Diamond vs Silicon Dioxide: Two Giant Covalent Structures

Exam boards like to pair diamond and silicon dioxide, and with good reason. Draw a small piece of each and they look like cousins: every atom in the middle of the drawing is joined by strong covalent bonds to its neighbours, the bonds point to the corners of tetrahedra, and the network carries on in three dimensions until it reaches the edge of the crystal. Both are hard, both have very high melting points, neither conducts electricity. A student who learns one can nearly write out the other.

“Nearly” is the important word. Diamond is a single element; silicon dioxide is a compound with two different atoms playing two different roles. That changes the bond angles, the hardness, the density, the way each responds to heat and chemicals, and even the way you should phrase your answer when a mark scheme asks you to “describe the structure”. This comparison lays the two side by side so the similarities and the differences are both easy to see.

Two networks built from tetrahedra

In diamond, every carbon atom forms four single covalent bonds to four other carbon atoms. The four bonds point to the corners of a regular tetrahedron, 109.5° apart, which is what you expect for sp³ carbon. Each of those neighbours is itself bonded to four carbons, and so on. There are no separate molecules; the whole crystal is effectively one molecule.

In silicon dioxide (as quartz, the common crystalline form), every silicon atom forms four single covalent bonds to four oxygen atoms, again pointing to the corners of a tetrahedron. Each oxygen atom is bonded to two silicon atoms. So the structure is a network of SiO₄ tetrahedra that share every corner with a neighbouring tetrahedron.

That corner-sharing is the key to the formula. Each silicon “owns” four oxygens, but each oxygen is shared between two silicons, so each silicon effectively has 4 × ½ = 2 oxygens. The ratio is 1 : 2, hence SiO₂, even though there is no such thing as a discrete SiO₂ molecule in quartz. The formula is an empirical ratio, not a molecular formula.

Side-by-side comparison

Feature	Diamond	Silicon dioxide (quartz)
Type of substance	Element (an allotrope of carbon)	Compound
Atoms present	C only	Si and O
Bonds per atom	Each C bonded to 4 C	Each Si bonded to 4 O; each O bonded to 2 Si
Basic building unit	CC₄ tetrahedron	SiO₄ tetrahedron
Bond type	Non-polar C–C single bonds	Polar Si–O single bonds
Bond angle at the central atom	109.5° at every C	About 109.5° at Si; the Si–O–Si angle at O is bent and wide
Lone pairs in the network	None	Two on each oxygen
Structure	Giant covalent (3D network)	Giant covalent (3D network)
Hardness (Mohs scale)	10, the reference top of the scale	7
Melting behaviour	Does not melt at normal pressure; converts or sublimes at extreme temperature	Melts at about 1,710 °C
Electrical conductivity	Insulator	Insulator
Thermal conductivity	Exceptionally high	Much lower
Burns in oxygen?	Yes, to CO₂ at high temperature	No, it is already fully oxidised
Common forms	Gemstones, industrial abrasives	Quartz, sand, flint; amorphous silica in glass

What the two have in common

The similarities come from the same root cause: a continuous network of strong covalent bonds in three dimensions.

• High melting point. Melting or vaporising either solid means breaking a huge number of covalent bonds, not just separating molecules. That takes a great deal of energy.
• Hardness. To scratch or deform the surface you have to break bonds that are locked in every direction. Nothing can slip.
• No electrical conduction. Every outer electron is tied up in a localised σ bond (or, in SiO₂, a lone pair on oxygen). There are no delocalised electrons and no ions, so no charge carriers.
• Insolubility. Water molecules cannot pull individual atoms out of a covalent network, so neither dissolves in water.
• Brittleness. Covalent bonds are directional. They resist deformation right up to the point where they break, and then the crystal cleaves or chips rather than bending.

If a question just asks why both have high melting points, “giant covalent structure, many strong covalent bonds must be broken” earns the marks. The general pattern for this structure type is covered in giant covalent structures.

Difference 1: one element versus two

Diamond is pure carbon, so every bond is identical and non-polar: both ends of a C–C bond attract the shared pair equally.

Silicon dioxide is a compound of silicon and oxygen. Oxygen is much more electronegative than silicon, so each Si–O bond is polar, with a partial negative charge on the oxygen. The bonds are still covalent (the electronegativity gap is not large enough for a simple ionic description to work well), but they have noticeable ionic character. That is why some textbooks describe silica as sitting part-way along the bonding continuum rather than as a textbook non-polar network.

Difference 2: the shape around each atom

Around every atom in diamond the geometry is the same: four bonds, tetrahedral, 109.5°.

In quartz, silicon is tetrahedral, but oxygen is not a copy of carbon. Oxygen forms only two bonds, to two silicons, and it keeps two lone pairs. The Si–O–Si link is therefore bent, like a stretched version of the angle in water. The angle is wider than in water, and it can flex a little, which lets the tetrahedra twist relative to each other. That flexibility is one reason silica can form many crystalline arrangements and can also be frozen into a disordered, glassy state. Diamond’s rigid network allows much less of that freedom.

Difference 3: hardness

Diamond defines the top of the Mohs scale at 10; quartz sits at 7. Both are hard enough to scratch window glass, but diamond scratches quartz and not the other way round.

The reasons are partly geometric. Diamond’s C–C bonds are short (154 pm), strong, and packed densely in every direction, with every atom joined identically to its neighbours. Quartz’s network is more open: the bent Si–O–Si links act a little like hinges, and there are more “gaps” in the structure. You can see the openness in the densities, too: diamond is around 3.5 g/cm³ while quartz is about 2.65 g/cm³, even though silicon is a heavier atom than carbon.

Difference 4: what happens on heating

Silicon dioxide melts, at about 1,710 °C, into a very viscous liquid. Cool that liquid quickly and the network cannot reorganise into neat crystals; you get fused silica, a glass. Ordinary window glass is silica with other oxides added to lower the working temperature.

Diamond behaves differently. At normal atmospheric pressure carbon does not form a liquid at all; it sublimes at extremely high temperature. The melting point stored for carbon in our element data is 3,823 K (about 3,550 °C), but that figure only has meaning under high pressure, so it is not a fair comparison with quartz’s normal melting point. In practice, diamond heated strongly in the absence of air tends to convert into graphite first, and heated in air it simply burns to carbon dioxide. Silica cannot burn: its silicon is already bonded to oxygen.

Difference 5: heat conduction

Neither conducts electricity, but diamond is one of the best conductors of heat of any material, far better than quartz and better than copper at room temperature. Heat in diamond is carried by vibrations passing through the lattice, and a stiff network of light, identical atoms joined by very strong bonds transmits those vibrations extremely efficiently. Quartz’s heavier, mixed atoms and flexible links scatter the vibrations much more. This is why diamond is used to spread heat away from some high-power electronic devices, and why a real diamond feels cold to the touch compared with a glass imitation.

How to describe each in an exam

Diamond: “Giant covalent structure. Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. Many strong covalent bonds must be broken to melt it. No delocalised electrons, so it does not conduct electricity.”

Silicon dioxide: “Giant covalent structure. Each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms, forming a 3D network. Many strong covalent bonds must be broken to melt it. No delocalised electrons or free ions, so it does not conduct.”

Examiners frequently penalise answers that say “SiO₂ molecules” or “intermolecular forces” for either substance. Neither has molecules, so there are no intermolecular forces to overcome when it melts. Simple molecular vs giant covalent is worth reading if that distinction still feels shaky, and the giant structures practice set has questions to test it.

Common mistakes

• Treating silicon dioxide as a small molecule like CO₂. Carbon dioxide is a simple molecular gas; silicon dioxide is a giant network. Same group, very different structures, because silicon does not form stable double bonds to oxygen the way carbon does.
• Saying each oxygen bonds to four atoms. Silicon bonds to four; oxygen bonds to two.
• Quoting a melting point for diamond at normal pressure. It sublimes (or converts to graphite) instead.
• Linking thermal conductivity to free electrons. Diamond conducts heat without any free electrons, through lattice vibrations.
• Calling Si–O bonds ionic. They are polar covalent bonds within a covalent network.

Key takeaways

• Both diamond and silicon dioxide are giant covalent networks built from tetrahedral units, so both are hard, high-melting, insoluble in water and electrical insulators.
• Diamond: each C bonded to four C, non-polar bonds, C–C 154 pm, Mohs hardness 10, an outstanding heat conductor, burns in air.
• Silicon dioxide: each Si bonded to four O and each O to two Si, polar Si–O bonds, bent Si–O–Si links, Mohs hardness 7, melts at about 1,710 °C, does not burn.
• The formula SiO₂ is a ratio from corner-sharing tetrahedra, not a molecule.
• Never mention intermolecular forces when explaining the melting of either substance.