VSEPR

What is VSEPR theory?

Molecules aren't arranged randomly. In fact, they always take specific shapes. We call the shape of a molecule its geometry, and geometry depends on a molecule's valence electrons. More specifically, it depends on the number of lone and bonded pairs of electrons. This can be neatly summarised into a handy model known as VSEPR theory.

VSEPR is based on two key principles.

1. Electron pairs repel each other. Because of this, electron pairs around a central atom will try and take up positions as far away from each other as possible.
2. Lone pairs of electrons repel other electrons more than bonded pairs. Because of this, the presence of lone electron pairs will squash two bonded electron pairs closer to each other, changing the geometry of the molecule.

Let's look at the two above ideas in more detail.

Electron Pair Repulsion

Firstly, electron pairs repel each other. All electron pairs are negatively charged, and like charges don't get on - they always repel each other. As a result, electron pairs around a central atom try and stay as far away from each other as possible. This generally involves spacing out equally around the central atom. Because of this, molecules with certain numbers of electron pairs have a certain shape and certain angles between their bonds. In fact, all molecules with the same number of electron pairs have the same basic shape.

For example, say a central atom has just two pairs of electrons, both involved in single covalent bonds. The electron pairs are as far away from each other as possible when they lie on opposite sides of the atom. This results in a linear molecule, with an angle between the bonded pairs of 180°. Don't worry - we'll look at the names of differently shaped molecules in just a second.

Lone Pairs and Bonded Pairs

We now know that molecules with the same number of electron pairs all have the same basic shape. But when it comes to geometry and electron pair repulsion, not all electron pairs are equal. This alters the shape of the molecule ever so slightly. This is because lone pairs of electrons repel other electron pairs much more strongly than bonded pairs. If any lone pairs are present, they squeeze the bonded pairs closer together. This decreases the angle between the bonded pairs and changes the molecule's shape.

Examples of molecules with lone and bonded pairs of electrons. Source: Expii

For example, a molecule with four pairs of electrons around a central atom is always based on a tetrahedral shape. If all four pairs of electrons are bonded pairs, the angle between the bonds is approximately 109.5°. But if you swap one of the bonded pairs for a lone pair, the angle between the three remaining bonds decreases down to 107.0°, and the shape changes slightly, becoming trigonal pyramidal. Swap another bonded pair out, and the angle decreases down to 104.5°, making the molecule v-shaped.

Now that we've looked at the fundamentals of VSEPR theory, let's move on to the shapes of the molecules themselves.

VSEPR Shapes of Molecules and Geometry

We've learned that VSEPR uses the number and arrangement of valence electrons to predict the geometry of a molecule. We'll now focus on the different shapes caused by varying numbers of pairs of electrons, starting with molecules with just two pairs and working up to those with six. We'll start with the basic shape of each molecule, which occurs when all of the pairs of electrons are bonded pairs, before exploring the effect of swapping some of them for lone pairs.

Linear

Molecules with just two pairs of electrons have a linear shape. The two electron pairs, whether bonded or lone, position themselves as far away from each other as possible. This means ending up directly opposite each other. The angle between the two bonds is therefore 180°.

Two examples of linear molecules are beryllium chloride (BeCl₂) and carbon dioxide (CO₂). They consist of two atoms joined to a central atom by single or double covalent bonds. In both cases, the bond angle is 180°.

Beryllium chloride, a linear molecule. Anna Brewer, StudySmarter Original

Carbon dioxide, a linear molecule. Anna Brewer, StudySmarter Original

Trigonal Planar

Molecules with three bonded pairs of electrons have a trigonal planar shape. To picture this shape, imagine an equilateral triangle with the molecule's central atom directly in the middle. The three pairs of electrons point out towards the triangle's three corners. If all of the electron pairs are bonded pairs, this makes the angle between them 120°.

One example of a trigonal planar molecule is boron trifluoride, BF₃.

Boron trifluoride, a trigonal planar molecule. Anna Brewer, StudySmarter Original

Earlier on, we learned that lone electron pairs repel other electrons more strongly than bonded electron pairs. If we swap one of the bonded pairs of electrons in a trigonal planar molecule for a lone pair, the remaining two bonds get squeezed more closely together, reducing the bond angle to slightly less than 120°. This forms a version of a trigonal planar molecule called a bent molecule. An example of a bent molecule is sulfur dioxide, SO₂.

Sulfur dioxide, a bent molecule. Anna Brewer, StudySmarter Original

Tetrahedral

Molecules with four pairs of electrons have a tetrahedral basic shape, i.e. you now have to start thinking in 3D. If the molecule has four bonded pairs of electrons and no lone pairs, the angle between each of the bonds is 109.5°.

For example, methane, CH₄, consists of four hydrogen atoms joined to a central carbon atom by single covalent bonds. It is a tetrahedral molecule with bond angles of 109.5°.

Methane, a tetrahedral molecule. Created using images from commons.wikimedia.org

But like with trigonal planar molecules, this geometry changes slightly as we swap some of the bonded pairs for lone pairs:

• Swapping one bonded pair for a lone pair decreases the remaining bond angles slightly and forms a trigonal pyramidal molecule, such as ammonia.
• Swapping a second bonded pair for a lone pair decreases the one remaining bond angle even more and forms a v-shaped molecule, such as water.

Ammonia, and trigonal pyramidal molecule. Based on images from commons.wikimedia.org

Trigonal Bipyramidal

Molecules with five pairs of electrons are based on a trigonal bipyramidal shape. Like before, we'll start by looking at the shape formed when all of the electron pairs are bonded pairs. Three of the bonded pairs arrange themselves in a similar way to a trigonal planar molecule: they spread themselves out equally across a plane at 120° to each other. The other two bonded pairs arrange themselves directly above and below this plane, at 90° to the other three bonds.

Phosphorus pentachloride is a good example of a trigonal bipyramidal molecule. It contains five chlorine atoms joined to a central phosphorous atom by single covalent bonds.

Phosphorus pentachloride, a trigonal bipyramidal molecule. Based on images from commons.wikimedia.org

Once again, swapping some of the bonded pairs of electrons for lone pairs changes the shape of the molecule. it also changes the remaining bond angles.

• A molecule with four bonded pairs and one lone pair forms a see-saw molecule. An example is sulfur tetrafluoride.
• A molecule with three bonded pairs and two lone pairs forms a T-shaped molecule. One example is chlorine trifluoride.
• A molecule with just two bonded pairs and three lone pairs forms another type of linear molecule. An example is xenon difluoride.

Octahedral

Finally, let's look at molecules with six pairs of electrons. Their basic shape is octahedral. To picture an octahedral molecule with six bonded pairs, imagine that the central atom is placed directly in the middle of a square. Four of the bonds point towards the corners of the square; these are all at 90° to each other. The other two bonded electron pairs are found directly above and below the plane. This means that these bonds are also at 90° to all of the others. Sulfur hexafluoride is a common example of an octahedral molecule. All of the angles between its S-F single bonds are 90°.

Sulfur hexafluoride, an octahedral molecule. Based on images from commons.wikimedia.org

Swapping out some of the bonded pairs of electrons for lone pairs changes the geometry of this molecule and reduces the angle between the remaining bonds.

• Replacing one bonded pair with a lone pair creates a square pyramidal molecule, such as bromine pentafluoride.
• Replacing two bonded pairs with two lone pairs creates a square planar molecule, such as xenon tetrafluoride.

Xenon tetrafluoride, a square planar molecule. Based on images from commons.wikimedia.org

VSEPR Chart

By now, you should be familiar with the shapes of different molecules as dictated by VSEPR theory. To help you consolidate your knowledge, we've made a handy chart comparing the basic shapes of molecules, their numbers of bonded pairs of electrons, and their bond angles. We've also included the names of the shapes of variants of these molecules, that occur when you swap out some of the bonded pairs of electrons for lone pairs.

A VSEPR theory chart. Anna Brewer, StudySmarter Original

That's it for this article. You should now know what VSEPR theory is, and be able to use it to name, identify, and draw the shapes of molecules between two and six bonded pairs of electrons. You should also be able to explain the effect of lone pairs of electrons on molecule geometry.