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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)

Covalent Bonds and Lewis Formulas

Specification Reference S2.2.1

Quick Notes

  • A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
  • The octet rule: atoms tend to have 8 electrons in their valence shell (like noble gases).
  • Lewis structures (or electron dot structures) show:
    • Bonding pairs (shared electrons)
    • Lone pairs (non-bonding electrons)
  • Electron pairs can be represented using dots, crosses, or dashes.

Full Notes

What Is a Covalent Bond?

In a covalent bond, a pair of electrons is shared between two atoms.

The positively charged nuclei of both atoms get attracted to the shared pair of negatively charged electrons and this pulls them in closer together.

IB Chemistry diagram of covalent bonding showing nuclei attracting a shared pair of electrons.

The bond results from electrostatic attraction between the shared electrons and the positive nuclei of both atoms.

The Octet Rule

The octet rule describes how atoms (especially in Period 2) tend to gain, lose, or share electrons to achieve 8 electrons in their outer shell.

Note there are exceptions to this rule (such as for larger atoms).

Lewis Structures (Electron Dot Structures)

Lewis structures are diagrams used to how electrons are involved in the bonding between atoms. They show:

Electrons are represented as Dots (●).

Steps to draw a Lewis structure:

  1. Count total valence electrons around each atom.
  2. Connect atoms with single bonds (1 shared pair = 2 electrons).
  3. Distribute remaining electrons to complete octets.
  4. Use double or triple bonds if needed to meet the octet rule.

Examples of Lewis Structures

Molecules with single bonds (1 shared pair):

IB Chemistry examples of Lewis structures with single bonds including H2, Cl2, HCl, CH4, NH3 and C2H6.

Molecules with double bonds (2 shared pairs):

IB Chemistry examples of Lewis structures with double bonds including O2, CO2 and C2H4.

Lewis Structures with Fewer Than an Octet

Some atoms (mainly in Period 2) can be stable with less than 8 electrons.

For example:

IB Chemistry examples of Lewis structures with incomplete octets including BeCl2 and BF3.

These are exceptions to the octet rule and should be recognised. Sometimes, bonded atoms can have more than 8 electrons in their outer shell (such as SF6), this is referred to as an expanded octet.

Summary