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 display a popup showing 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 some links in no particular order: Univ of Waterloo, St. Olaf College. 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.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO
  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.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO
  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.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO
  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.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO
  5. If any atoms do not have octets, make multiple bonds (double and triple)by sharing electron pairs from atoms that do have octets.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO
  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.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO
  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.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO
  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.
    CH4 CF2Cl2 SO2 O3 PO43- H2SO4 NO

****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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Last updated: June 15, 2017
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