Metallic Bonding

You probably have sodium chloride, NaCl, in your house right now. You’ll know it as table salt. Although the crystals found in your salt shaker are generally small, you can grow larger crystals at home simply by heating up a saturated salt solution and allowing it to cool. These crystals will be hard and brittle. Pure sodium metal, on the other hand, is soft and malleable. You can easily cut it with a butter knife. What makes sodium so physically different when it is on its own?

What is metallic bonding?

Metals can form compounds with non-metals by donating their outer shell electrons (see Ionic Bonding for more information). The metals form positive ions whilst the non-metals, which accept the electrons, form negative ions. However, if a metal is on its own it cannot donate electrons because there is no non-metal atom that can accept them. Instead, it does something else: it bonds metallically.

Let’s break that down and explore each term together.

Delocalised electrons

When metal atoms bond with one another, their outer shell electron orbitals merge. The electrons are no longer confined to one particular atom and are free to move within the merged orbitals, which form a region that stretches throughout the whole metal. We say that these electrons are delocalised and that they form a sea of delocalisation.

Electrostatic attraction

The metal atoms form ions with a positive charge, since they are now no longer associated with their outer shell electrons.

The positive ions are then attracted to the negative sea of electrons by electrostatic attraction, much like the attraction in ionic compounds. The attraction spreads throughout the whole metal and so forms a giant lattice structure. 'Giant' simply means that it is made from a large but indeterminate number of atoms, and 'lattice' means that it contains a repeating arrangement.

Although the metal contains positive ions, overall no electrons have been lost. They are simply delocalised within the metal’s structure. Therefore, metals have a neutral charge and we represent them using just their chemical symbol. For example, the molecular formula of sodium is Na.

Fig. 1 - A diagram showing the bonding in a metal

Let’s go back to our example of sodium, Na. Sodium has the electron configuration . When sodium atoms bond with each other, their 3s orbitals merge and the valence electron within each atom’s orbital is free to move about in the newly merged region. This leaves positive ions with a charge of +1 surrounded by a sea of delocalised electrons, as shown below.

Fig. 2 - The bonding in sodium, Na. Each sodium ion is attracted to the sea of delocalisation around it by electrostatic attraction

Beryllium, on the other hand, has the electron configuration and has two valence electrons. Each beryllium atom loses two electrons from its outer shell to form ions with a charge of +2.

Fig. 3 - The bonding in beryllium

Factors affecting the strength of metallic bonding

Some metals are much stronger than others. This is because of the difference in levels of electrostatic attraction within the different metals. There are two factors that affect the strength of the metal bond, and we’ll explore them now.

A positive ion with a higher charge will be more attracted to the negative sea of electrons than one with a lower charge. Remember that a metal bond is simply the electrostatic attraction between positive metal ions and the sea of delocalised electrons, so this creates stronger bonding.

Aluminium, for example, loses three valence electrons to form an ion with a charge of +3. However, magnesium only loses two electrons to form an ion with a charge of +2, so has much weaker metallic bonds.

Size of ion

In metals with larger ions, the positive nucleus is a lot further away from the delocalised electrons. This weakens the electrostatic attraction between them. For example, the positive ions in magnesium and calcium both have the same charge, but calcium contains much larger ions and so has weaker metallic bonds.

Properties of metals

Because of their unique arrangement of positive ions within a sea of delocalised electrons, metals have certain properties that differentiate them from ionic and covalent compounds. We use copper, for example, to make wires and pipes. We wouldn’t use ionic compounds like sodium chloride for this. They would dissolve if dampened, and wouldn’t conduct electricity when solid. Furthermore, ionic compounds are brittle and break easily if stressed.

However, metals are quite different.

• They have high melting and boiling points. This is because of the strength of their electrostatic attraction which stretches throughout the entire metal. Any of the factors explored above that increase the strength of the metallic bonding increase the melting and boiling points of a metal.
• They are ductile, meaning they can be stretched out into wires, and malleable, meaning they can be hammered into shape. This is because the positive ions form regular rows within the sea of electrons which can roll over each other smoothly.
• They are not brittle and generally strong. Again, this is because the rows of metal ions still maintain their bonds with the delocalised electrons when they slide over each other.
• They are good conductors of heat and electricity, as the delocalised electrons are free to move throughout the metal and carry a charge. Metals that form ions with higher charges have more delocalised electrons and so are better conductors than metals with lower-charged ions.
• They are insoluble.

Alloys

We know that sodium is relatively soft. Pure iron is too. This causes problems when making useful products out of metals. Iron nails wouldn’t be of much use if you could easily bend and distort them. To make pure metals stronger, we turn them into alloys.

The different-sized atoms of the second element in an alloy disrupt the regular rows of metal ions, preventing them from sliding over each other as much, thus making them much harder. Iron often contains carefully controlled amounts of carbon, and steel is also a common alloy made from iron.

Fig. 4 - The atoms in an alloy. Here, the smaller atoms disrupt the regular lattice structure of the larger metal atoms and prevent them sliding over each other. This strengthens the compound