Lewis Dot Diagrams

When drawing chemical species, we often want to know the arrangement of their valence electrons. After all, the number of valence electrons often dictates a species' properties and how it reacts. But drawing out all of the electron shells can get quite tiresome and complicated, especially for larger molecules. Instead, we use drawings called Lewis dot diagrams.

Fig. 1 - An example of a Lewis dot diagram used to explain the reaction between two fluorine molecules.

• This article is about Lewis dot diagrams in chemistry.
• We'll start by looking at what Lewis dot diagrams are before turning our attention to some common examples.
• We'll then learn how to draw Lewis dot diagrams.
• Finally, we'll consider the importance of Lewis dot diagrams.

What are Lewis dot diagrams?

Above, we introduced you to the idea of a Lewis dot diagram: a simplified representation of a molecule's valence electrons. Lewis dot diagrams are also known as Lewis structures, Lewis dot structures or electron dot structures. Don't worry if you see any of these terms floating around - they all mean the same thing!

Lewis dot diagrams show the atoms in a molecule, be it a neutral compound or a charged ion. They also show the number, position and arrangement of the atoms' valence electrons, whilst assuming that all inner electron shells are complete. Specifically, they show any bonded pairs or lone pairs of electrons.

Note the following:

• Electrons are shown as dots. Two dots side by side represent a lone pair of electrons.
• Bonded pairs of electrons are shown as lines, signifying a covalent bond. A single line represents a single covalent bond with just one bonded pair of electrons, whilst two lines represent a double covalent bond with two bonded pairs of electrons.
• Lewis dot diagrams are based on the octet rule. Atoms are at their most stable when they have full outer shells and the octet rule states that for most atoms, this occurs when they have eight valence electrons.
• Electrons are generally spaced out evenly around all four sides of the atom, with two per side.

Want to see Lewis dot diagrams in practice? Let's now look at some examples.

Lewis dot diagram for oxygen

An oxygen molecule consists of two oxygen atoms, joined by a double covalent bond. Each oxygen atom also has two lone pairs of electrons. You can see this in the Lewis dot diagram below.

Fig. 1: A Lewis dot diagram for an oxygen molecule. Anna Brewer, StudySmarter Originals

How many valence electrons does each oxygen atom have? If you count them up, you'll see that each oxygen atom has four electrons from its two lone pairs, and four electrons from its double covalent bond. This adds up to eight valence electrons. This oxygen molecule, therefore, satisfies the octet rule.

Lewis dot diagram for carbon dioxide

Next up: carbon dioxide, CO₂. This molecule consists of two oxygen atoms joined to a central carbon atom by double covalent bonds. Again, each oxygen atom also has two lone pairs of electrons.

Counting the electrons, we can see that both oxygen atoms have four electrons from their two lone pairs, and four electrons from their double covalent bond. This adds up to eight valence electrons. The carbon atom doesn't have any lone pairs but features two double bonds. Each double bond supplies four electrons, and so carbon also has eight valence electrons.

Lewis dot diagram for water

Let's now look at the Lewis dot diagram for water, H₂O.

Fig. 3: A Lewis dot diagram for water. Anna Brewer, StudySmarter Originals

Here is another example: the Lewis dot diagram for methane, CH₄.

Fig. 4: A Lewis dot diagram for methane. Anna Brewer, StudySmarter Originals

Lewis dot diagram for carbon

We can also draw Lewis dot diagrams for uncombined atoms that have yet to form a molecule. One example is carbon. Here is its Lewis dot diagram:

Fig. 5: A Lewis dot diagram for carbon. Anna Brewer, StudySmarter Originals

Drawing Lewis dot diagrams for simple molecules like oxygen or methane is fairly straightforward. However, it can get a little more tricky for more complex molecules and so there is a set procedure that is helpful to follow. In the next section, we'll look at some more complicated examples of drawing Lewis dot diagrams.

How do you draw Lewis dot diagrams?

Now that you know what a Lewis dot diagram is and what each part represents, let's practice drawing some more complicated molecules. Here are the steps you should follow. Don't worry if it seems quite long-winded - we'll break each part down for you.

1. Use the chemical formula of the molecule and a periodic table to work out the molecule's total number of valence electrons. You'll use this number when adding up the electrons later on.
2. Draw the rough position of the atoms in the molecule. Remember that molecules like to be symmetrical - this should help you to work out which atom is the central atom, and which are outer atoms.
3. Connect the atoms using single covalent bonds, showing them with straight lines.
4. Add pairs of electrons to the outer atoms until they all have a full outer shell of electrons. This usually means that they need eight electrons, but there are some exceptions. Remember to count any bonds as a pair of electrons.
5. Count up how many electrons you've added to your molecule in total, and subtract this from the molecule's number of valence electrons that you worked out earlier. This will give you your number of remaining electrons. Once again, remember to count any bonds as pairs of electrons.
6. Add any remaining electrons to the central atom.
7. The central atom should hopefully now have a full outer shell of electrons. If it doesn't, use some of the lone pairs of electrons from the outer atoms to form double covalent bonds with the central atom, until all atoms have full outer shells.

Sound tricky? We'll start off with a basic molecule. Let's look at an example now.

First of all, we need to work out the number of valence electrons in this molecule, ammonia. We can do this by looking at the periodic table. Nitrogen is in group V and so has five valence electrons. Hydrogen is in group I and so has one valence electron. The total number of valence electrons is therefore 5 + 3(1) = 8.

Next, draw the rough position of the atoms in the molecule. In this case, the central atom is nitrogen, and it is surrounded by the three hydrogen atoms. Add a bond between nitrogen and each hydrogen atom, like so.

Ammonia. Anna Brewer, StudySmarter Originals

Next, add electrons to the outer atoms until they all have full outer shells. You should know that hydrogen can only have two electrons in its outermost shell. Here, each hydrogen atom already has two electrons, namely the bonded pair shared with nitrogen. We, therefore, don't need to add any more electrons to any of the hydrogens. We can move on to the next step: adding any remaining electrons to the central atom, nitrogen.

First, count up how many electrons you've already added to the structure. Here, we've added three bonding pairs, and so we've already added 3(2) = 6 electrons. We calculated earlier that ammonia has eight valence electrons, and so we still need to add 8 - 6 = 2 electrons. We add them to the central atom, nitrogen, as a lone pair.

Fig. 6: Ammonia. Anna Brewer, StudySmarter Originals

Let's count up the electrons around nitrogen. It has three bonding pairs and one lone pair, giving it 3(2) + 2 = 8 electrons. It, therefore, has a full outer shell - we don't need to form any more covalent bonds. Our Lewis dot diagram is complete.

Here's a slightly harder example for you to have a go at.

Firstly, let's count up the number of valence electrons. Carbon is in group IV, giving it four valence electrons. Oxygen is in group VI, giving it six valence electrons. However, this is a negative ion with a charge of 2-. It, therefore, has two extra electrons. This means that the total number of valence electrons is 4 + 3(6) + 2 = 24.

Next, draw out the different atoms. This molecule consists of a carbon atom surrounded by three oxygen atoms. Add bonds between each oxygen atom and the carbon atom, and you should end up with something like this:

Fig. 7: The carbonate ion. Anna Brewer, StudySmarter Originals

Next up, add electrons to the outer atoms until they all have full outer shells. Each oxygen atom already has two electrons in its outer shell, thanks to the bonded pair it shares with carbon, so we just need to add six more to each.

Fig. 8: The carbonate ion. Anna Brewer, StudySmarter Originals

Count up the total number of electrons that you've already added. Here, we have three bonding pairs and nine lone pairs, giving us 3(2) + 9(2) = 24 electrons. We know that the carbonate ion only has 24 valence electrons. Therefore, we can't add more to the central atom, carbon. But looking at this carbon atom, we can see that it doesn't quite have a full outer shell - it only has six valence electrons, and ideally, it needs eight. To solve this, we use a lone pair of electrons from one of the oxygen atoms to form another bonded pair, creating a double bond between carbon and oxygen.

All atoms now have full outer shells. There's just one final thing to do - don't forget to add charge symbols to the molecule, to show that it is a 2- ion! Here is your final Lewis dot diagram.

Why are Lewis dot diagrams important?

So far, we've found out what Lewis dot diagrams are and had a look at some examples for simple molecules. We've also had a go at drawing Lewis dot diagrams for more complicated molecules. But why do we need to know how to do this, and why are Lewis dot diagrams important?

Well, Lewis dot diagrams show us the number and type of valence electrons of each atom in a molecule. With a Lewis dot diagram, we can clearly see how many lone pairs and bonded pairs of electrons an atom has, and whether these are part of a single, double or triple bond. Knowing the arrangement of a molecule's valence electrons is useful for a few different reasons.

• Firstly, it allows us to predict its shape, alongside the angle between any bonds. It also allows us to predict the bond lengths. This is known as molecular geometry.
• Secondly, it also helps us to predict how the molecule reacts, as well as its properties.