The Octet Rule

A football team can only send 11 players onto the pitch at any one time. They're not allowed anymore players - if they go over 11, they have to send someone off. But too few players is also a disadvantage, and in that case, the manager will try to send someone else on. You could say that football teams are at their best with exactly eleven players. Atoms are quite similar. Generally, they prefer to have a certain number of electrons in their outer shell - no more, no less. This number is typically eight. We call this the octet rule.

• This article is about the octet rule in chemistry.
• We'll define the octet rule before looking at the octet rule and Lewis diagrams.
• After that, we'll explore some examples of the octet rule.
• We'll then learn about the limitations of the octet rule, including notable exceptions.

What is the octet rule?

Noble gases are historically known as inert gases. Their name gives us a clue about their behavior. They are odorless, colorless, have high ionization energies and low electron affinities, and are generally extremely unreactive. In fact, it was long believed that they didn't react with any other element at all (although we now know that this is incorrect, as you'll see later on in the article).

If we take a look at their electron configurations, we discover an explanation for the relative inertness of noble gases: they all have full outer shells of electrons. For all of the noble gases apart from helium, this means having eight electrons in their outer shell.

Having a full outer shell of electrons makes an atom extremely stable, which is highly desirable. Therefore, other atoms tend to try to gain or lose electrons until they have full outer shells, and because this usually involves having exactly eight valence electrons, this phenomenon is known as the octet rule.

Obeying the octet rule

Thanks to the octet rule, it is easy to predict how certain elements react when they bond. For example:

• Group I metals have one electron in their outer shell. They tend to lose this electron in order to achieve a noble gas electron configuration.
• Group II metals have two electrons in their outer shell. They tend to lose these two electrons in order to achieve a noble gas electron configuration.
• Group VI non-metals have six electrons in their outer shell. They tend to gain two electrons in order to achieve a noble gas electron configuration.
• Group VII halogens have seven electrons in their outer shell. They tend to gain one electron in order to achieve a noble gas electron configuration.

The octet rule and Lewis diagrams

We can use the octet rule to help us predict the likely structure of a molecule when faced with multiple different Lewis diagrams. It is also helpful when drawing Lewis diagrams from scratch.

Here's how you go about it:

1. Work out the molecule's total number of valence electrons.
2. Draw the rough position of the atoms in the molecule.
3. Join the atoms using single covalent bonds.
4. Add electrons to the outer atoms until they have full outer shells of electrons.
5. Count up how many electrons you have added, and subtract this from the molecule's total number of valence electrons that you calculated earlier. This tells you how many electrons you have left.
6. Add the remaining electrons to the central atom.
7. Use lone pairs of electrons from the outer atoms to form double covalent bonds with the central atom until all atoms have complete outer shells.

We'll start by applying this process to carbon dioxide, CO₂. First of all, we need to count up the molecule's number of valence electrons. Carbon is in group IV and so has four valence electrons. Oxygen is in group VI and so has six valence electrons. The molecule's total number of valence electrons is therefore 4 + 2(6) = 16.

We then draw the rough position of the atoms in the molecule and add single covalent bonds between them. Here, we draw a central carbon atom joined to two outer oxygen atoms using single covalent bonds.

Carbon dioxide. Anna Brewer, StudySmarter Original

Next, we use our first application of the octet rule. Each atom wants to have eight valence electrons, giving it a full outer shell. We start by looking at the outer atoms. In this case, these are the two oxygen atoms. Both currently only have two valence electrons from the covalent bond that they share with carbon. To get to a full outer shell, each oxygen needs to gain six more electrons. Let's draw them in.

If we add up the total number of valence electrons we've added to the molecule, we find 2(2) = 4 electrons from the two single bonds, and 6(2) = 12 electrons from the lone pairs. 12 + 4 = 16, which you might remember is the number of valence electrons that carbon dioxide is allowed to have. We can't add anymore electrons. However, our Lewis structure isn't complete. This is where the octet rule comes in again. The central carbon atom currently only has four electrons in its outer shell; to achieve a full outer shell and satisfy the octet rule, it needs eight. To give it four extra electrons, we use a lone pair of electrons from each oxygen atom to form two C=O double bonds.

Carbon dioxide. Anna Brewer, StudySmarter Original

All atoms now satisfy the octet rule. Our structure is complete.

Octet rule limitations and exceptions

Although the octet rule is a great model, it doesn't always hold up. In fact, it has some notable exceptions.

Odd number of electrons

Some molecules have odd numbers of electrons. Because of this, it is impossible for them to obey the octet rule. These include free radicals such as nitric oxide and nitrogen dioxide.

Nitric oxide, left, and nitrogen dioxide, right. Anna Brewer, StudySmarter Original

Incomplete octets

Other atoms, particularly the two members of period 1 and the smaller members of group III, are able to form molecules with incomplete octets. Hydrogen and lithium are both stable with just two valence electrons. When they bond, they take the electron configuration of helium, 1s². The outer electron shell contains just an s-subshell, meaning it only has space for a single electron pair. Therefore, having just two valence electrons satisfies their desire to have a full outer shell.

Boron and aluminium are also able to form stable molecules with incomplete octets. Take boron trifluoride, BF₃. It consists of a central boron atom joined to three outer fluorine atoms by single covalent bonds. The fluorine atoms have complete octets, but boron does not: it only has six electrons in its outer shell. However, boron finds this arrangement stable and so doesn't react any further. Aluminium trichloride acts in much the same way.

Molecules with incomplete octets: boron trifluoride and aluminium chloride. Anna Brewer, StudySmarter Original

Expanded octets

Some atoms, specifically those in period three and beyond, are able to form expanded octets. This means that the atom has more than eight electrons in its outer shell. These electrons go in the d-subshell, which is why elements in periods 1 and 2 can't form expanded octets - they don't have a d-subshell. This is also why some noble gases can form bonds. As we've explored, noble gases are generally unreactive because they already have a full outer shell of electrons. However, noble gases in period 3 and beyond can form bonds with other atoms; the extra bonded electrons go in the d-subshell.

One example of a molecule with an expanded octet is phosphorus pentachloride. This molecule consists of a central phosphorus atom bonded to five chlorine atoms using single covalent bonds. Phosphorus has ten electrons in its valence shell, giving it an expanded octet.

A molecule with an expanded octet: phosphorus pentachloride. Image credit: commons.wikimedia.org

Another example is xenon tetrafluoride, consisting of a central xenon atom joined to four fluorine atoms with single covalent bonds. It has twelve electrons in its outer shell.

That's it for this article. With any luck, you should now understand what we mean by the octet rule and be able to draw Lewis diagrams using the octet rule for guidance. You should also be able to identify exceptions to the octet rule.