In the world of chemistry, we can find different shapes of molecules. For example, water and carbon dioxide are two triatomic molecules that share some similarities. Both of them are partially made of oxygen and contain covalent bonds. However, they have different shapes. Carbon dioxide has a straight line of atoms, while water is bent. But why is this the case?

To understand this, we need to talk about VSEPR, which stands for valence shell electron pair repulsion theory. This theory is used by chemists to predict the shape of molecules. Here's how it works: electrons like to move around in pairs, and because they are charged particles, they repel each other. So, they try to stay as far away from each other as possible. This is where the valence shell comes in - it's the outer shell of electrons in an atom. In simple covalent molecules, electron pair repulsion determines the position of the bonds, which in turn dictates the shape of the molecule. So, VSEPR states that electron pairs repel each other and try to take up positions as far apart as possible. This helps to minimise repulsion and keep the molecule stable. By understanding the behaviour of electrons, we can predict the shape of simple covalent compounds. If you want to refresh your memory on how atoms share electrons to achieve stable electron configurations, check out Covalent and Dative Bonding. In summary, VSEPR is an important theory in chemistry that helps us to understand the shapes of molecules. By predicting the positions of electron pairs, we can predict the shape of simple covalent compounds, such as water and carbon dioxide.

How do you draw the shapes of molecules in 3D?

Before we dive into examples of covalent structures, let's first learn about how we can represent them. You might already know that covalent bonds are typically represented as a line between two atoms. This gives us a simple picture of the molecule, but it doesn't show us the full 3D shape. To better represent a molecule's 3D structure, we can use wedged and dotted lines. Wedged lines represent a bond coming out of the screen or page towards you, while dotted or dashed lines represent a bond going into the screen or page away from you. Lone pairs of electrons are shown as dots, and any standard straight lines simply indicate a planar bond. Using these different line types can help us to better visualise a molecule's shape and structure. So, when you see a diagram of a covalent molecule, look out for the wedged and dotted lines to get a better idea of its 3D shape. Now that we know how to represent covalent structures, let's take a look at some examples in the next section.

The methane molecule is a good example of this:

The various shapes of molecules

When all the pairs of valence electrons in an atom are bonded, they will repel each other equally. This results in the bonds being spaced evenly 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, it's important to note 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, and we'll explore their effects later.

Some common shapes of molecules include:

• Linear: two bonded electron pairs, bond angle of 180 degrees
• Trigonal planar: three bonded electron pairs, bond angle of 120 degrees
• Tetrahedral: four bonded electron pairs, bond angle of 109.5 degrees
• Trigonal bipyramidal: five bonded electron pairs, bond angles of 90 and 120 degrees
• Octahedral: six bonded electron pairs, bond angle of 90 degrees

These shapes can help us to understand the arrangement of atoms in a molecule, which in turn affects its properties and behaviour. However, as mentioned earlier, lone pairs of electrons can also affect a molecule's shape and properties, so we'll explore those effects further in the next section.

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, . 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°.

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.

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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 () has four valence electrons, and each electron is part of a pair bonded covalently to a hydrogen atom. It is a tetrahedral molecule.

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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.

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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.

Lone pairs of electrons

When a molecule has a lone pair of electrons, it can affect the molecule's shape and bond angles. Let's take a molecule with four electron pairs as an example.

If all of the electrons are part of bonding pairs, the molecule will be tetrahedral with bond angles of 109.5°. However, if one of the electron pairs is a lone pair instead, the bond angles are reduced to 107°. This is because lone pairs repel each other more strongly than shared pairs, which can squeeze the bonds together.

Each lone electron pair in a molecule with eight valence electrons reduces the bond angle by 2.5°. So, for example, 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 different combinations of bonded and lone pairs of electrons:

• Bonded pairs of electrons: repel each other equally
• Bonded and lone pairs of electrons: lone pairs repel bonded pairs more strongly than bonded pairs repel each other
• Lone pairs of electrons: repel each other more strongly than bonded pairs

By taking into account the repulsion between bonded and lone pairs of electrons, we can better understand the shape and bond angles of molecules with lone pairs. This is important for understanding the properties and behaviour of these molecules.

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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, . 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.

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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, .

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The following diagram summarises the different shapes of molecules.

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Examples of the shapes of molecules

Carbon dioxide, CO2, has a linear shape. This is because it has two double bonds, which 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 results in a linear molecule with a bond angle of 180°.

So while water has a V-shaped structure due to the effect of its lone electron pairs on the bonding pairs, carbon dioxide has a linear structure due to its two double bonds. Understanding the shapes of molecules is important for understanding their properties and behaviour, and can be particularly useful in fields such as chemistry and biochemistry.

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Another example is xenon tetrafluoride, . 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.

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Yes, those are some of the key takeaways from our discussion on the shapes of molecules. It's important to understand the VSEPR theory and how it influences molecular shapes, as well as how to represent covalent bonds and lone pairs in diagrams. Knowing the common shapes of molecules, both with and without lone pairs, is also useful for understanding their properties and behaviour.

Shapes of Molecules

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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