Simple Rules for drawing Lewis Structures for Molecules. Mr. Citta – Chemistry

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I All Single Bonds (Rules #1-4) Example: Draw a Lewis structure for iodomethane CH3I.

  1. Count the total number of valence electrons for the molecule. (Multiply the number of valence electrons in each atom by the number of atoms present in the molecule, it helps to draw Lewis dot notation for each atom present.)

CH3I = 1 3 1

Type of atom # of atoms x # of V.E. = # of Valence electrons due to each atom in molecule

C = 1 x 4 = 4 valence electrons from carbon

H = 3 x 1 = 3 valence electrons from hydrogen

I = 1 x 7 = 7 valence electrons from iodine

CH3I = 14 valence electrons total in molecule of CH3I.

  1. Arrange the atoms to form a skeleton structure for the molecule. The least electronegative element present is the central atom in the structure, except if:

i) ** hydrogen (if present) is never the central atom – it can only be connected by a single bond.

ii) ** carbon (if present) is always the central atom – it usually connects with four bonds.

Then connect all atoms by electron-pair bond. (--) or ( )

Preferred (lines represent bonds – electron pair) Alternative Version – (Dots as Bonds)

  1. Add unshared pairs so that each non-metal other than hydrogen is surrounded by eight electrons (octet rule). (remember non-metals only!!, notmetals or metalloids!! Never hydrogen!! (only has 1 bond pair)

(sp3 hybridization = 4 molecular orbitals)

Preferred (lines represent bonds – electron pair) Alternative Version – (Dots as Bonds)

  1. Count the electrons in the structure to be sure that the number of valence electrons used equals the number of valence electrons originally available from step #1.

If number of valence electrons in structure are equal to those available in step #1, you are done.

Example: 14 v.e. in structure in step #3 = 14 v.e. available from step #1

(rules 1-4 are complete if compound has all single bonds, octet rule is not exceeded.)

  1. If you have too many valence electrons in your structure. (Multiple bonds are present), follow steps #5 a+b. (Back Page)
  2. If you have too few valence electrons in your structure. (octet rule is exceeded, follow step #6)

(On to Back Page)

II Multiple Bonds (rules #1-4 followed by rule #5)

  1. Subtract two unshared electron pairs from two adjacent atoms (one each) and place a 2nd bond (double bond) between those two atoms.
  1. Repeat step until valence electrons in structure equal # of valence electrons available from step #1.
  2. Check the formal charge of each atom. (see instructions for formal charge assignment)

III Exceeding the Octet Rule- (8 electrons surrounding central atom)(rules #1-4 followed by rule #6)

– (only occurs for elements with a “d” orbital – period 3 up to period 7)

(**never happens for C, N, O, F – period 2 elements as they only have

(sp3d hybridization) or (sp3d2 hybridization) “s” and “p” orbitals available, 8 electrons maximum)

(5 orbitals) (6 orbitals)

(10 electrons) (12 electrons)

  1. Add electron in pairs to the central atom (d orbitals are used), until the number of valence electrons in your structure are equal to the number of valence electrons available in step #1.
  1. If number of electrons available are odd, the structure is a free radical (single electrons), add or subtract a single electron to make the number correct.
  2. Check the formal charge of each atom, to make sure structure is correct.

IV Notes on polyatomic ion Lewis Structures.

a.)Polyatomic cations – subtract electrons from your initial valence electron count to satisfy the charge present.

Ex. A charge of +2 has 2 less valence electrons in Lewis structure.

b.)Polyatomic anions -add extra electrons from your initial valence electron count to satisfy the charge present.

Ex. A charge of -3 has 3 more valence electrons in Lewis structure.

**remember to always check your Lewis structure for correctness by assigning formal charges**