Nature of Covalent Bonding
A concise revision guide to the nature of covalent bonding, shared electron density, orbital overlap, sigma bonds, pi bonds, multiple bonds and dot-and-cross representations for Cambridge International AS & A Level Chemistry.
GCSE Recap: Shared Pairs in Simple Molecules
Before you start, check that you can still count shared and unshared pairs the GCSE way.
What Is a Covalent Bond?
A covalent bond is a strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
In a hydrogen molecule, X-ray diffraction evidence shows a high concentration of negative charge between the two hydrogen nuclei. This shared negative charge is attracted to both nuclei, so the attractive forces are greater than the repulsive forces.
The evidence is drawn as an electron density map. Each contour joins points of equal electron density, in the same way that contour lines on a map join points of equal height, so contours packed tightly between the two nuclei show where the shared charge is concentrated.
The bond is therefore strongest when there is a significant density of shared electrons between the bonded atoms.
Definition: A covalent bond forms when atomic orbitals overlap and a pair of electrons is shared between two atoms. The shared electron density between the nuclei holds the atoms together.
The highest electron density is between the nuclei, which explains the attractive force in the covalent bond.
Quick Check: What Holds F₂ Together?
Apply the definition of a covalent bond to a molecule this page has not described.
Orbital Overlap and Single Covalent Bonds
Covalent bonds form when atomic orbitals overlap. The shared electrons occupy the overlapping region and are attracted by both nuclei.
In hydrogen, two 1s orbitals overlap to form a molecular orbital containing the shared pair of electrons. In chlorine, a pair of p orbitals overlap, with each chlorine atom contributing one unpaired electron.
A molecular orbital is the new orbital produced by the overlap. It belongs to both atoms and holds the shared pair.
The two orbitals that overlap need not be the same type, and the two atoms need not be the same element. An s orbital on one atom can overlap end-on with a p orbital on another atom, and the bond formed is still a single covalent bond. To decide which orbitals overlap, write the electron configuration of each atom and pick out the orbital that holds the unpaired electron.
| Molecule | Orbitals involved | Bond formed |
|---|---|---|
| H2 | Overlap of two 1s orbitals | A single covalent bond with electron density between the two nuclei |
| Cl2 | Overlap of two p orbitals, each containing one unpaired electron | A single covalent bond between the two chlorine atoms |
Sigma and Pi Bonds: End-On and Sideways Overlap
Bring two atoms together and watch their orbitals overlap: end-on between the nuclei for a σ bond, and sideways above and below the plane for a π bond.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
Key idea: The shared pair of electrons is not just drawn between atoms for convenience. It represents electron density in the overlapping region between the nuclei.
Quick Check: Which Orbitals Overlap?
Use electronic configurations to decide which orbitals overlap in three new molecules.
Sigma Bonds and Pi Bonds
A sigma, σ, bond is formed by direct orbital overlap along the line between two nuclei. It is the first covalent bond formed between two atoms.
After a sigma bond has formed, it is sometimes possible for a pi, π, bond to form. A pi bond is formed by the sideways overlap of p orbitals, creating electron density above and below the internuclear axis, the line joining the two nuclei.
In a planar molecule such as ethene that is the same as above and below the plane of the molecule, but it is not the same in every molecule. N≡N is linear and has no molecular plane, and its two pi bonds lie at right angles to each other. Quote the internuclear axis and the answer is right every time.
| Bond type | How it forms | Electron density |
|---|---|---|
| σ bond | Direct overlap of orbitals along the internuclear axis | Concentrated directly between the two nuclei |
| π bond | Sideways overlap of p orbitals after a σ bond has already formed | Above and below the internuclear axis, the line joining the two nuclei |
The sideways overlap of p orbitals produces two regions of electron density, one above and one below the molecular plane.
Bonding in Nitrogen, N2: Sigma and Pi Bonds
Drag to rotate, scroll or pinch to zoom. Build the molecule one layer at a time, from the orbitals each nitrogen brings through to the finished triple bond. The side view looks straight down the molecule and shows both pi bonds at right angles.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
Quick Check: Sigma or Pi?
Decide what type of bond forms when an s orbital overlaps end-on with a p orbital.
Double Bonds, Triple Bonds and Bond Strength
A double bond contains one σ bond and one π bond. For example, the carbon-carbon double bond in ethene contains one sigma bond and one pi bond.
A triple bond contains one σ bond and two π bonds. For example, the nitrogen molecule, N≡N, contains one sigma bond and two pi bonds.
Double and triple bonds place a greater electron density between the nuclei. This increases the electrostatic attraction between nuclei and bonding electrons, giving a shorter bond and a stronger bond.
A π bond is weaker than a σ bond. Sideways overlap is less effective than end-on overlap, so the electron density of a π bond lies further from the two nuclei and is held less strongly. That is why the enthalpy of a C=C bond is less than twice the enthalpy of a C-C bond, and why it is the π bond that breaks in the reactions of alkenes.
| Bond | Bond length / nm | Mean bond enthalpy / kJ mol⁻¹ |
|---|---|---|
| C-C | 0.154 | +347 |
| C=C | 0.134 | +612 |
| C≡C | 0.120 | +838 |
Exam focus: More shared electron pairs usually means greater electron density between nuclei, stronger attraction, shorter bond length and greater bond strength.
Multiple bonds contain more shared electron pairs, which changes both bond length and bond strength.
Quick Check: Count the Sigma and Pi Bonds
Count the σ and π bonds in four molecules and type the numbers.
Quick Check: Explain Bond Length and Strength
Write a short explanation using the data, then compare it with the mark points and the model answer.
Dot-and-Cross Diagrams and Electron Configuration
Dot-and-cross diagrams show how outer shell electrons are shared in covalent bonds. Only the outer shell electrons are normally shown.
Many atoms reach eight electrons in their outer shell when they bond, a stable arrangement often described as the octet rule. Hydrogen is an exception, because it becomes stable with two electrons in its outer shell.
The octet rule is a guide, not a law. Some central atoms finish with fewer than eight outer electrons, and some from Period 3 onwards finish with more, so count the pairs you have actually drawn instead of assuming eight.
Single bond
One shared pair of electrons is shown between the two atoms.
Double bond
Two shared pairs of electrons are shown between the two atoms.
Triple bond
Three shared pairs of electrons are shown between the two atoms.
Remember: Dot-and-cross diagrams show electron sharing, but orbital overlap explains why the shared electrons hold the nuclei together.
Quick Check: Pairs in New Dot-and-Cross Diagrams
Sketch each diagram on paper first, then type the number asked for.
Common Exam Points
Strong covalent bonding answers link electron sharing to electrostatic attraction, orbital overlap and the location of electron density.
Do not say electrons are simply transferred
Covalent bonding involves shared electron pairs. Electron transfer is used to explain ionic bonding.
Use attraction between nuclei and shared electrons
The bond is held by electrostatic attraction between the positively charged nuclei and the shared negative electron density.
Distinguish σ and π bonds clearly
A sigma bond forms by direct overlap along the internuclear axis. A pi bond forms by sideways overlap of p orbitals.
Link multiple bonds to electron density
Double and triple bonds have greater electron density between the nuclei, so they are generally shorter and stronger than single bonds between the same atoms.
Expanding the Octet: SO2, PCl5 and SF6
Elements in Period 2 can never hold more than eight electrons in their outer shell, because the second shell has only the 2s and 2p orbitals. Elements in Period 3 and below have empty 3d orbitals of similar energy, so the central atom can use more than eight electrons in bonding. Cambridge calls this expanding the octet and expects three examples.
| Molecule | Bonds to the central atom | Electrons around the central atom | Shape |
|---|---|---|---|
| SO2 | Two S=O double bonds plus one lone pair on sulfur | 10 (4 bonding pairs + 1 lone pair) | Bent, about 119° |
| PCl5 | Five P-Cl single bonds | 10 (5 bonding pairs) | Trigonal bipyramidal, 120° and 90° |
| SF6 | Six S-F single bonds | 12 (6 bonding pairs) | Octahedral, 90° |
In each case the dot-and-cross diagram shows the central atom sharing all of its outer electrons: sulfur uses all six in SO2 and SF6 (in SO2 two are left as a lone pair), and phosphorus uses all five in PCl5. Nitrogen, in the same group as phosphorus, cannot do this: NCl5 does not exist, although NCl3 does. Phosphorus forms PCl3 as well as PCl5, so expanding the octet is something a Period 3 atom can do, not something it must do.
Period 3 central atoms can hold 10 or 12 outer electrons because empty 3d orbitals are available.
Exam sentence: Sulfur and phosphorus are in Period 3, so they can expand their octet: sulfur has 10 outer electrons in SO2 and 12 in SF6, and phosphorus has 10 in PCl5.
Quick Check: Which Atom Expands Its Octet?
Count the pairs of electrons around each central atom.
QuickSnap
Covalent bonding is explained by shared electron density between bonded atoms. This electron density is attracted to both nuclei, holding the atoms together.
Covalent bond
A shared pair of electrons attracted to the nuclei of both bonded atoms.
Orbital overlap
Atomic orbitals overlap, and the shared electrons occupy the overlapping region.
σ bond
Forms by direct overlap along the internuclear axis, with electron density between the nuclei.
π bond
Forms by sideways overlap of p orbitals, with electron density above and below the internuclear axis. It is weaker than a σ bond.
Multiple bonds
Double bonds contain one σ and one π bond. Triple bonds contain one σ and two π bonds.
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Some ionic radii are shown.
| Ion | Ionic radius / nm |
|---|---|
| Na+ | 0.102 |
| K+ | 0.138 |
| F− | 0.133 |
| Cl− | 0.180 |
Which compound has the strongest ionic bonding?
Explain why the metallic bonding in magnesium is much stronger than that in sodium.
FAQs
Use these quick answers to check the common covalent bonding ideas that appear in Cambridge International AS & A Level Chemistry questions.
What is a covalent bond?
A covalent bond is a strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
Why is electron density important in covalent bonding?
The shared electron density lies between the nuclei and is attracted to both positively charged nuclei. This attraction holds the atoms together.
What is the difference between a sigma bond and a pi bond?
A sigma bond forms by direct orbital overlap along the internuclear axis. A pi bond forms by sideways overlap of p orbitals and has electron density above and below that same axis. A pi bond is weaker than a sigma bond.
How many sigma and pi bonds are in a double bond?
A double bond contains one sigma bond and one pi bond.
How many sigma and pi bonds are in a triple bond?
A triple bond contains one sigma bond and two pi bonds.
Copyright and author footprint: This OLS revision page was written for Online Learning System by Dr. Mohammed Al-Fatah. It is designed for A Level Chemistry revision and should not be copied or redistributed without permission.
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