Carbon and Its Inorganic Compounds

Lessons

Senior 3 · Chemistry

Carbon and Its Inorganic Compounds

5 lessons

Lesson 2 of 5

Reactions of Carbon and Carbon Monoxide

25 min

Black rock, silver metal

Rwanda mines cassiterite, a heavy black mineral that is mostly tin(IV) oxide, $\mathrm{SnO_2}$ . At Karuruma in Kigali, the LuNa Smelter turns this rock into tin ingots that are more than 99.9% pure metal. Something has to pull the oxygen away from the tin. In most smelters around the world that job is done by the element of this unit: carbon. To see why carbon is so good at it, look at how carbon reacts with oxygen and with the compounds that contain oxygen.

Carbon and oxygen

All forms of carbon burn in air or oxygen once they are hot enough; even diamond burns. With plenty of oxygen the product is carbon dioxide (complete combustion):

$\mathrm{C(s) + O_2(g) \rightarrow CO_2(g)}$

With too little oxygen, as in a smouldering charcoal stove, carbon monoxide forms as well (incomplete combustion):

$\mathrm{2C(s) + O_2(g) \rightarrow 2CO(g)}$

Both reactions are exothermic, which is why charcoal and coke are fuels. In S1 you saw why a charcoal stove must never be used in a closed room. This lesson explains what carbon monoxide is and why it is so dangerous.

Carbon and carbon dioxide

When carbon dioxide passes over carbon at red heat (about 1000 °C), the carbon takes one oxygen atom from each molecule and both end up as carbon monoxide:

$\mathrm{C(s) + CO_2(g) \rightarrow 2CO(g)}$

This happens in the lower part of a charcoal fire, where carbon dioxide made by burning rises through more hot charcoal. It is also a way to make carbon monoxide in the laboratory (Figure 2.1).

Definition

Reducing agent (reductant): a substance that removes oxygen from another substance (or gives it electrons). The substance that loses oxygen is reduced; the reducing agent gains oxygen and is itself oxidised. Example: in $\mathrm{2CuO(s) + C(s) \rightarrow 2Cu(s) + CO_2(g)}$ , carbon is the reducing agent: copper(II) oxide is reduced to copper and carbon is oxidised to carbon dioxide.

Carbon as a reducing agent

Experiment. Mix a little black copper(II) oxide powder with powdered charcoal, heat the mixture strongly in a hard-glass test tube, and lead the gas given off into limewater. The limewater turns milky, and when the tube cools, reddish-brown copper is left:

$\mathrm{2CuO(s) + C(s) \rightarrow 2Cu(s) + CO_2(g)}$

Carbon can remove oxygen from the oxides of metals that are below it in the reactivity series (zinc, iron, tin, lead and copper), but not from the oxides of metals above it, such as aluminium, magnesium or sodium, which hold their oxygen too strongly. This is why carbon, as coke or charcoal, is used to extract iron and tin:

$\mathrm{2Fe_2O_3(s) + 3C(s) \rightarrow 4Fe(s) + 3CO_2(g)}$

$\mathrm{SnO_2(s) + 2C(s) \rightarrow Sn(l) + 2CO(g)}$

In a blast furnace, most of the iron(III) oxide is actually reduced by carbon monoxide made from the coke: coke burns to carbon dioxide, the carbon dioxide meets more hot coke and becomes carbon monoxide, and the carbon monoxide removes the oxygen from the ore:

$\mathrm{Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(l) + 3CO_2(g)}$

Carbon and concentrated acids

Carbon does not react with dilute acids. Hot concentrated sulfuric acid and nitric acid are strong oxidising agents, and they oxidise carbon to carbon dioxide while being reduced themselves:

$\mathrm{C(s) + 2H_2SO_4(l) \rightarrow CO_2(g) + 2SO_2(g) + 2H_2O(l)}$

$\mathrm{C(s) + 4HNO_3(l) \rightarrow CO_2(g) + 4NO_2(g) + 2H_2O(l)}$

With nitric acid you see choking brown fumes of nitrogen dioxide; with sulfuric acid, the gas has the sharp smell of sulfur dioxide. Both gases are toxic, so these reactions are shown only by a teacher, in a fume cupboard.

Table 2.1: Reactions of carbon

Carbon reacts with	Conditions	Products	Carbon acts as
oxygen (plenty)	heat to start	$\mathrm{CO_2}$	fuel (it is oxidised)
oxygen (too little)	heat to start	$\mathrm{CO}$	fuel (it is oxidised)
carbon dioxide	red heat, about 1000 °C	$\mathrm{CO}$	reducing agent
copper(II), iron(III), tin(IV) oxides	strong heat	metal + $\mathrm{CO_2}$ or $\mathrm{CO}$	reducing agent
concentrated $\mathrm{H_2SO_4}$	hot	$\mathrm{CO_2}$ , $\mathrm{SO_2}$ , $\mathrm{H_2O}$	reducing agent
concentrated $\mathrm{HNO_3}$	hot	$\mathrm{CO_2}$ , $\mathrm{NO_2}$ , $\mathrm{H_2O}$	reducing agent
dilute acids	any	no reaction	not reactive

Carbon monoxide

Preparation. In the laboratory, carbon monoxide can be made by passing dry carbon dioxide over strongly heated charcoal (Figure 2.1). Some carbon dioxide passes through unchanged, so the gas is bubbled through sodium hydroxide solution, which absorbs carbon dioxide but not carbon monoxide:

$\mathrm{CO_2(g) + 2NaOH(aq) \rightarrow Na_2CO_3(aq) + H_2O(l)}$

Carbon monoxide is collected over water because it is almost insoluble in water, and because its density is so close to that of air ( $M_{r} = 28$ , against about 29 for air) that it cannot be collected by displacing air.

Figure 2.1: Making carbon monoxide from carbon dioxide and hot charcoal (a teacher demonstration).

Warning

Safety. Carbon monoxide is highly toxic, and you cannot see, smell or taste it. It is prepared only by a teacher, in a working fume cupboard, and any gas left over is burned off at a jet (it burns to harmless carbon dioxide). If anyone feels a headache, dizziness or sickness, stop, get everyone into fresh air and tell the teacher at once.

Properties of carbon monoxide

A colourless, odourless gas, almost insoluble in water, slightly less dense than air.
A neutral oxide: it does not react with water, acids or alkalis, so it does not turn limewater milky and has no effect on litmus.
It burns in air with a pale blue flame, forming carbon dioxide: $\mathrm{2CO(g) + O_2(g) \rightarrow 2CO_2(g)}$ . The blue flames that flicker on top of a glowing charcoal fire are carbon monoxide burning.
It is a reducing agent: passed over hot copper(II) oxide, it turns the black powder into reddish-brown copper: $\mathrm{CuO(s) + CO(g) \rightarrow Cu(s) + CO_2(g)}$ .
It is poisonous. In the lungs it combines with haemoglobin, the red pigment in blood that carries oxygen, more than 200 times more strongly than oxygen does. Haemoglobin that has taken up carbon monoxide cannot carry oxygen, so the brain and heart are starved of oxygen even though the person is still breathing.

Worked example

Problem. Iron(III) oxide is reduced by carbon: $\mathrm{2Fe_2O_3(s) + 3C(s) \rightarrow 4Fe(s) + 3CO_2(g)}$ . Calculate (a) the mass of iron and (b) the mass of carbon needed, for 1.00 tonne of iron(III) oxide. (Fe = 56, O = 16, C = 12; 1 tonne = $10^6\,\mathrm{g}$ )

Solution. $M(\mathrm{Fe_2O_3}) = 2 \times 56 + 3 \times 16 = 160\,\mathrm{g\,mol^{-1}}$ .

$n(\mathrm{Fe_2O_3}) = \frac{1.00 \times 10^{6}\,\mathrm{g}}{160\,\mathrm{g\,mol^{-1}}} = 6250\,\mathrm{mol}$

(a) From the equation, 2 mol $\mathrm{Fe_2O_3}$ give 4 mol Fe, so $n(\mathrm{Fe}) = 2 \times 6250 = 12\,500\,\mathrm{mol}$ and the mass is $12\,500 \times 56 = 700\,000\,\mathrm{g} = 0.700\,\mathrm{t}$ .

(b) 2 mol $\mathrm{Fe_2O_3}$ need 3 mol C, so $n(\mathrm{C}) = \frac{3}{2} \times 6250 = 9375\,\mathrm{mol}$ and the mass is $9375 \times 12 = 112\,500\,\mathrm{g} \approx 0.113\,\mathrm{t}$ (113 kg).

Worked example

Problem. A colourless gas X has no effect on limewater or on damp litmus paper. When X is lit, it burns with a pale blue flame, and the gas produced turns limewater milky. When X is passed over heated black powder Y, the powder turns reddish-brown. Identify X and Y and write equations for the two reactions.

Solution. X is neutral (no effect on litmus or limewater), burns, and forms a gas that turns limewater milky, so it contains carbon: X is carbon monoxide. It reduces Y to a reddish-brown metal, copper, so Y is copper(II) oxide.

$\mathrm{2CO(g) + O_2(g) \rightarrow 2CO_2(g)} \qquad \mathrm{CuO(s) + CO(g) \rightarrow Cu(s) + CO_2(g)}$

Tip

"Carbon grabs oxygen." Whenever carbon or carbon monoxide meets an oxide it can beat, it takes the oxygen: the oxide is reduced, and the carbon is oxidised (to $\mathrm{CO}$ or $\mathrm{CO_2}$ ). If the metal is above carbon in the reactivity series, carbon loses the tug-of-war.

FuseSchool: Extraction Using Heat (4 min). Watch how coke, carbon monoxide and iron ore react in the blast furnace, and write the equation for each step as it appears.

✦

Practice

Write balanced equations, with state symbols, for the complete and the incomplete combustion of carbon.
In the reaction $\mathrm{2CuO(s) + C(s) \rightarrow 2Cu(s) + CO_2(g)}$ , which substance is reduced and which is oxidised? Which is the reducing agent?
Can carbon be used to extract aluminium from aluminium oxide? Explain.
Give two physical and two chemical properties of carbon monoxide.
Carbon monoxide is called a neutral oxide. What does this mean, and what would you see if you bubbled it through limewater?
What mass of copper is obtained when 8.0 g of copper(II) oxide is completely reduced by carbon? (Cu = 64, O = 16)
On a cold night, Habimana carries a glowing charcoal stove into his bedroom and closes the door and window. Use two equations from this lesson to explain why this could kill him, and suggest what he should do instead.

Show answersHide answers

$\mathrm{C(s) + O_2(g) \rightarrow CO_2(g)}$ ; $\mathrm{2C(s) + O_2(g) \rightarrow 2CO(g)}$ .
Copper(II) oxide is reduced (loses oxygen); carbon is oxidised (gains oxygen). Carbon is the reducing agent.
No. Aluminium is above carbon in the reactivity series, so it holds oxygen more strongly than carbon can take it away (aluminium is extracted by electrolysis).
Physical (any two): colourless; odourless; almost insoluble in water; slightly less dense than air. Chemical (any two): burns with a blue flame to $\mathrm{CO_2}$ ; reduces metal oxides such as CuO; neutral, no reaction with acids or alkalis; combines with haemoglobin.
It does not react with acids or with alkalis (or water). Nothing: the limewater stays clear.
$M(\mathrm{CuO}) = 80$ ; $n = \frac{8.0}{80} = 0.10\,\mathrm{mol}$ CuO, which gives 0.10 mol Cu; mass $= 0.10 \times 64 = 6.4\,\mathrm{g}$ .
In the closed room the oxygen is soon used up, so the charcoal burns incompletely: $\mathrm{2C(s) + O_2(g) \rightarrow 2CO(g)}$ ; carbon dioxide rising through hot charcoal also becomes carbon monoxide: $\mathrm{C(s) + CO_2(g) \rightarrow 2CO(g)}$ . The colourless, odourless carbon monoxide combines with his haemoglobin so his blood cannot carry oxygen, and he may die in his sleep. He should never bring a burning charcoal stove into a bedroom or any closed room: he should leave it outside and keep warm with blankets.

Sources and credits

Written by Ganzaa. The facts about the Karuruma smelter come from About LuNa Smelter, LuNa Smelter Ltd (read 30 September 2026).
Photo: Cassiterite - Mt Bischoff mine, Waratah, Tasmania, Australia.jpg, Ralph Bottrill, licensed under CC BY 3.0, via Wikimedia Commons.
Video: Extraction Using Heat, FuseSchool – Global Education, embedded from YouTube.
The diagram, table, worked examples and Practice items were written and drawn by Ganzaa.

1Black rock, silver metal

2Carbon and oxygen

3Carbon and carbon dioxide

4Carbon as a reducing agent

5Carbon and concentrated acids

6Carbon monoxide