### Theory:

Electronegativity:
An element's electronegativity measures its atom's tendency to attract the shared pair of electrons towards itself in a covalent bond.
• Consider the $$HCl$$ molecule. The hydrogen and chlorine atoms give one electron each to form a covalent bond.
• A chlorine atom has a higher electronegativity, and hence it pulls the shared electrons towards itself more strongly than hydrogen. Thus, the bonding electrons are left with chlorine, forming $$H^+$$ and $$Cl^–$$ ions when the bond breaks.
Relative electronegativity of $$H$$ and $$Cl$$.

• Electronegativity is determined by various experimental data, including bond energy, ionisation potential, electron affinity, etc.
• The Pauling scale is the most used scale for determining electronegativity, which predicts the type of bonding (ionic or covalent) between atoms in a molecule.
Some of the elements' electronegativity is listed below.
F = $$4.0$$, Cl = $$3.0$$, Br = $$2.8$$, I = $$2.5$$, H = $$2.1$$, Na = $$1$$
• If the electronegativity difference between two elements is $$1.7$$, then the bond has $$50$$% ionic character and $$50$$% covalent character.
• The bond is considered more covalent if the difference is less than $$1.7$$.
• If the difference is more than $$1.7$$, the bond is considered more ionic.
• From left to right in the periodic table, the electronegativity increases due to an increase in nuclear charge, which in turn attracts electrons more strongly.
• Conversely, on moving down a group, the electronegativity of the elements decreases because of the increased number of valence shells.
 Periodic Property In Periods In Groups Atomic radius Decreases Increases Ionic radius Decreases Increases Ionisation energy Increases Decreases Electron affinity Increases Decreases Electronegativity Increases Decreases