Periodic table
The Periodic Table of elements is a table displaying all the chemical elements. The elements
are arranged in order of their proton number (atomic number).
Development of the Periodic Table
Some elements such as gold, silver and copper have been known since ancient times. However,
it was Aristotle (in ancient Greece) who first proposed that everything was made up of four
basic substances; earth, air, fire and water. These were named elements by Plato. This theory
was dismissed when more elements started being discovered.
The first scientific discovery of an element occurred in 1649 when Hennig Brand discovered the
element phosphorus. In 1661, Robert Boyle came up with the definition; an element is any
substance that cannot be broken down into simpler substances.
By 1869, a total of 63 elements have been discovered. As the number of
elements grew and their properties became known, scientists tried to devise
ways to classify the elements. Although many scientists came up with
different forms of ‘periodic tables’, Siberian born-Russian chemist, Dmitri
Dmitri Ivanovich Mendeleev
Ivanovich Mendeleev is credited as the father of the modern Periodic Table.
Mendeleev collected a huge amount of data about the elements, some by research, and some
by writing to scientists around the world. He made a card for each of the 63 elements known,
with all the data written on it. He arranged the cards on a table, first in order of increasing
atomic mass, then into groups with similar behaviour, resulting in the Periodic Table,
published in 1869. Part of Mendeleev’s original Periodic Table is shown below.
Period 1
Period 2
Period 3
H
Li
Na
K
Cu
Rb
Ag
Group 2
Be
Mg
Ca
Zn
Sr
Cd
Group 3
B
Al
*
*
Y
In
Group 4
C
Si
Ti
*
Zr
Sn
Group 5
N
P
V
As
Nb
Sb
Group 6
O
S
Cr
Se
Mo
Te
Group 7
F
Cl
Mn
Br
*
I
Group 1
The (*) show gaps Mendeleev deliberately left .
Period 4
Period 5
Mendeleev’s Periodic Table was arranged in order of increasing mass. He left gaps in his table
for elements that were not yet discovered at that time. From the gaps present in his table,
Mendeleev predicted the existence and properties of unknown elements which he called ekaaluminum, eka-boron, and eka-silicon. The elements gallium, scandium and germanium were
found later to fit his predictions quite well.
Fifty years after Mendeleev, the British scientists Henry Moseley discovered that the number
of protons in the nucleus of a particular type of atom was always the same. When the atoms
were arranged according to increasing atomic number, the few problems with Mendeleev’s
Periodic Table disappeared.
Today the Periodic Table has grown a lot since Mendeleev. The current standard table contains
118 elements (as of March, 2009). But not all are found in nature. Out of the 118, 94 are
naturally occurring and the rest made artificially in nuclear reactions.
Differences between Mendeleeve’s periodic table and modern periodic table
Mendeleev’s periodic table
Modern periodic table
Elements arranged in order of increasing mass Elements arranged in order of increasing
number
atomic number
There were about 63 elements
There are about 117 elements
Gaps were left.
No gaps are left
Groups are horizontals rows and periods are Periods are horizontals rows and groups are
vertical columns
vertical columns
The periodic table
In the modern Periodic Table the elements are arranged in order of increasing atomic
number. The elements are arranged in periods and groups.
GROUPS
PERIODS
1
2
3
4
5
6
7
8
1
H
2
Li
B
B
C
N
O
F
Ne
3
Na
Mg
Al
Si
P
S
Cl
Ar
4
K
Ca
Ga
Ge As
Se
B
Kr
5
Rb
Sr
In
Sn
Sb
Te
I
Xe
6
Cs
Ba
Tl
Pb
Bi
Po
At
Ra
7
Fr
Ra
He
Sc Ti
V
Cr Mn Fe
Co
Ni
Cu
Zn
Across the period – size of the atom decreases
Periods
A horizontal row of elements in the Periodic Table is called a period. Each period is numbered
1, 2, 3 etc. The first period contains only 2 elements, hydrogen and helium.
The number of shells an element has corresponds to the period it is found in. Hence, any
element found in period 2 will have 2 shells and any element found in period 4 will have 4
shells.
Ca
C
Carbon (6C) is in Period 2 and has 2 shells
Calcium (20Ca) is in Period 4 and has 4 shells
Groups
A vertical column of the Periodic Table is called a group. The groups are numbered from I to VII
(1 – 7) and then Group 0 or Group VIII (8). Some groups have names as shown in the table.
Group number
Group name
I
Alkali metals
II
Alkaline earth metals
VII
Halogens
0
Noble gases
The group number corresponds to the number of electrons in the outer most shell of an
element. For example any element in Group I will have 1 electron in its outer most shell and
any element in Group VII will have 7 electrons in its outer most shell. As a result of this,
elements in the same group have similar chemical properties.
Na
Li
Lithium (3Li) and sodium (11Na) are in Group 1 and have 1 electron in the outer most shell.
Trends in the Periodic Table
Metals and Non–Metals
Elements can be classified as metals or non-metals. The dividing line between metals and nonmetals run diagonally through the Periodic Table. The elements on the left side of this line are
metals while the elements on the right side of this line are non-metals. Elements which are
close to this line have some properties of metals and some properties of non-metals and are
called metalloids.
Left side
Metals




Proton number increases.
Number of electrons in the outer most shell increases.
Metallic nature decreases.
Changes from good electrical conductors to poor ones.
Right side
Non-metals
Atoms attains stability by achieving noble gas configuration,( outer shells of 2 or 8). And to
become stable, they either lose or gain electrons.
For metals
When metals react, they lose electrons in the outer most shell, to become stable (to attain a
full outer shell).
To lose the outer most electrons, energy is required to overcome the electrostatic force of
attraction between the electrons in the outer shell and the positively charged nucleus.
For non-metals
When non- metals react, they gain electrons to become stable (to attain a full outer shell).
Across the Periods
Going across the periods, from left to right,
• Metals get less reactive
As we go from left to right of the periodic table, the number of electrons in their outer shells
increases. Thus for metals, the number of electrons it has to lose to become stable also
increases. Thus reactivity of metals decreases from left to right of a period.
• Non-metals get more reactive
As the number of electrons in the outer most shells increases from left to right of the
periodic table, the number of electrons, non-metals have to gain to become stable also
decreases. Hence, reactivity of non-metals increases from left to right of a period.
• Atomic size decreases
This is because nuclear charge (number of protons) increases from left to right of the
periodic table and electrons are being added to the same outer shell. (Same period = same
number of shells). Thus electrons are attracted more towards the nucleus with the higher
nuclear charge.
Period 3
Down the group
For metals
Reactivity of metals increases DOWN the group because;
• As we go down the group, the number of shells of atom increases. As a result the atomic
size increases.
• When atomic size increases the electrostatic force (the force of attraction) between the
positively charged nucleus and outer most electrons, which are negatively charged,
becomes less.
• As the outer electrons get further away from the nucleus it becomes easier to lose
electrons (REMEBER: to become stable metals lose electrons).
• Hence the reactivity of metals increases down the group.
For non-metals
Reactivity of non metals decreases DOWN the group because;
• As we go down the group, the atomic size increases .As a result the force of attraction
between the nucleus and outer most electron become less.
• To become stable the non–metals gain electrons and the incoming electrons has to be
strongly held in the shell and this occurs best when the force of attraction is stronger.
• As the outer shell get further away from the nucleus it becomes more difficult to gain
incoming electrons (REMEBER: to become stable non-metals gain electrons).
• Hence the reactivity of non metals decreases down the group.
Metalloids
• Metalloids are elements which have the properties of both metals and non metals
• They often look like metals but are brittle like non metals
• There are 8 metalloids; they are boron, silicon, germanium, antimony, tellurium, polonium
and astatine
Transition metals
• Transition metals are found between group II and III
• These are referred to as d-block elements
• Transition metal compounds are coloured Eg Cu2+ - blue, Fe2+ - green, Cr3+ - orange
ALKALI METALS (GROUP 1)
GENERAL PROPERTIES
 Soft silvery metals and all are easily cut with a knife. When freshly cut, the surface
shows a metallic silvery luster which quickly tarnishes on exposed to air.
 Good conductors of heat and electricity.
 Low density and float on water.
 Low melting points compared to other metals (hence they float).
 Alkali metal compounds are soluble in water.
 Generally very reactive (found at the top of the reactivity series).
-
Hence they tarnish when exposed to air.
-
They are stored in oil.
Melting point
Softness
Density
Reactivity
Li
Na
K
Rb
Cs
Fr
Increases increases
increases
CHEMICAL PROPERTIES
1. REACTION WITH WATER
Alkali metals react vigorously with
water
giving
the
metal
hydroxide
and
hydrogen gas. The trend in reactivity is
demonstrated by the vigor of the reaction of
these metals with water.
OBSERVATIONS:
 Lithium: Floats on the surface, moves around and reacts giving off hydrogen gas.
2Li(s) + 2H2O (l) → 2LiOH (aq) + H2 (g)
 Sodium: Floats on the surface, darts around and melts forming yellow sparks and giving off
hydrogen gas.
2Na(s) + 2H2O (l) → 2NaOH (aq) + H2 (g)
 Potassium: Floats on the surface, darts around, melts and forms a lilac flame giving off
hydrogen gas.
2K(s) + 2H2O (l) → 2KOH (aq) + H2 (g)
2. REACTION WITH AIR OR OXYGEN
Alkali metals exhibit a distinct gradation in reactivity
when exposed to air or oxygen. The surface of freshly cut
lithium quickly tarnishes in air, but the metals need to be
warmed before. It will react quite vigorously to produce
lithium oxide.
4Li(s) + O2 (g) → 2Li2O (s)
TRANSITION METALS
Transition metals are the elements found in between group 2 and group 3 of the periodic
table.
METALS
GENRAL PHYSICAL PROPERTIES
 They are hard and strong.
 They have high tensile strength.
 They are good conductors of heat and electricity.
 They are much less reactive than group 1 and 2 metals.
 They are malleable (can be bend) and ductile (can be drawn into wires).
 They have high melting points (due to strong metallic bonds).
 They are much denser than group 1 and 2.
In addition to this transition metals also have some distinctive properties, which make them more useful for
specific purpose. These include;
 Many of them form coloured compounds.
e.g.: Cu (ii) – Blue
Fe (ii) – Dirty green
Fe (iii) – Reddish brown
 They can have variable oxidation states in their compounds.
 The metals or their compounds can be used as catalysts.
 A few of them have magnetic properties.
COLOURED COMPOUNDS
Salts of group 1, 2 and 3 are usually white and form colourless solutions. But many of the
transition metals are coloured compounds and forms coloured solutions. Transition metals are the
major contributors of colours in our lives.
e.g.:
 Vanadium compounds in solution can be yellow, blue, green or purple.
 The coloured stones on the beach are probably caused by transition metals. Iron makes them
brown and copper makes them green.
 Impressive colours on glass stained windows are produced by transition metals.
 Transition metals produce the colours in gem stones such as sapphire and ruby. These stones are
naturally occurring as colourless crystalline form of aluminium oxide.
-
Traces of titanium and iron make sapphire blue.
-
Traces of chromium as Cr3+ makes ruby red.
The rose window at york minster.
The statue of liberty is made of copper
The colours of the stained glass are
and has become coated in a layer of green
due to the presence of transition metal
carbonate.
ions in the glass.
COLOURS OF SOME TRANSITION METAL IONS IN SOLUTION
Metal ions in solution
Colour
Name of ion
Formula
Copper(ii)
Cu2+
Blue
Iron(ii)
Fe2+
Green
Iron(iii)
Fe3+
Reddish brown
Chromium(iii)
Cr3+
Dark green
Cobalt(ii)
Co2+
Blue
Manganate(vii)
MnO4-
Purple
Chromate(vi)
CrO42-
Yellow
Dichromate(vi)
Cr2O72-
orange
Cu2O
Orange red
CuO
Black solid
Cu(OH)2
Green gelatinous
CuSO4
Blue crystalline solid
ZnO
White solid
Zn(OH)2
White precipitate
ZnSO4
White crystalline solid
PbO
Yellow
PbO2
Dark brown
PbI2
Bright yellow precipitate
PbSO4
White precipitate
FeSO4
Green crystals
Fe2O3
Red powder
Fe(OH)2
Pale-green precipitate
Fe(OH)3
Foxy red-brown precipitate
FeCl2
White solid
FeCl3
Red black solid
VARIABLE OXIDATION STATES
Metals in group 1, 2 and 3 only show one valency. Group 1 metals always show a valency of
1 only (Na+, K+ etc). Group 2 metals like Mg, Ca etc always show a valency of 2 as Ca2+, Mg2+ etc.
Similarly aluminium has a valency of 3. However the transition metals can form ions of more than
one valency. E.g.: Iron atoms can lose either 2 electrons to give iron (ii) ions, Fe2+ or it can lose 3
electrons to give iron (iii) ions, Fe3+. Fe2+ and Fe3+ compounds have different colours and different
properties.
THE OXIDATION STATES SHOWN BY THE FIRST TRANSITION SERIES
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
+3
+4
+5
+6
+7
+6
+4
+4
+2
+2
+3
+4
+3
+6
+3
+3
+2
+3
+2
+4
+2
+2
+2
+2
+1
+3
+2
USES OF TRANSITION METALS
Transition metals have certain properties that make them more useful metallic elements available to
us.
AS CATALYSTS
A catalyst is something used to speed up a chemical reaction which would otherwise be too slow.
Hence in industry a lot (over 90%) of process use catalysts. Many of these catalysts are either
transition metals themselves or their compounds. Catalytic property of transition metals is associated
with the variable oxidation states. Using a catalyst saves time and energy of the manufacture and it
lowers the energy demands. Hence, this is economically advantageous as it helps to conserve energy.
EXAMPLES OF CATALYSTS
Catalyst
Name of process used in
Iron
Haber process: to produce ammonia
Vanadium (v) oxide
Contact process: to produce sulphuric acid
Nickel
Hydrogenation of alkenes: to make margarine
Titanium chloride
Polymerization: manufacture of polythene
Platinum
Oxidation of ammonia to give nitric acid
MAGNETIC PROPERTIES
Very few transition metals have magnetic properties. These include iron, nickel and cobalt.
OTHER USES
 Making alloys: Stainless steel is made by mixing iron, nickel and chromium (all transition
metals) since they don’t go rusty, they are used to produce a lot of goods. E.g.: Knives.
 Titanium has a very great strength, very high melting point and comparatively low density. It is
also corrosion resistant. Hence it is used for the construction of supersonic aircrafts, space
vehicles etc.
 Also because of its very high melting point of 3410oC, it is used in filaments in light bulbs.
HALOGENS
 Group seven elements in the periodic table are known as halogens.
 The word “halogens” is derived from Greek word which means “salt producers”
 All the elements in group seven have 7 electrons in the outer most shell.
Element
State at
Colour
r.t.p
Fluorine
Gas
Pale yellow
Chlorine
Gas
Greenish yellow
Bromine
Liquid
Reddish brown
Iodine
Solid
Blue black solid at r.t.p, Brown solution, Purple vapour
Astatine
Solid
black
 They all have a valency of 1
 They all form a negative ion carrying a single negative charge, and are collectively known as
“halides”. Example: Cl- , Br- , I They produce a series of salts: chloride, bromides, iodides etc. Example: Sodium chloride,
Potassium iodide
 They exist as diatomic molecules (molecules consisting of two atoms. Example: F2 , Cl2 , Br2
Physical properties – Down Group 7
 Colour becomes darker, from pale yellow to black.
 State changes from gas to liquid to solid
 Melting point and boiling point increases. Fluorine has the lowest melting and boiling point
while astatine has the highest melting and boiling point.
Chemical properties
Reactivity trend in group VII
 Halogens react by gaining electrons. non-metals
 Reactivity decreases down the group.
 Fluorine is the most reactive element because the atomic size of fluorine is the smallest in
the group. Therefore the electrons in the outer most shell are closer to the nucleus. Hence
the electrostatic force of attraction between the nucleus and the electrons in the outer shell
is stronger, as a result it becomes easier for Fluorine to gain the electron compared to other
elements in the group.
 Down the group the size of atoms increases. As a result the electrostatic force of attraction
decreases, decreasing the ease of attracting an extra electron. Hence the reactivity decreases
down the group.
Displacement reactions
The more reactive halogen displaces the less reactive halogen
from a solution of its ions.
Cl2 > Br2 > I2
 Add aqueous bromine to aqueous potassium chloride.

No reaction because bromine is less reactive than chlorine.
 Therefore bromine cannot displace chloride ion from its solution.
KCl(aq) + Br2(aq)
No reaction
 Add aqueous bromine to aqueous potassium iodide
 Colourless solution turns brown, because bromine is more reactive than iodine.
 Therefore bromine displaces the iodide ion from its solution. Iodide ion is oxidized to
iodine, which is brown in solution
2KI(aq) + Br2(aq)
Chloride
Bromide
2KBr(aq) + I2 (aq)
Iodide
Chlorine
No reaction
Colourless to reddish brown
Colourless to brown
Bromine
Iodine
No reaction
No reaction
Colourless to brown
No reaction
No reaction
No reaction
Test for halide ions
Halide
ion
Cl -
Test
Acidify with dilute nitric acid and then add aqueous silver nitrate.
KCl (aq)
+ AgNO3 (aq)
AgCl (s) +
KNO3(aq)
Observation
White precipitate
is formed
(white ppt)
Acidify with dilute nitric acid and then add aqueous silver nitrate
Br
-
KBr (aq) +
AgNO3 (aq)
AgBr (s) +
KNO3(aq)
(cream ppt)
Acidify with dilute nitric acid and then add aqueous lead (II)
I-
nitrate
2KI (aq) +
Pb(NO3)2(aq)
Cream precipitate
is formed
Yellow precipitate
is formed
PbI2 (s) + 2KNO3(aq)
(yellow ppt)
Uses of Halogens
Chlorine
 Chlorine is used in the manufacture of plastics such as PVC, Hydrochloric acid,
insecticides and many other chlorinated compounds.
 Chlorine gas is used for sterilizing water and for making „bleach‟.
Bromine
 Used for the preparation of bromides, drugs and dyes.
Iodine
 Iodine is used in medicine because of its antiseptic properties. It is also used in dyes and in
photographic films.
NOBLE GASES
o Group VIII elements in the periodic table are called noble gases or inert gases or zero
group elements.
Elements
Electronic configuration
Helium
2
Neon
2 ,8
Argon
2, 8, 8
Krypton
Xenon
Radon (a radioactive element)
o In group 8, all the elements have completely filled shells. So they have stable electronic
configurations.
o Except helium all other elements in the group have 8 electrons in the outer most
shell(Helium has 2 electrons).
Physical properties
 Exist as mono atomic molecules.
Eg ; He , Ne
 All are colourless gases
Chemical properties
 They are highly unreactive as they have stable electronic configuration. Hence they are also
known as inert gases.
Uses of Noble gases
Helium
 To provide an inert atmosphere for welding
 As a coolant in nuclear reactors
 As a breathing gas used by deep sea divers, by mixing with oxygen.
 Used to fill weather balloons and airships
 Used in the Helium –Neon laser
Neon
Used in advertising sign boards and fog beacons. When electricity is passed through it ,it
glows red.
 Used in Helium –Neon laser
 In Geiger Muller tubes, which are used for detection of radioactivity
Argon
 Used to fill in electric bulbs. The bulbs have tungsten filament and argon prevents the
filament from getting burned.
 To provide an inert atmosphere :
 in the production of Titanium metal.
 in arc welding
 Used in industrial manufacture of steel
Krypton
 It produces intense white light when a current is passed through it.
 Used in airport runway lighting and light house lamps
 Used in special bulbs such as those used in miners lamp
Xenon
 It also produces intense white light when a current is passed through it.
 In lamps
 Used in photographic flash units, in stroboscopic lamps
Radon
 Used to trace gas leakages and treating some forms of cancer