CHEMISTRY
p-Block Elements (Halogens and Noble Gases) Introduction : Group 13 to 18 of the periodic table of elements constitute the p–block. The p–block contains metals, metalloids as well as non–metals. The p–block elements have general valence shell electronic configuration ns2 np1–6. The first member of each group from 13–17 of the p–block elements differ in many respects from the other members of their respective groups because of small size, high electronegativity and absence of d–orbitals. The first member of a group also has greater ability to form p–p multiple bonds to itself (e.g. C=C, CC, NN) and to element of second row (e.g C=O, C=N, CN, N=O) compared to the other members of the same group. The highest oxidation of p–block element is equal to the group number –10. Moving down the group, the oxidation state two less than the highest group oxidation state becomes more stable in groups 13 to 16 due to inert pair effect (reluctance of s-subshell electrons to participate in chemical bonding)
GROUP 17 ELEMENTS : THE HALOGEN FAMILY Fluorine, chlorine, bromine, iodine and astatine are members of Group 17. These are collectively known as the halogens (Greek halo means salt and genes born i.e., salt producers). The halogens are highly reactive non-metallic elements. Electronic Configuration
All these elements have seven electrons in their outermost shell (ns2 np5) which is one electron short of the next noble gas. Atomic and Ionic Radii The halogens have the smallest atomic radii in their respective periods due to maximum effective nuclear charge . Atomic and ionic radii increase from fluorine to iodine due to increasing number of quantum shells. Ionisation Enthalpy They have little tendency to lose electron. Thus they have very high ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy decreases down the group. Electron Gain Enthalpy Halogen have maximum negative electron gain enthalpy in the corresponding period. This is due to the fact that the atoms of these elements have only one electron less than stable noble gas configurations. Electron gain enthalpy of the elements of the group becomes less negative down the group. However, the negative electron gain enthalpy of fluorine is less than that of chlorine. It is due to small size of fluorine atom. As a result, there are strong interelectronic repulsions in the relatively small 2p orbitals of fluorine and thus, the extra electron (incoming) does not experience much attraction. Electronegativity They have very high electronegativity. The electronegativity decreases down the group. Fluorine is the most electronegative element in the periodic table Physical Properties Fluorine and chlorine are gases, bromine is a liquid whereas iodine is a solid. Their melting and boiling points steadily increase with atomic number. All halogens are coloured. This is due to absorption of radiations in visible region which results in the excitation of outer electrons to higher energy level. By absorbing different quanta of radiation, they display different colours. For example, F2, has yellow, Cl2, greenish yellow, Br2, red and I2, violet colour. Fluorine and chlorine react with water. Bromine and iodine are only sparingly soluble in water. But are soluble in organic solvents such as chloroform, carbon tetrachloride, carbon disulphide and hydrocarbons to give coloured solutions. Except the smaller enthalpy of dissociation of F2 compared to that of Cl2. The X-X bond disassociation enthalpies from chlorine onwards show the expected trend : Cl – Cl > Br – Br > F – F > I – I. The reason for the smaller enthalpy of dissociation of F2 is the relatively larger electronselectrons repulsion among the lone pairs in F2 molecule where they are much closer to each other than in case of Cl2.
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CHEMISTRY Table : 1 Atomic and physical properties Element
F
Cl
Br
I
Atomic Number
9
17
35
53
Atomic Mass
19 2
Electronic configuration
[He] 2s 2p 64
Covalent Radius / pm –
35.45 5
2
79.90 5
10
126.90 2
5
10
[Kr] 4d 5s 5p
99
114
133
133
184
196
220
1256
1142
1008
– 333
– 349
– 325
– 296
143
199
229
266
158.8
242.6
192.8
151.1
Electronegativity
4
3.2
3.0
2.7
Melting point / K
54.4
172
265.8
386.6
Boiling point / K
84.9
239.0
332.5
458.2
–1
Ionization enthalpy / (kJ mol ) –1
Electron gain enthalpy /(kJ mol ) Distance X -X/pm Enthalpy of dissociation (X2)/kJ mol
–1
5
[Ar] 3d 4s 4p
1680
Ionic Radius X / pm
2
[Ne] 3s 3p
Chemical Properties Oxidation states and trends in chemical reactivity All the halogens exhibit –1 oxidation state. However, chlorine, bromine and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states also. The higher oxidation states of chlorine, bromine and iodine are realised mainly when the halogens are in combination with the small and highly electronegative fluorine and oxygen atoms e.g., in interhalogens, oxides and oxoacids. The fluorine atom has no d orbitals in its valence shell and therefore cannot expand its octet. Being the most electronegative, it exhibits only – 1 oxidation state. All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group. The ready acceptance of an electron is the reason for the strong oxidising nature of halogens. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the solid phase. The decreasing oxidising ability of the halogens in aqueous solution down the group is evident from their standard electrode potentials. Fluorine oxidises water to oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids. The reactions of iodine with water is non- spontaneous . I– can be oxidised by oxygen in acidic medium; just the reverse of the reaction observed with fluorine. 2F2(g) + 2H2O() 4H+ (aq) + 4F– (aq) + O2(g) X2(g) + H2O () HX(aq) + HOX (aq) ; (where X = Cl or Br) 4I– (aq) + 4H+ (aq) + O2(g) 2 I2 (s) + 2H2O ()
Standard Reduction Potential (SRP) X2 + 2e– 2X– F2 + 2e– 2F–
° = + 2.87 V ; Cl2 + 2e– 2Cl–
° = + 1.36 V
Br2 + 2e – 2Br –
° = + 1.09 V ; 2 + 2e– 2–
° = + 0.54 V
More the value of the SRP, more powerful is the oxidising agent. Hence the order of oxidising power is F2 > Cl2 > Br2 > 2 Since SRP is the highest for F2 (among all elements of periodic table), it is a strongest oxidising agent. Example-1 Solution
Covalent radius of fluorine is 64 pm but the bond length is not equal to 128 pm and that is 143 pm and bond energy is found to be comparable to 2 . This may be attributed to.p – .p repulsions due to small size of F atom.
Example-2 Solution
Electron affinity of chlorine is more than F. Inspite of this F2 is the better oxidising agent. Why ? SRP of F2 is much higher than that of Cl2 on account of smaller bond dissociation energy.
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CHEMISTRY Hydration energy of X– Smaller the ion, higher is the hydration energy. F– Cl – 515 381
Br – 347
– 305
in kJ/mol
Anomalous behaviour of fluorine The anomalous behaviour of fluorine is due to its small size, highest electronegativity, low F- F bond dissociation enthalpy, and non availability of d orbitals in valence shell. Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with other elements). It forms only one oxoacid while other halogens form a number of oxoacids. Hydrogen fluoride is liquid (b.p. 293 K) due to strong hydrogen bonding. Other hydrogen halides are gases. (i)
Reactivity towards hydrogen: They all react with hydrogen to give hydrogen halides but affinity for hydrogen decreases from fluorine to iodine with increasing atomic number. They dissolve in water to form hydrohalic acids. The acidic strength of these acids increases in the order : HF < HCl < HBr < HI. The stability of these halides decreases down the group due to decrease in bond (H–X) dissociation enthalpy in the order : H – F > H – Cl > H –Br > H – I .
(ii)
Reactivity towards oxygen : Halogens form many oxides with oxygen but most of them are unstable. Fluorine forms two oxides OF2 and O2F2. However, only OF2 is the thermally stable at 298 K. These oxide are essentially oxygen fluorides because of the higher electronegativity of flurorine than oxygen . Both are strong fluorinating agents. O2F2 oxidises plutonium to PuF6 and the reaction is used in removing plutonium as PuF6 from spent nuclear fuel. Chlorine, bromine and iodine form oxides in which the oxidation states of these halogen vary from + 1 to + 7. A combination of kinetic and thermodynamic factors lead to the generally decreasing order of stability of oxides formed by halogens, I > Cl > Br. The higher oxides of halogens tend to be more stable than the lower ones. Chlorine oxides, Cl2O, ClO2, Cl2O6 and Cl2O7 are highly reactive oxidising agents and tend to explode. ClO2 is used as a bleaching agent for paper pulp and textiles and in water treatment. The bromine oxides, Br2O, BrO2, BrO3 are the least stable halogen oxides and exist only at low temperature. They are very powerful oxidising agents. The iodine oxides, I2O4, I2O5, I2O7 are insoluble solids and decompose on heating. I2O5 is very good oxidising agent and is used in the estimation of carbon monoxide.
(iii)
Reactivity towards metals : Halogen react with metals to form metal halides. For e.g., bromine reacts with magnesium to give magnesium bromide.
(iv)
Reactivity of halogen towards other halogens : Halogens combine amongst themselves to form a number of compounds known as interhalogen of the types AB, AB3, AB5 and AB7 where A is a larger size halogen and B is smaller size halogen.
FLUORINE (F2) : PREPARATION : (i)
Electrolytic method : CaF2 is treated with concentrated H2SO4 to give an aqueous mixture of HF. This is distilled, giving anhydrous liquid HF. A cooled solution of KHF2 in anhydrous HF (KHF2 (1 part) + HF (5 part)) is electrolysed using carbon as anode and steel as cathode in a vessel made of monel metal. electrolysis H2 + F2 CaF2 + H2SO4 CaSO4 + 2HF ; KF + HF K [HF2] Cathode : 2H+ + 2e– H2(g) Anode : 2F– F2 + 2e– The F2 gas thus evolved must be free from HF because it is more corrosive than fluorine. In order to make fluorine free from HF, the gas is passed through NaF which absorbs HF. On Electrolysis :
Anode of carbon should be free from graphite because F2 reacts with graphite easily to form a polymeric substance known as graphite fluoride. There should be no moisture present in the vessel otherwise fluorine will react with water according to the following reaction, 2F2 + 2H2O 4HF + O2
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CHEMISTRY Note : It is not possible to prepare fluorine by electrolysis of aqueous solution of NaF or KF. It is because when aqueous solution of KF is subjected to electrolysis, there will be following two oxidation in competition at anode, H2O 1/2O2 + 2H+ + 2e– SOP = – 1.23 V and F– 1/2F2 + e–
SOP = – 2.87 V
As a matter of rule that substance will be oxidise whose SOP is higher therefore water gets oxidise at anode and not F–. (ii)
Chemical method : From Potassium hexafluoromanganate (IV), K2 [MnF6]. K2 [MnF6] + 2 SbF5 2K [SbF6] + MnF3 +
via 1 1 F2 F2 ; MnF4 MnF3 2 2
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.
PROPERTIES : (i)
Diatomic, Pale green-yellow gas which appears to be almost colourless. It is heavier than air. It condenses to yellow liquid at – 1880C and yellow solid at – 2230C. It has pungent odour and is highly poisonous.
(ii)
Oxidising character : It is the most powerful oxidising agent. F2 + 2NaX 2NaF + X2 ; where (X = Cl, Br, ) (a) It can oxidise all other halide ions into halogen molecules (b) It can oxidise ClO3– into ClO4– and IO3– to IO4– F2 + ClO3– + H2O 2F– + ClO4– + 2H+ (c)
It can oxidise HSO4– into S2O82– 2HSO4– + F2 2F– + S2O82– + 2H+
(iii)
Reaction with NaOH solution : With dilute alkali forms oxygen difluoride and with concentrated alkali liberates O2. 2F2 + 2 NaOH (dilute) OF2 (g) + 2 NaF + H2O 2F2 + 4 NaOH (concentrated) O2 (g) + 4 NaF + 2H2O
(iv)
Reaction with NH3 : 2NH3 + 3F2 N2 + 6 HF It is the distinction from other halogens; other halogens form explosive NX3 with concentrated NH3 i.e. liquor ammonia.
(v)
Reaction with H2S : H2S + F2 SF6 + 2 HF
(vi)
Reaction with SiO2 : It attacks glass at about 1000C. SiO2 (s) + 2F2 (g) SiF4 (g) + O2 (g) The reaction is slow with dry F2 .
(vii)
Reaction with SO3 : 180 º C FSO OOSO F SO3 + F2 2 2 AgF
(viii)
Reaction with O2 : only in presence of O2 + F2 O2F2 silent electric disch arg e
(ix)
Reaction with metals and non–metals : It combines with most of the metals. Almost all non-metals except O2 & N2 ignite spontaneously in presence of F2 . 2 Ag + F2 2 AgF ; 2 Al + 3 F2 2 AlF3 C + 2 F2 CF4 ; Si + 2 F2 SiF4 Fluorocarbons are extremely inert. CF4 can be heated in air without burning. Fluorocarbons are inert to concentrated HNO3 , H2SO4 and to strong oxidising agents such as KMnO4 or O3 and to strong reducing agents such as LiAlH4 or C at 100ºC.
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CHEMISTRY CHLORINE (Cl2) : PREPARATION : (i)
By heating chloride with concentrated H2SO4 in presence of MnO2. 4H+ + MnO2 + 2X– X2 + Mn+2 + 2H2O Bromides and iodides also liberate Br2 and I2 respectively with concentrated H2SO4 and MnO2.
(ii)
(a)
NaCl + HNO3 NaNO3 + HCl × 3 HNO3 + 3HCl NOCl + Cl2 + 2H2O –––––––––––––––––––––––––––––––––––– 3NaCl + 4 HNO3 3 NaNO3 + NOCl (nitrosyl chloride) + Cl2 + 2H2O 2NOCl + O2 2NO2 + Cl2 ; NO2 + H2O HNO3 (to be recycled)
(b)
When Cl2 is used for the chlorination of hydrocarbon the by-product is HCl. The HCl is catalytically oxidised into H2O & Cl2 using copper powder mixed with rare earth chlorides. Cu powder rare earth chloride 4 HCl + O2 2H2O + 2Cl2
(d) (e)
PbO2 + 4 HCl PbCl2, + Cl2 + 2 H2O
(iii)
OCl
+ 2HCl CaCl2 + Cl2 + H2O Cl 2KMnO4 + 16 HCl 2 KCl + 2 MnCl2 + 5 Cl2 + 8 H2O
(c)
Ca
These methods are exclusively used only for chlorine.
Manufacture of chlorine : (a)
Deacon’s process : By oxidation of hydrogen chloride gas by atmospheric oxygen in the presence of CuCl2 (catalyst) at 723 K. 2 4 HCl + O2 2 Cl2 + 2 H2O Electrolytic process : Chlorine is obtained by the electrolysis of brine (concentrated NaCl solution). Chlorine is liberated at anode. It is obtained as a by–product in many chemical industries e.g.; in manufacturing of sodium hydroxide. NaX (aq) Na+ (aq) + X– (aq) Anode : 2X– X2 + 2e–
CuCl
(b)
PROPERTIES : (i)
It is a greenish–yellow gas with pungent and suffocating odour. It is about 2–5 times heavier than air. It can be liquefied into greenish–yellow liquid which boils at 239 K. It is soluble in water.
(ii)
At low temperature it forms a hydrate with water having formula Cl2 . 8H2O which is infact a clathrate compound.
(iii)
Oxidising & bleaching properties : Chlorine dissolves in water giving HCl and HOCl. Hypochlorous acid (HOCl) so formed, gives nascent oxygen which is responsible for oxidising and bleaching properties of chlorine. (a) It oxidises ferrous to ferric, sulphite to sulphate, sulphur dioxide to sulphuric acid and iodine to iodic acid. 2 FeSO4 + H2SO4 + Cl2 Fe2(SO4)3 + 2 HCl Na2SO3 + Cl2 + H2O Na2SO4 + 2 HCl SO2 + 2 H2O + Cl2 H2SO4 + 2 HCl I2 + 6 H2O + 5 Cl2 2 HIO3 + 10 HCl (b) Chlorine oxidises both Br– and l– to Br2 and I2 respectively. X– + Cl2 X2 + 2Cl– (c)
It is a powerful bleaching agent ; bleaching action is due to oxidation. Cl2 + H2O 2 HCl + O
Coloured substance + O Colourless substance It bleaches vegetable or organic matter in the presence of moisture. Bleaching effect of chlorine is permanent.
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CHEMISTRY Note : The bleaching action of SO2 is temporary because it takes place through reduction. SO2 + 2 H2O H2 SO4 + 2 H SO32– + Coloured material SO42– + Reduced colourless material.
O2 of air Reduced Colourless material Coloured material. (iv)
Reaction with NaOH : (a)
2 NaOH ( cold & dilute) + Cl2 NaCl + NaClO + H2O
(b)
6 NaOH (hot & concentrated) + 3 Cl2 5 NaCl + NaClO3 + 3 H2O
These reactions are also given by Br2 and I2. (v)
Reaction with dry slaked lime, Ca(OH)2 : To give bleaching powder. 2 Ca(OH)2 + 2 Cl2 Ca(OCl)2 + CaCl2 + 2 H2O
1. 2. 3. 4. 5.
USES : Cl2 is used for bleaching wood pulp (required for the manufacture of paper and rayon), bleaching cotton and textiles, in the manufacture of dyes, drugs and organic compounds such as CCl4 , CHCl3 , DDT, refrigerants, etc. in the extraction of gold and platinum. in sterilising drinking water and preparation of poisonous gases such as phosgene (COCl 2), tear gas (CCl 3NO 2), mustard gas (ClCH2CH2SCH2CH2Cl). BROMINE (Br2) : PREPARATION :
(i) (ii)
Common method : 2 NaBr + 3H2SO4 (concentrated) + MnO2 Br2 + MnSO4 + 2NaHSO4 + 2H2O
From Sea-water : NaCl is main component but NaBr is also present in some quantity in sea water. Cl2 gas is passed through sea water when vapours of bromine are evolved. 2 Br– (aq) + Cl2 2Cl– (aq.) + Br2 The Br2 is removed by a stream of air since Br2 is quite volatile. The gas is passed through a solution of Na2CO3 when the Br2 is absorbed forming a mixture of NaBr and NaBrO3. The solution is then acidified and distilled to give pure bromine. 3Br2 + 3Na2 CO3 5NaBr + NaBrO3 + 3CO2 5NaBr + NaBrO3 + 3H2SO4 5HBr + HBrO3 + 3Na2SO4 5HBr + HBrO3 3Br2 + 3H2O
PROPERTIES : (i)
Reddish brown liquid, fairly soluble in water. It also forms hydrate like Cl2 (Br2 . 8H2O) Clathrate compound
(ii)
Rest reactions are same as with Cl2
IODINE (I2) : PREPARATION : (i) (ii)
Common method : 2Na + 3H2SO4 (concentrated) + MnO2 2 + MnSO4 + 2NaHSO4 + 2H2O
From Caliche or Crude chile salt petre : The main source of iodine is NaO3 (sodium iodate) which is found in nature with NaNO3 (chile saltpetre). NaO3 is present in small amount. After crystallisation of NaNO3 , the mother liquor left contains NaO3 (soluble). To this solution NaHSO3 is added where upon 2 is precipitated. 2O3– + 6HSO3– 2I– + 6 SO42– + 6H+ 5– + IO3– + 6H+ 3I2 + 3 H2O
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CHEMISTRY (iii)
From sea-weeds : Certain marine plants absorb and concentrate I– selectively in presence of Cl– and Br–. Sea-weeds are dried and burnt in shallow pits, ash left is called kelp. Ash on extraction with hot water dissolves out chlorides, carbonates, sulphates and iodides of sodium and potassium. The solution on concentration separates out all leaving behind iodide in the solution. Solution is mixed with MnO2 and concentrated H2SO4 in iron retorts. Liberated iodine is condensed in series of earthenware known as aludels. 2Na+ MnO2 + 3H2SO4 2NaHSO4 + MnSO4 + 2 + 2H2O
(iv)
CuSO4 + 2K K2SO4 + Cu2 ;
2Cu2 Cu22 + 2
This 2 gets dissolved into K forming 3 , since 3– ions are yellow, therefore solution develops yellow colour.
(i)
It is a dark violet solid, undergoes sublimation, least soluble (among halogens) in water but much more soluble in K(aq.) due to formation of K3 . It is soluble in organic solvents like CHCl3, CCl4 etc. to get violet solutions. Reaction with hypo :
PROPERTIES :
(ii)
S4O62– (tetrathionate ions) + 2– S2 O32– (thiosulphate ions) + 2 This reaction is the basis of iodometric titration, which is carried out for the estimation of iodine using starch indicator.
(iii)
Reaction with NH3 : NH3 (g) + 2 No Reaction
NH3 (aq) + 2 (s) (Ammonia liquor)
(iv)
A slurry is formed which can be dried and on hammering it explodes causing sound (crakers)
N3 . NH3 + 3H an explosive (Nitrogentriiodide ammoniated)
8N3. NH3 5N2 + 9 2 + 6NH4 Reaction with KClO3 or KBrO3 :
2 KClO3 + 2 2 KIO3 + Cl2 ; 2 KBrO3 + 2 2KIO3 + Br2 Example-3
Identify (A) and (B). Br2 + OH– (hot) (A) + (B) (A) + (B) + H+ Br2 (A) gives yellow precipitate with AgNO3 which is completely soluble in concentrated NH3 solution.
Solution
Br2 + OH– (hot) Br– + BrO3– distillation Br– + BrO3– + H+ Br2
HALOGEN ACIDS (HCl, HBr & H) : PREPARATION : (i)
By direct combination of elements : H2 + Cl2 2HCl
(ii)
;
Pt H2 + Br2 2HBr ;
Pt , 450 C H2 + 2 2H
By heating a halide with concentrated acid : (a)
150 º C NaHSO4 + HCl NaCl + H2SO4
550 º C Na2SO4 + HCl NaHSO4 + NaCl This method is called as salt cake method as it involves the formation of NaHSO4 (salt cake). HCl cannot be dried over P2O5 (P4O10) or quick lime since they react with gas chemically. CaO + 2HCl CaCl2 + H2O P4O10 + 3HCl POCl3 + 3HPO3 HCl is, hence dried by passing through concentrated H2SO4 .
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CHEMISTRY (b)
HBr (or H) cannot be prepared by heating bromide (iodide) with concentrated H2SO4 because HBr and H are strong reducing agents and reduce H2SO4 to SO2 and get themselves oxidised to bromine and iodine respectively. KX + H2SO4 KHSO4 + HX H2SO4 + 2HX SO2 + X2 + 2H2O (X = Br or ) Hence, HBr and H are prepared by heating bromides and iodides respectively with concentrated H3PO4. 3KBr(K) + H3PO4 K3PO4 + 3HBr (H)
(iii)
By reaction of P4 (Laboratory Method) : P4 + 6Br2 (62) 4PBr3 (4P3) (product in situ) PBr3 (P3) + 3H2O 3HBr (H) + H3PO3
(iv)
By passing H2S / SO2 into solutions of halogens : H2S + X2 2HX + S SO2 + 2H2O + X2 2HX + H2SO4
PROPERTIES : (i) (ii) (iii) (iv)
(v)
These are colourless, pungent smelling gases with acidic tastes. These are neither combustible nor supporter of combustion. When perfectly dry, they have no action on litmus, but in presence of moisture, they turn blue litmus red, showing acidic nature. Among HX, H is the strongest and HF is the weakest acid. These are quite soluble in water. HCl ionises as below : HCl(g) + H2O () H3O+ (aq) + Cl– (aq) ; Ka = 107 It aqueous solution is called hydrochloric acid. High value of dissociation constant (Ka) indicates that it is a strong acid in water. When three parts of concentrated HCl and one part of concentrated HNO3 are mixed, aqua regia is formed which is used for dissolving noble metals, e.g., gold, platinum. Au + 4 H+ + NO3– + 4 Cl– [AuCl4]– + NO + 2 H2O 3 Pt + 16 H+ + 4 NO3– + 18 Cl– 3 [PtCl6]2– + 4 NO + 8 H2O Reducing property and stability of hydracids : HCl : It is quite stable and hence is oxidised by strong oxidising agents like MnO2, KMnO4, K2Cr2O7, PbO2, Pb3O4 . (i) MnO2 + 4HCl MnCl2 + 2H2O + Cl2 (ii) 2KMnO4 + 16HCl 2KCl + 2MnCl2 + 8H2O + 5Cl2 (iii) K2Cr2O7 + 14HCl 2KCl + 2CrCl3 + 7H2O + 3Cl2 (iv) PbO2 + 4HCl PbCl2 + 2H2O + Cl2 ; (v) Pb3 O4 + 8HCl 3PbCl2 + 4H2O + Cl2 Therefore, HCl is a weak reducing agent. HBr :
It is not very stable and hence more easily oxidised or acts as a strong reducing agents. In addition to above reducing properties of HCl, it also reduces H2SO4 to SO2 which is not done by HCl. H2SO4 + HBr SO2 + Br2 + 2H2O Aqueous HBr on exposure to atmospheric oxygen is oxidised to bromine (yellow) 4HBr + O2 2 Br2 + 2H2O
H :
It is least stable hydrogen halide. It is readily oxidised and thus acts as a powerful reducing agent. In addition to reaction shown by HCl, it shows following reactions also. H2SO4 + 2H SO2 + 2 + H2O ; H2SO4 + 6H S + 32 + 4H2O; H2SO4 + 8H H2S + 42 + 4H2O ; 2HNO3 + 2H 2NO2 + 2 + 2H2O (c) 2HNO2 + 2H 2NO + 2 + 2H2O HO3 + 5H 32 + 2H2O (e) K2S2O8 + 2H K2SO4 + 2 + H2SO4 2FeCl3 + 2H 2FeCl2 + 2 + 2HCl
(a) (b) (d) (f)
Aqueous solution of acid, if exposed to O2 is oxidised to iodine. 4H + O2 22 + 2H2O
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CHEMISTRY (vi)
Detection of cation: HCl :
AgNO3 + HCl AgCl (white) + HNO3 (CH3COO)2 Pb + 2HCl PbCl2(white) + 2CH3COOH Hg(NO3)2 + 2HCl Hg2 Cl2 (white) + 2HNO3
HBr :
AgNO3 + HBr AgBr (pale yellow) + HNO3 (CH3COO)2 Pb + 2HBr PbBr2 (white) + 2CH3COOH
H:
AgNO3 + H Ag (bright yellow) + HNO3 (CH3COO)2 Pb + 2H Pb2 (yellow) + 2CH3COOH HgCl2 + 2H Hg2 (scarlet red) + 2HCl
H reacts with CuSO4 liberating iodine via the formation of cupric iodide (not by HCl or HBr). 2CuSO4 + 4H 2Cu2 + 2H2SO4 ; 2Cu2 Cu2 2 + 2
USES : 1.
2. 3.
HCl is used in preparation of Cl2, chlorides, aqua regia, glucose, (from corn starch), medicines, laboratory as reagents, cleaning metal surfaces before soldering or electroplating. It is also used for extracting glue from bones and purifying bone black. HBr is used as laboratory reagent for preparing bromo derivatives like sodium bromides and potassium bromide. H is used as reducing agent in organic chemistry.
HYDROFLUORIC ACID [H2 F2, HF] : PREPARATION : H2 and F2 combine with each other very violently (even in dark) to form HF. So simple reaction cannot be used for its preparation, special methods are employed for its preparation. (i)
Laboratory Method : Anhydrous HF is obtained by heating dry potassium hydrogen fluoride in a copper retort connected with copper condenser. KHF2 KF + HF
(ii)
Industrial Method : HF is prepared by heating fluorspar (CaF2) with concentrated H2SO4. CaF2 + H2SO4 CaSO4 + 2HF
Aqueous HF being corrosive to glass, is stored in wax lined bottles or vessel made of copper or monel. In glass or silica bottles, it attacks them as follows: Na2SiO3 + 6HF Na2 SiF6 + 3H2O ; CaSiO3 + 6HF CaSiF6 + 3H2O SiO2 + 4HF SiF4 + 2H2O ; SiF4 + 2HF H2 SiF6 This action of HF on silica (silicates) is used for etching glass. The glass surface to be etched is coated with wax, the design, is scratched on glass through wax coating and this is then treated with 40% solution.
PROPERTIES : (i) (ii) (iii)
It is colourless, corrosive liquid with pungent smell with high boiling point due to hydrogen bonding. Dry HF does not attack metals under ordinary conditions (except K), but in presence of water it dissolves metals with liberation of hydrogen gas. It is a weak dibasic acid (due to strong HF bond) and forms two series of salt. NaOH + H2F2 NaHF2 + H2O ; NaHF2 + NaOH 2NaF + H2O
(iv)
HF also react reacts with CCl4 to form freons. CCl4 + HF CFCl3 + HCl
;
CFCl3 + HF CF2Cl2 + HCl.
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CHEMISTRY Example-4
(a)
When HCl reacts with finely powdered iron, it forms ferrous chloride and not ferric chloride. why ?
Solution
(b) (a)
Chlorine water turns blue litmus red but solution becomes colourless after sometime. It forms H2 gas. Fe + 2 HCl FeCl2 + H2 . Liberation of hydrogen prevents the formation of ferric chloride. Blue litmus change into red due to acidic nature (Cl2 + H2O HOCl + HCl) but it is bleaching agent also (oxidising agent),therefore, it decolourises the red litmus.
(b)
OXIDES OF CHLORINE : CHLORINE DIOXIDE (ClO2) : PREPARATION : (i)
traces of NaCl 2 NaClO3 + SO2 + H2SO4 2ClO2 + 2NaHSO4
Chlorates of sodium and potassium can be used
(ii)
90 º C 2KClO3 + 2 H2C2O4 2ClO2 (g) + 2CO2 (g) + K2C2O4 + 2H2O
(iii)
90 º C 2AgClO3 + Cl2 2AgCl (white) + 2ClO2 + O2
(iv)
Cl2O6 + N2O4
(v)
2HClO3 + 2HCl 2ClO2 + Cl2 + 2H2O
ClO2 + [NO2+] [ClO4–]
PROPERTIES : (i)
Yellow gas at room temperature dissolves in water evolving heat and giving a dark green solution.
light ClO2 ClO + [O] It kill bacteria better than Cl2. (ii)
H Cl O (yellow solid) + 2O Reaction with ozone : 2ClO2 + 2O3 2 6 2 dichlorine hexa oxide In the reaction O3 is behaving as an oxidising agent. ClO2 dissolves in NaOH forming chlorite and chlorate.
of
Cl2O6 (s) is a mixed anhydride of HClO3 and HClO4 because on dissolving in water it gives a mixture these two acids. 2Cl2O6 + 4NaOH 2NaClO3 + 2NaClO4 + 2H2O.
(iii)
In solid state, Cl2O6 exists as ClO2+ and ClO4–.
Reaction with alkaline H2O2 : In this reaction H2 O2 acts as a reducing agent. It reduces ClO2 into ClO2– H2O2 2H+ + O2 + 2e– e– + ClO2 ClO2– ] × 2 ––––––––––––––––––––––––––––––––––––––––––––––––– H2 O2 + 2ClO2 2ClO2– + 2H+ + O2 ––––––––––––––––––––––––––––––––––––––––––––––––– or H2O2 + 2ClO2 + 2OH– 2ClO2– + 2H2O + O2
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CHEMISTRY (iv)
Reaction with H : In this reaction H behaves as a reducing agent where it reduces ClO2 into Cl– and itself is oxidised to 2 . [5e– + 4 H+ + ClO2 Cl– + 2H2O] × 2 [2– 2 + 2e–] × 5 ––––––––––––––––––––––––––––––––––––––––––– 2ClO2 + 8H+ + 10 – 5I2 + 2Cl– + 4H2O –––––––––––––––––––––––––––––––––––––––––––
Dichlorine Monoxide (Cl2O) : PREPARATION : 300 º C HgCl2 . HgO + Cl2O 2Cl2 + 2HgO (Freshly prepared precipitate, brown)
PROPERTIES : Cl2O is yellow-brown gas very soluble in water. 3Cl2O + 10NH3 6NH4Cl + 2N2 + 3H2O Cl2O + 2NaOH 2NaOCl + H2O
BLEACHING POWDER (CaOCl2.H2O) Bleaching powder is also called calcium chlorohypochlorite because it is considered as a mixed salt of hydrochloric acid and hypochlorous acid.
Ca(OCl)Cl PREPARATION: Ca(OH)2 + Cl2
40 C Ca(OCl)Cl + H2O 0
Properties (i) (ii)
It is a pale yellow powder. It has a strong smell of chlorine. It is soluble in water but a clear solution is never formed due to the presence of impurities. On long standing, it undergoes auto–oxidation into calcium chlorate and calcium chloride. 6 CaOCl2 Ca(CIO3)2 + 5 CaCl2 2 2 CaOCl2 2 CaCl2 + O2
CoCl
(iii) (iv)
In presence of a little amount of a dilute acid, it loses oxygen. 2 CaOCl2 + H2SO4 CaCl2 + CaSO4 + 2 HClO HClO HCl + O On account of the liberation of nascent oxygen, it shows oxidising and bleaching properties. (a)
Oxidising properties CaOCl2 + H2S CaCl2 + H2O + S CaOCl2 + 2 KI + 2 HCl CaCl2 + 2 KCl + H2O + I2 3 CaOCl2 + 2 NH3 3 CaCl2 + 3 H2O + N2
It oxidises NO2– to NO3– , AsO33– to AsO43– and Fe2+ to Fe3+ (in acidic medium)
(b)
(v)
Bleaching action Coloured matter + [O] colourless product. When bleaching powder reacts with dilute acids or CO2 it liberates chlorine which is known as available chlorine. CaOCl2 + 2 HCl CaCl2 + H2O + Cl2 CaOCl2 + H2SO4 CaSO4 + H2O + Cl2 CaOCl2 + CO2 CaCO3 + Cl2 HNO3 is a strong oxidising acid to be avoided, here.
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CHEMISTRY ESTIMATION OF AVAILABLE CHLORINE : Let the weight of sample of bleaching powder be W g. Add into a beaker containing acetic acid solution and excess K. A yellow brown solution is formed (3–) 2 + – 3– Now few drops of starch solution is added into it. An intensive blue color is observed. Now hypo is used as the titrant. Note the volume where the blue colour disappear.
Reaction involved : CaOCl2 + 2CH3COOH (CH3COO)2 Ca + H2O + Cl2 Cl2 + 2K 2KCl + 2 2 + 2S2O32– S4O62– + 2–
[Mhypo Vhypo ] Calculation : (vi)
%Cl =
W
1 71 2 100
Bleaching powder converts acetone or ethyl alcohol into CHCl3 CaOCl2 + H2O Ca(OH)2 + Cl2 CH3COCH3 + 3 Cl2 CCl3COCH3 + 3 HCl 2 CCl3COCH3 + Ca(OH)2 (CH3COO)2Ca + 2 CHCl3
OXY-ACIDS OF Halogens : HOX SERIES : HYPO-FLOROUS ACID [HOF] : HOF has been prepared by trapping F2 and H2O in unreactive matrix of solid N2 at very low temperature and photolysing the gases. hv F2 + H2O HOF + HF Recent method is by passing, F2 over ice at 0ºC and removing the product into a cold trap.
H2 O + F 2
–40ºC
HOF + HF
HOCl, HOBr and HOl are not very stable and are known only in aqueous solution.
HYPO-CHLOROUS ACID [HClO] : PREPARATION : (i)
The acid is known only in solution, It is obtained by shaking precipitate of HgO with chlorine water. 2HgO + 2Cl2 + H2O Hg2 OCl2 (Oxychloride of mercury) + 2HClO
(ii)
Commercially, it is obtained by passing CO2 through suspension of bleaching powder and then distilling. 2CaOCl2 + H2O + CO2 CaCl2 + CaCO3 + 2HClO Maximum concentration obtained is 25% as in the process of distillation, the acid decomposes into its anhydrides, Cl2O. 2HOCl H2O + Cl2O
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CHEMISTRY PROPERTIES : (i)
It is a weak acid. Its concentrated solution is yellow in colour while dilute solution is colourless. It is unstable and decomposes. 2HClO 2HCl + O2
(ii)
It dissolves magnesium with evolution of hydrogen. Mg + 3HClO Mg(ClO)2 + H2
(iii)
With alkalies, it forms salts called hypochlorites.
(iv)
It acts as a powerful oxidising and bleaching agent. This is due to release of nascent oxygen easily. HClO HCl + O
HXO2 SERIES : CHLOROUS ACID [HClO2] : PREPARATION : It is obtained in aqueous solution when barium chlorite suspension in water is treated with H2SO4. The insoluble barium sulphate is filtered off. Ba(ClO2)2 + H2SO4 BaSO4 + 2HClO2
PROPERTIES: (i)
The freshly prepared solution is colourless but it soon decomposes to ClO2 which makes the solution yellow. 5 HClO2 4 ClO2 + HCl + 2H2O
(ii)
(iii)
Salts of HClO2 are called chlorite and prepared by one of the following methods. 2ClO2 + 2NaOH NaClO2 + NaClO3 + H2O 2ClO2 + Na2O2 2NaClO2 + O2 Chlorites are used as bleaches. They are stable in alkaline solution even when boiled, but in acid solution they disproportionate, particularly when heated. 5HClO2 4ClO2 + HCl + 2H2O and 4HClO2 2ClO2 + HClO3 + HCl + H2O The acid liberates iodine from K 4K + HClO2 + 2H2O 4KOH + HCl + 22
HXO3 SERIES CHLORIC ACID [HClO3] : PREPARATION : This acid is only known in solution. The acid is prepared by the action of the dilute H2SO4 on barium chlorate. Ba (ClO3)2 + H2SO4 BaSO4
+ 2HClO3.
After reaction, BaSO4 is removed by filtration, and the filtrate is evaporated in vacuum till 40 percent solution is obtained. However, further concentration by evaporation leads to decomposition. 3HClO3 HClO4 + Cl2 + 2O2 + H2O HBrO3 can be prepared by similar method using Ba(BrO3)2 .
PROPERTIES : Concentrated acid is colourless and pungent smelling liquid. It decomposes in light. However, it is stable in dark. It acts as a strong oxidising and bleaching agent in light. Organic substances like paper, cotton, wool, etc., catch fire in contact with the acid.
HClO3 oxidises SO2 to SO3 : HClO3 + 3SO2 HCl + 3SO3 HClO3 when evaporates to dryness decomposes giving ClO2. 4HClO3 4ClO2(g) + 2H2O(g) + O2(g) HBrO3 is not very stable, but is known in solution, and as salts. HlO3 is formed by oxidation of I2 with concentrated HNO3 or O3. 8H+ + 10NO3– + I2 2IO3– + 10NO2 + 4H2O lO3– oxidises I– to I2 : lO3– + 5I– + 6H+ 3l2 + 3H2O Iodic acid is reasonably stable and exists as a white solid.
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CHEMISTRY
150 º C 2KCl + 3O 2KClO3 MnO 2 2
low temperatur e 4KClO3 3KClO4 + KCl (in absence of catalyst)
2Zn(ClO3)2 2ZnO + 2Cl2 + 5O2 Chlorates are used in fire work.
PERCHLORIC ACID [HCIO4] : PREPARATION: (i)
(ii)
It is the most stable oxy-acid of chlorine. Anhydrous HClO4 is obtained by doing distillation of KCIO4 (potassium perchlorate), with 96-97.5% H2SO4 under low pressure at 90-160°C. KCIO4 + H2SO4 KHSO4 + HCIO4 An aqueous solution of the acid is obtained by reacting barium perchlorate with calculated quantity of dilute H2SO4. The insoluble barium sulphate is removed by filtration. Ba(CIO4)2 + H2SO4 BaSO4 + 2HClO4
(iii)
Electrolysis ClO4– + 2H+ + 2e– NaClO3 + H2O
(iv)
NH4ClO4 + HNO3 HClO4 + NH4NO3
(v) (vi)
4CIO3– 3ClO4– + CI– HClO4 . 2H2O + 2H2S2O7 HClO4 (obtained as anhydrous HClO4) + 2H2SO4
(i) (ii) (iii) (iv) (v)
PROPERTIES: Anhydrous HClO4 is a colourless liquid which turns dark on keeping. It fumes in moist air. It is one of the strongest acid and ionises as follows : HClO4 H+ + ClO4– It dissolves most of the metals. Zn + 2HClO4 Zn(ClO4)2 + H2 Hot concentrated acid (73%) behaves as a remarkable oxidising agent : 4HClO4 2Cl2 + 7O2 + 2H2O 2HClO4 + P2O5 2HPO3 + Cl2O7 Mg (ClO4)2 is used in dry batteries and is also an effective desiccant called anhydrone. KClO4 is used in fire works and flares.
Example-5
CaOCl2 in aqueous solution changes to Cl2 . What is the type of this change ?
Solution
CaOCl2 + H2O Ca(OH)2 + Cl2
Example-6
Identify A, B and C in the following
Redox reaction.
Solution
Example-7 Solution
Which of the following is a mixed anhydride (A) P4 O10 (B) SO3
(C) Cl2O6
(D) SO2
(C) Cl2O6 when dissolves in water produces HClO3 and HClO4, therefore, it is mixed anhydride.
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CHEMISTRY Example-8
Give appropriate reasons for each of the following : (a) Addition of Cl2 to KI solution gives it a brown colour but excess of Cl2 turns it colourless. (b) Perchloric acid is a stronger acid than sulphuric acid. (c) HI can not be prepared by heating NaI with concentrated H2SO4.
Solution
(a)
Cl2 being a stronger oxidising agent than I2, first oxidises KI to I2 which imparts brown colour to the solution. But when Cl2 is passed in excess, the I2 so formed gets further oxidised to HIO3 (colourless) 2KI (aq) + Cl2(g) 2KCl (aq) + I2(s) ; 5Cl2 + I2 + 6H2O 10HCl + 2HIO3
(b)
(c)
Oxidation state of Cl in HClO4 is + 7 and that of S in H2SO4 is + 6. (Cl is more electronegative than S). As a result, ClO3 part of HClO4 can break the O–H bond more easily to liberate a proton than SO2 part in H2SO4 . Thus HClO4 is a stronger acid then H2SO4. HI is a stronger reducing agent. It, therefore, reduces H2SO4 to SO2 and itself gets oxidised to I2. NaI + H2SO4 NaHSO4 + HI] × 2 H2SO4 + 2HI SO2 + I2 + 2H2O –––––––––––––––––––––––––––––––––––––––––––– 2NaI + 3H2SO4 2NaHSO4 + I2 + 2H2O + SO2
Some important order (a) (b)
(c)
Acid strength (i) H > HBr > HCl > HF Oxidising powder (i) F2 > Cl2 > Br2 > 2
(ii) HOCl > HOBr > HO
(iii) HClO4 > HClO3 >HClO2 > HClO
(ii) BrO4– > IO4– > ClO4– (According to electrode potential) Order of disproportionations 3 XO– 2X– + XO3– (hypohalite ion) ; O– > BrO– > ClO–
Pseudo halogens and pseudo halides : Some inorganic compounds consisting of two or more atoms of which at least one is N have been found to behave like halogens & they are known as pseudo halogen solids, e.g. (CN)2 cyanogen, (SCN)2 thiocyanogen, (SeCN)2 selenocyanogen, (SCSN3)2 azidocarbondisulphide. Similarly few ions are known that have properties similar to those of the halide ions. They are therefore called pseudohalide ions, e.g. (CN–) cyanide ion, (SCN–) thiocyanate ion, (SeCN) – selenocyanate ion, (OCN) – cyanate ion, (NCN)2– cyanamide ion, (N3)– azide ion etc.
INTERHALOGEN COMPOUNDS : We know that halogen atoms have different electronegativity. Due to this difference in electronegativity the halogen atoms combine with each other and give rise to the formation of binary covalent compounds, which are called interhalogen compounds. These are of four types. AB AB3 AB5 AB7 ClF ClF3 ClF5 lF7 BrF BrF3 BrF5 ICl ICl3 IF5 IF IF3
PREPARATION : (i)
By the direct combination of halogens : 473 K Cl2 + F2 (equal volumes) 2ClF ;
573 K Cl2 + 3F2 (excess) 2ClF3 ;
I2 + Cl2 2ICl ; (equimolar)
Diluted with water :
Br2 (g) + 3F2 2BrF3
F2 is diluted with N2 :
78 º C 2IF3 I2 + 3F2
F2 is taken in freon :
Br2 + 5F2 (excess) 2BrF5
IF7 can not be prepared by direct combination of I2 & F2 .
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CHEMISTRY (ii)
From lower interhalogens : ClF + F2 ClF3 ;
ClF5 ClF3 + F2 (excess) 350 º C
BrF5 ; BrF3 + F2 (excess) 200 º C
270 º C IF7 IF5 + F2
This method is generally applied for the preparation of halogen fluorides.
(iii)
Other methods : 6HCl + KIO3 + 2KI 2KCl + 3H2O + 3ICl 250 350 º C 200 º C KF + ClF5 Cl2 + ClF3 3CIF ; KCl + 3F2
3l2 + 5AgF 5Agl + lF5
; 8F2 + PbI2 PbF2 + 2lF7
PROPERTIES : (i) (ii) (iii) (iv) (v)
(vi)
These compounds may be gases, liquids or solids. Gases : ClF, BrF, ClF3 , IF7 ; Liquids : BrF3, BrF5 ; Solids : ICl, IBr, IF3, ICl3. Interhalogens containing fluorine are colourless but inter halogens consisting of heavier halogens are coloured. The intensity of colour increases with increase in the molecular weight of the compounds. All interhalogens are covalent molecules and are diamagnetic in nature since all the valence electrons present as bonding or non-bonding electrons are paired. The boiling points increases with the increase in the electronegativity difference between A and B atoms. Thermal stability of AB type interhalogen compounds decreases with the decrease in electronegativity difference between A and B atoms. IF > BrF > ClF > ICl > IBr > BrCl. More polar is the A – B bond more is the stability of interhalogen. Interhalogen compounds are more reactive than the parent halogens but less reactive than F2. ICl + 2Na NaI + NaCl The order of reactivity of some interhalogens is as follows : ClF3 > BrF3 > IF7 > BrF5 > BrF.
(vii)
Hydrolysis : All these undergo hydrolysis giving halide ion derived from the smaller halogen and a hypohalite (when AB), halite (when AB3), halate (when AB5), and perhalate (when AB7) anion derived from the larger halogen. AB + H2O HB + HOA BrCl + H2O HCl + HOBr ; ICl + H2O HCl + HIO ICl3 + 2H2O 3HCl + HIO2 ; IF5 + 3H2O 5HF + HIO3 IF7 + 6H2O 7HF + H5IO6 ; BrF5 + 3H2O 5HF + HBrO3
Oxidation state of A atom does not change during hydrolysis.
(viii)
Reaction with non-metallic and metallic oxides : 4BrF3 + 3SiO2 3SiF4 + 2Br2 + 3O2 ; 4BrF3 + 2WO3 2WF6 + 2Br2 + 3O2
(x)
Reaction with alkali metal halides : IBr + NaBr NaIBr2 ;
ICl3 + KCl KICl4
USES : These compounds can be used as non aqueous solvents. Interhalogen compounds are very useful fluorinating agents. ClF3 and BrF3 are used for the production of UF6 in the enrichment of 235U . U(s) + 3 ClF3 (l) UF6 (g) + 3 ClF (g) Example-9 Solution
The correct order of pseudohalide, polyhalide and interhalogen are : (A) Br2– , OCN– , F5 (B) F5 , Br2– , OCN– (C) OCN– F5 , Br2–
(D) OCN– Br2– , F5
OCN– is pseudohalide, Br2– polyhalide and IF5 is interhalogen. So option [D] is correct.
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CHEMISTRY
GROUP 18 ELEMENTS : (THE ZERO GROUP FAMILY) Group 18 consists of six elements: helium, neon, argon, krypton , xenon and radon . All these are gases and chemically unreactive. They form very few compounds . Because of this they are termed noble gases. Occurrence All the noble gases except radon occur in the atmosphere. Their atmospheric abundance in dry air is ~ 1% by volume of which argon is the major constituent. Helium and sometimes neon are found in minerals of radioactive origin e.g., pitchblende, monazite, cleveite. The main commercial source of helium is natural gas. Xenon and radon are the rarest elements of the group. Radon is obtained as a decay product of 226Ra.
226 88 Ra
222 86 Rn
+ 24He
Most abundant element in air is Ar. Order of abundance in the air is Ar > Ne > Kr > He > Xe. Electronic Configuration All noble gases have general electronic configuration ns2np6 except helium which has 1s2 . Many of the properties of noble gases including their inactive nature are ascribed to their closed shell structures. Ionisation Enthalpy Due to stable electronic configuration these gases exhibit very high ionisation enthalpy . However, it decreases down the group with increases in atomic size. Atomic Radii Atomic radii increase down the group with increase in atomic number. Electron Gain Enthalpy Since noble gases have stable electronic configurations, they have no tendency to accept the electron and therefore, have larger positive values of electron gain enthalpy. Physical properties All the noble gases are mono-atomic. They are colourless, and tasteless. They are sparingly soluble in water. They have very low melting and boiling points because the only type of interatomic interaction in these elements is weak dispersion forces,. Helium has the lowest boiling point (4.2K) of any known substance. It has a unusual property of diffusing through most commonly used laboratory materials such as rubber, glass or plastics.
Table : 2 Atomic and physical properties Element
He
Ne
Ar
Kr
Xe
Atomic Number
2
10
18
36
54
Atomic Mass
4
20.18 2
Electronic configuration Atomic Radius (pm) –1
Ionization enthalpy / (kJ mol ) –3
Density (at STP)/g cm
39.10
2
6
83.80
2
6
10
131.30 2
6
10
2
6
1s
[He] 2s 2p
[Ne] 3s 3p
[Ar] 3d 4s 4p
[Kr] 4d 5s 5p
120
160
190
200
220
2372
2080 –4
1.8 × 10
1520 –4
9.0 × 10
1351 –3
1.8 × 10
1170 –3
–3
3.7 × 10
5.9 × 10
Melting point / K
–
24.6
83.8
115.9
161.3
Boiling point / K
4.2
27.1
87.2
119.7
165.0
Chemical Properties : In general, noble gases are least reactive. Their inertness to chemical reactivity is attributed to the following reasons: (i) (ii)
The noble gases except helium (1s2) have completely filled ns2 np6 electronic configuration in their valence shell. They have high ionisation enthalpy and more positive electron gain enthalpy. The reactivity of noble gases has been investigated occasionally ever since their discovery, but all attempt to force them to react to form the compounds were unsuccessful for quite a few years. In March 1962, Neil Bartlett, then at the University of British Columbia, observed the reaction of a noble gas. First , he prepared a red compound which is formulated as O2+ PtF6–. He , then realised that the first ionisation enthalpy of molecular oxygen (1175 kJ mol –1) was almost identical with that xenon (1170 kJ mol –1). He made efforts to prepare same type of compound with Xe+ PtF6 – by mixing Pt F6 and Xenon. After this discovery, a number of xenon compounds mainly with most electronegative elements like fluorine and oxygen, have been synthesised.
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CHEMISTRY The compounds of krypton are fewer. Only the difluoride (KrF2) has been studied in detail. Compounds of radon have not been isolated but only identified (e.g., RnF2) by radiotracer technique. No true compounds of Ar, Ne or He are yet known .
If Helium is compressed and liquified it forms He() liquid at 4.2 K. This liquid is a normal liquid like any other liquid. But if it is further cooled then He() is obtained at 2.2 K, which is known as super fluid, because it is a liquid with properties of gases. It climbs through the walls of the container & comes out. It has very high thermal conductivity & very low viscosity. Clatherate compounds : Inert gas molecules get trapped in the cages formed by the crystal structure of water. During the formation of ice Xe atoms will be trapped in the cavities (or cages) formed by the water molecules in the crystal structure of ice. Compounds thus obtained are called clatherate compounds. There are no chemical bonds. They do not possess an exact chemical formula but approx it is 6 water molecules : 1 inert gas molecule. The cavity size is just smaller than the atom of the noble gas. Such compounds are also formed by the other organic liquids like dihydroxybenzene (for example quinol). The smaller noble gases He and Ne do not form clathrate compounds because the gas atoms are small enough to escape from the cavities. Clathrate provides a convenient means of storing radioactive isotopes of Kr and Xe produced in nuclear reactors.
Example-10 Solution
Example-11 Solution
Name the noble gas which (A) is most abundant in atmosphere. (B) has least boiling point. (A) Argon. (B) Helium ; Exists as mono-atomic molecules and are held together by weak van der Waal’s forces. These van der Waal’s forces increase with the increase in atomic size of the atom, and therefore, the boiling points increases from He to Rn. Hence He has least boiling point. What is the utility of the clatherate compounds? It can be used to separate mixture of inert gases containing say He and Xe. Process is much cheaper than say distillation as a means of separation.
COMPOUNDS OF XENON : XENON DIFLUORIDE (XeF2): PREPARATION: (i)
873 K ,1bar Xe + F2 XeF2 Ni Tube or monel metal (alloy of Ni)
Volume ratio should be 2 : 1 otherwise other higher fluorides tend to form.
(ii)
Xe + O2 F2 118 C XeF2 + O2
(iii)
Hg ( arc ) Xe + F2 XeF2
(iv)
Recently discovered method : K+ [AgF4]– [potassium tetrafluoroargentate ()] is first prepared and this is reacted with BF3 . BF3 K+ [AgF4]– AgF3 (red solid) + KBF4 2 AgF3 + Xe 2 AgF2 (Brown solid) + XeF2
PROPERTIES : (i) (ii) (iii) (iv)
Colorless crystalline solid and sublimes at 298 K. Dissolves in water to give a solution with a pungent odour. Much soluble in HF liquid. This is stored in a vessel made up of monel metal which is a alloy of nickel. Reaction with H2 : It reacts with hydrogen gas at 4000C XeF2 + H2 Xe + 2HF
(v)
Hydrolysis : (a) 2XeF2 + 2H2O 2Xe + 4HF + O2 (slow) The above is neither a cationic hydrolysis nor an anionic hydrolysis as seen in ionic equilibrium. It is a covalent compound and hydrolysis is like that of PCl5 . (b)
Hydrolysis is more rapid with alkali. XeF2 + 2 NaOH Xe + ½O2 + 2NaF + H2O (fast) The reaction (a) is slower probably due to dissolution of XeF2 in HF.
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CHEMISTRY (vi)
Oxidising properties : Higher the value of SRP better is the oxidising property of the species. The standard reduction potential for XeF2 is measured to be + 2.64 V. Therefore, it acts as a strong oxidising agent. 2e– + 2H+ + XeF2 Xe + 2HF; SRP = + 2.64 V This oxidises halides (except F–) to their respective halogens. XeF2 + 2 HCl Xe + 2 HF + Cl2 It oxidises 2Br – Br2 + 2e– & 2– 2 + 2e– Similarly it can oxidise BrO3– (bromate) which are themselves good oxidising agents to BrO4– (perbromate ions) and Ce3+ to Ce4+ ion.
(vii)
Oxidising as well as fluorinating properties : It can act as strong oxidising agent as well as fluorinating agent. C6 H5 + XeF2 C6 H5 F2 + Xe ; CH3 + XeF2 CH3 F2 + Xe
CH3 F2 exists as CH3+F2– ,F2– is analogous to 3– F
:
– I
hybridisation = sp3d :
:
F
F3– can not be formed as it has no d-orbitals to attain sp3d hybridisation.
(viii)
Reactions of XeF2 + HF (anhydrous) : HF PtF6 + 3Xe ; Pt + 3XeF2
S8 + 24 XeF2 8SF6 + 24 Xe
2CrF2 + XeF2 2CrF3 + Xe ; 2MoO3 + 6XeF2 2MoF6 + 6Xe + 3O2 Mo (CO)6 + 3 XeF2 Mo F6 + 3 Xe + 6CO 2
HF 2 + XeF2
+ Xe
8 NH3 + 3 XeF2 N2 + 6 NH4 F + 3 Xe 2NO2 + XeF2 2 NO2 F + Xe (nitronium fluoride) (ix)
Formation of addition compounds : XeF2 reacts with fluoride ion acceptors to form cationic species and fluoride ion donors to form fluoroanions. XeF2 + PF5 [XeF]+ [PF6]– F5 (lewis acid) + XeF2 [XeF]+ [F6]– ; 2SbF5 (lewis acid) + XeF2 [XeF]+ [SbF6]– Similar behaviour is shown by PF5 and AsF5 Structure :
Shape linear and geometry trigonal bipyramidal.
hybridisation = sp3d
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CHEMISTRY XENON TETRAFLUORIDE (XeF4) : PREPARATION : Xe
+
2F2
1
:
5
873 K, 7 bar , Nitube
XeF4
PROPERTIES : (i) (ii) (iii)
It is a colorless crystalline solid and sublimes at 298 K. It undergoes sublimation, soluble in CF3 COOH. It undergoes hydrolysis violently hence no moisture must be present during it’s preparation. Reaction with H2O : 6 XeF4 + 12 H2O 4 Xe + 2XeO3 + 24 HF + 3O2 XeO3 is white solid and explosive compound (dry), soluble in water (well behaved in water) XeO3 reacts with NaOH forming sodium xenate XeO3 + NaOH Na+ [HXeO4]– (sodium xenate) It disproportionates into perxenate ion in basic medium. 4– 2 [HXeO4]– + 2OH– [XeO6] + Xe + O2 + 2H2 O Xenic acid (H2XeO4) is a very weak acid.
H+ water XeO3 + O2 (slow decomposition)
(i) [XeO]64–
(ii) [XeO6]4– +Mn +2
–
MnO4 + XeO3
[XeO6]4– is obtainable as Na4 XeO6 . 8H2O (sodium perxenate) (iv)
Oxidising properties of XeF4 : XeF4 + 2 H2 Xe + 4 HF ; XeF4 + 2 Hg Xe + 2HgF2
(v)
Addition reactions : XeF4 reacts with fluoride ion acceptors to form cationic species and fluoride ion donors to form fluoroanions. XeF4
(vi)
+
SbF5 [XeF3]+ [SbF6]–
Fluorinating agent : XeF4 + Pt PtF4 + Xe ;
XeF4 + 4NO Xe + 4NOF (nitrosyl fluoride)
XeF4 + 4 NO2 Xe + 4 NO2 F (nitronium fluoride) ;
XeF4 + 2C6 H6 C6H5F + 2HF + Xe
Structure : Shape square planar & geometry octahedral
hybridisation = sp3d2
XENON HEXAFLUORIDE (XeF6) : PREPARATION: (i)
573 K , 60 – 70 bar Xe 3F2 XeF6 Nitube 1 : 20
(ii)
XeF4 + O2 F2 XeF6 + O2
PROPERTIES: (i) (ii) (iii)
Colourless crystalline solid and sublimes at 298 K. It gives yellow liquid on melting where as other form white liquids on melting (a point of difference) HF is a good solvent for all three fluorides.
(iv)
Hydrolysis : (a) Complete hydrolysis : XeF6 + 3H2 O XeO3 (white solid) + 2HF (b) Partial hydrolysis : XeF6 + H2O XeOF4 (colorless solid) + 2HF
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CHEMISTRY (v)
Reaction with silica (SiO2) : 2XeF6 + SiO2 2XeOF4 + SiF4
(vi)
Thermal decomposition (effect of heat) : 2XeF6 XeF2 + XeF4 + 3 F2
(vii)
XeF2 & XeF4 do not undergo decomposition
Formation of addition compounds : XeF6 + SbF5 [XeF5]+ [SbF6]– ;
XeF6 + BF3 [XeF5]+ [BF4]–
XeF6 + 3H2 6HF + Xe
(viii)
Reaction With H2 :
(ix)
Reaction of XeF6 with XeO3 : XeO3 + 2 XeF6 3 XeOF4
(x)
F– donating/ F– accepting properties : XeF6 reacts with fluoride ion acceptors to form cationic species and fluoride ion donors to form fluoroanions. XeF6 + MF M+ [XeF7]– (M = Na, K, Rb or Cs) donation { XeF6 + PtF5 (XeF5+) (PtF6–)
Example-12
Does the hydrolysis of XeF4 at – 80ºC lead to a redox reaction ?
Solution
– 80 º C XeF4 + H2O XeOF2 + 2HF.. The oxidation states of all the elements in the products remain the same as it was in the reacting state. hence, it is a not redox reaction .
XENON–OXYGEN COMPOUNDS : Hydrolysis of XeF4 and XeF6 with water gives XeO3. 6 XeF4 + 12 H2O 4 Xe + 2 XeO3 + 24 HF + 3 O2 XeF6 + 3 H2O XeO3 + 6 HF Partial hydrolysis of XeF6 gives oxyfluorides, XeOF4 and XeO2F2 . XeF6 + H2O XeOF4 + 2 HF XeF6 + 2 H2O XeO2F2 + 4 HF XeO3 is a colourless explosive solid and has a pyramidal molecular structure. XeOF4 is a colourless volatile liquid and has a square pyramidal molecular structure.
USES : Helium is a non–inflammable and light gas. Hence, it is used in filling balloons for meteorological observations. It is also used in gas–cooled nuclear reactors. Liquid helium (b.p.4.2 K) finds use as cryogenic agent for carrying out various experiments at low temperatures. It is used to produce and sustain powerful superconducting magnets which form an essential part of modern NMR spectrometers and Magnetic Resonance Imaging (MRI) systems for clinical diagnosis. It is used as a diluent for oxygen in modern diving apparatus because of its very low solubility in blood. Neon is used in discharge tubes and fluorescent bulbs for advertisement display purposes. Neon bulbs are used in botanical gardens and in green houses. Argon is used mainly to provide an inert atmosphere in high temperature metallurgical process (arc welding of metals or alloys) and for filling electric bulbs. It is also used in the laboratory for handing substances that are air–sensitive. Xenon and Krypton are used in light bulbs designed for special purposes.
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CHEMISTRY
MISCELLANEOUS SOLVED EXAMPLES 1.
Identify [A] [B] and [C] and gives the complete chemical reactions involved.
Ans.
[A] = Br– ; [B] = BrO3– ; [C] = concentrated H2SO4
Sol.
3Br2 + 6OH– 5Br– + BrO3– + 3H2O 5Br– + BrO–3 + 6H+ 3Br2 + 3H2O
2.
Comment on the following. (a) Electrolysis of ICN in pyridine solution. (b) Iodine dissolves in oleum. (c) Electrical conductivity of molten iodine.
Sol.
(a) Iodine is liberated at cathode indicating the ionisation of ICN into I+ and CN–. (b) Bright blue solution is formed which has been shown to have I2+ and I3+. 2 I2+ 6H2SO4. SO3 2I2+ + 2HS3O10– + 5H2SO4 + SO2 3 I2 + 6H2SO4. SO3 2I3+ + 2HS3O10– + 5H2SO4 + SO2. (c) It is due to the presence of (I3+ and I3–) species produced by self ionisation of iodine 3I2
3.
I3+ + I3–
Why anhydrous HF liquid is not electrolysed alone to get F2?
Sol.
Anhydrous HF is only slightly ionized and is, therefore a poor conductor of electricity Thus a mixture of KF and HF is electrolysed to increase the conductivity.
4.
Which of the following product(s) is/are obtained in the following reaction KBrO3 + F2 + KOH product(s) (A) KBrO4
(B) KF
(C) HOF
(D) Br2
Sol.
KBrO3 + F2 + 2KOH KBrO4 + 2KF + H2O. Ans (A,B)
5.
Match the following . Column - I
Column - II
(A) CIO2 CI2O3
(p) Boiling with NaOH solution.
(B) [AI (OH)4] AI(OH)3
qOn passing ozone.
CP4 PH3 + H2PO2–
(r) Reaction with hydrogen.
(D) XeF2 Xe
(s) On passing CO2 gas.
–
Sol.
(A q) ; (B s) ; (C p) ; (D p,r) (A)
2CIO2 + 2O3 CI2O6 + 2O2 oxidation by ozone.
(B)
CO2 + H2O H2CO3
CO32– + 2H+ 2AI3+ + 3CO32- + 3H2O 2AI(OH)3 + 3CO2
As acidic property of AI(OH)3 is very weak. (C)
boil / warm P4 + 3NaOH + 3H2O PH3 + 3NaH2PO2.
AIkaline hydrolysis. (D)
2XeF2 + 4OH- 2Xe + 4F– + 2H2O + O2 Alkaline hydrolysis. XeF2+ H2 2Xe + 2HF Reduction by hydrogen.
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CHEMISTRY 6.
Na2S2O3 may react with the compounds given in column (I). Na2S2O3 exhibits the properties of the type given in the column (II) Column - I
Column - II
(reactant)
(type of property shown)
(A) Chlorine (CI2)
(p) Complexing reagent
(B) Silver bromide
(q) Disproportionation
(C) Hydrochloric acid
(r) Only as reductant
(D) Iodine (I2)
(s) An-antichlor
Ans.
(A - r, s) ; (B - p) ; (C - q) ; (D - r)
Sol.
(A) Na2 S 2 O3 + 4CI2 + 5H2O 2NaH S O 4 + 8HCI
6
2
It destroys any excess of chlorine on fabric in bleaching industry. Thus it acts as antichlor. (B) Ag+ + 2S2O32– [Ag (S2O3)2]3– (soluble complex) 2
4
O
(C) Na2 S 2 O 3 +2HCI 2NaCI + S O2 + S + H2O 2
(D) 2Na2 S 2 O 3 + I2 Na2S4O6 + 2NaI
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