Lewis structures go by many names, including Lewis electron dot structures, Lewis dot diagrams, and electron dot structures. All these names refer to the same sort of diagram, which is intended to show the location of bonds and electron pairs.

Key Takeaways: Lewis Structure

• A Lewis structure is a diagram that shows the covalent bonds and lone electron pairs in a molecule.
• Lewis structures are based on the octet rule.
• While Lewis structures are useful for describing chemical bonding, they are limited in that they do not account for aromaticity, nor do they describe magnetic behavior accurately.

Lewis Structure Definition

A Lewis structure is a structural representation of a molecule where dots are used to show electron position around the atoms and lines or dot pairs represent covalent bonds between atoms. The purpose of drawing a Lewis dot structure is to identify the lone electron pairs in molecules to help determine chemical bond formation. Lewis structures may be made for molecules that contain covalent bonds and for coordination compounds. The reason is that electrons are shared in a covalent bond. In an ionic bond, it's more like one atom donates an electron to the other atom.

A Lewis structure is named for Gilbert N. Lewis, who introduced the idea in the article The Atom and the Molecule in 1916.

Also Known As: Lewis structures are also called Lewis dot diagrams, electron dot diagrams, Lewis dot formulas, or electron dot formulas. Technically, Lewis structures and electron dot structures are different because electron dot structures show all electrons as dots, while Lewis structures indicate shared pairs in a chemical bond by drawing a line.

How a Lewis Structure Works

A Lewis structure is based on the concept of the octet rule in which atoms share electrons so that each atom has 8 electrons in its outer shell. As an example, an oxygen atom has 6 electrons in its outer shell. In a Lewis structure, these 6 dots are arranged so an atom has two lone pairs and two single electrons. The two pairs would be opposite each other around the O symbol and the two single electrons would be on the other sides of the atom, opposite each other. In general, single electrons are written on the side of an element symbol. An incorrect placement would be (for example), four electrons on one side of the atom and two on the opposite side. When oxygen bonds to two hydrogen atoms to form water, each hydrogen atom has one dot for its lone electron. The electron dot structure for water shows the single electrons for oxygen sharing space with the single electrons from hydrogen. All 8 spots for dots around oxygen are filed, so the molecule has a stable octet.

How to Write a Lewis Structure

For a neutral molecule, follow these steps:

1. Determine how many valence electrons each atom in the molecule has. Like for carbon dioxide, each carbon has 4 valence electrons. Oxygen has 6 valence electrons.
2. If a molecule has more than one type of atom, the most metallic or least electronegative atom goes in the center. If you don't know the electronegativity, remember the trend is that electronegativity decreases as you move away from fluorine on the periodic table.
3. Arrange electrons so each atom contributes one electron to form a single bond between each atom.
4. Finally, count the electrons around each atom. If each has 8 or an octet, then the octet is complete. If not, proceed to the next step.
5. If you have an atom that is missing dots, re-draw the structure to make certain electrons form pairs to get the number on each atom to 8. For example, with carbon dioxide, the initial structure has 7 electrons associated with each oxygen atom and 6 electrons for the carbon atom. The final structure puts two pairs (2 sets of 2 dots) on each oxygen atom, two oxygen electron dots facing the carbon atom, and two sets of carbon dots (2 electrons on each side). There are 4 electrons between each oxygen and carbon, which are drawn as double bonds.

Sources

• Lewis, G. N. (1916). "The Atom and the Molecule", J. Am. Chem. Soc., 38 (4): 762–85. doi:10.1021/ja02261a002
• Weinhold, Frank; Landis, Clark R. (2005). Valency and bonding: A Natural Bond Orbital Donor-Acceptor Perspective. Cambridge: Cambridge University Press. p. 367. ISBN 0-521-83128-8.
• Zumdahl, S. (2005) Chemical Principles. Houghton-Mifflin. ISBN 0-618-37206-7.