Trends in Chemical Properties of Group 17 Elements and Their Compounds

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Senior 4 · Chemistry

Trends in Chemical Properties of Group 17 Elements and Their Compounds

5 lessons

Lesson 3 of 5

Chemical Properties of the Halogens

25 min

Electron thieves

Chlorine bubbled into a colourless solution of potassium bromide turns it orange; bubbled into potassium iodide it turns it brown. Iron wool burns in chlorine to make iron(III) chloride, but heated with iodine it gives only iron(II) iodide. All these observations have one explanation: the halogens are oxidising agents, they take electrons, and the smaller the halogen atom, the more strongly it takes them. This lesson ranks the halogens as oxidising agents and uses that order to explain their reactions with halide ions, metals, hydrogen, water and alkalis.

Definition

Oxidising agent: a substance that accepts electrons and is itself reduced. Each halogen is reduced to halide ions: $\mathrm{X_2 + 2e^- \rightarrow 2X^-}$ .

Displacement reaction (of halogens): a reaction in which a more reactive halogen oxidises the halide ions of a less reactive halogen, releasing it. Example: $\mathrm{Cl_2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br_2(aq)}$ .

Disproportionation: a reaction in which one element is simultaneously oxidised and reduced. Example: $\mathrm{Cl_2 + 2OH^- \rightarrow Cl^- + ClO^- + H_2O}$ .

The order of oxidising power

Standard electrode potentials measure how strongly each halogen pulls electrons (Figure 12.4): the more positive $E^\circ$ , the stronger the oxidising agent.

Figure 12.4: Standard electrode potentials, $\mathrm{X_2 + 2e^- \rightleftharpoons 2X^-}$ .

Oxidising power falls from fluorine to iodine because:

the atoms get bigger, so an incoming electron is further from the nucleus and more shielded, and is attracted less;
smaller halide ions are more strongly hydrated in water, which releases more energy (fluoride most of all);
fluorine's weak F–F bond (Lesson 1) makes its molecules easy to split.

Displacement reactions

A halogen oxidises the halide ions of any halogen below it in the group. The colours are clearest if a little hexane (or cyclohexane) is shaken with the mixture: the halogen dissolves in the upper hexane layer.

Halogen added	to KCl(aq)	to KBr(aq)	to KI(aq)
chlorine water	no reaction	orange solution; orange-red in hexane: $\mathrm{Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2}$	brown solution; violet in hexane: $\mathrm{Cl_2 + 2I^- \rightarrow 2Cl^- + I_2}$
bromine water	no reaction	no reaction	brown solution; violet in hexane: $\mathrm{Br_2 + 2I^- \rightarrow 2Br^- + I_2}$
iodine solution	no reaction	no reaction	no reaction

Reactions with metals

All the halogens react with metals to form halides; the reactions get less vigorous down the group. Sodium burns in chlorine with a yellow flame, $\mathrm{2Na + Cl_2 \rightarrow 2NaCl}$ . With a metal that has two oxidation states, the strength of the oxidising agent decides the product:

\mathrm{2Fe(s) + 3Cl_2(g) \rightarrow 2FeCl_3(s)} \qquad \mathrm{Fe(s) + I_2(s) \rightarrow FeI_2(s)}

Chlorine (and bromine) can oxidise iron all the way to +3; iodine is too weak an oxidising agent, so it stops at iron(II).

Reactions with hydrogen

The reaction $\mathrm{H_2 + X_2 \rightarrow 2HX}$ becomes less vigorous down the group: fluorine and hydrogen explode even in the dark and cold; chlorine and hydrogen explode in sunlight (Lesson 1); bromine reacts only when heated; iodine reacts slowly and reversibly on heating, because HI is easily decomposed.

Reactions with water

Fluorine oxidises water to oxygen: $\mathrm{2F_2 + 2H_2O \rightarrow 4HF + O_2}$ .
Chlorine dissolves and partly reacts, giving chloric(I) acid (hypochlorous acid), which kills bacteria and bleaches: $\mathrm{Cl_2(aq) + H_2O(l) \rightleftharpoons HCl(aq) + HOCl(aq)}$ Chlorine is oxidised (0 to +1 in HOCl) and reduced (0 to −1 in HCl) at the same time: a disproportionation. Blue litmus turns red (acid), then white (bleached by HOCl).
Bromine reacts in the same way but to a smaller extent; iodine is only very slightly soluble in water, though it dissolves well in potassium iodide solution as $\mathrm{I_3^-}$ .

Reactions with alkalis

With sodium hydroxide, chlorine disproportionates, and the products depend on the temperature:

\text{cold, dilute: } \mathrm{Cl_2 + 2NaOH \rightarrow NaCl + NaClO + H_2O}

\text{hot, concentrated: } \mathrm{3Cl_2 + 6NaOH \rightarrow 5NaCl + NaClO_3 + 3H_2O}

In the cold, chlorine goes to −1 (chloride) and +1 (chlorate(I), the active ingredient of household bleach); when hot, to −1 and +5 (chlorate(V)). Hot alkali makes chlorate(I) disproportionate further: $\mathrm{3ClO^- \rightarrow 2Cl^- + ClO_3^-}$ .

Worked example 1: why iron gives FeCl₃ but FeI₂

Use $E^\circ$ values to explain why chlorine and bromine oxidise iron(II) to iron(III) but iodine does not. $E^\circ(\mathrm{Fe^{3+}/Fe^{2+}}) = +0.77\ \mathrm{V}$ .

Solution. A reaction is feasible when $E^\circ_{\text{cell}} = E^\circ(\text{oxidising agent}) - E^\circ(\mathrm{Fe^{3+}/Fe^{2+}})$ is positive.

Chlorine: $1.36 - 0.77 = +0.59\ \mathrm{V}$ : feasible, $\mathrm{Cl_2 + 2Fe^{2+} \rightarrow 2Cl^- + 2Fe^{3+}}$ .
Bromine: $1.07 - 0.77 = +0.30\ \mathrm{V}$ : feasible.
Iodine: $0.54 - 0.77 = -0.23\ \mathrm{V}$ : not feasible. In fact the reverse happens: $\mathrm{Fe^{3+}}$ oxidises iodide to iodine, $\mathrm{2Fe^{3+} + 2I^- \rightarrow 2Fe^{2+} + I_2}$ .

Worked example 2: chlorine in hot alkali

21.3 g of chlorine is passed into excess hot, concentrated sodium hydroxide. (a) Give the oxidation numbers of chlorine in each product. (b) Calculate the mass of sodium chlorate(V) formed. $M$ : $\mathrm{Cl_2}$ 70.90, $\mathrm{NaClO_3}$ 106.44.

Solution. $\mathrm{3Cl_2 + 6NaOH \rightarrow 5NaCl + NaClO_3 + 3H_2O}$ .
(a) 0 in $\mathrm{Cl_2}$ ; −1 in NaCl (reduced); +5 in $\mathrm{NaClO_3}$ (oxidised). Check: five Cl atoms each gain one electron (5 electrons); one Cl atom loses five. Balanced.
(b) $n(\mathrm{Cl_2}) = \frac{21.3}{70.90} = 0.300\ \mathrm{mol} \qquad n(\mathrm{NaClO_3}) = \frac{0.300}{3} = 0.100\ \mathrm{mol} \qquad m = 0.100 \times 106.44 = 10.6\ \mathrm{g}$

Warning

Common mistakes. (1) Saying "bromine displaces chlorine from chloride": a halogen displaces only those below it. (2) Calling the brown colour of iodine in water "violet": iodine is brown in water and violet in hexane. (3) Writing $\mathrm{FeI_3}$ for iron and iodine: iodine cannot oxidise iron to +3. (4) Forgetting that both products in chlorine + alkali contain chlorine in different oxidation states (disproportionation).

Safety. Use chlorine water and bromine water only in small amounts in a fume cupboard; bromine burns skin and its vapour is toxic. Hexane is highly flammable: no flames nearby. Never mix household bleach with acids (such as some toilet cleaners): the mixture releases chlorine gas. If chlorine is breathed in, move the person to fresh air and get medical help.

Tip

"A halogen higher up pushes out the one below." F₂ > Cl₂ > Br₂ > I₂ as oxidising agents (2.87, 1.36, 1.07, 0.54 V). In alkali, chlorine splits both ways: cold → Cl⁻ + ClO⁻ (bleach); hot → Cl⁻ + ClO₃⁻.

Royal Society of Chemistry: Halogen Displacement Reactions, 14–16 Practicals (7 min). Watch which halogen solutions change colour when added to which halide solutions, and how the results give the order of reactivity.

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Practice

Arrange the halogens in order of oxidising power and give two reasons for the order.
Predict what you would see when bromine water is added to (a) potassium iodide solution; (b) potassium chloride solution. Write an ionic equation where a reaction occurs.
Explain why iron reacts with chlorine to give $\mathrm{FeCl_3}$ but with iodine to give $\mathrm{FeI_2}$ .
Write equations for the reactions of fluorine and of chlorine with water, and state the change in oxidation number of the halogen in each.
Describe what happens to damp blue litmus paper in chlorine, and explain.
Write equations for the reaction of chlorine with cold dilute and with hot concentrated sodium hydroxide. Which is used to make bleach?
Use $E^\circ$ values to decide whether bromine will oxidise iodide ions, and calculate $E^\circ_{\text{cell}}$ .
What volume of chlorine, at room temperature and pressure, is needed to make 7.45 g of sodium chlorate(I), NaClO, with cold sodium hydroxide? $M(\mathrm{NaClO}) = 74.44$ ; 24.0 dm³ mol⁻¹.

Show answersHide answers

$\mathrm{F_2 > Cl_2 > Br_2 > I_2}$ . Smaller atoms attract an incoming electron more strongly (less distance and shielding); smaller halide ions are more strongly hydrated; (F₂ also has a weak bond).
(a) Brown solution (violet in hexane): $\mathrm{Br_2 + 2I^- \rightarrow 2Br^- + I_2}$ . (b) No reaction; the orange colour of bromine water remains.
Chlorine ( $E^\circ = 1.36\ \mathrm{V}$ ) is a strong enough oxidising agent to oxidise iron to +3 (Fe³⁺/Fe²⁺ 0.77 V); iodine (0.54 V) is not, so it stops at +2.
$\mathrm{2F_2 + 2H_2O \rightarrow 4HF + O_2}$ : F from 0 to −1. $\mathrm{Cl_2 + H_2O \rightleftharpoons HCl + HOCl}$ : Cl from 0 to −1 and to +1.
It turns red (HCl and HOCl are acids), then white, as HOCl bleaches the dye.
Cold: $\mathrm{Cl_2 + 2NaOH \rightarrow NaCl + NaClO + H_2O}$ (bleach, NaClO); hot: $\mathrm{3Cl_2 + 6NaOH \rightarrow 5NaCl + NaClO_3 + 3H_2O}$ .
$E^\circ_{\text{cell}} = 1.07 - 0.54 = +0.53\ \mathrm{V}$ : positive, so bromine oxidises iodide.
$n(\mathrm{NaClO}) = \frac{7.45}{74.44} = 0.100$ mol = $n(\mathrm{Cl_2})$ ; $V = 0.100 \times 24.0 = 2.40\ \mathrm{dm^3}$ .

Quick check

Which mixture produces a reaction?

Key points

Oxidising power: $\mathrm{F_2 > Cl_2 > Br_2 > I_2}$ ( $E^\circ$ = 2.87, 1.36, 1.07, 0.54 V), because smaller atoms attract electrons more strongly and smaller ions are better hydrated.
A halogen displaces the halogens below it from their halides (colours clearest in hexane: $\mathrm{Br_2}$ orange-red, $\mathrm{I_2}$ violet).
Chlorine oxidises Fe to $\mathrm{FeCl_3}$ ; iodine only to $\mathrm{FeI_2}$ . $\mathrm{H_2}$ + $\mathrm{X_2}$ reactions slow down from F to I.
$\mathrm{F_2}$ oxidises water to $\mathrm{O_2}$ ; $\mathrm{Cl_2 + H_2O \rightleftharpoons HCl + HOCl}$ . In NaOH chlorine disproportionates: cold → $\mathrm{Cl^- + ClO^-}$ (bleach); hot → $\mathrm{Cl^- + ClO_3^-}$ .

Sources and credits

Image: Brominewater.jpg by Madowiki, via Wikimedia Commons, licensed under CC BY-SA 3.0.
Image: Iron(III) chloride anhydrate.jpg by Leiem, via Wikimedia Commons, licensed under CC BY-SA 4.0.
Standard electrode potentials from Wikipedia, Standard electrode potential (data page), CC BY-SA 4.0, adapted by Ganzaa.
First aid for chemical exposure: NIOSH Pocket Guide, first aid, CDC.
Video: Halogen Displacement Reactions, 14–16 Practicals, Royal Society of Chemistry, embedded from YouTube.
The text, chart, table, 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.

1Electron thieves

2The order of oxidising power

3Displacement reactions

4Reactions with metals

5Reactions with hydrogen

6Reactions with water

7Reactions with alkalis