## Lewis Dot Structure Definition

Lewis dot structure definition: a visual way to clearly depict the connection of atoms and the electrons present in a molecule. With a carbon Lewis dot structure, one can see how the atoms in a molecule are bonded together, which gives us more information about the structure than the molecular formula. The figure below shows two molecules’ molecular formulas and Lewis dot structures to depict how much more information can be obtained from a Lewis dot structure.

Remember the Lewis dot structure definition here that they are much more visual than the molecular formulas: To create a Lewis dot structure, go through the following steps:

#### 1. Count the number of valence electrons in the molecule.

Valence electrons are high energy electrons in the outermost electron shell where bonding typically occurs. The number of valence electrons can be easily identified by looking at the column on which the atom is located on the periodic table: Here, the periodic table is color coded. Each column corresponds to a certain number of valence electrons. For example, all the elements in the red column have 1 valence electron, all the orange elements have 2 valence electrons, and so on.

Add up the number of valence electrons for each atom in the molecule to find the total number of electrons. This sum is the number of electrons you must use in the Lewis dot structure.

We will follow each step of creating a carbon Lewis dot structure using CO2 as an example.

Example
Find the Lewis dot structure for CO2.
C has 4 valence electrons
O has 6 valence electrons (we must multiple this by 2 because there are 2 oxygen atoms)
Total number of valence electrons=4+6(2)
Total number of valence electrons=16
(this example is continued throughout each step below)

#### 2. Draw the arrangement of the atoms, placing the most electronegative atoms at the end (see section 4 of this chapter to learn more about electronegativity).

Example

Keep it Simple
Practice is the key to being able to solve Lewis dot structure questions. As you do more and more practice problems, patterns for certain atoms begin to emerge. For example, hydrogen and the halogens (F, Cl, Br, I) almost always occupy the end positions of a Lewis dot structure. These atoms also tend to bond only using single bonds.

Reminder: The arrangement of atoms is key. Remember, the visual connection of atoms is key in the Lewis dot structure definition.

Example

#### 4. Add full octets (meaning 8 electrons) to all of the atoms, except the central atom.

i.  EXCEPTIONS:
1. Hydrogen only needs 1 pair of electrons as it does not follow the octet rule.
2. Phosphorus (P) and sulfur (S) can exceed the octet rule. They can have up to 12 electrons (6 pair).

Example

#### 5. Count the number of electrons you have used to far. Subtract the total number of valence electrons from this number. The difference obtained is the number of electrons you have left that still need to be placed on the molecule.

i. Total valence electrons-Electrons used=Electrons remaining.

Example
Find the Lewis dot structure for CO2 (continued from step 4).
16-16=0
This tells us we have used all the electrons available to us in this molecule.

#### 6. Add the number of electrons from the difference obtained in step 5 to the central atom. If the difference in step 5 was zero, skip this step and go to step 7.

Example
Find the Lewis dot structure for CO2 (continued from step 5).
We skip this step because we have zero remaining electrons. We cannot add any more to the central atom

#### 7. Ensure all atoms have a full octet (remember the exceptions given in step 4). If any atom does not have the full octet, rearrange lone pairs into double or triple bonds to ensure that all atoms have a full octet of electrons.

Example
Find the Lewis dot structure for CO2 (continued from step 6).
Carbon does not have a full octet in the structure given in step 4, so let’s bring in one of the lone pairs from the left oxygen atom and see what happens: Here, we have a double bond between the oxygen and the carbon on the left. But the central carbon is still lacking a full octet, so we must rearrange further. Let’s move a lone pair from the oxygen on the right to form yet another double bond: Here, all atoms have a full octet, so this is our answer. Redrawn, the molecule looks like: 