Types of Bonding

Ionic, covalent and metallic bonding are types of intramolecular bond and are very strong. In each type of bond, you have an attraction between something positive and something negative - this is called an electrostatic attraction.

Chemical bonds

There are three types of strong chemical bonds: ionic, covalent and metallic. These bonds are found between atoms in a compound or a molecule so are known as intramolecular bonds. You need to be able to define each of these types of bonding.

1. Ionic bonding – this is the electrostatic interaction between oppositely charged ions. It forms between metals (which lose electrons to form positive ions) and non-metals (which gain electrons to form negative ions.

2. Covalent bonding – this is when atoms share electrons. A covalent bond is the electrostatic attraction between the positive nuclei and the shared pair of electrons. It occurs in most non-metallic elements and in compounds of non-metals.

3. Metallic bonding – this is the electrostatic attraction between the positive metal ions and the delocalised electrons. Metallic bonding occurs within metals and alloys, such as copper, iron or steel.

Ionic bonding

Atoms are always trying to get a complete set of electrons in their outer energy levels. If they are metallic elements located in the first few groups of the periodic table and only have a couple of electrons, they will happily give away these electrons to form positive ions. If it is a non-metal, with six or seven electrons in their outermost energy level, they will grab electrons from another atom to get a complete octet. Gaining electrons turns the atom into a negative ion. An ionic bond is the electrostatic attraction between positive and negative ions.

We can show how the metallic atom loses electrons and the non-metallic atom gains electrons by drawing a dot and cross diagram as shown below. Typically we only show the electrons present in the outer energy level and we use different symbols for the electrons on each atom, so we can clearly see the electron transfer.

This dot and cross diagram shows how sodium, which has one electron in its outer shell, will lose this outer electron to become an ion with a +1 charge. An atom of chlorine takes the electron from sodium, so that it now has a full outer shell of eight electrons. Since the chloride atom has gained an additional electron, its ion has a -1 charge.

You need to know the charges of certain ions which are easy to predict depending on which group they are in. If they are in group 1, they will lose this single outer electron to form a 1+ ion. However, if they are in group 6, they need two extra electrons to achieve a full octet so will form a 2- ion. The block of elements in the middle (the transition metals) are a little trickier and you may need to memorise these.

Ionic compounds

Ionic compounds can form giant lattices, with a large number of positive and negative ions forming numerous ionic bonds. These structures have high melting and boiling points because ionic bonds are strong and require a great deal of energy to break. They do not conduct electricity when solid, but can conduct electricity when molten or as part of an aqueous solution.

To conduct electricity, there needs to be charged molecules which are free to move. The positive and negative ions can only move when the ionic compound is melted or dissolved in solution.

Covalent bonding

A covalent bond is formed when two atoms share electrons. When this happens there is an electrostatic attraction between the nuclei (which are positively charged) and the pair of shared electrons (which are negatively charged). In the dot and cross diagram below, you can see how hydrogen forms a covalent bond by sharing its single electron with chlorine to form hydrogen chloride.

Simple covalent molecules have low melting and boiling points because the strong covalent bonds remain intact when these molecules change state. Instead, it is the weak intermolecular forces which need to be broken when converting a simple covalent molecule between a solid, liquid and gaseous state.

Although all simple covalent molecules will have low melting and boiling points, those with a large molecular mass will have a higher melting and boiling point than a molecule with a smaller mass. This is because the weak intermolecular forces that are found between molecules form from electrons. Heavier molecules have a larger number of electrons which means they can form more intermolecular forces, thereby increasing their melting and boiling point.

Giant covalent structures, such as diamond, graphite and silicon dioxide (silica) are made up of a large number of atoms covalently bonded in a huge network. They have very high melting and boiling points because the strong covalent bonds need to break when these molecules change states.

Covalent molecules are not made up of charged particles, so neither simple nor giant covalent structures are able to conduct electricity. An exception is graphite, which has delocalised electrons between layers of carbon atoms.

Although diamond and graphite are both giant covalent structures made up of only carbon atoms, they have very different properties due to the way that the carbon atoms are arranged. In diamond, each carbon atom bonds to four other carbon atoms to form a strong, rigid structure which makes diamonds one of the hardest materials on Earth and virtually unbreakable. In contrast, each carbon atom in graphite forms only three other bonds. As the fourth electron in carbon’s outer electron shell is unbonded, it is free to move (delocalised) and can conduct electricity. The electron moves between layers of carbon atoms which are held together with weak intermolecular forces which are easily broken. This makes graphite soft and slippery, useful as a lubricant or in pencils, since the layers can separate and slide over each other easily.

Metallic bonding

Metals consist of atoms arranged in neat rows or layers, stacked on top of one another. The outer electrons from each atom are able to abandon the nucleus and move freely throughout the metal - when this happens we say that the electrons are delocalised. Without the outer electrons spinning around the nucleus, each atom is now a positive ion. A metallic bond is the electrostatic attraction between a positive metal ion and the delocalised electrons.

The structure of a metal explains two important properties:

1. Good conductors of electricity - remember that anything that is charged and free to move can conduct electricity. The delocalised electrons fit this criteria and it’s this that makes metals such good conductors.

2. Malleable - a metallic lattice consists of layers that can slide over each other. This allows metals to be hammered or rolled into flat sheets without breaking.