In this lesson, we will learn:
- To understand the ball-and-stick method of describing molecules in 3d space.
- To understand the principles of VSEPR theory.
- To apply VSEPR theory to predict the shapes and bond angles of different molecules.
- Molecules and compounds alike are groups of atoms held together by chemical bonds between them.
- One way to imagine these molecules is a ball and stick model – the atoms are spherical balls connected by sticks (the chemical bonds) to each other to form the molecule.
- Think about the following facts:
We can use these to predict the shape of molecules by considering the number of valence (outer shell) electrons the central atom has around it. This can determine both the shape and bond angles around the central atom. This method is known as Valence Shell Electron Pair Repulsion (VSEPR) theory.
- Molecules are made of atoms 'attached' together by covalent bonds.
- A covalent bond is made of a pair of electrons shared between two atoms.
- Electrons repel each other.
- Because electrons repel, electron pairs of any sort will try to place themselves as far away from each other as possible to reduce repulsion.
- There are two main factors that determine the geometry (shape) of a molecule:
- The number of bonding electron pairs around the central atom. These outer shell or valence electrons will try and site themselves as far away from one another as possible to reduce electron repulsion. (Note: double and triple bonds count as one bonding pair for VSEPR).
- For example, if there was just two pairs of bonding electrons around one atom, they would position themselves on opposite sides of the central atom. The angle between the bonds they make would be 180°, which we call the bond angle. The atoms of the molecule would more or less be in a straight line with each other, so the shape is called linear. See below for BeH2, a linear shaped compound with two bonding pairs of electrons.
- If there were three pairs of bonding electrons around the central atom they would divide the same 360o area in three equally spaced angles – 120° each. The three directions the bonds now point and place the atoms they're bonded to form the edges of a flat triangle, so the shape is called trigonal planar. See below for BF3, a trigonal planar molecule with three bonding pairs.
- The number of pairs of non-bonding ("lone pair") electrons around the central atom. Non-bonding pairs cause greater repulsion than bonding pairs because they are localized on one atom rather than 'sitting between' two atoms in a chemical bond. This means that bond angles in molecules with lone pairs are decreased – around 2.5° for each lone pair present.
- For example, in a compound with two bonding pairs of electrons and two non-bonding pairs on the central atom, the normal bond angle of 109.5° for 4 bonding pairs (see below) is reduced by (2 x 2.5) = 5° for the angle between the two bonding pairs. The drop from 4 to 2 bonding pairs also means the shape changes – there's two less atoms to make a shape with! Both lone pairs repel the bonding pairs down into a 'V' shape, sometimes called angular or bent. See below for the geometry of water, H2O, which is a V-shaped molecule with 2 bonding and 2 lone pairs.
- Molecular shapes and their particular bond angles can be predicted from finding the number of bonding and lone pairs around a central atom. Many shapes are summarized in this table: