Creating Lewis Structures
First step: Determine the placement of the atoms and draw the Lewis structures for
each atom.
Elements that form many bonds (carbon for example) make good central
elements
Elements that tend to form one bond (hydrogen for example) tend to go on the
periphery
Often formulae are written to indicate the placement of the atoms. For example,
“CH3OCH2CH3” contains structural information and accurately describes one
molecule, whereas “C3H8O” contains no structural information and does not
describe a specific molecule.
Distribute the valence electrons around the nucleus.
Do not pair electrons until necessary.
For anionic molecules, add the necessary extra e–‘s 1 at a time to the most
electronegative atoms.
For cationic molecules, remove the necessary e–‘s one at a time from the least
electronegative atoms.
Second step: Circle two unpaired electrons on adjacent atoms (one on the first atom
and a second on the other atom) to indicate where bonds will form.
Continue circling unpaired electrons until all atoms have an octet of electrons.
If you run out of unpaired electrons rearrange the electrons that remain and continue
making bonds. Never put more than 8 e–‘s in the valence shell of a second period
(row two) element.
Third step: Redraw the molecule.
Replace each pair of circled electrons with a line indicating the presence of bonds.
Make certain all remaining electrons (called lone pair electrons) are paired.
Fourth step: Determine formal charges and indicate the charge of the molecule.
Formal Charge = (# e– 's element started with) - (# e– 's element ended up with)
For the purpose of determining formal charge an atom "ends up with" all of its
unshared electrons, and one half of the electrons in the bonds.
(# e–'s element ended up with) = (# unshared e –'s) + (1/2 # bonding electrons)
Fifth step: Determine resonance structures.
Resonance structures usually can be found by moving around electrons that are on
atoms adjacent to double bonds.
Another method
First step: Determine the total number of valence e–'s in the molecule.
For anionic molecules, remember to add the appropriate number of extra e–'s.
For cationic molecules, remember to subtract the appropriate number of e–'s.
Second step: Arrange the atoms in a way that makes sense.
Elements that form many bonds (carbon for example) make good central
elements
Elements that tend to form one bond (hydrogen for example) tend to go on the
periphery
Often formulae are written to indicate the placement of the atoms. For example,
“CH3OCH2CH3” contains structural information and accurately describes one
molecule, whereas “C3H8O” contains no structural information and does not
describe a specific molecule.
Third step: Connect the atoms with single bonds
Fourth step: Count the number of e–‘s used to make bonds.
Subtract the number of e–‘s used to make bonds from the total number of e–‘s
available
Distribute the remaining e–‘s around the molecule as lone pair e–‘s satisfying
the valence shells of as many atoms as possible
If all atoms have filled valence shells move on to step five.
If all atoms do not have filled valence shells, go back to step three and add
additional bonds between atoms that tend to form more than one bond, and
repeat step four. Continue adding multiple bonds at step 3 and repeating
step 4 until all elements have filled valence shells.
(Hint: the number of e– required to fill the empty shells is a hint to the number of
multiple bonds that are needed. If an unfilled shell needs 2 electrons, you would
convert 1 set of lone pair electrons to a bond to make 1 multiple bond. If 4 electrons
would be required to fill all the valence shells, 2 sets of lone pair e– 's would be
converted to bonds to make more multiple bond, etc)
Fifth step: Determine formal charges and indicate the charge of the molecule.
Formal Charge = (# e– 's element started with) - (# e– 's element ended up with)
For the purpose of determining formal charge an atom "ends up with" all of its
unshared electrons, and one half of the electrons in the bonds.
(# e–'s element ended up with) = (# unshared e –'s) + (1/2 # bonding electrons)
Sixth step: Determine resonance structures.
Resonance structures usually can be found by moving around electrons that are on
atoms adjacent to double bonds
Some simple rules...
For us H will ALWAYS have only one bond (it is possible to have more, but those
bonds are not classical 2-electron bonds, the kind we are talking about for
99.99% of this course. So, one bond for H.)
First row elements will have a maximum of four bonds (same rambling comment
as above).
Carbon will always have four bonds. Three bonds is possible, for example, in CN–
or CO. (Since you asked yes, pentavalent carbon is actually possible, but they
come with the same caveats as mentioned above.)
Oxygen will usually have two bonds. O can have 3, but will have positive charge
as, e.g., in H3O+. O can also have 1 bond, but it will have a negative charge as,
e.g., in OH–.
Nitrogen will usually have 3 bonds. N can have 4, but it will have a positive
charge as, e.g., in NH4+. N can also have 2 bonds, but it will have a negative
charge as, e.g., in H2N–.
Halogens will usually have 1 bond, but can have more when they are part of a
polyatomic ion as, e.g., in ClO4-.
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Tips…
If using the circle unpaired electrons method, when you draw the atoms on the page
try to draw unpaired electrons pointing towards each other.
••
O
•
•
•
•
H
O
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not
•
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H
H
••
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H
Circling electrons and drawing bonds will be easier this way.
Finish off the easiest elements first. If an element has 7 electrons find an unpaired
electron and make a bond, then you do not have to worry about that atom any
more.
Completely finish an element’s valence shell before moving on to the next element.