Furniture building requires following instructions to ensure that all the pieces fit together correctly. Similarly, molecules have specific shapes that are determined by an instruction manual called VSEPR theory. VSEPR stands for valence shell electron pair repulsion theory, which is a set of rules used in chemistry to predict the shape of a molecule. By understanding the molecule's number and arrangement of valence electrons, chemists can determine its geometry. This article is all about VSEPR theory in chemistry. We'll start by explaining what it is and then dive into the different shapes of molecules that it creates. You'll also learn how to name and describe these shapes based on the molecule's valence electrons.
Molecules don't just randomly arrange themselves. Instead, they always take specific shapes, which we call their geometry. This geometry is determined by the number of lone and bonded pairs of electrons in the molecule. VSEPR theory is a model that neatly summarises these principles.
VSEPR theory is based on two key principles. Firstly, electron pairs repel each other. This means that electron pairs around a central atom will always try to be as far away from each other as possible. Secondly, lone pairs of electrons repel other electrons more than bonded pairs. This means that the presence of lone electron pairs can squash two bonded electron pairs closer to each other, changing the geometry of the molecule.
For this article, we'll be focusing on molecules with the general formula ABn, which are made up of identical atoms bonded to a central atom. This means that all of the bonds are the same. In molecules made up of different atoms bonded to a central atom, other factors like bond length, strength, and electron density come into play, making it harder to predict the geometry of the molecule. That's why we stick to molecules with identical bonds.
Now, let's dive deeper into these two principles of VSEPR theory.
The first principle of VSEPR theory is that electron pairs repel each other. All electron pairs are negatively charged, and as like charges repel each other, they try to stay as far apart as possible. This results in electron pairs within a molecule spacing out equally around the central atom. Molecules with the same number of electron pairs have the same basic shape, with certain angles between their bonds.
For instance, if a central atom has two pairs of electrons involved in single covalent bonds, the electron pairs will be as far apart as possible when they lie on opposite sides of the atom. This creates a linear molecule with an angle of 180° between the bonded pairs. In the next section, we'll look at the names of molecules with different shapes.
In addition to having the same basic shape, molecules with the same number of electron pairs will have slightly different shapes depending on the presence of lone pairs. Lone pairs of electrons repel other electron pairs more strongly than bonded pairs, so their presence can change the angle between bonded pairs and alter the molecule's shape.
For example, a molecule with four pairs of electrons around a central atom always has a tetrahedral shape. If all four pairs of electrons are bonded pairs, the angle between the bonds is approximately 109.5°. However, if a bonded pair is replaced with a lone pair, the angle between the remaining three bonds decreases to 107.0°, resulting in a trigonal pyramidal shape. If another bonded pair is replaced with a lone pair, the angle decreases further to 104.5°, creating a v-shaped molecule.
Double and triple covalent bonds contain two and three bonded electron pairs, respectively, which are treated as one "super pair" in VSEPR theory. Molecules with these "super pairs" have the same geometry as molecules with standard bonded pairs of electrons.
Now that we've covered the basics of VSEPR theory, let's move on to exploring the different shapes of molecules.
Moving on to molecules with three pairs of electrons, we have two possible shapes. When all three pairs of electrons are bonded pairs, the molecule has a trigonal planar shape, with an angle of 120° between the bonded pairs. If one of the bonded pairs is swapped for a lone pair, the molecule becomes a bent shape, with an angle of less than 120° between the remaining two bonded pairs.
Molecules with four pairs of electrons around the central atom always have a tetrahedral shape when all four pairs are bonded pairs. The angle between the bonded pairs is approximately 109.5°. If one of the bonded pairs is swapped for a lone pair, the molecule becomes a trigonal pyramidal shape, with an angle of less than 109.5° between the remaining three bonded pairs. If two of the bonded pairs are swapped for lone pairs, the molecule becomes a v-shaped shape, with an angle of less than 109.5° between the remaining two bonded pairs.
Moving on to molecules with five pairs of electrons, we have two possible shapes. When all five pairs of electrons are bonded pairs, the molecule has a trigonal bipyramidal shape, with angles of 120° and 90° between the bonded pairs. If one of the bonded pairs is swapped for a lone
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°.
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.
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.
For your exams, you only need to know that lone pairs of electrons reduce the bond angle in a molecule - you don't need to know the exact number of degrees the lone pair reduces the angle by.
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°.
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.
That's correct! In a molecule of phosphorus pentachloride (PCl5), the central phosphorus atom has five pairs of electrons around it, including five single bonds with chlorine atoms. The molecule has a trigonal bipyramidal shape, with three chlorine atoms arranged in a plane at 120° to each other, and the other two chlorine atoms located above and below this plane at 90° to the other three bonds.
It's important to note that molecules with five electron pairs can also have lone pairs. For example, if one of the bonded pairs in PCl5 is replaced with a lone pair, the molecule becomes a seesaw shape, with four chlorine atoms arranged in a distorted tetrahedral shape, and the fifth chlorine atom located above or below the plane of the other four. If two of the bonded pairs are replaced with lone pairs, the molecule becomes a T-shape, with three chlorine atoms arranged in a plane at 120° to each other, and the other two chlorine atoms located above and below the plane at an angle of less than 90° to the other three bonds.
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.
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.
VSEPR ChartBy 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.
VSEPR theory is a set of rules used in chemistry used to predict the geometry of a molecule. It is based on the molecule's number and arrangement of valence electrons. VSEPR theory is built on two key principles: Electron pairs repel each other. Because of this, they try to position themselves as far apart from each other as possible, giving molecules with the same number of electron pairs the same basic shape. Lone pairs of electrons repel other electrons more than bonded pairs. Because of this, they reduce the bond angle in molecules, changing the molecule's shape slightly. Molecules with two pairs of electrons are based on linear molecules with a bond angle of 180°.Molecules with three pairs of electrons are based on trigonal planar molecules with a bond angle of 120°.Those with four pairs of electrons are based on tetrahedral molecules with a bond angle of 109.5°.Molecules with five pairs of electrons are based on trigonal bipyramidal molecules. They have bond angles of 90° and 120°.Finally, molecules with six pairs of electrons are based on octahedral molecules with a bond angle of 90°.
What does VSEPR stand for?
VSEPR stands for valence shell electron pair repulsion. It is a theory used to predict the geometry of molecules.
What is VSEPR theory?
VSEPR theory is a model used to predict the geometry of molecules, such as their shape and bond angles.
What does VSEPR predict?
VSEPR predicts the shape of molecules, including their shape and bond angles.
How does VSEPR affect the shape of molecules?
VSEPR dictates that electron pairs repel each other. Because of that, they try to space themselves out as far away from each other as possible. This causes molecules to always take certain shapes depending on their number of electron pairs.
What are repelled in VSEPR theory?
Electron pairs are repelled in VSEPR.
