Classification Of Elements And Periodicity In Properties Class 11 Notes Chemistry Chapter 3 - CBSE

Chapter : 3

What Are Classification Of Elements And Periodicity In Properties ?

Genesis Of Periodic Classification Dobereiner’s Triads

In 1829, Dobereiner arranged certain elements with similar properties in groups of three in such a way that the atomic mass of the middle element was nearly the same as the average atomic masses of the first and the third elements. A few triads proposed by him are listed:

Element Li Na K Ca Sr Ba Cl Br I
Atomic Weight 7 23 39 40 88 137 35.5 80 127

Limitations of Dobereiner’s Triads

Dobereiner’s triads were helpful in grouping some elements with similar characteristics together, but he could not arrange all the elements known at that time into triads.

Newlands’ Law Of Octaves

It states that when elements are arranged in order of increasing atomic masses, every eighth element has properties similar to the first. Newlands called it law of octaves because similar relationship exists in the
musical notes also. This can be illustrated as:

sa (do) re (re) ga (mi) ma (fa) pa (so) da (la) ni (ti)
H Li Be B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe
Co and Nl Cu Zn Y In As Se
Br Rb Sr Ce and La Zr - -

Limitations of Dobereiner’s Triads

  • This classification was successful only up to the element calcium. After that, every eighth element did not possess the same properties as the element lying above it in the same group.
  • The inclusion of noble gases (discovered at later stage) in the table disturbed the entire arrangement.

Mendeleev’s Periodic Table

Mendeleev’s Periodic Law

The physical and chemical properties of the elements are a periodic function of their atomic masses. Mendeleev arranged the elements known at that time in order of increasing atomic masses and this arrangement was called periodic table. Elements with similar characteristics were present in vertical rows called groups. The horizontal rows were known as periods.

Description of Mendeleev’s Periodic Table

  • In the periodic table, the elements are arranged in vertical rows called groups and horizontal rows known as
  • There are nine groups indicated by Roman Numerals as I, II, III, IV, V, VI, VII, VIII and zero. Group VIII consists of nine elements which are arranged in three triads. The zero group contains elements belonging to inert gases or noble gases and elements present have zero valency.
  • There are seven periods (numbered from 1 to 7) or, horizontal rows in the Mendeleev’s periodic table.

Importance of Mendeleev’s Periodic Table

  • This made the study of the elements quite systematic, i.e., if the properties of one element in a particular group are known, those of others can be predicted.
  • It helped to a great extent in the discovery of these elements at a later stage.
  • Mendeleev corrected the atomic masses of certain elements with the help of their expected positions and

Defects in Mendeleev’s Periodic Table

  • The position of hydrogen in the Mendeleev’s periodic table was not justified. Hydrogen has been placed in group IA along with alkali metals. But it also resembles halogens of group VII A in many properties.
  • Although the elements in the Mendeleev’s periodic table was arranged in order of their atomic masses, but in some cases the element with higher atomic mass precedes the element with lower atomic mass.
  • Since, periodic table has been framed on the basis of increasing atomic masses of the elements, different positions must have been allotted to all the isotopes of a particular element.
  • According to Mendeleev, the elements placed in the same group must resemble in their properties. But there is no similarity among the elements in the two sub-groups of a particular group.
  • In some cases, elements with similar properties have been placed in different groups.
  • Lanthanoids and actinoids were placed in two separate rows at the bottom of the periodic table without
    assigning a proper reason.
  • No proper explanation was given that why the elements placed in group show resemblance in their properties.

Modern Periodic Law

The physical and chemical properties of the elements are the periodic function of their atomic numbers. The long form of periodic table, also called Modem Periodic Table, is based on Modern periodic law. In this table,
the elements have been arranged in order of increasing atomic numbers.

Structural Features Of The Periodic Table


The long form of periodic table also consists of the vertical rows called groups. There are in all 18 groups in the periodic table. Unlike Mendeleev periodic table, each group is an independent group.

Characteristics of Groups

  • All the elements present in a group have same general electronic configuration of the atoms.
  • The elements in a group are separated by definite gaps of atomic numbers (2, 8, 8, 18, 18, 32).
  • The atomic sizes of the elements in group increase down the group due to increase the number of shells.
  • The physical properties of the elements such as m.p., b.p. density, solubility etc., follow a systematic pattern.
  • The elements in each group have generally similar chemical properties.


Horizontal rows in a periodic table are known as periods. There are in all seven periods in the long form of periodic table.

Characteristics of Periods

  • In all the elements present in a period, the electrons are filled in the same valence shell.
  • The atomic sizes generally decrease from left to right.

s-Block Elements

General electronic configuration: ns1-2

Characteristics of s-block elements

  • All the elements are soft metals.
  • They have low melting and boiling points.
  • They are highly reactive.
  • Most of them impart colours to the flame.
  • They generally form ionic compounds.
  • They are good conductors of heat and electricity.


General electronic configuration: ns2np1-6

Characteristics of p-block elements

  • The compounds of these elements are mostly covalent in nature.
  • They show variable oxidation states.
  • In moving from left to right in a period, the non-metallic character of the elements increases.
  • The reactivity of elements in a group generally decreases downwards.
  • At the end of each period is a noble gas element with a closed valence shell ns2 np6 configuration.
  • Metallic character increases as we go down the group


General electronic configuration: (n-1) d1-10 ns0-2. The d-block elements are known as transition elements because they have incompletely filled d-orbitals in their ground state or in any of the oxidation states.

Characteristics of d-block elements

  • They are all metals with high melting and boiling points.
  • The compounds of the elements are generally paramagnetic in nature.
  • They mostly form coloured ions, exhibit variable valence (oxidation states).


General electronic configuration: (n-2) f1-14 (n-1) d0-1 ns2 . They are known as inner transition elements because in the transition elements of d-block, the electrons are filled in (n-1) d sub-shell while in the inner transition elements of f-block the filling of electrons takes place in (n – 2) f sub-shell, i.e., inner subshell.

Characteristics of f-block elements

  • The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids Ce (Z=58)- Lu (Z=71) and Actinoids Th (Z=90) -Lr (Z=103).
  • These two series of elements are called Inner Transition Elements (f-Block Elements).
  • They are all metals. Within each series, the properties of the elements are quite similar.
  • Most of the elements of the actinoid series are radioactive in nature.


  • Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table.
  • Metals are solids at room temperature.
  • Metal usually have high melting and boiling points.
  • They are good conductors of heat and electricity.
  • They are malleable and ductile.


  • Non-metals are located at the top right hand side of the Periodic Table.
  • Non-metals are usually solids or gases at low temperature with low melting and boiling points.
  • They are poor conductors of heat and electricity.
  • The non-metallic character increases as one goes from left to right across the Periodic Table.
  • Most non-metallic solids are brittle and are neither malleable nor ductile.


The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) show the characteristic, of both metals and non-metals. These elements are also called semi-metal.

Noble Gases

  • These are the elements present in group 18.
  • Each period ends with noble gas element.
  • All the members are of gaseous nature and because of the presence of all the occupied filled orbitals, they have very little tendency to take part in chemical combination.
  • These are also called inert gases.

Representative Elements

These are the elements of group 1 (alkali metals), group 2 (alkaline earth metals) and group 13 to 17 constitute the representative elements. They are elements of s-block and p-block.

  • Transition Elements include all the d-block elements and they are present in the centre of the periodic table between s and p-block elements.
  • Lanthanoids (the fourteen elements after Lanthanum) and actinides (the fourteen elements after actinium) are called inner transition elements. They are also called f-block elements.
  • The elements after uranium are also called transuranic elements.

Periodic Trends In Properties Of Elements

Trends in Physical Properties

Atomic Radii

It is defined as the distance from the centre of the nucleus to the outermost shell containing the electrons. Depending upon whether an element is a non-metal or a metal, three different types of atomic radii are used. These are:

  • Covalent Radius: It is equal to half of the distance between the centres of the nuclei of two atoms held together by a purely covalent single bond.
  • Ionic Radius: It may be defined as the effective distance from the nucleus of an ion up to which it has an influence in the ionic bond.
  • Van der Waal’s Radius: Atoms of Noble gases are held together by weak Van der Waal’s forces of attraction. The Van der Waal’s radius is half of
    the distance between the centre of nuclei of atoms of noble gases.
  • Metallic Radius: It is defined as half of the internuclear distance between the two adjacent metal ions in the metallic lattice.


The removal of an electron from an atom results in the formation of a cation. The radius of cation is always smaller than that of the atom.


Gain of an electron leads to an anion. The radius of the anion is always larger than that of the atom.

Isoelectronic Species

These are atoms and ions which contain the same number of electrons. For example, O2-, F , Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges.

Ionisation Enthalpy

It is the energy required to remove an electron from an isolated gaseous atom in its ground state.

$$\text{M(g) + I.E}\xrightarrow{}\text{M}^{+}(g) + e^{\normalsize-}$$

The unit of ionization enthalpy is kJ mol-1 and the unit of ionisation potential is electron volt per atom.

Variation Of Atomic Radius In The Periodic Table

Along a period, the atomic radii of the elements generally decrease from left to right. On moving down the group, the atomic radii of the elements in every group of the periodic table increases. The ionic radii can be estimated by measuring the distances between cations and anion in ionic crystals. In general, Ionic radius increases as you move from top to bottom on the periodic table. Ionic radius decreases as you move across the periodic table, from left to right.

Electron Gain Enthalpy

It is the energy released when an electron is added to an isolated gaseous atom so as to convert it in to a negative ion. The electron gain enthalpy of some elements are listed in table below:

Electron Gain Enthalpy kJ mol-1
Fluorine -333
Chlorine -348
Bromine -324
Iodine -295
Astatine -270.1

Periodic Trends In Chemical Properties Along A Period

  • Metallic Character: Decrease across a period maximum on the extreme left (alkali metals).
  • Non-metallic Character: Increases from left to right in a period.
  • Basic nature of oxides: Decreases from left to right in a period.
  • Acidic nature of oxides: Increases from left to right in a period.

Variation From Top To Bottom On Moving Down A Group

  • Generally, metallic character increases because of increase in atomic size and hence there is a decrease in the ionisation energy of the elements in a group from top to bottom.
  • Generally, non-metallic character, decreases down a group as electronegativity of elements decreases from top to bottom in a group.
  • Since metallic character or electropositivity of elements increases in going from top to bottom in a group basic nature of oxidise naturally increases.
  • Generally, acidic character of oxides decreases as non-metallic character of elements decreases in going from top to bottom in a group.
  • Generally, reactivity of metals increases down a group as tendency to lose electron increases.
  • Generally, reactivity of non-metals. decreases down the group, Higher the electro-negativity of non-metals,
    greater is their reactivity. Since electronegativity of non-metals in a group decreases from top to bottom,
    their reactivity also decreases.


It is a qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself. Unlike ionization enthalpy and electron gain enthalpy, it is not a measurable quantity. However, a number of numerical scales of electronegativity of elements viz, Pauling scale, Milliken- Jaffe scale, Allred Kochow scale have been developed. The electronegativity of any given element is not constant; it varies depending on the element to which it is bound.

  • Across a Period: Electronegativity generally increases across a period from left to right.
  • In a Group: It decreases down a group.

Anomalous Properties Of Second Period Elements

The first element of each of the group 1 (lithium) and 2 (beryllium) and group 13-17 (boron to fluorine) differs in many respect from the other members of their respective groups. For example, lithium unlike other
alkali metals, and beryllium unlike other alkaline earth metals form compounds which have significant covalent character; the other members of these groups, predominantly form ionic compounds.

It has been observed that some elements of the second period show similarities with the elements of the third period placed diagonally to each other, though belonging to different groups.

For example,


This similarity in properties of elements placed diagonally to each other is called diagonal relationship.