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 4 of 5

Hydrogen Halides and Tests for Halide Ions

25 min

Same recipe, different results

Warm solid sodium chloride with concentrated sulphuric acid and steamy fumes of hydrogen chloride pour out of the tube. Do the same with sodium bromide and the fumes are tinged orange with bromine; with sodium iodide, purple clouds of iodine appear, the solid turns black and there is a smell of rotten eggs. The hydrogen halides are made in the same way, but the heavier ones are destroyed as they form. Why? This lesson compares HF, HCl, HBr and HI, and then gives the standard test that tells chloride, bromide and iodide ions apart.

Definition

Hydrogen halide: a compound HX of hydrogen with a halogen. In water it forms a hydrohalic acid, for example hydrochloric acid, $\mathrm{HCl(aq)}$ .

Thermal stability (of a hydrogen halide): its resistance to splitting into hydrogen and the halogen when heated, $\mathrm{2HX \rightarrow H_2 + X_2}$ .

Preparation

A hydrogen halide is driven out of its sodium salt by a less volatile acid. For HCl, concentrated sulphuric acid works:

\mathrm{NaCl(s) + H_2SO_4(conc) \rightarrow NaHSO_4(s) + HCl(g)}

Hydrogen fluoride is made the same way from fluorite, $\mathrm{CaF_2 + H_2SO_4 \rightarrow CaSO_4 + 2HF}$ , in plastic or metal vessels because it attacks glass. For HBr and HI, concentrated sulphuric acid fails: it is an oxidising agent, and bromide and iodide are reducing enough to be oxidised by it:

\mathrm{2HBr + H_2SO_4 \rightarrow Br_2 + SO_2 + 2H_2O} \qquad \mathrm{8HI + H_2SO_4 \rightarrow 4I_2 + H_2S + 4H_2O}

Iodide reduces sulphur all the way from +6 to −2 in $\mathrm{H_2S}$ (bromide only to +4 in $\mathrm{SO_2}$ ): the reducing power of the halide ions increases from $\mathrm{Cl^-}$ to $\mathrm{I^-}$ . To make HBr or HI, phosphoric acid, which is not an oxidising agent, is used instead: $\mathrm{NaI + H_3PO_4 \rightarrow NaH_2PO_4 + HI}$ .

Table 12.2: Concentrated sulphuric acid on solid halides

Halide	Observations	Products	Oxidation state reached by sulphur
NaCl	steamy fumes	HCl	stays +6
NaBr	steamy fumes; orange-brown vapour	HBr, $\mathrm{Br_2}$ , $\mathrm{SO_2}$	+4
NaI	some steamy fumes; purple vapour; black solid; yellow solid; rotten-egg smell	HI, $\mathrm{I_2}$ , $\mathrm{SO_2}$ , S, $\mathrm{H_2S}$	+4, 0 and −2

Properties of the hydrogen halides

All four are colourless gases at room temperature (HF boils at 19.5 °C) that fume in moist air and are very soluble in water.

Boiling points rise from HCl to HI as London forces grow, but HF boils far higher because its molecules are held by strong hydrogen bonds (Figure 12.5).
Thermal stability falls from HF to HI as the H–X bond gets longer and weaker (Figure 12.6): a red-hot wire plunged into HI gives violet iodine vapour, but HCl is unchanged.
Acid strength. HCl, HBr and HI are strong acids, fully ionised in water. HF is a weak acid ( $\mathrm{p}K_a = 3.17$ ), because the H–F bond is very strong and the small fluoride ion holds water molecules (and the proton) tightly. Even so, hydrofluoric acid dissolves glass: $\mathrm{SiO_2 + 4HF \rightarrow SiF_4 + 2H_2O}$ .
Reducing power increases from HCl to HI, for the same reason: the weaker H–X bond and the larger halide ion give up electrons more easily.
With ammonia, hydrogen chloride forms a white smoke of ammonium chloride: $\mathrm{NH_3(g) + HCl(g) \rightarrow NH_4Cl(s)}$ , a test for either gas.

Figure 12.5: Boiling points of the hydrogen halides.

Figure 12.6: H–X bond energies.

Tests for halide ions

Acidify the solution with dilute nitric acid (to remove carbonate and other ions that would also give precipitates), then add silver nitrate solution. The silver halide precipitates:

\mathrm{Ag^+(aq) + X^-(aq) \rightarrow AgX(s)}

Table 12.3: Silver halide precipitates

Ion	Precipitate with $\mathrm{AgNO_3(aq)}$	Dilute $\mathrm{NH_3(aq)}$	Concentrated $\mathrm{NH_3(aq)}$	In sunlight
$\mathrm{Cl^-}$	white, AgCl	dissolves	dissolves	turns purple-grey quickly
$\mathrm{Br^-}$	cream, AgBr	does not dissolve	dissolves	darkens slowly
$\mathrm{I^-}$	yellow, AgI	does not dissolve	does not dissolve	hardly changes
$\mathrm{F^-}$	none: AgF is soluble	–	–	–

The colours are close (white, cream, pale yellow), so the ammonia test confirms them. Ammonia dissolves a silver halide by forming the complex ion $\mathrm{[Ag(NH_3)_2]^+}$ , which works only for the more soluble halides: AgCl dissolves easily, AgBr needs concentrated ammonia, and AgI, the least soluble, does not dissolve at all. Nitric acid, not hydrochloric acid, is used to acidify, because hydrochloric acid would add chloride ions.

Worked example 1: identifying an unknown halide

A white solid X dissolves in water. Its solution, acidified with dilute nitric acid, gives a cream precipitate with silver nitrate, which does not dissolve in dilute ammonia but does dissolve in concentrated ammonia. Solid X, warmed with concentrated sulphuric acid, gives steamy fumes and an orange-brown vapour. Identify the anion in X and explain each observation.

Solution. The anion is bromide, $\mathrm{Br^-}$ .

Cream precipitate of AgBr: $\mathrm{Ag^+ + Br^- \rightarrow AgBr}$ ; soluble only in concentrated ammonia, as $\mathrm{[Ag(NH_3)_2]^+}$ .
With sulphuric acid: steamy fumes of HBr, $\mathrm{NaBr + H_2SO_4 \rightarrow NaHSO_4 + HBr}$ ; some HBr is oxidised to orange-brown bromine, $\mathrm{2HBr + H_2SO_4 \rightarrow Br_2 + SO_2 + 2H_2O}$ .

Worked example 2: chloride by mass

Excess acidified silver nitrate is added to 25.0 cm³ of a sodium chloride solution. The dried precipitate of silver chloride weighs 0.287 g. Calculate the concentration of chloride ions. $M(\mathrm{AgCl}) = 143.32\ \mathrm{g\,mol^{-1}}$ .

Solution. $n(\mathrm{AgCl}) = \frac{0.287}{143.32} = 2.00 \times 10^{-3}\ \mathrm{mol} = n(\mathrm{Cl^-}) \qquad c = \frac{2.00 \times 10^{-3}}{0.0250} = 0.0801\ \mathrm{mol\,dm^{-3}}$ The precipitate must be dried in the dark: light decomposes silver chloride and would change its mass.

Warning

Common mistakes. (1) Acidifying with hydrochloric acid before a halide test: it adds chloride. (2) Calling HF a strong acid because the other hydrogen halides are: HF is weak. (3) Explaining HF's high boiling point by its strong H–F bond: boiling breaks hydrogen bonds between molecules, not the H–F bond. (4) Trying to make HI with concentrated sulphuric acid: it oxidises HI; use phosphoric acid.

Safety. Hydrofluoric acid is extremely dangerous: it passes through skin, may cause little pain at first, and attacks bone and calcium in the blood; it is never used in schools. Hydrogen chloride and hydrogen bromide fumes attack the eyes and lungs, and hydrogen sulphide from iodides is very toxic: the sulphuric-acid tests are teacher demonstrations in a fume cupboard. Silver nitrate stains skin black. Rinse any splash with plenty of running water and get medical help.

Tip

Silver halides: white, cream, yellow — dilute NH₃, concentrated NH₃, nothing. Down the group the H–X bond weakens, so HX gets less stable, more reducing and (from HCl on) more acidic; HF is the odd one out: weak acid, high boiling point.

Sergei's Chemistry: Test for halide ions using silver nitrate, with ammonia to distinguish chloride, bromide and iodide (3 min). Watch the colours of the three silver halide precipitates and what happens to each when ammonia is added.

✦

Practice

Write the equation for making hydrogen chloride from sodium chloride and concentrated sulphuric acid.
Explain why concentrated sulphuric acid cannot be used to make hydrogen iodide, and name the acid that is used instead.
Describe what is seen when concentrated sulphuric acid is added to solid sodium iodide, and give the oxidation states reached by sulphur.
Explain why HF is a weak acid while HCl is a strong acid.
Explain the trend in thermal stability from HF to HI.
Describe how to show that a solution contains iodide ions rather than chloride ions, using silver nitrate and ammonia.
Why is dilute nitric acid, not hydrochloric acid, used in the halide test?
2.925 g of sodium chloride reacts completely with excess concentrated sulphuric acid. What volume of hydrogen chloride forms at room temperature and pressure? $M(\mathrm{NaCl}) = 58.44$ ; 24.0 dm³ mol⁻¹.

Show answersHide answers

$\mathrm{NaCl + H_2SO_4 \rightarrow NaHSO_4 + HCl}$ .
Concentrated sulphuric acid is an oxidising agent; it oxidises HI (a strong reducing agent) to iodine. Phosphoric acid, which is non-oxidising, is used.
Steamy fumes, purple iodine vapour, a black solid (iodine), yellow solid (sulphur), and a smell of bad eggs ( $\mathrm{H_2S}$ ). Sulphur goes from +6 to +4 ( $\mathrm{SO_2}$ ), 0 (S) and −2 ( $\mathrm{H_2S}$ ).
The H–F bond is very strong (574 kJ mol⁻¹) and the small $\mathrm{F^-}$ ion strongly holds the proton and water molecules, so HF ionises only slightly. The H–Cl bond is weaker, so HCl ionises fully.
Stability decreases (H–F 574, H–Cl 428, H–Br 363, H–I 294 kJ mol⁻¹): as the halogen atom gets bigger, the H–X bond gets longer and weaker, so less energy is needed to break it.
Acidify with dilute nitric acid and add silver nitrate: iodide gives a yellow precipitate that does not dissolve even in concentrated ammonia; chloride gives a white precipitate that dissolves in dilute ammonia.
Nitric acid removes interfering ions (such as carbonate) without adding halide ions; hydrochloric acid would add chloride and give a false positive.
$n(\mathrm{NaCl}) = \frac{2.925}{58.44} = 0.0500$ mol = $n(\mathrm{HCl})$ ; $V = 0.0500 \times 24.0 = 1.20\ \mathrm{dm^3}$ .

Quick check

An acidified solution gives a cream precipitate with silver nitrate that dissolves only in concentrated ammonia. Which ion is present?

Key points

HCl (and HF) are made from the salt and concentrated $\mathrm{H_2SO_4}$ ; HBr and HI need non-oxidising $\mathrm{H_3PO_4}$ , because $\mathrm{H_2SO_4}$ oxidises them (to $\mathrm{Br_2}$ , $\mathrm{I_2}$ , with $\mathrm{SO_2}$ , S and $\mathrm{H_2S}$ ).
Down the group the H–X bond weakens (574 → 294 kJ mol⁻¹): thermal stability falls, reducing power rises, and acid strength rises; HF is a weak acid with an anomalously high boiling point (hydrogen bonding).
Halide test: dilute $\mathrm{HNO_3}$ then $\mathrm{AgNO_3}$ : AgCl white (dissolves in dilute $\mathrm{NH_3}$ ), AgBr cream (concentrated $\mathrm{NH_3}$ ), AgI yellow (insoluble); fluoride gives no precipitate.

Sources and credits

Image: Ammonia smoke.JPG by Chemicalinterest, via Wikimedia Commons, public domain.
Image: AgBr Sample.jpg by Leiem, via Wikimedia Commons, licensed under CC BY-SA 4.0.
Data: boiling points and $\mathrm{p}K_a$ from the Wikipedia articles Hydrogen fluoride, Hydrogen chloride, Hydrogen bromide and Hydrogen iodide; H–X bond energies (from Greenwood and Earnshaw) from Halogen; all CC BY-SA 4.0, adapted by Ganzaa.
First aid for chemical splashes: NIOSH Pocket Guide, first aid, CDC.
Video: Test for halide ions using silver nitrate; ammonia to distinguish between chloride, bromide, iodide, Sergei's Chemistry, embedded from YouTube.
The text, charts, tables, 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.

1Same recipe, different results

2Preparation

3Properties of the hydrogen halides

4Tests for halide ions

Same recipe, different results

Preparation

Properties of the hydrogen halides

Tests for halide ions