Physical Properties

Physical properties of crystal structures

We now know that state of matter is a physical property, and we know that we can change a substance’s state by heating it. A solid’s particles will increase in kinetic energy, moving faster and faster until enough energy is supplied to break some of the bonds between them. This happens at a certain temperature - the melting point.

But different substances have very different melting points. Sodium chloride melts at 800 °C whereas chlorine gas will remain a liquid until -101.5 °C! This is just one example of their differing physical properties.

What causes these differences? To understand this, we need to look at different types of crystal structures as well as their forces and how they bond.

What is a crystal?

These forces could be intramolecular, such as covalent, metallic, or ionic bonds, or intermolecular, such as van der Waals forces, permanent dipole-dipole forces or hydrogen bonds. We’re interested in four different crystal types:

• Molecular crystals.
• Giant covalent crystals.
• Giant ionic crystals.
• Giant metallic crystals

Molecular crystals

Molecular crystals are made up of simple covalent molecules held together by intermolecular forces. Although strong covalent bonds within each molecule hold the atoms together, the intermolecular forces between molecules are weak and easy to overcome. This gives molecular crystals low melting and boiling points. They are also soft and break easily. An example is chlorine, . Although each chlorine molecule is made up of two covalently bonded chlorine atoms, the only forces between individual molecules are weak van der Waals forces. These do not require much energy to overcome, so chlorine is a gas at room temperature.

A chlorine crystal, made of many chlorine molecules. Each molecule is made from two chlorine atoms held together by a strong covalent bond. However, the only forces between molecules are weak intermolecular forces.commons.wikimedia.org

Another type of physical property is conductivity. Molecular crystals can’t conduct electricity - there are no charged particles free to move within the structure.

Giant covalent crystals

Giant covalent structures are also known as macromolecules.

Like molecular crystals, macromolecules contain covalent bonds, but in this case all the crystal’s particles are atoms covalently bonded together. Because these bonds are so strong, macromolecules are extremely hard and have high melting and boiling points.

An example is diamond (explore more in Carbon Structures). Diamond consists of carbon atoms, each one joined to four other atoms with covalent bonds. Melting diamond would involve breaking these extremely strong bonds. In fact, diamond doesn’t melt at all under atmospheric pressure.

Like molecular crystals, giant covalent crystals can’t conduct electricity, as there are no charged particles free to move within the structure.

A 3D representation of a diamond crystal.commons.wikimedia.org

Giant metallic crystals

When metals bond, they form giant metallic crystals. These consist of a lattice arrangement of positively charged metal ions in a sea of negative delocalised electrons. There is strong electrostatic attraction between the ions and the electrons, holding the crystal together. This gives metals high melting and boiling points.

Because they contain a free-moving sea of delocalised electrons, metals are able to conduct electricity. This is one way of distinguishing them from other structures.

Metallic bonding. There is a strong electrostatic attraction between the positive metal ions and the delocalised electrons. commons.wikimedia.org

Giant ionic crystals

Like metals, ionic lattices contain positive ions. But in this case, they are ionically bonded to negative ions with strong electrostatic attraction. Again, this makes ionic compounds hard and strong with high melting and boiling points.

In a solid state, the ions in ionic crystals are held together tightly in ordered rows. They cannot move out of position and only vibrate on the spot. However, when molten or in solution, the ions can move about freely and so carry a charge. Therefore, only molten or aqueous ionic crystals are good conductors of electricity.

An ionic lattice. commons.wikimedia.org

Comparing properties of structures

Let’s go back to our examples. Sodium chloride, , has a very high melting point. We now know that this is because it is an ionic crystal and its particles are held in position by strong ionic bonds. These require a lot of energy to overcome. We must heat sodium chloride a lot in order for it to melt. In contrast, solid chlorine, , forms a molecular crystal. Its molecules are held together by weak intermolecular forces which don’t require much energy to overcome. Therefore, chlorine has a much lower melting point than sodium chloride.

Sodium chloride, NaCl. The lines represent the strong ionic bonds between oppositely charged ions. Compare this to the chlorine crystal earlier in the article, which only has weak intermolecular forces between its particles.commons.wikimedia.org

The following table should help you summarise the differences in physical properties between the four types of crystal structure we’ve learnt about.

A table comparing the physical properties of different crystal structures.StudySmarter Originals

For more information on any of the types of bonding mentioned above, check out Covalent and Dative Bonding, Ionic Bonding and Metallic Bonding.

Physical properties of water

Like chlorine, solid water forms a molecular crystal. But unlike chlorine, water is liquid at room temperature. To understand why, let’s compare it to another simple covalent molecule, ammonia, . They both have similar relative masses. They are both molecular solids and also both form hydrogen bonds. We could therefore predict that they have similar melting points. Surely they experience similar intermolecular forces between their molecules? But in actual fact, water has a much higher melting point than ammonia. It requires more energy to overcome the forces between its particles. Water is also less dense as a solid than as a liquid, which you should know is unusual for any substance. Let's explore why. (If you aren’t familiar with hydrogen bonding, we’d recommend looking at Intermolecular Forces before continuing.)

Take a look at a water molecule. It contains one oxygen atom and two hydrogen atoms. Each oxygen atom has two lone pairs of electrons. This means that water can form up to four hydrogen bonds - one using each hydrogen atom and one using each of oxygen’s lone pairs of electrons.

Each water molecule can form up to four hydrogen bonds. commons.wikimedia.org

When water is a liquid, the molecules are constantly moving about. The hydrogen bonds between water molecules are constantly being broken and reformed. In fact, not all of the molecules have all four hydrogen bonds. However, when water is solid ice, all of its molecules form the maximum number of hydrogen bonds possible. This forces them into a lattice with all the molecules in a certain orientation, which affects water’s density and melting and boiling points.

Density

Water is less dense as a solid than a liquid. As we mentioned earlier, this is unusual. This is because the arrangement and orientation of the water molecules in their solid lattice pushes them slightly further apart than in a liquid.

Melting point

Water has a relatively high melting point compared to other simple covalent molecules with a similar relative mass. This is because its multiple hydrogen bonds between molecules require a lot of energy to overcome.

Hydrogen bonding in ice and liquid water. Note that each water molecule in ice forms four hydrogen bonds. This pushes the molecules apart into a regular lattice. commons.wikimedia.org

If we compare the structures of water and ammonia, we can explain the difference seen in melting points. Ammonia can only form two hydrogen bonds - one with the single lone pair of electrons on its nitrogen atom, and the other with one of its hydrogen atoms.

Hydrogen bonding between ammonia molecules. Note that each molecule can form a maximum of two hydrogen bonds. StudySmarter Originals

However, we now know that water can form four hydrogen bonds. Because water has twice as many hydrogen bonds as ammonia, it has a much higher melting point. The following table summarises the differences between these two compounds.

A table comparing water and ammonia. StudySmarter Originals