The polar C—F bonds pull charge in opposite directions and cancel each other out: Figure 2.27d.
Figure 2.27 (a) Overall molecular dipole in fluoromethane; (b) overall molecular dipole in difluoromethane; (c) overall molecular dipole in trifluoromethane; (d) tetrafluoromethane with no overall dipole but polar bonds.
Worked Example 2.10
Determine if the molecule propan‐2‐one (acetone), C3H6O, is polar and has an overall dipole moment.
Solution
Once the structure of acetone has been determined, using a dot‐and‐cross diagram, it is possible to see that it is trigonal planar in shape; there are two single carbon‐to‐carbon bonds, and there is a double bond between oxygen and carbon. The carbon‐carbon and carbon‐hydrogen bonds are not polar, because carbon and hydrogen have similar electronegativities, but the carbon-oxygen bond is polar because oxygen is more electronegative than carbon. This molecule is also polar overall, as charge is pulled towards the oxygen atom, leaving the central carbon atom slightly positively charged.
2.4 Intermolecular forces
The bonds that hold atoms together in a molecule are known generally as intramolecular forces because they are inside a molecule. They are very strong and not easily broken. For example, in the molecule water, H2O, the covalent O—H bonds are the intramolecular forces. It is very difficult to split water into its component atoms, oxygen and hydrogen. In fact, it takes roughly 51.5 kJ energy to split 1 g of water into its elements, oxygen and hydrogen. If we could easily separate water molecules into gaseous hydrogen and oxygen, we would have a low‐cost way of producing hydrogen as a fuel stock.
However, simple covalent molecules such as water don't exist in isolation; they are surrounded by many millions of other similar molecules with much weaker forces between them. The forces that exist between molecules are called intermolecular forces. They are much weaker than the forces holding atoms together but are very important. The intermolecular forces holding water molecules together in liquid water are around 1.3 kJ per gram, which is about one‐twentieth the strength of an O—H bond. If it weren't for these intermolecular forces, we wouldn't have any liquid water to drink – water would exist as a gas on the earth's surface.
Intermolecular forces determine the properties of a covalently bonded compound, such as its melting and boiling point. It is important to understand how they arise in order to have an idea of the strength of these forces.
Figure 2.28 shows the strong intramolecular forces within the methane, CH4, molecule and the weaker intermolecular forces between the methane molecules.
Figure 2.28 Comparison of inter‐ and intramolecular forces.
It is important that you understand the difference between inter‐ and intramolecular forces, as it is subtle but very important. Intermolecular forces are between two different molecules (like international flights are between two different countries); intramolecular forces are between the atoms that form a bond and are within a molecule
There are three main types of intermolecular force between molecules:
Instantaneous dipole–induced dipole or London dispersion forces
Permanent dipole–permanent dipole
Hydrogen bonding
The first two types of intermolecular forces that involve dipole‐to‐dipole interactions (both permanent and instantaneous) are called van der Waals forces. Van der Waals forces are attractive forces between slightly positively and slightly negatively charged areas of a molecule. The term van der Waals is reasonably general and does not take into account the type of dipoles that are interacting. The term London dispersion forces is more specific, and this name is used for instantaneous dipole to induced dipole interactions. The third type of intermolecular force, hydrogen bonding, is a special type of dipole–dipole interaction.
The origins of these interactions will be discussed in the following sections.
2.4.1 Permanent dipole–permanent dipole interactions
Permanent dipole–permanent dipole interactions are stronger than London dispersion forces but weaker than hydrogen bonds. These interactions occur through space and are purely electrostatic with a δ+ charge in one molecule interacting with a δ− charge in another molecule. Figure 2.29 shows the permanent dipoles set up in HCl molecules, where the chlorine atom is more electronegative than the hydrogen atom. Permanent dipole–permanent dipole interactions are formed between the oppositely charged ends of molecules.
Figure 2.29 Permanent dipoles in the hydrogen chloride molecule and resultant permanent dipole–permanent dipole interactions.
2.4.2 London dispersion forces (instantaneous dipole–induced dipole)
An instantaneous dipole can occur in a bond between any two elements, regardless of the electronegativities of the bonded atoms. As the name suggests, they are fleeting and so do not last very long, but they can have an impact upon other molecules that are nearby. The common name for an instantaneous dipole to induced dipole interaction is London dispersion forces.
If the atoms in a bond have similar electronegativities, the electron charge is evenly distributed between them. However, because electrons are constantly moving, there is still a chance that at any one moment, the electrons may suddenly be at one end of the bond, rendering that end of the bond slightly negatively charged (δ−) and the other end of the bond slightly positively charged (δ+) in comparison. This forms an instantaneous dipole. Once an instantaneous dipole has been set up, it induces a dipole in another bond in a nearby molecule: Figure 2.30.
Figure 2.30 (a) Chlorine molecule with even distribution of