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9.1: Lewis Electron Dot Structures

Difficulty Level: At Grade Created by: CK-12
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Lesson Objectives

  • Describe how a covalent bond forms, including the energy change involved in the process.
  • Use the octet rule to draw Lewis electron dot structures for simple molecules. Know how and when to incorporate double and triple bonds into the structures.
  • Understand how a coordinate covalent bond differs from other covalent bonds.
  • Be able to draw Lewis structures for polyatomic ions.
  • Understand the concept of resonance.
  • Know some common exceptions to the octet rule.
  • Relate bond energy to the stability and reactivity of molecules.

Lesson Vocabulary

  • bond energy
  • coordinate covalent bond
  • covalent bond
  • diatomic molecule
  • double covalent bond
  • Lewis electron dot structures
  • lone pair
  • resonance
  • single covalent bond
  • structural formula
  • triple covalent bond

Check Your Understanding

Recalling Prior Knowledge

  • What is a molecule?
  • What is shown in an electron dot diagram?
  • What are polyatomic ions?

In previous chapters, you have learned about several forms that matter can take. Pure metals exist as extended three-dimensional structures of close packed metal cations with mobile valence electrons. Ionic compounds are also crystalline in nature, but with alternating cations and anions held together by attractive electrostatic forces called ionic bonds. The noble gases are monatomic and exist as individual atoms. In this chapter, you will learn about the structure and bonding that occurs in molecular compounds.

The Covalent Bond

Energy and Bond Formation

Molecular compounds are those that take the form of individual molecules. A molecule is generally comprised of two or more nonmetal atoms. Familiar examples include water (H2O), carbon dioxide (CO2) and ammonia (NH3). Recall that a molecular formula shows the quantity of each atom that occurs in a single molecule of that compound. One molecule of water contains two hydrogen atoms and one oxygen atom. Hydrogen (H2) is an example of an element that exists naturally as a diatomic molecule. A diatomic molecule is a molecule that contains exactly two atoms.

Nature favors chemical bonding because most atoms attain a lower potential energy when they are bonded to other atoms than when they are isolated. Consider two hydrogen atoms that are separated by a distance large enough to prevent any interaction between them. At this distance, the potential energy of the system is said to be equal to 0 (Figure below).

The graph shows how the potential energy of two hydrogen atoms changes as a function of their separation distance. The potential energy is zero when they are completely isolated from one another. The energy reaches its minimum at the ideal bond distance and increases rapidly when the atoms come closer because of nuclear repulsion.

As the atoms approach one another, their electron clouds gradually begin to overlap, giving rise to several new interactions. For example, the single electrons possessed by each hydrogen atom begin to repel each other, causing the potential energy of the system to increase. However, attractive forces also begin to develop between each electron and the positively charged nucleus of the other atom, causing a decrease in potential energy.

As the atoms first begin to interact, the attractive force is stronger than the repulsive force, so the potential energy of the system decreases, as seen in Figure above. Remember, a lower potential energy is indicative of a more stable system. As the two hydrogen atoms move closer and closer together, the potential energy continues to decrease. Eventually, a position is reached where the potential energy is at its lowest possible point. If the hydrogen atoms move any closer together, the repulsive force between the two positively charged nuclei (a third type of interaction) begins to dominate. When like charges are forced this close together, the resulting repulsive force is very strong, as can be seen by the sharp rise in energy at the far left of the diagram.

The point at which the potential energy reaches its minimum represents the ideal distance between hydrogen atoms for a stable chemical bond to occur. This type of chemical bond is called a covalent bond. A covalent bond is a bond in which two atoms share one or more pairs of electrons. The single electrons from each of the two hydrogen atoms are shared when the atoms come together to form a hydrogen molecule (H2).

Lewis Electron-Dot Structures

In a previous chapter, you learned that the valence electrons of an atom can be shown in a simple way with an electron dot diagram. A hydrogen atom is shown as H• because of its one valence electron. The structures of molecules that are held together by covalent bonds can be diagrammed by Lewis electron-dot structures. The hydrogen molecule is shown in Figure below.

On the left is a single hydrogen atom with one electron. On the right is an H2 molecule showing the electron cloud overlap. The shared pair of electrons in the covalent bond can be shown in a Lewis structure by either a pair of dots or a dash.

The shared pair of electrons is shown as two dots in between the two H symbols (H:H). This is called a single covalent bond, when two atoms are joined by the sharing of one pair of electrons. The single covalent bond can also be shown by a dash in between the two symbols (H–H). A structural formula is a formula that shows the arrangement of atoms in a molecule and represents covalent bonds between atoms by dashes.

The Octet Rule and Covalent Bonds

When ions form, they conform to the octet rule by either losing or gaining electrons in order to achieve the electron configuration of the nearest noble gas. Similarly, nonmetal atoms share electrons by forming covalent bonds in such a way that each of the atoms involved in the bond can attain a noble-gas electron configuration. The shared electrons are “counted” for each of the atoms involved in the sharing. For hydrogen (H2), the shared pair of electrons means that each of the atoms is able to attain the electron configuration of the noble gas helium, which has two electrons.

For atoms other than hydrogen, the sharing of electrons will usually provide each of the atoms with eight valence electrons.

Single Covalent Bonds

A covalent bond forms when two singly occupied orbitals overlap with each other. For the hydrogen molecule, this can be shown as:

The shared electrons in the complete H2 molecule must have opposite spins, so they are also shown with opposite spins in the atomic 1s orbitals.

The halogens also form single covalent bonds to produce diatomic molecules. An atom of any halogen, such as fluorine, has seven valence electrons. Fluorine's unpaired electron is located in the 2p orbital.

Unpaired electrons in 2p orbitals from two adjacent fluorine atoms combine to form a covalent bond (Figure below).

On the left is a fluorine atom with seven valence electrons, and on the right is the F2 molecule. The shared electrons that form the single covalent bond are from the 2p sublevel.

The diatomic fluorine molecule (F2) contains a single shared pair of electrons. Each F atom also has three pairs of electrons that are not shared with the other atom. A lone pair is a pair of electrons in a Lewis electron-dot structure that is not shared between atoms. Each F atom has three lone pairs. When combined with the two electrons in the covalent bond, each F atom effectively has eight valence electrons, so both atoms follow the octet rule.

Sample Problem 9.1: Lewis Electron Dot Structures

Draw the Lewis electron dot structure for water.

Step 1: List the known quantities and plan the problem


  • molecular formula of water = H2O
  • 1 O atom = 6 valence electrons
  • 2 H atoms = 2 × 1 = 2 valence electrons
  • total number of valence electrons = 8

Use the periodic table to determine the number of valence electrons for each atom and the total number of valence electrons in the entire molecule. Arrange the atoms and distribute the electrons so that each atom follows the octet rule. The oxygen atom will have 8 electrons, while the hydrogen atoms will each have 2.

Step 2: Solve

The electron dot diagram for each atom is:

Each hydrogen atom will use its single electron to form a covalent bond with one of the unpaired electrons on the oxygen atom. The resulting Lewis electron dot structure is:

Step 3: Think about your result.

The oxygen atom follows the octet rule with two pairs of bonding electrons and two lone pairs. Each hydrogen atom follows the octet rule with one bonding pair of electrons.

Practice Problems
  1. Draw the Lewis electron dot structure for each atom.
    1. NH3
    2. CH4
    3. CH2Cl2

Multiple Covalent Bonds

Some molecules are not able to satisfy the octet rule by making only single covalent bonds between the atoms. Consider the compound ethene, which has a molecular formula of C2H4. The carbon atoms are bonded together, and each carbon is also bonded to two hydrogen atoms.

two C atoms = 2 × 4 = 8 valence electrons
four H atoms = 4 × 1 = 4 valence electrons
total of 12 valence electrons in the molecule

If the Lewis electron dot structure was drawn with a single bond between the carbon atoms and lone pairs were then added until all atoms satisfied the octet rule, it would look like this:

However, this Lewis structure is incorrect because it contains 14 valence electrons, and the atoms in this molecule only have a total of 12 valence electrons available. The Lewis structure can be corrected by eliminating one lone pair and moving another lone pair to a bonding position. Now, the two carbon atoms share two pairs of electrons instead of just one.

A double covalent bond is a covalent bond formed by atoms that share two pairs of electrons. The double covalent bond that occurs between the two carbon atoms in ethene can also be represented by a structural formula or a molecular model as shown in Figure below.

(A) The structural model for C2H4 consists of a double covalent bond between the two carbon atoms and single bonds to each of the hydrogen atoms. (B) Molecular model of C2H4.

Similarly, a triple covalent bond is a covalent bond formed by atoms that share three pairs of electrons. In its pure form, the element nitrogen exists as a diatomic gas. The majority of Earth’s atmosphere is made up of N2 molecules. A nitrogen atom has five valence electrons, which can be shown as one pair and three unpaired electrons. When combining with another nitrogen atom to form a diatomic molecule, the three single electrons on each atom combine to form three shared pairs of electrons.

Each nitrogen atom has one lone pair of electrons and six electrons that are shared with the other atom, so each atom obeys the octet rule.

Practice Problems
  1. Draw Lewis electron dot structures for the molecules below, each of which contains one or more multiple covalent bonds.
    1. CO2
    2. C2H2

Coordinate Covalent Bonds

For all of the covalent bonds that we have looked at so far, each of the atoms involved in the bond has contributed one electron to each shared pair. However, there is an alternate type of covalent bond in which one of the atoms provides both of the electrons in a shared pair. Carbon monoxide, CO, is a toxic gas that is released as a by-product during the burning of fossil fuels. The bonding between the C atom and the O atom can be thought of as follows:

At this point, a double bond has formed between the two atoms, with each atom providing one of the electrons to each bond. The oxygen atom now has a stable octet of electrons, but the carbon atom only has six electrons and is unstable. This situation is resolved if the oxygen atom contributes one of its lone pairs in order to make a third bond with the carbon atom.

The carbon monoxide molecule is correctly represented by a triple covalent bond between the carbon and oxygen atoms. One of the bonds is considered a coordinate covalent bond, which is a covalent bond in which one of the atoms contributes both of the electrons in the shared pair.

Once formed, a coordinate covalent bond is the same as any other covalent bond. The two "conventional" bonds in the CO molecule are not stronger or different in any other way from the coordinate covalent bond.

Polyatomic Ions

Recall that a polyatomic ion is a group of covalently bonded atoms that carries an overall electrical charge. For example, the ammonium ion (NH4+) is formed when a hydrogen ion (H+) attaches to the lone pair of an ammonia (NH3) molecule via a coordinate covalent bond.

When drawing the Lewis structure of a polyatomic ion, the charge of the ion is reflected in the total number of valence electrons in the structure. In the case of the ammonium ion:

1 N atom = 5 valence electrons
4 H atoms = 4 × 1 = 4 valence electrons
subtract 1 electron to give the ion an overall charge of 1+
total of 8 valence electrons in the ion

It is customary to put the Lewis structure of a polyatomic ion into a large set of brackets, with the charge of the ion as a superscript outside the brackets.

Sample Problem 9.2: Lewis Electron Dot Structure of a Polyatomic Ion

Draw the Lewis electron dot structure for the sulfate ion.

Step 1: List the known quantities and plan the problem


  • molecular formula of the sulfate ion = SO42−
  • 1 S atom = 6 valence electrons
  • 4 O atoms = 4 × 6 = 24 valence electrons
  • add 2 electrons to give the ion an overall charge of 2−
  • total of 32 valence electrons

The less electronegative sulfur atom is the central atom in the structure. Place the oxygen atoms around the sulfur atom, each connected to the central atom by a single covalent bond. Distribute lone pairs to each oxygen atom in order to satisfy the octet rule. Count the total number of electrons. If there are too many electrons in the structure, make multiple bonds between the S and O.

Step 2: Solve

Step 3: Think about your result.

The Lewis structure for the sulfate ion consists of a central sulfur atom with four single bonds to oxygen atoms. This yields the expected total of 32 electrons. Since the sulfur atom started with six valence electrons, two of the S-O bonds are coordinate covalent.

Practice Problems
  1. Draw the Lewis structure for the chlorate ion, ClO3. Chlorine is the central atom.


There are some cases in which more than one viable Lewis structure can be drawn for a molecule. An example is the ozone (O3) molecule. By distributing a total of 18 valence electrons in a way that allows each atom to satisfy the octet rule, both of the following structures can be drawn.

The structure on the left can be converted to the structure on the right by shifting electrons without altering the positions of the atoms.

It was once thought that the structure of a molecule such as O3 consisted of one single bond and one double bond that shifted rapidly back and forth, as shown above. However, further studies showed that the two bonds are identical. Additionally, the properties of each bond are in between those expected for a single bond and a double bond between two oxygen atoms. For example, a double covalent bond between two given atoms is typically stronger and shorter than a single covalent bond between those two atoms. Studies have shown that the two identical bonds in O3 are stronger and shorter than a typical O-O single bond but longer and weaker than an O-O double bond.

Resonance is the use of two or more Lewis structures to represent the covalent bonding in a molecule. Each of the valid structures is referred to as a resonance structure. It is now understood that the true structure of a molecule that displays resonance is an average or a hybrid of all the resonance structures. In the case of the O3 molecule, each of the covalent bonds between O atoms is best thought of as being “one and a half” bonds, as opposed to either a pure single bond or a pure double bond. This “half-bond” can be shown as a dotted line in both the Lewis structure and the molecular model.

Many polyatomic ions also display resonance. In some cases, the true structure may be an average of three valid resonance structures, as in the case of the nitrate ion, NO3.

The bond lengths between the central N atom and each O atom are identical, and each bond can be approximated as a "one and one-third" bond.

Exceptions to the Octet Rule

It is said that all rules are made to be broken, and this saying certainly applies to the octet rule. Exceptions to the octet rule generally fall into one of three categories: (1) an incomplete octet, (2) odd-electron molecules, and (3) an expanded octet.

Incomplete Octet

In some compounds, the number of electrons surrounding the central atom in a stable molecule is fewer than eight. Beryllium is an alkaline earth metal, so you might expect it to form ionic bonds. However, due to its small size and relatively high ionization energy (compared to other metals), beryllium forms primarily molecular compounds when combined with many other elements. Since beryllium only has two valence electrons, it does not typically attain a full octet by sharing electrons. The Lewis structure of gaseous beryllium hydride (BeH2) consists of two single covalent bonds between Be and two H atoms.

Boron and aluminum, with three valence electrons, also tend to have an incomplete octet when they form covalent compounds. The central boron atom in boron trichloride (BCl3) has six valence electrons as shown below.

Odd-Electron Molecules

There are a number of molecules whose total number of valence electrons is an odd number. It is not possible for all of the atoms in such a molecule to satisfy the octet rule. An example is nitrogen dioxide (NO2). Each oxygen atom contributes six valence electrons and the nitrogen atom contributes five, for a total of seventeen. Possible Lewis structures for NO2 are:

Expanded Octets

Atoms of elements in the second period cannot have more than eight valence electrons around the central atom. However, elements of the third period and beyond are capable of exceeding the octet rule. Starting with the third period, the d sublevel becomes available, so it is possible to use these orbitals in bonding, resulting in more than eight electrons around the central atom.

Phosphorus and sulfur are two elements that react with halogens to make stable compounds in which the central atom has an expanded octet. In phosphorus pentachloride, the central phosphorus atom makes five single bonds to chlorine atoms, so it is surrounded by a total of ten valence electrons. In sulfur hexafluoride, the central sulfur atom has twelve electrons from its six bonds to the fluorine atoms.

Phosphorus pentachloride.

Sulfur hexafluoride.

Bond Energy

As you saw in the first section of this lesson, the formation of a chemical bond results in a decrease in potential energy. Consequently, breaking a chemical bond requires an input of energy. Bond energy is the energy required to break a covalent bond between two atoms. A high bond energy means that a bond is strong, and a molecule containing such a bond is likely to be more stable and less reactive than similar molecules that contain weaker bonds. Reactive compounds often contain at least one bond that has a low bond energy. Some bond energies are listed in Table below.

Bond Bond Energy (kJ/mol)
H–H 436
C–H 414
C–C 347
C=C 620
C≡C 812
F–F 157
Cl–Cl 243
Br–Br 193
I–I 151
N≡N 941

The halogen elements all exist naturally as diatomic molecules (F2, Cl2, Br2, and I2). However, relatively small amounts of energy are required to break these bonds, which makes them very reactive.

Comparing the bond energies for various carbon-carbon bonds, you can see that double bonds are substantially stronger than single bonds, and triple bonds are even stronger. The triple bond that exists between the nitrogen atoms in nitrogen gas (N2) makes it very unreactive. All plants and animals require the element nitrogen, but the direct absorption of nitrogen gas from the atmosphere does not provide the element in a readily usable form, due to its strong, unreactive triple bond. However, some species of bacteria have the ability to convert nitrogen gas into other compounds, such as ammonium and nitrate ions, which are then absorbed by plants from the soil. By eating those plants, animals can obtain nitrogen in a form that can be used by the body to manufacture other more complex molecules.

Lesson Summary

  • The chemical bonding that occurs in molecular compounds is a sharing of valence electrons called covalent bonding. Covalent bonds occur primarily between nonmetal atoms. The formation of a covalent bond between two atoms decreases their potential energy, making them more stable than they were as isolated atoms.
  • Atoms tend to form covalent bonds in ways that satisfy the octet rule. Lewis electron dot structures are drawn to show how the atoms are arranged in a molecule. Covalent bonds may be single, double, or triple, depending on the number of shared electrons.
  • Coordinate covalent bonds occur when one of the atoms involved in the bond contributes both shared electrons.
  • The atoms in a polyatomic ion are held together by covalent bonding. Electrons are added to or subtracted from the structure according to the charge of the ion.
  • Resonance occurs when more than one valid Lewis structure can be drawn for a molecule. The true structure is an average of the possible resonance structures.
  • Exceptions to the octet rule include incomplete octets, odd-electron molecules, and expanded octets.
  • Bond energy is the energy required to break a covalent bond. The most stable and unreactive molecules contain bonds with high bond energies.

Lesson Review Questions

Reviewing Concepts

  1. Identify the compounds below as being most likely ionic or molecular.
    1. CaBr2
    2. PCl3
    3. H2S
    4. ZnO
  2. Describe the difference between ionic and covalent bonds.
  3. Describe the attractive and repulsive forces that occur as two atoms approach one another and form a covalent bond.
  4. How many electrons are shared between atoms in a single covalent bond? In a double covalent bond? In a triple covalent bond?
  5. How does the bond length between the oxygen atoms in an ozone molecule compare to the bond lengths of oxygen-oxygen single and double bonds? Explain.
  6. Which elements are capable of exceeding the octet rule when forming covalent bonds. Why?


  1. Draw Lewis structures for the following molecules, each of which follows the octet rule.
    1. H2S
    2. PCl3
    3. HCN
    4. H2CO
    5. OF2
    6. BrCl
    7. CS2
    8. C2H6
  2. Draw Lewis structures for the following polyatomic ions.
    1. SO32-
    2. OH-
    3. PO43-
  3. Draw all resonance structures for the carbonate ion, CO32−.
  4. Compounds that contain a C–N and/or a C–O bond are capable of forming a coordinate covalent bond with the H+ ion. Compounds that only contain C–C and C–H bonds cannot. Explain this observation.
  5. Draw Lewis structures for the following molecules. Which do not follow the octet rule?
    1. AlH3
    2. SF4
    3. BeCl2
    4. SCl2
    5. NO
    6. SiCl4
  6. The noble gas xenon is capable of making a few compounds with fluorine and oxygen. Draw Lewis structures for the following xenon-containing compounds. In all cases, xenon is the central atom.
    1. XeF4
    2. XeF6
    3. XeOF4

Further Reading / Supplemental Links

Points to Consider

The physical and chemical properties of various substances are dependent upon their structures and bonds. Some of these properties are related to the arrangement of the atoms within a molecule and the resulting molecular geometry.

  • How do the covalent bonds in a molecule allow one to predict molecular shapes?
  • How many different basic shapes are possible?

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