Did you know that diamond and graphite are made entirely of carbon yet behave so differently that one is used for cutting tools and the other can carry an electrical current? If you try to explain everything using the word “carbon”, the contrast seems mysterious. If you examine how the atoms are arranged and where the electrons can move, it becomes a powerful Chemistry lesson.
I use diamond, graphite and silicon dioxide with JC students because these materials reveal why chemical formula alone does not predict a substance’s physical properties. An exam may describe a solid that is hard, brittle, conductive or high melting, then ask you to infer its structure. The best answer begins with atoms, bonding and mobile charge carriers—not with a memorised product picture.
This H1 Chemistry Tutor Hougang guide helps JC1 and JC2 students connect giant covalent networks to melting points, conductivity, hardness and real-world applications. It also contrasts these materials with simple molecular substances and metals, providing a practical method for unseen structure-and-bonding questions under Singapore–Cambridge H1 Chemistry 8873.
The quick comparison
Diamond is a three-dimensional giant covalent network in which each carbon atom is bonded to four neighbours in an approximately tetrahedral arrangement. It is extremely hard, has a very high sublimation or melting-related temperature under appropriate conditions, and normally conducts electricity poorly.
Graphite is also a giant covalent form of carbon, but each carbon is bonded to three neighbours within approximately planar hexagonal layers. The fourth valence electron contributes to a delocalised electronic system, so graphite can conduct electricity, especially parallel to the layers. The layers can slide relative to one another.
Silicon dioxide, SiO₂, is a giant covalent network with silicon and oxygen atoms connected through a continuous structure. It is generally a high-melting electrical insulator in ordinary dry solid form. The properties follow from the extended bonding arrangement, not from molecules of SiO₂ floating individually in a crystal.
Why “giant covalent” matters
A simple molecular substance contains distinct molecules with covalent bonds inside each molecule and intermolecular attractions between molecules. Melting or boiling often primarily overcomes those attractions rather than the internal covalent bonds.
A giant covalent structure is different. Covalent bonds extend throughout a network. To disrupt the solid structure substantially, considerable energy is required to break or reorganise covalent connections.
Do not say that every giant covalent structure has exactly the same properties. The connectivity, dimensionality, strength of bonds and electron availability determine significant differences.
Diamond: four bonds per carbon
Carbon has four valence electrons. In diamond, each carbon forms four covalent bonds to neighbouring carbon atoms. The local bonding geometry is approximately tetrahedral, giving an extended three-dimensional network.
The structure has no easy, weakly connected layers equivalent to those in graphite. Deformation can require disrupting strong covalent connections, contributing to diamond’s great hardness.
A useful drawing shows one central carbon with four bonds directed into a three-dimensional arrangement. A flat square with four C–C bonds is not an accurate representation of its tetrahedral local geometry.
Why diamond is hard
Hardness is resistance to local deformation such as scratching or indentation. In diamond, a connected network of strong covalent bonds extends in three dimensions, making the structure highly resistant to the movement of atoms past one another.
This is not simply because a diamond contains “many atoms”. A large wax block also contains many atoms but is soft because its structure is held together differently at the scale relevant to deformation.
The examiner wants the correct structural cause: a continuous network of strong covalent bonds and no easily sliding planes analogous to graphite.
Why diamond does not ordinarily conduct electricity
In the simple bonding model, diamond’s four valence electrons per carbon are used in the covalent network. There is no significant population of mobile charge carriers in an ordinary pure diamond sample under typical conditions.
Its electrons are associated with the bonding network rather than being freely available to carry current as in a metal or in graphite’s delocalised electronic system.
There are specialised diamond materials, doping methods and surface effects that create useful electronic behaviour. Those extensions do not overturn the basic H1 comparison: ordinary diamond is an excellent electrical insulator.
Graphite: three bonds and layers
Each carbon in a graphite layer is connected to three other carbons in a hexagonal arrangement. The layers are networks of covalent bonds extending in two dimensions.
Graphite is often represented as flat sheets stacked together. Strong covalent bonding acts within each layer, while the interactions holding the layers together are much weaker and include dispersion forces.
This creates an important difference from diamond. Both have strong C–C bonds, but the arrangement of those bonds and the interactions between structural units differ.
What is the fourth electron doing?
In the school model, one carbon valence electron per graphite carbon contributes to a delocalised electronic system associated with the layer. This allows mobile electronic charge carriers along the planes.
Graphite therefore conducts electrical current along layers in ordinary conditions. Its conductivity is anisotropic: properties differ when measured parallel to a layer and across layers.
If a question asks why graphite conducts, do not say “because graphite is a metal”. It is giant covalent. The relevant mechanism is electron delocalisation in the layered carbon structure.
Why graphite can be slippery
The graphite layers are strongly bonded internally but interact more weakly with their neighbours. The layers can shift relative to one another under appropriate conditions, contributing to graphite’s use as a dry lubricant.
This contrasts with the strong three-dimensional network of diamond. Both contain carbon and covalent bonds, yet their deformation behaviours differ.
A common mistake is to say “graphite is soft because its C–C covalent bonds are weak”. The C–C bonds within each layer are strong. The easier movement concerns the planes and their weaker interlayer attractions.
Graphite does not conduct equally in every direction
The delocalised electrons are most mobile along graphite layers, so conductivity parallel to the planes is much greater than across them.
This is a useful H1 example of how structural direction can matter. A property can be directional even in a chemically uniform material.
For a simple written question, explaining the presence of delocalised electrons in the layers is normally the central point. A more detailed data question may give separate in-plane and perpendicular measurements and ask students to interpret the difference.
Silicon dioxide: a connected network, not separate SiO₂ molecules
Silicon dioxide contains silicon and oxygen in a 1:2 atomic ratio. In crystalline quartz, each silicon is associated with four oxygen neighbours in a tetrahedral network, and each bridging oxygen connects two silicon centres.
The formula SiO₂ gives a ratio of atoms rather than identifying independent little SiO₂ molecules inside quartz.
This structural picture explains why silica behaves unlike a simple molecular oxide. The extended Si–O covalent network requires substantial energy to disrupt.
Why silica has a high melting point
Breaking or reorganising a substantial three-dimensional network of strong Si–O covalent bonds requires considerable energy. That helps explain silicon dioxide’s high melting behaviour.
It is not sufficient to say “silica has a large relative molecular mass”. The solid is a giant network, not a collection of discrete SiO₂ molecules with one conventional molecular mass separating by weak attractions on melting.
The structure–property connection is the skill the question is testing.
Why silica is generally an electrical insulator
In common crystalline silicon dioxide, electrons are associated with its covalent bonding structure, and there is no comparable supply of mobile electronic carriers as in graphite or copper.
The dry solid therefore conducts electricity poorly under ordinary conditions. This makes silica useful in many electrical and electronic insulation applications when prepared in appropriate forms.
Do not automatically transfer the word “silicon” into a claim that silica is a semiconductor like elemental silicon. Silicon dioxide and silicon are different substances with different structures and electronic properties.
A comparison with elemental silicon
Elemental crystalline silicon has a giant covalent structure and semiconducting behaviour that can be controlled through doping and device engineering. Silicon dioxide is a compound of silicon and oxygen and is commonly used as an insulator.
They share the element silicon but differ in chemical composition and electronic structure. A student should never infer that silicon dioxide behaves exactly like elemental silicon because the names look related.
This is a real-world reminder that chemical formulas are not decoration. Adding oxygen can create a materially different substance.
A three-column structure audit
When you see an unfamiliar solid, write three headings: particles, bonding network, mobile charge carriers.
For diamond: carbon atoms, three-dimensional covalent network, no ordinary mobile charge carriers. For graphite: carbon atoms, covalently bonded layers with weaker interlayer interactions, delocalised carriers in planes. For silica: silicon and oxygen atoms, extended covalent network, no ordinary mobile charge carriers.
You can then infer hardness, conductivity and melting behaviour rather than recalling each property as a separate fact.
Why graphite conducts but molten salt also conducts
Graphite conduction arises from delocalised electrons. Molten sodium chloride conducts because positive and negative ions can move in the liquid.
Both materials can transport electric charge, but the mobile carriers are different. The chemical mechanisms must not be conflated.
A question that asks you to identify the mobile particle rather than just whether something conducts is designed to reveal whether you understand the structure.
Why copper is a third distinct case
Copper is a metal. It conducts through mobile electrons in a metallic electronic structure held together by attraction between positive metal ion cores and delocalised electrons.
Graphite has delocalised electrons too, but its atoms are covalently bonded into layers rather than forming the simple metallic ion-core lattice model.
When explaining conductivity, identify both the mobile carriers and the actual bonding category. A shared property need not imply an identical structure.
Worked question: identify three unknown solids
Suppose an exam provides an original table describing Solid A as very hard and nonconductive, Solid B as soft and electrically conductive along planes, and Solid C as high-melting and nonconductive with silicon and oxygen atoms in an extended network.
The likely identifications are A diamond-like carbon, B graphite-like carbon, and C silicon dioxide. To justify A, cite its three-dimensional covalent network. For B, cite layered structure and delocalised electrons. For C, cite the Si–O network and lack of ordinary mobile carriers.
This is not a universal identification method from three observations alone. Real materials can share properties. The provided composition and structural clues make the assignment defensible in the stated exercise.
Worked question: is graphite molecular?
A student claims graphite must contain separate C₆ molecules because the hexagonal rings contain six carbon atoms. That is incorrect.
The hexagons are repeating features of a sheet that extends through many covalently bonded carbon atoms. They are not discrete C₆ molecules whose boundaries terminate at every hexagon.
A six-membered ring drawn inside a lattice is a structural motif, not automatically a separate molecular entity.
Worked question: why does diamond not have easy slip planes?
A student may reason that a solid made only of carbon must consist of stacked layers. The correct distinction is connectivity.
Diamond’s covalent network extends in three dimensions, so the movement of large planes requires disrupting numerous covalent connections. Graphite’s layers, by contrast, can move relative to one another through comparatively weaker interlayer interactions.
The difference is not the atomic symbol. It is the way the same atoms are connected.
What about graphene?
Graphene is a single atom-thick sheet of carbon atoms arranged in a hexagonal lattice, closely related to an individual graphite layer. It has useful electronic, thermal and mechanical properties.
Graphene helps students connect the familiar graphite structure to modern materials research, but it is not necessary to replace a clear H1 explanation with complex quantum band calculations.
If a question supplies information about graphene, use that data. Distinguish it from a stack of graphite layers and from the three-dimensional network of diamond.
Why simple molecular iodine is different
Iodine at ordinary conditions consists of I₂ molecules arranged in a molecular solid. Strong covalent bonding acts within each I₂ molecule, while comparatively weaker intermolecular attractions hold the molecular solid together.
During ordinary melting, the I–I covalent bonds are not all broken. The molecules become able to move past one another.
That is the key contrast with giant covalent solids: the principal structural connections that must be disrupted during phase change are different.
The bond-versus-force mistake
Students often say “diamond melts at a high temperature because it has strong intermolecular forces”. That misidentifies its structure. Diamond is giant covalent, not a molecular substance held together principally by intermolecular forces.
Another student may say “graphite is soft because the covalent bonds are weak”. That ignores the strong bonds within graphite layers and the weaker interactions between them.
A good correction asks which atoms or groups are moving during the property being described. That selects the relevant type of bonding or attraction.
What the examiner means by “explain”
If asked “Explain why graphite conducts electricity”, a concise strong answer is: “Graphite contains delocalised electrons within its covalently bonded carbon layers. These electrons can move along the layers and carry electric charge.”
If asked “Explain why diamond is hard”, a suitable answer is: “Each carbon is covalently bonded to four neighbouring carbons in a three-dimensional network. Deforming this network requires disrupting many strong covalent bonds.”
The answers differ because the questions target different properties. There is no one sentence about carbon that answers every structural question.
A seven-day revision sequence
Day one: draw a diamond network and count each carbon’s four neighbours. Day two: draw a graphite layer and identify three neighbours per carbon. Day three: explain delocalised electrons and direction-dependent conductivity.
Day four: compare the ways diamond and graphite deform. Day five: draw the silicon dioxide network using Si–O–Si connections. Day six: compare giant covalent, ionic, metallic and simple molecular solids. Day seven: attempt unfamiliar structure–property questions without headings or diagrams supplied.
If a learner cannot distinguish bonds within layers from attractions between layers, pause and repair that idea before adding more data interpretation.
How parents can recognise genuine understanding
Ask why a pencil’s graphite can make a mark while diamond is suitable for a cutting tool. Then ask why graphite can conduct electricity but diamond usually cannot.
Look for the words “layers”, “three-dimensional network” and “delocalised electrons”, used in the correct places rather than as a collection of memorised keywords.
A Hougang H1 Chemistry tutor should be able to diagnose whether the student is confusing molecular, giant covalent or metallic structures. In a small group, the useful test is whether each student can explain a new material independently after correction.
A-Level and university connections
The official 2027 H1 Chemistry syllabus directory provides the current examination scope. Structure, bonding and properties sit at the heart of the H1 curriculum, and data-based questions may require students to use those ideas in unfamiliar material contexts.
These concepts matter in semiconductor devices, cutting tools, sensors, ceramics and advanced carbon materials. University entry requirements differ, and some science and engineering courses require specified H2 subjects. An H1 Chemistry grade alone should not be assumed to establish eligibility for all those programmes.
For now, the foundation is the ability to move from a drawn network to a property without an unexplained leap.
Frequently asked questions
Are diamond and graphite made of the same element?
Yes. Both are allotropes of carbon, with different atomic arrangements and bonding structures.
Why is diamond harder than graphite?
Diamond has a strong three-dimensional covalent network. Graphite contains strong covalent layers that can slide relative to one another because interlayer interactions are comparatively weaker.
Why does graphite conduct electricity?
It has delocalised electronic charge carriers associated with its carbon layers.
Why does diamond not conduct electricity well?
Ordinary pure diamond does not have a substantial population of mobile charge carriers at typical conditions.
Is silicon dioxide a molecular solid?
Crystalline quartz is a giant covalent network of silicon and oxygen atoms, not an ordinary collection of discrete SiO₂ molecules.
Why is silicon dioxide high melting?
Its extended network of strong Si–O covalent bonds requires substantial energy to disrupt.
Does graphite consist of separate six-carbon rings?
No. Its six-membered rings are repeated features of continuous covalently bonded sheets.
Does electrical conductivity always mean a material is metallic?
No. Graphite is giant covalent and conducts through delocalised electrons; ionic melts and solutions conduct through mobile ions.
Continue your Hougang H1 Chemistry studies
Read H1 metallic bonding, alloys and electrical conductivity to complete the comparison. For ionic structures, use H1 lattice energy, ionic radius and MgO versus NaCl. The JC1 H1 chemical bonding and physical properties article provides another entry point.
Families can review JC1 H1 Chemistry tuition routes and JC2 H1 Chemistry tutorials or use eduKate’s consultation information for current arrangements.
When you see a new solid described in Paper 2, ask what connects its atoms and what lets any charged particle move. The answer usually begins to appear before you remember a single product name.