Lewis Structures

Full Notes

A Lewis structure (also called a Lewis dot diagram) represents the arrangement of valence electrons in a molecule. It helps us understand bonding, molecular shape, and formal charges.

Purpose of Lewis Structures

• Show which atoms are bonded
• Show how many bonds exist between atoms (single, double, or triple)
• Indicate lone pairs of electrons not involved in bonding

These diagrams are foundational for predicting geometry, polarity, and reactivity.

General Rules for Drawing Lewis Structures

• Count total valence electrons
  • Add up all valence electrons from each atom
  • For polyatomic ions:
    • Add 1 electron for each negative charge
    • Subtract 1 electron for each positive charge
• Determine central atom
  • Usually the least electronegative (never hydrogen)
  • Hydrogen and halogens are usually terminal atoms
• Form single bonds between central and outer atoms
  • Each bond = 2 electrons
• Distribute remaining electrons to complete octets
  • Start with outer atoms first
  • Place any leftover electrons on the central atom
• Form double or triple bonds if needed
  • If central atom does not have an octet, convert lone pairs from outer atoms into bonding pairs

Special Cases

• Hydrogen: Only 2 electrons (1 bond)
• Boron and Beryllium: Often stable with fewer than 8 electrons
• Expanded octets: Possible for elements in period 3 or beyond (e.g. P, S, Cl)

Examples of Lewis Structures

Molecules with Single bonds (1 shared pair):

• H₂: each H shares 1 electron → H–H
• Cl₂: each Cl shares 1 electron → Cl–Cl
• HCl: H and Cl share 1 electron → H–Cl
• CH₄: C shares 4 electrons with 4 H atoms → 4 single bonds
• NH₃: N shares 3 electrons with 3 H atoms
• C₂H₆ (ethane): all single C–C and C–H bonds

Molecules with Double bonds (2 shared pairs):

• O₂: O=O → each O shares 2 electrons
• CO₂: O=C=O → two double bonds
• C₂H₄ (ethene): double bond between two carbon atoms

Lewis Structures with Fewer Than an Octet

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

• BeCl₂: Be forms 2 bonds, giving a total of 4 outer electrons
• BF₃: B forms 3 bonds, giving a total of 6 outer electrons

Expanded Octets

Some elements (usually in Period 3 or below) can hold more than 8 electrons. These are exceptions to the octet rule and are referred to as an expanded octet.

Examples:

• SO₂: sulfur forms 2 double bonds with O and expands its octet to 10 electrons
• PCl₅: phosphorus has 5 bonding pairs = 10 electrons
• SF₆: sulfur has 6 bonding pairs = 12 electrons Show all bonding pairs in dot-and-cross form.

Common Errors to Avoid

• Giving hydrogen more than 2 electrons
• Forgetting to subtract/add electrons for ions
• Forgetting to check for octets on all atoms
• Misplacing double or triple bonds (especially in O, N, and C-containing molecules)

Matt’s exam tip

Always count your total electrons before and after drawing your structure.

Summary

Lewis structures are a visual tool to represent how valence electrons are distributed in molecules or ions. Following a step-by-step process ensures that all atoms achieve a stable electron configuration (usually an octet), and the structure reflects bonding accurately. Lewis diagrams are essential for predicting molecular geometry, resonance, polarity, and reactivity.

Key points to remember:

• Count total valence electrons carefully
• Connect atoms using single bonds, then distribute lone pairs
• Use double/triple bonds if needed to complete octets
• Hydrogen gets 2 electrons; period 3+ elements can expand octets
• Ions require electron adjustments based on charge