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.
Valence 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.
How do you draw the shapes of molecules in 3D?
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.
- Wedged lines show a bond coming out of the screen or page towards you.
- Dotted or dashed lines show a bond going into the screen or page away from you.
- Lone pairs of electrons are shown as dots.
- Any standard straight lines simply show a planar bond.
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
The various shapes of molecules
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.
Linear
If a molecule only has two bonded electron pairs (and no lone pairs), it forms a linear molecule. The simplest example is beryllium chloride, \(BeCl_2\) . Although beryllium is a metal, it can bond covalently to chlorine. Beryllium only has two electrons in its valence shell and so forms two bonds. The electron pairs will repel each other equally, resulting in an angle between the two bonds of 180°.
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°.Vaia Originals
Trigonal planar
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°.Vaia Originals
Tetrahedral
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 \(CH_4\) has four valence electrons, and each electron is part of a pair bonded covalently to a hydrogen atom. It is a tetrahedral molecule.
Methane. The angle between each bond around the central atom is 109.5°.Vaia Originals
Trigonal bipyramidal
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.Vaia Originals
Octahedral
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°. Vaia Originals
Lone pairs of electrons
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. Vaia Originals
Let’s now look at the shapes formed by molecules with lone pairs.
Pyramidal
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, \(NH_3\) . The nitrogen atom contains five valence electrons. Three are covalently bonded to hydrogen atoms and the remaining two form a lone pair. This lone pair repels the bonding pairs more strongly than the bonding pairs repel each other, reducing the bond angle and forming a pyramidal molecule.
An ammonia molecule. Compared to a tetrahedral molecule with no lone pairs, the bond angle is reduced by 2.5°. Vaia Originals
V-shaped
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, \(H_2O\) .
A v-shaped water molecule. Vaia Originals
The following diagram summarises the different shapes of molecules.
A table summarising the shapes of different molecules. Vaia Originals
Examples of the shapes of molecules
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, \(CO_2\) , has two double bonds. These double bonds can be thought of as single units when it comes to shape. Like single bond electron pairs, these groups of four electrons will want to be as far apart from each other as possible. This forms a linear molecule with a bond angle of 180°.
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, \(XeF_4\) . Xenon contains eight electrons in its valence shell. Four form bonds with fluorine atoms and four remain as two lone pairs. This forms what is known as a square planar arrangement, with the lone pairs at 180° to each other, and the angle between the bonding pairs at 90°. Note its similarity to an octahedral arrangement.
Xenon tetrafluoride. The lone pairs of electrons are positioned above and below the plane.Vaia Originals
Shapes of Molecules - Key takeaways
- VSEPR, also known as valence shell electron pair repulsion theory, states that electron pairs repel each other and will try to take up positions as far away from each other as possible, in order to minimise repulsion. This influences the shapes of molecules.
- You can use straight lines to represent covalent bonds. Wedged lines show a bond protruding out of the page and dashed or dotted lines show a bond extending backwards.
- Lone pairs of electrons repel each other more strongly than bonding pairs. Each lone pair reduces the bond angle by 2.5° in molecules with four electron pairs.
- Common molecule shapes with no lone pairs of electrons include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.
- Common molecule shapes with lone pairs of electrons include pyramidal and v-shaped.
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