Carbon: allotropes and oxides: WASSCE Chemistry

The diamonds dug from the gravels of Kono and the charcoal burned in every coal pot are the same element: carbon. Its atoms can join in different ways, giving the hardest natural substance and a soft, black, slippery one. This lesson covers the allotropes of carbon and its two oxides, carbon(IV) oxide and the poisonous carbon(II) oxide.

Learning objectives

By the end of this lesson you should be able to:

describe the structures and properties of diamond and graphite and relate them to their uses
describe the amorphous forms of carbon and their uses
state the chemical properties of carbon
describe the laboratory preparation, test, properties and uses of carbon(IV) oxide
describe the preparation, properties and dangers of carbon(II) oxide
compare carbon(IV) oxide and carbon(II) oxide

Allotropes of carbon

Diamond: each carbon bonded to four others in a rigid three-dimensional network. Graphite: flat layers of hexagons held by weak forces, with delocalised electrons

Figure 1. The structures of diamond and graphite. WASSCE Tutor diagram.

Property	Diamond	Graphite
Structure	Each C bonded to 4 others; tetrahedral 3-D giant covalent network	Each C bonded to 3 others in flat hexagonal layers; layers held by weak forces
Hardness	Hardest natural substance	Soft and slippery: layers slide
Appearance	Transparent, sparkling crystals	Black, shiny, greasy
Electrical conductivity	Does not conduct: all electrons in bonds	Conducts: one delocalised electron per atom moves along the layers
Density	3.5 g/cm³	2.3 g/cm³
Melting point	Very high (about 3550 °C)	Very high
Uses	Jewellery; cutting, drilling and grinding tools	Pencil "lead" (with clay); lubricant; electrodes in electrolysis and dry cells; moderator in nuclear reactors

Both are giant covalent structures, so both have very high melting points and are insoluble in water. Burning either in excess oxygen gives only carbon(IV) oxide, which proves both are pure carbon.

Newer allotropes: fullerenes (such as C₆₀, "buckyballs") and graphene (a single graphite layer, very strong and highly conducting).

Amorphous carbon

Forms of carbon with no regular crystal structure (actually tiny graphite crystals):

Form	Made by	Use
Charcoal (wood charcoal)	Heating wood with little air	Cooking fuel; adsorbs gases and colours
Activated charcoal	Charcoal treated to give a huge surface area	Water filters; gas masks; removing colour from sugar; treating poisoning
Coke	Heating coal without air	Fuel and reducing agent in the blast furnace
Soot (lampblack)	Incomplete burning of oils	Printing ink, shoe polish, car tyres
Animal charcoal	Heating bones	Decolourising raw sugar

Chemical properties of carbon

1. Burns in oxygen: in plenty of oxygen it forms carbon(IV) oxide; in limited oxygen, carbon(II) oxide.

\text{C} + \text{O}_2 \rightarrow \text{CO}_2 \qquad 2\text{C} + \text{O}_2 \rightarrow 2\text{CO}

\begin{gathered}\text{C} + \text{O}_2 \\[6pt] \rightarrow \text{CO}_2 \\[6pt] 2\text{C} + \text{O}_2 \\[6pt] \rightarrow 2\text{CO}\end{gathered}

2. Reducing agent: carbon removes oxygen from the oxides of metals below it in the reactivity series (zinc, iron, lead, copper). This is the basis of metal extraction:

2\text{CuO} + \text{C} \rightarrow 2\text{Cu} + \text{CO}_2 \qquad 2\text{ZnO} + \text{C} \rightarrow 2\text{Zn} + \text{CO}_2

\begin{gathered}2\text{CuO} + \text{C} \\[6pt] \rightarrow 2\text{Cu} + \text{CO}_2 \\[6pt] 2\text{ZnO} + \text{C} \\[6pt] \rightarrow 2\text{Zn} + \text{CO}_2\end{gathered}

3. With steam at white heat, it forms water gas: C + H₂O → CO + H₂.

4. Oxidised by hot concentrated acids (nitric and sulphuric) to CO₂.

Carbon(IV) oxide, CO₂

Dilute hydrochloric acid on marble chips in a flask with a thistle funnel; carbon(IV) oxide collected by downward delivery

Figure 2. Preparing carbon(IV) oxide. WASSCE Tutor diagram.

Laboratory preparation: dilute hydrochloric acid on marble chips (calcium carbonate):

\text{CaCO}_3\text{(s)} + 2\text{HCl(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}

\begin{gathered}\text{CaCO}_3\text{(s)} + 2\text{HCl(aq)} \\[6pt] \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}\end{gathered}

Dilute sulphuric acid is not used: it forms insoluble calcium sulphate, which coats the marble and stops the reaction.

Collection: by downward delivery (denser than air), or over water (it is only fairly soluble).

Test: it turns lime water (calcium hydroxide solution) milky, because insoluble calcium carbonate forms. With excess CO₂, the milkiness clears, because soluble calcium hydrogencarbonate forms:

\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}

\begin{gathered}\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \\[6pt] \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}\end{gathered}

\text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)} \rightarrow \text{Ca(HCO}_3)_2\text{(aq)}

\begin{gathered}\text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)} \\[6pt] \rightarrow \text{Ca(HCO}_3)_2\text{(aq)}\end{gathered}

Properties:

colourless, odourless gas, about 1.5 times denser than air
fairly soluble in water, forming weak carbonic acid (H₂CO₃), which turns blue litmus faintly red
acidic oxide: reacts with alkalis, CO₂ + 2NaOH → Na₂CO₃ + H₂O
does not burn and does not support combustion of most substances. But burning magnesium continues to burn in it, leaving black specks of carbon: 2Mg + CO₂ → 2MgO + C
solid CO₂ ("dry ice") sublimes at −78 °C

Uses:

fire extinguishers (dense, does not support burning)
fizzy drinks (dissolved under pressure)
dry ice for refrigeration and stage "smoke"
baking: CO₂ from baking powder or yeast makes bread and cakes rise
photosynthesis: the source of carbon in all plant material

Carbon(II) oxide, CO

Preparation: pass carbon(IV) oxide over red-hot coke (CO₂ + C → 2CO), or warm methanoic acid with concentrated sulphuric acid (which removes water).

Properties:

colourless, odourless, very poisonous gas; about the same density as air
almost insoluble in water; neutral oxide
burns with a blue flame: 2CO + O₂ → 2CO₂
reducing agent: reduces metal oxides, as in the blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂

Why it is so dangerous: CO combines with haemoglobin in red blood cells about 200 times more strongly than oxygen does, forming carboxyhaemoglobin. Blood can then no longer carry oxygen. Because CO has no smell, victims do not notice, become drowsy and may die. It is produced by incomplete combustion in charcoal stoves, generators and car engines in enclosed spaces.

CO₂ and CO compared

Property	Carbon(IV) oxide, CO₂	Carbon(II) oxide, CO
Smell	Odourless	Odourless
Toxicity	Not poisonous (suffocates only at high levels)	Very poisonous
Density compared with air	Denser	About the same
Solubility in water	Fairly soluble	Almost insoluble
Nature	Acidic oxide	Neutral oxide
Burning	Does not burn	Burns with a blue flame
Lime water	Turns milky	No effect
Reducing agent?	No	Yes

In Sierra Leone

Sierra Leone's diamonds, mined in Kono and along rivers in the east, formed from carbon under enormous pressure deep in the earth and were brought up in volcanic rock. The charcoal sold at roadsides is made by burning wood with limited air in earth-covered kilns. It burns hotter and with less smoke than wood, but its large-scale production cuts down forests.

Common mistake

Graphite conducts electricity because of its delocalised electrons, not because it is a metal. Diamond, with every electron held in a bond, does not conduct.

Key points

Diamond: each C bonded to 4, 3-D network: hardest, non-conductor; used for jewellery and cutting.
Graphite: layers of hexagons with weak forces and delocalised electrons: soft, slippery, conducts; used for pencils, lubricants and electrodes.
Amorphous carbon: charcoal, activated charcoal, coke, soot.
Carbon is a reducing agent, used to extract zinc, iron, lead and copper.
CO₂: CaCO₃ + 2HCl; turns lime water milky; dense, acidic, does not support burning; extinguishers, drinks, baking.
CO: poisonous, odourless, burns with a blue flame, a reducing agent; forms by incomplete combustion.

WASSCE examination focus

WAEC tests this

Expect to compare diamond and graphite, relating structure to properties and uses, to describe the preparation and test of CO₂ (and why H₂SO₄ is not used), to explain the lime water test with excess CO₂, and to compare CO and CO₂, including why CO is poisonous.

Quick check

Why does graphite conduct electricity but diamond does not?
Why is dilute sulphuric acid not used with marble chips to prepare CO₂?
What happens when excess CO₂ is bubbled through lime water?
Why is carbon(II) oxide poisonous?

Answers:

Graphite has one delocalised electron per atom, free to move along its layers; in diamond all electrons are held in bonds.
It forms insoluble calcium sulphate, which coats the chips and stops the reaction.
It first turns milky (CaCO₃), then clears as soluble calcium hydrogencarbonate forms.
It combines with haemoglobin and stops the blood carrying oxygen.

Carbon: allotropes and oxides