# Covalent and coordinate bonding

> CIE A-Level Chemistry · 9701
> Source: https://www.owlsprep.com/study/cie-9701-u3-covalent-and-coordinate-bonding/

This module covers the formation of single, double and triple covalent bonds through electron pair sharing, and explains the structure, properties and common examples of coordinate (dative covalent) bonds for CIE A-Level Chemistry exam questions.

**Prerequisites:** [Atomic structure and electron configurations](https://www.owlsprep.com/study/cie-9701-u2-atomic-structure/); [Ionic bonding](https://www.owlsprep.com/study/cie-9701-u3-ionic-bonding/)

## Learning objectives

- Describe covalent bond formation in terms of electron pair sharing
- Distinguish between single, double and triple covalent bonds by bond order, length and energy
- Explain the formation of coordinate (dative covalent) bonds
- Draw dot-and-cross diagrams for molecules and ions containing covalent and coordinate bonds

## Formation of Covalent Bonds

**Covalent bond** — A chemical bond formed by the electrostatic attraction between two positively charged atomic nuclei and a shared pair of electrons between the atoms. Covalent bonds typically form between two non-metal atoms with similar electronegativity.

*Example:* Bond between two chlorine atoms in Cl₂

Covalent bonds form when the overall potential energy of the bonded atoms is lower than that of the separate atoms. This comes from a balance of attraction between the shared electrons and the two nuclei, and repulsion between the positively charged nuclei.

**Worked example:** Show the formation of a single covalent bond between two chlorine atoms using a dot-and-cross diagram

1. Identify the number of valence electrons for each chlorine atom: chlorine is in Group 17, so it has 7 outer shell electrons and needs 1 more to complete a stable octet.
2. Each chlorine atom donates one electron to the shared pair, forming one single covalent bond.
3. The completed dot-and-cross diagram shows: 1 shared pair of electrons (one dot from the first Cl, one cross from the second Cl), 3 lone pairs of electrons on each Cl atom.

> **Exam tip:** Always use different symbols (dots vs crosses) for electrons from different atoms in dot-and-cross questions to earn full marks

## Multiple Covalent Bonds

Covalent bonds are classified by their bond order, which is the number of shared electron pairs between two atoms. A clear trend exists between bond order, bond length and bond energy: as bond order increases, bond length decreases and bond energy increases, meaning higher order bonds are shorter and stronger.

**Bond order** — The number of shared electron pairs between two bonded atoms, directly related to bond strength and length

**Worked example:** Compare the bond length and bond energy of C-C, C=C and C≡C bonds

1. Identify the bond order for each bond: C-C (single bond) = 1, C=C (double bond) = 2, C≡C (triple bond) = 3
2. Apply the trend: increasing bond order → decreasing bond length, increasing bond energy
3. Final result: Bond length order: C-C > C=C > C≡C; Bond energy order: C-C < C=C < C≡C

## Coordinate (Dative Covalent) Bonding

**Coordinate (dative covalent) bond** — A covalent bond where both electrons in the shared pair are donated from a single atom (the donor, which has a lone pair of electrons), to an acceptor atom with an empty orbital.

*Example:* The N-H bond in NH₄⁺ formed between NH₃ and H⁺

Once formed, a coordinate bond is identical in strength, length and other properties to any other covalent bond. In structural diagrams, it is represented by an arrow pointing from the donor atom to the acceptor atom.

**Worked example:** Explain the formation of a coordinate bond in the ammonium ion NH₄⁺

1. Ammonia (NH₃) has a central nitrogen atom that forms three regular N-H covalent bonds, leaving one lone pair of electrons on nitrogen.
2. A hydrogen ion H⁺ has an empty 1s outer orbital, so it can accept the lone pair from nitrogen.
3. The lone pair from nitrogen is shared between N and H⁺, forming the coordinate bond. The overall ion has a +1 charge.

> **Exam tip:** Common exam examples to memorise: NH₄⁺, H₃O⁺, Al₂Cl₆, and coordinate bonds in transition metal complexes

## Representing Covalent and Coordinate Bonds

Exam questions regularly require you to draw dot-and-cross diagrams, structural formulas and displayed formulas for species containing both covalent and coordinate bonds. Following standard conventions will help you earn full marks.

**Worked example:** Show how to represent the coordinate bond in the Al₂Cl₆ dimer

1. Each Al atom in AlCl₃ has only 6 outer electrons after forming 3 regular covalent bonds to Cl atoms, so it can accept a lone pair.
2. A Cl atom on one AlCl₃ unit donates a lone pair to the Al atom on the other AlCl₃ unit, and vice versa, forming two coordinate bonds.
3. In the structural diagram, draw an arrow pointing from the Cl donor atom to the Al acceptor atom for each coordinate bond.

**Check your understanding**

Test your understanding:

1. Which of the following species contains at least one coordinate bond?

   - A) H₂O
   - B) NH₄⁺
   - C) CH₄
   - D) Cl₂

   *Why:* Correct. NH₄⁺ forms when NH₃ donates a lone pair to H⁺. All other options only contain regular covalent bonds.

## Common pitfalls

- **Wrong:** Claiming coordinate bonds are weaker than regular covalent bonds
  - Why it fails: This is a common misconception based on different formation, not different properties
  - Correct: State that coordinate bonds are identical in strength and properties to equivalent regular covalent bonds
- **Wrong:** Forgetting to add the overall charge to ions in dot-and-cross diagrams
  - Why it fails: Examiners penalise missing charges even if the electron arrangement is completely correct
  - Correct: Always add the overall charge outside square brackets for any ionic species
- **Wrong:** Drawing the coordinate bond arrow pointing from acceptor to donor
  - Why it fails: The arrow shows the direction of electron pair donation, so direction matters for marks
  - Correct: Point the arrow from the donor atom (with the lone pair) to the acceptor atom (with the empty orbital)
- **Wrong:** Using the same symbol for all electrons in dot-and-cross diagrams
  - Why it fails: Examiners require different symbols to show you understand which electron came from which atom
  - Correct: Use dots for electrons from one atom and crosses for electrons from the other atom

## Cheatsheet

| Bond Type | Electron Source | Bond Order | Common Examples |
| --- | --- | --- | --- |
| Single covalent | 1 from each atom | 1 | Cl-Cl, C-C (ethane) |
| Double covalent | 2 from each atom | 2 | O=O, C=C (ethene) |
| Triple covalent | 3 from each atom | 3 | N≡N, C≡C (ethyne) |
| Coordinate covalent | Both from donor | 1 | NH₄⁺, Al₂Cl₆, H₃O⁺ |

## What's next

Understanding covalent and coordinate bonding is the foundation for all advanced bonding topics in CIE A-Level Chemistry. Your next step is to learn about shapes of covalent molecules, which relies directly on your ability to count bonding pairs and lone pairs around central atoms, a skill built in this module. You will also encounter coordinate bonding again later in the syllabus when studying transition metals, where it forms the basis of ligand bonding in complex ions. Mastering the basics covered here will make these more advanced topics much easier to understand and help you earn full marks in exam questions.

- [Metallic bonding](https://www.owlsprep.com/study/cie-9701-u3-metallic-bonding/)
- [Shapes of molecules and ions](https://www.owlsprep.com/study/cie-9701-u3-shapes-of-molecules-and-ions/)
- [Bonding and physical properties](https://www.owlsprep.com/study/cie-9701-u3-bonding-and-physical-properties/)

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