Lewis Dot Diagrams


Lewis dot structures are used to show how the atoms of a compound bond together.  Bonding occurs between the valence electrons, electrons in the outermost level of the atom.  As an example we can look at this picture of the electron energy levels of sodium:



Sodium has 11 electrons, 2 electrons in the first energy level (orbit), 8 electrons in the second, and 1 in the third.  Since there is one electron in the outermost level, sodium has one valence electron.

The valence electrons follow a pattern similar to the charges and the electron configurations on the periodic table:


Group Number (Name)

Number of valence electrons

Group 1 (alkali metals)


Group 2 (alkali earth metals)


Group 13 (boron group)


Group 14 (carbon group)


Group 15 (nitrogen group)


Group 16 (oxygen group)


Group 17 (halogens)


Group 18 (noble gases)



A Lewis dot structure for an element shows two things: the symbol with an appropriate number of dots representing the proper number of valence electrons.




The standard way of dot placement in a Lewis dot structure has an electron placed in the four directions (North, South, East and West), and then when more than 4 are needed listed, the electrons begin pairing up:



Octet Rule

Atoms tend to gain, lose or share electrons in order to acquire a full set of valence electrons.  All atoms want to be stable, so they wish to achieve a stable electron configuration, like the noble gas.  Look at the following examples:


Example 1: Selenium has 34 electrons.  In order to become stable, selenium must obtain the same configuration as xenon.  To do this, selenium gains 2 electrons.  When this occurs, the addition of 2 electrons brings selenium’s total electrons up to 36.  Now if we look at the valence electrons, by adding 2 electrons to the 6 valence electrons, selenium obtains 8 valence electrons.



Example 2: Sodium works much in the same way.  The closest noble gas to sodium is neon.  To achieve the same configuration as neon, sodium must lose one electron.  By doing this, the outermost energy level becomes the 2nd energy level, which has 8 electrons.  In both cases, the object is to get to 8 valence electrons.


“Octet” comes from the idea that most atoms want to have eight electrons in its outermost level (shell).  An octet is achieved by gaining, losing or sharing electrons.


An octet will contain 0 or 8 valence electrons.  An octet is achieved with 0 valence electrons, because the next lower orbit becomes the outermost orbit, and this orbit will always have 8.


Exceptions to this rule are H, He and the transition metals.  Helium is a noble gas and stable with 2 electrons.  Hydrogen, having 1 electron, will desire to have the same configuration as helium, the closest noble gas.  Therefore, hydrogen is also stable with 2 valence electrons.  Transition metals can have varies levels of stability and various numbers of valence electrons, so as not to confuse us all, we’ll skip them for now.


Drawing Ionic Compounds

Using the Lewis structures, we can show the ionic bonding of compounds.

In ionic compounds, electron(s) transfer from one element to another, making cations (positive ions) and anions (negative ions) that attract one another.

Ex: Sodium Chloride (NaCl)


Explanation:  Na needs to lose one electron to obtain an octet.  Cl must gain one to get an octet.  The valence electron from the Na is moved to the Cl.  When this happens, Na achieves a +1 charge, and Cl become -1.  The charges pull the two atoms together.


Another example:  MgBr2



Drawing Covalent Compounds

Another way that compounds can form is when the electrons are shared between 2 or more atoms to complete an octet.  This results in a molecule with covalent bonds.  A molecule that combines covalently is ammonia NH3 :



Explanation:  N has 5 valence electrons.  Each H has 1 valence electron.  When H and N share electrons, both atoms can claim the electrons in the sharing.  Therefore, H can claim it has two electrons.  With the first bond, the valence electron count for N is 6.  The N then bonds with the other two H and achieves an octet.



Very often in covalent compounds, there are unshared pairs of electrons which belong exclusively to one of the atoms.  This is OK.  It is not necessary to have bonds coming off of all the valence electrons.


Multiple Bonds

Two atoms may form double or triple covalent bonds in which multiple pairs of shared electrons are used to satisfy the octet rule.

Example 1: Carbon Dioxide, CO2. Between one of the O and the C, there is a double bond.



Example 2:  HCN  Between the C and the N, there is a triple bond.



  Complete the worksheet: Lewis Dot Structures