Electronegativity and bonding

Definition (learn this exactly). Electronegativity is the power of an atom to attract the bonding pair of electrons (shared electrons) in a covalent bond towards itself.

• It is a property of an atom within a bond, not of an isolated atom.
• The most common scale is the Pauling scale, a dimensionless number — fluorine is the highest at $4.0$.
• A bigger electronegativity means the atom pulls shared electrons more strongly, making the bond polar.

Electronegativity is set by how strongly the nucleus can pull on the bonding electrons — the same three factors that control ionisation energy:

• Nuclear charge — more protons → stronger pull → higher electronegativity.
• Atomic radius — smaller atom → bonding pair closer to nucleus → higher electronegativity.
• Shielding — more inner shells → weaker net pull → lower electronegativity.

Across a period (left → right): electronegativity increases.

• Nuclear charge increases and atomic radius decreases (shielding roughly constant), so the nucleus attracts the bonding pair more strongly.

Down a group: electronegativity decreases.

• Atomic radius increases and shielding increases, outweighing the higher nuclear charge, so the bonding pair is attracted less strongly.

Overall pattern. Electronegativity is highest at the top right (excluding noble gases): the order to remember is F > O > N ≈ Cl.

The difference in electronegativity ($\Delta$EN) between two bonded atoms predicts the type of bond:

$\Delta$EN	Bond type	Example
Zero (same element)	Non-polar covalent	$\text{Cl}_2$, $\text{O}_2$
Small (roughly < 1.8)	Polar covalent	$\text{HCl}$, $\text{H}_2\text{O}$
Large (roughly > 1.8)	Ionic	$\text{NaCl}$, $\text{MgO}$

• Same electronegativity → electrons shared equally → pure (non-polar) covalent.
• Moderate difference → electrons pulled towards the more electronegative atom → polar covalent (a $\delta+$/$\delta-$ dipole).
• Very large difference → the more electronegative atom effectively takes the electrons → ionic.

Worked. Using Pauling values Na 0.9 and Cl 3.0: $\Delta\text{EN} = 2.1$ (large) → ionic (NaCl). For H 2.1 and Cl 3.0: $\Delta\text{EN} = 0.9$ (small) → polar covalent (HCl).

(Bonding is a continuum: there is no exact cut-off, so use the difference as a guide and quote the Pauling values given in the question. The presence of covalent character in ionic compounds is not assessed at 9701.)