Group 17 Elements: The Halogen Family

Q: What is bleaching powder and what is available chlorine?

Bleaching powder, roughly Ca(OCl)₂·CaCl₂·Ca(OH)₂·2H₂O, is made by passing chlorine over dry slaked lime. Excess dilute acid or carbon dioxide releases chlorine from it; the chlorine released, as a percentage of the sample's mass, is its available chlorine, usually 35 to 38 percent.

ChemistryThe p-Block Elements: Group 15, 16, 17 And 18For JEE aspirants

Group 17 elements, the halogens, are fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and the radioactive astatine (At), all with the valence configuration $n s^{2} n p^{5}$ , one electron short of an octet. The group 17 elements are the most reactive non-metals and strong oxidising agents, with oxidising power falling from $F_{2}$ to $I_{2}$ . This page covers their trends and key compounds: hydrogen halides, oxides and oxoacids of chlorine, bleaching powder and interhalogens. JEE Main and NEET test the trends and fluorine's unique behaviour; JEE Advanced also asks about the compounds.

On this page1At a glance2Trends3Anomalous F4Oxidising power5HX and oxides6Chlorine7F, Br, I8HCl9Oxoacids10Bleaching powder11Interhalogens12Solved examples

Key Formulas - Quick Reference

Valence shell $n s^{2} n p^{5}$ ; F shows only $- 1$ ; Cl, Br, I also show +1, +3, +5, +7 (and +4, +6 in oxides).
★ Must learn Electron gain enthalpy: $Cl > F > Br > I$ (most negative first); X-X bond enthalpy: $C l_{2} > B r_{2} > F_{2} > I_{2}$ .
★ Must learn Oxidising power: $F_{2} > C l_{2} > B r_{2} > I_{2}$ ( $E^{\circ}$ = +2.87, +1.36, +1.09, +0.54 V).
★ Must learn Acid strength: $HF ≪ HCl < HBr < HI$ ; HF is a liquid (hydrogen bonding) and etches glass: $Si O_{2} + 4HF \to Si F_{4} + 2 H_{2} O$ .
★ Must learn Chlorine in alkali: cold dilute $2NaOH + C l_{2} \to NaCl + NaOCl + H_{2} O$ ; hot conc. $6NaOH + 3C l_{2} \to 5NaCl + NaCl O_{3} + 3 H_{2} O$ .
Bleaching: $C l_{2} + H_{2} O \to HCl + HOCl$ , $HOCl \to HCl + [O]$ (permanent, by oxidation).
Oxoacids of chlorine: acid strength $HCl O_{4} > HCl O_{3} > HCl O_{2} > HOCl$ ; oxidising power in the reverse order.
Bleaching powder: $2Ca (OH)_{2} + 2C l_{2} \to Ca (OCl)_{2} + CaC l_{2} + 2 H_{2} O$ ; available chlorine is released by acids or $C O_{2}$ .
★ Must learn Interhalogens $X X_{n}^{'}$ ( $n$ = 1, 3, 5, 7): $Cl F_{3}$ T-shaped, $Br F_{5}$ square pyramidal, $I F_{7}$ pentagonal bipyramidal.
Polyhalide: $I_{2} + I^{-} ⇌ I_{3}^{-}$ (linear), which is why iodine dissolves in KI solution.

1. The Halogen Family at a Glance

Group 17 contains fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At); the synthetic element tennessine (Ts) also belongs here. They are called halogens, from the Greek hals (salt) and gennan (to produce), because they are literally salt formers. None of them occurs free in nature, because they are so reactive; astatine is radioactive and extremely rare.

Occurrence. Fluorine: fluorspar $Ca F_{2}$ , cryolite $N a_{3} Al F_{6}$ and fluorapatite $3C a_{3} (P O_{4})_{2} \cdot Ca F_{2}$ ; small amounts in soil, river water, bones and teeth. Chlorine: sodium chloride in sea water (about 2.5% by mass) and rock salt, with KCl, $MgC l_{2}$ and $CaC l_{2}$ . Bromine and iodine: bromides and iodides in sea water; iodine also in sea weeds and as sodium iodate $NaI O_{3}$ in Chile saltpetre.

Figure 1: Group 17 at a glance. Size grows and electronegativity falls down the group, but two values break the pattern at the top: chlorine, not fluorine, has the most negative electron gain enthalpy and the strongest X-X bond.

2. Trends in Physical Properties

Property	F	Cl	Br	I
Configuration	[He] $2 s^{2} 2 p^{5}$	[Ne] $3 s^{2} 3 p^{5}$	[Ar] $3 d^{10} 4 s^{2} 4 p^{5}$	[Kr] $4 d^{10} 5 s^{2} 5 p^{5}$
Covalent radius / pm	64	99	114	133
Ionic radius $X^{-}$ / pm	133	184	196	220
1st ionisation enthalpy / kJ mol $^{- 1}$	1680	1256	1142	1008
Electron gain enthalpy / kJ mol $^{- 1}$	$- 333$	$- 349$	$- 325$	$- 296$
Electronegativity	4.0	3.2	3.0	2.7
Hydration enthalpy of $X^{-}$ / kJ mol $^{- 1}$	$- 515$	$- 381$	$- 347$	$- 305$
X-X bond enthalpy / kJ mol $^{- 1}$	158.8	242.6	192.8	151.1
Melting / boiling point / K	54.4 / 84.9	172.0 / 239.0	265.8 / 332.5	386.6 / 458.2
Colour and state (298 K)	pale yellow gas	greenish-yellow gas	reddish-brown liquid	violet-black solid (violet vapour)

Atomic and ionic radii increase down the group; halogens have the smallest radii in their periods because of the high effective nuclear charge.
Ionisation enthalpies are very high, so halogens have little tendency to lose electrons; they decrease down the group.
Electron gain enthalpy is the most negative in each period. It becomes less negative down the group, except that fluorine's is less negative than chlorine's, because the added electron enters fluorine's small, crowded $2 p$ subshell.
Electronegativity decreases down the group; fluorine (4.0) is the most electronegative element.
Bond dissociation enthalpy falls from $C l_{2}$ to $I_{2}$ , but $F_{2}$ (158.8) is weaker than both $C l_{2}$ and $B r_{2}$ : the lone pairs on the two small F atoms repel strongly across the short F-F bond.
Colour: halogen molecules absorb visible light, which excites outer electrons to higher levels. $F_{2}$ absorbs high-energy violet light and looks yellow; $I_{2}$ absorbs lower-energy yellow light and looks violet.
State: $F_{2}$ and $C l_{2}$ are gases, $B r_{2}$ a liquid and $I_{2}$ a solid, because van der Waals forces grow with molecular size. Halogens dissolve sparingly in water but readily in organic solvents such as $CC l_{4}$ and $CHC l_{3}$ .

Exam Trick

Chlorine tops two lists. Most negative electron gain enthalpy: $Cl > F > Br > I$ . Strongest X-X bond: $C l_{2} > B r_{2} > F_{2} > I_{2}$ . In both lists fluorine is pushed down for the same reason: it is too small, so its electrons crowd each other.

3. Oxidation States and the Anomalous Behaviour of Fluorine

All the halogens show $- 1$ . Chlorine, bromine and iodine also show +1, +3, +5 and +7, using their vacant $d$ orbitals, mainly when bonded to the more electronegative fluorine and oxygen (in oxides, oxoacids and interhalogens); +4 and +6 occur in oxides such as $Cl O_{2}$ , $C l_{2} O_{6}$ , $Br O_{2}$ and $Br O_{3}$ (see Solved Example 5). Fluorine, the most electronegative element, shows only $- 1$ .

Why fluorine is different. Its small size, highest electronegativity, low F-F bond enthalpy and lack of $d$ orbitals give it unusual properties:

It is the most reactive element of the group and the strongest oxidising agent.
It shows only the $- 1$ state and forms only one oxoacid, hypofluorous acid HOF.
Its compounds are more ionic: metal fluorides are ionic, while most other halides are partly covalent.
HF is a liquid (b.p. 293 K) because of strong hydrogen bonding; the other hydrogen halides are gases.
AgF is soluble in water, while AgCl, AgBr and AgI are insoluble.
It cannot expand its octet, so it is never a central atom with more than one bond.

Positive halogen needs a stronger partner. A halogen shows a positive oxidation state only when it is bonded to something more electronegative: oxygen (oxides, oxoacids) or a lighter halogen (interhalogens). That is why fluorine, with nothing above it, is stuck at $- 1$ , and why the +7 state appears in $HCl O_{4}$ and $I F_{7}$ but never in a chloride.

Key idea

Fluorine is the odd one out: smallest, most electronegative, no $d$ orbitals. So it shows only

- 1

, forms the weak F-F bond, and is the strongest oxidising agent.

4. Chemical Reactivity of the Group

4.1 Reactivity and Oxidising Power

Fluorine is the most reactive halogen, and reactivity decreases down the group. The high reactivity of fluorine comes from its high electronegativity and low bond dissociation enthalpy; in general, halogens are reactive because the X-X bond is weak and the atoms have a strong affinity for electrons (Solved Example 4).

The halogens are strong oxidising agents because they readily accept electrons. Oxidising power decreases down the group, $F_{2} > C l_{2} > B r_{2} > I_{2}$ . Fluorine is the strongest even though its electron gain enthalpy is less negative than chlorine's; the full energy cycle (Figure 2) explains why.

Figure 2: Why

F_{2}

is a stronger oxidising agent than

C l_{2}

. The weak F-F bond and the large hydration enthalpy of the small

F^{-}

ion outweigh chlorine's slightly more negative electron gain enthalpy: the net change is

- 769

kJ for fluorine against

- 609

kJ for chlorine.

Each halogen therefore displaces (oxidises) the halide ions of the heavier halogens:

F_{2} + 2C l^{-} 2 F^{-} + C l_{2}

C l_{2} + 2B r^{-} 2C l^{-} + B r_{2}

C l_{2} + 2 I^{-} 2C l^{-} + I_{2}

B r_{2} + 2 I^{-} 2B r^{-} + I_{2}

With water the difference is striking. Fluorine oxidises water to oxygen (and some ozone); chlorine and bromine react to form hydrohalic and hypohalous acids; iodine does not oxidise water. Instead, oxygen oxidises iodide ions in acid:

2 F_{2} + 2 H_{2} O 4HF + O_{2}

3 F_{2} + 3 H_{2} O 6HF + O_{3}

X_{2} + H_{2} O X = Cl, Br HX + HOX

4 I^{-} + 4 H^{+} + O_{2} 2 I_{2} + 2 H_{2} O

Figure 3: Oxidising power of the halogens,

F_{2} > C l_{2} > B r_{2} > I_{2}

. A tick means the halogen in that row displaces the halide in that column; each halogen oxidises only the halides of the heavier halogens.

4.2 Hydrogen Halides (HX)

All the halogens combine with hydrogen to form volatile hydrides HX; their affinity for hydrogen decreases from fluorine (explosive even in the dark) to iodine (slow and reversible). HF is a liquid because of hydrogen bonding; HCl, HBr and HI are gases. All are covalent, and all act as acids in aqueous solution.

Property	HF	HCl	HBr	HI
Melting / boiling point / K	190 / 293	159 / 189	185 / 206	222 / 238
Bond length / pm	91.7	127.4	141.4	160.9
H-X bond enthalpy / kJ mol $^{- 1}$	574	432	363	295
$p K_{a}$	3.2	$- 7.0$	$- 9.5$	$- 10.0$

Acid strength increases $HF ≪ HCl < HBr < HI$ . The H-X bond becomes longer and weaker, so HI releases $H^{+}$ most easily. HF is a weak acid despite fluorine's electronegativity, because the H-F bond is very strong (574 kJ mol $^{- 1}$ ).
Reducing character increases in the same order, again because the H-X bond weakens; HI is a strong reducing agent and HF none.
Thermal stability decreases from HF to HI.

Hydrogen fluoride attacks glass and is stored in wax-lined or plastic bottles; this reaction is used to etch glass. HF also forms the hydrogen difluoride ion ( $KH F_{2}$ ) through hydrogen bonding:

Si O_{2} + 4HF Si F_{4} + 2 H_{2} O

Si F_{4} + 2HF H_{2} Si F_{6}

KF + HF KH F_{2}

Figure 4: Hydrogen halides. Hydrogen bonding links HF molecules into zig-zag chains, so HF is a liquid; the H-X bond weakens from HF to HI, so acid strength and reducing power increase in that order.

4.3 Halides

The halogens form compounds with almost all elements except He, Ne and Ar. These halides may be simple or complex.

With metals of low ionisation enthalpy (Na, K, Mg), they form ionic halides; ionic character decreases as the halogen becomes larger ( $MF > MCl > MBr > MI$ ).
With metals of higher ionisation enthalpy (Sn, Pb, Sb), the halides are largely covalent. A halide in a higher oxidation state is more covalent than one in a lower state: $SnC l_{4}, PbC l_{4}, SbC l_{5}$ are more covalent than $SnC l_{2}, PbC l_{2}, SbC l_{3}$ .
With non-metals (P, As, S) and with Al they form covalent halides such as $P F_{5}, AsC l_{3}, S_{2} C l_{2}$ and $AlC l_{3}$ .

4.4 Oxides

Halogens do not combine directly with oxygen, because both have high electron affinities, but their oxides can be made indirectly. Fluorine forms $O F_{2}$ and $O_{2} F_{2}$ ; since fluorine is more electronegative, these are oxygen fluorides. Only $O F_{2}$ is stable at 298 K; $O_{2} F_{2}$ oxidises plutonium to $Pu F_{6}$ , which removes plutonium from spent nuclear fuel.

Halogen	Oxides (oxidation state of the halogen)
F	$O F_{2}$ , $O_{2} F_{2}$ (oxygen fluorides)
Cl	$C l_{2} O$ (+1), $Cl O_{2}$ (+4), $C l_{2} O_{6}$ (+6), $C l_{2} O_{7}$ (+7)
Br	$B r_{2} O$ (+1), $Br O_{2}$ (+4), $Br O_{3}$ (+6)
I	$I_{2} O_{4}$ (+4), $I_{2} O_{5}$ (+5), $I_{2} O_{7}$ (+7)

Most halogen oxides are endothermic and unstable; they are powerful oxidising agents and may explode on heating or mechanical shock.
Stability: iodine oxides are the most stable, then chlorine, then bromine ( $I > Cl > Br$ ); the bromine oxides exist only at low temperatures. Higher oxides are more stable than lower ones.
$Cl O_{2}$ bleaches paper pulp and textiles and is used in water treatment. $I_{2} O_{5}$ is used to estimate carbon monoxide: $I_{2} O_{5} + 5CO \to I_{2} + 5C O_{2}$ .

Preparation of some oxides:

2 F_{2} + 2NaOH 2 % NaOH 2NaF + O F_{2} + H_{2} O

2C l_{2} + 2HgO 575 K HgO \cdot HgC l_{2} + C l_{2} O

2NaCl O_{3} + S O_{2} + H_{2} S O_{4} 2Cl O_{2} + 2NaHS O_{4}

2AgCl O_{3} + C l_{2} 363 K 2AgCl + 2Cl O_{2} + O_{2}

2NaCl O_{2} + C l_{2} 2NaCl + 2Cl O_{2}

2Cl O_{2} + 2 O_{3} C l_{2} O_{6} + 2 O_{2}

2HCl O_{4} + P_{2} O_{5} 263 K 2HP O_{3} + C l_{2} O_{7}

2HI O_{3} 513 K I_{2} O_{5} + H_{2} O

Reactions: the chlorine oxides are the anhydrides of the oxoacids. $Cl O_{2}$ and $C l_{2} O_{6}$ are 'mixed' anhydrides that disproportionate in water or alkali:

C l_{2} O + H_{2} O 2HOCl

C l_{2} O + 2KOH 2KOCl + H_{2} O

2Cl O_{2} + H_{2} O HCl O_{2} + HCl O_{3}

2Cl O_{2} + 2NaOH NaCl O_{2} + NaCl O_{3} + H_{2} O

C l_{2} O_{6} + H_{2} O HCl O_{3} + HCl O_{4}

C l_{2} O_{7} + H_{2} O 2HCl O_{4}

$Cl O_{2}$ has an odd electron and is paramagnetic; $C l_{2} O_{6}$ is molecular $Cl O_{3}$ (odd electron) in the vapour but ionic, $[Cl O_{2}]^{+} [Cl O_{4}]^{-}$ , in the solid.

Figure 5: Oxides of chlorine.

C l_{2} O

and

Cl O_{2}

are bent;

Cl O_{2}

has an odd electron and is paramagnetic. In

C l_{2} O_{7}

two

Cl O_{4}

tetrahedra share one oxygen.

Quick Recall: tap to check

Why is the electron gain enthalpy of F less negative than that of Cl?

The incoming electron enters the small, crowded 2p shell of F and is repelled by the electrons already there.

Can

B r_{2}

liberate

I_{2}

from KI?

Yes: a halogen oxidises the halides of the halogens below it.

Which hydrogen halide is a liquid at room temperature, and why?

HF (b.p. 293 K), because its molecules are joined by hydrogen bonds into zig-zag chains.

5. Chlorine ( $C l_{2}$ )

5.1 Preparation

In the laboratory, by oxidising hydrochloric acid with manganese dioxide or potassium permanganate, or by heating a chloride with $Mn O_{2}$ and concentrated sulphuric acid:

Mn O_{2} + 4HCl MnC l_{2} + 2 H_{2} O + C l_{2}

4NaCl + Mn O_{2} + 4 H_{2} S O_{4} MnC l_{2} + 4NaHS O_{4} + 2 H_{2} O + C l_{2}

2KMn O_{4} + 16HCl 2KCl + 2MnC l_{2} + 8 H_{2} O + 5C l_{2}

Industrially, (a) by the electrolysis of brine (aqueous NaCl), which gives chlorine at the anode (and NaOH and $H_{2}$ as by-products), or of molten NaCl; (b) by Deacon's process, oxidising HCl with air over a $CuC l_{2}$ catalyst:

4HCl + O_{2} CuC l_{2}, 723 K 2C l_{2} + 2 H_{2} O

5.2 Properties

Greenish-yellow gas with a pungent, suffocating smell; about 2.5 times heavier than air; poisonous. It liquefies easily to a greenish-yellow liquid (b.p. 239 K) and is fairly soluble in water.
It reacts with metals and non-metals to give chlorides, and with hydrogen explosively in sunlight.

H_{2} + C l_{2} sunlight 2HCl

2Na + C l_{2} 2NaCl

2Al + 3C l_{2} 2AlC l_{3}

2Fe + 3C l_{2} 2FeC l_{3}

P_{4} + 6C l_{2} 4PC l_{3}

2As + 3C l_{2} 2AsC l_{3}

2Sb + 3C l_{2} 2SbC l_{3}

S_{8} + 4C l_{2} 4 S_{2} C l_{2}

It has a great affinity for hydrogen, removing it even from hydrocarbons and $H_{2} S$ :

H_{2} S + C l_{2} 2HCl + S

C_{10} H_{16} + 8C l_{2} turpentine 16HCl + 10C

With ammonia the product depends on which is in excess (see Group 15):

8N H_{3} + 3C l_{2} excess N H_{3} 6N H_{4} Cl + N_{2}

N H_{3} + 3C l_{2} excess C l_{2} NC l_{3} + 3HCl

With alkalis chlorine disproportionates (0 to $- 1$ and +1, or $- 1$ and +5):

2NaOH + C l_{2} cold, dilute NaCl + NaOCl + H_{2} O

6NaOH + 3C l_{2} hot, conc. 5NaCl + NaCl O_{3} + 3 H_{2} O

2Ca (OH)_{2} + 2C l_{2} dry slaked lime Ca (OCl)_{2} + CaC l_{2} + 2 H_{2} O

With hydrocarbons, substitution or addition:

C H_{4} + C l_{2} UV light C H_{3} Cl + HCl

C_{2} H_{4} + C l_{2} 1,2-dichloroethane C_{2} H_{4} C l_{2}

Oxidising and bleaching action. With water, chlorine gives hypochlorous acid, which releases nascent oxygen; this oxidises many substances and bleaches coloured matter permanently:

C l_{2} + H_{2} O HCl + HOCl

HOCl HCl + [O]

2FeS O_{4} + H_{2} S O_{4} + C l_{2} F e_{2} (S O_{4})_{3} + 2HCl

N a_{2} S O_{3} + C l_{2} + H_{2} O N a_{2} S O_{4} + 2HCl

S O_{2} + 2 H_{2} O + C l_{2} H_{2} S O_{4} + 2HCl

I_{2} + 6 H_{2} O + 5C l_{2} 2HI O_{3} + 10HCl

Chlorine bleaches vegetable and organic matter only in the presence of moisture.

Figure 6: Chlorine reaction map. Chlorine combines with metals and hydrogen, oxidises ammonia (the product depends on which is in excess), substitutes into alkanes, disproportionates in alkali and bleaches through the nascent oxygen released by HOCl.

Exam Trick

Temperature decides the chlorine product in alkali. Cold and dilute: hypochlorite (Cl goes to +1). Hot and concentrated: chlorate (Cl goes to +5). Both are disproportionations, so each equation also produces chloride ( $- 1$ ): 1 : 1 for cold, 5 : 1 for hot.

Uses: bleaching wood pulp (for paper and rayon) and cotton and textiles; extracting gold and platinum; making dyes, drugs and organic compounds such as $CC l_{4}$ , $CHC l_{3}$ , DDT and refrigerants; sterilising drinking water; making poisonous gases such as phosgene ( $COC l_{2}$ ), tear gas ( $CC l_{3} N O_{2}$ ) and mustard gas ( $ClC H_{2} C H_{2} SC H_{2} C H_{2} Cl$ ).

Figure 7: Chlorine disproportionates in alkali. Cold, dilute alkali stops at hypochlorite (+1); hot, concentrated alkali goes on to chlorate (+5). In both, part of the chlorine is reduced to

C l^{-}

Key idea

Chlorine in alkali disproportionates: part goes down to

C l^{-}

(

- 1

) and part goes up, to

Cl O^{-}

(+1) when cold and dilute, to

Cl O_{3}^{-}

(+5) when hot and concentrated.

6. Fluorine, Bromine and Iodine

6.1 Fluorine

Fluorine cannot be made by chemical oxidation of fluorides in ordinary conditions, so it is made electrolytically. Moissan (1886) electrolysed a solution of KF in anhydrous HF in a platinum-iridium apparatus. Today a molten mixture of KF and HF (about 1 : 2) is electrolysed at about 350 K, with a mild steel cathode and a graphite-free carbon anode; the products are kept apart to avoid an explosive reaction between hydrogen and fluorine. KF is the electrolyte; HF only provides the ions.

2 H^{+} + 2 e^{-} cathode H_{2}

2 F^{-} anode F_{2} + 2 e^{-}

Fluorine is a pale yellow, very poisonous gas and the strongest known oxidising agent. Copper and nickel resist it because a protective fluoride layer forms (e.g. $Cu F_{2}$ ), which is why fluorine cells are made of copper, steel or Monel metal. Typical reactions:

2Al + 3 F_{2} 2Al F_{3}

I_{2} + 5 F_{2} 2I F_{5}

Si O_{2} + 2 F_{2} Si F_{4} + O_{2}

2 F_{2} + 4NaOH concentrated 4NaF + O_{2} + 2 H_{2} O

KCl O_{3} + F_{2} + H_{2} O KCl O_{4} + 2HF

2 K_{2} S O_{4} + F_{2} K_{2} S_{2} O_{8} + 2KF

6.2 Bromine

Bromine is made from sea water and natural brines by oxidising bromide with chlorine at a pH of about 3.5; air blows the bromine out of the solution. It is a dark red, dense liquid (b.p. 332 K) whose vapour attacks the skin, eyes and throat.

2B r^{-} + C l_{2} B r_{2} + 2C l^{-}

Bromine is a good oxidising agent, and its red colour is discharged by unsaturated compounds (a test for C=C bonds):

B r_{2} + H_{2} S 2HBr + S

N a_{2} S O_{3} + B r_{2} + H_{2} O N a_{2} S O_{4} + 2HBr

2HI + B r_{2} 2HBr + I_{2}

6FeS O_{4} + 3B r_{2} 2F e_{2} (S O_{4})_{3} + 2FeB r_{3}

KN O_{2} + B r_{2} + H_{2} O KN O_{3} + 2HBr

8N H_{3} + 3B r_{2} 6N H_{4} Br + N_{2}

C_{2} H_{4} + B r_{2} red to colourless C_{2} H_{4} B r_{2}

6.3 Iodine

Iodine is obtained from brines and sea weed by oxidising iodide with chlorine, and from sodium iodate (in Chile saltpetre) by reduction with sodium hydrogen sulphite; the iodide formed then reacts with more iodate:

2 I^{-} + C l_{2} I_{2} + 2C l^{-}

NaI O_{3} + 3NaHS O_{3} NaI + 3NaHS O_{4}

NaI O_{3} + 5NaI + 3 H_{2} S O_{4} 3 I_{2} + 3N a_{2} S O_{4} + 3 H_{2} O

Another route precipitates silver iodide from brine, converts it to iron(II) iodide with iron, and oxidises that with chlorine:

2AgI + Fe Fe I_{2} + 2Ag

Fe I_{2} + C l_{2} FeC l_{2} + I_{2}

Iodine is a violet-black, shiny solid that sublimes to a violet vapour. It is only sparingly soluble in water but dissolves in KI solution by forming the triiodide ion, and in $C S_{2}$ or $CHC l_{3}$ to give a violet solution. It turns starch blue.

I_{2} + I^{-} ⇌ I_{3}^{-}

7. Hydrogen Chloride (HCl)

Preparation: heating sodium chloride with concentrated sulphuric acid. At 420 K sodium hydrogen sulphate forms; at 823 K the rest of the chloride reacts. The gas is dried by passing it through concentrated $H_{2} S O_{4}$ .

NaCl + H_{2} S O_{4} 420 K NaHS O_{4} + HCl

NaHS O_{4} + NaCl 823 K N a_{2} S O_{4} + HCl

Colourless gas with a pungent smell; it liquefies to a colourless liquid (b.p. 189 K) and freezes to a white solid (159 K).
Extremely soluble in water and fully ionised ( $K_{a} \approx 1 0^{7}$ ): $HCl + H_{2} O \to H_{3} O^{+} + C l^{-}$ . It gives white fumes of $N H_{4} Cl$ with ammonia.
It decomposes salts of weaker acids, such as carbonates, hydrogen carbonates and sulphites.

N H_{3} + HCl N H_{4} Cl

N a_{2} C O_{3} + 2HCl 2NaCl + H_{2} O + C O_{2}

NaHC O_{3} + HCl NaCl + H_{2} O + C O_{2}

N a_{2} S O_{3} + 2HCl 2NaCl + H_{2} O + S O_{2}

Aqua regia (3 parts conc. HCl : 1 part conc. $HN O_{3}$ ) dissolves the noble metals gold and platinum; the chloride ions hold the metal ions as stable complexes:

Au + 4 H^{+} + N O_{3}^{-} + 4C l^{-} AuC l_{4}^{-} + NO + 2 H_{2} O

3Pt + 16 H^{+} + 4N O_{3}^{-} + 18C l^{-} 3PtC l_{6}^{2 -} + 4NO + 8 H_{2} O

Uses: making chlorine, $N H_{4} Cl$ and glucose (from corn starch); extracting glue from bones and purifying bone black; in medicine and as a laboratory reagent.

Cold, dilute NaOH

C l_{2} + 2NaOH \to NaCl + NaOCl + H_{2} O

: hypochlorite, Cl goes 0 to

- 1

and +1.

Hot, concentrated NaOH

3C l_{2} + 6NaOH \to 5NaCl + NaCl O_{3} + 3 H_{2} O

: chlorate, Cl goes 0 to

- 1

and +5.

8. Oxoacids of Halogens

Fluorine forms only one oxoacid, hypofluorous acid HOF, because of its high electronegativity and small size. Chlorine, bromine and iodine form four series, HOX, $HX O_{2}$ , $HX O_{3}$ and $HX O_{4}$ , though most are stable only in solution or as salts.

Series	F	Cl	Br	I	Salts
HOX (+1)	HOF	HOCl	HOBr	HOI	hypohalites
$HX O_{2}$ (+3)	-	$HCl O_{2}$	-	-	halites
$HX O_{3}$ (+5)	-	$HCl O_{3}$	$HBr O_{3}$	$HI O_{3}$	halates
$HX O_{4}$ (+7)	-	$HCl O_{4}$	$HBr O_{4}$	$HI O_{4}$	perhalates

Acid strength rises with the oxidation state of the halogen: $HCl O_{4} > HCl O_{3} > HCl O_{2} > HOCl$ (perhalic > halic > halous > hypohalous). Each extra O atom spreads the negative charge of the anion over more atoms by resonance, so the anion is more stable.
For the same series, acidity falls as the halogen gets larger and less electronegative: $HOCl > HOBr > HOI$ .
Hypohalous acids (HOCl, HOBr, HOI) are weak acids formed when the halogen disproportionates in water, $X_{2} + H_{2} O \to HOX + HX$ ; their salts are hypohalites, as in bleaching powder.
Halic acids: $HCl O_{3}$ and $HBr O_{3}$ exist only in solution, but iodic acid $HI O_{3}$ is a white solid, so stability increases with the size of the halogen. Salts are halates: $NaCl O_{3}$ is a powerful weed-killer and $KCl O_{3}$ is used in fireworks and matches.
Perhalic acids give perhalates.

Figure 8: Oxoacids of chlorine. From HOCl to

HCl O_{4}

each extra O spreads the negative charge of the anion over more atoms, so acid strength and anion stability rise while oxidising power falls.

8.1 Oxoacids of Chlorine in Detail

Acid	Preparation	Key properties	Important salt and use
HOCl (Cl +1)	$C l_{2}$ + $H_{2} O$ (disproportionation); shaking HgO with chlorine water; $C O_{2}$ into a bleaching powder suspension	weak acid, known only in solution; decomposes to HCl + $O_{2}$ ; strong oxidant and bleach (releases nascent O)	NaOCl: household bleach, disinfectant
$HCl O_{2}$ (Cl +3)	$Ba (Cl O_{2})_{2}$ + dil. $H_{2} S O_{4}$ ; the barium chlorite comes from $Ba (OH)_{2}$ + $H_{2} O_{2}$ + $Cl O_{2}$	unstable; decomposes to $Cl O_{2}$ ; disproportionates; liberates $I_{2}$ from KI	$NaCl O_{2}$ : bleaching textiles, removing NOx
$HCl O_{3}$ (Cl +5)	$Ba (Cl O_{3})_{2}$ + dil. $H_{2} S O_{4}$ ; chlorates from $C l_{2}$ + hot alkali	known only in solution; strong oxidant (sulphite to sulphate); decomposes giving $Cl O_{2}$	$KCl O_{3}$ : matches, fireworks, laboratory $O_{2}$
$HCl O_{4}$ (Cl +7)	$NaCl O_{4}$ + conc. HCl, or $KCl O_{4}$ + $H_{2} S O_{4}$ distilled under reduced pressure	one of the strongest acids; colourless oily liquid, fumes in air; explodes with organic matter; dehydrated by $P_{2} O_{5}$ to $C l_{2} O_{7}$	$N H_{4} Cl O_{4}$ : rocket propellant; $Mg (Cl O_{4})_{2}$ : desiccant (anhydrone)

Hypochlorous acid:

2HgO + 2C l_{2} + H_{2} O HgO \cdot HgC l_{2} + 2HOCl

2CaOC l_{2} + H_{2} O + C O_{2} CaC l_{2} + CaC O_{3} + 2HOCl

2HOCl 2HCl + O_{2}

Mg + 2HOCl Mg (OCl)_{2} + H_{2}

HOCl + 2F e^{2 +} + H^{+} C l^{-} + 2F e^{3 +} + H_{2} O

C l_{2} + 2O H^{-} making NaOCl C l^{-} + OC l^{-} + H_{2} O

Chlorous acid:

Ba (OH)_{2} + H_{2} O_{2} + 2Cl O_{2} Ba (Cl O_{2})_{2} + 2 H_{2} O + O_{2}

Ba (Cl O_{2})_{2} + H_{2} S O_{4} BaS O_{4} ↓ + 2HCl O_{2}

5HCl O_{2} 4Cl O_{2} + HCl + 2 H_{2} O

2HCl O_{2} HOCl + HCl O_{3}

4KI + HCl O_{2} + 2 H_{2} O 4KOH + HCl + 2 I_{2}

Chloric acid:

6Ba (OH)_{2} + 6C l_{2} 5BaC l_{2} + Ba (Cl O_{3})_{2} + 6 H_{2} O

Ba (Cl O_{3})_{2} + H_{2} S O_{4} BaS O_{4} ↓ + 2HCl O_{3}

4HCl O_{3} 4Cl O_{2} + 2 H_{2} O + O_{2}

Cl O_{3}^{-} + 3S O_{3}^{2 -} C l^{-} + 3S O_{4}^{2 -}

Perchloric acid:

NaCl O_{4} + HCl NaCl + HCl O_{4}

KCl O_{4} + H_{2} S O_{4} KHS O_{4} + HCl O_{4}

3HCl O_{3} HCl O_{4} + 2Cl O_{2} + H_{2} O

Zn + 2HCl O_{4} Zn (Cl O_{4})_{2} + H_{2}

4HCl O_{4} hot, 73 % 2C l_{2} + 7 O_{2} + 2 H_{2} O

NaCl O_{4} + N H_{4} Cl NaCl + N H_{4} Cl O_{4}

Exam Trick

More oxygen: stronger acid, weaker oxidant. Going from $Cl O^{-}$ to $Cl O_{4}^{-}$ the anion gains resonance forms and stability, so the acid gets stronger, but the anion becomes less eager to react, so oxidising power falls: $Cl O^{-} > Cl O_{2}^{-} > Cl O_{3}^{-} > Cl O_{4}^{-}$ .

9. Bleaching Powder

Bleaching powder is often written $CaOC l_{2}$ or $Ca (OCl) Cl$ , calcium chlorohypochlorite, a mixed salt of hydrochloric and hypochlorous acids. Its actual composition is closer to $Ca (OCl)_{2} \cdot CaC l_{2} \cdot Ca (OH)_{2} \cdot 2 H_{2} O$ . It is made by passing chlorine over dry slaked lime, industrially in Hasenclever's or Bachmann's plant:

Ca (OH)_{2} + C l_{2} Ca (OCl) Cl + H_{2} O

2Ca (OH)_{2} + 2C l_{2} Ca (OCl)_{2} + CaC l_{2} + 2 H_{2} O

A pale yellow powder with a strong smell of chlorine; it dissolves in water, but the solution is never clear because of impurities such as lime.
On heating it gives chlorate and chloride (autoxidation); with a cobalt chloride catalyst it releases oxygen.

6CaOC l_{2} Δ Ca (Cl O_{3})_{2} + 5CaC l_{2}

2CaOC l_{2} CoC l_{2} 2CaC l_{2} + O_{2}

Oxidising and bleaching action. A little dilute acid releases hypochlorous acid, whose nascent oxygen oxidises and bleaches:

2CaOC l_{2} + H_{2} S O_{4} limited dil. acid CaC l_{2} + CaS O_{4} + 2HOCl

CaOC l_{2} + H_{2} S CaC l_{2} + H_{2} O + S

CaOC l_{2} + 2FeS O_{4} + H_{2} S O_{4} F e_{2} (S O_{4})_{3} + CaC l_{2} + H_{2} O

CaOC l_{2} + KN O_{2} CaC l_{2} + KN O_{3}

3CaOC l_{2} + 2N H_{3} 3CaC l_{2} + 3 H_{2} O + N_{2}

Available chlorine. Excess dilute acid, or carbon dioxide of the air, releases all the chlorine. The chlorine obtained this way, as a percentage of the mass of the sample, is the available chlorine (35-38% for good commercial samples; the theoretical value for $CaOC l_{2}$ is about 56%):

CaOC l_{2} + 2HCl CaC l_{2} + H_{2} O + C l_{2}

CaOC l_{2} + H_{2} S O_{4} CaS O_{4} + H_{2} O + C l_{2}

CaOC l_{2} + C O_{2} CaC O_{3} + C l_{2}

Making chloroform. Bleaching powder in water supplies both chlorine and slaked lime; with acetone (or ethanol) it gives chloroform:

CaOC l_{2} + H_{2} O Ca (OH)_{2} + C l_{2}

C H_{3} COC H_{3} + 3C l_{2} CC l_{3} COC H_{3} + 3HCl

2CC l_{3} COC H_{3} + Ca (OH)_{2} (C H_{3} COO)_{2} Ca + 2CHC l_{3}

Uses: bleaching cotton and linen; disinfecting drinking water and swimming pools; making chloroform; making wool unshrinkable.

Key idea

More oxygen on chlorine means more resonance in the anion: acid strength rises (HOCl to

HCl O_{4}

) while oxidising power falls.

10. Interhalogen Compounds and Polyhalide Ions

Two different halogens combine to form interhalogen compounds $X X_{n}^{'}$ , where X is the larger, less electronegative halogen and $n$ = 1, 3, 5 or 7. They are made by direct combination under specific conditions:

Type	Examples	Shape (hybridisation of X)
$X X^{'}$	ClF, BrF, BrCl, ICl, IBr	linear
$X X^{'} 3$	$Cl F_{3}$ , $Br F_{3}$ , $I F_{3}$ , $IC l_{3}$	bent T-shape ( $s p^{3} d$ , 2 lone pairs)
$X X^{'} 5$	$Cl F_{5}$ , $Br F_{5}$ , $I F_{5}$	square pyramid ( $s p^{3} d^{2}$ , 1 lone pair)
$X X^{'} 7$	$I F_{7}$	pentagonal bipyramid ( $s p^{3} d^{3}$ )

C l_{2} + F_{2} 437 K equal volumes 2ClF

C l_{2} + 3 F_{2} 573 K excess F_{2} 2Cl F_{3}

Cl F_{3} + F_{2} hv Cl F_{5}

I_{2} + C l_{2} 2ICl

I_{2} + 3C l_{2} excess C l_{2} 2IC l_{3}

B r_{2} + 3 F_{2} 2Br F_{3}

B r_{2} + 5 F_{2} excess F_{2} 2Br F_{5}

They are covalent and diamagnetic, and mostly volatile liquids or solids at 298 K (ClF is a gas).
They are more reactive than the parent halogens (except fluorine), because the X-X' bond is weaker than the X-X bond in the halogen.
They are good oxidising agents. Their physical properties lie between those of the parent halogens, with melting and boiling points a little higher than expected because the bonds are polar.
They hydrolyse to a halide ion from the smaller halogen and a hypohalite (or oxohalide) from the larger one: $X X^{'} + H_{2} O \to H X^{'} + HOX$ .
$Cl F_{3}$ and $Br F_{3}$ are fluorinating agents; for example, uranium is converted into volatile $U F_{6}$ for enriching $^{235}$ U: $U + 3Cl F_{3} \to U F_{6} + 3ClF$ .

Chlorofluorocarbons (freons), used as refrigerants, are not interhalogens but carbon compounds of chlorine and fluorine: Freon-11 is $CC l_{3} F$ , Freon-12 $CC l_{2} F_{2}$ and Freon-13 $CCl F_{3}$ .

10.1 Polyhalide Ions

A halide ion can combine with a halogen or interhalogen molecule to form a polyhalide ion. The best known is the triiodide ion, which explains why iodine dissolves in KI solution:

I^{-} (aq) + I_{2} ⇌ I_{3}^{-} (aq)

Other polyhalide anions include $IC l_{2}^{-}$ , $IB r_{2}^{-}$ , $IC l_{4}^{-}$ , $Br F_{4}^{-}$ and $I F_{6}^{-}$ . Polyhalide cations such as $I_{3}^{+}$ and $Cl F_{2}^{+}$ are also known.

Figure 9: Shapes of interhalogens and the triiodide ion from VSEPR theory. Lone pairs (shaded lobes) take equatorial positions in

Cl F_{3}

and

I_{3}^{-}

, giving T-shaped and linear molecules;

Br F_{5}

is a square pyramid and

I F_{7}

a pentagonal bipyramid.

JEE Advanced

Pseudohalogens. A few small, stable groups behave chemically like halogens. Cyanogen $(CN)_{2}$ , thiocyanogen $(SCN)_{2}$ and cyanogen azide behave like $X_{2}$ , while $C N^{-}$ , $SC N^{-}$ and $OC N^{-}$ behave like $X^{-}$ : they form acids (HCN, HSCN), insoluble silver salts (AgCN, AgSCN), interpseudohalogens (ClCN, ICN) and disproportionate in alkali, $(CN)_{2} + 2O H^{-} \to C N^{-} + OC N^{-} + H_{2} O$ , exactly as chlorine does.

Figure 10: VSEPR in three counts. Lone pairs take equatorial positions in the trigonal bipyramid and opposite (trans) positions in the octahedron, which fixes every interhalogen shape.

Quick Recall: tap to check

What is the shape of

IC l_{4}^{-}

Square planar: 4 bond pairs + 2 lone pairs (trans) around I.

Why are interhalogens more reactive than the halogens (except

F_{2}

The X-X′ bond is weaker than the X-X bond of the parent halogen (except F-F).

Name the active ingredient that makes bleaching powder bleach.

Hypochlorite,

OC l^{-}

, which releases nascent oxygen.

Figure 11: Group 17 on one page. Two exceptions to learn (Cl has the most negative electron gain enthalpy,

C l_{2}

the strongest bond) and one rule to apply everywhere: oxidising power falls from

F_{2}

I_{2}

11. Solved Examples

Solved Example 1

Which one of the following forms a linear (zig-zag) polymeric structure through hydrogen bonding?
(A) HCl
(B) HF
(C)

H_{2} O

(D)

N H_{3}

Solution:

Answer: (B). Each H-F molecule has one H and one F, so each can form only one hydrogen bond on either side; the molecules link into zig-zag chains $(HF)_{n}$ . Water and ammonia form three-dimensional networks, and HCl does not hydrogen bond significantly.

Solved Example 2

'Fluorine can displace only halide ions from their salts, but no halogen can displace fluoride.' Explain.

Solution:

Fluorine is the strongest oxidising agent ( $E^{\circ} = + 2.87$ V), so it oxidises $C l^{-}$ , $B r^{-}$ and $I^{-}$ to the free halogens. The fluoride ion, in turn, is the hardest halide to oxidise, and no other halogen is a strong enough oxidant to remove its electron, so no halogen can displace fluoride.

Solved Example 3

Although fluorine has a less negative electron gain enthalpy than chlorine, it is the stronger oxidising agent. Explain.

Solution:

The $E^{\circ}$ value refers to the whole change $\frac{1}{2} X_{2} (g) + e^{-} \to X^{-} (aq)$ , whose enthalpy is the sum of three steps: $\frac{1}{2} Δ_{d i ss} H$ (energy in), $Δ_{e g} H$ and $Δ_{h y d} H$ (energy out).

Step (kJ mol $^{- 1}$ )	F	Cl
$\frac{1}{2} Δ_{d i ss} H$	+79.4	+121.3
$Δ_{e g} H$	$- 333$	$- 349$
$Δ_{h y d} H$	$- 515$	$- 381$
Net	$- 769$	$- 609$

Chlorine gains a little on electron gain, but fluorine gains much more from its weak F-F bond and the large hydration enthalpy of the small $F^{-}$ ion. The net change is far more negative for fluorine (Figure 2).

Solved Example 4

Why are the halogens highly reactive?

Solution:

Because the X-X bond dissociation enthalpy is low (the molecules split easily into atoms) and the atoms have a very negative electron gain enthalpy, being one electron short of a noble gas configuration.

Solved Example 5

Give examples of compounds in which halogens show positive oxidation states of +1, +3, +4, +5, +6 and +7.

Solution:

State	+1	+3	+4	+5	+6	+7
Example	$C l_{2} O$	$Cl F_{3}$	$Cl O_{2}$	$I F_{5}$	$C l_{2} O_{6}$	$I F_{7}$

Solved Example 6

Which of the following shows only negative oxidation states?
(A) chlorine
(B) bromine
(C) iodine
(D) fluorine

Solution:

Answer: (D). Fluorine is the most electronegative element and has no $d$ orbitals, so it shows only $- 1$ (in HOF, fluorine is $- 1$ and oxygen 0).

Solved Example 7

The shape of

Cl O_{3}^{-}

is
(A) pyramidal
(B) tetrahedral
(C) triangular planar
(D) triangular bipyramidal

Solution:

Answer: (A). Cl has three bond pairs and one lone pair ( $s p^{3}$ ), so the ion is trigonal pyramidal.

Solved Example 8

Solutions of the salts of hypohalous acids (for example NaOCl) are alkaline. Explain.

Solution:

NaOX + H_{2} O ⇌ N a^{+} + O H^{-} + HOX

Hypohalous acids are weak and stay largely un-ionised, so the hypohalite ion hydrolyses, taking a proton from water and leaving $O H^{-}$ ions in solution.

Solved Example 9

2.00 g of bleaching powder is treated with excess dilute acid and KI. The iodine set free needs 45.0 mL of 0.500 M sodium thiosulphate. Find the percentage of available chlorine.

Solution:

CaOC l_{2} + 2HCl CaC l_{2} + H_{2} O + C l_{2}

C l_{2} + 2KI 2KCl + I_{2}

I_{2} + 2N a_{2} S_{2} O_{3} 2NaI + N a_{2} S_{4} O_{6}

Moles of thiosulphate $= 0.0450 \times 0.500 = 0.0225$ mol, so $I_{2} = C l_{2} = 0.01125$ mol.

Mass of chlorine $= 0.01125 \times 71.0 = 0.799$ g. Available chlorine $= \frac{0.799}{2.00} \times 100 =$ 39.9%.

Solved Example 10

Chlorine reacts with hot, concentrated NaOH. Write the balanced equation and show the change in oxidation state of chlorine.

Solution:

6NaOH + 3C l_{2} 5NaCl + NaCl O_{3} + 3 H_{2} O

This is a disproportionation: of the six Cl atoms, five go from 0 to $- 1$ (gaining 5 electrons) and one goes from 0 to +5 (losing 5 electrons). Electrons lost equal electrons gained, which is why the ratio of NaCl to $NaCl O_{3}$ is 5 : 1.

Solved Example 11

Predict the products when ICl reacts with water, and explain the result.

Solution:

ICl + H_{2} O HCl + HOI

In ICl, chlorine is more electronegative and carries the partial negative charge, so it ends up as the chloride ion (HCl); iodine, the less electronegative partner (+1), becomes hypoiodous acid.

Solved Example 12

Why is HF the weakest of the hydrohalic acids although fluorine is the most electronegative halogen?

Solution:

Acid strength depends mainly on how easily the H-X bond breaks in water. The H-F bond is very short and strong (574 kJ mol $^{- 1}$ , against 432 for H-Cl), and $F^{-}$ is strongly hydrogen bonded to the undissociated HF. So HF ionises only slightly ( $p K_{a}$ 3.2), while HCl, HBr and HI are strong acids.

Practice Questions

(i) $C l_{2}$ bleaches permanently but $S O_{2}$ only temporarily. Why? (ii) Why is HF stored in wax-coated bottles? (iii) What is the hybridisation of Br in $Br F_{3}$ and $Br F_{5}$ ?Answer: (i) $C l_{2}$ bleaches by oxidation (nascent O from HOCl), which is irreversible; $S O_{2}$ bleaches by reduction, and air re-oxidises the dye. (ii) HF attacks glass ( $Si O_{2} + 4HF \to Si F_{4} + 2 H_{2} O$ ). (iii) $s p^{3} d$ in $Br F_{3}$ ; $s p^{3} d^{2}$ in $Br F_{5}$ .
Which of the hydrohalic acids reacts with glass?Answer: Hydrofluoric acid, HF (it forms $Si F_{4}$ and $H_{2} Si F_{6}$ ).
Suggest a method for the laboratory preparation of DCl. Write a balanced equation.Answer: Heat NaCl with deuterated sulphuric acid: $NaCl + D_{2} S O_{4} \to NaDS O_{4} + DCl$ (or hydrolyse $PC l_{3}$ with $D_{2} O$ ).
Unlike $PbC l_{4}$ , $Pb I_{4}$ and $PbB r_{4}$ are not found because (A) chlorine is more electropositive (B) iodine and bromine are larger (C) bromide and iodide reduce Pb(IV), so iodine and bromine cannot hold lead in the +4 state (D) the statement is wrongAnswer: (C): Pb(IV) is a strong oxidant because of the inert pair effect, and it oxidises $B r^{-}$ and $I^{-}$ .
Which of the following is not true? (A) Among halide ions, iodide is the most powerful reducing agent (B) Fluorine is the only halogen that does not show variable oxidation states (C) HOCl is a stronger acid than HOBr (D) HF is a stronger acid than HClAnswer: (D): HF is a much weaker acid than HCl.
$F_{2}$ is formed by reacting $K_{2} Mn F_{6}$ with (A) $Sb F_{5}$ (B) $Mn F_{3}$ (C) $KSb F_{6}$ (D) $Mn F_{5}$ Answer: (A): $2 K_{2} Mn F_{6} + 4Sb F_{5} \to 4KSb F_{6} + 2Mn F_{3} + F_{2}$ , the only chemical preparation of fluorine.
Arrange $Cl O^{-}$ , $Cl O_{2}^{-}$ , $Cl O_{3}^{-}$ and $Cl O_{4}^{-}$ in order of decreasing oxidising power.Answer: $Cl O^{-}$ > $Cl O_{2}^{-}$ > $Cl O_{3}^{-}$ > $Cl O_{4}^{-}$ .
Why is ICl more reactive than $I_{2}$ ?Answer: The polar I-Cl bond is weaker than the I-I bond, so ICl breaks up more easily.

Common Mistakes to Avoid

Watch out

Calling HF the strongest hydrohalic acid. It is the weakest: $HF ≪ HCl < HBr < HI$ .
Saying fluorine has the most negative electron gain enthalpy. Chlorine does ( $- 349$ vs $- 333$ kJ mol $^{- 1}$ ).
Taking $F_{2}$ as the halogen with the strongest bond. The order is $C l_{2} > B r_{2} > F_{2} > I_{2}$ .
Giving fluorine a positive oxidation state (e.g. +1 in HOF). Fluorine is always $- 1$ ; in HOF oxygen is 0.
Assuming the strongest oxoacid is also the strongest oxidant. $HCl O_{4}$ is the strongest acid but $Cl O^{-}$ is the strongest oxidant.
Mixing up the alkali products: cold dilute NaOH gives NaOCl, hot concentrated NaOH gives $NaCl O_{3}$ .
Drawing $Cl F_{3}$ as trigonal planar. It is T-shaped: two lone pairs sit in equatorial positions.
Writing that iodine oxidises water. It cannot; instead oxygen oxidises iodide in acid.

Frequently Asked Questions

Why are group 17 elements called halogens?

The name comes from the Greek hals, salt, and gennan, to produce, because these elements react with metals to form salts such as sodium chloride. Group 17 elements, fluorine, chlorine, bromine, iodine and astatine, have the $n s^{2} n p^{5}$ configuration and need one electron to complete an octet.

Why is fluorine a stronger oxidising agent than chlorine?

Although chlorine has the more negative electron gain enthalpy, fluorine's F-F bond is weak and the small fluoride ion releases much more energy on hydration. The overall change from half a mole of $F_{2}$ gas to aqueous fluoride is about -769 kJ, against -609 kJ for chlorine, so fluorine is the stronger oxidant.

Why is hydrofluoric acid a weak acid?

The H-F bond is short and very strong, 574 kJ per mole, so it breaks with difficulty in water, and the fluoride ion formed stays hydrogen bonded to HF molecules. As a result HF ionises only slightly, with a pKa of about 3.2, while HCl, HBr and HI are strong acids.

What is bleaching powder and what is available chlorine?

Bleaching powder, roughly $Ca (OCl)_{2}$ · $CaC l_{2}$ · $Ca (OH)_{2}$ ·2 $H_{2} O$ , is made by passing chlorine over dry slaked lime. Excess dilute acid or carbon dioxide releases chlorine from it; the chlorine released, as a percentage of the sample's mass, is its available chlorine, usually 35 to 38 percent.

What are interhalogen compounds?

Interhalogens are covalent compounds of two different halogens, of the types XX', XX'3, XX'5 and XX'7, for example ClF, $Cl F_{3}$ , $Br F_{5}$ and $I F_{7}$ . They are more reactive than the parent halogens (except fluorine) because the polar X-X' bond is weaker, and they are used as fluorinating agents.

Why does acid strength increase from hypochlorous to perchloric acid?

As oxygen atoms are added, chlorine's oxidation state rises from +1 to +7 and the negative charge of the anion is spread over more oxygen atoms by resonance. The perchlorate ion is therefore the most stable anion, so perchloric acid gives up its proton most easily and is the strongest acid.

What does NEET ask from group 17 elements?

NEET covers the general trends of group 17 elements: electronic configuration, oxidation states, atomic size, ionisation and electron gain enthalpy, electronegativity, bond enthalpy and the anomalous behaviour of fluorine. Expect questions on oxidising power, acid strength of HX and oxoacids, and interhalogen shapes.

Which group 17 topics matter most for JEE Advanced?

JEE Advanced lists hydrohalic acids, oxides and oxoacids of chlorine, and bleaching powder. The most tested points are acid strength and oxidising power trends, disproportionation of chlorine in alkali, structures of chlorine oxides, oxoacids and interhalogens, and available chlorine calculations.

Previous year questions on Group 17 Elements: The Halogen Family

10 questions from past papers, each with a step-by-step solution.

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Practice questions →Take a mock test

1. The Halogen Family at a Glance

2. Trends in Physical Properties

3. Oxidation States and the Anomalous Behaviour of Fluorine

4. Chemical Reactivity of the Group

4.1 Reactivity and Oxidising Power

4.2 Hydrogen Halides (HX)

4.3 Halides

4.4 Oxides

5. Chlorine (Cl2​)

5.1 Preparation

5.2 Properties

6. Fluorine, Bromine and Iodine

6.1 Fluorine

6.2 Bromine

6.3 Iodine

7. Hydrogen Chloride (HCl)

8. Oxoacids of Halogens

8.1 Oxoacids of Chlorine in Detail

9. Bleaching Powder

10. Interhalogen Compounds and Polyhalide Ions

10.1 Polyhalide Ions

11. Solved Examples

Common Mistakes to Avoid

Frequently Asked Questions

Why are group 17 elements called halogens?

Why is fluorine a stronger oxidising agent than chlorine?

Why is hydrofluoric acid a weak acid?

What is bleaching powder and what is available chlorine?

What are interhalogen compounds?

Why does acid strength increase from hypochlorous to perchloric acid?

What does NEET ask from group 17 elements?

Which group 17 topics matter most for JEE Advanced?

Previous year questions on Group 17 Elements: The Halogen Family

Ready to master The p-Block Elements: Group 15, 16, 17 And 18?

5. Chlorine ( $C l_{2}$ )