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*Revision Materials and Past Papers* 1 Atomic Structure and the Periodic Table 2 Bonding and Structure 3 Redox I 4 Inorganic Chemistry and the Periodic Table 5 Formulae, Equations and Amounts of Substance 6 Organic Chemistry I 7 Modern Analytical Techniques I 8 Energetics I 9 Kinetics I 10 Equilibrium I 11 Equilibrium II 12 Acid-base Equilibria 13 Energetics II 14 Redox II 15 Transition Metals 16 Kinetics II 17 Organic Chemistry II 18 Organic Chemistry III 19 Modern Analytical Techniques II RP Required Practicals

2 Bonding and Structure

2.1 Ionic Bonding and Ion Formation 2.2 Covalent and Dative Bonding 2.3 Molecular Shapes and Bond Angles 2.4 Electronegativity and Bond Polarity 2.5 Intermolecular Forces and Hydrogen Bonding 2.6 Solubility and Choice of Solvents 2.7 Metallic Bonding 2.8 Structures and Physical Properties

Electronegativity and Bond Polarity

Specification Reference Topic 2, points 13–15 (Edexcel A-Level Chemistry)

Quick Notes

  • Electronegativity is an atom’s ability to attract the bonding electrons in a covalent bond.
  • Trends in electronegativity:
    • Increases across a period (more protons, same shielding)
    • Decreases down a group (more shielding and distance)
  • Ionic and covalent bonding are the extremes of a continuum of bonding type.
  • A polar bond forms when two atoms in a bond have different electronegativities.
    • More electronegative atom gains δ⁻ (partial negative charge)
    • Less electronegative atom gains δ⁺ (partial positive charge)
  • A molecule may contain polar bonds but still be non-polar overall if the dipoles cancel due to symmetry.
    • CO2 has polar bonds but is non-polar overall (linear, dipoles cancel)
    • H2O has polar bonds and is polar overall (bent, dipoles don’t cancel)

Full Notes

Electronegativity

Electronegativity is defined as the ability of an atom to attract the bonding electrons in a covalent bond.

It is measured on a relative scale called the Pauling scale, where:

Trends in Electronegativity have been covered in more detail here.

Electronegativity increases across a period and decreases down a group.

Edexcel A-Level Chemistry graph showing trends in electronegativity across periods and down groups.

Across a period: More protons in the nucleus and similar shielding means greater attraction to bonding electrons.

Down a group: Atomic radius increases and shielding increases means weaker attraction to bonding electrons.

This trend explains why atoms like oxygen and nitrogen tend to form polar bonds when joined to hydrogen or carbon.

Polar Covalent Bonds

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

The bonding electrons are pulled closer to the more electronegative atom, making the electrons unevenly shared. This creates partial charges (δ⁺ and δ⁻) at either end of the bond:

Example Hydrogen Chloride (HCl)

Edexcel A-Level Chemistry diagram of HCl showing chlorine with δ⁻ and hydrogen with δ⁺ partial charges.

Chlorine (Cl) is more electronegative than Hydrogen (H). The bonding electrons are pulled closer to Cl, giving it a partial negative charge (δ⁻). H loses electron density, giving it a partial positive charge (δ⁺).

If both atoms are the same (e.g. H2, O2), the bond is non-polar because electrons are shared equally.

Bonding Continuum

Ionic and covalent bonding are the extremes of a continuum of bonding type. In reality, most bonds exist between the two extremes.

The greater the difference in electronegativity between the bonding atoms, the more ionic character the bond has and the smaller the difference in electronegativity, the more covalent character the bond has.

Examples

Edexcel A-Level Chemistry diagram showing bonding continuum from ionic NaCl, polar covalent HCl, to non-polar covalent Cl₂.

Polar and Non-Polar Molecules

A molecule is polar or non-polar depending on its symmetry.

Non-Polar Molecules (No Permanent Dipole)

If polar bonds are arranged symmetrically then dipoles cancel out and the molecule is non-polar.

Edexcel A-Level Chemistry diagram of CO₂ molecule showing linear shape and cancellation of bond dipoles.

Example CO₂

Each C=O bond is polar, but the molecule is linear, so dipoles cancel. No overall dipole = non-polar molecule.

Example CCl₄ (Tetrachloromethane)

Each C-Cl bond is polar, but tetrahedral shape means dipoles cancel. CCl₄ is non-polar despite having polar bonds.

Polar Molecules (Have a Permanent Dipole)

If dipoles do not cancel due out, the molecule is polar.

Edexcel A-Level Chemistry diagram of H₂O and CHCl₃ molecules showing asymmetry and net dipole moments.

Example H₂O

O-H bonds are polar and form a bent shape (104.5°). Dipoles do not cancel meaning water is polar.

Example CHCl₃ (Chloroform)

The C-H and C-Cl bonds have different polarities. Dipoles do not cancel meaning CHCl₃ is polar.

Summary