Molecular Geometry- Practice Questions and Answer Key

Mastering Molecular Geometry: Practice Questions That Actually Work

Molecular geometry isn't a subject you can memorize your way through. You either understand how electron pairs repel each other, or you don't. This guide cuts through the confusion with real practice problems and straight answers.

What You Need to Know First

Before touching these questions, make sure you have VSEPR theory locked down. Valence Shell Electron Pair Repulsion is the engine driving molecular shapes. Electron pairs—bonding and nonbonding—arrange themselves to maximize distance. That's it. Everything else flows from that one principle.

The Core Concepts

• Central atom — the atom bonded to multiple others, usually the least electronegative element
• Electron domains — regions of electron density around the central atom (includes bonds and lone pairs)
• Bonding pairs — electrons shared between atoms
• Lone pairs — electrons belonging to only one atom
• Molecular geometry — the actual shape based on atom positions
• Electron geometry — the arrangement including lone pairs

Here's where students consistently trip up: lone pairs occupy more space than bonding pairs. A molecule with 4 electron domains might be tetrahedral in electron geometry but trigonal pyramidal in molecular geometry. Know the difference.

Common Molecular Shapes at a Glance

Practice Questions

Question 1

Carbon dioxide (CO₂) has two carbon-oxygen double bonds. What is its molecular geometry? What are the bond angles?

Question 2

Sulfur tetrafluoride (SF₄) has one sulfur atom bonded to four fluorine atoms with one lone pair. Draw the shape and name the molecular geometry.

Question 3

Compare the bond angles in CH₄ (methane), NH₃ (ammonia), and H₂O (water). Which has the largest? Which has the smallest? Explain why.

Question 4

Phosphorus pentachloride (PCl₅) exists as a trigonal bipyramidal molecule. Explain why it adopts this geometry rather than a square planar arrangement.

Question 5

Xenon tetrafluoride (XeF₄) has four bonding pairs and two lone pairs. What shape does it exhibit? Identify the bond angles.

Question 6

A molecule has 3 electron domains with 1 lone pair. Another molecule has 4 electron domains with 2 lone pairs. Both are described as "bent." Are they the same shape? What's different and why?

Question 7

Determine the molecular geometry for:
a) BF₃
b) SO₂
c) XeF₂
d) BrF₅

Question 8

Why does a molecule with 5 electron domains and 5 bonding pairs (like PCl₅) have two distinct bond angles, while a molecule with 6 electron domains and 6 bonding pairs (like SF₆) has only one type of bond angle?

Question 9

Draw the electron geometry and molecular geometry for a molecule with 6 electron domains, 4 bonding domains, and 2 lone pairs.

Question 10

Which molecule has a bond angle closest to 90°: SF₄, XeF₄, or IF₅? Explain your reasoning.

Answer Key

Answer 1

Linear geometry. 180° bond angle.

CO₂ is a linear molecule because carbon has two electron domains (two double bonds). The domains arrange on opposite sides to maximize distance. There are no lone pairs on the central carbon, so molecular geometry matches electron geometry.

Answer 2

Seesaw geometry. The lone pair occupies an equatorial position to minimize repulsions.

In trigonal bipyramidal electron geometry, lone pairs always go to equatorial positions. This gives SF₄ a distorted seesaw shape. Bond angles are approximately 120° in the equatorial plane and 90° between axial and equatorial bonds.

Answer 3

CH₄ has the largest (109.5°). H₂O has the smallest (~104.5°). NH₃ falls in between (~107°).

Lone pairs repel more strongly than bonding pairs. More lone pairs means more repulsion pushing bonds closer together, reducing angles. Water has two lone pairs, ammonia has one, methane has none.

Answer 4

Trigonal bipyramidal is more stable because it minimizes electron pair repulsions more effectively than square planar.

In trigonal bipyramidal geometry, three positions are in a plane (equatorial) at 120° to each other, and two are axial at 90° to the plane. This creates five distinct positions with maximum separation. Square planar would force bonds closer together, increasing repulsions.

Answer 5

Square planar geometry. 90° bond angles.

XeF₄ has six electron domains: four bonding pairs and two lone pairs. The lone pairs occupy opposite positions (axial in trigonal bipyramidal terms), forcing the four fluorine atoms into a square planar arrangement.

Answer 6

They're both bent, but the angles differ. The first has ~117° angles (3 domains, 1 lone pair). The second has ~104.5° angles (4 domains, 2 lone pairs).

The difference comes from electron geometry. One has trigonal planar electron geometry; the other has tetrahedral electron geometry. More lone pairs compress the bonds more. Same name, different physics.

Answer 7

a) BF₃ — Trigonal planar (3 bonding pairs, no lone pairs)
b) SO₂ — Bent (3 electron domains, 1 lone pair)
c) XeF₂ — Linear (5 electron domains, 3 lone pairs)
d) BrF₅ — Square pyramidal (6 electron domains, 1 lone pair)

Answer 8

PCl₅ has two types of positions: axial and equatorial. Axial bonds are at 90° to equatorial bonds, but equatorial bonds are at 120° to each other. That's two distinct angles.

SF₆ has all positions equivalent in octahedral geometry. Every bond is 90° to every adjacent bond. No variation.

Answer 9

Electron geometry: Octahedral. Molecular geometry: Square planar.

With 6 electron domains and 2 lone pairs, the lone pairs take opposite positions. The remaining 4 bonding pairs form a flat square around the central atom.

Answer 10

XeF₄ has bond angles closest to 90°.

XeF₄ is square planar with all bonds at exactly 90°. SF₄ has seesaw geometry with angles of 90° and 120°. IF₅ is square pyramidal with angles slightly above 90° due to lone pair repulsion from the single lone pair pushing bonds down.

How to Solve Any Molecular Geometry Problem

Follow this sequence every time. No exceptions.

Step 1: Find the Central Atom

Usually the least electronegative element (except hydrogen). Carbon, nitrogen, sulfur, and phosphorus are common central atoms.

Step 2: Count Total Valence Electrons

Add up valence electrons from all atoms. Then account for charges. Add electrons for negative charge, subtract for positive.

Step 3: Count Electron Domains

Each bond (single, double, triple) counts as one domain. Each lone pair counts as one domain. Each unpaired electron counts as half a domain (rare in stable molecules).

Step 4: Place Lone Pairs

In trigonal bipyramidal geometry, lone pairs go equatorial. In octahedral, lone pairs go opposite each other first if there are two.

Step 5: Name the Shape

Match your arrangement to the chart. Remember: molecular geometry excludes lone pairs, electron geometry includes them.

Common Mistakes to Avoid

• Forgetting double bonds count as one domain — CO₂ has 2 domains, not 4
• Confusing electron geometry with molecular geometry — always specify which one the question asks for
• Placing lone pairs in axial positions in trigonal bipyramidal — equatorial is always better
• Memorizing without understanding — you need to know WHY shapes form, not just what they are
• Ignoring lone pair repulsion strength — lone pairs compress bond angles more than bonding pairs do

Quick Reference: Lone Pair Effects

The more lone pairs you add, the more the bonds get pushed together. That's the whole game.