Lewis+Structures

Lewis Structures
Lewis structures are drawings that show the valence electrons in an atom by showing bond formation through using dots that represent each valence electron. The best way to go about finding the element's number of valence electrons is writing out the electron configuration. (A good thing to do before moving on to Lewis Structures would be to go back and review how to write the electron configuration.)

For example, Oxygen. The electron configuration of Oxygen would be:

 From this, you would take the electrons in the outermost energy level, so in this case, there would be six valence electrons. Then you write the chemical symbol (O) and distribute the valence electrons so that all are used, but only two per side. It should look like this when complete:   This same path should be taken for virtually any element.

Here is an easy overview of Lewis Structures which are helpful for understanding this concept:  media type="youtube" key="EJMnyHCP0H4" height="385" width="480"  This video also presents writing Lewis Structures for compounds. Basically, the same path is taken, but slightly modified. Before moving on, it is important to remember **all atoms' outer valence shells are complete with 8 electrons**, so that is a major concept to keep in mind when writing these structures for compounds.

First, Ionic compounds such as NaCl: It is important to write out the Lewis structures for both Na and Cl before trying to do anything else. and  Now remembering that 8 electrons makes an atom "happy", you can observe that taking one valence electron from Na and giving it to Cl would make Cl have 8 valence electrons.



<span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"><span style="font-family: 'Trebuchet MS',Helvetica,sans-serif;"> In the case of Ionic compounds such as this, the electrons are being **transferred**. This is shown by taking the transferred electron and adding it to the Cl structure, but you must show this through the use of brackets. In the end, it would look like:

This same process can be followed for other Ionic compounds.

Covalent compounds follow a different set of rules. Unlike Ionic ones, these compounds share valence electrons equally. An easy example would be CH 4. First, count up the valence electrons for both C and the 4 H's... following the same process as before, but it is not absolutely imperative that you draw the Lewis Structures quite yet. Now add up them up to get the total number of valence electrons for the compound which would be 8. It would be drawn with the Carbon in the middle with the Hydrogen's surrounding it, so that Hydrogen gets two electrons because that is all Hydrogen needs to be stable. (Review special circumstances.)



An important thing to keep in mind when writing the Lewis Structures for Covalent Compounds is the **Octet Rule**. (All atoms are stable with 8 electrons except for Hydrogen as mentioned earlier.) Some break the Octet Rule!

Those that __**NEVER**__ break the Octet Rule are: Carbon, Nitrogen, Oxygen and Fluorine. The goal is not to break it unless it is necessary... For clarification on going over or under the normal 8 valence electrons, see: [|Exceptions to the Octet Rule]

Certain atoms of elements require double or triple bonds as opposed to the usual single bonds. In this case, Carbon, Nitrogen, Oxygen, Phosphorous and Sulfur are the **only** ones that make double or triple bonds, so any one of the elements listed will be on either side of the bond. Halogens **never** make double or triple bonds.

For example, CO 2. This has a total of 16 electrons which need to be distributed among the C and the 2 O's which seems impossible because each is stable with 8, but with double bonds, it is possible.

Because of double bonds, each has the 8 electrons, so each is stable. As you can see below, the eight are circled for each.



The number of shared electrons between atoms plays a part in the distance between the atoms. As the number of shared electrons increases, the distance between the atoms decreases. This distance is known as the bond length which is measured between the nuclei of the bonded atoms.

For information on **Resonance Structures**, see Valence Bond Theory. Another important thing to view when studying Lewis Structures is the **[|VSEPR Theory]** of Molecular Geometry.

Citations: http://www.fordhamprep.org/gcurran/sho/sho/lessons/lesson38.htm