Covalent Bonding and Coordinate (Dative Covalent) Bonding

Full Notes

Covalent Bonding

Covalent bonding is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.

Each atom shares one or more electrons to achieve a full outer shell. The shared electrons are electrostatically attracted to both nuclei, creating a strong bond.

Sometimes, two atoms share more than one pair of electrons, forming double and triple bonds:

• Double bonds = two shared pairs (e.g. O=O, CO₂)
• Triple bonds = three shared pairs (e.g. N≡N)

These bonds are stronger and shorter than single bonds.

Covalent Bonding in Common Molecules

Single bonds (1 shared pair)

• H₂: each H shares 1 electron → H–H
• Cl₂: each Cl shares 1 electron → Cl–Cl
• HCl: H and Cl share 1 electron → H–Cl
• CH₄: C shares 4 electrons with 4 H atoms → 4 single bonds
• NH₃: N shares 3 electrons with 3 H atoms
• C₂H₆ (ethane): all single C–C and C–H bonds

Double bonds (2 shared pairs):

• O₂: O=O → each O shares 2 electrons
• CO₂: O=C=O → two double bonds
• C₂H₄ (ethene): double bond between C atoms

Triple bonds (3 shared pairs):

• N₂: each N shares 3 electrons → N≡N (very strong bond)

Expanded Octets in Period 3 Elements

Some atoms in Period 3 and below can have more than 8 electrons in their outer shell.

ExamplesCommon examples include:

• SO₂: S forms 2 double bonds with O → 10 electrons
• PCl₅: P forms 5 single bonds → 10 outer shell electrons
• SF₆: S forms 6 single bonds → 12 outer shell electrons

Coordinate (Dative Covalent) Bonding

A dative covalent bond is a type of covalent bond where both bonding electrons come from the same atom.

Once formed, it is identical to a normal covalent bond in strength and length. Represented using an arrow (→) from the donor to the acceptor atom.

Examples

• Ammonium ion (NH₄⁺): NH₃ donates a lone pair on N to H⁺, forming NH₄⁺.

  

• Aluminium Chloride (Al₂Cl₆): contains coordinate bonds between Al and Cl atoms.

  

σ and π Bonds

There are different ways orbitals can overlap to form covalent bonds.

σ (sigma) bonds: formed by direct overlap of orbitals along the bond axis. Present in all covalent bonds.

π (pi) bonds: formed by sideways overlap of adjacent p orbitals. Found in double and triple bonds.

Examples of σ and π Bonds

• H₂: 1 σ bond
• C₂H₆: C–C contains 1 σ bond
• C₂H₄: C=C contains 1 σ and 1 π bond
• HCN: C≡N includes 1 σ and 2 π bonds
• N₂: N≡N has 1 σ and 2 π bonds

Hybridisation: sp, sp², and sp³

Hybridisation is the mixing of atomic orbitals to form new, equivalent hybrid orbitals.

Carbon is a commonly used as an example to show hybridisation.

Hybridisation helps explain why atoms bond as they do (e.g. carbon forms four covalent bonds).

Bond Energy and Bond Length

Bond energy: the energy required to break 1 mole of a specific bond in the gaseous state.

Bond length: the average distance between nuclei of two bonded atoms.

General trend: shorter bonds = stronger bonds = higher bond energy.

Order of strength: Triple bond > Double bond > Single bond.

Matt’s exam tip

Remember that double bonds aren’t twice as strong as single bonds. A double bond is a sigma + pi bond. The pi bond is weaker since its electrons are further from the positively charged nuclei, making it easier to break.

Molecules with lower bond energies are generally more reactive, as their bonds are easier to break. Longer bonds are usually weaker and more reactive.

Summary

• Covalent bonding = attraction between nuclei and shared electrons.
• Single, double, triple bonds differ in strength and length.
• Period 3 elements can expand their octet.
• Coordinate bonding = both electrons from one atom.
• σ bonds from direct overlap, π bonds from sideways overlap.
• Hybridisation explains bonding (sp, sp², sp³).
• Shorter bonds = stronger; higher bond energy = less reactive.