Covalent Bonding and Dative Covalent Bonds — AQA A-level Chemistry

SPEC 3.1.3.2 • BONDING

Covalent & Dative Covalent Bonding

Shared electron pairs, multiple bonds, and coordinate bonding

What is Covalent Bonding?

Definition

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

Covalent bonds form between non-metal atoms. Unlike ionic bonding where electrons are transferred, in covalent bonding electrons are shared between atoms. No electrons are transferred — only shared.

The positive nucleus of each atom has an attraction for the bonding electrons shared in the covalent bond. This electrostatic attraction holds the atoms together.

Why Do Atoms Share?

Sharing electrons allows each of the bonded atoms to achieve an electron configuration similar to a noble gas. This makes each atom more stable. By sharing, both atoms can count the shared electrons as part of their outer shell.

AQA Exam Tip

The definition must include: (1) electrostatic attraction, (2) shared pair of electrons, and (3) nuclei of both atoms. Missing any part loses marks!

Dot and Cross Diagrams

Dot and cross diagrams are used to represent covalent bonding. They show how electrons are shared between atoms and help us track which atom each electron originally came from.

They show just the outer shell of the atoms involved. Electrons are shown in pairs on dot-and-cross diagrams.

Key Principle

To differentiate between the two atoms involved, electrons from one atom are shown as dots (•) and electrons from the other atom are shown as crosses (×). This shows the origin of each electron in the bond.

What to Show

Bonding pairs: Electrons shared between atoms (shown in overlap region)

Lone pairs: Electrons not involved in bonding (stay on original atom)

Type of BondRepresentationNumber of Electrons Shared

SingleC—C2

DoubleC=C4

TripleC≡C6

AQA Exam Tip

Always use dots for ONE atom and crosses for the OTHER. Never mix them on the same atom. The circles represent electron shells and should overlap where bonding occurs.

Single Covalent Bonds: Hydrogen (H₂)

A single covalent bond contains one shared pair of electrons (2 electrons). Each atom contributes one electron to the bond.

Each hydrogen has 1 electron in its outer shell. The outer shell (n=1) has a maximum capacity of 2 electrons, so by sharing, both H atoms achieve a full outer shell.

Chlorine Molecule (Cl₂)

Chlorine has 7 outer electrons and needs 1 more to complete its octet. Two Cl atoms each contribute one electron to form a shared pair.

After bonding, each chlorine atom has 8 electrons in its outer shell — this makes for a stable molecule.

Water Molecule (H₂O)

Oxygen has 6 outer electrons and needs 2 more. It forms 2 single bonds with hydrogen atoms.

Common Mistake

Don't forget the 2 lone pairs on oxygen! Total: 2 bonding pairs (4e⁻) + 2 lone pairs (4e⁻) = 8 electrons around O.

Ammonia Molecule (NH₃)

Nitrogen has 5 outer electrons and needs 3 more. It forms 3 single bonds with hydrogen atoms.

Methane Molecule (CH₄)

Carbon has 4 outer electrons and needs 4 more. It forms 4 single bonds with hydrogen atoms.

Ethane Molecule (C₂H₆)

Ethane has a C—C single bond between two carbon atoms, with each carbon also bonded to 3 hydrogen atoms.

Key Point

The C—C bond is formed by sharing one electron from each carbon atom. This is the basis of all organic chemistry!

Worked Example01

Draw the dot-cross diagram for phosphine (PH₃), showing all bonding pairs and lone pairs.

1

Count outer electrons: P has 5 (Group 15), each H has 1

2

P needs 3 more electrons → forms 3 single bonds

3

P uses 3 electrons in bonds, leaving 1 lone pair (2 electrons)

P has 8 electrons: 3 bonding pairs (6e⁻) + 1 lone pair (2e⁻). PH₃ has the same structure as NH₃.

Double Covalent Bonds

Definition

A double bond contains two shared pairs of electrons (4 electrons total).

Electron Count

Each O has 8 electrons: 4 in bonding pairs (double bond) + 4 in lone pairs (2 lone pairs). Once bonded, each oxygen atom's octet is complete.

Important

Double bonds are stronger than single bonds because more electrons are shared between the atoms, creating a greater electrostatic attraction.

Triple Covalent Bonds

Definition

A triple bond contains three shared pairs of electrons (6 electrons total). Each nitrogen atom shares 3 electrons to form the triple bond.

Each nitrogen atom has 5 electrons in its outer shell. It requires 3 more electrons to complete its octet, which it achieves by forming a triple bond.

Common Mistake

N has 5 outer electrons: 3 in bonding pairs + 2 as lone pair = 5. After bonding, each N effectively has 8 electrons (3 shared pairs + 1 lone pair).

Bond Strength & Length

The N≡N triple bond is one of the strongest bonds in chemistry (945 kJ mol⁻¹). This is why N₂ is so unreactive.

Bond Strength & Bond Length

Key Relationship

The more electrons shared between atoms, the shorter and stronger the bond.

Why? More shared electrons means greater electron density between the nuclei. This increases the electrostatic attraction, pulling the atoms closer together.

Bond TypeElectrons SharedRelative LengthRelative StrengthExample

Single2LongestWeakestC—C (154 pm)

Double4ShorterStrongerC=C (134 pm)

Triple6ShortestStrongestC≡C (120 pm)

AQA Exam Tip

Remember: More shared electrons → Shorter bond → Stronger bond. Triple bonds are hardest to break (highest bond enthalpy).

Worked Example02

Draw the dot-cross diagram for carbon dioxide (CO₂).

1

Outer electrons: C has 4, each O has 6

2

C needs 4 more; each O needs 2 more

3

C forms 2 double bonds (one with each O)

Line notation: O = C = O

In CO₂, each oxygen atom has 2 non-bonded electron pairs — these are called 'lone pairs'. 4 electrons are shared between the carbon atom and oxygen in each C=O double covalent bond.

Dative Covalent Bonds

Definition

A dative covalent bond (or coordinate bond) is a covalent bond in which both electrons in the shared pair come from the same atom.

For a dative bond to form:

Donor atom: Must have a lone pair of electrons that can be donated

Acceptor atom: Must be electron-deficient — an atom with an unfilled outer orbital

In dative covalent bonding, both electrons in the covalent bond are shared by one atom.

Normal Covalent Bond

A — B

Each atom contributes 1 electron

Dative Covalent Bond

A → B

Arrow points FROM donor TO acceptor

Drawing Convention

A dative covalent bond is drawn using an arrow from the donated pair of electrons to the electron-deficient atom. The arrow shows the direction electrons are donated.

Important

Once formed, a dative bond is identical to a normal covalent bond. The arrow just shows where the electrons originally came from.

Worked Example03

Draw the dot-cross diagram for carbon monoxide (CO) and explain its bonding.

1

Outer electrons: C has 4, O has 6

2

C needs 4 more electrons; O needs 2 more

3

C and O share 2 pairs (double bond) → C still needs 2 more

4

O donates a lone pair to C → forms a dative bond

5

Result: Triple bond (2 covalent + 1 dative)

Line notation: C ≡ O or C ← O (showing dative)

The triple bond consists of 2 normal covalent bonds (each atom contributes 1 electron) plus 1 dative bond (both electrons from oxygen's lone pair).

Why CO is Special

CO demonstrates that dative and covalent bonding can occur in the same molecule. The dative bond allows carbon to achieve 8 outer electrons despite only having 4 to share.

Formation of Ammonium Ion (NH₄⁺)

The lone pair on nitrogen in ammonia can be donated to a proton (H⁺) to form an ammonium ion.

The hydrogen ion, H⁺, is electron-deficient and has space for two electrons in its shell. The nitrogen atom in ammonia has a lone pair of electrons which it can donate to the hydrogen ion to form a dative covalent bond.

Dative Bond

In NH₄⁺, the fourth N-H bond is a dative covalent bond where both electrons in the shared pair come from nitrogen's lone pair. When bonded to a 4th hydrogen, nitrogen now has a positive charge.

Aluminium Chloride (Al₂Cl₆)

Aluminium chloride is also formed using dative covalent bonding.

At high temperatures, aluminium chloride can exist as a monomer (AlCl₃). The molecule is electron-deficient and needs two electrons to complete the aluminium atom's outer shell.

At lower temperatures, two molecules of AlCl₃ join together to form a dimer (Al₂Cl₆). Two of the chlorine atoms donate a lone pair to each aluminium atom, forming two dative covalent bonds. This is more stable than AlCl₃.

Exam Tip

Al has only 3 outer electrons, so it's electron deficient. The two bridging Cl atoms each donate a lone pair to complete Al's octet via dative bonds.

Octet Rule & Exceptions

Octet Rule

Atoms tend to achieve 8 electrons in their outer shell. Hydrogen needs only 2.

In some instances, the central atom of a covalently bonded molecule can accommodate more or less than 8 electrons in its outer shell.

Being able to accommodate more than 8 electrons in the outer shell is known as 'expanding the octet rule'

Accommodating less than 8 electrons in the outer shell means the central atom is 'electron deficient'

Exception TypeExamplesElectrons on Central Atom

Electron deficientBF₃, AlCl₃6 electrons (less than 8)

Expanded octetSO₂, PCl₅, SF₆10 (S in SO₂), 10 (P in PCl₅), 12 (S in SF₆)

Odd electronNO, NO₂Odd total electrons cannot all pair

AQA Key Point

Only elements in Period 3 and below can expand their octet (they have d orbitals available). Period 2 elements (C, N, O, F) never exceed 8 electrons.

Practice

Question 1: Draw the dot-cross diagram for hydrogen chloride (HCl).

H: 1 outer electron (blue dot) | Cl: 7 outer electrons (red crosses)

Both need 1 more electron → share one pair to complete their shells

Both atoms achieve stable configurations: H has 2e⁻, Cl has 8e⁻

Practice

Question 2: Explain why BF₃ can act as an acceptor in dative bond formation.

•

Boron has 3 outer electrons

•

It forms 3 covalent bonds with F atoms

•

This gives boron only 6 electrons in its outer shell

•

BF₃ is electron deficient - no octet

•

Boron has an empty p orbital

•

A lone pair donor (e.g., NH₃) can donate into this empty orbital

Practice

Question 3: State the difference between a covalent bond and a dative covalent bond. How can you distinguish between them once formed?

Difference:

•

Covalent: each atom contributes one electron

•

Dative: both electrons come from one atom

Once formed:

•

They are indistinguishable

•

Same bond length and strength

•

Arrow (→) only shows electron origin

Summary

Key Definitions

Bond TypeElectrons SharedNotation

Single bond1 pair (2e⁻)—

Double bond2 pairs (4e⁻)=

Triple bond3 pairs (6e⁻)≡

Dative bond1 pair (both from donor)→

AQA Exam Skills

Draw dot-cross diagrams with overlapping circles

Use dots for one atom, crosses for another

Show dative bonds with arrow from donor to acceptor

Identify electron-deficient molecules as acceptors