3.1 The Octet Rule and Lewis Dot Diagrams

Learning Objectives

  1. Understand the octet rule and how to read and write Lewis dot diagrams
  2. Utilise Lewis dot diagrams to predict the bonding of simple molecules.

In 2.2 ‘Electronic Configuration’, we discussed how valence electrons are essential to determining the bonding of elements. Generally speaking, elements want to achieve 8 electrons in their outer shell — the same configuration seen in noble gases — and will bond with other elements to make this occur. This rule of thumb is known as the octet rule, and is true for elements up to the [latex]d[/latex] block.

The Lewis electron-dot model is based on the fact that atoms tend to combine to give compounds that achieve a noble gas valence electron configuration of [latex]s^{2}p^{6}[/latex] (or [latex]s^{2}[/latex] for hydrogen). Elements within this model are represented through their chemical symbol and dots, which depict valent [latex]s[/latex] and [latex]p[/latex] electrons. Following the rules of electronic orbitals, electrons are placed surrounding each of the four sides of the element before paring up (see Figure 3.1.1).

Picture showing the Lewis electron-dot diagrams of neutral atoms. Dots represent electrons, and may appear on the four sides of each elemental symbol. The number of dots represents the number of valence electrons. The first 4 electrons cover each side, and then additional electrons pair up with lone electrons. For instance, oxygen has 6 valence electrons, with 2 paired up on 2 sides, with the other 2 sides having one lone electron each.
Figure 3.1.1: Lewis Electron Dot Diagrams. Image attribution: Lewis dot structure of elements 1 to 18. Periodic table © Zizo – stock.adobe.com. Transcript.

To understand how dot diagrams can help us predict bonding, let’s have a look at a compound containing sodium ([latex]\ce{Na}[/latex]) and chlorine ([latex]\ce{Cl}[/latex]). We will begin by determining the number of valence shell electrons in each:

Table 3.1.1: Electron configurations of sodium and chlorine.
Atom Atomic number Electron configuration Valence shell electrons
[latex]\ce{Na}[/latex] 11 [latex]1s^{2}2s^{2}2p^{6}3s^{1}[/latex] 1
[latex]\ce{Cl}[/latex] 17 [latex]1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}[/latex] 7

Sodium contains 1 valence electron, while chlorine has 7. For both to achieve full valence shells, sodium can donate its electron to chlorine.

The atom that can form most bonds is considered the central atom. If two different atoms can form the same number of bonds, then the least electronegative atom should be the central atom. Among sodium and chlorine, sodium is the least electronegative atom. Let us draw the two elements as Lewis Dot Diagrams (see Figure 3.1.2):

Lewis dot diagrams of sodium and chlorine. Sodium has one dot on the right side. Chlorine has 7 dots, with all sides paired up except the left, with one lone electron.
Figure 3.1.2: Sodium and chlorine in Lewis dot diagram.

To create this neutral compound, a single bond is formed between sodium and chlorine to donate electrons. This is depicted by combining the two dots together or by drawing a line (see Figure 3.1.3):

Lewis dot diagram of sodium and chlorine. Sodium and chlorine share two electrons in the middle of each symbol, and chlorine now has a full valence shell.

A lewis dot diagram of sodium and chlorine. A line is drawn between the sodium and chlorine with the two lone dots removed, indicating a bond between the elements.
Figure 3.1.3: Sodium and chlorine form a bond by sharing an electron in a Lewis dot diagram. The bond can be shown as a line or two electrons in Lewis dot diagrams.

Chemistry Is Everywhere: Salt

The element sodium is a very reactive metal; given the opportunity, it will react with the sweat on your hands and form sodium hydroxide, which is a very corrosive substance. The element chlorine is a pale yellow, corrosive gas that should not be inhaled due to its poisonous nature. Bring these two hazardous substances together, however, and they react to make the ionic compound sodium chloride, known simply as salt. Sodium, chlorine, and sodium chloride can be seen in Figure 3.1.4 below:


Three-part image showing common sodium compounds in different states: the first features metallic sodium chunks in a transparent jar with a white lid, the second displays a clear bottle containing yellowish chlorine gas, and the third shows a pile of sodium chloride crystals against a black background.
Figure 3.1.4 “Sodium Chloride.” (a) Sodium is a very reactive metal. (b) Chlorine is a pale yellow, noxious gas. (c) Together, sodium and chlorine make sodium chloride — salt — which is necessary for our survival. Image attribution: “Sodium metal chunks in oil” by Wilco Oelen © CC BY-SA (Attribution-ShareAlike)“Chlorine in bottle” by Wilco Oelen © CC BY-SA (Attribution-ShareAlike). “Salt Crystals” by Mark Schellhase © CC BY-SA (Attribution-ShareAlike).

Salt is necessary for life. Na+ ions are one of the main ions in the human body and are necessary to regulate the fluid balance in the body. Cl ions are necessary for proper nerve function and respiration. Both of these ions are supplied by salt. The taste of salt is one of the fundamental tastes; salt is probably the most ancient flavouring known and one of the few rocks we eat.

The health effects of too much salt are still under debate, although a 2010 report by the US Department of Agriculture concluded that “excessive sodium intake … raises blood pressure, a well-accepted and extraordinarily common risk factor for stroke, coronary heart disease, and kidney disease.”[1] It is clear that most people ingest more salt than their bodies need, and most nutritionists recommend curbing salt intake. Curiously, people who suffer from low salt (called hyponatremia) do so not because they ingest too little salt but because they drink too much water. Endurance athletes and others involved in extended strenuous exercise need to watch their water intake so their body’s salt content is not diluted to dangerous levels.


This completes the diagram for sodium chloride. Let’s have a look at a more complex molecule, ammonia (NH3). Ammonia contains nitrogen and hydrogen. As before, we will begin by considering the valence electrons.

Table 3.1.2: Electron configuration of nitrogen and hydrogen.
Atom Atomic number Electron configuration Valence shell electrons
[latex]\ce{H}[/latex] 1 [latex]{1s^{1}}[/latex] 1
[latex]\ce{N}[/latex] 7 [latex]1s^{2}2s^{2}2p^{3}[/latex] 5

Hydrogen contains 1 valence electron, while nitrogen has 5. To achieve a full valence shell, 3 hydrogens are needed to donate their electrons with nitrogen (see Figure 3.1.5).

Lewis dot diagram of a nitrogen atom in the center surrounded by three hydrogen atoms. Each hydrogen atom has a single dot representing its single valence electron, positioned to illustrate the potential bonding with the nitrogen atom, which has five valence electrons: one lone pair at the top and one electron each ready to form a bond with the surrounding hydrogen atoms.
Figure 3.1.5: Nitrogen and hydrogen in Lewis dot diagram.

To create this neutral compound, a single bond is formed between each of the hydrogen and the lone nitrogen to donate electrons (see Figure 3.1.6):

Lewis dot diagram with a central nitrogen atom connected to three hydrogen atoms. The nitrogen has five valence electrons: two in a pair at the top, and the remaining three electrons are each forming a bond with the surrounding hydrogen. The hydrogen atoms are shown with their single electrons bonded to the nitrogen, and this sharing of electrons has formed stable covalent bonds.
Figure 3.1.6: Nitrogen and hydrogen in Lewis dot diagram, forming bonds through sharing electrons.

Lewis dot diagrams are vital in understanding the geometry of molecules. As we will explore in VSEPR theory — the location of lone pairs of electrons are of great importance. Let us review one more example: carbon dioxide (CO2).

Table 3.1.3: Electron configuration of carbon and oxygen.
Atom Atomic number Electron configuration Valence shell electrons
[latex]\ce{C}[/latex] 6 [latex]1s^{2}2s^{2}2p^{2}[/latex] 4
[latex]\ce{O}[/latex] 8 [latex]1s^{2}2s^{2}2p^{4}[/latex] 6
A lewis dot diagram with carbon in the centre of two oxygens atoms, each on the left and right. The oxygen atoms have six valence electrons, two pairs and two single dots. The carbon has four dots on each side.
Figure 3.1.7: Carbon and oxygen in Lewis dot diagram.

In this compound, carbon is the central atom (see Figure 3.1.7). Carbon requires 4 more electrons to complete its valence shell. We can begin by donating one electron from each neighbouring oxygen atom (see Figure 3.1.8):

A lewis dot diagram with carbon in the centre of two oxygens atoms, each oxygen on the left and right. Each oxygen has two pairs and a single valence electron. Carbon has two single valence electrons above and below, and a line connecting it to each oxygen atom representing single bonds shared between the carbon and two oxygen atoms.
Figure 3.1.8: Carbon and oxygen in Lewis dot diagram, forming a bond through sharing electrons.

After completing this, we can see how all atoms still have unpaired electrons present. Two on the carbon and one on each of the oxygens. For a stable compound to form, all electrons must be paired. This, therefore, must mean that the oxygens are double bonded to the carbon, where two pairs of electrons are transferred/shared between elements (see Figure 3.1.9). Within chemistry, triple bonds are also a possibility.

A lewis dot diagram with carbon in the centre of two oxygens atoms, on the left and to the right. Each oxygen now has two pairs of valence electrons, and is sharing two valence electrons with Carbon, which has double lines connecting it to each oxygen atom representing the double bonds shared between the shared valence electrons.
Figure 3.1.9: Carbon and oxygen in Lewis dot diagram, forming a double bond through sharing additional electrons.

Some molecules do not follow the octet rule. Atoms such as sulphur, phosphorus and chlorine deviate from the octet rule by having more than eight electrons in the valence shell. This is due to the availability of vacant d orbitals in the third shell. These exceptions are fully explored in tertiary chemistry studies.

Key Takeaways

  • Main group elements want to achieve 8 electrons in their outer shell. This is known as the octet rule.
  • Lewis electron-dot diagrams allow for us to visually represent valent electrons and predict bonding.
  • Some molecules require the formation of double and triple bonds to complete their octet.
  • Not all compounds follow the traditional octet rule, such as sulphur and phosphorous.


Practice Questions



  1. U.S. Department of Agriculture Committee for Nutrition Policy and Promotion, “Report of the Dietary Guidelines Advisory Committee on the Dietary Guidelines for Americans,” accessed January 5, 2010, https://www.dietaryguidelines.gov/sites/default/files/2019-05/2010DGACReport-camera-ready-Jan11-11.pdf.


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