Classification of elements
and Periodicity in properties
Classification of elements
● The earlier attempt to classify elements involve dividing them into metals and
nonmetals.
● The classification was done by Antoine Laurent Lavoisier (Father of
chemistry).
● This method failed because most of the elements were metals and a very few
were non metals.
● Some elements resembled metals and nonmetals and hence they could not
be classified.
Dobereiner's triads
● In the year 1817, Johann Wolfgang Dobereiner,
a German chemist classified elements having
similar properties into groups called triads.
● When the elements were arranged in groups of 3, in the
increasing order of atomic masses, the middle element
has the atomic mass and properties roughly the average
of the other two elements.
Alkali metal group
Alkaline earth metal group
Element
Element
At: mass
At: mass
Li
7
Ca
40
Na
23
Sr
88
K
39
Ba
137
Halogen group
Element
At: mass
Cl
35.5
Br
80
I
127
Newlands law of Octaves - John Newland
When elements are arranged in the increasing order of their atomic masses (from
H to Th), the properties of the 8th element resembled with the properties of the
first element. This was compared to the octaves found in music. Hence called
Newlands law of octaves.
Mendeleev’s Periodic Law - Dmitri Ivanovich
Mendeleev
When Mendeleev started his work , 63 elements were known.
He examined the relationship between the atomic masses of
the elements and their physical and chemical properties.
“ The properties of elements are a periodic function of their
atomic masses”
Features of Mendeleev’s Periodic table.
● The vertical columns were called groups and horizontal
rows were called periods. There were 6 periods and 8
groups in the table.
● The elements in each group resembles with each other in
many properties.
● All the elements in the group show the same valency.
● The properties like atomic size, MP and BP changed
gradually in any period while going from left to right.
Advantages of Mendeleev’s Periodic table.
● Mendeleev could classify all the 63 elements discovered at that time
on the basis of similarity in properties.
● He left gaps for elements yet to be discovered. He named them by
prefixing the sanskrit numeral, Eka (One) to the name of the
preceding element in the same group. Eg: Scandium, gallium and
germanium have properties similar to Eka-Boron, Eka-Aluminium and
Eka-Silicon respectively.
● He predicted the properties of the undiscovered elements and helped
in the discovery of these elements.
● His periodic table could accommodate noble gases when they were
discovered.
Modern Periodic Law
Henry Moseley showed that atomic number is a more
fundamental property than atomic mass.
“ Properties of elements are periodic function
of their atomic number”
● Atomic number (Z) is the number of protons present in an
atom.
● Z increases by one in going from one element to the next.
● Prediction of properties of elements was more precise
with the arrangement of elements in the increasing order
of atomic number.
Notation for IUPAC Nomenclature of Elements
Nomenclature of elements with atomic number above 100
General features of long form periodic table
● The elements are arranged in the increasing order of
atomic number.
● There are 7 horizontal rows called periods.
● There are 18 vertical columns called groups.
● Elements having similar valence electrons are placed in
the same group.
● Each period starts with the filling of a new shell.
● The 14 elements after Lanthanum are called lanthanoids.
● The 14 elements after Actinium are called actinoids
● To limit the size of the periodic table and to keep the elements
with similar properties together, Lanthanoids and Actinoids are
placed at the bottom of the periodic table.
● The elements of group1, 2 and 13-17 (s-block and p-block) are
called representative elements or main group elements. They
have completely filled inner shells but incomplete valence shell.
● The elements of group 3-12(d-block) are called transition
elements. They have incomplete outer shells and penultimate
shells.
● The elements placed at the bottom of the periodic table (fblock) are called inner transition elements. They are also called
rare earth metals because they occur rarely in the earth's
crust.
● Group 18 elements are called noble gases or inert gases.
Their outer orbits are completely filled and are unreactive.
● A zig-zag line separates metals from nonmetals by some
elements called metalloids or semimetals (B, Si, Ge, As, Sb,
Te and Po)
● Metals are found on the left side of the periodic table and
nonmetals on the right side.
Important points to note
● Elements in any group have same number of valence
electrons.
● Size of the atom decreases across a period and increases
down a group.
● Each period starts with a new shell.
● Maximum number of electrons in each shell is given by
the formula 2n2, where n is the number of the given shell.
(K=1, L=2, M=3, ….)
Periodicity
The periodic repetition of elements with similar properties after certain
regular intervals when the elements are arranged in the order of
increasing atomic number.
● 1st period consists of 2 elements (1s)
● 2nd period consists of 8 elements (2s,2p)
● 3rd period consists of 8 elements (3s,3p)
● 4th period consists of 18 elements (4s,3d,4p)
● 5th period consists of 18 elements (5s,4d,5p)
● 6th period consists of 32 elements (6s,4f,5d,6p)
Groups
● Group I and 2 - s- block elements
● Group 3 to 12 - Transition elements ( d-block elements)
● Groups 13 to 18 - p-block elements
● Lanthanides and actinides (Placed below - f - block elements)
● Those elements in which the filling of electrons takes place in
4f orbital is called lanthanoid series.
● Those in which 5f orbitals are filled are called actinoid series.
Different blocks of elements
S-block elements ( ns1-2 )
● ns1 - Alkali metals and ns2 - Alkaline earth metals.
1. They are all metals
2. Possess low ionisation energy since they have 1 or 2
electrons in their valence shell.
3. They form ionic compounds.
4. Metallic character increases down the group.
p - block elements ( ns2 np1-6)
● p block elements consists of groups 13 to 18 together with
the s block elements are called representative elements
or main group elements.
1. Group 16 - ns2 np4 - Chalcogens(Most of them are found
as oxide or sulphide)
2. Group 17 - ns2 np5 - Halogens
3. Group 18 - ns2 np6 - Noble gases or inert gases
d - block elements Transition elements ( (n-1)d1-10 ns1-2)
● They are called d block elements since they are characterised
by the filling of inner d orbitals.
1. They are all metals
2. They form coloured ions
3. Exhibit variable valency
4. They form complex compounds (Co(NH3)6)Cl3
Zn, Cd and Hg do not show the properties of transition elements
since they have completely filled d-orbitals.
f - block elements Inner transition elements
( (n-2)f1-14 (n-1)d0-1 ns2)
1. They are metals
2. Most of the actinoids are radioactive.
3. The elements coming after Uranium are called transuranium elements.
● Elements are also classified as metals, nonmetals and
metalloids or semi metals.
● Semi metals are those that show the properties
characteristic of both metal and nonmetals.
● 113 - Nh -Nihonium - Japan
● 115 - Mc - Moscovium - Moscow region
● 117 - Ts - Tennessine - tennessee
● 118 - Og - Oganesson - Yuri Oganessian
Periodic Trends in properties
1. Size:
● Atomic radius - It is defined as the radius of the sphere or the
distance between the centre of the nucleus and the outermost shell of
electrons.
● Covalent radius - It is the half the distance between two atoms when
they are joined by a single bond in a covalent molecule.
● Metallic radius - It is taken as the half of the internuclear distance
separating the metal ions in the metallic crystal.
Variation
● From left to right - Size decreases due to the increase in nuclear
charge. So the electrons are pulled more closer to the nucleus. Group
one elements (Alkali metals) are the largest in size.
● Top to bottom - Size increases as new shells are added up. Thereby
attraction decreases.
● Size of the cations : Cations are formed by the loss of electrons.
Cations are smaller than the neutral atom because when the
electrons are lost the ion has greater number of protons. So there are
more protons acting on lesser number of electrons and hence
attraction increases.
● Size of the Anions : Anions are formed by the gain of electrons. Size
of the anion is greater than the neutral atom because there are more
electrons in anions than the number of protons. Hence attraction
decreases.
● Iso electronic ions: Isoelectronic ions are those atoms or ions which
contains the same number of electrons.
Eg: O2- , F- , Na+ , and Mg2+. Cation with greater positive charge will have
smaller size due to greater attractive forces.
Size is Mg2+ <
Na+ < F- < O2-
2. Ionisation enthalpy
● It is the energy required to remove the most loosely bound electron from a
gaseous neutral atom in the ground state.
Factors affecting Ionisation energy
a. Size of the atom: Smaller the size greater is the attractive force and hence
higher IE
b. Effective nuclear energy: Greater the charge greater the attractive forces and
higher the IE values.
c. Shielding of the inner orbitals : The reduction in the force of attraction by the
electrons present between the nucleus and the valence electrons is called
screening or shielding effect. Weaker the shielding effect greater is the
attractive force and higher is the IE.
● Ionisation energy increases with the increase in atomic number across a
period, due to the increase in nuclear charge and decrease in atomic size.
● Ionisation energy decreases down a group due to the increase in atomic size
and increase in the shielding effect.
● Noble gases have high IE due to stable octet configuration.
● Second IE is always greater than the first, after removing one electron nuclear
charge increases.
● IE of N and P is high due to the stable half filled electronic configuration
resulting from symmetrical distribution of electrons and high exchange
energy.
● IE of Be is high due to the completely filled s subshell.
● IE of alkali metals is low , since they have one electron in their valence shell
and easily loses it to attain octet configuration.
First IE of elements of the second period against Z
Anomalies
● The first ionisation enthalpy of boron (Z=5) is slightly less
than Be(Z=4) even though B has greater nuclear charge.
(s - electron is more attracted towards the nucleus than pelectrons)
● The first ionisation enthalpy of oxygen (Z=8) is less than
Nitrogen (Z=7).
3. Electron gain enthalpy (Electron affinity)
● It is the energy released when an electron is added to a
gaseous neutral atom to produce a negative ion.
X(g) + e-
X-(g) + Energy
● When energy is released electron gain enthalpy is
negative.
1. Electron gain enthalpy generally increases across a period due
to the decrease in the atomic size and increase in nuclear
charge.( Across a period it becomes more negative)
2. It decreases down a group due to the increase in size of the
atom. (Down a group it becomes less negative)
3. The elements of group 2, 15 and 18 has relatively very low
values of electron affinity, due to stable electronic configuration
(Completely filled and half filled outer shells).
● Chlorine has higher electron affinity than fluorine. Why?
In fluorine, the electron is added to the second shell which has 7
electrons. The added electron will suffer high repulsion from
electrons already present in the small shell. But in chlorine, an
electron is added to the third shell which is much bigger than the
second shell and hence the repulsion is less.
(Electron gain enthalpy of O is less than that of S)
● Chloride ion is more stable than bromide ion. Why?
The stability of a negative ion depends on the electron affinity.
Higher the electron affinity more stable the negative ion.
4. Valency
Valency - It is the combining capacity of an element. It
denotes the number of electrons lost or gained during a
chemical reaction.
● On moving from left to right in a period, valency of the
elements first increases from 1 to 4 and then decreases to
zero.
● All the elements in a group has same valency, as they
have same number of valence shell electrons.
5. Electronegativity - Linus Pauling
It is the tendency of the bonded molecules to attract the shared pair of
electrons more towards itself.
● Electronegativity increases across a period due to the increase in
nuclear charge and decrease in size with the increase in atomic
number in a period.
● Electronegativity decreases down a group due to the increase in size
of the atoms.
● Among the stable elements, Caesium has the lowest electronegativity
and fluorine has the highest electronegativity.
6. Metallic character:
● Metals are found on the left hand side of the table and
nonmetals on the right hand side.
● Metals has a tendency to lose electrons - electropositive.
● Nonmetals has a tendency to gain electrons electronegative.
● On moving from left to right in a period, the electropositive
character decreases and electro negative character
increases.
● On moving from left to right in a period, the tendency of
● On moving down a group the metallic character of the
elements increases or the electropositive character of the
element increases.
● As we go down in group 17, the electronegative character
(nonmetallic character) of elements decreases.
● Fluorine is the most electronegative element.
7. Nature of oxides.
● On moving from left to right in a period the basic nature of
oxides decreases and the acidic nature of oxides
increases.
● On moving down a group there is no change in the nature
of oxides of elements
● All elements of group 1 form basic oxides
● All the elements of group 17 form acidic oxides
Anomalous properties
● The first element in each group differs from the rest due to
➔ Small size, high charge to size ratio, high electronegativity of element
and non availablity of d-orbitals.
Diagonal relationship
● The similarity in properties of an element with another element lying
diagonally opposite to it in second period is referred as diagonal
relationship.
● This is due to similar charge to size ratio. Thus Li resembles Mg, Be
resembles Al.
To predict the position of an element in periodic table
1. Write the electronic configuration
2. Period - Principal quantum number
3. Block - The subshell which receives the electrons
4. Group - s block (No: of valence electron)
p block ( 10 + No: of valence electron)
d block ( 2+ No: of (n-1)d electrons)