3. METALS AND NON-METALS
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.
PHYSICAL
PROPERTIES
Metals
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.
Hardness:
- 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.
Malleability:
-
It is the ability of metals to be beaten into thin
sheets.
-
Gold and silver are the most
malleable metals.
Ductility:
- 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.
Sonority:
- It
is the ability to produce sound on striking hard surface.
- Metals
are sonorous. So they are used to make bells.
Non-metals
- 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 & Symbol |
Type of
surface |
Hardness |
Malleability |
Ductility |
Conducts
Electricity |
Sonority |
|
Graphite (C) |
Non-lustrous |
Hard |
No |
No |
Yes |
No |
|
Coal (C) |
Non-lustrous |
Hard |
No |
No |
No |
No |
|
Sulphur (S) |
Non-lustrous |
No |
No |
No |
No |
No |
|
Iodine (I) |
Lustrous |
No |
No |
No |
No |
No |
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.
CHEMICAL
PROPERTIES OF METALS
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.
HOW DO
METALS AND NON-METALS REACT?
-
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 |
Element |
Atomic
number |
Number of
electrons in shells |
|||
|
K |
L |
M |
N |
|||
|
Noble gases |
Helium (He) |
2 |
2 |
|
|
|
|
Neon (Ne) |
10 |
2 |
8 |
|
|
|
|
Argon (Ar) |
18 |
2 |
8 |
8 |
|
|
|
Metals |
Sodium (Na) |
11 |
2 |
8 |
1 |
|
|
Magnesium (Mg) |
12 |
2 |
8 |
2 |
|
|
|
Aluminium (Al) |
13 |
2 |
8 |
3 |
|
|
|
Potassium (K) |
19 |
2 |
8 |
8 |
1 |
|
|
Calcium (Ca) |
20 |
2 |
8 |
8 |
2 |
|
|
Non-metals |
Nitrogen (N) |
7 |
2 |
5 |
|
|
|
Oxygen (O) |
8 |
2 |
6 |
|
|
|
|
Fluorine (F) |
9 |
2 |
7 |
|
|
|
|
Phosphorus (P) |
15 |
2 |
8 |
5 |
|
|
|
Sulphur (S) |
16 |
2 |
8 |
6 |
|
|
|
Chlorine (Cl) |
17 |
2 |
8 |
7 |
|
|
- 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) |
|
NaCl |
1074 |
1686 |
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.
OCCURRENCE
OF METALS
- 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)
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+
+ e– → Na
At anode: 2Cl– → Cl2 + 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
CORROSION
-
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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