Covalent bonding
Section: 3 Chemical Bonding | Syllabus: Cambridge AS Level Chemistry (9701)
Covalent Bonding
Covalent Bonding The electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.Covalent bonds form between non-metal atoms. Each atom contributes one electron to a shared pair (bonding pair). The shared electrons are attracted to both nuclei simultaneously, holding the atoms together.
- A single bond = 1 shared pair of electrons
- A double bond = 2 shared pairs of electrons
- A triple bond = 3 shared pairs of electrons
- Lone pairs = non-bonding pairs on an atom
Dot-and-Cross Diagrams for Covalent Molecules
Only outer (valence) shell electrons are shown. Shared pairs sit between the two atoms; lone pairs sit on one atom only.
| Molecule | Bond type | Bonding pairs | Lone pairs on central atom |
|---|---|---|---|
| H₂ | Single H–H | 1 | 0 |
| Cl₂ | Single Cl–Cl | 1 | 3 on each Cl |
| HCl | Single H–Cl | 1 | 3 on Cl |
| O₂ | Double O=O | 2 | 1 on each O (+ 1 unpaired - free radical character) |
| N₂ | Triple N≡N | 3 | 1 on each N |
| CO₂ | Two double C=O | 2 per bond | 2 on each O |
| NH₃ | Three single N–H | 3 | 1 on N |
| CH₄ | Four single C–H | 4 | 0 |
| C₂H₆ (ethane) | C–C and four C–H per C | 7 total | 0 |
| C₂H₄ (ethene) | C=C and two C–H per C | 6 total | 0 |
Figure 3.6: Dot-and-Cross Diagrams for H₂, Cl₂, HCl, O₂, N₂
(Row of five dot-and-cross diagrams. H₂: two H atoms sharing one pair. Cl₂: two Cl atoms with 3 lone pairs each, sharing one bonding pair. HCl: H and Cl sharing one pair, 3 lone pairs on Cl. O₂: two O atoms sharing 2 pairs, 2 lone pairs each. N₂: two N atoms sharing 3 pairs, 1 lone pair each.)
Figure 3.7: Dot-and-Cross Diagrams for NH₃, CH₄, CO₂, H₂O
(NH₃: N central with 3 N–H bonds and 1 lone pair on N. CH₄: C central with 4 C–H bonds, no lone pairs. CO₂: O=C=O with 2 lone pairs on each O. H₂O: O central with 2 O–H bonds and 2 lone pairs on O.)
Figure 3.8: Dot-and-Cross Diagrams for C₂H₆ and C₂H₄
(C₂H₆: two C atoms each bonded to 3 H and to each other by single C–C bond, no lone pairs. C₂H₄: two C atoms joined by a double bond (C=C), each also bonded to 2 H atoms, no lone pairs.)
Expanded Octet - Period 3 Elements
Elements in Period 3 (and beyond) have available 3d orbitals that can be used for bonding, allowing them to accommodate more than 8 electrons in their valence shell - an expanded octet.
| Compound | Central atom | Valence electrons around central atom | Octet expanded? |
|---|---|---|---|
| SO₂ | S | 10 (one double bond to O, one dative bond, one lone pair) | Yes (10 electrons) |
| PCl₅ | P | 10 (five P–Cl bonds) | Yes (10 electrons) |
| SF₆ | S | 12 (six S–F bonds) | Yes (12 electrons) |
Figure 3.9: Dot-and-Cross Diagrams - PCl₅ and SF₆
(PCl₅: P central with 5 Cl atoms bonded by single bonds. Each Cl has 3 lone pairs. P has 10 electrons total - expanded octet. SF₆: S central with 6 F atoms, each F has 3 lone pairs. S has 12 electrons total.)
Key PointElements in Period 2 (C, N, O, F) cannot expand their octet - they have no available d orbitals in the second shell. Only Period 3+ elements (P, S, Cl, etc.) can do so.
Coordinate (Dative Covalent) Bonding
Coordinate (Dative Covalent) Bond A covalent bond in which both electrons of the shared pair are donated by the same atom. The donor atom must have a lone pair; the acceptor atom must have a vacant orbital.Once formed, a coordinate bond is identical to an ordinary covalent bond - the distinction is only in how it was formed. It is represented by an arrow (→) pointing from the donor to the acceptor.
Example 1: Formation of the Ammonium Ion, NH₄⁺
- NH₃ has a lone pair on the nitrogen atom.
- H⁺ (a proton) has a vacant orbital and no electrons.
- The lone pair on N is donated to H⁺, forming a new N→H coordinate bond.
- The resulting NH₄⁺ ion has four identical N–H bonds - all equivalent despite one being dative in origin.
Figure 3.10: Formation of NH₄⁺ by Coordinate Bonding
(Left: NH₃ with lone pair on N and H⁺ with empty orbital. Arrow from lone pair on N to H⁺. Right: NH₄⁺ in square brackets with + charge, all four N–H bonds shown equivalently. The coordinate bond may be shown with an arrow N→H.)
Example 2: Al₂Cl₆ (Aluminium Chloride Dimer)
- AlCl₃ has only 6 electrons around Al - an electron-deficient molecule with a vacant orbital.
- A lone pair from Cl of one AlCl₃ molecule donates into the vacant orbital of Al in another AlCl₃.
- Two coordinate bonds form, creating the dimer Al₂Cl₆ (two bridging Cl atoms).
Figure 3.11: Structure of Al₂Cl₆
(Two AlCl₃ units joined by two bridging Cl atoms. Each Al bonded to 2 terminal Cl and 2 bridging Cl. Coordinate bonds shown as arrows from bridging Cl lone pairs to Al. Al now has 8 electrons - octet complete.)
Sigma (σ) and Pi (π) Bonds
σ Bond (Sigma Bond) Formed by the direct (head-on) overlap of orbitals along the internuclear axis. Electron density is concentrated between the two nuclei. All single bonds are σ bonds. π Bond (Pi Bond) Formed by the sideways overlap of adjacent p orbitals, above and below the σ bond axis. Electron density is concentrated above and below the internuclear axis. π bonds only exist alongside a σ bond.| Molecule | Bond(s) | σ bonds | π bonds |
|---|---|---|---|
| H₂ | H–H single bond | 1 (s-s overlap) | 0 |
| C₂H₆ (ethane) | C–C + C–H | 7 | 0 |
| C₂H₄ (ethene) | C=C + C–H | 5 | 1 |
| N₂ | N≡N triple bond | 1 | 2 |
| HCN | H–C + C≡N | 2 | 2 |
Figure 3.12: σ and π Bond Formation in C₂H₄
(Left: sp² hybridised C atoms. The C–C σ bond shown as direct orbital overlap along the axis. The remaining unhybridised p orbitals (one per C, perpendicular to the plane) shown sideways overlapping above and below the C–C axis to form the π bond. Right: resulting C=C double bond with electron density regions labelled σ (between nuclei) and π (above and below).)
Key PointA double bond = 1 σ + 1 π. A triple bond = 1 σ + 2 π. Free rotation is possible about a σ bond but not about a π bond - the sideways overlap would be broken by rotation, which is why C=C double bonds cause restricted rotation and geometric (E/Z) isomerism.
Hybridisation
Hybridisation is the mixing of atomic orbitals on the same atom to form a new set of equivalent hybrid orbitals suitable for bonding.
sp³ Hybridisation
- One s + three p orbitals combine → four equivalent sp³ orbitals.
- Arranged tetrahedrally (109.5° apart).
- Examples: CH₄, NH₃, H₂O - all sp³ on the central atom.
- Each sp³ orbital forms a σ bond (or holds a lone pair in NH₃/H₂O).
sp² Hybridisation
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