Electronegativity and Bonding

Full Notes

Electronegativity has been outlined with in more detail with background theory at this page.
This page is just what you need to know for CIE A-level Chemistry :)

Definition of Electronegativity

Electronegativity is the power of an atom to attract electrons to itself. It is a relative value and does not have units.

The most commonly used scale is the Pauling scale, where:

• Fluorine (F) = 4.0 (most electronegative)
• Group 1 metals = ~0.7 to 1.0 (least electronegative)

Electronegativity values help predict whether a bond is likely to be:

• Ionic
• Polar covalent
• Non-polar covalent

Factors Affecting Electronegativity

The electronegativity of an element depends on three main factors:

• Nuclear charge: More protons in nucleus = greater attraction to bonding electrons.

• Atomic radius: Smaller atoms = bonding pair is closer to the nucleus and more strongly attracted.
• Shielding: More inner shells (and sub-shells) reduce the attraction between the nucleus and bonding electrons.

We can use these factors to explain the periodic trends in electronegativity.

Periodic Trends in Electronegativity

Electronegativity increases across a period (left to right):

• Atomic radius decreases and nuclear charge increases.
• Shielding stays roughly the same.

Electronegativity decreases down a group (top to bottom):

• Atomic radius increases and shielding increases.
• Outer electrons are further from the nucleus and less strongly attracted.

Predicting Bond Type Using Electronegativity

We can predict a bond type by comparing the Pauling electronegativity values of the two bonded atoms:

• Large Difference(typically > 1.7) = bond is ionic.
• One atom essentially takes the bonding electrons for itself.
• The atom with the higher electronegativity becomes a negatively charged ion and the atom with a lower electronegativity becomes a positively charged ion.
• Moderate Difference (between ~0.5 and 1.7) = bond is polar covalent.
• The bonding electrons are shared unequally.
• The atom with the higher electronegativity has a partial negative charge (δ⁻) and the other atom a partial positive (δ⁺).

• Small or zero difference = the bond is non-polar covalent.
• The bonding electrons are shared equally.

Matt’s exam tip

Always remember it is the difference in electronegativity between two bonding atoms that matters when determining whether a bond will be ionic, polar covalent or non-polar covalent.

For Example:

• NaCl: large difference → ionic
• HCl: moderate difference → polar covalent
• Cl₂: no difference → non-polar covalent

Summary

• Electronegativity: the power of an atom to attract electrons to itself.
• Depends on nuclear charge, atomic radius, and shielding.
• Increases across a period, decreases down a group.
• Bond type is determined by the difference in electronegativity:
  • Large difference → ionic
  • Moderate difference → polar covalent
  • No difference → non-polar covalent