GCSE Chemistry Carbon & nanoparticles revision

Carbon sits in Group 4 with four outer electrons, so it forms four strong covalent bonds. Few elements combine that bonding power with the ability to bond to itself in endless chains and rings, and carbon exploits it to the full: the allotropes of carbon are different structures built from the very same atoms. A property like hardness, conductivity or melting point belongs to the structure, not the element — which is why the soft, black pencil core and the transparent gemstone on a ring are chemically identical.

In diamond, every carbon atom is covalently bonded to four neighbours arranged tetrahedrally, and the network repeats through the whole crystal. The consequences all flow from one fact: no electrons are left over. Every outer electron is locked into a bond, so there are no charged particles free to move — diamond does not conduct electricity. Melting it would mean breaking covalent bonds, so its melting behaviour sits at the top of the scale (it sublimes near 3600 °C under ordinary pressure). The rigid three-way network also makes diamond the hardest natural substance — the standard for cutting tools and drill bits — and a fine conductor of heat, because lattice vibrations pass efficiently through the stiffly joined atoms.

In graphite, each carbon atom bonds to only three neighbours, forming flat hexagonal layers. The fourth outer electron of every atom is delocalised: freed from bonding, it roams along its layer. These mobile electrons let graphite conduct electricity — unusual for a non-metal — which is why graphite electrodes carry current through molten ionic compounds in electrolysis. Within each layer the covalent bonds are strong, so graphite keeps diamond’s very high melting behaviour; the difference lies between the layers.

No covalent bonds hold the layers to each other — only weak intermolecular forces. Press a pencil and whole layers shear off onto the paper; slide a machine part lubricated with graphite and the layers glide over one another. The weak inter-layer forces also leave small gaps that let atoms slide in between the sheets — the trick behind intercalation chemistry in batteries.

Graphene is a single one-atom-thick sheet of the graphite lattice — the first two-dimensional material ever isolated. It inherits graphite’s delocalised electrons, so it conducts electricity superbly, and its honeycomb of strong C–C bonds makes it phenomenally strong for its mass. Applications chase those two properties: flexible electronics and touchscreens, and light composite materials where a whisper of graphene adds enormous strength.

The fullerenes close the sheet into hollow shapes. The most famous, buckminsterfullerene (C₆₀), is a football-shaped cage of 60 atoms; related cages come in many sizes, and each can trap other molecules or metal atoms inside — the idea behind drug delivery, where a fullerene ferries a medicine to its target, and behind lubricants, where ball-shaped molecules roll between surfaces. Fullerenes also act as catalysts and their large surface area makes them efficient at it.

Carbon nanotubes roll the sheet into a cylinder a few nanometres wide but potentially centimetres long. Every atom along the tube wall is bonded as in graphene, giving high tensile strength with almost no weight — nanotubes stitch through composite materials to reinforce them, from sports equipment to aircraft parts, and their conductivity promises wiring at molecular scale.

Particle sizes come in three bands. Coarse particles (dust, pollen) are above 2500 nm; fine particles (like many combustion soots) run from 100 to 2500 nm; and nanoparticles sit between 1 nm and 100 nm across — each particle a clump of a few hundred to a few thousand atoms. A nanometre is a billionth of a metre: a nanoparticle is to a football roughly what the football is to the whole Earth. The chemistry-only material in the rest of this chapter asks what happens when matter is engineered at that scale.

A material’s chemistry happens at its surface, so what matters is not the surface area alone but the surface area to volume ratio (SA:V). Shrink a cube and the ratio explodes. A cube of side 1 cm has SA 6 cm², volume 1 cm³, ratio 6:1. Cut the side to 1 mm: SA is 6 × (0.1 × 0.1) = 0.06 cm², volume 0.001 cm³, ratio 60:1 — ten times the ratio from a ten-times smaller side. Halve a side and the ratio doubles; nanoparticles of a few nanometres have ratios thousands of times a lump’s.

The consequence: in a nanoparticle, an enormous fraction of the atoms sit on the surface, exposed and ready to interact. That is why nanoparticle catalysts (where only surface atoms work) are startlingly effective from tiny masses, why nano-sunscreens spread clear yet block ultraviolet light, and why materials as ordinary as gold begin to show entirely new optical or catalytic behaviour at nanometre size.

Industry exploits both signatures. Big surface area makes nanoparticles superb catalysts and gives cosmetics their invisibility: nano titanium dioxide and zinc oxide screen ultraviolet light in sunscreens without leaving a white film; nano-scale silver fights bacteria in deodorants and dressings. Tiny size lets nanoparticles pass through cell membranes, promising targeted drug delivery straight into diseased cells, and lets electronics pack transistors ever smaller. Added in small percentages, nanoparticles also modify materials — stronger plastics, stain-proof fabrics, self-cleaning glass.

The same two signatures carry the risks. Because particles that small can cross biological boundaries, they may reach places bulk versions never could — deep in the lungs, through the skin, into cells — and nobody yet knows the long-term effects of that exposure, for people or for ecosystems. Nanoparticle waste could accumulate in waterways and food chains. Sensible use therefore means researching effects before scaling up, and the exam expects exactly that balanced judgement: valuable uses, genuinely unknown risks, regulation still catching up.