Lewis Dot Structure

How To Draw The Lewis Dot Structure Of Nitrogen Trifluoride

By looking at the Lewis Dot structure of Nitrogen Trifluoride we can understand how it takes on its triangular shape, as seen to the here. It also explains why the molecule is not drawn as flat but rather has a pyramidal look to it.

To build a picture of Nitrogen Trifluoride, we need to start with the electron dot diagrams for the elements involved in the molecule. In this case those are Nitrogen and Fluorine. There are many ways to fit Nitrogen and Fluorine together in theory so we use the chemical formula to tell us the number of each type of atom involved. In this case the formula for Nitrogen Trifluoride is NF3, meaning there is one Nitrogen atom and three Fluorine atoms in each NF3 molecule, as seen below.

Putting The Lewis Dot Structure Together

Once we have the electron dot diagrams for these elements, it is then a simple matter of putting the puzzle together. We need to fit the four atoms into a single molecule and have no gaps left over. This means that when we look at each atom in the finished molecule it must have eight electrons surrounding it, regardless of how many other atoms it is sharing them with.

Each Fluorine atom has only one gap in its outer shell, so it can bond in only one place. Nitrogen has three gaps in its outer shell so it makes sense to place the Nitrogen in the middle of the molecule and attach the Fluorine atoms to it.

There is one simple rule of thumb for figuring out which atom goes into the center of the molecule. If you look on the Periodic Table (free download available there), it is the element furthest to the left that goes in the middle.

Note that the eight electrons in the outer shell can be moved around to suit the build of the molecule, though they must be evenly spread across the four quarters of the atom as outlined in the electron dot structure page. Building NF3 looks like this:
When we join the pairs of electrons up, the molecule becomes complete. If we count the number of electrons around each atom we can see that each atom is surrounded by eight electrons. This means that its outer electron shell is full and the molecule is (relatively) stable.

From Lewis Dots To Lines

When drawing molecules, it is traditional to replace each pair of electrons with a single solid line joining the atoms together. In the case of the electron pairs that do not join to another atom, the line just stops in space. These are called lone pairs. They are often left out of Lewis Dot Structure drawings, as shown below, but they do have an effect on the structure of the molecule.

It is for this reason I ask my students to always include the lone pairs, as failing to do so can develop bad molecule drawing habits that can be very hard to correct. The actual shape is important for many more advanced concepts, especially in organic chemistry.

remove the lone pairs if you dare

3D Shape: Like Charges Repel

The electron pairs, which are shown as lines, all have negative charges. This means they push away from each other which is the basis of the VSEPR rules. Doing this gives us the three dimensional picture of Nitrogen Trifluoride without lone pairs, as shown below (image courtesy of Wikimedia Commons):

a nice 3D picture of Nitrogen Trifluoride

Note how it seems to be sitting with the Nitrogen atom elevated. If you can imagine the lone pair of electrons sticking out of the top of the blue sphere in the center of the molecule, you will see it has a pyramid-like 3D shape. For a more technical discussion of Nitrogen Trifluoride,, have a read of the Wikipedia page. It’s pretty heavy going though.