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Water and carbon dioxide are both triatomic molecules. The similarities go even further - they are both partially formed from oxygen and they both contain covalent bonds. However, their molecules are very different in shape. Whilst the atoms in carbon dioxide are held in a straight line, water is a bent molecule. To understand why this is the case, we need to take a look at electron pair repulsion, VSEPR.
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Send FeedbackValence shell electron pair repulsion theory, or VSEPR, is a model used in chemistry to predict the shape of molecules.
If we break that term down a little, we can understand what it means.
You should know that electrons tend to go around in pairs. This is because orbitals, which are regions of space where electrons can be found 95 percent of the time, can contain at most two electrons (check out Electron Shells, Subshells, and Orbitals for a refresher). Because electrons are charged particles, electron pairs will repel each other and try to be as far away from each other as possible. An atom’s outer shell of electrons is known as its valence shell. Because the valence electrons in a simple covalent molecule are the bonded electrons, electron pair repulsion determines the way in which the bonds are positioned. This dictates the shape of the molecule.
VSEPR states that electron pairs all repel each other and will try to take up positions as far away from each other as possible, in order to minimise repulsion. It simply uses our knowledge of the behaviour of electrons to predict the shape of simple covalent compounds. Check out Covalent and Dative Bonding to remind yourself of how atoms share electrons in order to achieve stable electron configurations.
Before we look at any examples of covalent structures, we need to learn how you can represent them. You might remember that we can draw covalent bonds as a line between two atoms. This gives a simple picture of molecules. However, if we want to better show a molecule’s 3D shape, we can use wedged and dotted lines.
The methane molecule is a good example of this:
A methane molecule, CH4. The central wedged bond protrudes out of the screen whilst the right hand dashed bond extends backwards. commons.wikimedia.org
If all the pairs of valence electrons in an atom are bonded, they will all repel each other mutually. This results in bonds spaced equally far apart. The number of bonded electron pairs affects the shape of the molecule and the angle between the bonding pairs.
Let’s take a look at some of the most common shapes. However, you should bear in mind that these rules only apply to molecules with no lone pairs of electrons. Lone pairs of electrons are unshared pairs that aren’t covalently bonded. We’ll explore their effect further later.
If a molecule only has two bonded electron pairs (and no lone pairs), it forms a linear molecule. The simplest example is beryllium chloride,
Beryllium chloride. Each of beryllium's valence electrons forms a covalent bond with a chlorine atom. The angle between the bonds around the central beryllium atom is 180°.StudySmarter Originals
Molecules with three bonded electron pairs are known as trigonal planar. This is because the bond angle between each bond is 120°, so the bonds lie flat on a plane. You could stack the molecules up one on top of the other like sheets of paper. Boron trifluoride is an example.
Boron trifluoride. The bonds are held apart at an angle of 120°.StudySmarter Originals
Molecules with four bonded electron pairs and no lone pairs form a tetrahedral shape. This is a regular triangular-based pyramid. All the bond angles are 109.5°. For example, the carbon in methane
Methane. The angle between each bond around the central atom is 109.5°.StudySmarter Originals
Molecules with five bonded electron pairs form a trigonal bipyramid. This shape is similar to a trigonal planar molecule but with two further bonds held at 90° extending above and below the plane. Phosphorus(V) pentachloride is a good example.
Phosphorus(V) pentachloride. Three planar bonds have angles of 120° between them, whilst two further bonds are held at right angles to the plane.StudySmarter Originals
If a molecule has six bonding pairs around a central atom, it forms an octahedral structure. All of the bonds are at right angles to each other, as shown in sulfur hexafluoride.
Sulfur hexafluoride has six bonded electron pairs. All bond angles are 90°. StudySmarter Originals
All of our above examples use molecules that don’t have any lone pairs of electrons. All their valence electrons are bonded. But what happens if a molecule does have a lone pair? Let’s take a molecule with four electron pairs as an example.
We now know that if all of the electrons are part of bonding pairs, the molecule will be tetrahedral and have bond angles of 109.5°. However, if one of the electron pairs is in fact a lone pair, the bond angles are reduced to 107°. This is because lone pairs repel each other more strongly than shared pairs, squeezing the bonds together. Each lone electron pair in a molecule with eight valence electrons reduces the bond angle by 2.5°, so a molecule with two bonding pairs and two lone pairs will have a bond angle of 104.5°. The following table shows the relative strength of repulsion between combinations of bonded and lone pairs of electrons.
A table comparing the strength of repulsion between bonded and lone pairs of electrons. StudySmarter Originals
Let’s now look at the shapes formed by molecules with lone pairs.
A molecule with three bonded electron pairs and one lone electron pair around a central atom has an angle of 107° between each bond. An example is ammonia,
An ammonia molecule. Compared to a tetrahedral molecule with no lone pairs, the bond angle is reduced by 2.5°. StudySmarter Originals
A molecule with two lone pairs and two bonding pairs has its bond angle reduced even further to 104.5°. This forms a v-shaped molecule, such as water,
A v-shaped water molecule. StudySmarter Originals
The following diagram summarises the different shapes of molecules.
A table summarising the shapes of different molecules. StudySmarter Originals
Let’s go back to our original molecules, water and carbon dioxide. We’ve already discovered that water has a v-shaped structure due to the effect of its lone electron pairs on the bonding pairs. But what sort of shape does carbon dioxide have?
By drawing a dot and cross diagram we can see that carbon dioxide,
Carbon dioxide. Although it contains four bonding pairs of electrons, the pairs are arranged as two double bonds. Each double bond is considered as a single unit, so the molecule is linear. commons.wikimedia.org
Another example is xenon tetrafluoride,
Xenon tetrafluoride. The lone pairs of electrons are positioned above and below the plane.StudySmarter Originals
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What is the shape of a water molecule?
Water molecules are v-shaped.
What is the shape of the DNA molecule called?
DNA forms a double helix shape.
What is the shape of the methane molecule?
Methane molecules are tetrahedral in shape.
What is the shape of the xenon tetrafluoride molecule?
Xenon tetrafluoride is square planar in shape.
How do you work out the shape of a molecule?
To work out the shape of a molecule, identify how many lone and bonding pairs of electrons it has. This dictates its shape. For example, the oxygen atom in a water molecule has two lone pairs and two bonding pairs. This gives it a v-shaped structure.
How can the shape of a molecule affect its polarity?
Molecules with polar bonds are often polar molecules. However, if the molecule is symmetrical, the charges of the polar bonds cancel out and the molecule is non-polar overall.
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