AP | A-Level | IB | NCERT 11 + 12 – FREE NOTES, RESOURCES AND VIDEOS!
S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

S2.2 - The covalent model

2.2.1 Covalent Bonds and Lewis Formulas 2.2.2 Bond Types 2.2.3 Co-coordination (Dative) Bonds 2.2.4 VSEPR Shapes of Molecules 2.2.5 Electronegativity and Bond Polarity 2.2.6 Polarity and Dipole Moments 2.2.7 Covalent Network Structures and Allotropes 2.2.8 Intermolecular Forces 2.2.9 Physical Properties of Covalent Substances 2.2.10 Chromatography and Intermolecular Forces 2.2.11 Resonance Structures (AHL) 2.2.12 Benzene and Resonance (AHL) 2.2.13 Expanded Octet and VSEPR (AHL) 2.2.14 Formal Charge (AHL) 2.2.15 Sigma and Pi Bonds (AHL) 2.2.16 Hybridization (AHL)

Electronegativity and Bond Polarity

Specification Reference S2.2.5

Quick Notes

  • Electronegativity is an atom’s ability to attract the shared electrons in a covalent bond.
  • A polar bond occurs when there is a difference in electronegativity between the bonded atoms.
  • The greater the difference, the more polar the bond.
  • We can use electronegativity values (from the IB data booklet) to:
    • Determine if a bond is non-polar or polar
    • Identify the direction of the bond dipole
  • Bond dipoles can be shown using:
    • δ⁺ / δ⁻ partial charges
    • Vector arrows pointing toward the more electronegative atom

Full Notes:

What Is Electronegativity?

Electronegativity is a measure of how strongly an atom attracts shared electrons in a covalent bond.

Example: Fluorine (F) has the highest electronegativity (≈ 4.0). Cesium (Cs) has a very low electronegativity.

When Is a Bond Polar?

A bond is polar when there is a difference in electronegativity between the two bonded atoms.

IB Chemistry diagram showing polar bond with unequal electron sharing and partial charges δ⁺ and δ⁻.

The more electronegative atom becomes δ⁻ and the less electronegative atom becomes δ⁺

Example:Hydrogen Chloride (HCl):

IB Chemistry diagram showing bond polarity in HCl with electron density pulled toward chlorine.

How to Deduce Bond Polarity

Follow these steps:

  1. Look up the electronegativity values in the data booklet.
  2. Subtract to find the difference (ΔEN).
  3. Interpret the result:
ΔEN (Electronegativity Difference) Bond Type Explanation
> 1.7 Ionic One atom takes the bonding electrons completely, forming ions.
0.5 – 1.7 Polar Covalent Electrons shared unequally, producing δ⁺ and δ⁻ charges.
0 – 0.4 Non-polar Covalent Electrons shared equally between atoms.

If the difference is large (typically > 1.7), the 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.

If the difference is moderate (between ~0.5 and 1.7), the 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⁺).

If the difference is very small or zero, the bond is non-polar covalent.
The bonding electrons are shared equally.

Photo of Matt
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.

Examples:

IB Chemistry diagram showing comparison of ionic, polar covalent, and non-polar covalent bonding using NaCl, HCl, and Cl₂.

Representing Bond Polarity

We can show bond dipoles in two ways:

IB Chemistry diagram showing methods of representing bond polarity using δ⁺/δ⁻ notation and vector arrows.

Summary