241

11 Group 17

The physical and chemical properties of the elements of Group 17 (the halogens) are
introduced in this topic.

11.1 Physical properties of the Group 17 elements

11.2 The chemical properties of the elements and their hydrides

11.3 Some reactions of the halide ions

11.4 The reactions of chlorine with aqueous sodium hydroxide

11.5 Some important uses of the halogens and of halogen compounds 


GROUP 17

CEDAR COLLEGE GROUP 17

 Cambridge International AS and A Level Chemistry 9701 syllabus Syllabus content

242

11 Group 17
The physical and chemical properties of the elements of Group 17 (the halogens) are introduced in this
topic.

Learning outcomes
Candidates should be able to:

11.1 Physical properties a) describe the colours and the trend in volatility of chlorine, bromine and
of the Group 17 iodine
elements b) interpret the volatility of the elements in terms of van der Waals’ forces

11.2 The chemical a) describe the relative reactivity of the elements as oxidising agents (see
properties of the also Section 6.3(f))
elements and their b) describe and explain the reactions of the elements with hydrogen
hydrides
c) (i) describe and explain the relative thermal stabilities of the hydrides
(ii) interpret these relative stabilities in terms of bond energies

11.3 Some reactions of a) describe and explain the reactions of halide ions with:
the halide ions (i) aqueous silver ions followed by aqueous ammonia
(ii) concentrated sulfuric acid

11.4 The reactions a) describe and interpret, in terms of changes of oxidation number, the
of chlorine with reaction of chlorine with cold and with hot aqueous sodium hydroxide
aqueous sodium
hydroxide

11.5 Some important a) explain the use of chlorine in water purification

uses of the b) state the industrial importance and environmental significance of the
halogens and halogens and their compounds (e.g. for bleaches, PVC, halogenated
of halogen hydrocarbons as solvents, refrigerants and in aerosols. See also
compounds Section 16.2)

Back to contents page www.cie.org.uk/alevel 31

CEDAR COLLEGE GROUP 17

1 Group 17 243

INTRODUCTION
dy the trends
the Group 17
Learning outcomes
The Halogens consist of the following elements:
ens. They are By the end of this topic you should be able to:
w decreasing • 11.1a)
Fluorine (F), the
describe Chlorine (Cl),the
colours and Bromine (Br), Iodine
trend in volatility (I) andbromine
of chlorine, Astatine
and (At).
asing proton iodine
All the halogens
11.1b) interpret consist of diatomic
the volatility molecules,
of the elements in terms of , linked
X2van by aforces
der Waals’ single
11.2a) bond.
covalent describe the relative
They are allreactivity of the elements as oxidising agents (see also
volatile.
Topic 23)
11.2b) describe
Fluorine (F)and explain
1s2 2sthe reactions
2 2p 5 of the elements with hydrogen
11.2c) describe and explain the relative thermal stabilities of the hydrides, and interpret
these relative
Chlorine (Cl) stabilities
1s2 2s2 in
2pterms of bond
6 3s2 3p 5 energies
11.3a) describe and explain the reactions of halide ions with aqueous silver ions
Bromine
followed(Br) 1s2 2sammonia,
by aqueous
2 2p6 3s2 3p6 3d10 4s2 4p5
and with concentrated sulfuric acid
11.4a) describe and interpret, in terms of changes of oxidation number, the reaction of
Iodine (I) 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p5
chlorine with cold, and with hot, aqueous sodium hydroxide
11.5a) explain the use of chlorine in water
1 purification
11.5b) state the industrial importance and environmental significance of the halogens
and their compounds (e.g. for bleaches, PVC, halogenated hydrocarbons as
solvents, refrigerants and in aerosols) (see also Topic 15).

11.1 Physical properties of the elements

PHYSICAL PROPERTIES
cal properties of the Halogen Chlorine Bromine Iodine
he importance of the
tials is discussed in Melting point /∞C −101 −7 114
Boiling point /oC −34 59 184
Appearance at room yellow-green gas red-brown liquid nearly black solid
temperature
E(X¬X)/kJ mol-1 242 193 151
-1
E(H¬X)/kJ mol 431 366 299
Standard electrode +1.36 +1.07 +0.54
The potential
van der Waals’
/V
forces between the molecules arise from the attractions between
temporary and induced dipoles. The greater the number of electrons in a molecule,
the more polarisable it is and the greater the possibility for temporary induced dipoles.
ThisMelting
explains why points and boiling
the boiling temperatures points
increase down the group of halogens.
The halogens exist as diatomic molecules,2 X2. The two atoms are bonded together
covalently in the solid, liquid and gaseous states. On melting, the van der Waals’
forces between the molecules are partially overcome, and on boiling they are almost
completely broken. On descending Group 17, the melting points and boiling points
increase. This is because the intermolecular forces become stronger as the number of
easily polarisable electrons rises (see section 3.17).
CEDAR COLLEGE GROUP 17
 yrtsimehC leveL SA lanoitanretnI egdirbmaC These are reactions betw
containing halide ions an
small amount of a solutio
244 SNOITCAER TNEMECA LPSID a solution containing a h
n o i t u l o s n i s e l u c e l o m n e g o l a h e h t f o s r u o l oc e h T gnisidixo eht ro( ytivitcaer eht egduj osla nac eW
(see Table 4.2). Potassium
eseht ni ylevitisop y fitnedi ot tluci ffid era tnemecalpsid rieht ta gnikool yb snegolah eht fo )rewop colourless. The colours o
snegolah eht ,revewoH .snoit caer tnemecalpsid .noitulos ni snoi edilah rehto htiw snoitcaer
shown in Figure 4.18.
ni elbicsimmi si hcihw( enaxeholcy c ni llew evlossid
snegolah ehT .)sreyal etarapes ow t gnimrof ,retaw
ssel a ecalpsid nac negolah evitcaer erom A
tnereffid ylt cnitsid era enaxeholcy c ni devlossid ssel eht fo noitulos edilah a morf negolah evitcaer
COLOURS OF SOLUTIONS
c i n a g r o s i h t f o e m o s g n i d d a e r o f e r e h T . s r u o l oc Orange colour, thedue to
.negolah evitcaer
,snoitulos edilah/negolah eht gnixim refta tnevlos production of bromine.
e l t t e s ot t i g n i w o l l a n e h t d n a e r u t x i m e h t g n i k a h s
Halogens in a non-polar
tneserp si negolah hcihw ylraelc swohs sreyal otni Chlorine solution is pale yellow-
,)qa(2lC ,retaw enirolhc nehW .elpmaxe na ta kool s’teL
.)3 .21 erugiF( selucelom cimotaid s ti sa gniniatnoc ,edimorb muidos fo noitulos a ot dedda si
organic solvent. green, bromine solution
hsiwolley a ot sRed-brown theis orange &
egnahc noitulos colour, a(−rBto
eht ,snoi )qdue
ecneserp eht yb desuac si ruoloc ehT .ruoloc nworb
ni dnuof sa ,)qa(2rBproduction oforbiodine.
iodine solution is red brown.
,selucelom enim devlossid fo
sekat taht noitcaer tnemecalpsid ehT .retaw enimorb
:si ecalp
)qa(2rB + )qa(lCaN2 )qa(rBaN2 + )qa(2lC The reactions that occur
enimorb decalpsid sah enirolhc taht yas eW
.noitulos morf Cl2(aq) + 2KBr(a
siht rof noitauqe cinoi eht ni desirammus si sihT Ionic equation: Cl2(aq) +
:noitcaer tnemecalpsid 471

)qa(2rB + )qa(−lC2 )qa(−rB2 + )qa(2lC Cl2(aq) + 2KI(aq

naht evitagenortcele erom era smota enirolhc ehT Ionic equation: Cl2(aq) +
mrof ot ycnednet regnorts a evah os smota enimorb
.snoi degrahc ylevitagen
Br2(aq) + 2KI(aq
na morf enidoi ecalpsid lliw enimorb ,esiwekiL
:noitulos edidoi
Ionic equation: Br2(aq) +
)qa(2I + )qa(rBaN2 )qa(IaN2 + )qa(2rB The more reactive haloge
.retaw fo pot no reyal a smrof enaxeholcyC 3.21 erugiF 3 :noitauqe cinoi na sa ro
,reyal cinagro reppu siht ni elprup si enidoi devlossiD
halogen from solution. Th
.neerg elap y rev si enirolhc dna egnaro si enimorb )qa(2I + )qa(−rB2 )qa(−I2 + )qa(2rB ions from solution, and b

negordyh htiw snoitcaeR NOITSEUQ

negordyh htiw sedil ah negordyh mrof snegol ah eTh
negordyh fo noitcaer eht si siht fo elpmaxe enO .sag muissatop fo noitulos a htiw dexim si retaw enirolhC 2
.) 371 egap ees( edirolhc negordyh evig ot enirolhc htiw dedda si enaxeholcyc fo 3mc 2 .ebut tset a ni edidoi
neht ,nekahs dna dereppots si hcihw ,ebut tset eht ot
snoitcaer rieht yb detartsulli si y tivitcaer ni dnert eTh
.dnats ot dewolla
.) 3.21 elbaT( Relate trends in the properties
etats gnidulcni ,noitauqe lobmys decnalab a etirW a
.srucco tahof
t noithe
tcaer oxides
eht rof ,sloof
bmyperiod
s 3
.noitcaer siht elements
rof noitauqe cito
noi ntheir
a etirW structure
b and
fo dne eht ta reyal enaxeholcyc eht si ruoloc tahW c Table 4.3 shows the form
bonding
HALOGENS AS OXIDANTS ?tnemirepxe eht
Describe the reactions of period
period 3 elements.
3 oxides with water
The halogens are powerful oxidising agents. Apart from fluorine, chlorine
These all have giant ionic
is the strongest oxidising agent among these elements. structure by strong electr
melting points. MgO and
The tendency to form negative ions decreases down the group. The both have melting points
ionic liquids are formed.
shielding effect of the inner electrons means that the effective nuclear and they all conduct elec
charge that attracts the outer electron is 7+ for all of the atoms.

Down the group the outer electrons get further and further away from the
same effective nuclear charge and so they are held less strongly and the
tendency to gain an extra electron decreases. Hence reactivity of
halogens decreases down the group.
4

CEDAR COLLEGE GROUP 17

 245

HALOGENS WITH HYDROGEN

All the halogens can react with hydrogen, but the conditions for the reaction vary
going down the group.

The reactivity with hydrogen decreases down the group because the decreasing
strength of the H—X (HCl, HBr, HI) bond makes the reaction less exothermic.

Halogen Reaction Condition Equation H—X Bond Energy

explodes when exposed to
Chlorine H2 (g) + Cl2 (g) ⟶ 2HCl (g) 431 kJ/mol
ultraviolet light

Bromine slow reaction when heated H2 (g) + Br2 (g) ⟶ 2HBr (g) 366 kJ/mol

reversible reaction when

Iodine H2 (g) + I2 (g) ⇌ 2HI (g) 299 kJ/mol
heated
5

HYDROGEN HALIDES
The hydrogen halides are all colourless gases
Halogen Bond Energy
with the formula HX, where X stands for Cl, Br or
I. The bonds between hydrogen and the H—Cl 431 kJ/mol
halogens are polar. Hydrogen chloride, hydrogen
bromide and hydrogen iodide are similar in that
H—Br 366 kJ/mol
they are very soluble in water to form acidic
solutions that ionise completely in water.
H—I 299 kJ/mol

CEDAR COLLEGE GROUP 17

 246

HYDROGEN HALIDES
The bond energies decrease going down Group Halogen Bond Energy
17, making it easier to break the H—X bond. This is
because the iodine atom is the largest atom, so H—Cl 431 kJ/mol

the overlap of its outer shell with a hydrogen atom

gives a much longer bond length than with the H—Br 366 kJ/mol
other smaller halogen atoms. The longer the
bond, the weaker it is, and the less energy H—I 299 kJ/mol
required to break it. Hence HI is less thermally
stable than HCl.

Worked example 7

Use the bond energies given in Table 11.1, and a value of 436 kJ mol −1 for the H¬H bond
energy in H2(g), to calculate the enthalpy change for the reaction between hydrogen and
chlorine, and between hydrogen and iodine.
on for the
and bromine Answer
e trend For H2 + Cl2 → 2HCl, ∆H = 436 + 242 − 2 × 431 = −184 kJ mol -1
For H2 + I2 → 2HI, ∆H = 436 + 151 − 2 × 299 = −11 kJ mol -1

Thermal stability of the hydrogen

HYDROGEN halides
HALIDES
The strength of the H¬X bonds also determines the thermal stability of the hydrogen
The strength
halides. Thisofisthe H—Xillustrated
readily bonds determines the thermalthat
by the observations stability of the
are made hydrogen
when halides.
they are
This is readily
heated illustratedmethod
(a convenient by the observations thatisare
in the laboratory to made
insert when they are
a hot metal heated.
wire into a test
tube of the gas). All the hydrogen halides are colourless gases. Hydrogen chloride
Hydrogen chloride does not decompose on heating. Hydrogen bromide decomposes
does not decompose on heating, so the test tube remains colourless; hydrogen
slightly
bromideanddecomposes
the gas turns brown.
slightly andHydrogen iodide
the gas turns decomposes
brown; hydrogenalmost
iodidecompletely
decomposesand
almost
the completely
gas turns purple.and the gas turns purple (see Table 11.4).

the hydrogen Halogen Reaction (if any) Comment

chlorine HCl(g) no decomposition
bromine 2HBr(g) H2(g) + Br2(g) some decomposition and gas appears
slightly brown
iodine 2HI(g) → H2(g) + I2(g) almost complete decomposition and gas
appears deep purple

Reactions of chlorine with sodium hydroxide

When chlorine gas is passed into cold, dilute aqueous sodium hydroxide, the
following reaction takes place and two salts are formed:

Cl2(g) + 2NaOH(aq) → NaCl(aq) + H2O(l) + NaClO(aq)

CEDARThe
COLLEGE
ionic equation is: GROUP 17

Cl2(g) + 2OH−(aq) → Cl−(aq) + H2O(l) + ClO−(aq)

 247

HALIDES
Halide ions are the ions of the halogen elements in oxidation state −1. They include the
chloride, Cl−, bromide, Br−, and iodide, I− , ions.

In Group 17, a more reactive halogen oxidises the ions of a less reactive halogen. And
as halogens are oxidising agents, a halogen that is a more powerful oxidising agent
displaces one that is less powerful. For example, bromine reacts with a solution of an
iodide to produce iodine and a bromide. This is because bromine has a stronger
tendency to gain electrons and turn into ions than iodine.

Br2 (aq) + 2I− (aq) ⟶ 2Br−(aq) + I2 (s)

But iodine will not displace bromide or chloride ions as iodine is the weakest of the
three oxidising agents.
9

HALIDE REACTIONS WITH SILVER IONS

Silver nitrate solution can be used to distinguish between halides. Silver halides
are insoluble. Adding silver nitrate to a solution of one of these halide ions
produces a precipitate.
Ag+ (aq) + X−(aq) ⟶ AgX (s)

Silver chloride is a white solid. This distinguishes it from silver bromide, which is a
creamy colour, and silver iodide, which is a brighter yellow.
The colour changes are not very distinct, but a further test with ammonia helps to
distinguish the precipitates. Silver chloride easily dissolves in dilute ammonia
solution. Silver bromide dissolves in concentrated ammonia solution, but silver
iodide does not dissolve in ammonia solution at all.
10

CEDAR COLLEGE GROUP 17

 in Table 11.5.

of the halide ions Halide ion Appearance and formula Solubility of precipitate in NH3(aq)
248
of precipitate
Cl− white, AgCl soluble in dilute NH3(aq)
−
Br cream, AgBr soluble in concentrated NH3(aq)
−
I yellow, AgI insoluble in concentrated NH3(aq)
HALIDE REACTIONS WITH SILVER IONS
f the silver halides
de ions.

AgCl AgBr AgI

Worked example 11

Write fully balanced, and ionic, equations for the reactions between:
a HCl and AgNO3
b KI and AgNO3

Answer
onic, equations a HCl + AgNO3 → AgCl + HNO3
aqueous solutions Cl−(aq) + AgCambridge
+
(aq) → AgCl(s)International AS Level Chemistry
MgBr2, and silver b KI + AgNO3 → AgI + KNO3
I−(aq) + Ag+(aq) → AgI(s)

HALIDE REACTIONS WITH SILVER IONS

TESTING FOR HALIDE IONS (CONTINUED)
The solubilities in water of the precipitates AgCl, AgBr and AgI decrease from AgCl to
AgI. This explains why they react differently with aqueous ammonia (see The general
Table 11.5). equation for the precipitatio
+
Aqueous silver ions react with ammonia to form the ion Ag(NH3)2 . By with Le silver nitrate solution is:
Chatelier’s Principle, in order to dissolve the silver halide precipitate, enough
− AgX(s) + NO3−(
ammonia has to be present to drive the following two equilibria over to theAgNO3(aq) + X (aq)
right-hand side: where X− represents the halide ion.
The aqueous nitrate ions can be left out
AgX(s) Ag+(aq) + X −(aq)
the ionic equation; they are spectator ion
Ag+(aq) + 2NH3(aq) Ag(NH3)2+(aq)
not get involved in the reaction:
If AgX is very insoluble, the first equilibrium lies well over to the left-hand side, and
Ag+(aq) + X−(aq)
so a large concentration of NH3(aq) is needed to pull the Ag+ ions into solution. AgX(s)
addition of
) to dissolve. Silver iodide is so insoluble that even in concentrated ammonia solution, Thetheadded ammonia can form complex i
concentration of ammonia is not high enough to pull the Ag+ ions into solution.
soluble:For
AgCl AgBr AgI
Figure 12.4 Colours of the silver halide precipitates: silver
■ silver chloride forms complex ions with dil
chloride (on the left), silver bromide and silver iodide (on 205
the right). ■ silver bromide forms complex ions with
12
concentrated ammonia.

Halide ion Colour of silver halide Effect on precipitate of Effect on precipit

29/09/14 8:39 PM
precipitate on addition of adding dilute ammonia adding concentra
silver nitrate solution solution ammonia solutio
chloride, Cl−(aq) white dissolves dissolves
bromide, Br−(aq) cream remains insoluble dissolves
176
CEDAR COLLEGE GROUP 17
iodide, I−(aq) pale yellow remains insoluble remains insoluble
Table 12.5 The results of positive tests for halide ions.
 249

HALIDE REACTIONS WITH SILVER IONS

Halide AgNO3 (aq) dilute NH3 (aq) concentrated NH3 (aq)

Chloride white ppt. Soluble -

Bromide cream ppt. Insoluble Soluble

Iodide pale yellow ppt. Insoluble Insoluble

The precipitates are due to the formation of silver halides.

Ag+ (aq) + CI− (aq) ⟶ AgCl (s)

A diammine complex, [Ag(NH3)2]+ is formed, which dissolves the precipitate.

AgCl (s) + 2NH3 (aq) ⟶ [Ag(NH3)2]+ (aq) + CI− (aq)
13

HALIDES AS REDUCTANTS
Halogen anions (halides) are powerful reducing agents.

This is due to halides requiring to lose only one electron to form halogens.

Down the group, reducing power increases. Hence, Fluoride ion is the
weakest reducing agent.

This is due to the fact that down a group, an increased ionic radius means
an electron can easily be removed from the halide ion.

14

CEDAR COLLEGE GROUP 17

 250

REDOX REACTIONS
The forward reaction is reduction of halogens, and
the backward reaction is the oxidation of halides.

Chlorine is completely reduced to chloride ions, Cl2 + 2e ⇌ 2Cl−

and iodide ions are completely oxidised to iodine.
Br2 + 2e ⇌ 2Br−
Since oxidation power is based on the tendency to
gain electrons, and reduction power is based on I2 + 2e ⇌ 2I−
the tendency to lose electrons:

a strong oxidising halogen has a weaker reducing

halide, and vice versa.
15

REACTIONS WITH CONC. SULFURIC ACID

Halides can also be distinguished from each other by reacting them with
concentrated sulfuric acid:

• Add concentrated sulfuric acid to a solid halide.

• Hydrogen halides are displaced from their salts, all of which give white fumes in
air.

• Down the group halides are strong reducing agents, they can further react to
reduce the sulfuric acid into many oxidation states.

• These reactions demonstrate that the strength of the halides as reducing agents
is in the order I— > Br— > Cl—
16

CEDAR COLLEGE GROUP 17

 251

REACTIONS WITH CONC. SULFURIC ACID

When concentrated sulfuric acid is added to a solid metal halide, the hydrogen
halide is produced.
With sodium chloride, steamy fumes of hydrogen chloride are evolved and sodium
hydrogensulfate is formed:

NaCl (s) + H2SO4 (aq) ⟶ NaHSO4 (aq) + HCl (g)

On heating the acid gives off two H+ and forms Na2SO4.

With sodium bromide and sodium iodide, similar reactions occur initially:

NaBr (s) + H2SO4 (aq) ⟶ NaHSO4 (aq) + HBr (g)

Nal (s) + H2SO4 (aq) ⟶ NaHSO4 (aq) + Hl (g)

17

REACTIONS WITH CONC. SULFURIC ACID

Concentrated sulfuric acid is a strong oxidising agent. It is unable to oxidise HCl, but
it can oxidise HBr and HI. Hence, only a little HBr is formed when concentrated
sulfuric acid is added to sodium bromide, the principal products being bromine and
sulfur dioxide

NaBr (s) + H2SO4 (aq) ⟶ NaHSO4 (aq) + HBr (g)

2HBr (g) + H2SO4 (aq) ⟶ 2H2O (l) + SO2 (g) + Br2 (g)

HI is so easily oxidised that only a trace of it is found. The sulfuric acid is reduced not
only to SO2 (+4), but further to sulfur (0), and even to hydrogen sulphide (—2).

18

CEDAR COLLEGE GROUP 17

 252

REACTIONS WITH CONC. SULFURIC ACID

Halide Observations Product Oxidation State of S Reaction type

Chloride White fumes HCl —1 Displacement of halide

White fumes HBr —1 Displacement of halide

Bromide Brown vapour Br2 0 Oxidation of halide

Colourless gas SO2 +4 Reduction of acid

White fumes HI —1 Displacement of halide

Purple vapour I2 0 Oxidation of halide

Iodide
yellow solid S8 0 Reduction of sulfur

H2S —2

19

DISPROPORTIONATION OF CHLORINE
Chlorine oxoanions form when chlorine reacts with water and alkalis. When chlorine
dissolves in water, it reacts reversibly to form a mixture of weak chloric(I) acid and
strong hydrochloric acid.
Cl2 + H2O ⇌ HCl + HOCl

This is an example of a reaction in which the same element both increases and
decreases its oxidation number.
In other words, some of the element is oxidised while the rest of the element is
reduced. Reactions of this type are called disproportionation reactions.
Bromine reacts in a similar way but to a much lesser extent. Iodine is almost insoluble
in water and hardly reacts at all.
20

CEDAR COLLEGE GROUP 17

 253

DISPROPORTIONATION OF CHLORINE
When chlorine dissolves in sodium (or potassium) hydroxide solution at room
temperature, it produces chlorate(I) and chloride ions.

2NaOH + Cl2 ⟶ NaCl + NaClO + H 2O

2OH— + Cl2 ⟶ Cl— + ClO— + H 2O

Chlorine itself, being an element, has the oxidation number 0. In the chloride ion, it
has an oxidation number of −1. In the ClO− ion, the chlorine has an oxidation
number of +1 in order to give an overall charge of −1 on the ion. This is another
example of disproportionation.

21

DISPROPORTIONATION OF CHLORINE
When chlorine is passed into hot, concentrated aqueous sodium
hydroxide, it produces sodium chloride and sodium chlorate(V).

6NaOH + 3Cl2 ⟶ 5NaCl + NaClO3 + 3H2O

6OH— + 3Cl2 ⟶ 5Cl— + ClO3— + 3H2O

22

CEDAR COLLEGE GROUP 17