Lewis Dots and the Bonds They Make
All discussions of bonding, include the two main types of bonds formed between atoms. We’re talking , of
course, about ionic and covalent bonds. Now I know there are other types of bonds, metallic bonds and
coordination complex bonds, but these are special and we won’t be dealing with them here. If you want
more detail on metallic bonding for instance, this is a very interesting extension of covalent bonding, but I
can reserve that for any interested individuals to ask me directly. Let’s get back to the bonding issue.
I’ve already covered much of the ideas surrounding they types of bonding in the summary “Bonding,
Shapes and Forces”, but I’ve only touched on Lewis structures and how we figure out how a covalent
molecule can be formed. This should be a brief write up since It’s just meant to provide some extra detail.
Below are some examples of Lewis structures highlighting where the electrons come from, and how the
bonds are formed
Fluorine (F2)
Ammonia (NH3) a.k.a. Nitrogen trihydride
Hydrogen Sulphide (H2S) a.k.a. Dihydrogen sulphide
Carbon tetrachloride (CCl4)
The purpose of these Lewis dot structures is to show how each atom in the structure is bringing their
normal compliment of valence electrons and making enough bonds with their unpaired electrons to
complete their octet (or duet in the case of Hydrogen). Using black or red for the electrons in the Lewis
diagram for each of the atoms in the molecules, we should see that the valence electrons pair up like
velcro, or lego and hold the atoms together.
Each of the atoms still has the same number of electrons it started with, and all bonds are made from two
electrons, so there are no charges /ions formed in this type of bonding
The goal with a Lewis structure is to find a way to use all the available unpaired electrons to make bonds.
In the case of Carbon (Group 4A), we’ve said in the past that it normally doesn’t form ions to complete it’s
octet (either bringing on 4 or getting rid of 4). Group 4A elements normally share their four unpaired
electrons with unpaired electrons from other atoms and form 4 bonds1. Hydrogen only has one electron to
share, so it normally only forms one bond with another atom2. The halogens just need that extra electron to
fill their octets, so they normally only form one bond as well3.
But, we set this up to suggest that only one bond can exist between two atoms. This is not necessarily
true. What I’ve said is that each atom tries to form as many bonds as it has unpaired electrons.
When we learned to write Lewis dot diagrams for an element, we learned that the diagram was just the
symbol surrounded by the number of valence electrons (its group number)
I also suggested that you should never pair the electrons when writing a Lewis structure until you had to,
i.e. when you reached group 4 and have one on each side, the next electron for group 5 has to be paired.
Going top, bottom, right then left ensures that you always have paired and unpaired as far apart as
possible. Then when you look at a Lewis dot diagram you can see how many Octet Rule compliant bonds
it will likely make (from Group 4 on). It also becomes clear why Group 8A doesn’t make any bonds4. Each
one has 8 paired electrons (or two in the case of He)
I know this is in the book too, so I should mention that Be and B, don’t always get an octet. They can form
just enough bonds to pair up their unpaired valence shell electrons. I know, they’re in Group 2 and 3, and
that should mean they should be 2+ and 3+ and form ionic compounds. But just in case you saw BeCl2 in
the book, or BCl3, written as covalent compounds, it’s just an(other) exception.
1
This is really only true for Carbon in Group 4A, since Silicon can actually make more than 4 bonds. But
that is another story and one that might give you unnecessary anxiety. So let’s forget about Octet
Expansion for the elements from the 3rd period and below
2
I hate to be a pain, but it is possible for hydrogen to form two bonds in a “three centre two electron
bond.” But that’s just being showy...
3
Yah, about that. Again, once we get below the second period, some exceptions happen as mentioned
above. There can be some expansion of the octets, but we’ll just pay a little naive on that.
4
OK, last time for this, but in actual fact some of the “Noble Gases” aren’t so noble and can actually form
compounds - like Xenon in XeF6. Now just forget I said it...
But I want to get to the idea that some elements can form multiple bonds with the another atom. Let’s
forget about the exceptions, they’re just confusing distractions, and lets deal with something more common
and important, since we deal with these compounds all time.
In order for Carbon to bond with two Oxygen in CO2, each of the Carbons and Oxygens must be sharing
two of their unpaired electrons, forming double bonds. In O2, the same must be true, because each
Oxygen needs to make two bonds, having two unpaired electrons. Since Nitrogen has three unpaired
electrons, it normally would need to form 3 bonds - just like in ammonia on the first page. In order to do
that in N2, each Nitrogen would need to form 3 bonds with the other - forming a triple bond
You may notice that these multiple bonds don’t look very elegant, and we should move electrons around so
that they aren’t so close to each other, but the point being made here is that each of the atoms has
managed to get it’s part time octet by sharing more than one unpaired electron with another atom. How
they actually do this in a way that makes it elegant...well that may be more than you signed up for.
...more quantum mechanics...
The other interesting issue that will come out of this is the shapes that molecules will take on to avoid the
obvious crowding of these electrons (remember electrons are negatively charged and will repel each other).
But hopefully I’ve introduce that concept - the infamous “valence shell electron pair repulsion theory” or
VSEPR for short - in the bonding, shapes and forces summary. Here I just wanted to show how covalent
bonding is still about getting an octet, but with co-operation and sharing - two of the most important things
we learned from our Kindergarten teachers...