# Covalent bonding

> IB Chemistry SL · IB SL Chemistry 2025 Syllabus
> Source: https://www.owlsprep.com/study/ib-chemistry-sl-u3-covalent-bonding/

This module covers the fundamentals of covalent bonding, where nonmetal atoms share valence electrons to form stable compounds. You will learn bond properties, how to draw Lewis structures, and recognize common exceptions to the octet rule.

**Prerequisites:** [Valence electrons and electron configuration](https://www.owlsprep.com/study/ib-chemistry-sl-u2-electron-configuration/)

## Learning objectives

- Describe the nature of covalent bond formation between nonmetal atoms
- Relate bond order to bond length and bond enthalpy
- Draw correct Lewis structures for simple molecules and polyatomic ions
- Identify common exceptions to the octet rule

## Nature of a Covalent Bond

A covalent bond forms when two nonmetal atoms share valence electrons to achieve stable full outer electron shells. The bond is held together by electrostatic attraction between the positively charged nuclei of the bonded atoms and the negatively charged shared electron pair.

**Covalent Bond** — A chemical bond formed by the electrostatic attraction between shared pairs of electrons and the nuclei of two bonded nonmetal atoms

*Notation:* Represented as a single line (-) for one shared pair

*Example:* Single covalent bond between two H atoms in $H_2$

**Worked example:** Explain why two hydrogen atoms form a stable covalent bond rather than remaining separate atoms.

1. Each hydrogen atom has 1 valence electron and needs 1 more electron to fill its 1s orbital (the first energy level holds a maximum of 2 electrons).
2. When two H atoms approach, each electron is attracted to the positively charged nucleus of the opposite atom.
3. The atoms share both electrons, giving each a full valence shell. The net electrostatic attraction between nuclei and shared electrons forms a stable $H_2$ molecule.

*Calculator:* forbidden

## Bond Order, Length and Enthalpy

Bond order describes the number of shared electron pairs between two atoms, and directly determines two key bond properties: bond length and bond enthalpy (strength).

**Bond Order** — The number of covalent bonding pairs shared between two atoms: single bonds = 1, double = 2, triple = 3

- As bond order increases, more shared electrons increase electrostatic attraction between nuclei
- Higher attraction pulls nuclei closer, so **bond length decreases as bond order increases**
- More attraction requires more energy to break the bond, so **bond enthalpy increases as bond order increases**

**Worked example:** Arrange the following carbon-carbon bonds in order of increasing bond enthalpy, and explain your answer: C-C (ethane), C=C (ethene), C≡C (ethyne).

1. First identify the bond order for each bond: C-C = 1, C=C = 2, C≡C = 3.
2. Bond enthalpy increases with increasing bond order, because higher bond order means more shared electron pairs and stronger electrostatic attraction between carbon nuclei and shared electrons.
3. Order of increasing bond enthalpy: $\text{C-C} < \text{C=C} < \text{C≡C}$

> **tip**
>
> The inverse relationship between bond length and bond enthalpy is one of the most frequently tested concepts in Paper 1 for this topic.

*Calculator:* forbidden

## Drawing Lewis Structures

Lewis structures show all valence electrons in a covalent molecule or polyatomic ion, distinguishing between bonding (shared) electrons and non-bonding (lone pair) electrons.

**Exam command terms**

- **Draw** — Show all atoms, bonds, and lone pairs of electrons *(Typically 2 marks total: 1 for connectivity, 1 for all lone pairs)*

- **State** — Name the number of bonding/lone pairs, no diagram required

1. Count total valence electrons: add 1 electron per negative charge, subtract 1 per positive charge for ions
2. Arrange atoms: the least electronegative atom (always except H) is the central atom
3. Draw one single bond between each connected atom pair
4. Distribute remaining electrons as lone pairs starting from terminal atoms, then the central atom
5. Convert terminal lone pairs to multiple bonds if the central atom needs an octet

**Worked example:** Draw the Lewis structure of the carbonate ion, $\text{CO}_3^{2-}$.

1. Calculate total valence electrons: C has 4, each O has 6, add 2 for the 2- charge. Total: $4 + (3 \times 6) + 2 = 24$ electrons.
2. C is the least electronegative, so it is the central atom bonded to 3 O atoms. Three single bonds use 6 electrons, leaving 18 electrons remaining.
3. Add 3 lone pairs (6 electrons) to each terminal O atom, which uses all 18 remaining electrons.
4. Central C only has 6 electrons. Convert one O lone pair to a bonding pair, forming a C=O double bond. All atoms now have full octets.
5. Enclose the structure in square brackets and add the 2- charge at the top right.

*Calculator:* forbidden

## Exceptions to the Octet Rule

The octet rule states that main group atoms tend to share electrons until they have 8 valence electrons, but three common classes of exceptions exist in covalent bonding.

**Octet Rule** — A general rule that main-group atoms achieve a noble gas electron configuration with 8 valence electrons in covalent compounds

- **Incomplete octet**: Central atom has < 8 electrons (common for Be, B, Al)
- **Expanded octet**: Central atom has > 8 electrons, only possible for period 3+ elements with empty d-orbitals
- **Odd electron species**: Free radicals with an unpaired electron cannot have all atoms with full octets

**Worked example:** Explain why phosphorus pentachloride ($PCl_5$) is an exception to the octet rule.

1. Calculate total valence electrons: P has 5, each Cl has 7: total = $5 + (5 \times 7) = 40$ electrons.
2. P is the central atom bonded to 5 Cl atoms, so there are 10 bonding electrons around P.
3. Each Cl has 3 lone pairs, so total electrons used: $10 + (5 \times 6) = 40$, matching the total.
4. Phosphorus is in period 3, so it has empty 3d orbitals that can accommodate more than 8 valence electrons. This is an expanded octet exception.

*Calculator:* forbidden

## Common pitfalls

- **Wrong:** Forgetting to add/subtract electrons for the charge when counting valence electrons for an ion
  - Why it fails: This leads to an incorrect total number of electrons and wrong Lewis structure
  - Correct: Always add 1 electron per unit negative charge, subtract 1 per unit positive charge before drawing
- **Wrong:** Mixing up the relationship between bond length and bond strength
  - Why it fails: Students often assume longer bonds are stronger, which is incorrect
  - Correct: Shorter bond length = stronger bond (higher bond enthalpy), longer bond = weaker bond
- **Wrong:** Putting the most electronegative atom as the central atom
  - Why it fails: This leads to incorrect connectivity and lost marks
  - Correct: The least electronegative atom (excluding hydrogen, which is always terminal) is the central atom
- **Wrong:** Allowing period 2 elements to have expanded octets
  - Why it fails: Period 2 elements only have s and p orbitals that can hold a maximum of 8 electrons
  - Correct: Only period 3 and below elements can have expanded octets, as they have empty d-orbitals
- **Wrong:** Counting shared electrons only once when checking for an octet
  - Why it fails: Shared electrons count towards the octet of both bonded atoms
  - Correct: Count both electrons in a bonding pair towards the octet of each bonded atom

## Cheatsheet

| Bond Type | Bond Order | Approx Length (pm) | Approx Enthalpy (kJ/mol) |
| --- | --- | --- | --- |
| C-C | 1 | 154 | 347 |
| C=C | 2 | 134 | 614 |
| C≡C | 3 | 120 | 839 |
| O-O | 1 | 148 | 146 |
| O=O | 2 | 121 | 498 |
| N-N | 1 | 147 | 163 |
| N≡N | 3 | 110 | 945 |
| C-O | 1 | 143 | 358 |
| C=O | 2 | 120 | 745 |

## What's next

Mastering covalent bonding is the foundation for all subsequent topics in chemical bonding and beyond. Lewis structures and the arrangement of bonding vs non-bonding electrons directly determine the shape of covalent molecules, which in turn determines molecular polarity and intermolecular forces. Bond enthalpy values from this topic are also used to calculate enthalpy changes for chemical reactions in energetic topics. Build on this knowledge with the following topics to prepare for exam questions.

- [Shapes of molecules and ions](https://www.owlsprep.com/study/ib-chemistry-sl-u3-shapes-of-molecules-and-ions/)
- [Intermolecular forces](https://www.owlsprep.com/study/ib-chemistry-sl-u3-intermolecular-forces/)
- [Metallic bonding](https://www.owlsprep.com/study/ib-chemistry-sl-u3-metallic-bonding/)

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