Properties
Metals
Metal + Oxygen  Metal
oxide
Metal + Water  Metal
oxides / Metal Hydroxides +
Hydrogen
Metal + Acid  Metal
Salt + Hydrogen
Non-Metals
Non-metal + Oxygen  Non-
metal Oxide
Do not react with Water to evolve
Hydrogen. They cannot give
electron to Hydrogen in water
Non-Metals do not react
with acids to release the
gas.

12. Metals vs Non-metals – Chemical Properties
Metals
Metal + Salt Solution  More
reactive metal will displace
less reactive metal
Metal + Chlorine  Metal
chloride
Metal + Hydrogen  Metal
Hydride
Non-Metals
Non-Metal + Salt Solution 
More reactive non-metal will
displace less reactive non-
metal
Non-Metal + Chlorine  Non-
Metal chloride (Covalent bond is
formed)
Non-Metal + Hydrogen 
Non-Metal Hydride

13. Metals with Oxygen
• Almost all metals combine with oxygen to form metal oxides
Metal + Oxygen Metal Oxide
For example, when is heated in air, it combines with oxygen to form
copper (II) oxide, a black oxide.
2Cu + O2 2CuO
(Copper) (Oxygen) (Copper (II) oxide)
Aluminium forms Aluminium Oxide.
4Al + 3O2 2Al2O3
(Aluminium) (Oxygen) (Aluminium oxide)

14. Nature of Metal oxides
NatureofMetalOxides
Basic in nature e.g. Copper Oxide
Amphoteric i.e. reacts with both acids &
base e.g. Aluminium Oxide, Zinc Oxide
Al2O3 + 6 HCl 2 AlCl3 + 3 H2O
Al2O3 + 2 NaOH 2 NaAlO3 + H2O
(Sodium Aluminate)

15. Nature of Metal oxides
NatureofMetalOxides
Insoluble in water
Only certain Metal oxides dissolve in water to form
alkalis e.g. Sodium oxide and Potassium Oxide
Na2O (s) + H2O (l) 2 NaOH (aq)
K2O (s) + H2O (l) 2KOH (aq)

16. Activity 3.9 – Observation of Burning Metals
• Burn metals with the help of pair of tongs to record the following
observations:
S.No. Metal Colour of Flame Colour of
Burnt Metal
Solubility
in Water
1 Na Yellow White Soluble
2 Ca Brick Red colour White Partially
Soluble
3 Mg White White Soluble in
hot water
4 Al White flame White Insoluble
5 Zn No colour White Insoluble
6 Fe No colour Reddish Insoluble
7 Cu Greenish Blue Black Insoluble

17. Non-Metals with Oxygen
• Non-metals react with oxygen to form non-metal oxide
Non-Metal + Oxygen Non-Metal Oxide
• Non-metal oxides are acidic.
C (s) + O2 (g) CO2 (g)
S (s) + O2 (g) SO2 (g)
N2 (g) + 2O2 (g) 2NO2 (g)
• CO and H2O are neutral oxides.
• Non-metal oxides are soluble in water. They dissolve in water to
form acids.
SO2 (g) + H2O (l) H2SO3 (aq) (Sulphurous acid)
CO2 (g) + H2O (l) H2CO3 (aq) (Carbonic acid)

18. Metals with Water
• Not all metals react with water.
• Metals that react with water, produce metal oxide and
hydrogen. If metal oxide is soluble in water, it produces metal
hydroxide.
Metal + Water Metal Oxide + Hydrogen
Metal Oxide + Water Metal Hydroxide
• Potassium and Sodium react with cold water violently.
Reaction is exothermic and Hydrogen evolved catches fire.
2K (s) + 2H2O (l) 2KOH (aq) + H2 (g) + Heat
2Na (s) + 2H2O(l) 2NaOH (aq) + H2 (g) + Heat
• Reaction of Calcium with water is less violent. Hydrogen does
not catch fire. Calcium starts floating as Hydrogen sticks to the
metal
Ca (s) + 2H2O (l) Ca(OH)2 (aq) + H2 (g)

19. Metals with Water
• Magnesium does not react with Cold water. It reacts only with
hot water. It also starts floating as evolved Hydrogen sticks to
the metal surface.
∆
Mg (s) + H2O (l) MgO (s) + H2 (g)
• Aluminium, Iron and Zinc do not react with cold or hot
water. But they can react with Steam to form Metal
oxides and Hydrogen.
2Al (s) + 3H2O (g) Al2O3 (s) + 3H2 (g)
3Fe (s) + 4H2O (g) Fe3O4 (s) + 3H2 (g)

20. Activity 3.10 – How metals react with Water
• Put Samples of small pieces of metals in beakers half-filled with
cold water and observe.
Metals that reacted with cold water Na, K, Ca
Metals that produced Fire Na and K
Metals that started floating after some
time
Ca and Mg
Metals that reacted with Hot Water Mg
Metals that reacted with Steam Al, Zn and
Fe
Metals that did not react with Steam
also
Pb, Cu, Ag
and Au
Reactivity of Metals:
K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Hg > Ag > Au

21. Metals with Acid
• Metals react with acids to give a salt and Hydrogen
gas.
Metal + Dilute Acid Salt + Hydrogen
• Reaction of Metals with dilute Hydrochloric acid is
explained in Activity 3.11
• When metals react with Nitric Acid, Hydrogen gas is
not evolved.
• This is because Nitric acid is an oxidizing agent
which oxidizes Hydrogen to form water and it gets
reduced to any of the Nitrogen Oxides.
• Only Magnesium (Mg) and Manganese (Mn) react
with very dilute HNO3 to evolve H2 gas.

22. Activity 3.11 – Observation of Metals reacting with dilute
Acid
• Put Samples of small pieces of metals (except Na and K) in
separate test tubes containing dilute Hydrochloric acid
Metal Reaction with
HCl
Reaction Rate of Reaction
Mg dil. HCl Mg + 2 HCl → MgCl2 + H2 Very fast
Al dil. HCl 2Al + 6 HCl → 2AlCl3 + 3H2 Fast
Zn dil. HCl Zn + 2 HCl → ZnCl2 + H2 Fast
Fe dil. HCl Fe + 2HCl → FeCl2 + H2 Slow
Cu dil. HCl No reaction

23. Metals with Salt solution of other metals
• Not all metals are equally reactive. Reactive metals
can displace less reactive metals from their
compounds in solution or molten form.
Metal A + Salt Solution Salt Solution + Metal B
of Metal B of Metal A
• In the above reaction, Metal A is more reactive than
Metal B. This reaction is called displacement
reaction.

24. Activity 3.12 – Reaction of Metals with Salt solution
Displacement Reaction
•
• Conclusion: Iron is more reactive than Copper, there it
displaces copper from its salt solution.
Test Tube A
• Iron nail dipped in CuSO4 Solution
• Fe + CuSO4 (aq) → FeSO4 (aq) + Cu
• Blue Colour → Green Colour
• Reddish Brown Deposit of Cu on
Iron Metal
Test Tube B
• Copper wire B dipped in FeSO4
solution
• No reaction

25. Reactivity Series of Metals
K Potassium More Reactive
Na Sodium
Ca Calcium
Mg Magnesium
Al Aluminium
Zn Zinc Reactivity Decreases
Fe Iron
Pb Lead
Cu Copper
Hg Mercury
Ag Silver
Au Gold Least Reactive

26. Reaction of Metal and Non-Metal
Metals Non-Metals
Cations Anions
+ve
ions
-ve
ions
Ionic
Bond
Ionic
Compound
Lose
electrons
Gain
electrons

27. Reaction of Metal and Non-Metal
• For example, Sodium (Metal) and Chlorine (Non-metal) bond with
each other as follows:
Element Sodium Chlorine
Symbols Na Cl
Atomic No 11 17
Electronic
Configuration
2, 8, 1 2, 8, 7
Ion formed Na+ Cl-
Configuration of
Ion
2, 8 2, 8, 8

28. Ionic Bond

29. Properties of Ionic Compounds
Physical
Nature
Ionic Solids are hard due to strong force of attraction
and generally brittle
Melting
Point &
Boiling Point
Have high Melting point and Boiling point
This is because large amount of heat is required to break the
strong Ionic attraction
Solubility
Soluble in water and insoluble in kerosene and petrol.
Conduction
of electricity
Ionic compounds in solid state do not conduct
electricity

30. Activity 3.13 – Properties of Ionic Compounds
Salt Physical
State
Colour of
Flame on
heating
Solubility in Conduction of
ElectricityH2O Petrol Kerosene
NaCl Solid Yellow √ x x Conducts
electricity
KI Solid Pale violet √ x x Conducts
electricity
BaCl2 Solid Pink to white √ x x Conducts
electricity

31. Occurrence of Metals
OccurrenceofMetals
(ItoccursinEarth’scurst,
seawater)
Minerals
Elements or compounds occurring naturally in Earth’s
crust.
Ores
Minerals that contain very high percentage of a
particular metal. These metals can be extracted
economically on large scale.
e.g. Bauxite → Aluminium. Haematite → Iron

32. Occurrence of Metals in Nature
Metal Reactivity Series Occurrence Ore Metallurgy
K
Most Reactive Combine State Electrolysis
Na
Ca
Mg
Al
Zn
Moderately
Reactive
Combine State Sulphide,
Oxide ores
Calcination,
Roasting
Fe
Pb
Cu
Least Reactive
Free and
Combine State
Carbonate
Ores
RoastingHg
Free State Sulphide or
Oxide Ore
Ag
Au

33. Extraction of Metals

34. Conversion of Ore to Metal Oxide
Calcination Roasting
It is done for Carbonate Ores It is done for Sulphide Ores
Heating of carbonate ores in absence
of oxygen
Heating of Sulphide ores in presence
of oxygen
CO2 gas is released and metal oxide
is obtained
SO2 gas is released and metal oxide
is obtained
ZnCO3 (s) → ZnO (s) + CO2 (g)
Heat
ZnS(s) + 3O2 → 2ZnO (s) + 2SO2(g)
Heat

35. Extraction of Metals Low in Activity Series
2Cu2S (s) + 3O2 (g) → 2Cu2O (s) + 2SO2 (g)
2Cu2O (s) + Cu2S → 6 Cu (s) + SO2 (g)
2HgS (s) + 3O2 (g) → 2HgO (s) + 2SO2 (g)
2HgO (s) → 2Hg (l) + O2 (g)
Cinnabar
Ore (HgS)
Mercuric
Oxide (HgO)
Mercury
Reduction (on
further heating)
Roasting
Heating
Heating
Copper
Sulphide
(Cu2S)
Copper
Oxide
(Cu2O)
Copper
Reduction (on
further heating)
Roasting
Heating
Heating

36. Extraction of Metals in the Middle of Activity Series
ZnCO3 (s) → 2ZnO (s) + 2SO2 (g)
2ZnO (s) + C (s) → Zn (s) + CO2 (g)
2ZnS (s) + 3O2 (g) → 2ZnO (s) + 2SO2 (g)
2ZnO (s) + C (s) → Zn (s) + CO2 (g)
Zinc
Sulphide
(ZnS)
Zinc Oxide
(ZnO)
Zinc
Reduction (on
heating with Coke)
Roasting
Heat
Heat
Zinc
Carbonate
(ZnCO3)
Zinc Oxide
(ZnO)
Zinc
Reduction (on
heating with coke)
Calcination
Heat
Heat

37. Thermite Reaction
• Apart from using coke to reduce metal oxides to metals,
displacement reactions can also be used.
• Highly reactive metals such as Sodium, Calcium, Aluminium
are used as reducing agents and displace metals of lower
reactivity than them from the metal oxides.
• For example, when Manganese dioxide is heated with
Aluminium powder, Aluminium displaces Manganese as
follows.
3MnO2 (s) + 4Al (s) 3 Mn (l) + 2Al2O3 (s) + Heat
• These displacement reactions are highly exothermic that the
metals produced will be in molten state.
Fe2O3 (s) + 2Al (s) 2 Fe (l) + Al2O3 (s) + Heat
• This phenomenon is used to join railway tracks.

38. Extraction of Metals in Top of Activity series
HighlyReactiveMetals
They have more affinity for
Oxygen than Carbon.
They cannot be obtained
from their compounds by
heating with carbon.
For e.g. Carbon cannot
reduce oxides of Sodium,
Magnesium, Calcium etc.
These metals are obtained
by electrolytic reduction.
At cathode Na+ + e- → Na
At anode Cl- → Cl + e-
Molten Chlorides
of Highly reactive
metals
Metals are
deposited at
cathode
Chlorine is
liberated at
anode
Electrolytic
Reduction

39. Refining of Metals - Electrolysis
• A strip of impure metal (to be refined) is taken as anode.
• Pure metal (Small Strip) of same material is taken as Cathode.
• Electrolyte of same metal solution is used.
• When electric current is passed, pure metal is deposited at
Cathode.
• Impurities are collected at the bottom of anode (anode mud)

40. Corrosion
• Metals when left open in air gets corroded.
• Silver + Sulphur in air Silver Sulphide (black coat)
• Copper + Carbon dioxide in air Copper carbonate
(green coat)
• Iron + Air and Moisture Iron oxide (rust) (Brown flaky)

41. Prevention of Corrosion
• By applying paint
• By applying oil / grease
• Galvanizing a metal Coating of zinc on metal
• Anodizing a metal Coating a layer of Aluminium
Oxide on Aluminium
• By making alloys Steel - (Fe + Ni + Cr)
Brass - (Cu + Zn)
Bronze – (Cu + Sn)

42. Activity 3.14 – Rusting of Iron Nails
• Take three test tubes A, B, C and place clean Iron
nails.
Test tube A – water
Test tube B – Distilled water
+ 1 ml of oil
Test tube C – anhydrous
Calcium chloride

43. Activity 3.14 – Rusting of Iron Nails
• Leave these test tubes for a few days and then observe.
Observation:
 Nails in Test tube A rusts.
 Iron Nails in Test tubes B & C does not rust
Reason:
 Nails in Test tube A  exposed to air and water
 Nails in Test Tube B  does not get exposed to air
due to layer of oil
 Nails in Test Tube C  Does not get both air and
water
Conclusion: Both air and water are required for iron nails
to form rust.