+2 UNIT 8 PAGE-3
GROUP 16 ELEMENTS
Oxygen, sulphur, selenium, tellurium and polonium constitute the group 16 of the periodic table of elements. The first four elements are collectively called chalcogens. The name derives from the Greek word for bronze and points to the association of sulphur and its congeners with copper. Polonium is radioactive and derives its name from Poland Sulphur
Atomic and physical properties
The atomic and physical properties of Group 16 elements are listed in TABLE.
Property | Oxygen | Selenium | Tellurium | Polonium | |
Atomic number | 8 | 16 | 34 | 52 | 84 |
Atomic mass | 12.01 | 28.09 | 72.60 | 118.71 | 207.2 |
Electronic configuration | [He]2s22p4 | [Ne]3s23p4 | [Ar]3d104s24p4 | [Kr]4d105s25p4 | [Xe]4f145d106s26p4 |
Covalent radius (pm) | 74 | 103 | 119 | 142 | 168 |
Ionic radus M2- (pm) | 140 | 184 | 198 | 221 | 230 |
Ionisation enthalpy(kJ/mol) I II | 1314 3388 | 1000 2251 | 941 2045 | 869 1790 | 813 - |
Electronegativity | 3.50 | 2.44 | 2.48 | 2.01 | 1.76 |
Density (g/cm3) at 298K | 1.32 | 2.06 (at melting point) | 4.19 (hexagonal grey allotrope) | 6.25 | - |
Melting point (K) | 54 | 393 (monoclinic form at 673 K) | 490 | 725 | 520 |
Boiling point (K) | 90 | 718 | 958 | 1260 | 1235 |
The valence shell, electronic configuration of the Group 16 elements is ns2np4, with each element two electron short of the noble gas configuration. The ionisation enthalpy decreases while the atomic and ionic radii increase as we descend the group. Unlike oxygen, which exists as as a diatomic gas , the other Group 16 elements exist as solids.
Occurrence , Extraction and uses
Selenium and tellurium are less abundant than sulphur (0.05 ppm in earth’s crust for selenium 0.002 ppm for Te) and
occurs as selenides and tellurides in sulphide ores. The principal source of Se and Te is the ‘anode slime’ deposited during the electrolytic refining of Cu.
Extraction of sulphur
Frash process for extraction of sulphur
Uses of S, Se and Te
Most of the sulphur produced is used for the manufacture of sulphuric acid and other industrially important sulphur compounds. The most important application of selium is as a photoconductor in photocopying (Xerox) machines, though the major use is as decoluriser of glass. Tellurium is mostly used as an additive in metallurgy (manufacture of iron and steel , non-ferrous metals and alloys). Tellurium and polonium are highly toxic , the latter more so because of its intense radioactivity.
Allotropy of sulphur and Selenium
S8 ring of rhombic sulphur
S6 allotrope
Several other modifications of sulphur containing 6 – 20 sulphur atoms per ring have been synthesised in the last two decades. In cyclo-S6 , the ring adopts the chair form and the molecular dimensions are shown in Fig (b above). In addition , various chain polymers called catena-Sn are known. Unstable small molecules , Sn ( n = 2 – 5) exist in liquid sulphur at elevated temperatures and in sulphur vapour. At 100 K , S2 is the dominant species. Like dioxygen (O2), S2 is paramagnetic.
Selenium exists in eight allotropic forms of which three are red monoclinic forms containing Se8 rings. The thermodynamically stable form is grey hexagonal ‘metallic’ selenium which consists of polymeric helical chains. The common form of the element is the amorphous black selenium. Grey selenium is the only allotrope of selenium which conducts electricity. Tellurium has only one crystalline form with a chain structure similar to that of grey Se.
Problem
04. Why does sulphur in vapour state exhibit paramagnetic behaviour ?
Oxidation states and Trends in Chemical Reactivity
The important oxidation states observed for S, Se and Te are -2, +2, +4 and +6. The chalcogenide dianions ( E2- ) exist only in their compounds with most electropositive elements. By and large, chalcogens with exception of polonium , display non-metallic covalent chemistry. Some trends observed in the chemistry of Group 16 elements may be summarised as follows :
(a) After oxygen, there is a steep drop in electronegativity for S, Se, Te and Po and consequently their compounds have less ionic character.
(b) The metallic character increases as we descend the group. Thus, sulphur is a typical non-metal and is an insulator ; Se and Te are metalloids and are semiconductors. Polonium shows metallic character.
(c) The thermal stability of the hydrides decreases in the order :
H2O > H2S > H2Se > H2Te > H2Po
(d) The tendency for catenation decreases appreciably as we move down the group.
(e) Sulphur
(f) The tendency to engage in multiple bonding decreases as we ascend the group. Thus S = C = S is moderately stable : Se = C = Se decomposes readily ; Te = C = Te is unknown.
HYDRIDES
Group 16 elements form binary hydrides H2R ( R = O, S, Se, Te or Po ).
Preparation
Water is prepared by burning hydrogen in an atmosphere of oxygen, while the hydrides of S, Se and Te are readily obtained by the action of acids on metal sulphides, selenides and tellurides, e.g.,
K2Se + H2SO4 ® K2SO4 + H2Se
Properties
(a) Thermal stability of these hydrides decreases from oxygen to polonium.
(b) Unlike water , the other hydrides are unpleasant foul smelling poisonous gases.
(c) The volatility increases markedly from H2O to H2S and then decreases. The abnormally high boiling point of water is due to the association of molecules through hydrogen bonding.
(d) All the hydrides, except water, are reducing agents. The reducing character increases on moving from H2S to H2Te.
(e) The bond angles H- R - H in H2O, H2S, H2Se and H2Te are 104.5°, 92.5°, 90° and 89.0° respectively
(f) H2S , H2Se and H2Te are weak diprotic acids in aqueous solutions ; the acidity increases in the series
H2O < H2S < H2Se < H2Te < H2Po
(g) Metal sulphides except those of Groups 1 and 2 are sparingly soluble in water.
Some physical properties are listed below :
Property | H2O | H2S | H2Se | H2Te |
m.p (K) | 273 | 188 | 208 | 222 |
b.p (K) | 373 | 213 | 232 | 269 |
M-H distance(pm) | 104 | 92 | 91 | 90 |
MHH angle(°) | 104.5 | 92.5 | 90 | 89 |
DfH° (kJ/mol) | - 286 | 20 | 73 | 100 |
E(M-H) (kJ/mol) | 463 | 347 | 276 | 238 |
Dissociation Constant(298K) HM- K1 M2- | 1.8 x 10-16 - | 1.3 x 10-7 1.8 x 10-15 | 1.3 x 10-4 1.8 x 10-11 | 2.3 x 10-3 1.6 x 10-11 |
HALIDES
(a) Sulphur
Halides of Group 16 elements
Elements | Fluorides | Chlorides | Bromides | Iodides |
Oxygen | F2O2, F2O | Cl2O, ClO2 Cl2O6, Cl2O7 | Br2O, BrO2, BrO3 | I2O, I2O4, I2O5 |
S2F2, SF4,SF6, S2F10 | S2Cl2, SCl2, SCl4 | S2Br2 | - | |
Selenium | Se2F2, SeF4, SeF6 | Se2Cl2, SeCl2 | Se2Br2, SeBr4 | - |
Tellurium | TeF4, Te2F10, TeF6 | TeCl2, TeCl4 | TeBr2, TeBr4 | TeI4 |
Polonium | - | PoCl2, PoCl4 | PoBr2, PoBr4 | PoI4 |
Since florine is more electronegative than oxygen, its compounds with oxygen are called fluorides.
(b) The stability of the halides decreases in the order:
F > Cl > Br > I.
The highest oxidation state is realised only in the fluorides. With
iodine , only a tetraiodide , TeI4 is known.
(a) The general preparative routes for chalcogen halides involve the direct reaction of the chalcogen with the respective halogen. Some typical reactions are given below :
(1/8) S8(s) + 3 F2(g) ® SF6(g)
(1/4) S8(ℓ) + Cl2(g) ® S2Cl2(ℓ)
Te(s) + 2 Cl2(g) ® TeCℓ4
Te(s) + 2 I2 ® TeI4 (s)
(b) Direct fluorination of elemental sulphur yields mainly SF6 along with traces of SF4 . SF4 is prepared by the fluorination of SCl2 with NaF in acetonitrile at 350 K.
(c) SF6 is inert, non-toxic gas at room temperature. Its inertness is due to the presence of sterically protected sulphur atom which does not allow thermodynamically favourable reactions like hydrolysis to take place. In contrast, the less sterically hindred SeF6 and SF4 undergo hydrolysis readily. Because of its inertness and good electrical properties, SF6 is used as gaseous insulator in high voltage generators. Both SF4 and SeF4 are used as fluorinating agents for the conversion of –COOH into –CF3 and C=O and P=O groups into CF2 and PF2 groups.
(d) The structure of SF4 and SF6 are shown in Fig.
Structure of (a) SF4 (b) SF6(e) Compared to sulphur halides, the halides of Se and Te , in +4 oxidation state , adopt oligomeric or polmeric structures. Thus SeCl4 , SeBr4, TeCl4 , TeBr4 and TeI4 exist as tetramers while TeF4has a polymeric structure.
OXIDES
The elements of Group 16 form several oxides as shown in the TABLE.
Name | Selenium | Tellurium | Polonium | |
Monoxide | SO | - | TeO | PoO |
Dioxide | SO2 | SeO2 | TeO2 | PoO2 |
Trioxide | SO3 | SeO3 | TeO3 | - |
Heptoxide | S2O7 | - | - | - |
The oxides of sulphur are the most important commercially for the manufacture of sulphuric acid.
The oxides result from the direct union of the elements :
S8(s) + 8 O2 ® 8 SO2 (g)
SO2 is a gas at room temperature and has an angular structure with a bond angle 119° and S-O distance of 143 pm. It exists as discrete SO2 molecules even in the solid state. SeO2 in the gas phase has a structure similar to that of SO2 but but in the solid state has an infinite covalent structure (Fig). TeO2 in the solid state has a layer structure consisting of TeO4 units.
The structures of SeO2 in the gas and solid phases.
SO3 in the gas phase exists as planar triangular molecular species (Fig).
The structures of (a) gaseous SO3 (b) cyclic trimer of solid SO3 (c) linear chain form of solid SO3 and (d) cyclic tetrameric form of SeO3.
In solid state , SO3 exists in several modifications having a linear cyclic tetrameric or a polymeric chain structures (Fig above). SeO3 , on the other hand , exists as a cyclic tetramer (Se4O12) in solid (Fig above). TeO3 is also a solid with a net work structure in which TeO6 octahedra share all vertices.
OXOACIDS OF SULPHUR
Name | Formula | O.S | Structure |
Sulphurous acid | H2SO3 | +4 |  |
Sulphuric acid | H2SO4 | +6 |  |
Thiosulphuric acid | H2S2O3 | +2 |  |
Dithionic acid | H2S2O6 | +5 |  |
Polythionic acid | H2SnO6 | +6 |  |
Peroxomono sulphuric acid | H2SO5 | +5 |  |
Peroxodi- sulphuric acid | H2S2O8 | +6 |  |
Pyrosulphuric acid (Disulphuric acid) | H2S2O7 | +6 |  |
Description of some of the oxo-acids are given below :
1. Sulphurous acid H2SO3
The acid is obtained by dissolving SO2 in water, the reaction being represented as :
The acid is known only in solution. It is a fairly strong acid and being diprotic , it ionises in two stages as :
The acid acts as a strong reducing agent on account of the ease with which it gets oxidised to H2SO4 even by atmospheric oxygen.
Sulphurous acid (or sulphur dioxide in prsence of water) acts as a bleaching agent. The bleaching action is due to its reducing properties.
The structure of the unionised acid is in doubt. One of the forms suggested has pyramidal structure as shown.
Sulphurous acid
In this case sulphur utilises sp3 hybridised orbitals for sigma bonding and one of the sp3 hybrid is occupied by a lone pair of electrons.
SULPHURIC ACID H2SO4
Sulphuric acid is is manufactured by contact process which involves three stages :
(a) Burning of sulphur or sulphide ores in air to generate SO2.
(b) Coversion of SO2 to SO3 by reaction with oxygen in presence of a catalyst , and
(c) Absorption of SO3 in H2SO4 to give oleum which is formulated as H2S2O7.
A flow diagram for the manufacture of sulphuric acid is shown in Fig.
1. Sulphur burners : Sulphur
S + O2 ® SO2
4 FeS2 + 11 O2 ® 2 Fe2O3 + 8 SO2
2. Purification unit : The gaseous mixture coming out of sulphur burners is generally impure. The gases are purified as follows :
(a) Dust chamber : Steam is introduced to remove dust particles.
(b) Coolers : The hot gases are cooled to about 373 K by passing them through cooling pipes.
(c) Scrubber : Gases are introduced into a washing tower (packed with quartz ) also known as scrubber which dissolves mist and any other soluble impurities.
(d) Drying tower : It absorbs water vapours from gases and they become completely dry.
(e) Arsenic purifier : This is a small chamber fitted with shelves containing freshly precipitated ferric hydroxide. The impurities of arsenic oxide present in the gases are absorbed by ferric hydroxide.
3. Testing Box : The gases coming out of purification unit are tested in this box with the help of a strong beam of light. If some impurities are present, they will scatter light and the path will become visible. In case the gases are impure, they are passed through the purifying unit again.
4. Contact chamber converter : The pure gases are then , heated to about 723 – 823 K in a preheater. These are then introduced in the contact chamber. It is a cylindrical iron chamber fitted with iron pipes. Each pipe is packed with the catalyst. In this chamber sulphur dioxide is oxidised to sulphur trioxide.
2 SO2 + O2 ® 2 SO3 ; DH = - 196.6 kJ
As the forward reaction is exothermic, the preheating of incoming gases is stopped once the oxidation reaction is started. The heat produced in the reaction is sufficient to maintain the temperature of the reaction.
5. Absorption tower : It is cylindrical tower packed with acid proof flint. Sulphur Sulphur
H2SO4 + SO3 ® H2S2O7
oleum
Oleum is then diluted with calculated amount of water to get acid of desired concentration.
H2S2O7 + H2O ® 2 H2SO4
It may be noted that sulphur trioxide is not directly absorbed in water to form sulphuric acid because the process is accompanied by the formation of dense fog of the acid particles. Therefore it becomes quite inconvenient for the workers.
Properties and Reactions of Sulphuric acid
Sulphuric acid is a colourless , dense, oily liquid with specific gravity of 1.84 g cm-3 at 298 K. The acid freezes at 283 K and boils at 590 K. It dissolves in water with the evolution of a large quantity of heat. Hence , care must be taken while preparing sulphuric acid solution from concentrated sulphuric acid. The concentrated acid must be added slowly into water with constant stirring.
The chemical reactions of sulphuric acid are a result of the following characteristics : (a) low volatility (b) strong acid character (c) strong affinity for water and (d) ability to act as an oxidising agent. In aqueous solution, sulphuric acid ionizes in two steps :
H2SO4 ® H+(aq) + HSO4-(aq) : K1 = very large
HSO4-(aq) ® H+(aq) + SO42-(aq) : K2 = 1.2 x 10-2
The acid forms two series of salts ; normal sulphates (such as sodium sulphate and copper sulphate) and acid bisulphates (e.g. sodiumbisulphate). The sulphate ion , SO42- is tetrahedral with an S - O bond length 149 pm.
Sulphuric acid , because of its low volatility , can be used to manufacture more volatile acids from corresponding salts.
2 M X + H2SO4 ® 2 HX + M2SO4 ( X = F-, Cl- , NO3- )
Concentrated sulphuric acid is a strong dehydrating agent. Many wet gases can be dried by passing these through sulphuric acid, provided the gases do not react with the acid. Sulphuric acid removes water from organic compounds as shown by its charring action on carbohydrates.
Hot concentrated sulphuric acid is a moderately strong oxidising agent. In this respect , it is itermediate between phosphoric and nitric acids. Both metals and non-metals are oxidised by concentrated sulphuric acid, which is reduced to SO2.
With zinc , which is a strong reducing agent than copper, the reduction of concentrated sulphuric acid goes further to give sulphur or H2S.
Uses
Sulphuric acid is the most important heavy chemical and finds extensive use in industry and laboratory. It is used :
(i) In fertliser industry : It is used in the preparation of ammonium
phosphate, superphosphate of lime , etc.
(ii) In petroleum refining to remove unsaturated compounds.
(i) In the manufacture of important chemicals such as hydrochloric acid, nitric acid, sulphates of metals, alums, ethers etc.
(ii) In metallurgical processes for the purification of metals by electrolysis where sulphuric acid is commonly used in the bath.
(iii) In preparing paints and pigments. Many of the pigments used in paints are sulphates, e.g., barium and lead sulphates.
(iv) In the manufacture of explosives such as nitroglycerine, gun cotton , T.NT , etc.
(v) It is used as a drying and dehydrating agent.
(vi) It is used for cleaning the surfaces of metals (picking) before electrolysis.
GROUP 17 ELEMENTS
THE HALOGEN FAMILY
The Group 17 of the periodic table contains five elements : Fluorine (F), chlorine (Cl) , bromine (Br), iodine (I) and astatine (At). These are named as halogens. The name halogen is derived from the Greek word meaning sea salt producers because the first three members occur as salts (chlorides, bromides and iodides) in sea water. Halogens are among the most reactive non-metals. There are greater similarities within the halogen group compared to other groups in in the periodic table with the exception of the alkali metals (group 1). Astatine is radioactive and is therefore , not of any practical importance.
Atomic and Physical Properties
The important atomic and molecular properties of halogens are given in Table.
Property | Fluorine | Chlorine | Bromine | Iodine | Astatine |
Atomic number | 9 | 17 | 35 | 53 | 85 |
Atomic mass | 19.00 | 35.45 | 79.90 | 126.90 | 210 |
Electronic configuration | [He]2s22p5 | [Ne]3s23p5 | [Ar]3d104s24p5 | [Kr]4d105s25p5 | [Xe]4f145d106s26p5 |
Covalent radius (pm) | 64 | 99 | 114 | 133 | - |
Ionic radus X- (pm) | 133 | 184 | 196 | 220 | - |
Ionisation enthalpy(kJ/mol) | 1680 | 1256 | 1142 | 1008 | - |
Electronegativity | 4 | 3.2 | 3.0 | 2.7 | 2.2 |
Density (g/cm3) | 1.51 | 1.66 | 3.19 | 3.19 | - |
Melting point (K) | 54.4 | 172.0 | 265.8 | 386.6 | - |
Boiling point (K) | 84.9 | 239.0 | 332.5 | 458.2 | - |
Distance X -X (pm) | 143 | 199 | 228 | 266 | - |
Enthalpy of dissociation X2(g) ® 2 X(g) (kJ/mol) | 158.8 | 242.6 | 192.8 | 151.1 | - |
Eo (V) | 2.87 | 1.36 | 1.09 | 0.54 | - |
1. Atomic and ionic radii
The halogens have smallest atomic radii in their respective periods due to the maximum effective nuclear charge. Among themseves , the atomic and ionic radii increase with increase in atomic number. This is due to increase in the number of electron shells.
2. Ionisation enthalpies
The ionisation enthalpies of halogens are very high. This indicates that they have very little tendency to lose electrons. However, on going down the group from fluorine to astatine, the ionisation enthalpy decreases. This is due to gradual increase in atomic size which is maximum with iodine. Consequently, it has the least ionisation enthalpy in the family.
3. Melting and boiling points
The melting and boiling points of halogens increase with increase in atomic number as we go down in the group.
Explanation : The forces existing between these molecules are weak van der Waal’s forces which increase down the group. This is also clear from the change of state from fluorine to iodine. At room temperature, fluorine and chlorine are gases , bromine is a liquid while iodine and astatine are solids.
4. Electron gain enthalpies
(i) All these have maximum electron gain enthalpies in their respective periods. This is due to the fact that the atoms of these elements have only one electron less than the stable noble gas (ns2np6) configurations. Therefore , they have maximum tendency to accept an additional electron.
(ii) In general, electron gain enthalpy decreases from top to bottom in a group. This is due to the fact that the effect of increase in atomic size is much more than the effect of increase in nuclear charge and thus, the additional electron feels less attraction by the large atom. Consequently, electron affinity decreases.
(iii) Fluorine has unexpectedly less electron gain enthalpy than chlorine. Therefore, chlorine has the highest electron gain enthalpy in this group. The lower electron gain enthalpy of fluorine as compared to chlorine is due to small size of the fluorine atom. As a result, there are strong inter-electronic repulsions in relatively small 2p subshell of fluorine and thus the incoming electron does not feel much attraction. Therefore , its electron gain enthalpy is small.
Thus , electron gain enthalpy among halogens varies as:
F < Cl > Br > I
5. Electronegativity
Halogens have large electronegativity values. The values decreases down the group from fluorine to iodine because the atomic size increases and the effective nuclear charge decreases. Fluorine is the most electronegative element in the periodic table.
6. Metallic and non-metallic character
Because of very high ionisation energy values, all halogens are non-metallic in character. The non-metallic character decreases as we go down the group. Therefore, the last element, iodine is a solid with a metallic lustre and forms positive ions such as I+ and I3+.
7. Colour
Fluorine and chlorine are gases with pale yellow and greenish yellow colours respectively. Bromine is a deep reddish-brown liquid with a high vapour pressure. Iodine is a lustrous greyish black crystalline solid which sublimes readily when heated to form a deep violet vapour.
Explanation : The colour of halogens is due to the fact that their molecules absorb radiations from visible light and the outer electrons are easily excited to higher energy levels. The amount of energy required for excitation depends upon the size of the atom. Fluorine atom is the smallest and the force of attraction between the nucleus and outer electrons is very large. As a result, it requires large excitation energy and absorbs violet light (high energy) and therefore , appears pale yellow. On the other hand, iodine needs very less excitation energy and absorbs yellow light of low energy. Thus it appears dark violet. Similarly, we can explain the greenish yellow colour of chlorine and reddish brown colour of bromine.
OXIDATION STATES
Halogens have one electron less than the next noble gas. Therefore, they can get the noble gas configuration either by gaining one electron to form uni-negative ion, X-, or by sharing electrons with other atoms. Thus, they show an oxidation state of - 1 or +1. Since fluorine is the most electronegative element , it always shows an oxidation state of - 1. It does not show any positive oxidation state.
The other elements also show positive oxidation states of +1, +3, +5 and +7. The higher oxidation states of chlorine, bromine and iodine are due to the presence of d-orbitals in their valency shells. As a result the outer s- and p-electrons can easily be promoted to the vacant d-orbitals as shown below :
Thus, the halogens exhibit the following oxidation states :
Element | F | Cl | Br | I |
Oxidation states | - 1 | - 1, +1, +3, +5, +7 | - 1, + 1 , +3 , +5 , +7 | - 1 , +1 , + 3 , + 5 , + 7 |
TRENDS IN CHEMICAL REACTIVITY
The halogens are the most reactive elements as a family. Fluorine is the most reactive of all the halogens. The reactivity of the halogens decreases down the group . The high reactivity of halogens is due to the following reasons.
(i) Low dissociation energies
All the halogens have very low dissociation energies. As a result, they can redily dissociate into atoms and react with other substances. As shown below , the dissociation energies of halogens are quite low in comparison to common molecules as H2, O2 and N2.
Molecule | F2 | Cl2 | Br2 | I2 | H2 | O2 | N2 |
Dissociation energy (kJ/mol) | 159 | 243 | 193 | 151 | 458 | 495 | 941 |
Fluorine is having a smaller enthalpy of dissociation than chlorine. In other words , the F – F bond energy is smaller than that of Cl – Cl , whereas , the X – X bond energies from chlorine onwards show the expected trend :
Cl – Cl > Br – Br > I – I .
A reason for this anomaly is the relatively larger electron – electron repulsions of lone pairs in F2 where they are much closer to each other than in Cl2.
(ii) Higher electron gain enthalpy : Halogens have very high electron affinity values and therefore have very strong tendency to gain electron. Thus, halogens are very reactive elements due to their low dissociation energies and high electron affinity values. As is clear from the values of bond dissociation energies , fluorine has the lowest bond dissociation energy. This is due to weak F– F bond because of the repulsion between the non-bonding electrons in the small molecule. Therefore, it is most reactive among the halogens.
The reactivity of halogens can also be understood in terms of their oxidising power. Halogens have high electron affinity values and therefore they have strong tendency to take up electron :
½ X2 + e-® X -
As a result, they act as powerful oxidising agents. Fluorine is the strongest oxidising agent and oxidises other halide ions in solution or even in solid phase. In general, a halogen of lower atomic number will oxidise halide ion of higher atomic number and therefore , will liberate them from their salt solutions as given below :
F2 + 2 X- ® 2 F -+ X2 ( X = Cl, Br , I )
Cl2 + 2 X- ® 2 Cl -+ X2 ( X = Br , I )
Br2 + 2 X-® 2 Br -+ I2 ( X = I )
The decreasing oxidising power of the halogens as we go down the group is shown by their reduction potentials.
Species | F | Cl | Br | I |
Eo(Volts) | 2.87 | 1.36 | 1.09 | 0.52 |
Occurrence and Isolation
Halogens are very reactive and hence do not occur in free or elemental state in nature. They are found mainly in the form of metal halides although iodine also occurs as iodate (IO3- ). Chlorine is the most abundant of the halogens and its main commercial source is NaCl (common salt) , which occurs in vast quantities in the oceans , salt lakes and rock beds. The sources of halogens and their abundances are shown in TABLE.
Halogen | Abundance | Main source | |
Crustal rocks | Oceans | ||
Fluorine | 5.44 x 10-2 | 7.0 x 10-6 | Fluorite (CaF2), Fluorapatite [Ca5(PO4)3F] |
Chlorine | 1.29 x 10-2 | 1.9 | Sea water, salt wells, salt beds (NaCl, KCl, MgCl2, CaCl2) |
Bromine | 0.25 x 10-3 | 6.5 x 10-3 | Sea water, salt lakes(NaBr, KBr, MgBr2) |
Iodine | 4.6 x 10-5 | 5.0 x 10-6 | Brine wells, sea weeds(I- ), |
Because fluorine is the strongest chemical oxidising agent, only practical method of preparing elemental fluorine is by electrolytic method developed by Moissan in 1886. Electrolysis of a molten mixture of potassium fluoride and anhydrous hydrogen fluoride at 350 K using mild steel and carbon rod free from graphite as cathodes and anodes respectively gives elemental fluorine :
Cathode : 2 H+ + 2 e-® H2(g)
Anode : 2 F- ® F2(g) + 2 e-
A diaphragm made of teflon is used to separate the cathode and anode to avoid the accidental explosive mixing of hydrogen and florine.
Chlorine is prepared industrially by the electrolysis of natural brine or aqueous solutions of NaCl. Sodium hydroxide and hydrogen are also produced in this method.
In the laboratory , chlorine can be prepared by the oxidation of HCl by MnO2 or KMnO4.
MnO2(s) + 4 HCl(aq) ® MnCl2(aq) + 2 H2O(ℓ) + Cl2(g)
2 KMnO4(s) + 16 HCl (aq) ® 2 KCl(aq) + 2 MnCl2(aq)
+ 8 H2O(ℓ)+ Cl2(g)
Bromine is commercially prepared by the oxidation of bromide ions in natural brine with chlorine :
2 Br- (aq) + Cl2(g) ® Br2 (ℓ) + 2 Cl - (aq)
Similarly, iodine is manufactured from natural brines or sea weed by the oxidation of iodide ion with chlorine. Iodine is recovered from iodates (occuring in chile salt petre) by reduction with sodium hydrogen sulphite, NaHSO3 by the sequence of reactions :
NaIO3(aq) + 3 NaHSO3(aq) ® NaI(aq) + 3 NaHSO4(aq)
NaIO3(aq) + 5 NaI (aq) + 3 H2SO4(aq) ® 3I2 (s)+ 3 H2O (ℓ)
+ 3 Na2SO4(aq)
Chemical method for preparation of fluorine
Recently, a chemical method for preparation of fluorine has been developed. This involves the following reaction :
K2MnF6 + 2 SbF5 ® 2 KSbF6 + MnF3 + ½ F2
In this reaction, the stronger Lewis acid SbF5 displaces the weaker one, MnF4 from its salt. MnF4 is unstable and readily decomposes to give MnF3 and fluorine.
Problem
05. Chlorine vapours are evolved when concentrated sulphuric acid is added to a mixture of sodium chloride and manganese dioxide. Write a balanced equation for the reaction.