Metallic Bonding

Metallic Bonding

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Sodium is a metal with a unique property. It has a strong metallic bond that gives it a high melting and boiling point. This makes it hard and brittle when in crystal form, like table salt, but soft and malleable when in its pure form. This strong bond is caused by the delocalized electrons that surround the cations, making the effective nuclear charge on the electrons higher and the cation size smaller. This bond is so strong that it requires a lot of energy to break, which is why sodium has such a high melting and boiling point.

What is metallic bonding?

Metals are able to bond with non-metals by giving away their outer shell electrons, which results in positive metal ions and negative non-metal ions. But when a metal is by itself, it can't give away electrons because there is no non-metal atom to accept them. In this case, it forms a metallic bond.

A metallic bond happens when there is an attraction between positive metal ions that are arranged in a lattice structure and a sea of electrons that are free to move around within the lattice. These electrons are delocalised, which means they're not attached to any specific metal ion. Instead, they move around the lattice, creating a type of glue that holds the metal ions together. A metallic bond is responsible for many of the physical properties of metals, such as their high melting and boiling points, their ability to conduct electricity and their malleability.

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

When metals bond metallically, they give up their outer shell electrons and become positively charged ions, which are called cations. These cations are then attracted to a sea of negatively charged electrons by electrostatic attraction, much like in ionic compounds.

This attraction spreads throughout the entire metal and forms a giant lattice structure. The lattice structure is made up of a large but indeterminate number of atoms that are arranged in a repeating pattern.

Despite the formation of positive ions, no electrons are actually lost. They are still present in the metal’s structure, but they're now delocalised. This means they're free to move around within the metal's lattice structure. As a result, metals have a neutral charge and are represented by their chemical symbol alone, such as Na for sodium.

Metallic bond Cu
Metallic bond Cu

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.

Each sodium ion is attracted to the sea of delocalisation around it by electrostatic attraction
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.

The bonding in beryllium
The bonding in beryllium

Factors affecting the strength of metallic bonding

The strength of a metal bond is determined by two main factors: the number of protons and the number of delocalized electrons per atom. The more protons, the stronger the bond, and the more delocalized electrons, the stronger the bond. Additionally, the size of the ion also affects the strength of the bond; the smaller the ion, the stronger the bond. For example, aluminium has a higher charge than magnesium because it loses three valence electrons to form an ion with a charge of +3, whereas magnesium only loses two electrons to form an ion with a charge of +2, resulting in 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

Metals have unique properties that set them apart from ionic and covalent compounds. For instance, copper is used to make wires and pipes because of its ductile and malleable properties. It can be easily stretched out into wires or hammered into shape. On the other hand, ionic compounds, like sodium chloride, are brittle and easily breakable under stress.

Metals also have high melting and boiling points. This is due to the strength of their electrostatic attraction, which extends throughout the entire metal. The stronger the metallic bonding, the higher the melting and boiling points of the metal. Metals are good conductors of heat and electricity. This is because the delocalised electrons are free to move throughout the metal, carrying a charge. Metals that form ions with higher charges have more delocalised electrons, making them better conductors than metals with lower-charged ions. Finally, metals are insoluble, meaning they do not dissolve in water or other solvents.


To make pure metals stronger, we turn them into alloys. Alloys are mixtures of two or more elements, at least one of which is a metal. The addition of a second element disrupts the regular rows of metal ions, preventing them from sliding over each other as much and making them much harder. Iron often contains carefully controlled amounts of carbon, and steel is a common alloy made from iron.

In summary, metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons. The strength of metallic bonding is affected by factors such as the charge and size of the ion. Metals are generally strong, not brittle, good conductors of heat and electricity, insoluble, and have high melting and boiling points. Alloys are stronger than pure metals due to the disruption of the regular rows of metal ions by the addition of a second element.

Metallic Bonding

What is metallic bonding?

Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons.

Do metals conduct electricity?

Metals can conduct electricity because the sea of delocalised electrons found within the metal are free to move and carry a charge.

How are metal bonds formed?

 Metals form bonds by merging their outer shell electron orbitals. The electrons within delocalise and are not attached to any particular metal atom. This forms positive metal ions within a sea of delocalised electrons. A metallic bond is simply the electrostatic attraction between the two.

Why are metallic bonds so strong?

Metallic bonds are strong because the strong electrostatic attraction between the positive metal ions and the negative sea of delocalised electrons extends throughout the entire metal.

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