Why Do Molecules Have Different Shapes?

Atoms in molecules are arranged in specific three-dimensional shapes.

A molecule is not simply a flat collection of atoms. The atoms occupy positions in three-dimensional space, and their arrangement determines the molecule's molecular geometry.

Molecular shape can affect:

• Polarity.
• Solubility.
• Boiling point.
• Intermolecular forces.
• Chemical reactions.
• Biological activity.

This means that two substances containing similar atoms can behave very differently because their atoms are arranged differently.

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Electron Pairs Determine Shape

The shape of many simple molecules can be predicted using the idea that electron pairs repel each other.

Electrons have negative charge.

Because like charges repel, regions containing electron pairs around a central atom tend to move as far apart as possible.

This produces a three-dimensional arrangement that minimizes repulsion.

This idea is called:

Valence Shell Electron Pair Repulsion theory

or:

VSEPR theory

The basic rule is:

Electron pairs around a central atom arrange themselves as far apart as possible.

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Bonding Pairs and Lone Pairs

There are two important types of electron pairs around a central atom.

Bonding pairs

A bonding pair is a pair of electrons shared between two atoms in a covalent bond.

Lone pairs

A lone pair is a pair of outer-shell electrons that is not involved in bonding.

Both bonding pairs and lone pairs repel other electron pairs.

However, lone pairs generally repel more strongly because their electron density is concentrated closer to the central atom.

A useful general relationship is:

lone pair–lone pair repulsion > lone pair–bond pair repulsion > bond pair–bond pair repulsion

This can change bond angles and therefore affect molecular shape.

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Linear Molecules

A molecule with two bonding regions around its central atom and no lone pairs is often linear.

The two bonding regions move as far apart as possible.

The bond angle is:

180°

A common example is carbon dioxide:

CO₂

Its arrangement can be represented as:

O=C=O

The three atoms form a straight line.

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Why Is Carbon Dioxide Linear?

The central carbon atom has two regions of bonding electron density.

These regions repel each other.

The greatest possible separation is:

180°

Therefore, they position themselves on opposite sides of the carbon atom.

The resulting molecular shape is:

linear

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Trigonal Planar Molecules

Three bonding regions around a central atom arrange themselves in a flat triangular pattern.

This shape is called:

trigonal planar

The bond angles are approximately:

120°

A common example is boron trifluoride:

BF₃

The three bonding regions spread equally around the central boron atom.

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Tetrahedral Molecules

Four bonding regions around a central atom arrange themselves in a tetrahedral shape.

A common example is methane:

CH₄

The bond angle is approximately:

109.5°

A flat cross would not separate four electron pairs as effectively.

The tetrahedral arrangement allows the bonding regions to spread out in three dimensions.

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Understanding a Tetrahedron

In methane, the carbon atom is at the center.

The four hydrogen atoms point toward the corners of an imaginary tetrahedron.

This demonstrates why molecular models are useful.

A two-dimensional formula such as:

CH₄

does not show the actual three-dimensional arrangement of the atoms.

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Lone Pairs Change Molecular Shape

Consider a central atom surrounded by four electron regions.

If all four regions are bonding pairs, the molecular shape is:

tetrahedral

But if one or more of those regions is a lone pair, the positions of the atoms change.

The electron-pair arrangement may still be based on a tetrahedron, but the molecular shape is named according to the positions of the atoms, not the invisible lone pairs.

This distinction is very important.

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Trigonal Pyramidal Molecules

Ammonia has the formula:

NH₃

Around the nitrogen atom there are:

• Three bonding pairs.
• One lone pair.

The four electron regions arrange themselves approximately tetrahedrally.

However, only three positions contain atoms.

The resulting molecular shape is:

trigonal pyramidal

The H–N–H bond angle is approximately:

107°

This is slightly smaller than the tetrahedral angle of 109.5° because the lone pair repels the bonding pairs more strongly.

---

Bent Molecules

Water has the formula:

H₂O

Around the oxygen atom there are:

• Two bonding pairs.
• Two lone pairs.

The four electron regions are arranged approximately tetrahedrally.

However, only two positions contain hydrogen atoms.

The molecular shape is therefore:

bent

or:

V-shaped

The H–O–H bond angle is approximately:

104.5°

The two lone pairs produce stronger repulsion, pushing the O–H bonds closer together.

---

Comparing CH₄, NH₃, and H₂O

These three molecules are especially useful for understanding the effect of lone pairs.

Notice the pattern:

More lone-pair repulsion → smaller bond angles between bonded atoms

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Trigonal Bipyramidal Molecules

Five bonding regions around a central atom can form a:

trigonal bipyramidal

shape.

A common example is:

PCl₅

The arrangement contains:

• Three positions around the middle.
• Two positions above and below.

Some bond angles are:

120°

while others are:

90°

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Octahedral Molecules

Six bonding regions around a central atom can produce an:

octahedral

shape.

A common example is:

SF₆

The six bonding regions point in six different directions around the central sulfur atom.

Adjacent bonds have angles of approximately:

90°

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Summary of Common Molecular Shapes

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Double and Triple Bonds

When predicting basic molecular geometry, a:

• Single bond.
• Double bond.
• Triple bond.

is normally treated as one region of electron density.

For example, carbon dioxide contains two double bonds:

O=C=O

There are still only two bonding regions around the carbon atom.

Therefore, CO₂ is linear.

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Worked Example: Predicting the Shape of CH₄

Consider methane:

CH₄

Identify the central atom

Carbon is the central atom.

Count the electron regions

Carbon forms four C–H bonds.

Therefore:

4 bonding regions

There are no lone pairs on the carbon atom.

Arrange the electron regions

Four regions repel each other and spread out as far as possible.

Determine the shape

The shape is:

tetrahedral

Bond angle:

109.5°

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Worked Example: Predicting the Shape of NH₃

Consider ammonia:

NH₃

Nitrogen has:

• Three N–H bonds.
• One lone pair.

Total:

4 electron regions

The electron regions arrange approximately tetrahedrally.

However, molecular shape describes the positions of the atoms.

Therefore:

NH₃ is trigonal pyramidal.

The bond angle is approximately:

107°

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Worked Example: Predicting the Shape of H₂O

Water has:

• Two O–H bonds.
• Two lone pairs.

Total:

4 electron regions

The electron regions have a tetrahedral arrangement.

However, only two regions contain bonded atoms.

Therefore:

H₂O is bent.

Bond angle:

approximately 104.5°

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Electron Geometry Versus Molecular Geometry

These terms are related but not identical.

Electron geometry considers:

• Bonding regions.
• Lone pairs.

Molecular geometry considers:

• The positions of the atoms.

For methane:

Electron geometry = tetrahedral

Molecular geometry = tetrahedral

For ammonia:

Electron geometry = tetrahedral

Molecular geometry = trigonal pyramidal

For water:

Electron geometry = tetrahedral

Molecular geometry = bent

This difference occurs because lone pairs affect the arrangement but are not atoms.

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Interpreting Molecular Models

Molecular models help us visualize three-dimensional structures.

Common representations include:

Ball-and-stick models

Atoms are represented as balls and bonds as sticks.

These clearly show:

• Bonding.
• Bond angles.
• Molecular geometry.

Space-filling models

Atoms are shown using overlapping spheres.

These better represent the approximate space occupied by atoms.

Structural formulas

These show which atoms are connected but may not clearly show three-dimensional shape.

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Wedge-and-Dash Diagrams

Chemists can represent three-dimensional molecules on a flat page using different types of bonds.

A normal line represents a bond approximately in the plane of the page.

A solid wedge represents a bond pointing toward the viewer.

A dashed wedge represents a bond pointing away from the viewer.

These diagrams allow three-dimensional molecular geometry to be represented in two dimensions.

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Molecular Shape and Polarity

Molecular shape can influence whether a molecule is polar.

A polar bond occurs when electrons are shared unevenly between atoms.

However, the presence of polar bonds does not automatically mean the entire molecule is polar.

The shape of the molecule determines whether the effects of individual bond polarities cancel.

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Carbon Dioxide and Polarity

Carbon dioxide contains polar C=O bonds.

However:

CO₂ is linear.

The two C=O bonds point in opposite directions.

Their effects largely cancel.

Therefore:

CO₂ is overall nonpolar.

This shows why molecular shape matters.

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Water and Polarity

Water also contains polar bonds.

However:

H₂O is bent.

The O–H bond polarities do not point directly opposite each other.

They therefore do not cancel.

As a result:

water is a polar molecule.

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Shape Can Affect Solubility

A useful general idea in chemistry is:

substances with similar polarity often mix more readily.

Water is polar, so many ionic and polar substances dissolve well in water.

Nonpolar substances often dissolve poorly in water.

Molecular shape can therefore indirectly affect solubility because shape influences molecular polarity.

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Shape and Intermolecular Forces

Molecular shape can also affect how molecules interact with each other.

These interactions are called intermolecular forces.

The strength and type of these forces can influence:

• Boiling point.
• Melting point.
• Solubility.
• Viscosity.
• Physical state.

Molecular geometry is therefore connected to the physical properties we observe.

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Shape and Biological Molecules

Molecular shape is extremely important in biology.

Many biological molecules interact because their shapes allow them to fit together.

Examples include:

• Enzymes and substrates.
• Hormones and receptors.
• Antibodies and antigens.
• Drugs and receptor proteins.

A change in molecular shape can sometimes dramatically change how a molecule behaves.

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Why Three-Dimensional Shape Matters

Imagine a key and a lock.

Having the correct material is not enough—the key must also have the correct shape.

Molecules behave similarly.

For some chemical and biological interactions, molecules must approach each other in particular orientations.

Therefore:

structure affects function.

This is one of the most important ideas connecting chemistry and biology.

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Predicting Simple Molecular Shapes

A useful approach is:

Identify the central atom

↓

Determine the number of bonding regions

↓

Determine the number of lone pairs

↓

Arrange electron regions to minimize repulsion

↓

Identify the positions of the atoms

↓

Name the molecular shape

This process works well for many simple molecules.

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Worked Comparison

Consider four molecules:

CO₂

2 bonding regions, 0 lone pairs on the central atom.

Shape:

linear

---

BF₃

3 bonding regions, 0 lone pairs.

Shape:

trigonal planar

---

CH₄

4 bonding regions, 0 lone pairs.

Shape:

tetrahedral

---

NH₃

3 bonding regions, 1 lone pair.

Shape:

trigonal pyramidal

The number and type of electron regions determine the final geometry.

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Common Mistakes

Assuming Molecules Are Flat

Many molecules are three-dimensional.

A structural formula drawn on paper does not necessarily represent the actual molecular shape.

Counting a Double Bond as Two Regions

A double bond normally counts as one region of electron density when predicting basic molecular geometry.

Ignoring Lone Pairs

Lone pairs repel bonding pairs and can significantly affect molecular shape.

Calling Water Linear

H₂O has two lone pairs on oxygen.

Its molecular shape is bent, not linear.

Calling Ammonia Tetrahedral

NH₃ has a tetrahedral electron-pair arrangement, but its molecular shape is trigonal pyramidal.

Assuming Four Bonds Form a Flat Square

Four bonding regions usually arrange tetrahedrally because this provides greater separation in three dimensions.

Assuming Polar Bonds Always Produce a Polar Molecule

Molecular geometry must also be considered.

Symmetrically arranged bond polarities may cancel.

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Check Your Understanding

1. What determines the three-dimensional shape of many simple molecules?

2. What does VSEPR stand for?

3. Explain why electron pairs repel each other.

4. What is the difference between a bonding pair and a lone pair?

5. Identify the shape of CO₂.

6. State the approximate bond angle in a linear molecule.

7. Identify the shape of BF₃.

8. State the approximate bond angle in BF₃.

9. Identify the shape of CH₄.

10. State the approximate bond angle in methane.

11. Why is NH₃ trigonal pyramidal rather than trigonal planar?

12. Why is the bond angle in NH₃ slightly smaller than in CH₄?

13. Explain why H₂O has a bent shape.

14. Compare the molecular shapes of CH₄, NH₃, and H₂O.

15. How is a double bond treated when counting regions of electron density?

16. Explain the difference between electron geometry and molecular geometry.

17. Why is CO₂ overall nonpolar even though its C=O bonds are polar?

18. Why is water polar?

19. Explain how molecular shape can influence solubility.

20. Give one example of why molecular shape is important in biology.

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Key Terms

• Molecular shape – three-dimensional arrangement of atoms within a molecule.
• Molecular geometry – another term describing the three-dimensional arrangement of atoms.
• VSEPR theory – model predicting molecular geometry from repulsion between electron regions.
• Bonding pair – pair of electrons shared between bonded atoms.
• Lone pair – pair of outer-shell electrons not involved in bonding.
• Linear – molecular shape with atoms arranged along a straight line.
• Trigonal planar – arrangement with three bonding regions approximately 120° apart in one plane.
• Tetrahedral – three-dimensional arrangement of four bonding regions around a central atom.
• Trigonal pyramidal – shape produced by three bonded atoms and one lone pair around a central atom.
• Bent – angular molecular shape such as that of water.
• Trigonal bipyramidal – arrangement associated with five electron regions.
• Octahedral – arrangement associated with six electron regions.
• Bond angle – angle between two bonds around a central atom.
• Electron geometry – arrangement of all electron regions around a central atom.
• Polarity – uneven distribution of electrical charge within a bond or molecule.

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Key Takeaways

• Molecules have three-dimensional shapes.
• Molecular shape depends largely on the arrangement of electron regions around atoms.
• Electron pairs repel each other and arrange themselves to minimize repulsion.
• This principle forms the basis of VSEPR theory.
• Bonding pairs and lone pairs both influence molecular geometry.
• Lone pairs generally repel more strongly than bonding pairs.
• CO₂ is linear with a bond angle of 180°.
• BF₃ is trigonal planar with bond angles of approximately 120°.
• CH₄ is tetrahedral with bond angles of approximately 109.5°.
• NH₃ is trigonal pyramidal with bond angles of approximately 107°.
• H₂O is bent with a bond angle of approximately 104.5°.
• Double and triple bonds usually count as one electron region when predicting basic geometry.
• Electron geometry includes lone pairs, while molecular geometry describes the positions of atoms.
• Molecular models help us visualize structures that cannot be shown fully by simple formulas.
• Molecular shape can influence polarity.
• Polarity can influence solubility and intermolecular forces.
• Molecular geometry can therefore affect physical and chemical properties.
• Molecular shape is also extremely important in biological interactions.
• The arrangement of atoms in space is an important part of understanding how molecules behave.

What Is Molecular Polarity?

Molecular polarity describes how electrical charge is distributed across a molecule.

In some molecules, electrons are distributed relatively evenly. These molecules are nonpolar.

In other molecules, electrons are distributed unevenly, producing slightly positive and slightly negative regions. These molecules are polar.

Whether a molecule is polar depends mainly on two things:

• The polarity of its bonds.
• The three-dimensional shape of the molecule.

This means that having polar bonds does not automatically make the entire molecule polar.

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Electronegativity and Bond Polarity

Atoms do not always attract bonding electrons equally.

Electronegativity describes an atom's ability to attract shared electrons in a covalent bond.

If two bonded atoms have different electronegativities, the bonding electrons are attracted more strongly toward one atom.

This produces a polar covalent bond.

The atom attracting the electrons more strongly develops a slight negative charge:

δ−

The other atom develops a slight positive charge:

δ+

These are called partial charges.

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Example: The H–Cl Bond

In hydrogen chloride:

H–Cl

chlorine attracts the shared electrons more strongly than hydrogen.

Therefore:

Hδ+ — Clδ−

The bond is polar.

Because HCl contains only two atoms, there is no second bond pointing in another direction to cancel this polarity.

Therefore:

HCl is a polar molecule.

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Nonpolar Covalent Bonds

If two identical atoms share electrons, they have the same electronegativity.

The electrons are therefore shared equally.

Examples include:

H₂

O₂

N₂

Cl₂

These molecules contain nonpolar covalent bonds and are themselves nonpolar.

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Bond Dipoles

A polar bond creates a bond dipole.

A bond dipole has a direction:

toward the more electronegative atom

We can think of each polar bond as producing a small pull of electron density in a particular direction.

Whether the whole molecule is polar depends on how all of these bond dipoles combine.

This is where molecular shape becomes extremely important.

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Shape and Molecular Polarity

Consider two molecules:

CO₂

and:

H₂O

Both contain polar bonds.

However:

CO₂ is nonpolar

while:

H₂O is polar

Why?

Their molecular shapes are different.

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Carbon Dioxide: Polar Bonds, Nonpolar Molecule

Carbon dioxide has the structure:

O=C=O

Each C=O bond is polar because oxygen attracts electrons more strongly than carbon.

However, CO₂ is:

linear

with a bond angle of:

180°

The two bond dipoles point in opposite directions.

Because they are equal and opposite, they cancel.

Therefore:

CO₂ is nonpolar overall.

This is an important example because it demonstrates that:

polar bonds do not necessarily produce a polar molecule.

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Water: A Polar Molecule

Water contains two polar O–H bonds.

Oxygen attracts the shared electrons more strongly than hydrogen.

Therefore:

• Oxygen has a partial negative charge.
• The hydrogen atoms have partial positive charges.

If water were linear, the bond dipoles could oppose each other.

But water is bent.

Its bond angle is approximately:

104.5°

Because the O–H bond dipoles point in different directions and do not cancel, water has an overall dipole.

Therefore:

H₂O is polar.

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Methane: A Nonpolar Molecule

Methane has the formula:

CH₄

Its molecular shape is:

tetrahedral

The four bonds are arranged symmetrically around the carbon atom.

Any small bond dipoles are distributed symmetrically and cancel overall.

Therefore:

CH₄ is nonpolar.

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Ammonia: A Polar Molecule

Ammonia has the formula:

NH₃

It contains:

• Three N–H bonds.
• One lone pair on nitrogen.

Its molecular shape is:

trigonal pyramidal

The molecule is not symmetrical enough for the bond dipoles to cancel.

Therefore:

NH₃ is polar.

The lone pair is important because it affects the molecular geometry.

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Comparing Common Molecules

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Boron Trifluoride

Boron trifluoride:

BF₃

contains three strongly polar B–F bonds.

However, BF₃ has a:

trigonal planar

shape.

The three identical bonds are arranged symmetrically at approximately:

120°

The bond dipoles cancel because of this symmetry.

Therefore:

BF₃ is nonpolar overall.

Again:

polar bonds + symmetrical geometry can produce a nonpolar molecule.

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Symmetry and Polarity

Symmetry is often useful when predicting molecular polarity.

If a molecule:

• Has identical surrounding atoms.
• Has a symmetrical shape.
• Has bond dipoles arranged equally around the central atom.

the bond dipoles may cancel.

The molecule may therefore be nonpolar.

Examples include:

CO₂

BF₃

CH₄

SF₆

However, symmetry should be used together with an understanding of the bonds and molecular geometry rather than as a rule by itself.

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Asymmetrical Molecules

If polar bonds are arranged asymmetrically, their dipoles usually do not cancel completely.

The molecule then has a net dipole.

Examples include:

H₂O

NH₃

HCl

These molecules are polar.

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Lone Pairs Can Affect Polarity

Lone pairs often influence molecular polarity because they affect molecular shape.

Consider:

CH₄

Four bonds and no lone pairs.

Shape:

tetrahedral

Overall:

nonpolar

Now consider:

NH₃

Three bonds and one lone pair.

Shape:

trigonal pyramidal

Overall:

polar

And:

H₂O

Two bonds and two lone pairs.

Shape:

bent

Overall:

polar

Lone pairs can therefore prevent bond dipoles from cancelling by changing molecular geometry.

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A Method for Predicting Molecular Polarity

When deciding whether a molecule is polar or nonpolar, use the following process.

Identify the bonds

Ask:

Are any of the bonds polar?

If there are no significant bond dipoles, the molecule is generally nonpolar.

Determine the molecular shape

Use the arrangement of bonding regions and lone pairs to determine the three-dimensional shape.

Consider the direction of the bond dipoles

Imagine each bond dipole as an arrow pointing toward the more electronegative atom.

Decide whether the dipoles cancel

If they cancel completely:

nonpolar molecule

If they do not cancel:

polar molecule

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Worked Example: CO₂

Bonds

C=O bonds are polar.

Shape

CO₂ is linear.

Dipole arrangement

The two dipoles are equal and point in opposite directions.

Result

They cancel.

CO₂ is nonpolar.

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Worked Example: H₂O

Bonds

O–H bonds are polar.

Shape

H₂O is bent.

Dipole arrangement

The two bond dipoles do not point directly opposite each other.

Result

They do not cancel.

H₂O is polar.

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Worked Example: NH₃

Bonds

N–H bonds are polar.

Shape

NH₃ is trigonal pyramidal.

Dipole arrangement

The three bond dipoles do not cancel because of the molecular geometry.

Result

NH₃ is polar.

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Worked Example: BF₃

Bonds

B–F bonds are polar.

Shape

BF₃ is trigonal planar.

Dipole arrangement

The three identical bond dipoles are arranged symmetrically.

Result

They cancel.

BF₃ is nonpolar.

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Molecular Dipole

When bond dipoles do not cancel, the molecule has an overall:

molecular dipole

One region of the molecule is relatively:

δ−

while another is relatively:

δ+

This uneven charge distribution influences how the molecule interacts with other molecules.

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Polar Molecules Attract One Another

Polar molecules can attract each other because opposite partial charges attract.

The δ+ region of one molecule can attract the δ− region of another.

These attractions are called:

dipole-dipole forces

These are intermolecular forces because they act between molecules.

They are not the same as the covalent bonds holding atoms together inside a molecule.

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Hydrogen Bonding

Some highly polar molecules can form particularly strong intermolecular attractions called hydrogen bonds.

Hydrogen bonding commonly occurs when hydrogen is bonded directly to:

• Nitrogen.
• Oxygen.
• Fluorine.

Water is an important example.

Hydrogen bonding between water molecules contributes to many of water's unusual properties.

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Polarity and Solubility

Polarity strongly influences whether substances dissolve in one another.

A useful general rule is:

"like dissolves like."

This means:

• Polar substances tend to dissolve more readily in polar solvents.
• Nonpolar substances tend to dissolve more readily in nonpolar solvents.

This is a useful guideline rather than an absolute rule.

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Water as a Polar Solvent

Water is strongly polar.

It can interact with:

• Ions.
• Polar molecules.

This helps explain why many salts and polar substances dissolve well in water.

When sodium chloride dissolves, for example, polar water molecules surround the Na⁺ and Cl⁻ ions.

The oxygen side of water is attracted toward positive ions.

The hydrogen side is attracted toward negative ions.

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Why Oil and Water Separate

Water is polar.

Most oils consist largely of nonpolar molecules.

The interactions between water molecules are much stronger and more favorable than interactions between water and many oil molecules.

As a result, oil and water tend to separate into different layers.

This familiar observation is an everyday example of molecular polarity.

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Soap and Polarity

Soap and detergent molecules have an interesting structure.

They usually contain:

• A polar or ionic hydrophilic head.
• A nonpolar hydrophobic tail.

The head interacts with water.

The tail interacts with oils and grease.

This allows soap molecules to form structures called micelles, trapping grease so that it can be carried away by water.

Polarity therefore explains an important part of how soaps and detergents work.

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Polarity and Boiling Point

Molecular polarity can affect boiling point.

If molecules attract each other strongly, more energy is required to separate them into the gas phase.

Polar molecules may experience:

• Dipole-dipole attractions.
• Hydrogen bonding in appropriate molecules.

Nonpolar molecules mainly experience London dispersion forces.

However, boiling point also depends on factors such as:

• Molecular size.
• Shape.
• Number of electrons.
• Strength of all intermolecular forces.

Therefore, polarity alone does not determine boiling point.

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Polarity and Biological Molecules

Polarity is extremely important in biology.

Cell membranes contain molecules called phospholipids.

A phospholipid contains:

• A polar head.
• Nonpolar tails.

In water, phospholipids naturally arrange themselves into a bilayer.

The polar heads face the watery environments.

The nonpolar tails point inward, away from the water.

This arrangement is a fundamental feature of cell membranes.

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Polarity and Molecular Interactions

Polarity can influence:

• Solubility.
• Boiling point.
• Melting point.
• Surface tension.
• Intermolecular attractions.
• Biological membrane behavior.
• Protein interactions.
• Drug interactions.

This demonstrates an important chemistry relationship:

molecular structure → polarity → intermolecular forces → properties

---

Comparing Polar and Nonpolar Molecules

These are general patterns. Actual physical properties also depend on other aspects of molecular structure.

---

Predicting Polarity from a Molecular Model

Suppose you are shown a molecular model.

Ask:

What atoms are bonded together?

↓

Which bonds are polar?

↓

What is the three-dimensional molecular shape?

↓

Is the arrangement symmetrical?

↓

Do the bond dipoles cancel?

↓

If yes:

NONPOLAR

If no:

POLAR

This method is much more reliable than simply looking for polar bonds.

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Challenge Example: CCl₄

Carbon tetrachloride has the formula:

CCl₄

Each C–Cl bond is polar.

Its shape is:

tetrahedral

The four identical chlorine atoms are arranged symmetrically around carbon.

The bond dipoles cancel.

Therefore:

CCl₄ is nonpolar overall.

This is another example of a molecule containing polar bonds but having no overall molecular dipole.

---

Challenge Example: CH₃Cl

Now replace one chlorine atom in CCl₄ with hydrogen:

CH₃Cl

The molecule is still approximately tetrahedral around carbon.

However, the surrounding atoms are no longer identical.

The bond dipoles do not cancel completely.

Therefore:

CH₃Cl is polar.

This comparison shows how changing just one atom can change the polarity of a molecule.

---

Real-World Connection: Removing Grease

Water alone often does not remove grease effectively.

Water is:

polar

Grease is largely:

nonpolar

They interact poorly.

Soap provides a connection between them because it contains both polar and nonpolar regions.

The nonpolar tails interact with grease while the polar heads interact with water.

This allows greasy material to be dispersed and washed away.

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Real-World Connection: Food

Polarity affects many processes in food chemistry.

For example:

• Sugar dissolves readily in water.
• Many oils do not.
• Emulsifiers help mixtures containing water and oil remain mixed.

Food products such as mayonnaise rely on molecules that help stabilize mixtures containing both polar and nonpolar substances.

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Common Mistakes

Assuming Every Molecule with Polar Bonds Is Polar

CO₂, BF₃, CCl₄ and other symmetrical molecules can contain polar bonds while remaining nonpolar overall.

Ignoring Molecular Shape

Polarity cannot always be predicted from a molecular formula alone.

Three-dimensional geometry matters.

Assuming CO₂ Is Polar Because Oxygen Is Electronegative

The C=O bonds are polar, but their dipoles cancel because CO₂ is linear.

Assuming Water Is Linear

Water is bent because oxygen has two lone pairs.

Its bond dipoles therefore do not cancel.

Confusing Bond Polarity With Molecular Polarity

A bond can be polar while the entire molecule is nonpolar.

Assuming Nonpolar Means No Electrons Are Moving

Electrons are still present and moving. Nonpolar means there is no permanent overall separation of charge across the molecule.

Assuming Polarity Is the Only Factor Affecting Properties

Molecular size, shape, hydrogen bonding and other intermolecular forces also matter.

---

Check Your Understanding

1. Define a polar molecule.

2. Define a nonpolar molecule.

3. What is electronegativity?

4. Explain how electronegativity differences produce polar bonds.

5. What do δ+ and δ− represent?

6. Explain what is meant by a bond dipole.

7. Why is HCl polar?

8. Why is O₂ nonpolar?

9. Explain why CO₂ is nonpolar even though it contains polar bonds.

10. Explain why H₂O is polar.

11. Why does molecular shape affect polarity?

12. Is CH₄ polar or nonpolar? Explain.

13. Is NH₃ polar or nonpolar? Explain.

14. Explain why BF₃ is nonpolar.

15. Predict whether CCl₄ is polar or nonpolar and explain your reasoning.

16. Why is CH₃Cl polar even though it has an approximately tetrahedral shape?

17. Explain the phrase "like dissolves like."

18. Why do oil and water tend to separate?

19. Explain how soap uses both polar and nonpolar regions to remove grease.

20. Describe how polarity contributes to the structure of cell membranes.

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Key Terms

• Polar molecule – molecule with an uneven distribution of electrical charge and an overall molecular dipole.
• Nonpolar molecule – molecule with no permanent overall separation of charge.
• Electronegativity – ability of an atom to attract shared bonding electrons.
• Polar covalent bond – covalent bond in which electrons are shared unequally.
• Partial charge – slight positive or negative charge produced by unequal electron distribution.
• δ+ – partial positive charge.
• δ− – partial negative charge.
• Bond dipole – separation of charge across a polar bond.
• Molecular dipole – overall separation of charge across a molecule.
• Dipole-dipole force – attraction between oppositely charged regions of polar molecules.
• Hydrogen bond – particularly strong intermolecular attraction involving hydrogen bonded to certain highly electronegative atoms.
• Hydrophilic – attracted to or interacting favorably with water.
• Hydrophobic – tending to avoid interaction with water.
• Symmetry – balanced arrangement of parts around a molecule.
• Intermolecular force – attraction acting between separate molecules.

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Key Takeaways

• Molecular polarity describes how electrical charge is distributed across a molecule.
• Polar molecules have an uneven distribution of charge.
• Nonpolar molecules have no permanent overall separation of charge.
• Differences in electronegativity can produce polar covalent bonds.
• Polar bonds contain partial positive and partial negative regions.
• Molecular polarity depends on both bond polarity and molecular shape.
• Polar bonds do not automatically make a molecule polar.
• Bond dipoles can cancel when arranged symmetrically.
• CO₂ contains polar bonds but is nonpolar because its linear bond dipoles cancel.
• BF₃ contains polar bonds but is nonpolar because of its symmetrical trigonal planar geometry.
• CH₄ is nonpolar because of its symmetrical tetrahedral geometry.
• H₂O is polar because its bent geometry prevents its bond dipoles from cancelling.
• NH₃ is polar because its trigonal pyramidal geometry produces a net molecular dipole.
• Lone pairs can influence polarity by changing molecular geometry.
• Polar molecules can experience dipole-dipole attractions.
• Some polar molecules can form hydrogen bonds.
• Polarity strongly influences solubility.
• Polar substances often interact well with polar solvents such as water.
• Nonpolar substances often interact more readily with other nonpolar substances.
• Soap works because its molecules contain both water-attracting and oil-attracting regions.
• Polarity is important in chemistry, biology, environmental science, food science and everyday life.
• To predict molecular polarity, examine the bonds, determine the molecular shape, and decide whether the bond dipoles cancel.

From Molecular Structure to Physical Properties

Molecular substances are made of individual molecules held together by intermolecular forces.

Within each molecule, atoms are joined by strong covalent bonds. Between the molecules, weaker intermolecular attractions act.

These intermolecular forces strongly influence physical properties such as:

• Melting point
• Boiling point
• Physical state
• Evaporation rate
• Volatility
• Viscosity
• Surface tension
• Solubility

A useful relationship is:

molecular structure → intermolecular forces → physical properties

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Molecular Substances

A molecular substance consists of separate molecules.

Examples include:

• H₂O
• CO₂
• CH₄
• NH₃
• O₂
• I₂
• Ethanol
• Glucose

The atoms within each molecule are held together by covalent bonds.

The molecules themselves interact through intermolecular forces.

This distinction is essential when explaining physical properties.

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Intermolecular Forces

The main intermolecular attractions considered for molecular substances are:

London dispersion forces

Present between all atoms and molecules.

They result from temporary fluctuations in electron distribution.

Dipole-dipole attractions

Occur between polar molecules.

The partially positive region of one molecule attracts the partially negative region of another.

Hydrogen bonding

A particularly strong intermolecular attraction occurring when hydrogen is bonded to:

• N
• O
• F

Water is an important example.

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Strength of Intermolecular Forces

For small molecules of broadly similar size, a useful general pattern is:

London dispersion < dipole-dipole < hydrogen bonding

However, this is not an absolute rule.

London dispersion forces can become very strong in large molecules containing many electrons.

Therefore, when comparing molecular substances, we should consider:

• Polarity
• Hydrogen bonding
• Molecular size
• Number of electrons
• Molecular shape

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Melting

Melting is the change:

solid → liquid

In a molecular solid, molecules are held in relatively fixed positions by intermolecular attractions.

When the substance is heated:

• Molecules gain kinetic energy.
• Molecular vibrations increase.
• Intermolecular attractions become less able to hold molecules in fixed positions.
• Molecules begin moving past one another.

The substance becomes a liquid.

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Melting Does Not Normally Break Covalent Bonds

When a molecular substance melts, the molecules themselves generally remain intact.

For example:

H₂O(s) → H₂O(l)

Ice contains H₂O molecules.

Liquid water also contains H₂O molecules.

The O–H covalent bonds are not normally broken during melting.

Instead, the arrangement and intermolecular interactions between water molecules change.

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Melting Point

The melting point is the temperature at which a solid changes into a liquid under a specified pressure.

Stronger intermolecular attractions often require more energy to overcome.

Therefore, stronger intermolecular forces can contribute to:

higher melting points

However, melting points can be difficult to predict from intermolecular force strength alone because the way molecules pack together in a solid crystal is also important.

Molecular symmetry and shape can strongly influence crystal structure.

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Boiling

Boiling is the change:

liquid → gas

For molecules to enter the gas phase, they must become sufficiently separated from neighboring molecules.

Energy is required to overcome intermolecular attractions.

Therefore:

stronger intermolecular forces generally → higher boiling point

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Boiling Point and Molecular Attractions

Compare two similarly sized substances.

Substance A

Weak intermolecular forces.

Molecules separate relatively easily.

Expected:

lower boiling point

Substance B

Strong intermolecular forces.

More energy is needed to separate the molecules.

Expected:

higher boiling point

This relationship is one of the most useful ways to connect microscopic molecular structure with observable properties.

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Boiling Does Not Break Molecules Apart

When water boils:

H₂O(l) → H₂O(g)

The molecules separate from one another.

The O–H covalent bonds remain intact.

Therefore, boiling mainly involves overcoming:

intermolecular forces

not:

covalent bonds

If the covalent bonds were broken, water would undergo a chemical change rather than simply boiling.

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Example: Methane and Water

Consider:

CH₄

and:

H₂O

Methane is:

• Nonpolar
• Small
• Mainly affected by relatively weak London dispersion forces

Water is:

• Polar
• Capable of hydrogen bonding

Therefore, water molecules experience much stronger intermolecular attractions.

This helps explain why methane is a gas at room temperature while water is a liquid.

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Hydrogen Bonding and Boiling Point

Hydrogen bonding can have a major effect on boiling point.

Water molecules form extensive hydrogen bonds with neighboring molecules.

A significant amount of energy is required to separate water molecules sufficiently for them to enter the gas phase.

This contributes to water having a much higher boiling point than might be expected for such a small molecule.

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Molecular Size and Boiling Point

Polarity is not the only factor affecting intermolecular forces.

Molecular size also matters.

Larger molecules generally have:

• More electrons.
• Larger electron clouds.
• More easily distorted electron distributions.

This can produce stronger London dispersion forces.

Therefore, within a similar family of molecules:

larger molecules often have higher boiling points.

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Example: The Halogens

Consider:

F₂, Cl₂, Br₂, I₂

These molecules are nonpolar.

Their main intermolecular attractions are London dispersion forces.

Moving down the group:

• Molecular size increases.
• Number of electrons increases.
• Polarizability increases.
• London dispersion forces become stronger.

This contributes to their different physical states around room temperature:

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Molecular Shape Matters

The shape of a molecule can influence how closely molecules approach one another.

Long or relatively flat molecules may sometimes have larger areas of contact than compact molecules.

Greater contact can strengthen London dispersion interactions.

Therefore:

molecular shape can influence boiling point and other physical properties.

This is particularly useful when comparing molecules with similar molecular masses.

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Evaporation

Evaporation occurs when molecules escape from the surface of a liquid into the gas phase.

Unlike boiling, evaporation can occur:

below the boiling point

and only occurs at the:

surface of the liquid

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Why Do Some Molecules Evaporate?

Molecules in a liquid have different kinetic energies.

Some molecules at the surface have enough energy to overcome intermolecular attractions and escape.

These molecules enter the gas phase.

Therefore, evaporation depends partly on:

• Intermolecular force strength.
• Temperature.
• Surface area.
• Air movement.
• Humidity.

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Intermolecular Forces and Evaporation

If intermolecular forces are weak:

molecules escape more easily

Therefore:

evaporation tends to be faster.

If intermolecular forces are strong:

more energy is required for molecules to escape

Therefore:

evaporation tends to be slower.

So, in general:

stronger intermolecular forces → slower evaporation

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Temperature and Evaporation

Increasing temperature increases the average kinetic energy of molecules.

At a higher temperature, a larger fraction of surface molecules have enough energy to escape.

Therefore:

higher temperature → faster evaporation

This is why wet clothing usually dries more quickly in warm conditions.

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Surface Area and Evaporation

Evaporation occurs at the liquid surface.

Increasing the surface area exposes more molecules to the air.

Therefore:

larger surface area → faster evaporation

For example, water spread across a wide tray usually evaporates faster than the same volume of water in a narrow container.

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Air Movement and Evaporation

Molecules that evaporate accumulate in the air above a liquid.

Moving air carries some of these vapor molecules away.

This helps maintain conditions that allow continued evaporation.

Therefore:

greater air movement often → faster evaporation

This helps explain why clothes dry faster on a windy day.

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Humidity and Evaporation

Humidity describes the amount of water vapor present in the air.

When humidity is high, the air already contains a large amount of water vapor.

Net evaporation from wet surfaces tends to be slower.

When humidity is low, evaporation can occur more readily.

Therefore:

lower humidity generally → faster evaporation of water

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Factors Affecting Evaporation

These factors explain many familiar observations.

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Volatility

Volatility describes how readily a substance vaporizes.

A highly volatile liquid evaporates readily.

Highly volatile molecular substances generally have:

• Relatively weak intermolecular forces.
• Relatively high vapor pressure.
• Relatively low boiling points.

Therefore:

higher volatility generally corresponds to easier evaporation.

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Vapor Pressure

Some molecules continuously escape from a liquid surface.

In a closed container, some gas molecules also return to the liquid.

Eventually, a dynamic equilibrium can develop between:

evaporation ⇌ condensation

The gas above the liquid produces a vapor pressure.

Liquids with weaker intermolecular forces generally have higher vapor pressures at the same temperature.

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Connecting Vapor Pressure and Boiling Point

A liquid boils when its vapor pressure becomes equal to the external pressure.

Therefore, substances with weaker intermolecular attractions often:

• Develop higher vapor pressures more easily.
• Reach the boiling condition at lower temperatures.

This gives the general relationship:

weaker intermolecular forces → higher vapor pressure → lower boiling point

and:

stronger intermolecular forces → lower vapor pressure → higher boiling point

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Viscosity

Viscosity is a fluid's resistance to flow.

Water has relatively low viscosity compared with substances such as:

• Honey
• Syrup
• Thick oils

Stronger intermolecular attractions can make it more difficult for molecules to move past one another.

Therefore, stronger intermolecular forces can contribute to:

greater viscosity

However, molecular size and shape are also important.

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Molecular Shape and Viscosity

Long molecules may become entangled with one another.

This can make them resist flowing past each other.

Therefore, viscosity depends on both:

• Intermolecular attractions.
• Molecular structure.

Large, long molecules can sometimes produce highly viscous liquids even when other factors are similar.

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Surface Tension

Surface tension is the tendency of a liquid surface to resist being stretched or disrupted.

Molecules inside a liquid experience attractions from molecules around them.

Surface molecules experience an overall inward attraction.

Strong intermolecular attractions can therefore produce high surface tension.

Water has relatively high surface tension because of hydrogen bonding.

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Physical State at Room Temperature

Intermolecular forces help determine whether a molecular substance is:

• Solid.
• Liquid.
• Gas.

If intermolecular attractions are very weak, molecules may separate easily and the substance may be a gas.

Stronger attractions can allow molecules to remain together as liquids or solids.

However, physical state also depends on:

• Temperature.
• Pressure.
• Molecular size.
• Molecular shape.

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Comparing Molecular Substances

Consider three hypothetical molecular substances:

This prediction assumes the molecules are reasonably similar in size.

If Substance A were dramatically larger than B or C, its London dispersion forces could become very significant.

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Predicting Properties from Structure

When given a molecular structure, use this process:

Determine whether the molecule is polar

↓

Identify the intermolecular forces

↓

Consider hydrogen bonding

↓

Consider molecular size and number of electrons

↓

Consider molecular shape

↓

Compare overall intermolecular attraction

↓

Predict physical properties

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Worked Example: CH₄ and NH₃

Methane, CH₄

Shape:

tetrahedral

Polarity:

nonpolar

Main intermolecular attraction:

London dispersion

Expected properties:

• Low boiling point.
• High volatility.
• Gas at room temperature.

Ammonia, NH₃

Shape:

trigonal pyramidal

Polarity:

polar

Important intermolecular attraction:

hydrogen bonding

Expected:

• Stronger intermolecular attraction than methane.
• Higher boiling point than methane.

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Worked Example: H₂O and H₂S

Water and hydrogen sulfide have somewhat similar molecular structures.

However:

H₂O forms strong hydrogen bonds.

H₂S does not form comparably strong conventional hydrogen bonds.

Therefore, water has much stronger intermolecular attractions.

This contributes to water having a much higher boiling point.

The difference cannot be explained simply by molecular mass.

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Worked Example: Br₂ and Cl₂

Both:

Br₂

and:

Cl₂

are nonpolar.

Both rely mainly on London dispersion forces.

However, Br₂:

• Has more electrons.
• Has a larger electron cloud.
• Is more polarizable.

Therefore:

Br₂ has stronger London dispersion forces.

This contributes to Br₂ having a higher boiling point than Cl₂.

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Comparing Evaporation

Suppose three similar liquids have the following intermolecular forces:

Liquid A: weak London forces

Liquid B: dipole-dipole attractions

Liquid C: strong hydrogen bonding

At the same temperature, we would generally expect:

A to evaporate fastest

and:

C to evaporate slowest

because the molecules in C require more energy to escape from the liquid.

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Comparing Boiling Points

Suppose two molecules have similar molecular masses.

Molecule X is nonpolar.

Molecule Y can form hydrogen bonds.

We would generally predict:

Molecule Y has the higher boiling point.

Reason:

Molecule Y has stronger intermolecular attractions, so more energy is required to separate its molecules.

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Real-World Connection: Sweating

Sweating cools the body through evaporation.

The fastest-moving water molecules can escape from the skin.

Energy is required to overcome intermolecular attractions between water molecules.

That energy comes partly from the skin.

As higher-energy molecules leave, the average kinetic energy of the remaining surface decreases.

The skin cools.

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Real-World Connection: Drying Clothes

Several factors affect how quickly clothes dry.

Warm weather

Higher temperature increases molecular kinetic energy.

Wind

Moving air carries water vapor away.

Low humidity

Dry air allows greater net evaporation.

Spreading clothes out

Greater surface area exposes more water to the air.

This is why clothes generally dry faster when they are:

warm + spread out + exposed to moving, relatively dry air

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Real-World Connection: Perfume

Many perfume ingredients are relatively volatile.

They evaporate from the skin or clothing and enter the air.

Gas molecules then spread through the surroundings.

For a fragrance to reach your nose, at least some of its molecules must enter the gas phase.

The volatility of the substances therefore influences how a fragrance behaves.

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Real-World Connection: Refrigeration

Refrigeration systems use substances that repeatedly:

• Evaporate.
• Condense.

During evaporation, the refrigerant absorbs energy.

During condensation, energy is released elsewhere.

The physical properties of the refrigerant must therefore be suitable for these repeated phase changes.

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Intermolecular Forces and Solubility

Physical properties are interconnected.

The same molecular features that affect boiling point and evaporation can also affect solubility.

Polar molecules can often interact strongly with polar solvents.

Nonpolar molecules often interact more favorably with nonpolar solvents.

Therefore, understanding molecular structure can help predict several properties at once.

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A Useful Property Pattern

For comparable molecular substances:

stronger intermolecular forces

often mean:

↑ boiling point

↓ volatility

↓ evaporation rate

↓ vapor pressure at a given temperature

and sometimes:

↑ viscosity

↑ surface tension

Melting point may also increase, but crystal packing makes melting-point trends less straightforward.

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Structure-to-Property Reasoning

A strong chemistry explanation should connect several steps.

Instead of saying:

“Water has a high boiling point because it is polar.”

A stronger explanation is:

Water molecules are polar and can form hydrogen bonds with neighboring water molecules. These relatively strong intermolecular attractions require considerable energy to overcome, contributing to water's relatively high boiling point.

This reasoning pattern can be used for many molecular substances.

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Example of a Full Prediction

Suppose a molecule:

• Is small.
• Is nonpolar.
• Cannot form hydrogen bonds.

We predict that its main intermolecular attractions will be relatively weak London dispersion forces.

Therefore, compared with a similarly sized molecule capable of hydrogen bonding, it will probably have:

• Lower boiling point.
• Higher volatility.
• Faster evaporation.
• Higher vapor pressure.

This is how molecular structure can be used to predict macroscopic properties.

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Common Mistakes

Saying Boiling Breaks Covalent Bonds

Boiling molecular substances normally overcomes intermolecular forces, not covalent bonds within the molecules.

Assuming Molecular Mass Is the Only Factor

Polarity, hydrogen bonding and molecular shape also matter.

Assuming Hydrogen Bonding Occurs in Every Molecule Containing Hydrogen

For the molecules studied here, hydrogen bonding generally requires H to be directly bonded to:

N, O, or F

Saying Stronger Forces Mean Faster Evaporation

The opposite is generally true.

Stronger intermolecular forces → slower evaporation

Confusing Evaporation and Boiling

Evaporation occurs at the surface and can occur below the boiling point.

Boiling occurs throughout the liquid when the vapor pressure equals the external pressure.

Assuming Melting Point Always Follows Boiling-Point Trends

Crystal packing can strongly affect melting point.

Forgetting London Dispersion Forces

All molecular substances experience London dispersion forces.

Assuming Large Nonpolar Molecules Always Have Weak Attractions

Large electron clouds can produce substantial London dispersion forces.

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Check Your Understanding

1. What holds atoms together within a molecular substance?

2. What holds separate molecules together?

3. Name three important intermolecular forces.

4. Explain why stronger intermolecular forces generally produce higher boiling points.

5. What happens to intermolecular attractions when a molecular substance boils?

6. Are covalent bonds normally broken during boiling? Explain.

7. Why does hydrogen bonding increase the boiling point of many substances?

8. Explain why water has a much higher boiling point than methane.

9. Why do London dispersion forces generally become stronger as molecular size increases?

10. Explain why Br₂ has stronger intermolecular forces than Cl₂.

11. Define evaporation.

12. Explain how intermolecular force strength affects evaporation rate.

13. How does increasing temperature affect evaporation?

14. Why does increasing surface area increase evaporation?

15. Explain why wind can make wet clothes dry faster.

16. How does humidity affect the evaporation of water?

17. Define volatility.

18. How are volatility and boiling point generally related?

19. Explain how intermolecular forces affect vapor pressure.

20. Why can molecular shape affect physical properties?

21. Explain how intermolecular forces can influence viscosity.

22. Why does water have relatively high surface tension?

23. A small nonpolar molecule and a similarly sized hydrogen-bonding molecule are compared. Predict which will have the higher boiling point and explain why.

24. A liquid evaporates extremely quickly at room temperature. What might this suggest about its intermolecular forces?

25. Explain the relationship:

molecular structure → intermolecular forces → physical properties

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Key Terms

• Molecular substance – substance consisting of separate covalently bonded molecules.
• Intermolecular force – attraction between separate molecules.
• London dispersion force – intermolecular attraction caused by temporary fluctuations in electron distribution.
• Dipole-dipole attraction – attraction between permanent dipoles of polar molecules.
• Hydrogen bonding – relatively strong intermolecular attraction involving H bonded to N, O, or F.
• Melting point – temperature at which a solid becomes a liquid at a specified pressure.
• Boiling point – temperature at which the vapor pressure of a liquid equals the external pressure.
• Evaporation – escape of molecules from the surface of a liquid.
• Volatility – tendency of a substance to vaporize readily.
• Vapor pressure – pressure produced by vapor above a liquid in a closed system.
• Viscosity – resistance of a fluid to flow.
• Surface tension – tendency of a liquid surface to resist disruption.
• Polarizability – ease with which an electron cloud can be distorted.
• Physical property – characteristic that can be observed or measured without changing a substance's chemical identity.

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Key Takeaways

• Molecular substances consist of separate molecules.
• Covalent bonds hold atoms together within molecules.
• Intermolecular forces act between molecules.
• Intermolecular forces strongly influence physical properties.
• Stronger intermolecular attractions generally lead to higher boiling points.
• Melting points are influenced by intermolecular forces as well as molecular packing.
• Boiling and melting normally do not break covalent bonds within molecules.
• London dispersion forces occur in all molecular substances.
• Larger, more polarizable molecules generally experience stronger London dispersion forces.
• Polar molecules can experience dipole-dipole attractions.
• Suitable molecules containing H bonded to N, O, or F can form hydrogen bonds.
• Hydrogen bonding can strongly influence physical properties.
• Evaporation occurs at the surface of a liquid and can occur below its boiling point.
• Weak intermolecular forces generally produce faster evaporation and greater volatility.
• Temperature, surface area, air movement and humidity can affect evaporation.
• Stronger intermolecular forces generally produce lower vapor pressure at a given temperature.
• Molecular size and shape can influence intermolecular attractions.
• Intermolecular forces can affect viscosity and surface tension.
• Molecular structure can therefore be used to predict observable physical behaviour.
• Understanding physical properties requires connecting molecular structure to intermolecular forces and then connecting those forces to the behaviour of the substance.