Molecular Structure and Intermolecular Forces
1. Molecular Shapes
Learning outcomes
- I can describe how atoms are arranged in molecules.
- I can identify common molecular shapes.
- I can explain how electron pairs influence molecular geometry.
- I can interpret simple molecular models.
- I can relate molecular shape to molecular properties.
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.
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.
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.
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.
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
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.
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.
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.
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.
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.
| Molecule | Bonding pairs | Lone pairs | Shape | Approximate bond angle |
|---|---|---|---|---|
| CH₄ | 4 | 0 | Tetrahedral | 109.5° |
| NH₃ | 3 | 1 | Trigonal pyramidal | 107° |
| H₂O | 2 | 2 | Bent | 104.5° |
Notice the pattern:
More lone-pair repulsion → smaller bond angles between bonded atoms
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°
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°
Summary of Common Molecular Shapes
| Electron arrangement | Example | Molecular shape | Typical bond angle |
|---|---|---|---|
| 2 bonding regions | CO₂ | Linear | 180° |
| 3 bonding regions | BF₃ | Trigonal planar | 120° |
| 4 bonding regions | CH₄ | Tetrahedral | 109.5° |
| 3 bonds + 1 lone pair | NH₃ | Trigonal pyramidal | about 107° |
| 2 bonds + 2 lone pairs | H₂O | Bent | about 104.5° |
| 5 bonding regions | PCl₅ | Trigonal bipyramidal | 90° and 120° |
| 6 bonding regions | SF₆ | Octahedral | 90° |
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.
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°
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°
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°
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.
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.