This document provides an overview of metals and non-metals. It discusses their physical and chemical properties, including how metals react with oxygen, water, acids, and other substances. Metals are solid, malleable, ductile, and good conductors of heat and electricity. They form basic oxides and react vigorously with acids. Non-metals do not have these properties and are usually gases or solids. The document also covers extraction methods for metals and corrosion prevention.

1. Standard/ Class/ Grade - X SSC, CBSE; - VIII ICSE
Metals and Non-metals
Gurudatta K Wagh

2. Elements
• Metals
• Non-metals
• Metalloids
Metalloids possess both the properties of
metals as well as non-metals e.g. silicon Si,
germanium Ge, antimony/ stibium Sb, etc.

4. Overview of the Chapter
Metals - Physical properties
Non-metals - Physical properties
Metals - Chemical properties
Reaction of metals with oxygen
Reaction of metals with water
Reaction of metals with acids
Reaction of metals with solutions of other metal salts
Reaction of metals with non-metals
Properties of ionic compounds
Occurrence of metals
Activity series
Extraction of metals of high reactivity
Extraction of metals of medium reactivity
Extraction of metals of low reactivity
Refining of metals
Corrosion of metals
Prevention of corrosion

5. Metals - Physical properties
1. Solids: At room temperature. Exceptions: Mercury/
hydrargyrum Hg and gallium Ga are in liquid state at
room temperature
2. Lustre: Shine in their pure state, polished to give a
highly reflective surface
3. Malleability: Can be made into thin sheets

6. Metals - Physical properties
4. Ductility: Drawn into thin wires. Most ductile: Gold/
aurum Au, silver/ argentum Ag. One gram of gold can
be drawn into a wire of 2 km length
5. Conduction of heat: Good conductors of heat; have
high melting points. Silver Ag and copper/ cuprum Cu
best conductors of heat. Lead/ plumbum Pb and
mercury poor conductors of heat
6. Conduction of electricity: Good conductors of
electricity. Electric wires are made of Cu having a outer
insulation of PVC (polyvinyl chloride)

7. Metals - Physical properties
7. Hardness: Generally hard, differs from metal to metal.
Alkali metals sodium Na and potassium K are soft
metals and can be cut easily with a knife
8. Melting point and Boiling point: High. Tungsten/
wolfram W highest melting point; sodium Na and
potassium K low melting points
9. Sonorous: Produce a sound on striking a hard surface
e.g. bells

8. Non-metals - Physical properties
1. Either in solid or gaseous state. Exception: Bromine Br
is in liquid state
2. Usually do not have lustre. Exception: Iodine I
3. Do not possess the property of hardness. Exception:
Carbon C in the form of diamond. It is the hardest
substance known which also has a high melting and
boiling point
4. Do not conduct electricity. Exception graphite
(allotrope of carbon)

9. Since there are exceptions in each of the physical
properties, elements can be classified more clearly as
metals and non-metals on the basis of their chemical
properties.

10. Metals - Chemical properties
• Reaction of metals with oxygen
Almost all metals react with oxygen to form metal oxide,
but the reactivity differs for different metals
Sodium and Potassium are the most reactive metals.
Sodium reacts with oxygen in air at room temperature to
form sodium oxide. Hence, sodium is stored under
kerosene oil to prevent its reaction with oxygen, moisture
and carbon dioxide
4Na + O2 → 2Na2O

11. Most metal oxides are insoluble in water
but
some of these dissolve in water to form alkalis, and are
basic in nature
e.g. sodium oxide Na2O and potassium oxide K2O dissolve
in water to produce alkalis
Na2O + H20 → 2NaOH
K20 + H20 → 2KOH

12. Magnesium does not react with oxygen at room
temperature
but on heating,
magnesium burns in air with intense light and heat to form
magnesium oxide MgO
2Mg + O2 heat → 2MgO
Zinc burns in air only on strong heating to form zinc oxide
2Zn + O2 heat → 2ZnO
Iron does not burn even on strong heating, but iron filings
burn vigorously when sprinkled in the flame of the burner
3Fe + 202 heat → Fe304

13. Copper is least reactive. It does not burn,
but on heating, the hot metal is coated with a black
coloured layer of copper oxide
2Cu + O2 heat → 2CuO
Aluminium develops a thin oxide layer when exposed to
air
4Al + 302 heat → 2Al2O3

14. Anodising
Anodising is a process of forming a thick oxide layer of
aluminium. Aluminium develops a thin oxide layer when
exposed to air. This aluminium oxide coat makes it
resistant to further corrosion. The resistance can be
improved by making the oxide layer thicker
In this technique aluminium article is used as an anode.
Electrolyte used is dilute sulphuric acid. The anode
reaction results in formation of a black coloured thin film
of aluminium oxide on the surface of anode

16. By putting appropriate dyes in the electrolytic bath,
coloured surface with decorative finish can be achieved
Kitchen articles like anodised pressure cookers, anodised
pans and also frames of sliding windows are the
applications of anodising techniques

17. Amphoteric oxides
Metal oxides are usually basic in nature
but
some metal oxides such as aluminium oxide Al2O3 and zinc
oxide ZnO react both with acids as well as bases to
produce salt and water
Al2O3 + 6HCl → 2AlCl3 + 3H2O
Al2O3 + 2NaOH → 2NaAlO2 (sodium aluminate) + H20

19. • Reaction of metals with water
Metals like potassium and sodium react vigorously with
cold water
Sodium reacts with water to evolve hydrogen which
immediately catches fire producing a lot of heat
(Exothermic reaction)
2K + 2H2O → 2KOH + H2 + heat energy
2Na + 2H2O → 2NaOH + H2 + heat energy

20. Reaction of Na with water

21. Calcium reacts with water less vigorously. The heat evolved
is not sufficient for hydrogen to catch fire
Instead, calcium starts floating because the bubbles of
hydrogen gas formed stick to the surface of the metal
Ca + 2H2O → Ca(OH)2 + H2
Magnesium reacts with hot water to form magnesium
hydroxide Mg(OH)2 and hydrogen H2. Magnesium also
starts floating since the bubbles of hydrogen gas stick to its
surface.

22. Metals like aluminium, iron and zinc do not react either
with cold or hot water but they react with steam to form
metal oxide and hydrogen (g)
2Al + 3H2O → Al2O3 + 3H2
3Fe + 4H2O → Fe3O4 + 4H2
Metals like gold, silver, and copper do not react with water
at all

23. • Reaction of metals with acids
Metals react with acids to give salt and hydrogen gas
Dilute hydrochloric acid HCl
Formation of metal chloride and hydrogen gas
The reactivity decreases in the order,
Mg > Al > Zn > Fe
Mg + 2HCl → MgCl2 + H2↑
2Al + 6 HCl → 2AlCl3 + 3H2↑
Zn + 2HCl → ZnCl2 + H2↑
Fe +2HCl → FeCl2 + H2↑
No bubbles are seen in case of copper. This shows that
copper does not react with dilute hydrochloric acid

24. Sulphuric acid H2SO4
Formation of metal sulphate and hydrogen gas
Fe + H2SO4 → FeSO4(aq) + H2(g)
Iron Fe reacts vigorously and exothermically with sulfuric
acid to produce iron(II) sulfate
Zn + H2SO4 → ZnSO4(aq) + H2(g)
Mg + H2SO4 → MgSO4(aq) + H2(g)

25. Nitric acid HNO3
Hydrogen gas is not evolved when a metal reacts with
nitric acid, as it is a strong oxidizing agent
It oxidizes the hydrogen to water and itself gets reduced to
any of the nitrogen oxides (N20 nitrous oxide, NO nitric
oxide, NO2 nitrogen dioxide)
But magnesium and manganese react with dilute HNO3 to
evolve hydrogen gas
Mg(s) + 2HNO3(aq) → Mg(NO3)2(aq) + H2(g)
magnesium + nitric acid → magnesium nitrate + hydrogen

26. Element Reaction with Air Reaction with Water Reaction with Acid
Potassium K
Sodium Na
Burns vigorously
to form oxides
With cold water forms
hydrogen gas and
alkaline hydroxide
solution. React with
decreasing vigour down
the series till Ca
Violent reaction to
give hydrogen gas
and salt solution
Calcium Ca
Magnesium Mg
Aluminium Al
Zinc Zn
Iron Fe
Burn with
decreasing
vigour down the
series
Mg reacts with hot
water to fomi
magnesium hydroxide
and hydrogen gas. For
Al, Zn and Fe, no
reaction with cold and
hot water. With steam
forms metal oxide and
hydrogen gas
React to form
hydrogen gas and
salt solution with
decreasing vigour
down the series
Reactions of metals with air, water, and acids

27. Aquaregia
Aquaregia is a highly corrosive as well as fuming liquid and
is one of the few reagents that is able to dissolve gold and
platinum
It is a freshly prepared mixture of concentrated
hydrochloric acid and concentrated nitric acid in the ratio
of 3:1
Aqua regia (Latin, lit. "royal water" or "king's water") was
so named by alchemists because it can dissolve the noble
metals gold and platinum

28. • Reactions of metals with solutions of other metal salts
Put an iron nail in the solution of copper sulphate
Fe(s) + CuSO4(aq) → Cu(s) + FeSO4(aq)
metal A + salt solution B → salt solution A + metal B
The iron nail gets coated with a reddish brown colour
copper and the blue colour of copper sulphate solution
fades out
In this reaction more reactive iron has displaced copper
which is less reactive from the copper sulphate solution

29. This reaction is known as displacement reaction. The
brown coating on the iron nail shows that copper is
deposited on the iron nail by displacing iron
The greenish colour of the solution in the test tube shows
that Fe2+ ions are present in the solution. This shows that
iron is more reactive than copper, as Fe2+ ions have
displaced Cu2+ ions from copper sulphate solution

30. This is a single displacement reaction in which copper has
been displaced from iron from copper sulphate solution
A more reactive metal can displace a less active metal from
its compound in a solution

31. • Reaction of metals with non-metals
Let us look at the electronic configuration of some metals
and non-metals. It will help us to understand the reactivity
of different elements and formation of different compounds
We have learnt that noble gases have a completely filled
valance shell, hence are chemically inactive

32. Electronic configuration of some metals and non-metals
Type of
element
Element Atomic
number
Electronic configuration
K L M N
Metals Sodium (Na)
Magnesium (Mg)
Aluminium (Al)
Potassium (K)
Calcium (Ca)
l l
12
13
19
20
2
2
2
2
2
8
8
8
8
8
1
2
3
8
8
1
2
Non-metals Nitrogen (N)
Oxygen (O)
Fluorine (F)
Phosphorus (P)
Sulphur (S)
Chlorine (Cl)
7
8
9
15
16
l7
2
2
2
2
2
2
5
6
7
8
8
8
5
6
7

33. Sodium, a metal, is a silver coloured metal that reacts so
violently with water that flames are produced due to
formation of hydrogen gas
Chlorine, a non-metal, is a greenish coloured gas which is
so poisonous that it was used as a weapon in World War l
When chemically bonded together, these two dangerous
substances form a compound sodium chloride so safe that
we eat it every day (common table salt)

35. Sodium (2, 8, 1) atom has one electron in its outermost
shell. If it loses one electron from its "M" shell then its "L"
shell becomes the outermost shell to acquire a stable
octet. The nucleus of this atom still has 11 protons but
the number of electrons has become 10, so there is a net
positive charge giving us a sodium cation (Na+)
On the other hand chlorine (2, 8, 7) has 7 electrons in its
outermost shell and requires one more electron to
complete its octet

36. The electron lost by sodium is taken up by chlorine. After
gaining one electron, its K, L and M shells have
altogether 18 electrons, but the nucleus still has 17
protons
This leads to the formation of chloride anion (Cl-). Both
these elements have a give and take relation between
them

37. Ionic or electrovalent bond/ compound
Sodium and chloride ions being oppositely charged attract each
other and are held by strong electrostatic forces of attraction to
exist as sodium chloride (NaCl), resulting in formation of an
electrovalent bond or an ionic bond
Sodium chloride exists as aggregates of oppositely charged ions in
definite geometrical shape
Such a bond formed by the give and take of electrons is called as
ionic or electrovalent bond. Compounds formed in this manner by
the transfer of electrons from a metal Na to a non-metal Cl are
known as ionic compounds or electrovalent compounds
Covalent compound
The chemical compound that is formed by mutual sharing of one or
more pairs of electrons between the two combining atoms is called
a covalent compound e. g. water, ammonia, etc.

38. Properties of ionic compounds
1. Ionic compounds are solids and hard due to strong
force of attraction between positive and negative ions
2. They are generally brittle and break into pieces when
pressure is applied
3. Ionic compounds have high melting and boiling points,
as a considerable amount of energy is required to break
the strong inter molecular attraction
4. They are generally soluble in water and insoluble in
solvents such as kerosene, petrol, etc.

39. 5. Ionic compounds in the solid state do not conduct
electricity because the movement of ions in the solid
state is not possible due to their rigid structure, but
they conduct electricity in the molten state. The
conduction of electricity through a solution involves
the movements of charged particles. A solution of an
ionic compound in water contains ions, which move to
the opposite electrodes when electricity is passed
through the solution

40. Occurrence of metals
The most unreactive metals (not affected by air and
water) e. g. silver, gold and platinum are generally found in
free or native state
Most metals however are found in combined state in the
form of their oxide ores, carbonate ores, or sulphide ores,
etc.
Minerals The naturally occurring compounds of metals
along with other impurities are known as minerals
Ores The minerals from which metals are extracted
profitably and conveniently are called as ores

41. Gangue Ores contain metal compounds with some of the
impurities like soil, sand, rocky material, etc. These
impurities are called as gangue
Metallurgy Metals can be extracted from the ores by
employing different separation techniques. The process
used for extraction of metals in their pure form from their
ores is called metallurgy

42. Activity series and related metallurgy
K
Na
Ca
M
g
Al
Zn
Fe
Pb
Cu
Hg
Ag
Au
Decreasing order
of reactivity
Top
↓
↓
↓
↓
↓
Bottom
The arrangement of metals in
the decreasing order of their
reactivity in the form of series
is called the reactivity series
of the metals
Thus the most reactive metal
is potassium and is placed in
the top of the list and least
reactive metal is gold which is
placed at the bottom of the
list

43. Three categories of metals
*Metals of high reactivity
*Metals of medium reactivity
*Metals of low reactivity

44. •Extraction of metals of high reactivity
Metals high up in the reactivity series are very reactive e.g.
sodium, potassium, calcium, aluminium, etc. These metals
are obtained by electrolytic reduction
Sodium, magnesium and calcium are obtained by
electrolysis of their molten chlorides
The metals are deposited at the cathode (-vely charged
electrode), whereas, chlorine is liberated at the anode
(+vely charged electrode). The reaction of sodium is as
follows:
At cathode : Na+ + e- → Na
At anode : 2Cl- → Cl2 + 2e-

45. Extraction of aluminium Al (silvery white, atomic no. 13;
electronic configuration 2,8,3; valency 3)
Aluminium is a reactive metal. Al is extracted from its main
ore bauxite (Al2O3.H2O). It contains 30 % - 70 % aluminium
oxide Al203
Aluminum is obtained by the electrolytic reduction of
aluminium oxide
The remaining portion is gangue made of sand, silica (SiO2),
iron oxide (Fe2O3) etc.

47. Extraction of Al involves two steps
Step 1 Concentration of ore i.e. conversion of bauxite into
alumina (by Bayer's process)
* Crude bauxite contains impurities like iron oxide Fe2O3
and silica SiO2. These impurities are removed by Bayer's
process. In this process, the ore is first crushed and then
treated with hot concentrated caustic soda (NaOH)
solution under high pressure for 2-8 hours at 140 °C to 150
°C in a tank called digester. Aluminium oxide being
amphoteric in nature dissolves in aqueous sodium
hydroxide to form water soluble sodium aluminate
Al2O3 + 2NaOH → 2NaAlO2 + H2O

48. The iron oxide in the gangue does not dissolve in aqueous
sodium hydroxide and is removed by filtration. However
silica from the gangue dissolves in aqueous sodium
hydroxide forming water soluble sodium silicate. Diluting
sodium aluminate with water and then cooling to 50 °C, it
is hydrolysed to give aluminium hydroxide as a precipitate
NaAlO2 + 2H2O → NaOH + Al(OH)3↓
The precipitate is filtered, washed, dried and ignited at
1000 °C to get Alumina (Al2O3)
2Al(OH)3 heat→ Al2O3 + 3H2O

50. Step 2 Electrolytic reduction of alumina
In this process a molten mixture of pure alumina (MP >
2000 °C) is electrolysed in a steel tank. This tank is lined
inside with carbon ( graphite) which acts as a cathode, and
a set of carbon (graphite) rods dipped in the molten
electrolyte act as anode. Cryolite (AlF3-3NaF) and
fluorspar (CaF2) is also added to the mixture to reduce the
melting point to about 1000 °C

51. On passing the current, aluminium is formed at the
cathode. The molten aluminium being heavier than the
electrolyte used, sinks to the bottom of the tank from
where it is removed periodically. On the other hand
oxygen is liberated at the anode
The electrode reactions are:
Anode: 2O2- - 4e- → O2
Cathode: Al3+ + 3e- → Al
The oxygen gas liberated, reacts with carbon anode and
forms carbon dioxide. As the anode gets oxidized during
the electrolysis of alumina, it has to be replaced from time
to time

52. •Extraction of metals of medium reactivity
The metals in the middle of activity series such as iron, zinc, lead,
copper, etc. are moderately reactive. These are usually present as
sulphides or carbonates in nature. It is easier to obtain a metal from
its oxide, as compared to its sulphides and carbonates. The sulphide
ores are first converted into oxides by heating strongly in excess of
air. This process is known as roasting. The carbonate ores are
changed into oxides by heating strongly in limited air. This process is
known as calcination. The chemical reaction that takes place during
roasting and calcination of zinc ores is:
Roasting 2ZnS + 302 → 2ZnO + 2SO2
Calcination ZnCO3 → ZnO + C02
This zinc oxide is then reduced to zinc by using a suitable reducing
agent such as carbon
ZnO + C → Zn + CO

53. Besides using carbon to reduce metal oxides to metals,
highly reactive metals like sodium, calcium, aluminium, etc.
are also used as reducing agents because they can displace
the metal of lower reactivity from their compounds. e.g.
when manganese dioxide is heated with aluminium
powder, the following reaction takes place
3MnO2 + 4Al → 3Mn + 2Al203 + heat
The amount of heat evolved in such reactions is so large
that the metals are produced in the molten state. Another
example of this type of reaction is the thermit reaction in
which iron oxide reacts with aluminium to give iron and
aluminium oxide evolving lot of heat
Fe2O3 + 2Al → 2Fe + Al203 + heat

54. •Extraction of metals of low reactivity
The metals at the bottom of the activity series are least
reactive. They are often found in free state e.g. gold, silver
and copper. But copper and silver are also found in
combined state as their sulphide or oxide ores. e. g. copper
which is found as Cu2S in nature can be obtained from its
ore by just heating in air
2Cu2S + 3O2 heat→ 2Cu2O +2SO2
2Cu2O + Cu2S heat→ 6Cu + SO2

55. Similarly, cinnabar (HgS) is an ore of mercury. When it is
heated in air, it is first converted into mercuric oxide
(HgO). Mercuric oxide is then reduced to mercury on
further heating
heat 2HgS + 3O2 heat→ 2HgO + 2SO2
heat 2HgO heat→ 2Hg + O2

56. Refining of metals
The metals produced by various reduction processes
described above are not very pure. They contain impurities,
which should be removed, to obtain pure metals. The most
widely used method for refining impure metals is
electrolytic refining
Corrosion of metals
Corrosion is degradation of materials due to reaction with
its environment. The major problem of corrosion occurs
with iron, since it is used as a structural material in
constructions, bridges, ship building, automobile etc.

57. Have you ever observed
*Old iron grills in the buildings?
*Copper vessels which are not cleaned for a long time?
*Silver ornament or idols which are kept exposed to air for
a long time?
*Old worn out cars?
*Iron reacts with moist air to acquire a coating of brown
flaky substance called rust.
*Copper reacts with moist carbon dioxide with air and
slowly loses its shine to gain a green coat of copper
carbonate.

58. *Silver articles become black after some time when
exposed to air because it reacts with hydrogen sulphide in
the air to form a coating of silver sulphide
*Aluminum undergoes an oxidation reaction forming a
thin layer of aluminum oxide, which acts as a barrier to
oxygen and water preventing the further oxidation and
resistant to further corrosion. This resistance can be
improved by making the oxide layer thicker. example
anodizing

59. Prevention of corrosion
Corrosion of metals can be prevented if the contact
between metal and air is cut off. This is done in a number
of ways. Some of the methods are given below:
(1) Corrosion can be prevented if the metal is coated with
something which does not allow moisture and oxygen to
react with it
(2) Coating of metals with paint, oil, grease or varnish
prevents the corrosion of metals e. g rusting of iron can be
prevented by this method

60. Coating of corrosive metals with non-corrosive metals also
prevents corrosion. Some of the methods by which metals
can be coated with non-corrosive metals are:
Galvanizing: It is process of giving a thin coating of zinc on
iron or steel to protect them from corrosion. e.g. shiny iron
nails, pins, etc.
Tinning: It is the process of giving a coating of tin, i.e.,
molten tin over other metal. e.g. Cooking vessels made of
copper and brass get a greenish coating due to corrosion.
This greenish coating is poisonous. Therefore they are
given a coating of tin to prevent corrosion (Kalhai)

62. Electroplating: In this method a metal is covered with
another metal using electrolysis. Silver-plated spoons,
gold-plated jewellery, etc. are electroplated
Anodizing: In this method metals like copper and
aluminum are electrically coated with a thin strong film of
their oxides. This film protects the metals from corrosion

64. Alloying: An alloy is an homogenous mixture of two or
more metals or a metal and a non-metal in definite
proportion. The resultant metals called alloys do not
corrode easily, e.g. stainless steel. Examples,
Brass (copper and zinc)
Bronze (copper and tin)
Stainless steel (iron, nickel and chromium, carbon )

66. If one of the metals is mercury, then the alloy is known as
an amalgam
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 parts of pure gold is alloyed with 2 parts of either
copper or silver
**********

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