# Lewis Structures

Last updated: October 28th, 2022 |

There are essentially three uses for Lewis structures. They’re good at helping you get acquainted with the placement of electrons around atoms, helping to visualize molecular geometry, and finally to remember the location of the lone pairs. As you’ll see in Org 1, the geometry of molecules as well as the placement of lone pairs has a huge impact on reactivity, so using Lewis structures to get yourself reacquainted with these principles would be a worthwhile exercise.

What are Lewis structures?

They are a way of drawing molecules that show the placement of electrons between atoms. For instance, here are the lewis dot structures of beryllium dichloride, borane, methane, ammonia, water, and hydrofluoric acid.

The advantage of the full Lewis is that it allows you to see where all the electrons are and to determine if each atom obeys the octet rule. You can see clearly that molecules can have both bonding electrons, which are shared between atoms and non-bonding electrons, otherwise known as lone pairs.

The full Lewis is kind of like training wheels on a bicycle. It’s useful when you are just getting started and feeling uneasy about this business of atoms, electrons, and molecules, and want to determine for yourself that the octet rule is indeed a widespread phenomenon that (most) molecules abide by [the exceptions here being beryllium and boron, which are electron-deficient].

However, you will quickly realize that it  is actually kind of a pain to draw a full Lewis structure. Once you become familiar with the basics of drawing molecules, you’ll find it’s much less work to simply draw a line where the bond is, as shown below.  This also has the advantage that it is much easier to show geometry using the line bonds, because it’s less cluttered.

This brings us to our second point. Electron pairs repel – this applies both to bonding electrons and electrons in lone pairs.  So molecules will adopt a geometry which maximizes the distance between them. This is why methane is tetrahedral, (internal angles 109°) and not square planar (internal angles 90°), and water is bent, not linear. You might recall this is referred to as VSEPR (valence shell electron pair repulsion).

Drawing out molecular geometry using a full Lewis makes for an extremely cluttered drawing. That’s why we drop the full Lewis for the half-Lewis, move the lines around, and just leave the electron pairs in.

Now there is even a second tier of laziness. It’s less work to just drop drawing the electron pairs altogether. This is by far the most common way molecules are drawn.  Let’s look at a few examples.

You’re supposed to know that the lone pairs are still there, even though they’re not drawn in. Think of them like chemical stick figures. For instance, xkcd tends not to draw faces, feet, or hands, but that doesn’t mean the characters he draws are implied to be faceless amputees. It’s just quicker to draw stick figures.This is important because in organic chemistry, lone pairs often don’t just sit around. As you’ll see later, lone pairs are nucleophiles – they participate in a host of chemical reactions. So it’s crucial to know that they’re there, even if they’re not drawn in.

## Comment section

### 13 thoughts on “Lewis Structures”

1. Bismuth says:

the sense of humour used in these lessons and how fun they teach are what I love about this site

2. Krieger says:

I love Randall Munroe! Glad to see you do too, James!

3. Farhina says:

this is really helpful. thanks a lot:)

4. Stephanie says:

Just curious, how do we learn the shapes? Should we memorize every compound/ion and their shapes?

1. There are really only 3 shapes. Tetrahedral, trigonal planar, and linear. The “extra” shapes (such as “bent”, “trigonal pyramidal”) just arise from the fact that a lone pair is present instead of a bond to an atom. Don’t memorize every compound/ion. Just learn to count the number of electron pairs and go from there.

5. Karmen says:

The first of the “half-Lewis” structures is incorrect, I think. It should be Beryllium diflouride instead of dichloride?

These lessons are great, by the way. I wish my professors taught like this.

1. james says:

Oh, thank you!

1. khushboo says:

Why can it not be beryllium(II) chloride?

1. James Ashenhurst says:

It could be the chloride, I arbitrarily picked fluoride.

6. Katharine says:

I’ve been finding some conflicting information regarding methyamine (CH3NH2) and its geometry/bond angles related to the central N atom. My textbook states that the bond angle is 109 degrees–I understand that it is sp3 hybridized and tetrahedral molecular geometry, but shouldn’t the bond angle be ~ 107 degrees (trigonal pyramidal electron geometry) because of the lone pair? I’ve found a few sites confirming this, but others state what my text does. If it is supposed to be 109 degrees, why?
Thanks. :)

1. james says:

I would agree that it should be around 107 degrees.

2. lirun says:

In my opinion, both carbon and nitrogen are sp3 hybridized and tetrahedral molecular geometry, the different is that carbon is symmetrical sp3 hybridized, thus endowing the bond angle 109 degree, while nitrogen is asymmetrical sp3 hybridized, thereby giving the bond angle 107 degree.
In summary, from the perspective of carbon or nitrogen, we get variety of bond angle, with the figure being 109 and 107, respectively.

3. Carrie Wolfe says:

In Organic chemistry we are not so careful with the bond angles. Often we state if the angle is close to 109 degrees or close to 120 or 180 degrees. It its sp3 we don’t worry about 107–we just say 109 (that is , “close to” 109).

This site uses Akismet to reduce spam. Learn how your comment data is processed.