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Drawing Lewis Structures

With some practice, you’ll be able to create Lewis structures for almost any molecule. Use the procedure below as a guide.

Let's take nitromethane (CH₃NO₂) as an example. The condensed structure is shown on the left, and your goal is to draw the corresponding Lewis structure on the right.

  1. Calculate the total number of valence electrons in the molecule.

    (1) C     1 x 4 =  4
    (1) N     1 x 5 =  5
    (2) O     2 x 6 = 12
    (3) H     3 x 1 =  3
                 total  =  24 valence electrons

  2. Determine the central Atom

    1. The central atom is often written first in the chemical formula
    2. C, B, P and Si are almost always the central atoms.  H and F are NEVER central atoms.
    3. The element with the lowest electronegativity (furthest from F on periodic table) is usually the central atom.
    4. If elements are in the same group (column), or have similar electronegativity, the element with the higher atomic number will be the central atom.  

  3. Adjust the valence electron count for charge

    For negative charges add the corresponding number of electrons, for positive charged atoms subtract electrons.

    In the case of nitromethane, since the molecule is neutral, we do not need to adjust.

  4. Draw the molecule with single bonds only

  5. Subtract the number of electrons in single bonds from the total valence electrons

    We have 6 sigma bonds and each bond holds 2 electrons, so we have 12 electrons in these bonds.  So 24 - 12 = 12 remaining electrons.

  6. Add the non-bonding electrons

    Start with the most electronegative elements filling them first, then to the next most electronegative.

  7. Calculate the formal charges

  8. Create any π bonds

    Use non-bonding pairs on negative atoms to create π bonds with adjacent positively charged atoms.  Be sure to recalculate the formal charges.



Take Note

When drawing Lewis structures for relatively small molecules and polyatomic ions, the structures are generally more stable when they are compact and symmetrical rather than arranged in extended chains of atoms. Compact, symmetrical arrangements minimize electron repulsion and often reflect the molecule's true 3D shape, leading to greater stability. For instance, carbon dioxide (CO₂) is drawn with a linear, symmetrical structure, which is more stable compared to any extended or bent configuration.