Ms. Burnside's Idiot-proof* method
for drawing correct Lewis Structures

1. Look at the chemical formula of the molecule, and determine the total number of valence electrons availablefrom all the atoms present.

  • If you're drawing the structure for a polyatomic ion, add or subtract electrons to account for the ion's charge, and enclose the ion's Lewis structure in brackets [ ] with the charge outside the upper right corner.

2. Arrange the atom's symbolsinto a logical skeleton structure. If Carbon is present, it will be the central atom. Hydrogen and Fluorine will never be central atoms -- they only ever form ONE covalent bond.

If you're wondering which atom to put in the center:

  • An atom generally will form the samenumber of covalent bonds as the number of electronsit needs to complete its octet (C forms 4 bonds, N forms 3, O forms 2, etc.) -- but this is just a guideline, not a hard-and-fast rule.
  • Highly electronegative atoms are unlikely to form more bonds than they need to, unless they're bonding to something even more electronegative than themselves. For example, oxygen, being the second-most electronegative element, usually will form two covalent bonds; but it will sometimes form only one covalent bond and will have three unshared e- pairs. Fluorine, being the most electronegative atom of all, never forms more than one bond.
  • If the central atom must hold an expanded octet (i.e., more than 8 e-), it must be an element from period 3, 4, 5, or 6 -- these elements have empty d-orbitals available to hold the extra electrons.Elements in periods 1 or 2 can never accommodate an expanded octet.
  • If you're still confused about which atom to put in the center, draw the structure in all possible arrangements and then use formal charge (see next page) to determine which is the best once you're done.

3. Using a "bar" or "stick" to represent a shared electron pair (i.e., one covalent bond), draw single bonds to connect the atoms in your skeleton structure.For each bond you drew, subtract two electrons from the total number of valence electrons you calculated in step 1. The remaining electronswill be usedto satisfy the atoms' octets... but before putting them in place, complete step 4!

4.COUNT how many electrons arestill needed to satisfy the octets of all the atoms in your structure now that single bonds are in place.

  • IF the number of e- needed is the exactly the same as the number of e-left over, go ahead and put them in; show them as pairs of dots (..) around the atoms' symbols. You're done!

But what ifyou need more electrons for octets than you have left????

  • If you need two more electrons than you have (i.e., you're one pair short), use a DOUBLE BOND between two atoms.Subtract two more e- from the total; then use what's left to satisfy octets.
  • If you need FOUR more electrons than you have (i.e., you're two pairs short), use either a TRIPLE BOND, orTWO DOUBLE BONDS between atoms.Subtract FOUR more e- from your total, and use what's left to satisfy octets. Remember, never put a double or triple bond on a hydrogen or fluorine atom!

IF you have extra electrons left over after satisfying all octets:

  • Place extra electrons on thecentral atom (as long as it's from period 3,4,5, or 6).ALL ELECTRONS MUST BE USED UP.

5. If there is more than one possible way to draw the molecule's Lewis structure, use formal charge to determine which is most ideal.

  • Calculate the formal charge for each atom in the structure by the following formula:

A - B - C = Formal Charge

where
A = How many valence electrons that atom normally has when unbonded (i.e., its group number)
B = The number of lone (unshared) electrons residing on that atom in this particular Lewis structure
C = The number of bondsresiding on that atom in this particular Lewis structure.

  • Ideally, in a perfect molecule, all the formal charges will be zero. If not all atoms' formal charges = 0, then the best structure will have the most negative formal charge on the most electronegative atom. The sum of all the formal charges must equal the overall charge on the molecule (i.e., zero for a neutral molecule, or the charge of a polyatomic ion).

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*You don’t want to be the person who proves this wrong. Trust me.