Page 1 : p-BLOCK ELEMENTS, , p-BLOCK ELEMENTS, GROUP 15 ELEMENTS : THE NITROGEN FAMILY, Group 15 includes nitrogen phosphorus, arsenic, antimony and bismuth. As we go down the group, there is a shift from nonmetallic to metallic through metalloidic character. Nitrogen and phosphorus are non-metal, arsenic and antimony metalloid, and bismuth is a typical metal., Electronic Configuration :, The valence shell electronic configuration of these element is ns2 np3 the s orbital in these element is completely filled and, p orbitals are half- filled, making their electronic configuration extra stable., Atomic and Ionic Radii :, Covalent and ionic (in a particular state) radii increase in size down the group. There is a considerable increase in covalent, radius from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence, of completely filled d and / or f orbitals in heavier members., Ionisation Enthalpy:, Ionisation enthalpy decreases down the group due to gradual increase in atomic size. Because of the extra stable half- filled, p-orbital electronic configuration and smaller size, the ionisation enthaply of the group 15 element is much greater, than of group 14 elements in the corresponding periods. The order of successive ionisation enthalpies, as expected is, iH1 < iH2 < iH3, Electronegativity :, The electronegativity value, in general, decreases down the group with increasing atomic size. However, amongst the heavier, elements, the difference is not that much pronounced., Physical Properties:, All the elements of this group are polyatomic. Dinitrogen is a diatomic gas while all others are solids. Metallic character, increases down the group. Nitrogen and phosphours are non – metals , arsenic and antimony metalloids and bismuth is a, metal. This is due to decrease in ionisation enthalpy and increase in atomic size. The boiling points , in general , increase, from top to bottom in the group but the melting point increases upto arsenic and then decreases upto bismuth. Except, nitrogen , all the elements show allotropy., ATOMIC & PHYSICAL PROPERTIES, Element, , N, , P, , As, , Sb, , Bi, , Atomic Number, , 7, , 15, , 33, , 51, , 83, , Atomic Mass, , 14.01, , 30.97, , 74.92, , 121.76, , 208.98, , Electronic configuration, , [He] 2s2 2p3, , [Ne] 3s2 3p3, , [Ar] 3d10 4s2 4p3, , [Kr] 4d10 5s2 5p3, , [Xe] 4f14 5d10 6s2 6p3, , Covalent Radius / pm, Ionic Radius / pm, a = M3–, b = M+3, , 70, , 110, , 120, , 140, , 150, , 171a, , 212a, , 222a, , 76b, , 103b, , Ionization enthalpy, / (kJ mol–1), Electronegativity, , , , 1402, , 1012, , 947, , 834, , 703, , , , 2856, , 1903, , 1798, , 1595, , 1610, , , , 4577, , 2910, , 2736, , 2443, , 2466, , 3.0, , 2.1, , 2.0, , 1.9, , 1.9, , Chemical Properties :, Oxidation States and trends in a chemical reactivity :, The common oxidation states of these elements are – 3 , + 3 and + 5. The tendency to exhibit – 3 oxidation state decreases, down the group , bismuth hardly forms any compound in –3 oxidation state. The stability of + 5 oxidation state decreases, down the group. The only well characterised Bi (V) compound is BiF5 .The stability of + 5 oxidation state decreases and that, of +3 state increases (due to inert pair effect) down the group. Nitrogen exhibits +1 , + 2 , + 4 oxidation states also when it, reacts with oxygen. Phosphours also shows + 1 and + 4 oxidation states in some oxoacids., In the case of nitrogen , all oxidation states from +1 to +4 tend to disproportionate in acid solution. For example ,, 3 HNO2  HNO3 + H2O + 2 NO, Similarly , in case of phosphorus nearly all intermediate oxidation states disproportionate into +5 and –3 both in alkali and, acid. However +3 oxidation state in case of arsenic , antimony and bismuth become increasingly stable with respect to, disproportionation.

Page 2 : p-BLOCK ELEMENTS, Nitrogen is restricted to a maximum covalency of 4 since only four (one s and three p) orbitals are available for bonding. The, heavier elements have vacant d orbitals in the outermost shell which can be used for bonding (covalency) and hence , expand, their covalence as in PF6– ., Anomalous properties of nitrogen :, Nitrogen differs from the rest of the members of this group due to its smaller size , high electronegativity , high ionisation, enthalpy and non – availability of d orbitals. Nitrogen has unique ability to form p – p multiple bonds with itself and with, other elements having small size and high electronegativity (e.g., C ,O). Heavier elements of this group do not form p – p, bonds as their atomic orbitals are so large and diffuse that they cannot have effective overlapping. Thus, nitrogen exists as a, diatomic molecule with a triple bond (one s and two p) between the two atoms. Consequently , its bond enthalpy (941.1 kJ, mol–1) is very high. On the contrary , phosphorus , arsenic and antimony form metallic bonds in elemental state. However ,, the single N – N bond is weaker than the single P – P bond because of high interelectronic repulsion of the non – bonding, electrons , owing to the small bond length. As a result the catenation tendency is weaker in nitrogen. Another factor which, affects the chemistry of nitrogen is the absence of d orbitals in its valence shell. Besides restricting its covalency to four ,, nitrogen cannot form d – p bonds as the heavier elements can e.g., R3P = O or R3P = CH2 (R = alkyl group). Phosphours, and arsenic can form d – p bond also with transition metals when their compounds like P(C2H5)3 and As(C6H5)3 act as, ligands., (i), Reactivity towards hydrogen :, All the elements of Group 15 form hydrides of the type EH3 where E = N , P, As, Sb or Bi. Some of the properties of these, hydrides are shown in Table. The hydrides show regular gradation in their properties. The stability of hydrides decreases, from NH3 to BiH3 which can be observed from their bond dissociation enthalpy. Consequently , the reducing character of the, hydrides increases. Ammonia is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the, hydrides. Basicity also decreases in the order, NH3 > PH3 > AsH3 > SbH3  BiH3 ., Properties of Hydrides of Group 15 Elements, , Property, , NH3, , PH3, , AsH3, , SbH3, , BiH3, , Melting point / K, , 195.2, , 139.5, , 156.7, , 185, , –, , Boiling point / K, , 238.5, , 185.5, , 210.6, , 254.6, , 290, , (E – H) Distance / pm, , 101.7, , 141.9, , 151.9, , 170.7, , –, , 107.8, , 93.6, , 91.8, , 91.3, , –, , – 46.1, , 13.4, , 66.4, , 145.1, , 278, , 389, , 322, , 297, , 255, , –, , 0, , HEH angle ( ), fH– / kJ mol–1, dissH (E – H) / kJ mol, –, , (ii), , (iii), , (iv), , –1, , Reactivity towards oxygen :, All these elements form two types of oxides : E2O3 and E2O5 . The oxide in the higher oxidation state of the element is more, acidic than that of lower oxidation state. Their acidic character decreases down the group. The oxides of the type E2O3 of, nitrogen and phosphours are purely acidic , that of arsenic and antimony amphoteric and those of bismuth is predominantly, basic., Reactivity towards halogens :, These elements react to form two series of halides : EX3 and EX5 . Nitrogen does not form pentahalide due to non –, availability of the d – orbitals in its valence shell. Pentahalides are more covalent than trihalides. All the trihalides of these, elements except those of nitrogen are stable. In case of nitrogen , only NF3 is known to be stable. Trihalides except BiF3 are, predominantly covalent in nature., Reactivity towards metals :, These elements react with metals to form their binary compounds exhibiting –3 oxidation state , such as , Ca 3N2 (calcium, nitride) Ca3P2 (calcium phosphide) , Na3As2 (sodium arsenide) , Zn3Sb2 (zinc antimonide) and Mg3Bi2 (magnesium, bismuthide).

Page 10 : p-BLOCK ELEMENTS, It is a polymeric substance forming linear chains like this., , (iii), , Black phosphorus:, It has two forms -black phosphorus and -black phosphorous, (a) -black phosphorous, ,    P(-black), P(red) tube, 803 K, insulated, , -black phosphorous structure is not definite and is non conductor of electricity., (b) -black phosphorous, 473 K, , ,  P(-black), P(white) High, pressure, , -black phosphorous is an electrical conductor resembling graphite in this respect and also in its flakiness and, luster. It is insoluble in CS2. It has a layered structure like graphite., P, , P, P, , P, , P, , P, P, , P, P, , P, P, , P, , P, P, , P, , (iv), , P, P, , P, P, , Brown phosphorus:, Above 1600oC, P4 molecules begin to dissociate into P2 molecules. Rapid cooling of this vapour gives brown phosphorus, which probably contains P2 molecules., , CHEMICAL PROPERTIES OF PHOSPHORUS :, , , (i), , (ii), , (iii), , Reactivity of the various allotropic forms of phosphorus towards other substances decreases in the order:, Brown > white > red > black, the last one being almost inert., Apart from their reactivity difference, all the forms are chemically similar., Action of air :, White phosphorus burns in air to form phosphorus trioxide and pentoxide., P4 + 5O2  2P2O5 ; P4 + 3O2  2P2O3, Red and other forms of phosphorus also burn in air or oxygen but on heating., Action of non-metals:, When heated with non-metals phosphorus forms compounds PX3, PX5, P2S3 and P2S5., 2P + 5X2  2PX5 (where X = Cl, Br, and I.), 2P + 3X2  2PX3,, Action with metals:, Alkali metals when heated with white phosphorus in vacuum produce alkali metal phosphide, which react with water to form, phosphine gas., , (iv), , , , M3P M3P + 3H2O , 3MOH + PH3 { where M = Na, K etc.}, 3M + P , Action of NaOH:, When white phosphorus is heated with NaOH solution, phosphine gas is evolved., , (v), , P4 + 3NaOH + 3H2O  3NaH2PO2 + PH3, Action of conc. HNO3 :, When heated with conc. HNO3, phosphorus is oxidized to H3PO4., P + 5NHO3  H3PO4 + 5NO2 + H2O

Page 14 : p-BLOCK ELEMENTS, (iii), , Neutralization with alkalies or bases:, NaOH, NaOH, NaOH, ,  NaH PO (pri. phosphate), ,  NaHPO (sec. phosphate) , ,  Na PO (tert. phosphate), H3PO4, H2O, H2O, H2O, 2, 4, 4, 3, 4, , , , , NaPO3 + H2O, NaH2PO4 , , USES:, It is used as a laboratory reagent and in manufacture of medicines., , GROUP SIXTEEN ELEMENTS : THE OXYGEN FAMILY, Oxygen, sulphur, selenium, tellurium and polonium constitute group 16 of the periodic table. This is sometimes known as, group of chalcogens. The name is derived from the greek word for brass and points to the association of sulphur and its, congeners with copper. Most copper minerals contain either oxygen or sulphur and frequently the other members of the, group., Occurrence :, Oxygen is the most abundant of all the elements on the earth. Oxygen forms about 46.6% by mass of earth’s crust . Dry air, contains 20.946% oxygen by volume., However, the abundance of sulphur in the earth’s crust is only 0.03-0.1%. Combined sulphur exists primarily as sulphates, such as gypsum CaSO4.2H2O, epsom salt MgSO4 .7H2O, baryta BaSO4 and sulphides such as galena PbS, zinc blende ZnS,, copper pyrites CuFeS2 . Traces of sulphur occur as hydrogen sulphide in volcanoes., Selenium and tellurium are also found as metal selenides and tellurides in sulphide ores. Polonium occurs in nature as a, decay product of thorium and uranium minerals., Electronic Configuration :, The elements of group 16 have six electrons in the outermost shell and have ns2 np4 general electronic configuration., Atomic and Ionic Radii :, Due to increase in the number of shells , atomic and ionic radii increase from top to bottom in the group. The size of oxygen, atoms is however, exceptionally small ., Ionisation Enthalpy :, Ionisation enthalpy decrease down the group. It is due to increase in size. However, the element of this group have lower, ionisation enthalpy values compared to those of group 15 in the corresponding periods. This is due to the fact that group 15, element have extra stable half-filled p orbitals elelctronic configurations., Electron Gain Enthalpy :, Because of the compact nature of oxygen atom, it has less negative electron gain enthalpy than sulphur. However from, sulphur onwards the value again becomes less negative upto polonium., Electronegativity :, Next to fluorine, oxygen has the highest electronegativity value amongst the elements. Within the group, electronegativity, decrease with an increase in atomic number. This implies that the metallic character increase from oxygen to polonium., Physical Properties :, Oxygen and sulphur are non-metal, selenium and tellurium metalloids, whereas polonium is a metal. Polonium is radioactive, and is short lived (Half-life 13.8 days). All these element exhibit allotropy. The melting and boiling points increase with an, increase in atomic number down the group. The larger difference between the melting and boiling points of oxygen and, sulphur may be explained on the basis of their atomicity; oxygen exist as diatomic molecules (O2) whereas sulphur exists as, polyatomic molecule (S8).

Page 15 : p-BLOCK ELEMENTS, ATOMIC & PHYSICAL PROPERTIES :, Element, , O, , S, , Se, , Te, , Atomic Number, , 8, , 16, , 34, , 52, , Atomic Mass, , 16, , 32.06, , 78.96, , 127.6, , Electronic configuration, , [He] 2s2 2p4, , [Ne] 3s2 3p4, , [Ar] 3d104s2 4p4, , [Kr] 4d105s2 5p4, , Covalent Radius / pm, , 74, , 103, , 119, , 142, , Ionic Radius X–2 / pm, , 140, , 184, , 198, , 221, , , , 1314, , 1000, , 941, , 869, , , , 3388, , 2251, , 2045, , 1790, , 3.5, , 2.44, , 2.48, , 2.01, , Ionization enthalpy / (kJ mol–1), Electronegativity, –3, , 1.32, , 2.06, , 4.19, , 6.25, , Melting point / K, , 54, , 393, , 490, , 725, , Boiling point / K, , 90, , 718, , 958, , 1260, , Density/[g cm, , (293 K)], , Chemical Properties :, Oxidation states and trends in chemical reactivity :, The elements of group 16 exhibit a number of oxidation states. The stability of -2 oxidation state decreases down the group., Polonium hardly shows -2 oxidation states. Since electronegativity of oxygen is very high, it shows only negative oxidation, states as -2 except in the case of OF2 where its oxidation states is + 2. Other elements of the group exhibit + 2 + 4 + 6, oxidation states but + 4 and + 6 are more common. Sulphur, selenium and tellurium usually show + 4 oxidation in their, compounds with oxygen and +6 oxidations state with fluorine. The stability of +6 oxidation state decreases down the group, and stability of + 4 oxidation state increases (inert pair effect). Bonding in + 4 and + 6 oxidation states are primarily covalent., Anomalous behaviour of oxygen :, The anomalous behaviour of oxygen, like other member of p-block present in second period is due to its small size and high, electronegativity. One typical example of effects of small size and high electronegativity is the presence of strong hydrogen, bonding in H2O which is not found in H2S., The absence of d orbitals in oxygen limits its covalency to four and in practice, rarely exceeds two. On the other hand, in case, of other elements of the group, the valence shell can be expanded and covalence exceeds four., (i), Reactivity with hydrogen : All the elements of group 16 form hydrides of the type H2E (E = S, Se., Te, Po). Some properties, of hydrides are given in Table. Their acidic character increases from H2O to H2Te. The increase in acidic character can be, explained in terms of decrease in bond (H-E) dissociation enthalpy down the group. Owing to the decrease in bond (H-E), dissociation enthalpy down the group, the thermal stability of hydrides also decreases from H2O to H2Po. All the hydrides, except water possess reducing property and this character increases from H2S to H2Te., Table : Properties of Hydrides of Group 16 Elements, , Property, , H2 O, , H2 S, , H2Se, , H2Te, , m.p./K, , 273, , 188, , 208, , 222, , b.p./K, , 373, , 213, , 232, , 269, , H-E distance/pm, , 96, , 134, , 146, , 169, , HEH angle (º), , 104, , 92, , 91, , 90, , fH/kJ mol-1, , -286, , -20, , 73, , 100, , diss H (H-E)/kJ mol-1, , 463, , 347, , 276, , 238, , 1.8 × 10-16, , 1.3 × 10-7, , 1.3 × 10-4, , 2.3 × 10-3, , Dissociation constanta

Page 22 : p-BLOCK ELEMENTS, , (iii), , When this reaction is carried out in dark, it is accompanied by emission of light (yellow coloured). It is an example of, chemiluminescene., An acidified solution of titanium salt gives yellow or orange colour with H2O2., Ti+4 + H2O2 + 2H2O  H2TiO4 (yellow/orange) + 4H+, , USES :, (i), , In bleaching of delicate materials such as silk, wool, cotton, ivory etc., , (ii), (iii), , As a valuable antiseptic and germicide for washing wounds, teeth and ears under the name perhydrol., As ‘antichlor’ to remove traces of chlorine and hypochlorite., , (iv), , As oxidising agent in rocket fuels, , 4., , SULPHUR (S) :, , Sulphur Allotropic Froms :, Sulphur forms numerous allotropes of which the yellow rhombic ( - sulphur) and monoclinic ( - sulphur) forms are the, most important. The stable forms at room temperature is rhombic sulphur, which transfroms to monoclinic sulphur when, heated above 369 K., Rhombic sulphur (- sulphur) :, This allotrope is yellow in colour , m.p. 385.8 K and specific gravity 2.06. Rhombic sulphur crystals are formed on evaporating, the solution of roll sulphur in CS2. It is insoluble in water but dissolved to some extent in benzene, alcohol and ether. It is, readily soluble in CS2 ., Monoclinic sulphur ( - sulphur) :, Its m.p. is 393 K and specific gravity 1.98. It is soluble in CS2. This form of sulphur is prepared by melting rhombic sulphur, in a dish and cooling till crust is formed. Two holes are made in the crust and the remaining liquid poured out. On removing, the crust, colourless needle shaped crystals of  - sulphur are formed. It is stable above 369 K and transforms into  - sulphur, below it . Conversely, - sulphur is stable below 369 K and transfroms into  - sulphur above this. At 369 K both the forms, are stable. This temperature is called transition temperature., Both rhombic and monoclinic sulphur have S8 molecules these S8 molecules are packed to give different crystal structures., The S8 ring in both the forms is puckered and has a crown shape. The molecular dimensions are given in figure., S, S, , 20, 4, , pm, , S, , S, 107o, , S, , 205.7 pm, , S, , o, , 102.2, , S, , S, S, , S, , S, S, , (a), , S, (b), S, , Fig. : The structures of (a) S8 ring in rhombic sulphur and (b) S6 form, 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 moleculatr dimension are as shown in fig. (b) At elevated temperatures (~, 1000 K ), S2 is the dominant species and is paramagnetic like O2., , 3., , COMPOUNDS OF SULPHUR :, , (A) SODIUM THIOSULPHATE (Na2S2O3 .5H2O) :, , PREPARATION :, (i), , Na2SO3 + S, , Na2S2O3, , (ii), (iii), , Na2CO3 + 2SO2 (excess) + H2O  2NaHSO3 + CO2 ; 2NaHSO3 + Na2CO3  2Na2S2O3 + H2O + CO2, 2 NaHS + 4NaHSO3  3Na2S2O3 + 3H2O

Page 28 : p-BLOCK ELEMENTS, (vi), , Dehydrating agent :, Sulphuric acid acts as a powerful dehydrating agent because it has a great affinity for water, (a), , H2SO 4, , ,  CO + CO2 + H2O, , , (b), (vii), (a), , Miscellaneous reactions :, Sulphonation of aromatic compounds, , + H2SO4 , , (b), , 12C, , C12H22O11 (cane sugar), , + H2O, , Benzene, Benzene sulphonic acid, Reaction with PCl5 :, , OH, , Cl, + PCl5  O2S, , O2S, OH, , OH, Chlorosulphonic acid, , OH, , Cl, + 2PCl5  O2S, , O2S, , + POCl3 + HCl, , + 2 POCl3 + 2HCl, , OH, , (c), (d), (e), , Cl, Sulphury chloride, K4[Fe(CN)6] + 6H2SO4 + 6H2O  2K2SO4 + FeSO4 + 3(NH4)2SO4 + 6CO, , 3 KHSO4 + HClO4 + 2ClO2 + H2O, 3KClO3 + 3H2SO4 , H2SO4 + P2O5  2HPO3 + SO3, , USES :, (i), (ii), (iii), (iv), , For the manufacture of fertilizer such as ammonium sulphate and super phosphate of lime., As an important laboratory reagent., In storage batteries., In leather, textile, paper and dyeing industries., , 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 dereases down the group., Electron Gain Enthalpy :, Halogen have maximum negative electrons 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 incoming electron does not experience much attraction.

Page 29 : p-BLOCK ELEMENTS, 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 and iodine is a solid. Their melting and boiling points steadily increase, with atomic number. All halogen 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 whereas X-X bond dissocitation enthalpies from chlorine onwards show the expected trend: Cl –, Cl > Br – Br > I – I . A reason for this anomaly is the relatively larger electrons- electron repulsion among the lone pairs in, F2 molecule where they are much closer to each other than in case of Cl2., , ATOMIC & PHYSICAL PROPERTIES :, Element, , F, , Cl, , Br, , I, , Atomic Number, , 9, , 17, , 35, , 53, , Atomic Mass, , 19, 2, , Electronic configuration, Covalent Radius / pm, –, , 35.45, 5, , 2, , 79.90, 5, , 10, , 126.90, 2, , 5, , [He] 2s 2p, , [Ne] 3s 3p, , [Ar] 3d 4s 4p, , [Kr] 4d105s2 5p5, , 64, , 99, , 114, , 133, , 133, , 184, , 196, , 220, , Ionization enthalpy / (kJ mol ), , 1680, , 1256, , 1142, , 1008, , Electron gain enthalpy /(kJ mol–1), , – 333, , – 349, , – 325, , – 296, , Ionic Radius X / pm, –1, , 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, , Distance X -X/pm, Enthalpy of dissociation (X2)/kJ mol, , –1, , 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, halogen 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 ()

Page 30 : p-BLOCK ELEMENTS, Standard Reduction Potential (SRP), X + 2e–  2X–, 2, , , , , 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 (algebraically) oxidising agent. Hence the order of oxidising power is F2 >, Cl2 > Br2 > 2, Since SRP is the highest for F2 (among all elements of P.T.), it is a strogenst oxidising agent., , F2 is more powerful oxidising agent than O3 [Inspite of 3 ‘O’s in O3 ], E.A. and .E. values pertain to atoms in gas phase where as redox phenomena occurs in gaseous medium., Hence properties in the gas phase cannot reflect parallely in solution phase., Electrode potential values would be the monitoring parameter in solution phase because they are experimental (based on the, correct situation)., Hydration energy of X–, Smaller the ion, higher is the hydration energy, F–, Cl –, Br –, –, 515 kJ/mol, 381, 347, 305, Anomalous behaviour of fluorine :, The anomalous behviour 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 halogen 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. They, dissolve in water to form hydrohalic acids. The acidic strength of these acids varies in the order : HF < HCl < HBr < HI. The, stability of these halides dereases 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 reactions 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 range from + 1 to + 7. A combination of kinetic and thermodynamic factors lead to the generally decreasing order of stability 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 and B is smaller size halogen., Note :

Page 39 : p-BLOCK ELEMENTS, H reacts with CuSO4 liberating iodine via the formation of cupric iodide (not by HCl or HBr)., , (xi), , 2CuSO4 + 4H  2Cu2 + 2H2SO4 ; 2Cu2  Cu2 2 + 2, Formation of aqua-regia :, 3 parts of conc. HCl and 1 part of conc. HNO 3 is known as aqua-regia. This is used for dissolving, noble metals like Au (Gold) and Pt (Platinum)., Au + 3Cl  AuCl, 3HCl + HNO  NOCl + 2H O + 2Cl (reactive) ;, 2, , 3, , 3, , USES :, (i), , HCl is used in preparation of Cl2, chlorides, aqua regia, glucose, (from corn starch), medicines, laboratory reagents, cleaning, metal surfaces before soldering or electroplating. It is also used for extracting glue from bones and purifying bone black., , (ii), , HBr is used as laboratory reagent for preparing bromo derivatives like sodium bromides and potassium bromide. These are, used in medicines as sedatives., , (iii), , H is used as reducing agent in organic chemistry., , HYDROFLUORIC ACID [H2 F2, HF] :, PREPARATION :, , (i), , (ii), , 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., Laboratory Method :, Anhydrous HF is obtained by heating dry potassium hydrogen fluoride in a copper retort connected with copper condenser., , , KF + HF, KHF2 , , Industrial Method :, HF is prepared by heating fluorspar (CaF2) with conc 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:, Na SiO + 6HF  Na SiF + 3H O ; CaSiO + 6HF  CaSiF + 3H O, 2, , 2, , 3, , 6, , 2, , 3, , 6, , 2, , 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 this is then treated with 40% solution., , PROPERTIES:, (i), (ii), , (iii), (iv), , 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 + H F  NaHF + H O ; NaHF + NaOH  2NaF + H O, 2, , 2 2, , 2, , 2, , 2, , Concentrated acid reacts with oxides, hydroxides and carbonates., 2NaOH + H2 F2  2NaF + 2H2O : CaCO3 + 2HF  CaF2 + H2O + CO2, , OXY-ACIDS OF CHLORINE :, 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 + 2Cl + H O  Hg OCl (Oxychloride of mercury) + 2HClO, 2, , 2, , 2, , 2

Page 45 : p-BLOCK ELEMENTS, 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 monoatomic. 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., , ATOMIC & PHYSICAL PROPERTIES :, Element, , He, , Ne, , Ar, , Kr, , Xe, , Atomic Number, , 2, , 10, , 18, , 36, , 54, , Atomic Mass, , 4, , Electronic configuration, Atomic Radius (pm), –1, , Ionization enthalpy / (kJ mol ), –3, , Density (at STP)/g cm, , 20.18, , 39.10, , 2, , [Ne] 3s 3p, , [Ar] 3d 4s 4p, , [Kr] 4d105s2 5p6, , 120, , 160, , 190, , 200, , 220, , 2080, –4, , 1.8 × 10, , 2, , 6, , 1520, –4, , 9.0 × 10, , 10, , 131.30, , [He] 2s 2p, , 2372, , 6, , 83.80, , 1s, , 2, , 2, , 1351, –3, , 1.8 × 10, , 6, , 1170, –3, , 3.7 × 10, , 5.9 × 10–3, , 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), The noble gases expect helium (1S2) have completely filled ns2 np6 electronic configuration in their valence shell., (ii), 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., The compounds of krypton are fewer. Only the difluoride (KrF2) has been studied in detali. Compounds of radon have not, , , , (i), , 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., Compounds of inert gases are of following two types., Physical compounds (possess no proper bonding), Physical compounds may be (A) compounds whose existence is on the basis of spectroscopic studies (temporary phase not, isolated) and (B) clatherate compounds.

Page 46 : p-BLOCK ELEMENTS, 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)., Xe, Kr or Rn (g), (Not He or Ne), , H2 O, (cool), , (ii), , True chemical compounds (posses proper bonding)., , COMPOUNDS OF XENON :, (a) XENON DIFLUORIDE (XeF2):, , PREPARATION:, 873 K ,1 bar,            XeF2, NiTube or monel metal (alloy of Ni), , (i), , Xe + F2, , , , 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., , 1, O + 2NaF + H2O (fast), 2 2, The reaction (a) is slower probably due to dissolution of XeF2 in HF., XeF2 + 2 NaOH  Xe +

Page 50 : p-BLOCK ELEMENTS, , , , Order of oxidising power :, XeF2 > XeF4 > XeF6, , STRUCTURE : Shape caped octahedral (distorted oactahedral), , (d) 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 ballons for meterological 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.