ch-3 Classification of Elements. grade 11

UNIT-3
CLASSIFICATION OF
ELEMENTS AND PERIODICITY
IN PROPERTIES

An arrangement of all the known
elements according to their properties so
that similar elements fall within the
same vertical column and dissimilar
elements are separated.
PERIODIC TABLE

EARLIER CLASSIFICATION
OF ELEMENTS

1. DOBEREINER’S LAW OF TRIADS
2. NEWLANDS’ LAW OF OCTAVES
3. MENDELEEV’S PERIODIC TABLE
4. MODERN PERIODIC TABLE
CLASSIFICATION OF ELEMENTS

Johan Dobereiner (1829)
DOBEREINER’S LAW OF TRIADS

✔In 1829, Dobereiner arranged the known
elements of at that time in the ascending
order of atomic masses.
✔He found out three elements group called
triad.
✔In a triad, the properties of the middle
elements are the average of the other two.
✔This law is known as Dobereiner’s law of
triads.

James Newlands (1865)
NEWLANDS’ LAW OF OCTAVES

✔In 1865, Newlands arranged all the known
elements of at that time in the ascending
order of atomic masses.
✔He observed that the properties of the
eighth elements are the simple repetitions of
the first one like eighth note in an octave in
music.
✔This law is known as Newlands law of
octaves.

Dimitri Mendeleev
MENDELEEV’S PERIODIC TABLE

✔In 1869, a Russian chemist Mendeleev
arranged the known elements of at that
time in the ascending order of atomic
masses.
✔He observed that same properties are
repeated in regular intervals and
proposed a law known as Mendeleev’s
periodic law.

.
The law states that “the physical and
chemical properties of elements are
periodic functions of their atomic masses”
Dimitri Mendeleev
MENDELEEV’S PERIODIC LAW

✔Gallium and Germanium were unknown at
the time Mendeleev published his periodic
table.
✔He left a gap under aluminium and a gap
under silicon.
✔He called these elements Eka-Aluminium
and Eka-Silicon.
NOTE

✔Elements with dissimilar properties are found in
same group.
✔He could not give an exact position for hydrogen.
✔He could not give exact position for Lanthanoids
and Actinoids and also for isotopes.
✔Did not strictly obey the increasing order of
atomic weights.
DEMERTIS OF MENDELEEV’S PERIODIC TABLE

Henry Moseley
MODERN PERIODIC TABLE

✔Moseley’s work on the x-ray spectra of
the elements reveals that atomic number
is a more fundamental property than
atomic mass.
✔On the basis of this, he put forward the
modern periodic law.

The law states that “the physical and
chemical properties of elements are
periodic functions of their atomic
numbers”.
MODERN PERIODIC LAW

✔The horizontal rows present in the modern periodic
table are called periods.
✔There are seven periods.
✔The first period consists of 2 elements.
✔Second and third period consists of 8 elements each.
✔Fourth and fifth period consists of 18 elements.
✔Sixth period consists of 32 elements.
✔The last seventh period is an incomplete period.
PERIODS

✔The vertical columns present in the
modern periodic table are called groups.
✔There are 18 vertical columns.
✔Therefore 18 groups are present in the
modern periodic table.
GROUPS

MODERN CLASSIFICATION
OF ELEMENTS

✔
In the modern periodic table, elements
are classified into four blocks.
✔
They are s, p d and f block elements.
✔
Classification is based on the orbital in
which the last electron of the atom of the
element enters.

✔ The elements in which the last electron enters
the s orbital of their valence shell are called s
block elements.
✔ It consists of elements of group 1 and group 2.
✔ The ground state configuration of the valence
shell is ns1
or ns2
i.e., (ns1 -2
).
S-BLOCK ELEMENTS

✔ The elements in which the last electron
enters the p orbitals of their valence shell
are called p block elements.
✔ It consists of group 13―18 except He.
✔ The ground state configuration of the
valence shell is ns 2 np1
to ns2
np6
.
P-BLOCK ELEMENTS

✔
The elements in which the last electron
enters the d orbitals are called d block
elements.
✔
It consists of groups 3―12.
✔
The general electronic configuration is
(n―1)d 1―10
ns 1―2
.
d-BLOCK ELEMENTS


The elements in which the last electron
enters the f orbitals are called f block
elements.

Their general electronic configuration is
(n―2)f 1―14
(n―1)d 0―1
ns 2
.
f-BLOCK ELEMENTS

All the elements of the s and p block
elements together constitute the
representative elements.
REPRESENTATIVE ELEMENTS

The elements of the 18th
group
are called noble gases or inert
gases or rare gases.
NOBLE GASES

✔ The d block elements i.e., elements of
group 3―12 are called transition elements.
✔ They are placed in between s and p block
elements.
TRANSITION ELEMENTS

✔The f block elements are called inner transition
elements.
✔It consists of Lanthanides and actinides.
✔The elements coming after Lanthanum are called
lanthanides.
✔The elements coming after actinium are called
actinides.
INNER TRANSITION ELEMENTS

CLASSIFICATION OF ELEMENTS INTO
METALS, NON METALS AND METALLOIDS

✔ More than 75% of all known elements
are metals.
✔ Appear on the left side of the periodic table.
✔ Usually solids at room temperature.
✔ Have high melting and boiling points.
✔ Good conductors of heat and electricity.
✔ Malleable and ductile.
METALS

✔ Non-metals are located at the top right hand
side of the periodic table.
✔ Usually exists as solids or gases at room
temperature.
✔ Low melting and boiling points.
✔ Bad conductors of heat and electricity.
NON METALS

✔Metalloids or semi metals are elements
which show both the properties of metals and
non metals.
✔Eg: Boron, Silicon, Germanium, Arsenic,
Antimony, Selinium, Tellurium and Polonium.
METALLOIDS

✔The metallic character increases from
top to bottom of a group.
✔Non metallic character increases from
left to right across a period.
NOTE

✔The elements are named using the numerical
roots for 0 and numbers 1-9.
✔The roots are put together in the order of
digits which make up the atomic number.
✔‘ium’ is added at the end.
NOMENCLATURE OF ELEMENTS
WITH ATOMIC NUMBER GREATER THAN 100

The IUPAC names for the elements with Z
above 100 are shown below.

Properties which are directly or indirectly
related to the electronic configuration of
the elements and show a regular
gradation when we move from left to
right across a period or from top to
bottom in a group are called periodic
properties.

✔The minimum amount of energy required
to remove the most loosely bound electron
from an isolated gaseous atom.
✔Ionisation energy is also known as
Ionisation Potential.
IONISATION ENERGY

✔The energy required to remove the first
electron is called first Ionisation energy
(IE1
).
✔The energy required to remove the second
electron is called second ionisation energy
(IE2
).
✔In general, IE2
> IE1
.

FACTORS INFLUENCING
IONISATION ENERGY

The larger the atomic size, smaller the
ionisation energy.
Smaller the atomic size, larger the ionisation
energy.
Ionisation energy increases with increase
in nuclear charge.
1. ATOMIC SIZE
2. NUCLEAR CHARGE

✔The inner electrons repel the outer electrons and
cut down the attractive force between the nucleus
and the valence shell.
✔This effect is known as shielding effect or
screening effect.
✔As the shielding increases the ionisation energy
decreases.
3. SHIELDING EFFECT

If an atom has half filled or completely filled
sub shells, its ionisation energy is higher
than that expected from its position in the
periodic table.
4. EFFECT OF HALF FILLED AND
COMPLETELY FILLED SUB SHELLS

The energy released when an isolated
gaseous atom changed into an anion by
accepting an electron.
ELECTRON AFFINITY

FACTORS INFLUENCING
ELECTRON AFFINITY

Larger the size of the atom, the smaller will
be the electron affinity and vice versa.
Greater the nuclear charge, greater the
electron affinity.
1. ATOMIC SIZE
2. NUCLEAR CHARGE

✔When the electronic configuration of the atom is stable,
the less will be the tendency of the atom to accept an
additional electron and hence lower will be the electron
affinity.
✔The electron affinity values of halogens are very high
because of their strong tendency to accept an electron to
attain the stable noble gas configuration.
3. ELECTRONIC CONFIGURATION

The tendency of an atom to attract the
shared pair of electrons towards itself.
3. ELECTRONEGATIVITY

✔Small atoms are more electronegative because
they attract electrons more strongly than the larger
ones.
✔Atoms with nearly filled shells will have higher
electronegativities than those with less densely
filled ones.
✔NOTE: The least electronegative element is cesium
and the most electronegative element is fluorine.

✔Most commonly used scales are
✔Pauling’s Scale
✔Mulliken’s Scale
✔Sanderson’s Scale
✔Allred-Rochow’s Scale
SCALES OF ELECTRONEGATIVITY

PAULING’S SCALE
This scale is based on an empirical relation
between the energy of a bond and the
electronegativities of bonded atoms.
MULLIKEN’S SCALE
According to this scale, electronegativity could
be regarded as the average of the ionization
energy and electron affinity of an atom.

PERIODIC TRENDS
ON
IONISATION ENERGY,
ELECTRON AFFINITY AND
ELECTRONEGATIVITY

The Ionisation energy, Electron Affinity and
Electronegativity increases from left to right along a
period. This is because
i) The decrease in atomic size of the elements along a
period.
ii) The increase in nuclear charge on moving along a
period.
iii) Decrease in shielding effect.
ACROSS A PERIOD

The ionisation Energy, Electron Affinity and
Electronegativity decreases down the group.
This is because along a group
i) The size of the atom increases.
ii) The nuclear charge decreases.
iii) Increase in shielding effect.
WITHIN A GROUP

✔It is one half of the distance between the
centres of the nuclei of two bonded atoms
of the same element.
✔Eg: The inter nuclear distance between
the covalently bonded Hydrogen atoms is
74 pm.
✔The covalent radius of Hydrogen is 37
pm.
1. COVALENT RADIUS

It is one half of the distance between the
centres of the nuclei of two non bonded
atoms of the adjacent molecules of the
element in the solid state.
2. VANDER WAALS RADIUS

It is half the inter-nuclear distance
separating the metal atoms in the metallic
crystal.
Eg: The distance between two adjacent
copper atoms in solid copper is 256 pm.
The metallic radius of copper is 128pm.
3. METALLIC RADIUS

The effective distance from the centre of
the nucleus of an ion up to which it has
an influence on the electron cloud.
4. IONIC RADIUS

Atoms and ions containing same number of
electrons.
Eg: Na +
is isoelectronic with F―
.
O2―
is isoelectronic with Mg 2+
.
NO3
―
is isoelectronic with CO3
2―
.
ISO ELECTRONIC SPECIES

1. A cation is smaller than its parent atom but an anion is
larger than its parent atom. Give reason.
✔A cation is smaller than its parent atom.
✔It has fewer electrons while its nuclear charge remains
✔ the same.
✔An anion is larger than the corresponding parent atom
✔The addition of one or more electrons would result in
increased repulsion among the electrons and decrease
✔in effective nuclear charge.

2. The electron affinity of chlorine is higher than that of
fluorine. Why?
✔Fluorine atom is much smaller than chlorine atom.
✔Due to this, there is much crowding of electrons in small space
around the fluorine nucleus.
✔Due to this crowding, fluorine atom has less attraction for the
outside electron in comparison to chlorine in which the crowding
of electrons is less due to the bigger size of chlorine atom.
✔As a result of this, electron affinity of fluorine is less than that
of chlorine.

3. The Ionisation Energy of Nitrogen is greater than
that of Oxygen. Why?
✔ The electronic configuration of Nitrogen is 1s2
,
2s2
, 2p3
✔The electronic configuration of Oxygen is 1s2
, 2s2
,
2p4
.
✔ In the case of Nitrogen atom, the p orbitals are half
filled.
✔Atoms with half-filled electronic configurations
have extra stability.
✔Therefore, the ionization energy of Nitrogen is
greater than that of Oxygen.

ch-3 Classification of Elements. grade 11

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