Limitations of Lewis Dot Structure

Lewis structures: Definition

You might have already met Lewis dot structures in the article "Lewis dot diagrams". We'd recommend checking out that article before you read on any further. But before we move on to the main focus of this article, what actually are Lewis structures?

Lewis dot structures are also known as Lewis structures, Lewis dot diagrams, dot and cross diagrams and electron dot diagrams. These five terms all refer to the same thing: a diagram showing a molecule's atoms, valence electrons and bonding. In these diagrams, electrons are shown as dots and covalent bonds are shown as lines.

Lewis dot structures are built around the idea that atoms obey the octet rule, meaning that they try to have eight electrons in their outer shell. However, as we'll see later on, this isn't always the case.

Examples of Lewis dot structures

It might be helpful to look at a few examples of Lewis dot structures before we discuss their drawbacks. First up, let's consider water, H₂O. Here is its Lewis structure:

Fig. 1. Water Lewis dot diagram.

What does this tell us?

• The molecule consists of two hydrogen atoms joined to a central oxygen atom by single covalent bonds. These are represented by straight lines.
• The oxygen atom has two lone pairs of electrons. These are represented by pairs of dots.

Next, let's consider the Lewis structure of ozone, O₃.

Fig. 2. Ozone Lewis dot diagram.

What can we infer?

• The molecule consists of three oxygen atoms joined by a single and a double covalent bond.
• The oxygen atoms have varying numbers of lone pairs of electrons.

However, this structure isn't completely accurate. The two bonds between the oxygen atoms are actually equivalent. Instead of there being one single and one double bond, both bonds can be thought of as one-and-a-half bonds. This is because ozone shows resonance. This is an example of one of the limitations of Lewis dot structures. Let's look at these now.

Limitations of the Lewis dot structure

As handy as the Lewis dot structure is, it has its limitations. We've already touched on a few of them. Here are some more:

• It doesn't show the length of bonds or the size of atoms.
• It doesn't show the different types of orbitals.
• It doesn't accurately represent resonance.
• It doesn't show geometry.
• It assumes that all atoms follow the octet rule.

Limitations of Lewis structures: Bond length and atom size

First of all, Lewis dot structures aren't scaled diagrams. They don't show the relative sizes of atoms or the lengths of bonds. In Lewis structures, all atoms are shown as being the same size and all bonds are shown as being the same length, whereas, in reality, this isn't the case.

Fig. 3. Ethene. Source: commons.wikimedia.org

Limitations of Lewis structures: Orbitals

When atoms bond covalently, their electron orbitals overlap. But before this happens, the atom sometimes alters some of its orbitals to make them all equal. This process is called hybridisation.

For example, sp² orbitals are made by one s orbital and two p orbitals rearranging themselves to form three identical orbitals. We find them in ethene, for example. On the other hand, sp³ orbitals are made by one s orbital and three p orbitals rearranging themselves to form four identical orbitals. We find them in ethane.

However, Lewis diagrams don't distinguish between different electron orbitals and show all covalent bonds as being the same.

Fig. 4. Ethene, left, and ethane, right. Source: commons.wikimedia.org

Limitations of Lewis structures: Resonance

Earlier in the article, we looked at the Lewis dot structure of ozone. It contained an O-O single bond and an O=O double bond. In our Lewis structure, the double bond was on the right-hand side, but we can also draw an equally valid Lewis structure with the double bond on the left.

Fig. 5. Resonance structures of ozone.

In actual fact, ozone shows resonance. This means that it can't be accurately represented by either of these two Lewis structures, known as resonance structures. Instead, it takes the form of a hybrid molecule, which is an average of the two resonance structures. Rather than having one O-O single bond and one O=O double bond, it has two equal one-and-a-half bonds. You can clearly see that this is not shown in the Lewis diagrams of ozone's resonance structures. They are an inaccurate representation of ozone, and of resonance in general.

Limitations of Lewis structures: Geometry

Next up, let's consider Lewis structures and geometry. In fact, Lewis dot structures are a poor representation of a molecule's geometry. They don't show any bond angles or positions.

Take water, for example. We've already seen its Lewis structure. Water is a v-shaped molecule, meaning that the angle between its two bonds is 104.5°. But you can't tell this from a Lewis diagram.

Fig. 6. Water Lewis diagram.

Fig. 6. Ammonia, shown with its Lewis structure, left, and geometry, right. Modified from: commons.wikimedia.org

Limitations of Lewis structures: Octet rule

When you draw Lewis diagrams, you have to assign pairs of electrons to atoms. You do this by assuming that all atoms follow the octet rule.

This means that we try to make sure that all atoms have eight valence electrons. However, for some atoms and molecules, this isn't the case.

Fig. 7. Boron trifluoride.

Lewis diagrams also tell us that noble gases like xenon can't form any bonds - they already have eight valence electrons. Actually, this isn't the case. Xenon can form molecules, such as xenon tetrafluoride.

Fig. 8. Xenon tetrafluoride Lewis dot diagram. Image credits: commons.wikimedia.org

Advantages and disadvantages of Lewis structures

Here's a table to help you understand the advantages and disadvantages of Lewis structures:

We're done with this article. By now you should be able to draw and interpret Lewis structures for different molecules. You should also be able to describe and explain some of the limitations of Lewis structures.