METALS AND NON-METALS.pdf

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METALS AND NON
-
METALS




Metals:

The elements, which have the properties of luster, malleability and ductility,
high thermal and electrical conductivities. The me
tals have a good tendency to loose
electrons.


Non
-
metals:

The elements, which do not have the above, mentioned properties.
These have a good tendency to gain electrons.


Metalloids:
The elements having the properties of metals as well as non
-
metals.



Dif
ference between metals and non
-
metals


Properties

Metals

Non
-
metals

Physical Properties



1.

State

Metals are
solids

at ordinary
temperature. (except
mercury, which is a liquid.)

Non
-
metals exist in all the three
states, that is,
solid, liquid and gas
.

2.

Lustre

They possess
lustre or
shine
.

They possess
no lustre
.

(except Iodine and graphite.)

3.

Malleability and
Ductility

Metals are generally
malleable and ductile
.

Non
-
metals are
neither malleable
nor ductile
.

4.

Hardness

Metals are generally
hard
.

A
lkali metals are exception.

Non
-
metals possess
varying
hardness
. Diamond is an
exception. It is the hardest
substance known to occur in nature.

5.

Density

They have
high
densities.

They generally possess
low

densities.

6.

Conductivity


(Heat & Electricit
y)

Metals are
good
conductors

of heat and
electricity.

Non
-
metals are
poor conductors

of
heat a
nd

electricity. The only
exception is graphite which is a good
conductor of electricity.

7.

Melting and boiling
point

They usually have
high

melting and boiling

point.

Their melting and boiling point are
usually
low
. The exceptions are
boron, carbon and silicon.


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Chemical Properties of Metals and Non
-
Metals

Properties

Metals

Non
-
Metals

1. Reaction with
oxygen

Metal + oxygen
→
Metal oxide

Example:

4Na + O
2
→
2Na
2
O

Metal oxides are basic (Na
2
O,
CaO, K
2
O, etc.) or Amphoteric (Zno
and Al
2
O
3
)

Non
-
metal + oxygen
→

non
-
metallic oxide

Example
:
S+O
2

→

SO
2

Non
-
metallic oxides are
acidic (SO
2
, CO
2
, etc) or
neutral (H
2
O, CO, N
2
O).

2. Reaction with
water

Metal + oxygen
→
Metal

oxide (Al,
Zn, Fe)
or
Metal hydroxide (K, Na,
Ca, Mg)

Example : (i) 2K +2H
2
O
→
2KOH+H
2

(ii) 2Al +2H
2
O
→
Al
2
O
3
+3H
2

Active metals (K, Na, Ca) react with
coldwater, moderate metals (Mg
react with warm water and reactive
metals (Al, Zn, Fe) react with steam.

Non
-
metals do not react
with water. Non
-
metals
are electronegative
hence do not lose
electrons

Non
–
metal + H
2
O
→

No
reaction


3. Reaction with
acids

Metal + Dilute acid
→

Salt +
Hydrogen

Example :

(i) Mg+ HCl (dil)
→

2Na
C
l + H
2

(ii) Mg+ H
2
SO
4

→

MgSO
4
+ H
2

N
itric acid (oxidizing agent) oxidize
s

H
2

to H
2
O and it self gets reduced to
NO
, N
2
O,
o
r NO
2

Except for Mg and Mn where nitric
acid forms metal nitrate and
liberates H
2

Non
-
metal + Acid
→

No
reaction

Non
-
metals do not
displace hydrogen from
acids.


4. Reac
tion
with salt
s
olutions

More active Metal A + Salt solution
of less active metal of B
→
Salt
solution of metal A + metal B.

Example :

(i) Zn (s) + CuSO
4
(aq)
→

ZnSO
4

(aq) + Cu(s)

More reactive non metal
A + Salt solution of less
reactive non
-
metal
B
→
Salt s
olution of non
-
metal
A
+ non
-
metal B

Example :I) 2 NaBr + Cl
2

→
2NaCl + Br
2

5. Reaction with
chlorine

Metal + chlorine
→

M
etal Chloride

Example : i) Mg + Cl
2

→

MgCl
2

ii) 2Fe+ 3Cl
2

→

2FeCl
3

Non
-
metal+ Chlorine
→

Non
-
metallic chloride

Example: i) H
2
+Cl
2

2HCl
sunlight
diffused
⎯
⎯
⎯
⎯
→
⎯

ii) P
4
+6Cl
2

3
4PCl
⎯
→
⎯

6. Reaction with
Hydrogen

Metal + Hydrogen
→

metal Hydride

Example : i) 2Na + H
2

→

2NaH

Only active metals l
ike Na, K and
Ca reacts with hyd
rogen

Non
-
metal+ Hydrogen
→

Non
-
metallic hydride

Example: i)

2H
2
+O
2

→

2H
2
O

ii) N
2
+3H
2

→

2NH
3

7. Oxides

Oxides of metals are either basic or
amphoteric.

i) Basic oxides turn
red litmus blue
and show

neutralization reaction
with acids or acidic oxides.

Example :

Oxides of n
on
-
metals are
either acidic or n
e
u
tral.

i) Acidic oxides t
urn blue
litmus red and show
neutralization reaction
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i)
water
2
salt
Acid
Basic
2
O
H
NaCl
2
2HCl

O
Na
+
→
+

ii) Amphoteric oxide show

neutralization reaction with acids as
well as base

O
H
2AlCl
6HCl
O
Al
O
H
2NaAlO
2NaOH
O
Al
2
3
3
2
2
2
3
2
+
→
+
+
→
+


with base or basic
oxides.

Example:

i) SO
2
+2NaOH
→

Na
2
S
O
3
+H
2
O

ii) Neutral oxides
d
o not
show neutralization
reaction with either acids
or bases

Example: N
2
O, CO, H
2
O

8
.

Electrochemical
behaviour

Metals are electropositive in
character. They form cations in
solution and are deposited on the
cathode when elect
ricity is passed
through their solution.

Non
-
metals are
electronegative in
character. They form
anions in solution and are
liberated at the anode
when their salt solutions
are subjected to
electrolysis. Hydrogen in
an exception. It usually
forms positive i
ons and is
liberated at cathode.

9
.

Oxidising or
reducing
behaviour

Metals behave as reducing agents.
This is because of their tendency to
lose electrons.

Na
⎯
→

Na
+

+ e
–

Non
-
metals generally
behave as oxidising
agents since they have
the tendency to gain

electrons.

–
2
Cl
e
Cl
2
1
→
+
−


•

Metal + Metal

→


No reaction

•

Metal + Non
-
metal

→

Electrovalent or lonic compound by complete transfer of
electrons from metallic atom to non
-
metallic and forming
corresponding positive and negative ions.

•

Nonmetal

+ Non
-
metal

→

Covalent compound by sharing of electrons



DIFFERENCE BETWEEN IONIC AND COVALENT COMPOUNDS

Ionic compounds

Covalent compounds

1.

Ionic compounds are usually
crystalline solids.

1.

Covalent compounds are usually
liquids or gases. Only some
of them
are solids.

2.

Ionic compounds have high melting
points and boiling points. That is,
ionic compounds are non
-
volatile.

2.

Covalent compounds have usually
low melting points and boiling points.
That is, covalent compounds are
usually volatile.

3.

Ionic compounds conduct electricity
when dissolved in water or melted.

3.

Covalent compounds do not
conduct electricity.

4.

Ionic compounds are usually soluble
in water.

4.

Covalent compounds are usually
insoluble in water (except, glucose,
sugar, urea, e
tc.).

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5.

Ionic compounds are insoluble in
organic solvents (like alcohol, ether,
acetone, etc.).

5.

Covalent compounds are soluble in
organic solvents.



Relative activities or reactivities of metals


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


Reactivity series of metals :


Element

Potassium

Sodium

Barium

Calcium

Magnesium

Aluminium

Zinc

Iron

Nickel

Tin

Lead

Copper

Mercury

Silver

Gold

Platinum

Symbol

K

Na

Ba

Ca

Mg

Al

Zn

Fe

Ni

Sn

Pb

Cu

Hg

Ag

Au

Most Reactive

Least Reactive

Pt


Hydrogen



H

Metals less
Reactive

than Hydrogen

Reactivity decreases downward

Metals More Reactive than Hydrogen



Act
ivity series of
non
-
metals

(Halogens)


F >

Cl > Br > I



Occurrence of metals


Free state

Combined or compound state


Least reactive metals

Reactive metals



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Mineral :

Metals occurring naturally in the earth’s crust in their inorganic elemental or
compound form are called Minerals.

Gangue:

The earthy, sandy and rock impurities
associated with minerals are called
gangue or matrix.

Ores:

The minerals from which the metals can be extracted conveniently and
profitably are called an ore:



Minerals

ores

concentration of ore


Highly reactive metals

Moderately reactive metals

Less reactive metals

Electrolysis of

molten ore

Sulphide ore

Pure metal

Roasting


Carbonate ore

Sulphide ore


Calcination

Roa
sting

Oxide of metal

Reduction to metal

Refining

Reduction to metal

Refining




Corrosion:

The slow eating up (natural reaction of oxidation) of m
etals or metallic
objects by the action of air, moisture etc. is called corrosion.



•

Corrosion of iron is called rusting.


•

Corrosion is mostly harmful but sometimes it is beneficial too


•

Necessary conditions for corrosion:



(i)

Presence of air (or o
xygen) and



(ii)

Presence of moisture (or water)



Prevention of rusting
:


(i)

by painting


(ii)

by applying grease or oil


(iii)

by galvaniz
ation


(iv)

by electroplating or chromoplating


(v)

by alloying



Alloys:

Homogenous mixture of two or more metals

or metal and a non
-
metal


Amalgam:

An alloy in which one of the constituent (metal) is mercury.


Objectives of alloy preparation:


(
i)

To increase the hardness: example
-

Gold alloyed with copper or silver.

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(ii)

To increase the tensile strength : example

–
Magnalium


(iii)

To increase resistance to corrosion: example
–
stainless steel


(iv)

To modify chemical reactivity : example
-

Sodium amalgam.


(v)

To lower the melting point example: solden.






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