Chlorine and Its Inorganic Compounds

Lessons

Senior 3 · Chemistry

Chlorine and Its Inorganic Compounds

4 lessons

Lesson 1 of 4

Preparing and Testing Chlorine

25 min

A pale green gas from a black mineral

In 1774 the Swedish chemist Carl Wilhelm Scheele warmed the black mineral pyrolusite with hydrochloric acid and collected a choking, yellow-green gas. In 1810 Humphry Davy showed that the gas was an element and named it chlorine, from the Greek chloros, "pale green". Pyrolusite is manganese(IV) oxide, and the reaction Scheele used is still the standard laboratory preparation of chlorine. Chlorine is so reactive that it is never found free in nature; almost all of it is locked up as chloride ions, above all in sodium chloride, common salt.

A glass bottle filled with chlorine gas. The colour is a pale yellow-green; you can see it only because the bottle is full of the gas. Photo: W. Oelen, CC BY-SA 3.0, via Wikimedia Commons

Definition

Halogens: the elements of Group 17 of the periodic table: fluorine, chlorine, bromine and iodine. Their atoms have seven outer electrons, so they gain one electron to form ions with a charge of −1, called halide ions (fluoride, chloride, bromide, iodide). Example: chlorine, $\mathrm{Cl_2}$ , forms chloride ions, $\mathrm{Cl^-}$ , in sodium chloride.

Preparing chlorine in the laboratory

Chlorine is made by oxidising the chloride in concentrated hydrochloric acid. Manganese(IV) oxide is the oxidising agent, and the mixture is warmed (Figure 1.1):

$\mathrm{MnO_2(s) + 4HCl(aq) \rightarrow MnCl_2(aq) + Cl_2(g) + 2H_2O(l)}$

Potassium manganate(VII) is a stronger oxidising agent, and with it no heating is needed:

$\mathrm{2KMnO_4(s) + 16HCl(aq) \rightarrow 2KCl(aq) + 2MnCl_2(aq) + 5Cl_2(g) + 8H_2O(l)}$

The gas carries some hydrogen chloride fumes from the concentrated acid. Bubbling it through water removes them, because hydrogen chloride is extremely soluble in water while chlorine is only moderately soluble.
It is then bubbled through concentrated sulfuric acid to dry it. The order matters: if the gas were dried first, it would pick up water again in the first bottle.
Chlorine is collected by downward delivery, because it is about 2.5 times as dense as air ( $M_{r} = 71$ ). Its colour shows when the jar is full.

Figure 1.1: Laboratory preparation of dry chlorine (a teacher demonstration).

Warning

Safety. Chlorine is poisonous: it attacks the lining of the nose, throat and lungs, and was used as a weapon in the First World War. Prepare it only in a working fume cupboard, in small amounts, with eye protection. Concentrated hydrochloric acid gives off corrosive fumes. If anyone breathes chlorine, move them into fresh air at once, keep them calm and upright, and get medical help.

Testing for chlorine

Table 1.1: Tests for chlorine gas

Test	Result	Why
look at the gas	pale yellow-green	chlorine is the only common coloured gas of this colour
damp blue litmus paper	turns red, then white	chlorine forms acids with water (red), and one of them, hypochlorous acid, bleaches the dye (white)
damp starch–iodide paper	turns blue-black	chlorine displaces iodine, which colours starch blue-black

The starch–iodide test works because chlorine is more reactive than iodine and takes the electrons from iodide ions:

$\mathrm{Cl_2(g) + 2KI(aq) \rightarrow 2KCl(aq) + I_2(aq)}$

Physical properties of chlorine

A pale yellow-green gas with a choking smell; poisonous.
About 2.5 times as dense as air.
Moderately soluble in water, giving a pale green solution, chlorine water.
Easily liquefied under pressure (it boils at −34 °C), so it is transported and stored as a liquid in steel cylinders.

Worked example

Problem. 8.7 g of manganese(IV) oxide reacts completely with excess concentrated hydrochloric acid. (a) What volume of chlorine is produced at room temperature and pressure? (b) What mass of hydrogen chloride reacts? (Mn = 55, O = 16, H = 1, Cl = 35.5; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Solution. $M(\mathrm{MnO_2}) = 55 + 32 = 87\,\mathrm{g\,mol^{-1}}$ , so $n = \frac{8.7}{87} = 0.10\,\mathrm{mol}$ .

(a) 1 mol $\mathrm{MnO_2}$ gives 1 mol $\mathrm{Cl_2}$ : volume $= 0.10 \times 24 = 2.4\,\mathrm{dm^3}$ .

(b) 1 mol $\mathrm{MnO_2}$ reacts with 4 mol HCl: $n(\mathrm{HCl}) = 0.40\,\mathrm{mol}$ , mass $= 0.40 \times 36.5 = 14.6\,\mathrm{g}$ . Only half of this chlorine (0.20 mol of Cl atoms out of 0.40) ends up as chlorine gas; the rest stays as chloride in manganese(II) chloride.

Worked example

Problem. What mass of potassium manganate(VII) is needed to make $1.2\,\mathrm{dm^3}$ of chlorine at room temperature and pressure, using the reaction without heating? ( $\mathrm{KMnO_4} = 158$ ; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Solution. $n(\mathrm{Cl_2}) = \frac{1.2}{24} = 0.050\,\mathrm{mol}$ . From the equation, 2 mol $\mathrm{KMnO_4}$ give 5 mol $\mathrm{Cl_2}$ :

$n(\mathrm{KMnO_4}) = 0.050 \times \frac{2}{5} = 0.020\,\mathrm{mol} \qquad \text{mass} = 0.020 \times 158 = 3.16\,\mathrm{g}$

Tip

Water first, then acid; then down. Wash out hydrogen chloride with water, dry with concentrated sulfuric acid, collect by downward delivery. Test: damp blue litmus goes red, then white.

Rugby School Chemistry: Preparations of Chlorine (7 min). Compare the two methods: potassium manganate(VII) from 1:40 (no heating) and manganese(IV) oxide from 3:11 (heating needed).

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Practice

Write the equation for the preparation of chlorine from manganese(IV) oxide and concentrated hydrochloric acid. What is the role of the manganese(IV) oxide?
In Figure 1.1, what is removed by (a) the water, (b) the concentrated sulfuric acid? Why must the water bottle come first?
Give two reasons why chlorine is collected by downward delivery.
Describe two tests for chlorine and their results.
Write the equation for the reaction of chlorine with potassium iodide solution. What colour does the solution turn?
Why can chlorine not be collected over water?
What volume of chlorine, at room temperature and pressure, can be made from 17.4 g of manganese(IV) oxide? ( $\mathrm{MnO_2} = 87$ ; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Show answersHide answers

$\mathrm{MnO_2(s) + 4HCl(aq) \rightarrow MnCl_2(aq) + Cl_2(g) + 2H_2O(l)}$ . It is the oxidising agent: it oxidises chloride ions to chlorine.
(a) Hydrogen chloride fumes; (b) water vapour. If the gas were dried first, it would become wet again when it passed through the water.
It is much denser than air ( $M_{r}$ 71 against 29); it dissolves in water (and it is coloured, so you can see when the jar is full).
Damp blue litmus paper turns red and then white (bleached); damp starch–iodide paper turns blue-black. (Also: its yellow-green colour.)
$\mathrm{Cl_2(g) + 2KI(aq) \rightarrow 2KCl(aq) + I_2(aq)}$ ; the solution turns brown (iodine).
It is moderately soluble in water and reacts with it, so much of it would be lost.
$n(\mathrm{MnO_2}) = \frac{17.4}{87} = 0.200\,\mathrm{mol}$ , which gives 0.200 mol $\mathrm{Cl_2}$ ; volume $= 0.200 \times 24 = 4.80\,\mathrm{dm^3}$ .

Sources and credits

Chemistry 2e, 18.11: Occurrence, Preparation, and Properties of Halogens (the preparation of chlorine by oxidising chloride, and the properties of the halogens), Paul Flowers, Klaus Theopold, Richard Langley and William R. Robinson, OpenStax, via LibreTexts, licensed under CC BY 4.0. Adapted by Ganzaa.
Photo: Chlorine in bottle.jpg, W. Oelen, licensed under CC BY-SA 3.0, via Wikimedia Commons.
Video: Preparations of Chlorine, Rugby School Chemistry, embedded from YouTube.
The diagram, table, worked examples and Practice items were written and drawn by Ganzaa.

1A pale green gas from a black mineral

2Preparing chlorine in the laboratory

3Testing for chlorine

4Physical properties of chlorine

A pale green gas from a black mineral

Preparing chlorine in the laboratory

Testing for chlorine

Physical properties of chlorine