Classroom Secrets: Teaching the Properties of Covalent Molecules

The Properties of Covalent Molecules and Giant Covalent structures

Just like the properties of ionic properties, covalent molecules are an important part of IGCSE Chemistry. Many students can recall that covalent bonds involve sharing electrons, yet they struggle to explain why methane is a gas while diamond is one of the hardest substances on Earth.

The problem isn’t understanding covalent bonding—it’s recognising that not all covalent substances have the same structure.

This is a complete guide on what to include when teaching the properties of covalent substances. I’ll share my experience and thoughts on what students find most difficult and include key visual aids to support your teaching. This guide follows the Edexcel IGCSE Chemistry specification.

Why do Students Find This Topic Difficult?

Students firstly struggle with identifying covalent substances. These simple rules will help them decide whether they are talking about ionic or covalent substances:

“A metal + a non-metal = Ionic (e.g. NaCl)

A non-metal + a non-metal = Covalent (CH₄)

Common Misconceptions

Examiners love testing the properties of covalent substances. With so many key words and points to remember, it is an area in the course filled with misconceptions. Look out for these early on in your teaching and address as you move through your lesson sequence.

A lot of students aim to just memorise a list of facts on this topic. We as teachers can support them by using lots of visual aids including diagrams drawn live on the board and moly mod kids. The more your students see the structures, the more they will be able to link structure to properties.

Misconception	Why Students Think This	How To Correct This.
All covalent substances have low melting and boiling points.	Students learn about water, methane and carbon dioxide first, then assume all covalent substances behave the same way.	Introduce simple molecular substances before giant covalent structures. Compare methane and diamond side by side and emphasise that structure determines properties.
Covalent bonds are weak.	Students confuse covalent bonds with the weak intermolecular forces between molecules.	Draw both on the board using solid lines for covalent bonds and dotted lines for intermolecular forces. Repeatedly ask, “Which force is actually being overcome?”
The covalent bonds break when a substance melts.	Students assume melting breaks every type of bond.	Explain that in simple molecular substances, only the intermolecular forces are overcome. The covalent bonds inside each molecule remain intact.
Diamond is made of giant molecules.	Students use the word “molecule” for every covalent substance.	Explain that diamond is a giant covalent lattice, not a collection of molecules. Ask, “Where does the molecule end?” to show there is no individual molecule.
Graphite is soft because its covalent bonds are weak.	Students link softness with weak bonding.	Show that the covalent bonds within each layer are very strong. It is the weak forces between the layers that allow them to slide over one another.
Graphite conducts electricity because carbon is a metal.	Students associate electrical conductivity only with metals.	Reinforce that carbon is a non-metal. Graphite conducts because it contains delocalised electrons that are free to move through the layers.
Diamond conducts electricity because it contains carbon atoms.	Students think all forms of carbon behave in the same way.	Compare diamond and graphite. Explain that in diamond, all four outer electrons form covalent bonds, leaving no mobile electrons to carry charge.
All giant covalent structures have the same properties.	Students overgeneralise after learning about diamond.	Compare diamond, graphite. Highlight both their similarities (strong covalent bonds, high melting points) and their differences (hardness and electrical conductivity).
Fullerenes have the same properties as diamond and graphite.	Students assume every carbon allotrope behaves similarly.	Introduce fullerenes as different carbon structures with different properties and applications. Reinforce that different structures lead to different properties.

Before You Start: Check Their Prior Knowledge

Before diving into the properties of covalent substances, it’s worth taking a few minutes to check that students have the necessary prior knowledge. I’ve found that many misconceptions later in the lesson can be traced back to gaps in their understanding of basic bonding concepts.

Students should already be confident with:

• The structure of the atom, including protons, neutrons and electrons.
• Electron arrangement and the idea of a full outer shell.
• Why atoms form chemical bonds.
• How covalent bonds are formed through the sharing of electrons.
• Drawing simple dot-and-cross diagrams for covalent molecules.

A quick retrieval activity at the start of the lesson can help identify any misconceptions before introducing more challenging ideas. This doesn’t need to take long—five minutes is often enough to refresh students’ memories and ensure everyone is ready to move on. Here are a few retrieval questions you could use:

What you need to include in your lessons on the properties of covalent substances

Step 1: Start With the Big Idea

Instead of listing properties immediately, introduce one key question:

Why do substances with covalent bonds have such different properties?

For example, carbon dioxide has a melting point of -78.2 °C but diamond melts at 4,027°C.

Introduce the 2 categories:

• Simple molecular substances
• Giant covalent structures

Explain that structure determines properties.

Step 2: Teach Simple Molecular Substances First

Once your students understand how covalent bonds form through the sharing of electrons, it’s tempting to jump straight into discussing melting points and electrical conductivity. However, I’ve found that taking a step back and focusing on structure first makes the rest of the lesson much easier.

Start by introducing familiar examples such as methane, water, oxygen and carbon dioxide. These are all simple molecular substances made up of small, discrete molecules. Draw a few simple molecular diagrams on the board and ask students what they notice. They should recognise that the atoms within each molecule are held together by strong covalent bonds.

This is a good opportunity to pause and ask an important question:

If covalent bonds are strong, why do substances like methane and oxygen have such low melting and boiling points?

Most students will answer that the covalent bonds must be weak. This is one of the biggest misconceptions you’ll encounter, so don’t correct it immediately. Instead, use it as a teaching opportunity.

Weak intermolecular forces break, not the covalent bonds.

Explain that when a simple molecular substance melts or boils, the covalent bonds inside the molecules are not broken. Those bonds remain intact. Instead, it is the weak intermolecular forces between neighbouring molecules that are overcome. Because these forces are relatively weak, only a small amount of energy is needed for the molecules to move further apart, giving simple molecular substances their low melting and boiling points.

Once students understand this distinction, the other properties become much easier to explain. Since simple molecular substances do not contain charged particles or delocalised electrons that are free to move, they cannot conduct electricity. Whether they are solid or liquid, there is no way for an electric current to pass through the substance.

Summary on properties of simple covalent molecules.

Each property links to a particular structural point. Share this with your students. I like to use a table.

Structural point	Property it links to.
There are weak intermolecular forces between the molecules	Low melting points
There are weak intermolecular forces between the molecules	Low boiling points
No ions or delocalised electrons in simple covalent molecules	Poor electrical conductivity

Step 3: Introduce Giant Covalent Structures

Once your students are confident explaining the properties of simple molecular substances, it’s time to introduce the idea that not all covalent substances are the same. This is often the turning point in the lesson because students begin to realise that the type of structure is just as important as the type of bonding.

I like to begin with a simple comparison of diamond and methane. Ask your students to say what they see. If they need scaffolding, pose simple questions like:

• Which one has more bonds? How do you know?
• Which one looks like a molecule? What makes you say this?
• How do you think more bonds will affect melting point?
• Do we see any charged particles?

Summary on structure and properties of DIAMOND (a giant covalent substance)

After the class discussion, these are the key points students must know about the structure of diamond.

• It is a giant covalent structure
• It is made of carbon atoms only
• Continuous network of covalent bonds (Each carbon atom bonds to 4 other carbon atoms).
• No individual molecules in the structure
• Contains many strong covalent bonds

Again, link each structural point to a specific property:

Structure	Property it links to
Many strong covalent bonds	High melting point
Many strong covalent bonds	High boiling point
Many strong covalent bonds	Very hard
No delocalised electrons or flowing ions	Does not conduct electricity

Step 4: Don’t Rush Graphite – It’s Where Students Usually Get Stuck

If there’s one part of this topic that deserves extra teaching time, it’s graphite. Time and again, I’ve found that students can happily describe diamond as being “hard with a high melting point,” but as soon as graphite appears, their confidence disappears.

The reason is simple: graphite seems to break all of the rules. It’s a non-metal, yet it conducts electricity. It has strong covalent bonds, yet it’s soft enough to be used in pencils. Unless students understand why this happens, they’re left trying to memorise isolated facts instead of seeing the underlying pattern.

I like to begin by reminding students that graphite still has a giant covalent structure. At first glance, it looks similar to diamond because both are made entirely of carbon atoms joined by strong covalent bonds. However, this is where the similarities end.

Draw the layered structure on the board and ask students what they notice. Encourage them to spot that each carbon atom is bonded to only three others, leaving one electron from each carbon atom free to move throughout the layers. These delocalised electrons are the reason graphite can conduct electricity, despite being a non-metal.

Next, focus on the layers themselves. Explain that while the covalent bonds within each layer are very strong, the forces between the layers are weak. This means the layers can slide over one another when a force is applied, making graphite soft and slippery. Linking this directly to its use in pencil “lead” helps students connect the chemistry to something familiar.

Summary on structure and properties of GRAPHITE (a giant covalent substanc).

Structure	Property it links to
Many strong covalent bonds	High melting point (but lower than diamond)
Many strong covalent bonds	High boiling point (but lower than diamond)
Contains layers which can slide	Soft
Contains delocalised electrons between the layers	Conducts electricity.

Step 5: Don’t Forget Fullerenes

Fullerenes are often taught alongside diamond and graphite. They are members of the giant covalent structures; family, but they different in that they exist as a molecular substance. This means they have lower melting and boiling points than both diamond and graphite.

Although a very short section in the Edexcel IGCSE specification, they are still important and allow your students to really show their understanding of the topic.