A Lewis structure is a diagram that shows the chemical bonds between atoms in a molecule and the valence electrons or lone pairs of electrons. The diagram is also called a Lewis dot diagram, Lewis dot formula, or electron dot diagram. Lewis structures take their name from Gilbert N. Lewis, who introduced valence bond theory and dot structures in the 1916 article The Atom and the Molecule.
A Lewis structure shows how electrons are arranged around atoms, but it doesn’t explain how the electrons are shared between atoms, how chemical bonds form, or what the geometry of a molecule is. Here is how to draw a Lewis structure, with examples and a look at both the importance and limitations of the diagrams.
Parts of a Lewis Structure
Lewis structures are drawn for molecules and complexes. A Lewis structure consists of the following parts:
- Element symbols
- Dots that indicate valence electrons
- Lines that indicate chemical bonds (one line for a single bond, two for a double bond, etc.)
- The dots and lines satisfy the octet rule.
- If the structure carries a net charge, brackets enclose it and the charge is listed in the upper righthand corner
Note: Sometimes the terms “Lewis structure” and “electron dot structure” are used interchangeably. Technically, they are a bit different. A Lewis structure uses lines to indicate chemical bonds, while an electron dot structure only uses dots.
Steps to Draw a Lewis Structure
There are only a few steps to draw a Lewis structure, but it can take some trial and error to get it right.
- Find the total number of valence electrons for all atoms in the molecule. For a neutral molecule, this is the sum of the valence electrons in each atom. The number of valence electrons for an element is usually the same as its group number of on the periodic table (except for helium and the metals). If the molecule has a charge, subtract one electron for each positive charge or add one electron for each negative charge. For example, for NO3–, you have 5 electrons for the nitrogen atom and 3 x 6 = 18 electrons for the oxygen atoms, plus one valence electron for the net charge, giving a total of 24 valence electrons (5 + 18 + 1).
- Draw the skeleton structure of the molecule. At this point, assume the atoms are connected by single bonds. Usually, the atom that has the most bonding sites is the central atom (so carbon would be central over oxygen).
- Determine how many electrons are needed to satisfy the octet rule. The valence electron shell of hydrogen and helium fill with 2 electrons. For other atoms, up to period 4 of the periodic table, the valence shell fills with 8 electrons. Each chemical bond requires two electrons, so use two valence electrons to form each bond between atoms in the skeleton structure. For NO3–, 6 electrons were used to draw the single bonds for the skeleton. So, 18 electrons remain. Starting with the most electronegative atom, distribute these electrons to try to fill the octets of the atoms.
- Distribute the remaining valence electrons. Draw these non-bonding electrons as dots around the atoms to satisfy the octet rule.
- Draw the chemical bonds in the molecule. If all of the octets aren’t filled, make double bonds or triple bonds. To do this, use a lone pair of electrons on an electronegative atom and make it into a bonding pair shared with an electropositive atom that lacks electrons.
- Check to make sure you have the lowest formal charge for each atom. Don’t violate the octet rule. The formal charge is the number of valence electrons, minus half the number of bonding electrons, minus the number of lone electrons. So, for each single-bonded oxygen it’s 6 – 1 – 6 = -1; for nitrogen it’s 5 – 4 – 0 = +1; for the double-bonded oxygen it’s 6 – 2 – 4 = 0. There are two single-bonded oxygen atoms, one nitrogen, and one double-bonded oxygen, so the net formal charge is -1 + -1 + 1 + 0 = -1. Either indicate the formal charges separately or else draw a bracket around the structure and add – or -1 as a superscript.
Different Ways to Draw Lewis Structures
There is more than one “right” way to draw a Lewis structure. If you are drawing the structures for a chemistry class, be sure to know what your instructor expects. For example, some chemists prefer to see skeletal structures that do no show any geometry, while other prefer to see shapes (e.g., the bent shape of water, with nonbonding electron pairs at an angle on one side of the oxygen atom). Some like to see atoms and their electrons in color (e.g., oxygen and its electrons in red, carbon and its atoms in black).
Why Lewis Structures Are Important
Lewis structures help describe valence, chemical bonding, and oxidation states because many atoms fill or half-fill their valence shell. The behavior described by the structures closely approximates real behavior of lighter elements, which have eight valence electrons. So, they are particularly helpful in organic chemistry and biochemistry, which relies on the behavior of carbon, hydrogen, and oxygen. Although Lewis structures do not necessarily show geometry, they are used to predict geometry, reactivity, and polarity.
Limitations of Lewis Structures
While useful for some applications, Lewis structures aren’t perfect. They don’t work well when molecules contain atoms with more than eight valence electrons, such as the lanthanides and actinides. Inorganic and organometallic compounds employ bonding schemes beyond those described by Lewis structures. In particular, molecular orbitals may be fully delocalized. Lewis structures do not account for aromaticity. Even with lighter molecules (O2, ClO2, NO), the predicted structures differ from real behavior enough that Lewis structures might lead to incorrect predictions about bond length, magnetic properties, and bond orders.
- IUPAC (1997). “Lewis formula”. Compendium of Chemical Terminology (the “Gold Book”) (2nd ed.). Blackwell Scientific Publications. ISBN 0-9678550-9-8.
- Lewis, G. N. (1916), “The Atom and the Molecule”. J. Am. Chem. Soc. 38 (4): 762–85. doi:10.1021/ja02261a002
- Miburo, Barnabe B. (1993). “Simplified Lewis Structure Drawing for Non-science Majors”. J. Chem. Educ. 75 (3): 317. doi:10.1021/ed075p317
- Zumdahl, S. (2005) Chemical Principles. Houghton-Mifflin. ISBN 0-618-37206-7.