Bonding and Elemental Properties

What are the Types of Bonding?

There are three main types of bonding found in chemistry:

• Covalent
• Ionic
• Metallic

We have articles specifically dedicated to each of these three types. However, we'll provide you with a quick recap now to get you up to speed, before we look at their properties.

Covalent Bonding

Atoms form bonds with each other in order to achieve a full outer shell of electrons. Non-metals typically do this by forming covalent bonds. These occur when the valence electron orbitals of two atoms overlap, forming a pair of electrons that both atoms share between them.

Atoms can form multiple covalent bonds, as well as double or even triple bonds. For example, ammonia contains a nitrogen atom bonded to three hydrogen atoms with single covalent bonds, whereas the cyanide ion contains a carbon atom bonded to a nitrogen atom using a triple covalent bond.

Multiple covalent bonds, shown in ammonia and the cyanide ion. Anna Brewer, StudySmarter Original

Ionic Bonding

Above, we saw that two non-metals bond by sharing electrons. In contrast, a metal and a non-metal bond by donating electrons. The metal gives up its outer shell electrons and transfers them over to the non-metal, which gains the electrons. This forms two ions which both have full outer shells of electrons. The ions are pulled together by strong electrostatic attraction, forming a giant ionic lattice that stretches in all directions. We call this attraction an ionic bond.

Sodium chloride is a great example of ionic bonding. In order to achieve a full outer shell, sodium needs to lose one electron. It does this by donating it to chlorine, which needs to gain one electron. This forms positive sodium ions and negative chloride ions, which are attracted to each other by strong electrostatic attraction.

Ionic bonding in sodium chloride. Anna Brewer, StudySmarter Original

Metallic Bonding

We now know how two non-metals bond, and how metal and non-metal bond. But how do two metals bond? As the name suggests, they use something called metallic bonding.

In order to achieve a full outer shell, each metal atom gives up its valence electrons and becomes a positive metal ion. The electrons form a sea of delocalization that surrounds the metal ions, and the whole structure is held together by electrostatic attraction between the positive ions and negative electrons.

A typical example is sodium. It has one electron in its outer shell. When sodium atoms bond, each atom loses its outer shell electron to form a positive sodium ion. The electrons form a sea of delocalization that surrounds these metal ions.

Metallic bonding in sodium. Anna Brewer, StudySmarter Original

Bonding and Properties

We've seen how metals and non-metals bond both with themselves and with each other. Let's now turn our attention to how this bonding affects their properties.

Properties of Covalent Bonding

Structures made from covalent bonds can either form covalent-network solids or simple covalent molecules. In covalent-network solids, hundreds upon thousands of atoms are joined together by multiple covalent bonds, forming a giant lattice that stretches in all directions. They don't form molecules.

Simple covalent molecules, on the other hand, are molecules that consist of just a handful of atoms bound together by covalent bonds. The molecules are held together by weak intermolecular forces. In a solid-state, we call them molecular solids. These two types of structures feature the same type of bonding but have different properties. Why is this the case?

Covalent bonds are very strong. They require a lot of energy to overcome. Because of this, covalent network solids have very high melting and boiling points - in order to melt the structure, you need to overcome the covalent bonds between atoms. But although simple covalent molecules also contain covalent bonds, they have low melting and boiling points and are usually gaseous at room temperature. This is because you don't need to overcome the covalent bonds within the molecule in order to melt the substance, but rather overcome the weak intermolecular forces between the molecules.

Similarly, covalent network solids are hard and strong, thanks to the strong covalent bonds holding them together. Furthermore, both covalent network solids and simple covalent molecules are usually poor conductors of electricity. This is because there aren't any charged particles free to move and carry a charge within the structures. Covalent network solids are also insoluble in water.

Properties of Ionic Bonding

Although covalent bonding and ionic bonding exist on the same spectrum, they show very different properties. Let's look at the properties of ionic compounds now.

First of all, ionic structures form giant lattices, not molecules. The ions in an ionic compound are attracted to all of the oppositely charged ions around them, forming an ionic lattice of alternating positive and negative ions that stretches in all directions. This actually creates a crystal structure. The electrostatic attraction is strong, giving ionic compounds high melting and boiling points.

Ionic compounds are also hard and brittle. If you jar them, you knock the lattice out of place; two ions with the same charge might suddenly find themselves directly next to each other. They repel each other, breaking apart the lattice.

As solids, ionic compounds are poor conductors of heat and electricity. The charged ions are held firmly in place by the strong electrostatic attraction and can't move around. But when molten or aqueous, they aren't held firmly in place and can move around, carrying charges. Because of this, molten and aqueous ionic compounds are good conductors of electricity. Luckily for us, most ionic compounds are highly soluble in water. Some of them are even soluble in organic liquids.

Properties of Metallic Bonding

Next up, let's focus on the properties of metallic bonding.

As with ionic compounds, metals form giant lattices. We tend to talk about them as an array of positive metal ions in a sea of delocalized electrons, held together by strong electrostatic attraction. Because of this strong attraction, metals have medium to high melting and boiling points.

However, unlike ionic compounds, metals are often malleable and ductile. The positive metal irons are arranged in rows, cushioned by the sea of electrons, and can slide over each other. This means that metals can be hammered into shape and drawn into wires.

The sea of delocalized electrons also helps metals conduct. The electrons are free to move around the array of positive metal ions, so they can carry a charge. Metals are therefore good conductors of heat and electricity. They're also insoluble in water.

Finally, the bonding of metals contributes to their shiny, lustrous appearance. Shining a light on a metal excites some of the delocalized electrons on the outside of the structure. To get back to their ground state, they then release energy as light, giving off a lustrous gleam.

Comparing Bonding and Properties

Looking at a species' properties is often a useful indication of its type of bonding. This is particularly handy when looking at species that stray from the familiar trends.

For example, beryllium chloride, BeCl₂, consists of beryllium and chlorine atoms. Beryllium is a metal, and chlorine is a non-metal. Because of that, you might expect them to bond ionically. However, they actually bond covalently, forming a simple covalent molecule. We can infer this when we look at the molecule's properties: beryllium chloride has low melting and boiling points.

Here's a table that compares the different types of bonding, and the structures and properties associated with them.

Periodicity of Bonding

If we knew an element's position on the periodic table, could we predict how it bonds? In fact, we often can. Bonding shows periodicity.

Here's the periodic table.

The periodic table. Image credit: commons.wikimedia.org

We know that metals, found on the left-hand side of the periodic table, bond using metallic bonding. As you move across to the right-hand side of the periodic table, you encounter the non-metals. We know that these tend to bond using covalent bonding. Moving down a column in the periodic table, known as a group, you encounter elements with very similar properties. This happens because the elements in the same group bond in similar ways. In fact, electron configuration and location on the periodic table are good indicators of an element's bonding.

That's it! By now, you should be able to explain the differences between the three types of bonding, and compare the properties they give a species.