Shapes of molecules and ions

Q: What is the core principle of Valence Shell Electron Pair Repulsion (VSEPR) theory?

Electron pairs in the outer shell of a central atom repel each other and push as far apart as possible to minimise mutual repulsion.

Q: What is the order of repulsive strength between different types of electron pairs?

Lone pair – Lone pair > Lone pair – Bonding pair > Bonding pair – Bonding pair.

Q: What is the shape and bond angle of a molecule with 2 bonding pairs and 0 lone pairs around the central atom?

Linear with a bond angle of \( 180^\circ \) .

Q: What are the shape and bond angle for 3 bonding pairs and 0 lone pairs, versus 2 bonding pairs and 1 lone pair?

3 bp, 0 lp gives Trigonal Planar (\( 120^\circ \)); 2 bp, 1 lp gives Bent / V-shaped (\( \approx 117.5^\circ \) or \( < 120^\circ \)).

Q: How do bond angles and shapes change across the 4 electron pair series (4 bp, 3 bp + 1 lp, 2 bp + 2 lp)?

4 bp, 0 lp is Tetrahedral (\( 109.5^\circ \)); 3 bp, 1 lp is Pyramidal (\( 107^\circ \)); 2 bp, 2 lp is Bent (\( 104.5^\circ \)).

Q: What are the shapes and bond angles of molecules with 5 and 6 bonding pairs (0 lone pairs)?

5 pairs: Trigonal Bipyramidal (\( 120^\circ \) and \( 90^\circ \)); 6 pairs: Octahedral (\( 90^\circ \)).

Q: How do you account for charge when determining the total electron pairs in polyatomic ions?

Subtract electrons for positive charges (\( + \)) and add electrons for negative charges (\( - \)).

Q: What three key points must you include to score full marks when explaining molecular shapes in CCEA exams?

1. State number of bonding and lone pairs; 2. State the repulsion principle (and relative lp vs. bp strengths); 3. State the correct shape name and exact bond angle.

Chemistry 1110 · CCEA A Level · 9 min read · Updated: 31 Aug 2026

Shapes of Molecules and Ions

Welcome to one of the most visual and interesting topics in AS Chemistry! Have you ever wondered why water (\( \text{H}_2\text{O} \)) is shaped like a bent "V", while methane (\( \text{CH}_4 \)) forms a perfect 3D pyramid with a triangular base? In this chapter, we will learn how to predict and draw the exact three-dimensional shapes of molecules and polyatomic ions, and explain their bond angles.

Understanding molecular shapes is essential because the 3D shape of a molecule determines how it interacts with other molecules — from how enzymes in your body recognize medicines to the physical properties of everyday materials.

Don't worry if 3D geometry sounds a bit intimidating at first! By following a simple, foolproof set of rules, you will be able to work out the shape and bond angle of any molecule or ion on your CCEA exam.

1. The Core Idea: Valence Shell Electron Pair Repulsion (VSEPR) Theory

The foundation of molecular shapes is called Valence Shell Electron Pair Repulsion (VSEPR) Theory. While the name sounds long, the main concept is very simple:

1. Electrons are negatively charged. Because like charges repel, pairs of electrons in the outer (valence) shell of a central atom push as far away from each other as possible to minimise mutual repulsion.

2. The final shape depends on the number and type of electron pairs. The total number of electron pairs around the central atom determines the overall geometric layout.

A Helpful Analogy: Tying Balloons Together

Imagine tying two inflated balloons together by their strings. They naturally point in opposite directions along a straight line (\( 180^\circ \)). If you tie three balloons together, they push apart into a flat triangle (\( 120^\circ \)). Tie four together, and they pop into a 3D tripod shape (\( 109.5^\circ \)). Electron pairs behave in the exact same way!

Bonding Pairs vs. Lone Pairs

There are two types of electron pairs around a central atom:

• Bonding Pairs (bp): Electrons shared between the central atom and bonded atoms. They are held tightly between two nuclei.
• Lone Pairs (lp): Unshared outer-shell electron pairs held only by the central atom's nucleus. Because they are closer to the central nucleus, their electron cloud is fatter and more concentrated, exerting a stronger repulsive force than bonding pairs.

The Order of Repulsion Strength

Different types of electron pairs repel with different strengths. You must memorize this order for your CCEA exams:

Lone pair – Lone pair repulsion > Lone pair – Bonding pair repulsion > Bonding pair – Bonding pair repulsion

Rule of Thumb for Bond Angles: Each lone pair pushes bonding pairs closer together, reducing the standard bond angle by approximately \( 2.5^\circ \).

Key Takeaway: Electron pairs repel each other to achieve maximum separation and minimum repulsion. Lone pairs repel more strongly than bonding pairs, squeezing bond angles smaller.

2. A Step-by-Step Method to Determine Shape and Bond Angles

Whenever you are asked to deduce the shape of a molecule or ion, follow these 5 steps:

Step 1: Identify the central atom and determine the number of electrons in its outer shell (its Group number in the Periodic Table).
Step 2: Count how many atoms are bonded to the central atom. Each singly-bonded atom contributes 1 electron.
Step 3: Adjust for any electrical charge:
• If the ion is positive (\( + \)), subtract electrons.
• If the ion is negative (\( - \)), add electrons.
Step 4: Add these electrons together and divide by 2 to find the total number of electron pairs.
Step 5: Compare the number of bonding pairs (equal to the number of bonded atoms) with the total pairs to find the number of lone pairs (\( \text{lone pairs} = \text{total pairs} - \text{bonding pairs} \)).

3. Summary of Shapes (From 2 to 6 Electron Pairs)

A. Two Electron Pairs

• Total Pairs: 2 (2 bonding pairs, 0 lone pairs)
• Name of Shape: Linear
• Bond Angle: \( 180^\circ \)
• Example: Beryllium chloride (\( \text{BeCl}_2 \)), Carbon dioxide (\( \text{CO}_2 \))
• Explanation: The two bonding pairs push as far apart as possible to opposite sides of the central atom.

B. Three Electron Pairs

Case 1: 3 Bonding Pairs, 0 Lone Pairs
• Name of Shape: Trigonal Planar
• Bond Angle: \( 120^\circ \)
• Example: Boron trifluoride (\( \text{BF}_3 \)), Aluminium chloride (\( \text{AlCl}_3 \))
• Explanation: The three pairs spread out equally in a flat, 2D plane.

Case 2: 2 Bonding Pairs, 1 Lone Pair
• Name of Shape: Bent / V-shaped / Non-linear
• Bond Angle: \( \approx 117.5^\circ \) (or \( < 120^\circ \))
• Example: Sulfur dioxide (\( \text{SO}_2 \))
• Explanation: The lone pair exerts greater repulsion on the bonding pairs, pushing the angle below \( 120^\circ \).

C. Four Electron Pairs

Case 1: 4 Bonding Pairs, 0 Lone Pairs
• Name of Shape: Tetrahedral
• Bond Angle: \( 109.5^\circ \)
• Example: Methane (\( \text{CH}_4 \)), Ammonium ion (\( \text{NH}_4^+ \))
• Explanation: The four bonding pairs repel equally into a 3D tetrahedral shape.

Case 2: 3 Bonding Pairs, 1 Lone Pair
• Name of Shape: Pyramidal / Trigonal Pyramidal
• Bond Angle: \( 107^\circ \) (calculated as \( 109.5^\circ - 2.5^\circ \))
• Example: Ammonia (\( \text{NH}_3 \)), Hydronium ion (\( \text{H}_3\text{O}^+ \))
• Explanation: The lone pair repels the three bonding pairs more strongly than they repel each other, pushing them closer together.

Case 3: 2 Bonding Pairs, 2 Lone Pairs
• Name of Shape: Bent / V-shaped / Non-linear
• Bond Angle: \( 104.5^\circ \) (calculated as \( 109.5^\circ - 2 \times 2.5^\circ \))
• Example: Water (\( \text{H}_2\text{O} \)), Hydrogen sulfide (\( \text{H}_2\text{S} \))
• Explanation: Two lone pairs exert powerful repulsion, compressing the \( \text{H-O-H} \) bond angle down to \( 104.5^\circ \).

D. Five Electron Pairs

• Total Pairs: 5 (5 bonding pairs, 0 lone pairs)
• Name of Shape: Trigonal Bipyramidal
• Bond Angles: \( 120^\circ \) (equatorial) and \( 90^\circ \) (axial)
• Example: Phosphorus pentachloride (\( \text{PCl}_5 \))
• Explanation: Three chlorine atoms form a flat equatorial triangle with \( 120^\circ \) angles, while two axial chlorine atoms sit directly above and below at \( 90^\circ \).

E. Six Electron Pairs

• Total Pairs: 6 (6 bonding pairs, 0 lone pairs)
• Name of Shape: Octahedral
• Bond Angle: \( 90^\circ \)
• Example: Sulfur hexafluoride (\( \text{SF}_6 \))
• Explanation: Six bonding pairs point to the six corners of a regular octahedron, with all adjacent angles equal to \( 90^\circ \).

Did you know? Even though the shape is called octahedral, it has only 6 bonded atoms! The prefix "octa-" refers to the 8 faces of the 3D geometric solid formed by joining the outer points.

Key Takeaway: Four pairs give a tetrahedral framework (\( 109.5^\circ \)), which drops to \( 107^\circ \) with 1 lone pair (pyramidal), and \( 104.5^\circ \) with 2 lone pairs (bent).

4. Working with Polyatomic Ions

Applying VSEPR to ions works identically, as long as you adjust for the ionic charge when counting electrons.

Example 1: The Ammonium Ion (\( \text{NH}_4^+ \))

• Central atom: Nitrogen (Group 5) = \( 5 \text{ electrons} \)
• 4 bonded hydrogen atoms contribute: \( 4 \times 1 = 4 \text{ electrons} \)
• Positive charge of \( +1 \): Subtract 1 electron
• Total electrons = \( 5 + 4 - 1 = 8 \text{ electrons} \)
• Total electron pairs = \( \frac{8}{2} = 4 \text{ pairs} \)
• Bonding pairs = 4, Lone pairs = 0
• Shape: Tetrahedral, Bond Angle: \( 109.5^\circ \)

Example 2: The Hydronium / Oxonium Ion (\( \text{H}_3\text{O}^+ \))

• Central atom: Oxygen (Group 6) = \( 6 \text{ electrons} \)
• 3 bonded hydrogen atoms contribute: \( 3 \times 1 = 3 \text{ electrons} \)
• Positive charge of \( +1 \): Subtract 1 electron
• Total electrons = \( 6 + 3 - 1 = 8 \text{ electrons} \)
• Total electron pairs = \( \frac{8}{2} = 4 \text{ pairs} \)
• Bonding pairs = 3, Lone pairs = 1
• Shape: Pyramidal, Bond Angle: \( 107^\circ \)

Key Takeaway: Treat polyatomic ions just like molecules, but remember to subtract electrons for positive charges and add electrons for negative charges.

5. How to Score Full Marks in CCEA Exam Questions

CCEA exam questions often ask you to "Predict the shape of the molecule and explain your reasoning." To guarantee full marks, always structure your answer using this standard 3-part template:

1. State the number of electron pairs: State clearly how many bonding pairs and how many lone pairs surround the central atom.
2. State the repulsion principle: State that "Electron pairs repel each other to be as far apart as possible" and specify that "Lone pairs repel more strongly than bonding pairs" (if lone pairs are present).
3. State the shape and bond angle: Name the correct shape and quote the exact bond angle in degrees.

Model Exam Answer for Ammonia (\( \text{NH}_3 \)):

"There are 3 bonding pairs and 1 lone pair of electrons around the central nitrogen atom. Electron pairs repel to get as far apart as possible, and lone pair–bonding pair repulsion is greater than bonding pair–bonding pair repulsion. Therefore, the shape is pyramidal with a bond angle of \( 107^\circ \)."

Drawing 3D Molecules (Wedge and Dash Conventions)

When asked to sketch 3D shapes (such as tetrahedral \( \text{CH}_4 \)):
• A solid straight line (\( - \)) represents a bond in the plane of the paper.
• A solid wedge represents a bond coming forward out of the page towards you.
• A dashed line / hatched wedge represents a bond going backward into the page away from you.

Common Mistakes to Avoid

• Forgetting to adjust for charge: Remember: \( + \) means subtract an electron, \( - \) means add an electron!
• Confusing the shape name of water: Water has 4 electron pairs, so its electron arrangement is tetrahedral, but its actual molecular shape is bent / non-linear.
• Missing units on angles: Always include the degree symbol (e.g., write \( 104.5^\circ \), not just 104.5).
• Leaving out lone pairs: Always check whether the central atom has unbonded valence electrons left over.

Quick Review Summary Table

• 2 bp, 0 lp: Linear | \( 180^\circ \) | e.g. \( \text{BeCl}_2 \)
• 3 bp, 0 lp: Trigonal Planar | \( 120^\circ \) | e.g. \( \text{BF}_3 \)
• 4 bp, 0 lp: Tetrahedral | \( 109.5^\circ \) | e.g. \( \text{CH}_4 \), \( \text{NH}_4^+ \)
• 3 bp, 1 lp: Pyramidal | \( 107^\circ \) | e.g. \( \text{NH}_3 \), \( \text{H}_3\text{O}^+ \)
• 2 bp, 2 lp: Bent / Non-linear | \( 104.5^\circ \) | e.g. \( \text{H}_2\text{O} \)
• 5 bp, 0 lp: Trigonal Bipyramidal | \( 120^\circ \) & \( 90^\circ \) | e.g. \( \text{PCl}_5 \)
• 6 bp, 0 lp: Octahedral | \( 90^\circ \) | e.g. \( \text{SF}_6 \)

Quick check

Can you answer these now?

Open each question to check the key ideas from this chapter.

What is the core principle of Valence Shell Electron Pair Repulsion (VSEPR) theory?

Electron pairs in the outer shell of a central atom repel each other and push as far apart as possible to minimise mutual repulsion.

What is the order of repulsive strength between different types of electron pairs?

Lone pair – Lone pair > Lone pair – Bonding pair > Bonding pair – Bonding pair.

What is the shape and bond angle of a molecule with 2 bonding pairs and 0 lone pairs around the central atom?

Linear with a bond angle of \( 180^\circ \).

What are the shape and bond angle for 3 bonding pairs and 0 lone pairs, versus 2 bonding pairs and 1 lone pair?

3 bp, 0 lp gives Trigonal Planar (\( 120^\circ \)); 2 bp, 1 lp gives Bent / V-shaped (\( \approx 117.5^\circ \) or \( < 120^\circ \)).

How do bond angles and shapes change across the 4 electron pair series (4 bp, 3 bp + 1 lp, 2 bp + 2 lp)?

4 bp, 0 lp is Tetrahedral (\( 109.5^\circ \)); 3 bp, 1 lp is Pyramidal (\( 107^\circ \)); 2 bp, 2 lp is Bent (\( 104.5^\circ \)).

What are the shapes and bond angles of molecules with 5 and 6 bonding pairs (0 lone pairs)?

5 pairs: Trigonal Bipyramidal (\( 120^\circ \) and \( 90^\circ \)); 6 pairs: Octahedral (\( 90^\circ \)).

How do you account for charge when determining the total electron pairs in polyatomic ions?

Subtract electrons for positive charges (\( + \)) and add electrons for negative charges (\( - \)).

What three key points must you include to score full marks when explaining molecular shapes in CCEA exams?

1. State number of bonding and lone pairs; 2. State the repulsion principle (and relative lp vs. bp strengths); 3. State the correct shape name and exact bond angle.

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