Lewis Structures are visual representations of the bonds between atoms and illustrate the lone pairs of electrons in molecules. They can also be called Lewis dot diagrams and are used as a simple way to show the configuration of atoms within a molecule.
Between 1916 and 1919, Gilbert Newton Lewis, Walther Kossel, and Irving Langmuir came up with a theory to explain chemical bonding. This theory would be later called Lewis Theory and it is based on the following principles:
A Lewis Symbol for an element is composed of a chemical symbol surrounded by dots that are used to represent valence electrons. An example of a Lewis symbol is shown below with the element Carbon, which has the electron configuration of 1s22s22p2:
This Lewis symbol shows that carbon has four valence electrons in its outer orbital and these four electrons play a major role in bonding of carbon molecules.
Lewis symbols differ slightly for ions. When forming a Lewis symbol for an ion, the chemical symbol is surrounded by dots that are used to represent valence electrons, and the whole structure is placed in square brackets with superscript representing the charge of the ion. An example of a Lewis symbol for the cation and anion of Carbon is shown below:
Cation of Carbon
Anion of Carbon
To construct Lewis Structures one can generally abide by the following steps:
When constructing the structures keep in mind the following:
Lewis Structures can differ based on whether the electrons are shared through ionic or covalent bonds.
An example of ionic bonding can be seen below in the instance of the reaction of Sodium and Chlorine:
Sodium has one valence electron and Chlorine has seven valence electrons; the two elements together form the noble gas configuration. The Chlorine atom takes the valence electron from the Sodium atom leaving the Chlorine atom with one extra electron and thus negatively charged and the Sodium atom without an electron and thus positively charged. The two atoms then become ions and because of their opposite charges the ions are held together in an ionic bond.
An example of covalent bonding can be seen below with the reaction of Hydrogen and Fluorine:
Hydrogen has one valence electron and Fluorine has seven valence electrons; together the elements form the noble gas configuration. The Hydrogen atom shares its electron with Fluorine atom so that the Hydrogen atom has 2 electrons and the Fluorine atom has 8 electrons. Therefore both atoms have their outermost shells completely filled.
|Example 1: The Chlorate Ion|
Try the Chlorate ion: (ClO3-)
First, lets find the how many valence electrons chlorate has:
ClO3- : 7 e-(from Cl) + 3(6) e-(from 3 O atoms) + 1 (from the total charge of -1) = 26
There are 26 valence electrons.
Next lets draw the basic framework of the molecule:
The molecule uses covalent bonds to hold together the atoms to the central Chlorine. The remaining electrons become non-bonding electrons. Since 6 electrons were used for the bonds, the 20 others become those un-bonding electrons to complete the octet:
|Example 2: Formaldehyde|
Constructing the Lewis Structure of the formaldehyde (H2CO) molecule.
First find the valence electrons:
H2CO: 2(1) e- (from the H atoms) + 4 e- (from the C atom) + 6 e- (from the O atom)
There are 12 valence electrons. Next draw out the framework of the molecule:
To satisfy the octet of Carbon, one of the pairs of electrons on Oxygen must be moved to create a double bond with Carbon. Therefore our Lewis Structure would look as it does below:
The Hydrogen atoms are each filled up with their two electrons and both the Carbon and the Oxygen atoms' octets are filled.
The charge on each atom in a molecule is called the formal charge. The formal charge can be calculated if the electrons in the bonds of the molecule are equally shared between atoms. This is not the same thing as the net charge of the ion.
In calculating formal charge, the following steps can be extremely helpful:
Formal charge = (number of valence electrons) - (number of non-bonding electrons + 1/2 number of bonding electrons)
In Lewis structures, the most favorable structure has the smallest formal charge for the atoms, and negative formal charges tend to come from more electronegative atoms.
An example of determining formal charge can be seen below with the nitrate ion, NO3-:
There are times when more than one acceptable Lewis structure can be drawn for a molecule and no single structure can represent the molecule entirely. When this occurs the molecule/ion is said to have resonance. The combination of the various plausible Lewis structures is called a resonance hybrid.
Some rules for drawing resonance structures are as follows:
|Example 3: Nitrate Ion|
Consider the nitrate ion, NO3-
2. All atoms in water have a formal charge of 0.
4. The formal charge for each H is 0 and for N is 0.
An NSF funded Project