Lewis Structures for Ions and Ionic Compounds

For ions the charge is always shown. Thus for metal ions such as calcium the Lewis Structure simply becomes the symbol for the ion. For negative ions such as we see for oxygen (2-) we enclose the ion and its electrons in brackets to indicate that the electrons are all "owned" by the oxygen - it does not share. Notice that the Lewis Structures of monoatomic ions are isoelectronic with the nearest Noble gas. Thus sodium loses an electron to leave a kernel isoelectronic with neon, whereas bromine gains an electron to become isoelectronic with krypton. Examples:

• Sodium ion: Na⁺
• Bromide ion: We need the brackets to show that the bromide ion "owns" all of the electrons rather than sharing them, and that the charge is distributed over the entire structure - it is not associated with any particular electron or locale.

Brackets are particularly important when we make ionic compounds:

• Sodium chloride
• Potassium bromide
• Aluminum chloride
  or Al³⁺ 3
• Barium iodide

Covalent Bonds

Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons.

Let's look at the formation of HCl as an example of the creation of a covalent bond:

H₂ + Cl₂ 2 HCl

In this case can consider that we get two equations each involving a homo dissociation to give radicals, that is atoms with unpaired electrons:

• H₂ 2 H^.
• Cl₂ 2 Cl^.These radical then combine to form a bond with these two electrons shared between the two atoms.
• H^. + Cl^. H:Cl (This is not a proper Lewis structure, I have only shown the bonding pair of electrons.)
• Lewis Structure:

Note Lewis Structure homework set in Discussion Manual.

Lewis Structures for Covalent Molecules

Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons.

• Covalent Compound Lewis Structure Examples:
  • Water: first write formula = H₂O, then determine central atom, O, then add electrons to get octet and draw L.S.
  • Ammonia = NH₃
  • Ammonium ion = NH₄⁺

  Extra "practice" structures below:

  • Methane (CH₄)
  • Hydrogen sulfide
  • Arsenic triiodide
  • Hydrogen selenide
  • Oxygen difluoride

Multiple Bonds & Resonance

Recall we must show an octet (or duet for Period I) in the outer-most shell (valence electrons). When this does not occur with single electron pairs (bonds) between atoms can sometimes make it happen with multiple bonds. You might find "Clark's Method" useful for determining the bonding patterns of various molecules:

• Clark's Method (abbreviated) for determining bonding in covalent Lewis Structures:

  • Add up all of the valence electrons in the structure (remember to add one electron for each negative charge, or subtract one for each positive charge)
    • If e^- = 6y + 2 where y = # atoms other than H, then octet rule is followed with single bonds only.
    • If e^- < 6y + 2 then probably have multiple bonding with the number of multiple bonds = /2 (remember a triple bond is 2 multiple bonds!).
    • If e^- > 6y + 2 then have an expanded valence shell.
      • If you can draw more than one structure, then chose the most symmetrical.
        • If two or more structures are equally symmetrical, then you probably have resonance and should show all structures connected by double arrows.
  • Multiple bond examples:
    • Carbon monoxide, CO
      • valence electrons = 4 + 6 = 10
      • 6y + 2 = 14, thus 4 fewer electrons than required for all single bonds, 4/2 = 2 multi-bonds (2 double or 1 triple)
      • LS = :C:::O:
    • Carbon dioxide, CO₂
      • valence electrons = 4 + 2x6 = 16
      • 6y + 2 = 20, thus 4 fewer electrons than required for all single bonds, 4/2 = 2 multi-bonds (2 double or 1 triple)
      • LS: from symmetry C will be central atom, therefore= :O::C::O:

    Practice example:

    • Carbon disulfide
      • Hint: sulfur is immediately below oxygen in Periodic Table, so it should be similar to carbon dioxide.
  • Resonance example:
    • Carbonate ion - CO₃^2-
      • valence electrons = 4 + 3 (6) + 2 = 24
      • 6y + 2 = 26, but e^- = 24, therefore expect one multiple bond.
      • LS =
      • However, other equally symmetrical structures are possible, so:

• Expanded Valence Shell Example:
  • SF₄
  • valence electrons = 6 + 4(7) = 30
  • 6y + 2 = 32, but e^- = 30, therefore expect expanded valence shell with one extra electron pair.
  • LS =

Additional exercises on Lewis Structures are available in the Lewis Structure Module.

For a modern view of bonding illustrated with QuickTime movies based on quantum calculations you may enjoy the Supplement.

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© R A Paselk

Last modified 27 March 2009