Writing Lewis Structures

This is how we do it:

• Draw skeletal structure of compound showing what atoms are bonded together.  Put the least electronegative element in the center.
• Count the total number of valence electrons.  Add 1 for each negative charge.  Subtract 1 for each positive charge.
• Draw single bonds between the central atom and the surrounding ones.  Complete the octet for all atoms bonded to the central atom except hydrogen.
• If central atom has less than eight electrons, form double and triple bonds on central atom as needed using lone pairs from neighboring atoms, to complete the octet on the central atom.

Example 1: Write the Lewis Structure of nitrogen trifluoride (NF3)

• N is less electronegative than F, so put N in the center.
• Count the valence electrons.  N has 5, and F has 7 each.  
• 5 + (7 x 3) = 26 valence electrons
• Draw single bonds between N and F atoms and complete octets on N and F atoms.
• Check, are the number of electrons in the structure equal to number of valence electrons?

Example 2: Write the Lewis structure of the carbonate ion (CO32-)

• C is less electronegative than O, put C in center.
• Count valence electrons.  C has 4 and O has 6.  Also, note the 2- charge.  4 + (6 x 3) + 2 = 24 valence electrons
• Draw single bonds between C and O atoms and complete octet on O atoms.  Which gives us this structure to the right.  But there's a problem.  Carbon, the central atom, hasn't formed a complete octet.  How do we correct this?
• In this situation we take one of the lone pairs of the oxygen atoms and make a double bond with the carbon atom.

So how about when there's more than one skeletal structure of a compound such as formaldehyfe (CH2O)?

An atom's formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.  Formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom - total number of nonbonding electrons - 1/2 (total number of bonding electrons).  The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule of ion.

So let's look at the above structure.  It appears to be pretty legit.  The hydrogen atoms both have two electrons.  The carbon atom has one non-bonding pair of electrons, a single bond, and a double bond.  That comes out to be a complete octet.  The oxygen atom appears to also be a complete octet.  It has one lone pair, a single bond, and a double bond.  In total there are 12 valence electrons.  Let's plug all of this into the formal charge recipe I have in rainbow colors in the paragraph above this one.

The total number of valence electrons in the free atom (C) is 4.  Apparently this is the number of valence electrons an element, in this case Carbon, has.  The total number of nonbonding electrons is 2 aka that lone pair I mentioned.  There are 6 remaining electrons that are bonded, which we cut in half because that's what the formula says to do.  Formal charge on C = 4 - 2 - (1/2 x 6) = -1

We can do this again with O.  The total number of valence electrons in this case is 6.  It has the same number of non-bonding electrons as C, 2.  It also happens to have the same number of bonded electrons, 6.  That's pretty easy.  Formal charge on O = 6 - 2 - (1/2 x 6) = +1

Now let's do the same thing with this alternative structure and see if anything changes and if there's any significance to doing any of this at all.

Formal charge on C = 4 - 0 - (1/2 x 8) = 0  (It's different!!!!!!)

Formal charge on O = 6 - 4 - (1/2 x 4) = 0 

This second structure is favorable because it has a charge of 0 on the atoms.  

Guidelines for Formal Charge and Lewis Structures:

• For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.  Just like we saw above!
• Lewis structures with large formal charges are less plausible than those with small formal charges.
• Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.

Resonance Structures

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.  These are the variations you can make by flipping around single and double bonds.  It's not that scary.

Exceptions to the Octet Rule

• The Incomplete Octet - when there are simply not enough electrons present to form an octet.
• Example: BeH2, BF3
• Odd-Electron Molecules - you just can't form an octet with an odd number =(
• Example: NO
• The Exapanded Octet (central atom with principal quantum number n > 2)
• Example: SF6