Covalent bonding

IB Chemistry SL· Structure 3: Chemical Bonding· 6 min read

1. Nature of a Covalent Bond★★☆☆☆⏱ 10 min

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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.

📘 Definition

Covalent Bond

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

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

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.

2. Bond Order, Length and Enthalpy★★☆☆☆⏱ 15 min

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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).

📘 Definition

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: $C-C < C=C < C\equivC$

3. Drawing Lewis Structures★★★☆☆⏱ 20 min

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

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

📐 Worked Example

Draw the Lewis structure of the carbonate ion, $CO_{3}^{2 -}$ .

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.

4. Exceptions to the Octet Rule★★★☆☆⏱ 15 min

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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.

📘 Definition

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 ( $P C l_{5}$ ) 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.

5. Common Pitfalls

Wrong move:

Forgetting to add/subtract electrons for the charge when counting valence electrons for an ion

Why:

This leads to an incorrect total number of electrons and wrong Lewis structure

Correct move:

Always add 1 electron per unit negative charge, subtract 1 per unit positive charge before drawing

Wrong move:

Mixing up the relationship between bond length and bond strength

Why:

Students often assume longer bonds are stronger, which is incorrect

Correct move:

Shorter bond length = stronger bond (higher bond enthalpy), longer bond = weaker bond

Wrong move:

Putting the most electronegative atom as the central atom

Why:

This leads to incorrect connectivity and lost marks

Correct move:

The least electronegative atom (excluding hydrogen, which is always terminal) is the central atom

Wrong move:

Allowing period 2 elements to have expanded octets

Why:

Period 2 elements only have s and p orbitals that can hold a maximum of 8 electrons

Correct move:

Only period 3 and below elements can have expanded octets, as they have empty d-orbitals

Wrong move:

Counting shared electrons only once when checking for an octet

Why:

Shared electrons count towards the octet of both bonded atoms

Correct move:

Count both electrons in a bonding pair towards the octet of each bonded atom

6. Quick Reference 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

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2023 · 1
Bond order/enthalpy relationship MCQ
2022 · 2
Draw Lewis structure for ion
2021 · 1
Identify octet rule exception

Going deeper

syllabusIB Chemistry SL 2025 SyllabusOfficial syllabus reference for Structure 3

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.