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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)

Bond Types: Single, Double, and Triple Bonds

Specification Reference S2.2.2

Quick Notes

  • Single bond = 1 shared pair of electrons
  • Double bond = 2 shared pairs
  • Triple bond = 3 shared pairs
  • As the number of shared pairs increases:
    • Bond length decreases
    • Bond strength increases
  • Therefore:
    • Triple bonds are shorter and stronger than double bonds
    • Double bonds are shorter and stronger than single bonds

Full Notes

Types of Covalent Bonds

As introduced in S2.2.1, covalent bonds involve the sharing of electron pairs between atoms.

Sometimes, two atoms will share more than one pair of electrons, forming double and triple bonds.

IB Chemistry diagram showing examples of single, double, and triple bonds with Cl2, O2, and N2 molecules.

These are referred to as bond orders:

Bond Length and Bond Strength

There is a clear relationship between the bond order and the physical properties of the bond:

Bond Type Bond Order Bond Length Bond Strength
Single 1 Longest Weakest
Double 2 Shorter Stronger
Triple 3 Shortest Strongest

More shared electrons = stronger attraction between the shared electrons (negatively charged) and positively charged nuclei of each atom.

IB Chemistry diagram showing how more shared electron pairs result in stronger attraction and shorter bond lengths.

This pulls the nuclei closer together, shortening the bond. More energy is required to break a stronger bond.

For Example: Carbon Bonding

IB Chemistry diagram showing carbon-carbon single, double, and triple bonds in different organic molecules.
Bond Bond Length (nm) Bond Energy (kJ mol⁻¹)
C–C single bond 0.154 348
C=C double bond 0.134 612
C≡C triple bond 0.120 837

This pattern is consistent across different types of covalent bonds, including C–C, C–N, and C–O.

Summary