Preparing the Halogens – Trends in Chemical Properties of Group 17 Elements and Their Compounds (Senior 4 Chemistry) | Ganzaa

Trends in Chemical Properties of Group 17 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 17 Elements and Their Compounds

5 lessons

Lesson 2 of 5

Preparing the Halogens

25 min

Taking electrons back

In nature the halogens are found as halide ions, which have already gained an electron. To get the element back, that electron must be taken away again: every method of making a halogen is an oxidation, $\mathrm{2X^- \rightarrow X_2 + 2e^-}$ . The only question is what does the oxidising. For iodide almost any oxidising agent will do; for fluoride nothing chemical is strong enough, and the job needs electricity. This lesson describes how each halogen is made in the laboratory and in industry, and why the methods differ.

Definition

Oxidation of a halide: removal of an electron from each halide ion to form the halogen: $\mathrm{2X^- \rightarrow X_2 + 2e^-}$ . The halide ion acts as a reducing agent.

Electrolysis: decomposition of a molten or dissolved ionic compound by an electric current. Example: in the electrolysis of brine, chloride ions are oxidised to chlorine at the anode.

Chlorine in the laboratory

Concentrated hydrochloric acid is oxidised by a strong oxidising agent:

\mathrm{MnO_2(s) + 4HCl(conc) \xrightarrow{\text{warm}} MnCl_2(aq) + Cl_2(g) + 2H_2O(l)}

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

Potassium manganate(VII) works without heating. The gas is passed through water (to remove hydrogen chloride that comes over with it), then through concentrated sulphuric acid (to dry it), and collected by downward delivery, because chlorine is denser than air (Figure 12.3).

Figure 12.3: Preparing dry chlorine.

Bromine and iodine in the laboratory

Bromide and iodide ions are more easily oxidised than chloride. Heating a bromide or iodide with manganese(IV) oxide and concentrated sulphuric acid releases the halogen, which distils off (bromine as a red-brown vapour that condenses to a liquid, iodine as a violet vapour that condenses to crystals):

\mathrm{2NaBr + MnO_2 + 3H_2SO_4 \rightarrow 2NaHSO_4 + MnSO_4 + Br_2 + 2H_2O}

Chlorine also displaces bromine and iodine from their salts (Lesson 3): $\mathrm{Cl_2 + 2KI \rightarrow 2KCl + I_2}$ .

The halogens in industry

Chlorine is made by the electrolysis of brine (concentrated sodium chloride solution) in a membrane cell, which also gives sodium hydroxide and hydrogen: $\text{anode: } \mathrm{2Cl^-(aq) \rightarrow Cl_2(g) + 2e^-} \qquad \text{cathode: } \mathrm{2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)}$ $\text{overall: } \mathrm{2NaCl(aq) + 2H_2O(l) \rightarrow 2NaOH(aq) + Cl_2(g) + H_2(g)}$
Fluorine cannot be made by any chemical oxidising agent, because $\mathrm{F_2}$ is itself the strongest oxidising agent of all. It is made by electrolysing potassium hydrogendifluoride, $\mathrm{KHF_2}$ , dissolved in anhydrous liquid hydrogen fluoride; water must be absent, because fluorine reacts with it. Henri Moissan first isolated fluorine this way in 1886, after many chemists before him had failed or been poisoned, and he received the 1906 Nobel Prize in Chemistry for it.
Bromine is extracted from sea water and salt-lake brines. The brine is acidified and chlorine is passed in, $\mathrm{Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2}$ ; the bromine is blown out with air or steam and condensed.
Iodine is obtained from brines by oxidising iodide with chlorine, or from the sodium iodate in Chilean nitrate deposits by reducing it with sodium hydrogensulphite.

Worked example 1: chlorine from manganese(IV) oxide

8.70 g of manganese(IV) oxide reacts with excess concentrated hydrochloric acid. (a) Give the oxidation numbers of manganese and chlorine before and after. (b) Calculate the volume of chlorine at room temperature and pressure (24.0 dm³ mol⁻¹). $A_{r}$ : Mn 54.94, O 16.00.

Solution. (a) Mn: +4 in $\mathrm{MnO_2}$ → +2 in $\mathrm{MnCl_2}$ (reduced). Cl: −1 in HCl → 0 in $\mathrm{Cl_2}$ for two of the four chlorine atoms (oxidised); the other two stay −1 in $\mathrm{MnCl_2}$ .
(b) $M(\mathrm{MnO_2}) = 86.94\ \mathrm{g\,mol^{-1}} \qquad n = \frac{8.70}{86.94} = 0.100\ \mathrm{mol} = n(\mathrm{Cl_2})$ $V = 0.100 \times 24.0 = 2.40\ \mathrm{dm^3}$

Worked example 2: chlorine by electrolysis

A membrane cell runs at a current of 1000 A for 1.00 hour. Calculate the mass of chlorine produced. Faraday constant $F = 96\,485\ \mathrm{C\,mol^{-1}}$ ; $M(\mathrm{Cl_2}) = 70.90$ .

Solution. Charge $Q = It = 1000 \times 3600 = 3.60 \times 10^{6}\ \mathrm{C}$ . $n(e^-) = \frac{3.60 \times 10^{6}}{96\,485} = 37.3\ \mathrm{mol} \qquad n(\mathrm{Cl_2}) = \frac{37.3}{2} = 18.7\ \mathrm{mol}$ (two electrons per $\mathrm{Cl_2}$ ). Mass $= 18.7 \times 70.90 = 1.32 \times 10^{3}\ \mathrm{g} \approx 1.32\ \mathrm{kg}$ . Real plants run many cells at far higher currents.

Warning

Common mistakes. (1) Collecting chlorine over water: it dissolves and reacts; collect it by downward delivery (or in a syringe). (2) Drying chlorine before washing it: the water bottle comes first, then the acid, otherwise the gas leaves wet. (3) Saying fluorine is made by "displacing it with another halogen": no oxidising agent is stronger than fluorine, so electrolysis is the only way. (4) Writing $\mathrm{MnO_2 + 2HCl \rightarrow MnCl_2 + H_2O + O}$ : the product is chlorine, and four HCl are needed.

Safety. Chlorine is toxic and attacks the lungs even at low concentration, so it is prepared only in small amounts in a working fume cupboard. Concentrated hydrochloric and sulphuric acids are corrosive; potassium manganate(VII) is a strong oxidiser. Anyone who breathes in chlorine should be moved to fresh air and given medical help; splashes on skin or in the eyes are rinsed with plenty of running water.

Tip

All halogen preparations are oxidations of halide ions. The harder the halide is to oxidise, the stronger the method: $\mathrm{I^-}$ and $\mathrm{Br^-}$ (MnO₂ + acid, or Cl₂), $\mathrm{Cl^-}$ (MnO₂ or KMnO₄ with conc. HCl; electrolysis of brine), $\mathrm{F^-}$ (electrolysis only).

Kigozi Brammer: Laboratory Preparation of Chlorine Using Manganese(IV) Oxide (9 min). Watch how the apparatus is arranged, why the gas is washed with water and dried with concentrated sulphuric acid, and how it is collected.

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Practice

Write the general half-equation for making any halogen from its halide ions, and state whether the halide ions are oxidised or reduced.
Write the equation for the laboratory preparation of chlorine from manganese(IV) oxide, and draw a labelled diagram of the apparatus for collecting dry chlorine.
Explain the order of the two wash bottles in Figure 12.3.
Write the electrode equations for the electrolysis of brine. Name the three products.
Explain why fluorine must be made by electrolysis, and why water must be absent.
Describe how bromine is obtained from sea water, with an equation.
Calculate the volume of chlorine, at room temperature and pressure, released when 3.16 g of potassium manganate(VII) reacts with excess concentrated hydrochloric acid. $M(\mathrm{KMnO_4}) = 158.04$ .
A brine cell produces 1.00 tonne of chlorine. What mass of sodium hydroxide is produced at the same time? $M$ : $\mathrm{Cl_2}$ 70.90, NaOH 40.00.

Show answersHide answers

$\mathrm{2X^- \rightarrow X_2 + 2e^-}$ ; oxidised.
$\mathrm{MnO_2 + 4HCl \rightarrow MnCl_2 + Cl_2 + 2H_2O}$ ; diagram as Figure 12.3: flask with tap funnel, water wash bottle, sulphuric acid wash bottle, gas jar filled by downward delivery.
Water first removes the HCl fumes (it would also wet the gas); concentrated sulphuric acid last dries the chlorine. In the reverse order the gas would leave wet.
Anode: $\mathrm{2Cl^- \rightarrow Cl_2 + 2e^-}$ ; cathode: $\mathrm{2H_2O + 2e^- \rightarrow H_2 + 2OH^-}$ . Chlorine, hydrogen and sodium hydroxide.
Fluorine is the strongest oxidising agent, so no chemical can oxidise fluoride ions; electricity can. Fluorine reacts with water (oxidising it to oxygen), so the electrolyte must be anhydrous.
Acidified brine is treated with chlorine, which oxidises bromide: $\mathrm{Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2}$ ; the bromine is blown out with air or steam and condensed.
$n(\mathrm{KMnO_4}) = \frac{3.16}{158.04} = 0.0200$ mol; $n(\mathrm{Cl_2}) = \frac{5}{2} \times 0.0200 = 0.0500$ mol; $V = 0.0500 \times 24.0 = 1.20\ \mathrm{dm^3}$ .
2 mol NaOH per mol $\mathrm{Cl_2}$ : $n(\mathrm{Cl_2}) = \frac{1.00 \times 10^6}{70.90} = 1.41 \times 10^4$ mol; NaOH $= 2.82 \times 10^4$ mol $\times 40.00 = 1.13 \times 10^6\ \mathrm{g} = 1.13\ \mathrm{t}$ .

Quick check

Why can fluorine be made only by electrolysis?

Key points

Every halogen is made by oxidising its halide: $\mathrm{2X^- \rightarrow X_2 + 2e^-}$ .
Laboratory chlorine: $\mathrm{MnO_2 + 4HCl \rightarrow MnCl_2 + Cl_2 + 2H_2O}$ (or $\mathrm{KMnO_4}$ + HCl); wash with water, dry with conc. $\mathrm{H_2SO_4}$ , collect by downward delivery, in a fume cupboard.
Industrial chlorine: electrolysis of brine (also NaOH and $\mathrm{H_2}$ ); fluorine: electrolysis of $\mathrm{KHF_2}$ in anhydrous HF (Moissan, 1886); bromine: $\mathrm{Cl_2}$ displaces it from brine; iodine: from brines or by reducing iodate.

Sources and credits

Image: Manganese-dioxide-sample.jpg by Benjah-bmm27, via Wikimedia Commons, public domain.
Image: Henri Moissan portrait.jpg, from Popular Science Monthly (1907), via Wikimedia Commons, public domain.
Fluorine's isolation and production: Wikipedia, Fluorine; bromine from brines: Bromine; both CC BY-SA 4.0, adapted by Ganzaa. Moissan's prize: The Nobel Prize in Chemistry 1906, NobelPrize.org.
First aid for chemical exposure: NIOSH Pocket Guide, first aid, CDC.
Video: Laboratory Preparation of Chlorine: Using Manganese(IV) Oxide Explained!, Kigozi Brammer, embedded from YouTube.
The text, apparatus diagram, worked examples and Practice items were written and drawn by Ganzaa.
This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.

1Taking electrons back

2Chlorine in the laboratory

3Bromine and iodine in the laboratory

4The halogens in industry

Taking electrons back

Chlorine in the laboratory

Bromine and iodine in the laboratory

The halogens in industry