Dr.
Gutow's
Lewis Structure Tutorial
To draw a Lewis Structure follow the steps below. At each step a
few molecules you can try are listed. Clicking on the highlighted
molecular formulas will bring up what you should have for each molecule
at each step. Try it yourself and then check as you go
along.
After you have mastered this there are also some quick
rules which can be used for some of the atoms in the periodic
table.
There are a lot of web sites with descriptions of how to draw Lewis
structures. You might find their explanations or methods helpful
as well. Here are a number of links in no particular order: thinkquest,
Univ
of Waterloo, St.
Olaf College, Univ.
Missouri-Rolla. You can find other links by doing a Google search for
"Lewis Structures".
1) Count the total number of valence electrons using the periodic
table as a guide. Remember that for the main group
elements
the number of electrons contributed by each atom can be determined from
the group the element is in. For groups 1 (1A) and 2 (2A) the
number
of valence electrons equals the group number. For groups 13 (3A)
- 18 (8A) the number of valence electrons is the number of the group
minus
10.
2) Next draw single bonds from each of the outer atoms to the central
atom
in the molecule (in the examples below the central atom is in boldface
type). The central atom is usually the atom with the lowest
electron affinity. However, it is sometimes difficult to tell
without
additional information. Subtract two electrons from the total
number
of electrons for each bond you have made. This tells you how many
electrons you have left to use elsewhere.
3) Remembering that each bond counts as two electrons, put electrons on
the outer atoms to give each atom a total of eight (an octet).
Remember
that H (hydrogen) only needs two electrons and B (boron) often only
takes
six electrons. Keep track of how many electrons you are
using.
If you run out of electrons before filling the outer atoms' octets,
stop.
4) Any electrons that were not used up in step 3 should be put on the
central
atom. You should now have no unused valence electrons.
5) If any atoms do not have octets, make multiple bonds (double and
triple)
by sharing electron pairs from atoms that do have octets.
6) Look for resonance structures. If you have made multiple bonds
or have odd electron species where all the atoms cannot have octets,
there
may be more than one way to arrange the multiple bonds or place the odd
electron. If so, the molecule is better modelled as an average of
all the possible structures.
7) Use formal charge to pick the resonance structure likely to
contribute the most to the structure. (Formal charge) = (number of
valence electrons atom starts with) - (number of electrons assigned to
the atom). (Number electrons assigned to atom) = (number of
non-bonding electrons) + 1/2(number of bonding electrons). Do
this calculation for each atom in each resonance structure.
Resonance structures with formal charges closest to zero are lower
energy. Negative formal charges on the more electronegative atoms
are also preferred.
8) If the central atom is from row 3 of the periodic table or
below, check for additional resonance structures where the central atom
has more than 8 electrons (expanded octet). Sometimes these
resonance structures produce better formal charges. This is most
commonly
encountered with compounds of S and P.
****Quick rules which work well for some
period two atoms (H, C,
N,
O, F) and the halogens (group 17) most of the time.*******
| Atom |
Number of Bonds |
Lewis Cartoon |
| H |
1 |
H- |
| C |
4 |
 |
| N |
3 |
 |
| O |
2 |
 |
| F (same for other halogens) |
1 |
 |
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Gutow's Home Page
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Last updated: December 21, 2006
Expires: ---