Carbon
Carbon is a Group IV, period 2 element — chemical symbol C, atomic number 6, electronic configuration 2,4.
Allotropes of Carbon
Allotropy is the ability of an element to exist in two or more different forms in the same physical state. Carbon’s crystalline allotropes are diamond and graphite; its non-crystalline (amorphous) forms include coal, coke, charcoal, carbon black (soot) and carbon fibres.
Diamond forms from carbon under very high pressure and is the hardest known natural material. Each carbon atom is covalently bonded to four neighbours in a rigid three-dimensional lattice, and the strength and uniformity of that bonding is what makes diamond so hard, so resistant to melting, and chemically inert.
Properties: a transparent crystal with a high refractive index; extremely dense (3.5 g/cm³); very unreactive and heat-resistant; does not conduct electricity, because all four of each carbon’s valence electrons are locked into covalent bonds, leaving none free to carry charge.
Uses: jewellery (its high refractive index gives diamond its sparkle); cutting glass and metal, and as an abrasive for sharpening hard tools and for mining drills — all exploiting its hardness.
Graphite is a black, opaque, slippery crystal with a hexagonal, layered structure: each carbon atom bonds to only three neighbours within its layer, and the layers themselves are held together only by weak Van der Waals forces — which is why graphite is soft and slippery (the layers slide over each other) even though it has a very high melting point. Because each carbon uses only three of its four valence electrons for bonding, the fourth is delocalised and free to move within a layer, which is what makes graphite — unlike diamond — a good conductor of heat and electricity.
Uses: pencil “lead” (soft and slippery); crucibles (very high melting point); electrodes in electroplating and dry cells (a good, chemically inert electrical conductor).
Reaction of Carbon with Metallic Oxides
Heated with the oxide of a metal below it on the reactivity (electrochemical) series, carbon reduces the oxide to the metal and is itself oxidised to carbon dioxide.
2ZnO(s) + C(s) → 2Zn(s) + CO₂(g)
2Fe₂O₃(s) + 3C(s) → 4Fe(s) + 3CO₂(g)
2CuO(s) + C(s) → 2Cu(s) + CO₂(g)
Carbon Dioxide (CO₂)
Preparation: reacting a dilute acid with marble (calcium carbonate).
2HCl(aq) + CaCO₃(s) → CaCl₂(aq) + H₂O(l) + CO₂(g)
Physical properties: colourless, odourless; slightly soluble in water, forming carbonic acid; turns moist blue litmus red; denser than air; at −78°C it sublimes directly to a solid (“dry ice”).
Chemical properties:
- With water: forms the weak acid carbonic acid. CO₂(g) + H₂O(l) → H₂CO₃(aq)
- With alkalis: forms a carbonate. CO₂(g) + 2NaOH(aq) → Na₂CO₃(aq) + H₂O(l). With excess CO₂, a hydrogencarbonate (bicarbonate) forms instead: CO₂(g) + NaOH(aq) → NaHCO₃(aq).
- With burning magnesium: magnesium is hot enough to split carbon dioxide into carbon and oxygen, and then burns in that liberated oxygen. CO₂(g) + 2Mg(s) → 2MgO(s) + C(s). (Carbon dioxide itself doesn’t support ordinary combustion — this is a special case where the burning metal is hot enough to decompose the gas first.)
Test: turns limewater milky, from a precipitate of calcium carbonate. CO₂(g) + Ca(OH)₂(aq) → CaCO₃(s) + H₂O(l). Excess CO₂ redissolves the precipitate, forming soluble calcium hydrogencarbonate: 2CO₂(g) + Ca(OH)₂(aq) → Ca(HCO₃)₂(aq).
Uses: fire extinguishers (denser than air, doesn’t support combustion); carbonating soft drinks; a raising agent in baking; photosynthesis (plants use it to build sugars, releasing oxygen); a refrigerant; food preservation.
Carbon Monoxide (CO)
Preparation: passing carbon dioxide over hot carbon (charcoal), which reduces it to carbon monoxide; any unreacted CO₂ is removed by passing the gas mixture through concentrated sodium or potassium hydroxide, and the CO is collected over water, in which it’s insoluble.
CO₂(g) + C(s) → 2CO(g)
An alternative method dehydrates methanoic (formic) acid or oxalic acid with concentrated sulphuric acid: HCOOH(l) —(conc. H₂SO₄)→ CO(g) + H₂O(l), or H₂C₂O₄(s) —(conc. H₂SO₄)→ CO(g) + CO₂(g) + H₂O(l).
Physical properties: colourless, odourless, tasteless; insoluble in water; slightly less dense than air; neutral to litmus; very poisonous even in fairly low concentrations in air.
Chemical properties:
- Combustion: burns with a blue flame to give carbon dioxide. 2CO(g) + O₂(g) → 2CO₂(g)
- As a reducing agent: reduces heated metal oxides, itself being oxidised to carbon dioxide. CuO(s) + CO(g) → Cu(s) + CO₂(g); Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g)
- With haemoglobin: binds to haemoglobin to form stable carboxyhaemoglobin, which can no longer carry oxygen — this is why carbon monoxide is so dangerous, and why burning charcoal in a poorly ventilated room is risky.
Uses: industrial extraction of metals from their ores; a component of fuel gases like producer gas and water gas; manufacturing organic compounds such as methanol and Perspex.
Carbonates
Carbonates are salts of carbonic acid, a dibasic acid that forms two kinds of salt: the normal carbonate (CO₃²⁻) and the acidic hydrogencarbonate (HCO₃⁻).
Preparing soluble carbonates: pass carbon dioxide into sodium or potassium hydroxide. CO₂(g) + 2NaOH(aq) → Na₂CO₃(aq) + H₂O(l); with excess CO₂, the hydrogencarbonate forms instead — CO₂(g) + NaOH(aq) → NaHCO₃(aq), or CO₂(g) + KOH(aq) → KHCO₃(aq).
Preparing insoluble carbonates: add sodium carbonate or sodium hydrogencarbonate to a solution of the metal’s salt, precipitating the insoluble carbonate.
CaCl₂(aq) + Na₂CO₃(aq) → CaCO₃(s) + 2NaCl(aq)
CaCl₂(aq) + 2NaHCO₃(aq) → CaCO₃(s) + 2NaCl(aq) + H₂O(l) + CO₂(g)
| Metal group | Solubility | Effect of heat | Reaction with dilute acid |
|---|---|---|---|
| K, Na | Soluble in water | Do not decompose on heating | All carbonates react with dilute acid to liberate carbon dioxide |
| Ca, Mg, Zn, Fe, Pb, Cu, Ag | Insoluble in water | Decompose on heating to the oxide + CO₂ (Ag₂CO₃ is the exception — it decomposes all the way to the metal, plus O₂ and CO₂) |
Polymorphs of calcium carbonate: polymorphism is a solid material’s ability to exist in more than one crystal form (related to allotropy, which specifically refers to elements). Calcium carbonate occurs naturally in several such forms — limestone, marble, seashells and coral.
Solvay process (industrial preparation of sodium carbonate): uses ammonia, carbon dioxide (from heating limestone), water and common salt — cheap, plentiful raw materials.
CO₂(g) + H₂O(l) + NH₃(g) + NaCl(aq) —(0–15°C)→ NaHCO₃(s) + NH₄Cl(aq)
2NaHCO₃(s) —(300°C)→ Na₂CO₃(s) + CO₂(g) + H₂O(l)
Uses of sodium carbonate: manufacturing glass; sewage treatment; water softening (washing soda).
Uses of sodium hydrogencarbonate: baking; relieving digestive problems and gastric hyperacidity; treating mouth (buccal) inflammations.
The Carbon Cycle, Global Warming and the Greenhouse Effect
Atmospheric carbon dioxide sits at a roughly constant level, maintained by a balance between processes that release it and processes that remove it.
Processes that release CO₂: burning of carbon-containing fuels; respiration in plants and animals; volcanic eruptions; fermentation of sugars — C₆H₁₂O₆(aq) → 2C₂H₅OH(aq) + 2CO₂(g); the action of dilute acid on carbonates.
Processes that remove CO₂: photosynthesis — 6CO₂(g) + 6H₂O(l) —(sunlight)→ C₆H₁₂O₆(aq) + 6O₂(g); dissolving in rainwater — CO₂(g) + H₂O(l) → H₂CO₃(aq); slow absorption by alkalis as mortar and whitewash harden — Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l).
A greenhouse warms up because sunlight passes through the glass, gets absorbed inside, and is re-radiated as longer-wavelength heat that can’t pass back out through the glass. Carbon dioxide behaves the same way in the atmosphere, trapping heat — which is essential for life (without it, Earth would be too cold), but rising CO₂ levels intensify the effect, driving global warming.
