+2 UNIT 8 PAGE-4

Some of the important chemical reactions of halogens are discussed below :
1. Formation of Hydrides
All the halogens react with hydrogen to form hydrides , HX known as halogen hydrides.
H2 + X2  2 HX
(a) Physical state : Hydrogen fluoride is a low boiling liquid (b.p. 292 K) while HCl, HBr and HI are gases. The anomalous property of HF is due to the presence of hydrogen bonding in the molecules. Due to hydrogen bonding in HF the molecules exist as associated molecule (HF)n.
…… H  F ….. H  F ….. H  F ….. H  F ….. or (HF)n.
(b) Thermal stability : The thermal stability of the hydrides decreases from HF to HI. HF is most stable whereas HI is least stable. For example, HF and HCl are stable upto 1500 K while HBr dissociate to extent of 10% and HI is dissociated to the extent of 20% at 700 K.
The decrease in stability of the hydrides is due to decrease in bond strength which decreases when we go down the group.
(c) Reducing character : The decreasing thermal thermal stability of the hydrogen halides from HF to HI indicates that the reducing character increases down the group as :
HF < HCl < HBr < HI Thus, HF is not a reducing agent at all. HCl is a weak reducing agent. HBr is a stronger reducing agent while HI is the strongest reducing agent among all the hydrides. (d) Acidic strength : In gaseous state, hydrogen halides are covalent. But in aqueous solutions, they ionise and behave as acids. The acid strenth of these acids decreases in the order HI > HBr > HCl > HF
Thus HF is the weakest acid and HI is the strongest acid
among these halogen acids.
Explanation : The above order of acidic strength is reverse of that expected on the basis of electronegativity. Fluorine is the most electronegative halogen, therefore , the electronegativity difference will be maximum in HF and and should decrease gradually as we move towards iodine through chlorine and bromine. Thus, HF should be more ionic in nature and consequently it should be strongest acid. Although many factors contribute towards the relative acidic strengths, the major factor is the bond dissociation energy. The bond dissociation energy decreases from HF to HI so that HF has maximum bond dissociation energy and HI has the lowest value.
Hydrogen halide HF HCl HBr HI
Bond dissociation energy (kJ/mol) 566 431 366 299
Since H I bond is weakest, it can be easily dissociated into H+ and I  ions while HF can be dissociated with maximum difficulty. Thus HI is the strongest acid while HF is the weakest acid among the hydrogen halides.
06. Suggest a method for laboratory preparation of DCl. Write a balanced equation for the reaction.
Fluorine forms two oxides F2O and F2O2 which are called oxygen fluorides. In this case fluorine is more electronegative than oxygen. On the otherhand the oxides of Cl, Br and I are called oxides. They form oxides from + 1 to + 7 oxidation states. They are listed in the TABLE.
Oxides of halogens in different oxidation states
Oxidation state Fluorine Chlorine Bromine Iodine
 1 OF2 - - -
+ 1 Cl2O Br2O -
+ 2 F2O2 - - -
+4 - ClO2 BrO2 I2O4
+5 - - - I 2O5
+6 - Cl2O6 BrO3 -
+7 - Cl2O7 - I2O7
All these oxides are powerful oxidising agents and decompose explosively when subjected to mechanical shock or heat. ClO2 and Cl2O are used as bleaching agents for paper, pulp, textiles and water treatment. The structures of some common oxides of chlorine are given.

Structures of oxides of Chlorine
In line with its high electronegativity and small size, fluorine forms only one oxoacid HOF known as fluoric(I)acid or hypofluorous acid. The other halogens form several oxoacids. Most of them cannot be isolated pure. They are stable only in aqueous solutions or in the form of their salts. The oxoacids of halogens are shown below :
The oxoacids of halogens are given in TABLE.

hypohalous acid) HOF
(hypofluorous acid) HOCl
(hypochlorous acid ) HOBr
(Hypobromous acid) HOI
(Hypoiodous acid)
(Halous acid) - HOClO
(chlorous acid) - -
(halic acid) - HOClO2
(chloric acid) HOBrO2
(bromic acid) HOIO2
(iodic acid)
(perhalic acid) - HOClO3
(perchloric acid) HOBrO3
(perbromic acid) HOIO3
(periodic acid)

The following important generalizations can be made regarding the acidic character of oxoacids of halogens :
(a) The acidic strength of oxoacids having the same oxidation number of the halogen atom decreases with increase in atomic number i.e., with decreasing electronegativity of the atom. For example , HClO is the strongest while HIO is the weakest of all the acids in which oxidation state of halogen atom is +1. This is evident from their dissociation constant (Ka) values :
Ka 2.9 x 108 5.0 x 109  1011
Explanation : This can be explained on the basis of the electronegativity of the halogen atom. The electronegativity of the halogen atom attached to oxygen decreases in the order Cl > Br > I. As a result the tendency to pull the electrons from the hydrogen decreases . Consequently , the release of H+ ion from HClO will be easier than HBrO. Similarly, the release of H+ ion from HBrO will be easier than from HIO. Thus, the acidic strength decreases in the order :
(b) With increase in oxidation number of a particular halogen atom, the acidic character of corresponding oxoacid increases. For example, the acid strength of oxoacids of chlorine increases in the order :
HClO < HClO2 < HClO3 < HClO4 Explanation : This can be explained on the basis of Lowry Bronsted concept. According to this concept, a strong acid has weak conjugate base and a weak acid has a strong conjugate base. Let us consider the stabilities of the conjugate bases, ClO, ClO2, ClO3 and ClO4 formed from these acids, HClO, HClO2, HClO3 and HClO4 respectively. These anions are stabilised to greater extent , it has lesser attraction for proton and therefore, will behave as weaker base (lesser tendency for the reaction to go in backward direction). Consequently, the corresponding acid will be strongest because weak conjugate base has strong acid and strong conjgate base has weak acid and vice versa. Now the charge stabilization will be minimum in ClOand maximum in ClO4. The charge stabilization increases in the order : ClO < ClO2 < ClO3 < ClO4 . This means that ClO will have minimum stability and therefore, will have maximum attraction for H+. In other words ClO will be strongest base and so its conjugate acid HClO will be the weakest acid. Similarly, in this series ClO4 is the weakest base (maximum stabilized) and its conjugate acid HClO4 is the strongest acid. Thus, the acidic strength increases in the order : HClO < HClO2 < HClO3 < HClO4 Chloric (I)acid HOCl Chloric(I) acid is formed by the disproportionation of chlorine in water : Cl2 (g) + H2O(ℓ)  HOCl(aq) + HCl(aq) Pure HOCl can be prepared by the addition of HgO or Ag2O to chlorine water. 2 Cl2 + 2 HgO + H2O  HgO. HgCl2 + 2 HOCl Chloric(I) acid is stable only in solution. The oxoanion ion is ClO , hypochlorite or [monoxochlorate(I). NaOCl is used as household bleach and is commercially produced by the electrolysis of cold dilute NaCl solution under conditions where chlorine liberated, reacts with hydroxide ion to produce the hypochlorite ion : 2 Cl2 + 2 OH (aq)  Cl  (aq) + OCl  (aq) + H2O(ℓ) Powdered calcium halate (I), Ca(OCl)2 2 H2O is used for disinfecting swimming pool water. For general bleaching and sanitation purposes, bleaching powder , produced by passing chlorine over slaked lime and having the composition Ca(OCl)2 CaCl2 Ca(OH)2 .2 H2O is used. Halate (I) salts react with acids to liberate chlorine , for example : Ca(OCl)2 (s) + 4 HCl (aq)  CaCl2 (aq) + 2 H2O(ℓ) + 2 Cl2 (g) Chloric(III)acid HOClO Chloric(III)acid HOClO is prepared by the reaction of Ba(ClO2)2 (formed by thre treatment of Ba(OH)2 , H2O2 and ClO2) with dilute sulphuric acid. Ba(OH)2 + H2O2 + 2 ClO2  Ba(ClO2)2 + 2 H2O(ℓ)+ O2 Ba(ClO2)2 + H2SO4  BaSO4 + 2 HOClO Among the salts of chloric (III)acid, NaClO2 is commercially important as a bleaching agent for textiles and for removal of (NO)x pollutants from industrial waste gases. Chloric(V) acid HOClO2 HOClO2 is stable only in solution. When chlorine is passed through hot concentrated alkali solutions, [(trioxochlorate(V)] salts are formed : 6 Cl2 (g) + 6 Ba(OH)2(aq)  Ba(ClO3)2 (aq) + 5 BaCl2(aq) + 6 H2O(ℓ) The free acid is obtained by treating the barium salt with dilute H2SO4. Ba(ClO3)2 + H2SO4  BaSO4 + 2 HClO3 Perchloric acid , [Tetraoxochloric(VII)] acid HClO4. HClO4 is obtained by treating anhydrous NaClO4 or Ba(ClO4)2 with concentrated HCl : NaClO4 + HCl  NaCl (s)+ HClO4(aq) The free acid HClO4 explodes above 365 K but it can be distilled at low temperatures under reduced pressure. Commercially NaClO4 is made by the electrolytic oxidation of NaClO3 . An important chlorate(VII) salt is ammonium perchlorate, NH¬4ClO4 which is used as an oxidiser in solid rocket propellants in missile and space programmes. It is produced frm NaClO4 by the reaction : NaClO4 (aq) + NH4Cl(aq)  NaCl (aq)+ NH¬4ClO4(s) Ammonium perchlorate is less soluble than sodium perchlorate and is precipitated as a solid. Sodium chlorate NaClO3 : Sodium chlorate is produced on a large scale by the electrolysis of brine in a cell where efficient mixing of chlorine produced at the anode with OH produced at the cathode takes place The structures of the oxoacids of chlorine are shown in Fig. INTERHALOGEN COMPOUNDS The compounds containing two or more halogen atoms are called interhalogen compounds. Each halogen combines with every other halogen to form interhalogen compounds. For example, ClF, ICl3, BrF5 etc. These are of two types : (a) Neutral interhalogen compounds : These are neutral molecules containing two or more halogen atoms. For example, ICl, BrF5, IF5, IF7 etc. (b) Interhalogen anions and cations : These are negatively charged interhalogen anions or polyhalide ions such as ICl2, ICl4, I3 and positively charged interhalogen cations such as ICl2+, ICl4+ etc. The different types of interhalogens of the type AX(diatomic), AX3(tetra atomic) , AX5 (hexa atomic) and AX7 (octa atomic) are given below : Types of interhalogen AX AX3 AX5 AX7 ClF chlorine fluoride ClF3 chlorine trifluoride ClF5 chlorine pentafluoride - Br F bromine fluoride BrF3 bromine trifluoride BrF5 bromine pentafluoride - IF* iodine fluoride IF3* iodine trifluoride IF5 iodine pentafluoride IF7 iodine heptafluoride BrCl bromine chloride - - - ICl iodine chloride ICl3 iodine trichloride - - * are unstable Interhalogen ions Type Examples AX2 ICl2, IBr2, BrCl2, ClF2, ICl2, BrF2+, ClF2+ AX4 ICl4, BrF4 , IF4+, ClF4+, BrF4+ AX6 IF6, BrF6, IF6+, BrF6+, ClF6+ All interhalogen compounds are covalent compounds. These are generally more reactive than the component halogens. It is due to the weakness of covalent bond between disimilar halogen atoms. These are prepared by direct combination or by the reactions of halogen with other interhalogen compounds. For example, Structures of some common interhalogen compounds (a) Neutral interhalogens : The structure of some common interhalogen compounds such as AX3, AX5 and AX7 are given below. These are easily understood on the basis of VSEPR theory. For example, (i) AX3 involves sp3d hybridisation of the central halogen atom and the molecule has trigonal bipyramidal geometry with two positions occupied by lone pairs. Its structure is T-shaped . For example, ClF3 has T-shaped structure. (ii) AX5 involves sp3d2 hybridisation of the central halogen atom and the molecule has octahedral geometry with one position occupied by a lone pair. Its structure is termed as square pyramidal. (iv) IF7 involves sp3d3 hybridisation of iodine atom and the molecule has pentagonal bipyramidal structure. For example IF7 has pentagonal bipyramidal structure. (b) Interhalogen ions Triatomic (AX2)  ions (i) AX2 ions involved sp3d hybridisation having trigonal bipyramidal geometry with three positions occupied by lone pairs. The structure is linear. For example ICl2 is linear. (ii) AX4 ions involve sp3d2 hybridisation having octahedral geometry with two positions occupied by lone pairs.. The two lone pairs are present on positions above and below the square plane and the structures are regarded as square planar. For example, ICl4 has square planar geometry. Evidence for cationic iodine All halogens are non-metallic due to high electronegativities and ionisation enthalpies. However, the last element iodine exhibits some properties showing the presence of positive ions. The positive iodine is found in unipositive (I+) and ( I3 +) tripositive states. (a) Iodine monochloride (ICl) conducts electricity in the molten state. On electrolysis, iodine is liberated at the cathode and while both iodine and chlorine are liberated at the anode. The liberation of iodine at the cathode indicates the presence of cationic iodine. The ionisation of ICl may be represented as : (b) ICl3 conducts electricity and on electrolysis, iodine and chlorine are liberated at both the electrodes. Therefore , the ionisation may be represented as : The ICl2+ and ICl4 contain I3+ ions . (c) A large number of compounds containing iodine as I3+ have been isolated which are ionic. For example, iodine triacetate I(CH3COO)3 , iodine phosphate IPO4 , iodine triperchlorate I(ClO4)3 . On electrolysis, these compounds in aqueous solution liberates iodine at the cathode as : Pseudohalides There are several uni-negative groups which show certain characteristics of the halide ions. These are called pseudohalides or pseudohalide ions. The important ones are given in TABLE. Pseudo halide ions and Pseudohalogens Pseudohalide ions Pseudohalogens Pseudohalide ion Formula Pseudohalogens Formula Cyanide CN Cyanogen (CN)2 Cyanate OCN - - Thiocyanate SCN Thiocyanogen (SCN) 2 Selenocyanate SeCN Selenocyanogen (SeCN) 2 Azide N3 - - Azidothiocarbonate SCSN3 Azidocarbon disulphide (SCSN3) 2 Isocyanate ONC  Pseudohalogens As each halide ion has a corresponding halogen, (chloride ion, for example, has chlorine as the corresponding halogen) attempts have been made to isolate corresponding pseudohalogens. So, far , only four pseudohalogens , have been isolated. Anomalous behaviour of Fluorine Like other elements of the second period, fluorine also differs from the rest of the family members in many characteristics. This behaviour is attributed to its : i) Very small size. ii) High electronegativity. iii) Absence of vacant d-orbitals in the valence shell. Due to the above reasons, fluorine exhibits many properties which are different from other halogens. For example, 1. Oxidation state : Fluorine shows oxidation state of 1 only, while other halogens show oxidation states such as +1, +3, +5 and +7 also. 2. Bond dissociation energy : The bond dissociation energy of fluorine molecule is less than that of other halogen molecules. This is due to the fact that the repulsion between non-bonding electrons in small fluorine molecule are very strong. As a result FF bond is weak and can be easily broken. 3. Anomalous behaviour of hydrogen fluoride : Due to high electronegativity of fluorine, the bonding pair in HF molecule is largely attracted towards fluorine and therefore, it forms hydrogen bonds. As a result of hydrogen bonding in HF, it exits as associated molecule (HF)n and is a liquid. On the other hand, HCl, HBr and HI are gases at room temperature. Due to association in HF molecules, its boiling point is high in comparison to other halogen acids. In aqueous solution, HF is much weaker acid than other hydrogen halides. 4. Ionic character of fluorides : Only form ionic fluorides. For example AlF3, SnF4 etc. are ionic in nature while the corresponding chlorides are covalent. 5. Solubility : Fluorides have abnormal solubilities than other halides. For example, AgF is soluble in water whereas AgCl is insoluble ; CaF2 is insoluble in water while CaCl2 is soluble. 6. Formation of polyhalide ions : Fluorine does not form polyhalide ions such as F3 while other halogens form polyhalide ions such as I3, Br3, I5 etc. Note The poly halide ion I3 which is formed when iodine dissolves in an aqueous solution of potassium iodide. I2(s) + I (aq)  I3 (aq) Problem 07. Deduce the molecular shape of BrF3 on the basis of VSEPR theory. 08. With what neutral molecule is ClO isoelectronic ? Uses of Halogens and their compounds Fluorine is used mainly for the manufacture of UF6 for nuclear power generation and SF6 for dielectrics. The important organic chemicals derived from HF are the chlorofluorocarbons and polytetrafluoro-ethylene (Teflon) . Chlorofluorocarbons known as freons are used as refrigerants (these are however, being phased out in favour of chlorofluoro hydrocarbons) and in aerosols and artificial blood. Minor uses of HF are in the glass industry as etching agent and in the manufacture of fluoride salts. Prominent among the fluorides is NaF used for the fluorination of water ; one part per million level of fluoride in drinking water prevents tooth decay . SnF2 is used in fluoride tooth pastes. Chlorine is used for bleaching of paper, pulp and textiles and as disinfectant for sterilising drinking water and in the production of organic compounds (e.g., polyvinyl chloride, chlorinated hydrocarbons, pharmaceuticals , herbicides, pesticides) and inorganic compounds (e.g., HCl, PCl3 and NaOCl). The main outlet for bromine is ethylene bromide which is used as a gasoline additive. Bromine is also used to make AgBr for photography. Iodide is necessary for the normal functioning of the thyroid gland. Insufficient iodine in the diet leads to goitre (enlargement of the thyroid gland). Hence , sodium or potassium iodate / iodide is added to table salt and this type of salt is known as ‘iodised’ salt. GROUP-18 ELEMENTS NOBLE GASES The Group-18 consists of elements Helium(He),Neon (Ne) Argon(Ar), Krypton(Kr) , Xenon (Xe) and radon(Rn). It constitutes zero Group of the periodic table. These gases at ordinary temperature do not have chemical reactivity and therefore, these were called inert gases. However, now-a-days, a number of compounds of these elements are prepared. Consequently , these gases are called noble gases instead of inert gases which signifies that these gases have some reactivity. Electronic Configuration : Except helium, the atoms of all noble gases have eight electrons in the valence shell. The general electronic configuration of noble gases(except helium) may be expressed as ns2np6. On the other hand, helium has 1s2 electronic configuration. Electronic configuration of Noble gases Element Symb-ol Atomic number Electronic configration He He 2 1s2 Neon Ne 10 1s2 2s22p6 Argon Ar 18 1s2 2s22p63s2 3 p6 Krypton Kr 36 1s2 2s22p63s23p63d10 4s2 4p6 Xenon Xe 54 1s22s22p63s2 3p63d10 4s24p6 4d10 5s2 5p6 Radon Rn 86 1s2 2s22p63s2 3p63d10 4s24p6 4d10 4f14 5s2 5d10 6s2 6p6 These configurations being stable, the noble gases neither have any tendency to gain nor lose electrons, and therefore, they do not enter into chemical combinations. It is therefore, reasonable to assume that inert nature of the noble gases is due to their stable electronic configurations. General characteristics of Noble gases 1. Existence : All the noble gases are monoatomic, colourless and odourless gases. The monoatomic state of these gases is due to stable electronic configuration of their atoms. As a result, they are not capable of combining even amongst themselves. 2. Atomic Radii : In the case of noble gases, the atomic radii correspond to van der Waal's radii. As we go down the Group, the van der Waal's radius increases due to the addition of new electronic shells. 3. Ionisation energies : The ionisation energies of noble gases are very high. This is attributed to the stable completely filled configurations of noble gases. However, the ionisation energies decrease with increase in atomic number from He to Rn due to increasing atomic size. 4. Electron gain enthalpies : Due to the stable ns2np6 electronic configurations, noble gas atoms have no tendency to accept additional electron. Therefore, their electron affinities are nearly zero. 5. Melting and boiling points : The melting and boiling points of noble gases are very low in comparison to those of other substances of comparable atomic and molecular masses. This indicates that only weak van der Waal's forces are present between atoms of the noble gase in the liquid or the solid state. These van der Waal's forces increase with the increase in atomic size of the atom and therefore, the boiling points and melting points increase from Helium to Radon. 6. Ease of liquefaction : The noble gases are not easily liquefied. This is due to the fact that there are only weak van der Waal's forces which hold atoms together. This is also clear from the low values heats of vapourisation of these gases. Due to increase in atomic size and therefore, increase in van der Waal's forces, the ease of liquefaction increases down the Group from Helium to Radon. 7. Solubility in water : The noble gases are only slightly soluble in water. The solubility , in general decreases down the Group from Helium to Radon. Chemical Properties In general, noble gases are not very reactive. Their inertness to chemical reactivity is attributed to the following reasons: i) The noble gases have completely filled ns2np6 electronic configurations in their valence shells. ii) The noble gases have very high ionisation energies. iii) The electron affinities of noble gases are almost zero. Therefore, they have neither tendency to gain nor to lose any electron and do not enter into chemical combinations. Before 1962, it was believed that noble gases do not combine at all and no compounds of noble gases were formed. In 1962, Bartlett noticed that platinum hexafluoride , PtF6 is a powerful oxidising agent which combines with molecular oxygen to form ionic compound, dioxygenyl hexafluoroplatinate(V), O2+[PtF6]. O2(g) + PtF6(g)  O2+[PtF6] This indicates that PtF6 has oxidised O2 to O2+. Now , oxygen and xenon have some similarities: i) The first ionisation energy of xenon gas(1170 kJ/mol) is fairly close to that of oxygen(1166 kJ/mol). ii) The molecular diameter of oxygen and atomic radius of xenon are similar(400 pm). On this assumption, Bartlett reacted xenon and platinum hexafluoride in gas phase and an orange yellow solid of composition XePtF6 was obtained. Xe(g) + PtF6 (g)  Xe+[PtF6](s) Orange yellow Once the reactivity of noble gases was established many more attempts were made to synthesis their compounds. Now many compounds of xenon and krypton are known with fluorine or oxygen. The compounds of krypton are fewer, only the difluoride , KrF2 has been studied in detail. The compounds of radon have not been isolated so far but have been identified by radiotracer techniques. Lastly, the compounds of He, Ne or Ar are not known. COMPOUNDS OF XENON Oxides, Fluorides and Oxyfluorides : Xenon forms three fluorides i.e., XeF2, XeF4 and XeF6. These can be obtained by the direct interaction between xenon and fluorine under appropriate experimental conditions. The oxide and oxyfluorides of xenon are obtained from the fluorides. Xenon trioxide can be obtained by the hydrolysis of XeF4 or XeF6. The oxyfluorides of xenon namely xenon oxydifluoride(XeOF2) and xenon oxytetrafluoride(XeOF4) are also obtained by the partial hydrolysis of XeF4 and XeF6 respectively. The fluorides of Xenon, i.e., XeF2 , XeF4 and XeF6 react with fluoride ion acceptors like PF5, SbF5, AsF5 to form adducts in which the flourides (XeF2, XeF4 and XeF6) change to cationic species whereas the fluoride ion acceptors form fluoroanions. Structure of Xenon Compounds The structure of xenon compounds can be explained on the basis of VSEPR theory as well as the concept of hybridisation. In the formation of fluorides of Xe, depending upon the the number of XeF covalent bonds to be formed e.g., two in XeF2, four in XeF4 and six in XeF6, the electrons in the valency shell of xenon get unpaired. These are promoted to vacant 5d-orbitals. Then depending upon the total pairs of electrons( bond as well as lone pairs) the molecule assumes a particular geometry. The above contention is illustrated in the following TABLE. Molecule Total electron pairs (bp+ ℓ p) Hybridisation Shape XeF2 5 (=2 + 3) sp3d Linear XeF4 6 (= 4 + 2) sp3d2 Square planar XeF6 7 (=6 + 1) sp3d3 Distorted octahedral* *On the basis of hybridisation, the molecule should have pentagonal bipyramidal structure. However, the structure has not yet been confirmed. Thus, it is preferably believed to have distorted octahedral structure. The structures of oxyfluorides and oxides can best be explained by the concept of hybridisation. The types of hybridisation in these molecules and their shapes are listed below : Structures of some Xenon compounds Molecule Hybridisation Geometry Actual shape XeOF2 Sp3d Trigonal bipyramidal T-shape XeOF4 Sp3d2 Square pyramidal - XeO3 Sp3 Tetrahedral pyramidal Problem 09. Give the formula of the noble gas species that is isostructural with : (a) ICl4 (bb) IBr2 (c) BrO3 Uses of Noble gases Argon is used mainly to provide an inert atmosphere in high temperature metallurgical processes (arc welding of metals or alloys) and for filling electric bulbs. It is also used in the laboratory for handling substances that are air sensitive. Helium is non-inflammable and light gas. Hence, it is used in filling balloons for materological observations. It is also used in gas-cooled nuclear reactors. Liquid helium (b.p 4.2 K) finds use as a cryogenic agent for carrying out various experiments at low temperatures. It is used to produce and substain powerful superconducting magnets which form essential part of modern NMR spectrometers and Magnetic Resonance Imaging (MRI) systems for clinical diagnosis. Neon is used in discharge tubes and fluorescent bulbs for adverisement display purposes. There is no significant uses of xenon and krypton. They are used in light bulbs designed for special purposes. Questions 1. Highlight the difference in the structures of boron chloride and anhydrous aluminium chloride. 2. Boron trifluoride is a strong Lewis acid. Explain. 3. What are boranes. Give chemical equations for the industrial preparation of diboranes. 4. Why do boron halides form addition compounds with amines. 5. Why are boron halides and diboranes referred to as electron-deficient compounds ? 6. BF3 is a weaker Lewis acid than BCl3. Give reason. 7. Aluminium is a good reducing agent. Explain. 8. Caustic alkalies like NaOH are not stored in aluminium vessels. Explain why ? 9. Why does aluminium not react with cold water under ordinary conditions ? 10. ‘Aluminium cannot be extracted by reduction of its ores by carbon.’ Explain briefly. 11. ‘Aluminium is more reactive than iron and yet there is less corrosion of aluminium, when both are exposed to air’. Explain. 12. Aluminium dissolves both in acidic and basic solutions. Explain. 13. Aluminium is used for electrical cables, though it is relatively less conducting than copper. Explain. 14. Boron forms electron deficient compounds. Give reason. 15. Give in brief, general trends in physical properties of group-14 elements. 16. Discuss the similarities and dissimilarities between carbon and silicon. 17. Give the points of similarities and dissimilarities between boron and silicon. 18. Discuss diagonal relationship between boron and silicon. 19. Contrast the structure and properties of carbon dioxide and silicon dioxide. 20. Carbon cannot form complexes ; whereas other elements of group-14 can form. Account for it. 21. Write a note on halides of Group-14 elements. 22. Account for the following observations : a) In trimethyl amine, the nitrogen atom has a pyramidal geometry, whereas in trisilylamineN(SiH3)3 , it has a planar geometry. b) Elemental silicon does not form a graphite-like structure as carbon. 74. Why is (SiH3) 3 N a weaker base than (CH3) 3N ? 75. Carbon tetrachloride is not hydrolysed with water ; but silicon tetrachloride is easily hydrolysed. Give reason. 76. Explain why silicon shows a higher covalence than carbon. 77. CO2 is a gas, but SiO2 is a high melting solid. Give reason. 78. Explain why boiling points of hydrides of group-16 elements changes in the order : H2 O > H2 Te > H2 Se > H2 S.
79. Carbon and silicon are almost tetravalent, but germanium , tin and lead show divalency. Explain why.
80. Why do tin(II)chloride a solid, but tin(IV) chloride a liquid ?
81. Carbon exhibits catenation ; whereas silicon does not. Why ?
82. What are silicones ? How are they manufactured ?
83. Discuss in brief, general physical properties of Group-15 elements.
84. Nitrogen does not pentahalides, but phosphorus does. Give reason.
85. Molecular nitrogen is not particularly reactive. Account for it.
86. Writethe shapes of hydrides of Group-15. Arrange them in the order of (i) decreasing basic strength (ii) increasing bond angle (iii) decreasing reducing character.
87. H3PO3 is dibasic. Explain.
88. In what ways nitrogen differs in its chemical behaviour from that of heavier congers ?
89. List the various oxidation states of nitrogen and give an example a compound or an ion for each case.
90. Draw the structures important oxides of nitrogen.
91. Illustrate how nitrogen compounds provide good examples of multiple bonding and reasonance.
92. Write a note on hydrides of Group-15 elements.
93. Explain the following observations :
(a) Ammonia has a higher boiling point than phosphine.
(b) Ammonia is a good complexing agent.
94. Explain why ammonia is :
(a) highly soluble in water.
(b) a Lewis base in liquid form.
(c) used as a non-aqueous solvent.
95. Ammonia is distinctly basic, but phosphine is weakly basic. Explain.
96. Why is ammonia soluble in water, but other hydrides of group-15 elements are insoluble in water ?
97. Ammonia is distinctly basic, but phosphine is weakly basic ; while the other hydrides of group-15 show no basic character at all. Explain why.
98. Discuss the stability of group-15 elements.
99. Draw the structures of :
(i) Phosphorus trihalide and
(ii) Phosphorus pentachloride.
100. Why does not elemental phosphorus exist as P2 under ordinary conditions ?
101. Give a method of preparation of orthophosphoric acid. What happens when it is heated ? Give its two uses.
102. Draw the structures of P2O3 and P2O5.
103. Nitrogen exists as a diatomic molecule(N2) ; whereas phosphorus exists as a traatomic molecule(P4). Give reasons.
104. Draw the structures of oxyacids of phosphorus.
105. Discuss in brief, the general trends in physical properties of Group-16 elements.
106. Describe the trends in the following properties of group-16 elements :
(i) Catenation
(ii) Stability of hydrides.
(iii) Metallic character
(iv) Capability to form halides.
(v) Acidic character of oxides.
107. What are chalcogens ? Why are they so called ?
108. Why oxygen does not exhibit +4 and +6 oxidation states like sulphur ?
109. The tendency for  2 oxidation state diminishes from sulphur downwards in the periodic table. Explain.
110. What is meant by catenation ? Give example with reference to Group-16 elements.
111. What has a higher boling point than hydrogen sulphide. Explain.
112. Water is a liquid, while hydrogen sulphide is a gas at ordinary temperature. Explain.
113. Oxygen exists as a gas; while sulphur exists as a solid at ordinary temperature. Explain.
114. Why is OF6 compound not known ?
115. Taking suitable examples, indicate the different oxidation states in which oxygen may exist in its compounds.
116. Give reason why oxygen molecules have the formula O2 ; whilst sulphur is S8.
117. State any four properties which make oxygen different from rest of elements of the same group. Give reasons for this anomalous behaviour.
118. Oxygen differs from the rest of the members of its family. Justify.
119. Amongst the hydrides of elements of oxygen family, water shows anomalous properties why ?
120. Compare the structures of : (i) SO2 and SeO2 (ii) SO3 and SeO3


Atoms and Molecules

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