METALS AND NON-METALS

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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

metals

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
.


Difference

between

metals

and

non
-
metals

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Properties

Metals

Non
-
metals

Physical

Properties

1.

State

Metals

are

solids

at

ordinary

temp
e
rat
u
re
.

(e
x
ce
p
t

mercury,

which

is

a

liquid
.
)

Non
-
metals
exist in all
the
three
states,

that

is,

solid,

liquid

and

gas
.

2.

Lustre

T
h
ey

p
o
s
sess

lu
s
tre

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
.
Alkali
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.

T
h
ey

g
e
n
e
ra
ll
y
densities.

possess

low

6. Conductivity
(Heat
&

Electricity)

Meta
l
s

are

g
o
o
d

conductors

of

heat

and

electricity
.

Non
-
metals

are

poor

conductors

of

heat

and

electricity
.

The

only

exception

is

graphite

which

is

a

good

conductor

of

electricity
.

7. Melting
point

and

boiling

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)

2
K

+
2
H
2
O

2
KOH+H
2

(ii)

2
Al

+
2
H
2
O

Al
2
O
3
+
3
H
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

e
l
e
c
tro
n
e
g
a
t
i
ve

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)


2NaCl
+

H
2

(ii)
Mg+ H
2
SO
4


MgSO
4
+

H
2

Nitric acid (oxidizing agent)

oxidizes

H
2
to
H
2
O
and
it self
gets reduced

to
NO,
N
2
O, or

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.
Reaction with salt
solutions

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 solution
of non
-

metal A + non
-
metal B
Example
:I)
2
NaBr
+

Cl
2


2NaCl
+
Br
2



5.
Reaction

with
chlorine


Metal +
chlorine


Metal
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



d

i
f
f

u
s

e

d


2
H
C
l

s
unl
i
g
ht

ii)
P
4
+6Cl
2
 


4 P C l
3


6.
Reaction

with
Hydrogen

Metal +
Hydrogen


metal

Hydride

Example :
i) 2Na
+
H
2


2NaH
Only active
metals
like
Na, K
and
Ca reacts with
hydrogen

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
non
-
metals
are
either acidic
or

neutral.

i) Acidic oxides
turn
blue
litmus
red and show
neutralization

reaction

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i)
Na
2
O


2HCl


2NaCl


H
2
O

Basic

Acid

salt

water

ii)
Amphoteric oxide show
neutralization reaction with
acids as
well
as base

Al
2
O
3


2NaOH


2NaAlO
2


H
2
O

Al

2
O
3



6HCl



2AlCl

3



H
2
O

with
base
or
basic
oxides.

Example:

i)
SO
2
+2NaOH



Na
2
SO
3
+H
2
O

ii)
Neutral oxides do

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

electricity

is

passed

through

their

solution
.

Non
-
metals

are

e
l
ect
r
o
n
e
g
ative

i
n

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

ions

and

is

liberated

at

cathode
.

9
.

Ox
i
d
i
s
i
ng

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
.

1

Cl



e




Cl
–

2

2

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•
Metal +

Metal



No reaction

•
Metal +

Non
-
metal



Electrovalent
or
lonic compound
by
complete
transfer

of

electrons
from metallic atom
to
non
-
metallic
and

forming


•
Nonmetal +

Non
-
metal




corresponding
positive and negative ions.

Covalent compound
by
sharing
of

electrons

DIFFERENCE BETWEEN IONIC
AND
COVALENT

COMPOUNDS

Ionic

compounds

Covalent
compounds

1
.

Io
n
i
c

c
o
mp
o
u
n
ds

a
r
e

us
u
a
ll
y
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
.

C
ov
a
l
e
n
t

c
o
mp
o
u
n
ds

do

n
o
t
conduct

electricity.

4. Ionic compounds are usually soluble
in water.

4
.

Covalent

compounds

are

usually

insoluble

in

water

(except,

glucose,

sugar,

urea,

etc
.
)
.

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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

their

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
Ma
g
nesium
Aluminium
Zinc

Iron
N
i
cke
l

Tin
Le
a
d

Copper
Mercury
Silver
Gold
P
latin
u
m

S
y
mbol

K

Na
Ba
Ca
Mg
Al
Zn
Fe
Ni
Sn
Pb

Cu

Hg

Ag

Au

Pt

Most

Reactive

Least

Reactive

Hydrogen

H

Metals less

Reactive
than

Hydrogen

Reactivity decreases

downward

Metals More Reactive than

Hydrogen

Activity
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

Roasting

Reduction
to

metal

Refining


Corrosion:
The
slow eating
up
(natural reaction
of
oxidation)
of metals 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 oxygen)

and

(ii)
Presence
of
moisture
(or

water)


Prevention of

rusting:

(i)
by

painting

(ii)
by
applying
grease or

oil

(iii)
by

galvanization

(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.

Reduction
to

metal

Oxide
of

metal

Refining

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(ii)

(iii)

(
i
v)

(v)

To increase the
tensile strength
:
example

–
Magnalium

To increase
resistance to corrosion: example
–
stainless

steel

To
modify chemical reactivity
:
example
-

Sodium
amalgam.
To
lower
the
melting point example:

solden.

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