Revision: Classification of Elements and Periodicity in Properties Chemistry Science (English Medium) Class 11 CBSE

It is the distance between the centre of the nucleus of an atom and its outermost shell.

The ionic radius is the average distance between the center of nucleus of an ion and its outermost shell containing electron.

The minimum amount of energy required to remove an electron from the isolated gaseous atom is called ionisation enthalpy.

The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is called its electronegativity.

• The modern periodic law states that the physical and chemical properties of elements are periodic functions of their atomic numbers.
• Atomic number is more fundamental than atomic mass.
• The modern periodic table is based on electronic configuration.
• It explains the periodic repetition of properties after regular intervals.
• Elements are arranged in increasing order of atomic number.
• This law removed the defects of Mendeleev’s periodic table.

• Classification arranges elements systematically according to similar properties.
• It makes the study of a large number of elements easier.
• Elements with similar chemical behaviour are placed in the same group.
• It helps in predicting properties of unknown or newly discovered elements.
• It shows periodic trends such as atomic size, ionisation enthalpy and electronegativity.
• Modern classification is based on atomic number and electronic configuration.

• Early classification was based on atomic masses and similar properties.
• Dobereiner arranged elements in triads.
• Newlands proposed the Law of Octaves.
• Mendeleev arranged elements in order of increasing atomic mass.
• The modern periodic table arranges elements according to increasing atomic number.
• The modern periodic table removed many defects of older classifications.

• Includes Group 1 and Group 2 elements.
• Group 1 elements are alkali metals.
• Group 2 elements are alkaline earth metals.
• General configuration: ns¹ or ns².
• They are highly electropositive and metallic.
• They form basic oxides and ionic compounds.
• They have low ionisation enthalpy.
• Reactivity generally increases down the group.

• Includes Groups 13 to 18.
• General configuration: ns²np¹ to ns²np⁶.
• Contains metals, non-metals and metalloids.
• Group 17 elements are halogens.
• Group 18 elements are noble gases.
• Shows variable oxidation states.
• Metallic character decreases from left to right.
• Non-metallic character increases from left to right.

• Elements of Groups 3 to 12 are d-block elements.
• General configuration: (n−1)d¹⁻¹⁰ns⁰⁻².
• They are called transition elements.
• They show variable oxidation states.
• They form coloured compounds.
• They often act as catalysts.
• They form complex compounds.
• They are generally hard metals with high melting points.

• The f-block includes lanthanoids and actinoids.
• Last electron enters the f-orbital.
• General configuration involves (n−2) f orbitals.
• Lanthanoids have atomic numbers 58–71.
• Actinoids have atomic numbers 90–103.
• They show variable oxidation states.
• Actinoids are mostly radioactive.
• They are placed separately to keep the periodic table compact.

• Metals are generally present on the left side of the periodic table.
• Non-metals are mainly on the right side.
• Metalloids lie along the zig-zag line.
• Metals lose electrons and form cations.
• Non-metals gain electrons and form anions.
• Metallic character decreases across a period.
• Metallic character increases down a group.
• Metalloids show intermediate properties.

• Atomic radius is the distance from the nucleus to the outermost shell of an atom.
• Across a period, atomic size decreases due to increase in effective nuclear charge.
• Down a group, atomic size increases due to addition of new shells.
• Atomic size depends on number of shells and nuclear charge.
• Cations are smaller than their parent atoms due to loss of electrons.
• Anions are larger than their parent atoms due to gain of electrons.
• In isoelectronic species, greater nuclear charge leads to smaller size.
• Noble gases show larger atomic size due to consideration of van der Waals radius.

Factors affecting ionisation enthalpy:

• Electronegativity is the ability of an atom to pull shared electrons; it is highest for fluorine (4.0).
• It increases across a period (left to right) and decreases down a group (top to bottom).
• Non-metals have high electronegativity (gain electrons), while metals have low electronegativity (lose electrons).
• Greater electronegativity difference increases ionic character.