Bonding

Imagine a world without bonding - without chemical bonding, to be precise. Atoms ignore each other and instead just drift around in empty space, minding their own business. The water in the oceans splits into hydrogen and oxygen gases, which expand to 1500 times their liquid volume. Chlorine atoms from salts dissolved in the water form a toxic cloud, poisoning all of life with deadly fumes. Not that there is any life, or water, or in fact much of anything at all.

All the reactions that take place and all the molecules that we find in life only exist because of bonding. The only elements we commonly find without bonds are noble gases.

There are two categories of bonding, known as primary and secondary bonding. Primary bonding is what most people think of when the word is mentioned. Primary bonds are also known as intramolecular forces. They take place between atoms within the molecule. They’re generally strong and hard to break.

In contrast, secondary bonds are a lot weaker. They are more commonly known as Intermolecular Forces, since they take place between molecules. When we mention bonding here, we mean primary bonding, unless otherwise stated.

Why do bonds form?

Atoms want to be stable - they like being in the lowest energy state possible. By combining with other atoms in various combinations, they can form different substances with much lower energy states.

Stability depends on how many electrons an atom has in its outer shell. To be at its most stable, an atom wants to have a full outer shell of electrons, much like that of a noble gas. This is why noble gases don’t readily form bonds with other atoms. They are already as stable as they can be! Instead, we find them as monatomic gases.

Atoms bonded together form molecules or compounds.

Types of bonding

We mentioned above that bonding occurs because of attraction between positive and negative charges. You should know from Atomic Structure and Fundamental Particles that atoms are made up of protons, neutrons, and electrons. The table below summarises their charges and locations within the atom.

Fig. 1 - A table comparing protons, neutrons and electrons

You can see that protons have a positive charge and electrons have a negative charge. These are the only charged particles within an atom. All attraction, and therefore all bonding, must be between protons and electrons.

We know that atoms try to form bonds in order to achieve a full outer shell of electrons. They do this by moving their electrons around between each other. They can do this in three different ways, which result in three different types of bonding:

• Sharing electrons results in covalent bonding.
• Donating electrons results in ionic bonding.
• Delocalising electrons results in metallic bonding.

Covalent bonding

Covalently bonded atoms share electrons with each other so that they all have full outer shells of electrons.

Only non-metals form covalent bonds. Electron orbitals from two different atoms overlap, and a shared electron pair is formed using an electron from each atom. The bond is held together by the attraction between the negative shared electron pair and the positive nuclei inside the two atoms.

Fig. 2 - A diagram showing covalent bonding. Each bond contains one electron from carbon and one electron from hydrogen

Ionic bonding

Ionic bonding occurs between metals and non-metals. Ionic bonds are formed when a metal atom donates electrons to a non-metal. This forms charged atoms known as ions, which are attracted to each other. An ionic bond is an electrostatic attraction between oppositely charged ions.

Fig. 3 - A diagram showing ionic bonding. Each sodium atom donates an electron to a chlorine atom, forming positive sodium ions and negative chloride ions

Metallic bonding

For a metal on its own to have a full outer electron shell, it must do something quite different. Its outer electrons delocalise and the metal forms positive metal ions. Unlike in ionic bonding, where the electrons are picked up by another atom, in metallic bonding the electrons float about freely within the structure. The attraction between the negative electrons and the positive metal ions holds the metal together.

Fig. 4 - Metallic bonding in sodium. Each sodium atom loses an electron to form a positive ion. The electrons are delocalised and move about within the structure

Intermolecular forces

As you now know, substances made from two or more atoms joined together by chemical bonds are called molecules. For example, two hydrogen atoms and one oxygen atom bond to form a water molecule. We know that bonds are found within molecules. What forces are found between molecules?

The answer is Intermolecular Forces, which can also be known as secondary bonds. There are three different types:

• Van der Waals forces.
• Permanent dipole-dipole forces.
• Hydrogen bonding.

Van der Waals forces

Van der Waals forces are the weakest type of intermolecular force. They occur between all molecules. The random movement of electrons within a molecule causes a temporary dipole, which induces a dipole in a neighbouring molecule. The attraction between the two dipoles holds the molecules together.

Permanent dipole-dipole forces

In some molecules, the electrons are permanently distributed in an uneven manner. This means that one side of the molecule is constantly more negative than the other, and we call this a permanent dipole. Oppositely charged dipoles attract each other. The forces are known as permanent dipole-dipole forces and are stronger than van der Waals forces.

Hydrogen bonding

Some molecules containing hydrogen atoms experience an even stronger type of intermolecular force that we call hydrogen bonding. It occurs between molecules that have a hydrogen atom bonded to an oxygen, nitrogen or fluorine atom.

The following diagram ranks the different types of primary and secondary bonds, also known respectively as intramolecular and intermolecular forces, according to their relative strength.

Fig. 5 - A diagram showing the relative strengths of intermolecular and intramolecular forces

Bonding and structure

You can probably guess just by looking at everyday objects around you that different types of bonding produce very different types of structures. Take a diamond ring, for example. The metal that makes up the ring is easily melted or beaten into its circular torus shape, but the diamond embedded in the centre is extremely hard and strong. In fact, diamond doesn’t melt at all under normal atmospheric conditions. If you heat it to extremely high temperatures, it simply sublime, i.e., turns straight into a gas.

This is because metals bond using metallic bonds whilst diamond uses covalent bonds. This gives the two substances very different structures and properties. However, oxygen is also a covalent molecule, but it behaves completely differently from a diamond! Look at their states of matter, for example. We just learnt that a diamond needs extreme temperatures before it sublimes, but oxygen, by contrast, is a gas at room temperature. We can therefore deduce that it isn’t just the type of bonding that affects a molecule’s properties, but also the structure and arrangement of atoms and the way the bonds hold the molecule together. The following table gives a summary of the different types of structures we find in chemistry.

Fig. 6 - A table comparing different structures caused by bonding

Bonding and shape

We learnt above about how metals and ionic substances form lattices. Some covalent substances do too.

For example, the ionic lattice sodium chloride alternates positive sodium ions and negative chloride ions. However, a simple covalent molecule doesn’t have a lattice structure. Instead, it forms a molecule with a specific shape, depending on the number of electron pairs and covalent bonds it contains.

Molecule shape is dictated by electron pairs. Imagine two magnets. If you hold the two south poles near each other, they will try to separate. This is because similar charges repel each other. Electron pairs are much the same. Put a group of electron pairs together in the shell of an atom and they’ll repel each other, trying to spread out as far apart as possible. If all the electron pairs are part of covalent bonds, the bonds will be equally spaced apart. But electron pairs that aren’t part of a bond, known as lone pairs, have a stronger repulsive force than bonded pairs. They repel other electron pairs more and squash the bonded pairs closer together.

An example is methane, or \(CH_4\) . It has four electron pairs in its outer shell. They are all bonded pairs and repel each other equally. The angle between each of the bonded pairs is 109.5°. Water \(H_2O\) also has four electron pairs in its outer shell. However, two of the pairs aren’t bonded - they are lone electron pairs. This squashes the bond angle to just 104.5°.

The following table summarises the shapes of different covalent molecules. There are also diagrams to help you consolidate your knowledge.

Fig. 7 - A table comparing the different shapes of molecules