JEE Main & Advanced Chemistry The p-block Elements-II Halogen Family

Halogen Family

Category : JEE Main & Advanced

Fluorine is the first member of group 17 or VIIA of the periodic table. It consists of five elements Fluorine (F), Chlorine (Cl), bromine (Br), iodine (I) and astatine (At). These are known as halogen because their salts are found in sea water. Halogen is a greek word meaning a sea salt.            

(1) Electronic configuration

Elements Electronic configuration (\[n{{s}^{2}}\ n{{p}^{5}}\])
\[_{9}F\] \[[He]\,2{{s}^{2}}2{{p}^{5}}\]
\[_{17}Cl\] \[[Ne]\,3{{s}^{2}}3{{p}^{5}}\]
\[_{35}Br\] \[[Ar]\,3{{d}^{10}}4{{s}^{2}}4{{p}^{5}}\]
\[_{53}I\] \[[Kr]\,4{{d}^{10}}5{{s}^{2}}5{{p}^{5}}\]
\[_{85}At\] \[[Xe]\,4{{f}^{14}}5{{d}^{10}}6{{s}^{2}}6{{p}^{5}}\]

Physical properties            

(1) Atomic and ionic radii : A halogen atom has the smallest radius as compared to any other element in its period. This is due to the increased effective nuclear charge which results in greater attraction of the electrons by the nucleus. The atomic radii. Increase from fluorine to iodine down the group due to increase in number of shells.

 

Element F Cl Br I
Covalent radius (pm) 72 99 114 133
Ionic radius (pm) 133 184 196 220

 

(2) Ionization energy : Ionization energy of these elements are higher than those of the corresponding elements of group 16 due to increased nuclear charge. these values decrease systematically as we move down the group from \[F\] to 1

 

Element F Cl Br I
I.E1  \[(kJ\text{ }mo{{l}^{1}})\] 1680 1256 1142 1008

 

Thus iodine which has a comparatively low value of I.E., has a tendency to lose an electon to form positive iodinium ion, I+ and thus shown electropositive or metallic character.            

(3) Electronegativity : Fluorine is the most electronegative element in the periodic table. With increase in atomic number down the group, the electronegativity decreases.

 

Element F Cl Br I At
Electronegativity 4.0 3.2 3.0 2.7 2.2

 

The decreasing order of electronegativity is \[F>Cl>Br>I\]            

(4) Electron affinity : Electron affinity of chlorine, bromine and iodine decrease as the size of the atom increases. The electron affinity of fluorine is, however, lower than that of \[Cl\] and \[Br\], because of its small size as a result of which inter-electronic repulsions present in its \[2p\] subshell are comparatively large. Thus chlorine has the highest electron affinity.

 

Element F Cl Br I
Electron affinity \[(kJ\text{ }mo{{l}^{1}})\] 333 348 325 296

 

The decreasing order of electron affinity is \[Cl>F>Br>I\]            

(5) Oxidation states : All the halogens show an oxidation state of –1. Fluorine being the most electronegative element always shows an oxidation state of ?1 while other halogens also show positive oxidation states up to a maximum of +7 (i.e. +1, +3, +5 and +7) due to the availability of vacant \[d\]-orbitals in the valence shell of these atoms. Some halogens also show +4 and +6 oxidation states in oxides and oxy acids.            

(6) Nature of bonds : All the halogens have seven electrons in the valence shell and hence require one more electron to acquire the nearest inert gas configuration either by gaining an electron from the metallic atom to form halide, \[{{X}^{-}}\] ion, or by sharing an electron with an electronegative element. Thus, halogens form both ionic and covalent compounds. The halides of highly electropositive metals are ionic while those of weakly electropositive metals and non-metals are covalent. The tendency to form ionic compounds decreases from \[F\] to I. Thus, \[F\] because of its high electronegativity forms ionic compounds even with less electropositive metals like \[Hg,\,Bi,\,Sn\] etc. while other halogens form only covalent compounds.            

(7) Non-metallic character : All the halogens are non-metallic in nature due to their high ionization energies. The non-metallic character gradually decreases down the group. However, iodine is, solid and has metallic lustre.            

(8) Atomicity and physical state : All the halogens exist as diatomic covalent molecules \[({{F}_{2}},\,C{{l}_{2}},\,B{{r}_{2}}\] and \[{{I}_{2}})\]. \[{{F}_{2}}\] and \[C{{l}_{2}}\] are gases at room temperature, \[B{{r}_{2}}\] is corrosive liquid and \[{{I}_{2}}\] is volatile solid.            

(9) Colour : All the halogens have characteristic colours. \[{{F}_{2}}\] is light yellow, \[C{{l}_{2}}\]is greenish yellow, \[B{{r}_{2}}\] is reddish brown and \[{{I}_{2}}\] is deep violet. The colour of halogens is due to the reason that their molecules absorb light in the visible region as a result of which electron are excited to higher energy levels. The amount of energy needed for excitation decreases progressively from \[{{F}_{2}}\] to \[{{I}_{2}}\] and consequently there is a progressive deepening of colour of the halogens from \[{{F}_{2}}\] to \[{{I}_{2}}\]. Since, fluorine atom requires large amount of energy for excitation of electrons and therefore absorbs violet light and apears yellow. On the other hand, iodine requires low energy for excitation of electrons (absorbs yellow light) and appears deep violet.            

(10) Bond dissociation energy : Bond dissociation energies of chlorine, bromine and iodine decrease down the group as the size of the atom increases. The bond dissociation energy of fluorine, is however, lower than those of chlorine and bromine because of inter electronic repulsions present in the small atom of fluorine

 

\[{{X}_{2}}\] \[{{F}_{2}}\] \[C{{l}_{2}}\] \[B{{r}_{2}}\] \[{{I}_{2}}\]
Bond dissociation energy \[(kJ\text{ }mo{{l}^{1}})\] 158 243 192 151

 

Hence bond energy decreases in the order \[C{{l}_{2}}>B{{r}_{2}}>{{F}_{2}}>{{I}_{2}}\]            

(11) Bond length in \[{{X}_{2}}\] molecule : As the size of the halogen atom increase, the bond length of \[X-X\] bond in \[{{X}_{2}}\] molecule increases from \[{{F}_{2}}\] to \[{{I}_{2}}\]  

          

X – X  bond F–F Cl–Cl Br–Br I–I
Bond length (pm) 143 199 228 266

 

Thus, the bond length increases in the order \[{{F}_{2}}<C{{l}_{2}}<B{{r}_{2}}<{{I}_{2}}\].            

(12) Melting points and boiling points : Melting points and boiling points of these elements increase as we move down the group from \[F\] to \[I\] due to an increase in the vander Waals forces of attraction which increase down the group as the size of the atom increases.

 

Element F Cl Br I
Melting point (K) 54 172 266 386
Boiling point (K) 85 239 332 458

 

Hence, the melting points and boiling points show the order as \[F<Cl<Br<I\].            

(13) Solubility : Halogens, being non polar in nature do not readily dissolve in a polar solvent like water. however, fluorine reacts with water vigorously even at low temperature (exothermally) forming a mixture of ozone and oxygen                      

\[2{{F}_{2}}+2{{H}_{2}}O\xrightarrow{{}}4HF+{{O}_{2}}\]                      

\[3{{F}_{2}}+3{{H}_{2}}O\xrightarrow{{}}6HF+{{O}_{3}}\]            

Chlorine and bromine are fairly soluble but iodine is very little soluble in water. chlorine, bromine and iodine are more soluble in organic solvents like \[CC{{l}_{4}},C{{S}_{2}}\] or \[CHC{{l}_{3}}\] and produce coloured solutions. Thus \[C{{l}_{2}},\,B{{r}_{2}}\] and \[{{I}_{2}}\] give yellow, brown and violet colour respectively. It is believed that in non-polar solvents, halogens exist as free molecules just as in the gas phase.            

In nucleophilic (electron donating) polar solvents like alcohols, ketones or liquid \[S{{O}_{2}}\], halogens produce brown solution. This colour is due to the complex formation (solvent \[\to \] halogen) which are charge transfer compounds.            

The solubility of iodine \[({{I}_{2}})\] in water increases with addition of \[KI\] or \[NaI\] due to the formation of polyhalide (triiodide, \[I_{3}^{-}\]) ion, \[KI+{{I}_{2}}\]\[\rightleftharpoons \]\[K{{I}_{3}}\]

However, this solution behaves as a simple mixture of \[KI\] and free \[{{I}_{2}}\] and contains \[{{K}^{+}}\] and \[{{I}^{-}}\] ions and free \[{{I}_{2}}\] molecules. It has a brown colour. The solution of iodine in water (due to its very little solubility) is also brown.            

(14) Oxidising power : All the halogens acts as strong oxidising agents since they have a strong tendency to attract electrons and have positive values of electrode potentials \[({{E}^{o}})\]. The oxidising power, however, decreases as we move down the group from \[F\] to \[I\]. i.e., \[{{F}_{2}}>C{{l}_{2}}>B{{r}_{2}}>{{I}_{2}}\]            

Since \[{{F}_{2}}\] is the strongest oxidising agent, it will oxidise all other halide ions to halogens.                      

\[{{F}_{2}}+2{{X}^{-}}\xrightarrow{{}}2{{F}^{-}}+{{X}_{2}}(X=Cl,\,Br,I)\]            

Similarly, \[C{{l}_{2}}\] will displace \[B{{r}^{-}}\] and \[{{I}^{-}}\] ions from their solutions while \[B{{r}_{2}}\] will displace \[{{I}^{-}}\] ions only.

\[C{{l}_{2}}+2{{X}^{-}}\xrightarrow{{}}2C{{l}^{-}}+{{X}_{2}}(X=Br,I)\]                      

\[B{{r}_{2}}+2{{I}^{-}}\xrightarrow{{}}2B{{r}^{-}}+{{I}_{2}}\]            

Hence \[{{F}_{2}}\] is the strongest and \[{{I}_{2}}\] is the weakest oxidising agent. This is also indicated by the decrease in the electrode potential \[({{E}^{o}})\] for the reaction \[{{X}_{2}}(aq)+2{{e}^{-}}\xrightarrow{{}}2{{X}^{-}}(aq)\] on moving down the group.      

                                  

\[{{X}_{2}}\] \[{{F}_{2}}\] \[C{{l}_{2}}\] \[B{{r}_{2}}\] \[{{I}_{2}}\] \[A{{t}_{2}}\]
\[{{E}^{o}}\] (volts) 2.87 1.36 1.09 0.53 0.3

 

The electron affinity of fluorine is less than that of chlorine but still it is the strongest oxidising agent. This is because of its low bond dissociation energy \[(158\,kJ\,mo{{l}^{-1}})\] and high heat hydration \[(510\,kJ\,mo{{l}^{-1}})\] as compared to chlorine (for which the values are 243  and \[372\,kJ\,mo{{l}^{-1}}\], respectively).            

(15) Heat of hydration : The heat of hydration of the halide ion \[({{X}^{-}})\] decreases as the size of the halogen decreases down the group from \[F\] to \[I\].

 

\[{{X}^{-}}\] ion \[{{F}^{-}}\] \[C{{l}^{-}}\] \[B{{r}^{-}}\] \[{{I}^{-}}\]
Heat of hydration \[(kJ\text{ }mo{{l}^{1}})\] 510 372 339 301

 

Thus, the decreasing order of heat of hydration of halides is as follows :                      

\[{{F}^{-}}>C{{l}^{-}}>B{{r}^{-}}>{{I}^{-}}\]            

Chemical properties              

(1) Reactivity : The halogen are most reactive elements due to their low bond dissociation energy, high electron affinity and high enthalpy of hydration of halide ion.  \[F>Cl>Br>I\]

(2) Reaction with \[{{H}_{2}}O\] : Halogens readily decomposes water. This tendency decreases on moving down the group. Fluorine decomposes water very energetically to give oxygen and ozone, 

\[2{{H}_{2}}O+2{{F}_{2}}\to 4HF+\underset{\text{Oxygen}}{\mathop{{{O}_{2}}}}\,\];  \[3{{H}_{2}}O+3{{F}_{2}}\to 6HF+\underset{\text{Ozone}}{\mathop{{{O}_{3}}}}\,\]            

Fluorine gives fumes in moist air. This is due to the formation of HF, which is a liquid and can absorb moisture to form liquid droplets and therefore, gives fumes with moist air. Chlorine and bromine react less vigorously,          

\[C{{l}_{2}}+{{H}_{2}}O\to HCl+\underset{\text{Hypochlorous acid}}{\mathop{HClO}}\,\]            

\[B{{r}_{2}}+{{H}_{2}}O\to HBr+\underset{\text{Hypobromous acid}}{\mathop{HBrO}}\,\]              

In the presence of sunlight, HClO (hypochlorous acid) HBrO (hypobromous acid) liberate oxygen.                      

\[2HClO\to 2HCl+{{O}_{2}}\];  \[2HBrO\to 2HBr+{{O}_{2}}\]            

Iodine is only slightly soluble in water. However, it dissolves in 10% aqueous solution of Kl due to the formation of \[I_{3}^{-}\] ions.

\[{{I}_{2}}+KI\rightleftharpoons K{{I}_{2}}+{{I}^{-}}\rightleftharpoons \underset{\text{Complex ion}}{\mathop{I_{3}^{-}}}\,\]

(3) Reaction with hydrogen : Form covalent halides.                      

\[{{H}_{2}}+{{F}_{2}}\xrightarrow{-{{200}^{o}}C}2HF\] (very violent)                      

\[{{H}_{2}}+C{{l}_{2}}\xrightarrow{\text{Sunlight}}2HCl\]

\[{{H}_{2}}+B{{r}_{2}}\underset{\text{pt}\text{. catalyst}}{\mathop{\xrightarrow{\text{Heat}}}}\,2HBr\]

\[{{H}_{2}}+{{I}_{2}}2HI\] (poor yield)            

  • Acidic strength in aqueous solution is in the order,                      

\[HI>HBr>HCl<HF\]

  • Reducing character of hydrides follow the order,

\[HI>HBr>HCl>HF.\]

Boiling point  \[HF>HI>HBr>HCl.\] Thermal stability,                      

\[HF>~HCl>HBr>HI.\]

HCl is also called Muriatic acid.

(4) Hydrides : All the halogens combine directly with hydrogen to form halogen acids but their reactivity progressively decreases from fluorine to iodine, \[{{H}_{2}}+\text{ }{{X}_{2}}\to 2HX\,\,(X=F,\text{ }Cl,\text{ }Br\,\,\text{or  }I).\] 

(i) Boiling points or volatility : In other words volatility decreases in the order : HCl > HBr > HI > HF as the boiling points increase in the order        :  HCl (189K)  <  HBr (206K)  <  HI (238K)  <   HF (292.5K).            

(ii) Thermal stability : Thermal stability of the hydrides decrease from HF to HI  i.e., HF > HCl > HBr > HI.            

(iii) Acidic strength : The acidic strength of halogen acids decreases from HI to HF i.e, HI > HBr > HCl > HF.            

(iv) Reducing properties : Since the stability of hydrides decreases from HF to Hl, their reducing properties increase in the order HF < HCl < HBr < HI.            

(v) Dipole moments : The dipole moments of hydrogen halides decrease in the order : HF > HCI > HBr > HI as the electro negativity of the halogen atom decreases form F to I.                      

                      HX             HF             HCl         HBr            Hl

Dipole moment (D)       1.74          1.07         0.78          0.38    

(5) Oxides : Halogens (except \[{{F}_{2}}\]) do not combine readily with oxygen. However, a number of compounds of halogens with oxygen have been prepared by indirect methods. Only two compounds of fluorine with oxygen, i.e. oxygen difluorine \[(O{{F}_{2}})\] and oxygen fluoride \[({{O}_{2}}{{F}_{2}})\] are known. Chlorine forms largest number of oxides i.e. \[C{{l}_{2}}O,\text{ }Cl{{O}_{2}},\text{ }C{{l}_{2}}{{O}_{6}}\] and \[C{{l}_{2}}{{O}_{7}}\] while iodine forms the least, i.e. \[{{I}_{2}}{{O}_{5}}\]. Bromine, however, forms three oxides \[(B{{r}_{2}}O,\text{ }Br{{O}_{2}}C\text{ }Br{{O}_{3}})\]. In all these compounds, bonds are largely covalent. All the oxides of halogens are powerful oxidizing agents. These compounds are very reactive and are unstable towards heat. The stability of oxides is greatest for iodine while bromine oxides are the least stable. For a particular halogen, higher oxides are more stable than the lower ones.

Iodine-oxygen bond is stable due to greater polarity of the bond (due to larger electro negativity difference between I and O) while in chlorine-oxygen bond, the stability is gained through multiple bond formation involving the d-orbital of chlorine atom. Bromine lacks both these characteristics and hence forms least stable oxides.

Oxides of chlorine, bromine and iodine are acidic and the acidic character increases as the percentage of oxygen increases in them.

Iodine also forms \[{{I}_{2}}{{O}_{4}}\] and \[{{I}_{4}}{{O}_{9}}\] compounds which are believed not to be true oxides but are basic iodyliodate, \[IO(I{{O}_{3}})\] and normal iodine triodate, \[I{{(I{{O}_{3}})}_{3}}\] having tripositive iodine as the cation.

\[O{{F}_{2}}\] is V-shaped having bond angle \[{{103}^{o}},\,C{{l}_{2}}O\] is also V-shaped with bond angle \[{{111}^{o}}\] while \[Cl{{O}_{2}}\] is angular with-bond angle \[{{118}^{o}}\]. It is paramagnetic due to odd number of electrons having three-electron bond. It is regarded as a mixed anhydride of chloric and chlorous acids.  \[2Cl{{O}_{2}}+{{H}_{2}}O\to HCl{{O}_{2}}+HCl{{O}_{3}}\]

(6) Oxoacids of halogens : Fluorine does not form any oxoacid since it is the strongest oxidizing agent. Chlorine, bromine and iodine mainly form four series of oxoacids namely hypohalous acid \[(HXO),\] halous acid \[(HX{{O}_{2}})\] halic acid \[(HX{{O}_{3}})\] and perhalic acid \[(HX{{O}_{4}})\]as given below :

 

Oxidation state Chlorine Bromine Iodine Thermal stability and acid strength  Oxidising power
+1 \[HClO\] \[HBrO\] \[HIO\]
+3 \[HCl{{O}_{2}}\]
+5 \[HCl{{O}_{3}}\] \[HBr{{O}_{3}}\] \[HI{{O}_{3}}\]
+7 \[HCl{{O}_{4}}\] \[HBr{{O}_{4}}\] \[HI{{O}_{4}}\]
  Acidity decreases  \[\to \]  

 

(i) Hybridized ion : In all these oxoacids, the halogen atom is \[s{{p}^{3}}\]-hybridized.

(ii) Acidic character : All these acids are monobasic containing an \[-OH\] group. The acidic character of the oxoacids increases with increase in oxidation number, i.e., \[HClO<HCl{{O}_{2}}<HCl{{O}_{3}}<HCl{{O}_{4}}\] and the strength of the conjugate bases of these acids follows the order,

\[Cl{{O}^{-}}~>ClO_{2}^{-}>ClO_{3}^{-}~>ClO_{4}^{-}\]

(iii) Oxidising power and thermal stability : The oxidizing power of these acids decreases as the oxidation number increases, i.e., \[HClO<HCl{{O}_{2}}<HCl{{O}_{3}}<HCl{{O}_{4}}\]. Stability of oxoacids of  chlorine  in  the  increasing order is, \[HClO<HCl{{O}_{2}}<HCl{{O}_{3}}<HCl{{O}_{4}}\] and the increasing stability order of anions of oxoacids of chlorine is, \[Cl{{O}^{-}}<ClO_{2}^{-}<ClO_{3}^{-}<ClO_{4}^{-}\].

As the number of oxygen atoms in an ion increases there will be a greater dispersal of negative charge and thus greater will be the stability of ion formed. For different halogen having the name oxidation number, the thermal stability decreases with increase in atomic number i.e., it is in the order \[HClO>HBrO>HIO\] and \[Cl{{O}^{-}}>Br{{O}^{-}}>l{{O}^{-}}\] However, in \[HX{{O}_{3}}\] is most stable. The stability order being \[HCl{{O}_{3}}<HBr{{O}_{3}}<HI{{O}_{3}}\].

(iv) Perhalates are strong oxidizing agents, the oxidizing power is in the order, \[BrO_{4}^{-}\ >\ IO_{4}^{-}\ >\ ClO_{4}^{-}\].

Thus \[Br{{O}_{4}}\] is the strongest oxidizing agent (though its reaction is quite slow) and \[ClO_{4}^{-}\] is the weakest.

(v) The acidity of oxoacids of different halogens having the same oxidation number decreases with increase in the atomic size of the halogen i.e. \[HCl{{O}_{4}}\ >\ HBr{{O}_{4}}>HI{{O}_{4}}\].

(7) Reaction with alkalies :

\[2{{F}_{2}}+\underset{\text{(cold dilute)}}{\mathop{2NaOH}}\,\to 2NaF+O{{F}_{2}}+{{H}_{2}}O\]

\[2F+\underset{\text{(hot conc}\text{.)}}{\mathop{4NaOH}}\,\to 4NaF+{{O}_{2}}+2{{H}_{2}}O\]

Halogen other than fluorine \[(C{{l}_{2}},B{{r}_{2}},{{I}_{2}})\] react with \[NaOH\] as follows,

\[{{X}_{2}}(g)+\underset{\text{(cold dilute)}}{\mathop{2O{{H}^{-}}}}\,\xrightarrow{{{15}^{o}}C}{{X}^{-}}+\underset{\text{(hypohalite  ion)}}{\mathop{O{{X}^{-}}+{{H}_{2}}O}}\,\]

\[{{X}_{2}}(g)+\underset{\text{(hot conc)}}{\mathop{6O{{H}^{-}}}}\,\xrightarrow{{{70}^{o}}C}5{{X}^{-}}+\underset{\text{(halate ion)}}{\mathop{XO_{3}^{-}}}\,+3{{H}_{2}}O\]

(8) Bleaching action of halogen : \[C{{l}_{2}}\] acts as bleaching agent, its bleaching action is permanent. \[C{{l}_{2}}\] water can also act as ink remover.

(9) Reaction with other halides

\[2KBr(aq.)+C{{l}_{2}}(g)\to 2KCl(aq.)+B{{r}_{2}}(aq.)\] 

\[2KI(aq.)+C{{l}_{2}}(g)\to 2KCl(aq.)+{{I}_{2}}(aq.)\]

(10) Inter halogen compounds : The compounds of one halogen with the other are called inter halogens or inter halogen compounds. The main reason for their formation is the large electronegativity and the size differences between the different halogens. Taking A as the less electronegative and B as the more electronegative halogen, they are divided into the following four types the less electronegative halogen (A) is always written first.

 

AB \[A{{B}_{3}}\] \[A{{B}_{5}}\] \[A{{B}_{7}}\]
ClF \[BrF,\,BrCl,\,ICl\] \[IBr,\,IF\] \[Cl{{F}_{3}},Br{{F}_{3}}\] \[I{{F}_{3}},\,IC{{l}_{3}}\] \[Br{{F}_{5}}I{{F}_{5}}\] \[I{{F}_{7}}\]

These interhalogen compounds are unstable and more reactive            

(i) General properties            

(a) Largest halogen always serves the central atom.            

(b) The highest interhalogen compound i.e. IF7  is obtained with iodine, the largest halogen attached to the smallest one            

(c) The bonds in interhalogen compounds are essentially covalent.            

(d) Thermal stability decreases as the size difference decreases and increases as the polarity of the bond increases. Thus ClF is thermally more stable as compared to IBr.            

(e) They ionize in solution or in the liquid state,                      

\[2ICl\]\[\rightleftharpoons \]\[{{I}^{+}}+ICl_{2}^{-}\,;\,\,\]\[2IC{{l}_{3}}\]\[\rightleftharpoons \]\[ICl_{2}^{+}+ICl_{4}^{-}\]

(f) Hydrolysis of interhalogen compounds always produces a halide ion derived from smaller halogen and oxyhalide derived from larger halogen, 

\[ICl+{{H}_{2}}O\to C{{l}^{-}}+O{{I}^{-}}+2{{H}^{+}}\];                                  

\[Br{{F}_{5}}3{{H}_{2}}O\to 5{{F}^{-}}+BrO_{3}^{-}+6{{H}^{+}}\]            

(g) They are strong oxidizing agents.            

(h) Largest number of interhalogens are formed by fluorine due to its smaller size and higher electronegativity or oxidizing power.            

(ii) Structure : Interhalogen compounds are,            

(a) AB type i.e. ICl, IBr, IF etc, are linear            

(b) \[A{{B}_{3}}\] type i.e. \[I{{F}_{3}},\text{ }Cl{{F}_{3}},\text{ }Br{{F}_{3}}\] have distorted trigonal bipyramidal (\[ds{{p}^{3}}-\]hybridization) structures of T-shape due to two lone pairs in equatorial positions \[IC{{l}_{3}}\] is dimeric, \[{{I}_{2}}C{{l}_{6}}\] and has a planar structure.

(c) \[A{{B}_{5}}\] types i.e. \[Br{{F}_{5}},\text{ }I{{F}_{5}}\] have distorted octahedral (\[{{d}^{2}}s{{p}^{3}}-\]hybridization) shapes or square pyramidal due to a lone pair one of the axial positions.

(d) \[A{{B}_{7}}\] type i.e. \[I{{F}_{7}},\] have pentagonal bipyramidal (\[{{d}^{3}}s{{p}^{3}}-\]hybridization) structures.

(11) Polyhalide ions : Halogens or interhalogens combine with halide ions to form polyhalide ions. The most common example of polyhalide ion formation is furnished by the increase in solubility of iodine in water in the presence of \[KI\] which is due to the formation of tri iodide ion, \[I_{3}^{-}\]

\[{{I}^{-}}+{{I}_{2}}\xrightarrow{{}}I_{3}^{-}\]

Many other examples of polyhalides ions are

(i) \[Cl_{3}^{-},Br_{3}^{-},\,ICl_{2}^{-},\,IBr_{2}^{-}\] including \[I_{3}^{-}\]. In these ions, one of the halogen atoms (in case of similar atoms) or halogen atom larger in size undergoes \[s{{p}^{3}}d\]-hybridization giving a linear shape with three lone pairs at equatorial positions.

(ii) \[Cl_{3}^{+},\,Br_{3}^{+},I_{3}^{+},ICl_{2}^{+},\,IBr_{2}^{+}\]. Here we find central atom \[s{{p}^{3}}\] hybridized giving a bent shape with two lone pairs of electrons on the central atom.

(iii) \[ICl_{4}^{-},\,BrF_{4}^{-},\,I_{5}^{-}\]. Here central atom involves \[s{{p}^{3}}{{d}^{2}}\] hybridization giving square planar shape with two lone pairs of electrons on axial positions.

(iv) \[ICl_{4}^{+},BrF_{4}^{+},\,I_{5}^{+}\]. In these ions central atom involves \[s{{p}^{3}}d\] hybridization giving a distorted tetrahedral structure with one lone pair of electrons on equatorial position.

(v) \[I_{7}^{-},\,IF_{6}^{-}\]. The central atom \[I\] undergoes \[s{{p}^{3}}{{d}^{3}}\] hybridization giving a distorted octahedral structure with one lone pair of electrons.

(vi) \[I_{7}^{+}\]. Here central \[I\] atom involves \[s{{p}^{3}}{{d}^{2}}\] hybridization giving an octahedral structure.

Fluorine due to its highest electronegativity (and only –1 oxidation state) does not form polyhalide ions where it acts as a central atom.

(12) Pseudohalogen and pseudohalides 

Pseudohalogen Pseudohalide
Cyanogen ?\[{{(CN)}_{2}}\] Cyanide ?\[C{{N}^{-}}\]
Oxocyanogen ? \[{{(OCN)}_{2}}\] Cyanate ?\[OC{{N}^{-}}\]
Thiocyanogen ? \[{{(SCN)}_{2}}\] Thiocyanate ?\[SC{{N}^{-}}\]
Selenocyanogen ? \[{{(SeCN)}_{2}}\] Selenocyanate ? \[SeC{{N}^{-}}\]

 

(13) Anomalous behaviour of fluorine : Fluorine differs from  rest  of the  elements  of  its family  due to  (i) its small size (ii) highest electronegativity, (iii) low bond dissociation energy and (iv) absence of d-orbitals in the valence shell. The main points of difference are :

 

(1) Fluorine is most reactive of all the halogens due to lower value of \[F-F\] bond dissociation energy \[{{F}_{2}}=158,\,C{{l}_{2}}=243\], bromine = 192 and iodine = 151 \[kJ\,\,mo{{l}^{-1}}\]) .

 

(2) Being the most electronegative element, it shows only an oxidation state of –1 and does not show positive oxidation states due to absence of \[d\]-orbitals in its valence shell. Other halogens show positive oxidation states of +1, +3, +5 and +7.

 

(3) Due to small atomic size and high electronegativity of \[F,\,HF\] undergoes strong \[H\]-bonding while other halogen acids do not. As a result,

 

(i) \[HF\] is a liquid (boiling point 292.5K), while other halogen acids are gases at room temperature (boiling point of \[HCl=189K\], \[HBr=206K\], \[HI=238K\]).

 

(ii) \[HF\] is weakest of all the halogen acids due to high strength of \[H-F\] bond.

 

(iii) Due to \[H\]-bonding, \[HF\] can form acid salts of the type \[KH{{F}_{2}}\], i.e., \[{{K}^{+}}[H-F.......{{F}^{-}}]\] while \[HCl,\,HBr\] and \[HI\] do not form such salts (i.e., no \[KHC{{l}_{2}},\,KHB{{r}_{2}}\] and \[KH{{I}_{2}}\] are known).

 

(4) Fluorides have the maximum ionic character. For example \[Al{{F}_{3}}\] is ionic while other halides of \[Al\] are covalent.

 

(5) Of all the halogens, fluorine has the highest positive electrode potential (\[{{F}_{2}}=2.87,\,C{{l}_{2}}=1.36,\,B{{r}_{2}}=1.09\] and \[{{I}_{2}}=0.53\] volt) i.e., it is most easily reduced and hence acts as the strongest oxidising agent. It brings about the highest oxidation of other elements with which it combines. For example with \[S\], it gives \[S{{F}_{6}}\], with \[{{I}_{2}}\] it gives \[I{{F}_{7}}\]. Other halogens do not always bring about the highest oxidation state. For example, with sulphur\[C{{l}_{2}}\] gives \[SC{{l}_{4}},\,B{{r}_{2}}\] gives \[SB{{r}_{2}}\] while \[{{I}_{2}}\] does not react at all. \[{{F}_{2}}\] is so powerful oxidising agent that it can even oxidise inert-gases.

 

(6) \[HF\] cannot be stored in glass bottles sicne it reacts with silicates to form fluorosilicates.

 

\[N{{a}_{2}}Si{{O}_{3}}+6HF\xrightarrow{{}}N{{a}_{2}}Si{{F}_{6}}+3{{H}_{2}}O\]

 

While other halogen acids (\[HCl,\,HBr\] and \[HI\]) do not react with silicates and hence can be stored in glass bottles.

 

(7) \[AgF\] is soluble in \[{{H}_{2}}O\] while all other silver halides i.e., \[AgCl,\,AgBr\] and \[AgI\] are insoluble in water. In constant, \[Ca{{F}_{2}}\] is insoluble while other calcium halides i.e., \[CaC{{l}_{2}},\,CaB{{r}_{2}},\,Ca{{I}_{2}}\] are soluble in \[{{H}_{2}}O\].

 

(8) Due to absence of \[d\]-orbitals, fluorine, does not form polyhalide ions while other halogens form polyhalides of the type \[I_{3}^{-},\,Br_{3}^{-},\,I_{5}^{-}\] etc.

Preparation of halogens and its uses

(1) Fluorine

 

(i) Occurrence of fluorine : Fluorine does not occur free in nature but occurs mostly as fluorspar \[Ca{{F}_{2}}\], cryolite, \[N{{a}_{3}}Al{{F}_{6}}\] and fluorapatite, \[Ca{{F}_{2}}.3C{{a}_{3}}{{(P{{O}_{4}})}_{2}}\]. Traces of fluoride occur in sea water, bones, teeth, blood, milk etc.

 

(ii) Difficulties encountered during its isolation : (a) \[{{F}_{2}}\] attacks all the materials of the apparatus such as glass, platinum, carbon and other metals, (b) \[{{F}_{2}}\] is the strongest oxidising agent and hence no oxidising agent can oxidise \[{{F}^{-}}\] ions to \[{{F}_{2}}\]. (c) \[{{F}_{2}}\] cannot be prepared even by electrolysis of an aqueous solution of \[HF\] because \[{{F}_{2}}\] formed reacts violently with water. If also cannot be prepared by electrolysis of anhydrous \[HF\] because it is not only poisonous, corrosive and volatile but also is a bad conductor of electricity.

 

(iii) Preparation : \[{{F}_{2}}\] is now prepared by electrolysis of a solution of \[KH{{F}_{2}}\] (1 part) in anyhydrous \[HF\](5 parts) in a vessel (modern method) made of \[Ni-Cu\] alloy or \[Ni-Cu-Fe\] alloy called the monel metal using carbon electrodes. During the electrolysis following reactions occur.

 

\[KH{{F}_{2}}\xrightarrow{{}}KF+HF\];\[KF\xrightarrow{{}}{{K}^{+}}+{{F}^{-}}\]

 

At cathode :  \[{{K}^{+}}+{{e}^{-}}\xrightarrow{{}}K\]; \[2K+2HF\xrightarrow{{}}2KF+{{H}_{2}}\uparrow \]

 

At anode :     \[{{F}^{-}}\xrightarrow{{}}F+{{e}^{-}}\]; \[F+F\xrightarrow{{}}{{F}_{2}}\]

 

(iv) Properties : It is the most reactive of all the halogens. It Combines with metals as well as non-metals to form fluorides. It decomposes water forming \[{{O}_{2}}\] and \[{{O}_{3}}\] and reacts vigorously with hydrogen of hydrocarbons leaving behind fluorinated hydrocarbons.

 

\[2{{H}_{2}}O+2{{F}_{2}}\underset{\text{Oxidation}}{\mathop{\xrightarrow{\text{Cold}}}}\,4HF+{{O}_{2}}\]

 

\[3{{H}_{2}}O+3{{F}_{2}}\underset{\text{Oxidation}}{\mathop{\xrightarrow{\text{Hot}}}}\,6HF+{{O}_{3}}\]

 

(\[HF\] being a volatile liquid fumes in air)

 

\[C{{H}_{4}}\xrightarrow{{{F}_{2}}}C{{H}_{3}}F\xrightarrow{{{F}_{2}}}C{{H}_{2}}{{F}_{2}}\xrightarrow{{{F}_{2}}}CH{{F}_{3}}\xrightarrow{{{F}_{2}}}C{{F}_{4}}\]

 

It is a strong oxidising agent and oxidises \[KCl{{O}_{3}}\] to \[KCl{{O}_{4}},\,Kl{{O}_{3}}\] to \[Kl{{O}_{4}}\] and bisulphates to peroxy sulphates.

 

\[KCl{{O}_{3}}+{{F}_{2}}+{{H}_{2}}O\xrightarrow{{}}KCl{{O}_{4}}+{{H}_{2}}{{F}_{2}}\]

 

\[2NaHS{{O}_{4}}+{{F}_{2}}\xrightarrow{{}}N{{a}_{2}}{{S}_{2}}{{O}_{8}}+2HF\]

 

It reacts with \[N{{H}_{3}}\] to form nitrogen and with \[{{H}_{2}}S\] forming \[S{{F}_{6}}\].

 

\[2N{{H}_{3}}+3{{F}_{2}}\xrightarrow{{}}{{N}_{2}}+6HF\](oxidation reaction)

 

\[{{H}_{2}}S+4{{F}_{2}}\xrightarrow{{}}S{{F}_{6}}+2HF\]

 

Fluorine reacts with cold and dilute sodium hydroxide solution to give oxygen difluoride \[(O{{F}_{2}})\]

\[2{{F}_{2}}+2NaOH\](cold, dil)\[\xrightarrow{{}}2NaF+{{H}_{2}}O+O{{F}_{2}}\]

 

However, with hot and concentrated sodium hydroxide solution it gives oxygen

\[2{{F}_{2}}+4NaOH(\text{Hot, conc}\text{.)}\xrightarrow{{}}4NaF+2{{H}_{2}}O+{{O}_{2}}\]

 

Since \[{{F}_{2}}\] is the strongest oxidising agent, it is always reduced and hence does not show disproportionation reactions while others halogens do.

 

\[{{F}_{2}}\] oxidises all other halide ions to the corresponding halogens \[({{F}_{2}}+2{{X}^{-}}\xrightarrow{{}}2{{F}^{-}}+{{X}_{2}})\]; \[(X=Cl,\,Br\] or \[I)\]

 

(v) Uses of fluorine : Fluorine is used in the manufacture of \[U{{F}_{6}}\] (which is used for nuclear power generation), \[S{{F}_{6}}\] (which is used as an electrical insulator), chlorofluorocarbons, teflon, cryolite and \[HF\].

 

(vi) Fluorocarbons are the derivatives of hydrocarbons in which \[H\]-atoms are replaced by \[F\]-atoms. these are obtained by fluorination of hydrocarbons with \[{{F}_{2}}\] diluted with an inert gas such as \[{{N}_{2}}\] in presence of \[Cu{{F}_{2}}\] as catalyst. Fluorocarbons are widely used in industry because of their extreme inertness (non-in-flammability and extreme stability). Freon \[(C{{F}_{2}}C{{l}_{2}})\] is used as a refrigerant, tetrafluoroethylene \[({{F}_{2}}C=C{{F}_{2}})\] is used for the manufacture of teflon which is highly non-inflammable, has high thermal stability and is chemically inert i.e., is not attacked by acids and corrosive chemicals. It is used for making pipes, surgical tubes, non-stick utensils and as an electrical insulator.

 

(2) Chlorine

 

(i) Occurrence : Chlorine mainly occurs as rock salt \[(NaCl)\] Carnallite, \[(KCl,\,MgC{{l}_{2}}.6{{H}_{2}}O)\] and Calcium chloride. \[(CaC{{l}_{2}})\].

 

(ii) Preparation : On a commercial scale chlorine is prepared by electrolysis of an aqueous solution of sodium chloride (brine solution) (Nelson cell, Castner and Kellner’s cell for the manufacture of \[NaOH\]) when \[C{{l}_{2}}\] is evolved at the anode and \[{{H}_{2}}\] is evolved at the cathode.

 

\[2NaCl+2{{H}_{2}}O\xrightarrow{\text{Electrolysis}}2NaOH+C{{l}_{2}}\uparrow +{{H}_{2}}\uparrow \]

 

It can also be prepared by electrolysis of molten \[NaCl\](Down’s cell for the manufacture of metallic sodium). When \[C{{l}_{2}}\] is evolved at the anode and sodium metal at the cathode.

 

\[2NaCl\xrightarrow{\text{Electrolysis}}2Na+C{{l}_{2}}\uparrow \]

 

In the laboratory, \[C{{l}_{2}}\] is prepared by the action of \[Mn{{O}_{2}}\] or \[KMn{{O}_{4}}\] or\[{{K}_{2}}C{{r}_{2}}{{O}_{7}}\] on conc. \[HCl\] or a mixture of \[NaCl\] and Conc. \[{{H}_{2}}S{{O}_{4}}\]

\[Mn{{O}_{2}}+4HCl\xrightarrow{{}}MnC{{l}_{2}}+C{{l}_{2}}+2{{H}_{2}}O\]

 

\[2KMn{{O}_{4}}+16HCl\xrightarrow{{}}2KCl+2MnC{{l}_{2}}+5C{{l}_{2}}+8{{H}_{2}}O\]

 

\[{{K}_{2}}C{{r}_{2}}{{O}_{7}}+14HCl\xrightarrow{{}}2KCl+2CrC{{l}_{3}}+7{{H}_{2}}O+3C{{l}_{2}}\]

 

\[Mn{{O}_{2}}+2NaCl+3{{H}_{2}}S{{O}_{4}}\xrightarrow{{}}2NaHS{{O}_{4}}+MnS{{O}_{4}}+2{{H}_{2}}O+C{{l}_{2}}\]

 

\[2KMn{{O}_{4}}+10NaCl+13{{H}_{2}}S{{O}_{4}}\xrightarrow{{}}\]

 

\[10NaHS{{O}_{4}}+{{K}_{2}}S{{O}_{4}}+2MnS{{O}_{4}}+8{{H}_{2}}O+5C{{l}_{2}}\]

 

Other oxidising agents such as \[Pb{{O}_{2}},\,P{{b}_{3}}{{O}_{4}},\,CaOC{{l}_{2}},\,{{O}_{3}}\] etc. also react with \[HCl\] to liberate \[C{{l}_{2}}\].

 

(iii) Properties : It combines with metals and non metals to form chlorides. it decomposes water forming \[HCl\] and \[HClO\] (hypochlorous acid) which is unstable and decomposes giving nascent oxygen which is responsible for oxidising and bleaching action of chlorine.

 

\[C{{l}_{2}}+{{H}_{2}}O\xrightarrow{{}}HCl+HClO\]; \[HClO\xrightarrow{hv}HCl+[O]\]

 

Coloured matter \[+O\xrightarrow{{}}\]Colourless matter.

 

The bleaching action is permanent and colour is not restored on standing. However, it cannot be used for bleaching delicate articles such as straw, silk, wool etc. which are damaged by it.

 

\[C{{l}_{2}}\] oxidises \[B{{r}^{-}}\] and \[{{I}^{-}}\] ions to \[B{{r}_{2}}\] and \[{{I}_{2}}\] respectively.

 

\[C{{l}_{2}}+2{{X}^{-}}\xrightarrow{{}}2C{{l}^{-}}+{{X}_{2}}(X=Br\]or \[I)\].

 

It combines with alkalies forming hypochlorite and chlorate salts in cold and hot conditions respectively.

 

\[2NaOH(\text{dil}\text{.)}+C{{l}_{2}}\xrightarrow{\text{Cold}}NaCl+NaClO+{{H}_{2}}O\]

 

\[6NaOH(\text{Conc}\text{.})+3C{{l}_{2}}\xrightarrow{\text{Heat}}5NaCl+NaCl{{O}_{3}}+3{{H}_{2}}O\]

 

During these reactions, halogen is simultaneously reduced to \[{{X}^{-}}\] ion and is oxidised to either hypohalite \[(X{{O}^{-}})\] or halate \[(XO_{3}^{-})\] ion. Such reactions are called disproportionation reactions.

 

With slaked lime, \[C{{l}_{2}}\] gives bleaching powder \[(CaOC{{l}_{2}})\]

 

\[Ca{{(OH)}_{2}}+C{{l}_{2}}\xrightarrow{{}}CaOC{{l}_{2}}+{{H}_{2}}O\]

 

With ammonia, \[C{{l}_{2}}\] reacts as follows :

 

\[8N{{H}_{3}}(excess)+3C{{l}_{2}}\xrightarrow{{}}6N{{H}_{4}}Cl+{{N}_{2}}\uparrow \]

 

\[N{{H}_{3}}+3C{{l}_{2}}(excess)\xrightarrow{{}}NC{{l}_{3}}+3HCl\]

 

With \[S{{O}_{2}}\] and \[CO\], addition compounds are formed

 

\[S{{O}_{2}}(dry)+C{{l}_{2}}\xrightarrow{{}}S{{O}_{2}}C{{l}_{2}}\](Sulphuryl chloride)

 

\[CO+C{{l}_{2}}\xrightarrow{{}}COC{{l}_{2}}\](Carbonyl chloride or phosgene)

 

\[C{{l}_{2}}\] is strong oxidising agent. It oxidises \[FeC{{l}_{2}}\] to \[FeC{{l}_{3}}\], moist \[S{{O}_{2}}\] to \[{{H}_{2}}S{{O}_{4}},SO_{3}^{2-}\] to \[SO_{4}^{2-}\], thiosulphate to sulphate and sulphur.

 

(iv) Uses of chlorine : It is used in the manufacture of \[HCl,\,NaOCl\], bleaching powder, chlorates, vinyl chloride, insecticides such as DDT, chlorinated organic solvents like \[CHC{{l}_{3}},CC{{l}_{4}}\]. It is also used in sterilisation of drinking water, in the extraction of \[Au\] and \[Pt\] and as a bleaching agent for paper, pulp and textiles.

 

(3) Bromine

 

(i) Occurrence : It mainly occurs in sea water and salt lakes as \[NaBr,\,KBr\] and \[MgB{{r}_{2}}\].

 

(ii) Preparation : On a commercial scale, bromine is prepared either from sea water (containing \[NaBr,\,KBr\] and \[MgB{{r}_{2}}\]) or the mother liquor (containing \[MgB{{r}_{2}}\]) left after crystallisation of chlorides from carnallite. On passing \[C{{l}_{2}}\] gas through these solutions, bromides get oxidised to bromine which is cooled and condensed to \[B{{r}_{2}}\] liquid.

 

\[2B{{r}^{-}}+C{{l}_{2}}\xrightarrow{{}}2C{{l}^{-}}+B{{r}_{2}}\]

 

In the laboratory, bromine can be prepared by heating \[NaBr\] with \[Mn{{O}_{2}}\] and conc. \[{{H}_{2}}S{{O}_{4}}\].

\[2NaBr+Mn{{O}_{2}}+3{{H}_{2}}S{{O}_{4}}\xrightarrow{{}}2NaHS{{O}_{4}}+MnS{{O}_{4}}+2{{H}_{2}}O+B{{r}_{2}}\]

 

It is also obtained by adding \[HCl\] to a mixture containing potassium bromide and potassium bromate.

 

\[5KBr+KBr{{O}_{3}}+6HCl\xrightarrow{{}}6KCl+3B{{r}_{2}}+3{{H}_{2}}O\]

 

(iii) Properties : Bromine is a reddish brown heavy liquid.

 

Its reaction with water, oxidising and bleaching action, reaction with alkalies, \[N{{H}_{3}}\], metals and non metals are similar to that of chlorine. \[B{{r}_{2}}\] oxidises only iodide ions to \[{{I}_{2}}\]. Bromine water reacts with mercuric oxide to form mercury oxy bromide

 

\[2HgO\underset{\text{Bromine water}\,\,\,\,\,\,\,\,\,\,\,\,}{\mathop{+2B{{r}_{2}}+{{H}_{2}}O}}\,\xrightarrow{{}}\underset{\text{Mercury oxy bromide}}{\mathop{HgB{{r}_{2}}.HgO+}}\,2HBrO\]

 

(iv) Uses of bromine : The main use of bromine is in the manufacture of ethylene bromide which is used as an additive to leaded petrol. It is also used to prepare \[AgBr\], bromine water, dyes, drugs and benzyl bromide (an effective tear gas).

 

(4) Iodine

 

(i) Occurrence : It mainly occurs in sea weeds or alkali metal iodides. Caliche (crude chile salt petre) which is mainly sodium nitrate contains iodine as sodium iodate \[(NaI{{O}_{3}})\].

 

(ii) Preparation of iodine : On a commercial scale iodine is prepared from sea weeds and caliche.

 

(a) From sea weeds : Sea weeds (Laminaria variety) are dried, burnt and ash (called kelp constains about 1%\[{{I}_{2}}\] as iodides of alkali metals besides chlorides and sulphates) is extracted with hot water. sulphates and chlorides are separated by fractional crystallisation, the mother liquor is treated with \[C{{l}_{2}}\] gas or heated with \[Mn{{O}_{2}}\] and conc. \[{{H}_{2}}S{{O}_{4}}\] to liberate \[{{I}_{2}}\] which is cooled and condensed to give violet crystals.

 

\[2NaI+C{{l}_{2}}\xrightarrow{{}}2NaCl+{{I}_{2}}\]

 

\[2NaI+Mn{{O}_{2}}+3{{H}_{2}}S{{O}_{4}}\xrightarrow{{}}2NaHS{{O}_{4}}+MnS{{O}_{4}}+2{{H}_{2}}O+{{I}_{2}}\]

 

(b) From Caliche : The mother liquor left after crystallisation of \[NaN{{O}_{3}}\] is treated with \[NaHS{{O}_{3}}\] to liberate \[{{I}_{2}}\] from \[NaI{{O}_{3}}\].

\[2NaI{{O}_{3}}+5NaHS{{O}_{3}}\xrightarrow{{}}3NaHS{{O}_{4}}+2N{{a}_{2}}S{{O}_{4}}+{{H}_{2}}O+{{I}_{2}}\]

 

In the laboratory, \[{{I}_{2}}\] is prepared by heating a mixture of potassium iodide and \[Mn{{O}_{2}}\] with conc. \[{{H}_{2}}S{{O}_{4}}\].

\[2KI+Mn{{O}_{2}}+3{{H}_{2}}S{{O}_{4}}\xrightarrow{{}}2KHS{{O}_{4}}+MnS{{O}_{4}}+{{H}_{2}}O+{{I}_{2}}\]

 

(iii) Properties : It is a dark violet shining solid which sublimes on heating. It is least soluble in water. However, its solubility can be increased by adding \[10%KI\] solution due to the formation of \[I_{3}^{-}\] complex ion in which \[{{I}^{-}}\] ion acts as a lewis base (ligand) and \[{{I}_{2}}\] molecule behaves as a lewise acid (central atom) which accommodates lone pair of electrons donated by \[{{I}^{-}}\] ion in the antibonding sigma \[{{p}_{z}}\] molecular orbital.

 

\[{{I}_{2}}+{{I}^{-}}\xrightarrow{{}}I_{3}^{-}\] (complex ion)

 

The aqueous solution containing \[I_{3}^{-}\] complex ion has a brown colour. It is soluble in many organic solvents. Its solution in \[C{{S}_{2}},\,CHC{{l}_{3}}\] and \[CC{{l}_{4}}\] is violet while in strong donor solvents like alcohols, ethers and amines is brown.

 

With cold, dilute \[NaOH\], iodine gives hypoiodous acid

 

\[NaOH+{{I}_{2}}\xrightarrow{\text{Cold}}NaI+HIO\]

 

However, with hot, conc. solution of \[NaOH\], the reaction is similar to that of \[C{{l}_{2}}\] or \[B{{r}_{2}}\].

 

Iodine does not displace chlorine and bromine from chlorides and bromides respectively, but it displaces them from their oxy salts

 

\[2KCl{{O}_{3}}+{{I}_{2}}\xrightarrow{{}}2KI{{O}_{3}}+C{{l}_{2}}\]

 

\[2KBr{{O}_{3}}+{{I}_{2}}\xrightarrow{{}}2KI{{O}_{3}}+B{{r}_{2}}\]

 

With \[N{{a}_{2}}{{S}_{2}}{{O}_{3}}\], iodine solution is decolourised due to the formation of colourless iodide and tetrathionate ions.

 

\[2N{{a}_{2}}{{S}_{2}}{{O}_{3}}+{{I}_{2}}\xrightarrow{{}}2NaI+N{{a}_{2}}{{S}_{4}}{{O}_{6}}\]

 

With ammonia it reacts as follows

 

\[2N{{H}_{3}}+3{{I}_{2}}\xrightarrow{{}}\underset{\text{(explosive)}}{\mathop{N{{I}_{3}}.N{{H}_{3}}}}\,+3HI\]

 

\[8N{{I}_{3}}.N{{H}_{3}}\xrightarrow{{}}5{{N}_{2}}+9{{I}_{2}}+6N{{H}_{4}}I\]

 

With strong oxidising agents such as \[HN{{O}_{3}},{{O}_{3}}\] and \[C{{l}_{2}}\], iodine gives iodic acid \[(HI{{O}_{3}})\]

 

\[{{I}_{2}}+10HN{{O}_{3}}\xrightarrow{{}}2HI{{O}_{3}}+10N{{O}_{2}}+4{{H}_{2}}O\]

 

\[{{I}_{2}}+{{H}_{2}}O+{{O}_{3}}\xrightarrow{{}}2HI{{O}_{3}}+5{{O}_{2}}\]

 

\[{{I}_{2}}+5C{{l}_{2}}+6{{H}_{2}}O\xrightarrow{{}}2HI{{O}_{3}}+10HCl\]

 

(iv) Uses of iodine : It is used to prepare tincture of iodine (2% solution of \[{{I}_{2}}\] in alcohol), iodex, iodoform, \[KI\], iodised salt (which contains \[KI\] or \[NaI\], 0.5 g per \[kg\] of \[NaCl\]) and as a laboratory reagent.

 

(5) Hydrogen halides : All the halogens combine with hydrogen to form hydrogen halides \[(HX)\].

 

(i) Preparation of HF and HCl : These are prepared by heating fluorides and chlorides respectively with conc. \[{{H}_{2}}S{{O}_{4}}\].

 

\[Ca{{F}_{2}}+{{H}_{2}}S{{O}_{4}}\xrightarrow{\text{Heat}}CaS{{O}_{4}}+2HF\]

 

\[2NaCl+{{H}_{2}}S{{O}_{4}}\xrightarrow{\text{Heat}}N{{a}_{2}}S{{O}_{4}}+2HCl\]

 

(ii) Preparation of HBr and HI :These are prepared by heating bromides and iodides respectively with phosphoric acid

 

\[3NaBr+{{H}_{3}}P{{O}_{4}}\xrightarrow{\text{Heat}}N{{a}_{3}}P{{O}_{4}}+3HBr\]

 

\[3NaI+{{H}_{3}}P{{O}_{4}}\xrightarrow{\text{Heat}}N{{a}_{3}}P{{O}_{4}}+3HI\]

Conc. \[{{H}_{2}}S{{O}_{4}}\] cannot be used for the preparation of \[HBr\] and \[HI\] because these being strong reducing agents reduced \[{{H}_{2}}S{{O}_{4}}\] to \[S{{O}_{2}}\] and are themselves oxidised to \[B{{r}_{2}}\] and \[{{I}_{2}}\] respectively.

 

\[2HBr+{{H}_{2}}S{{O}_{4}}\xrightarrow{{}}S{{O}_{2}}+B{{r}_{2}}+2{{H}_{2}}O\]

 

(6) Bleaching powder is obtained by the action of chlorine on dry slaked lime (Hasenclever method).

 

\[Ca{{(OH)}_{2}}+C{{l}_{2}}\xrightarrow{313K}CaOC{{l}_{2}}+{{H}_{2}}O\]

 

An aqueous solution of bleaching powder gives tests for \[C{{l}^{-}}\] and \[Cl{{O}^{-}}\]ions. On long standing, it undergoes auto-oxidation to form calcium chlorate. However, when heated, in presence of \[CoC{{l}_{2}}\], it gives \[{{O}_{2}}\]

 

\[6CaOC{{l}_{2}}\xrightarrow{{}}5CaC{{l}_{2}}+Ca{{(Cl{{O}_{3}})}_{2}}\]

 

\[2CaOC{{l}_{2}}\xrightarrow{CoC{{l}_{2}}}2CaC{{l}_{2}}+{{O}_{2}}\]

 

It is used for bleaching cotton, wood pulp etc., as a disinfectant, as a germicide for sterilization of drinking water, in the manufacture of chloroform and for making wood unshrinkable.

Other Topics


You need to login to perform this action.
You will be redirected in 3 sec spinner