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Rules for drawing Lewis dot structures

1. Count the number of valence e^- each atom brings into the molecule. For ions, the charge must be taken into account.

   How many valence electrons in BeCl₂?

   

   How many valence electrons in NO₂^- and NO₂⁺?

   

   
   
   
   
   
2. Put electron pairs about each atom such that there are 8 electrons around each atom (octet rule), with the exception of H, which is only surrounded by 2 electrons. Sometimes it's necessary to form double and triple bonds. Only C, N, O, P and S (rarely Cl) will form multiple bonds.

   Draw the Lewis dot structure for CF₄.

   The number of valence electrons is 4 + 4 ( 7 ) = 32 electrons.

   So, we obtain:

   

   Draw the Lewis dot structure for CO.

   The number of valence electrons is 4 + 6 = 10 electrons or 5 pairs. Since both C and O allow multiple bonds we can still follow the octet and write:

   

3. If there is not enough electrons to follow the octet rule, then the least electronegative atom is left short of electrons.

   Draw the Lewis dot structure for BeF₂.

   In BeF₂ number of valence e^- = 2+ 2(7) = 16 e^- or 8 pairs. Since neither Be or F form multiple bonds readily and Be is least electronegative we obtain:

   

4. If there are too many electrons to follow the octet rule, then the extra electrons are placed on the central atom.

   Draw the Lewis dot structure for SF₄.

   In SF₄ the number of valence electrons is 6 + 4 ( 7 ) = 34 electrons or 17 pairs. Placing the extra electrons on S we obtain:

   

How can the octet rule be violated in this last example? The octet rule arises because the s and p orbitals can take on up to 8 electrons. However, once we reach the third row of elements in the periodic table we also have d-orbitals, and these orbitals help take the extra electrons. Note that you still need to know how the atoms are connected in a polyatomic molecule before using the Lewis-Dot structure rules.