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Unit Objectives The p -Block 7 After studying this Unit, you will be able to...

Unit Objectives The p -Block 7 After studying this Unit, you will be able to Elements appreciate general trends in the chemistry of elements of groups 15,16,17 and 18; Diversity in chemistry is the hallmark of p–block elements manifested learn the preparation, properties in their ability to react with the elements of s–, d– and f–blocks as and uses of dinitrogen and well as with their own. phosphorus and some of their important compounds; describe the preparation, In Class XI, you have learnt that the p-block elements properties and uses of dioxygen are placed in groups 13 to 18 of the periodic table. and ozone and chemistry of some Their valence shell electronic configuration is ns2np1–6 simple oxides; 2 (except He which has 1s configuration). The properties know allotropic forms of sulphur, of p-block elements like that of others are greatly chemistry of its important influenced by atomic sizes, ionisation enthalpy, electron compounds and the structures of gain enthalpy and electronegativity. The absence of d- its oxoacids; orbitals in second period and presence of d or d and f describe the preparation, orbitals in heavier elements (starting from third period properties and uses of chlorine onwards) have significant effects on the properties of and hydrochloric acid; elements. In addition, the presence of all the three types know the chemistry of interhalogens and structures of of elements; metals, metalloids and non-metals bring oxoacids of halogens; diversification in chemistry of these elements. enumerate the uses of noble Having learnt the chemistry of elements of Groups gases; 13 and 14 of the p-block of periodic table in Class XI, appreciate the importance of you will learn the chemistry of the elements of these elements and their subsequent groups in this Unit. compounds in our day to day life. 7.1 Group 15 Group 15 includes nitrogen, phosphorus, arsenic, antimony, bismuth Elements and moscovium. As we go down the group, there is a shift from non- metallic to metallic through metalloidic character. Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids, bismuth and moscovium are typical metals. 7.1.1 Occurrence Molecular nitrogen comprises 78% by volume of the atmosphere. In the earth’s crust, it occurs as sodium nitrate, NaNO3 (called Chile saltpetre) and potassium nitrate (Indian saltpetre). It is found in the form of proteins in plants and animals. Phosphorus occurs in minerals Chemistry 170 2022-23  of the apatite family, Ca9(PO4)6. CaX2 (X = F, Cl or OH) (e.g., fluorapatite Ca9 (PO4)6. CaF2) which are the main components of phosphate rocks. Phosphorus is an essential constituent of animal and plant matter. It is present in bones as well as in living cells. Phosphoproteins are present in milk and eggs. Arsenic, antimony and bismuth are found mainly as sulphide minerals. Moscovium is a synthetic radioactive element. Its symbol is Mc, atomic number 115, atomic mass 289 and electronic 14 10 2 3 configuration [Rn] 5f 6d 7s 7p. Due to very short half life and availability in very little amount, its chemistry is yet to be established. Here, except for moscovium, important atomic and physical properties of other elements of this group along with their electronic configurations are given in Table 7.1. Table 7.1: Atomic and Physical Properties of Group 15 Elements Property N P As Sb Bi Atomic number 7 15 33 51 83 –1 Atomic mass/g mol 14.01 30.97 74.92 121.75 208.98 2 3 2 3 10 2 3 10 2 3 14 10 2 3 Electronic configuration [He]2s 2p [Ne]3s 3p [Ar]3d 4s 4p [Kr]4d 5s 5p [Xe]4f 5d 6s 6p Ionisation enthalpy I 1402 1012 947 834 703 –1 (∆iH/(kJ mol ) II 2856 1903 1798 1595 1610 III 4577 2910 2736 2443 2466 Electronegativity 3.0 2.1 2.0 1.9 1.9 a Covalent radius/pm 70 110 121 141 148 b b b c c Ionic radius/pm 171 212 222 76 103 d e Melting point/K 63* 317 1089 904 544 d f Boiling point/K 77.2* 554 888 1860 1837 –3 g h Density/[g cm (298 K)] 0.879 1.823 5.778 6.697 9.808 a EIII single bond (E = element); b E 3–; c E 3+; d White phosphorus; e Grey α-form at 38.6 atm; f Sublimation temperature; g At 63 K; hGrey α-form; * Molecular N2. Trends of some of the atomic, physical and chemical properties of the group are discussed below. 2 3 7.1.2 Electronic The valence shell electronic configuration of these elements is ns np. Configuration The s orbital in these elements is completely filled and p orbitals are half-filled, making their electronic configuration extra stable. 7.1.3 Atomic and Covalent and ionic (in a particular state) radii increase in size Ionic Radii 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. 7.1.4 Ionisation Ionisation enthalpy decreases down the group due to gradual increase Enthalpy in atomic size. Because of the extra stable half-filled p orbitals electronic configuration and smaller size, the ionisation enthalpy of the group 15 elements is much greater than that of group 14 elements in the corresponding periods. The order of successive ionisation enthalpies, as expected is ∆iH1 < ∆iH2 < ∆iH3 (Table 7.1). 171 The p-Block Elements 2022-23 7.1.5 The electronegativity value, in general, decreases down the group with Electronegativity increasing atomic size. However, amongst the heavier elements, the difference is not that much pronounced. 7.1.6 Physical All the elements of this group are polyatomic. Dinitrogen is a diatomic gas Properties while all others are solids. Metallic character increases down the group. Nitrogen and phosphorus 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. 7.1.7 Chemical Oxidation states and trends in chemical reactivity Properties The common oxidation states of these elements are –3, +3 and +5. The tendency to exhibit –3 oxidation state decreases down the group due to increase in size and metallic character. In fact last member of 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. Besides +5 oxidation state, nitrogen exhibits + 1, + 2, + 4 oxidation states also when it reacts with oxygen. However, it does not form compounds in +5 oxidation state with halogens as nitrogen does not have d-orbitals to accommodate electrons from other elements to form bonds. Phosphorus 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, 3HNO2 → HNO3 + H2O + 2NO 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 becomes increasingly stable with respect to disproportionation. 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 small size, high electronegativity, high ionisation enthalpy and non-availability of d orbitals. Nitrogen has unique ability to form pπ -p π 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 pπ -pπ 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 –1 (941.4 kJ mol ) is very high. On the contrary, phosphorus, arsenic and antimony form single bonds as P–P, As–As and Sb–Sb while bismuth forms metallic bonds in elemental state. However, the single Chemistry 172 2022-23 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 dπ –pπ bond as the heavier elements can e.g., R3P = O or R3P = CH2 (R = alkyl group). Phosphorus and arsenic can form dπ –dπ 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 7.2. 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. Due to high electronegativity and small size of nitrogen, NH3 exhibits hydrogen bonding in solid as well as liquid state. Because of this, it has higher melting and boiling points than that of PH3. Table 7.2: Properties of Hydrides of Group 15 Elements Property NH 3 PH 3 AsH3 SbH 3 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 – HEH angle (°) 107.8 93.6 91.8 91.3 – V –1 ∆f H /kJ mol –46.1 13.4 66.4 145.1 278 V –1 ∆dissH (E–H)/kJ mol 389 322 297 255 – (ii) 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 phosphorus are purely acidic, that of arsenic and antimony amphoteric and those of bismuth predominantly basic. (iii) 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. This is due to the fact that in pentahalides +5 oxidation state exists while in the case of trihalides +3 oxidation state exists. Since elements in +5 oxidation 173 The p-Block Elements 2022-23  state will have more polarising power than in +3 oxidation state, the covalent character of bonds is more in pentahalides. 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. (iv) Reactivity towards metals: All these elements react with metals to form their binary compounds exhibiting –3 oxidation state, such as, Ca3N2 (calcium nitride) Ca3P2 (calcium phosphide), Na3As (sodium arsenide), Zn3Sb2 (zinc antimonide) and Mg3Bi2 (magnesium bismuthide). Though nitrogen exhibits +5 oxidation state, it does not form Example 7.1 pentahalide. Give reason. Nitrogen with n = 2, has s and p orbitals only. It does not have d Solution orbitals to expand its covalence beyond four. That is why it does not form pentahalide. PH3 has lower boiling point than NH3. Why? Example 7.2 Unlike NH3, PH3 molecules are not associated through hydrogen bonding Solution in liquid state. That is why the boiling point of PH3 is lower than NH3. Intext Questions 7.1 Why are pentahalides of P, As, Sb and Bi more covalent than their trihalides? 7.2 Why is BiH3 the strongest reducing agent amongst all the hydrides of Group 15 elements ? 7.2 Dinitrogen Preparation Dinitrogen is produced commercially by the liquefaction and fractional distillation of air. Liquid dinitrogen (b.p. 77.2 K) distils out first leaving behind liquid oxygen (b.p. 90 K). In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite. NH4CI(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl (aq) Small amounts of NO and HNO3 are also formed in this reaction; these impurities can be removed by passing the gas through aqueous sulphuric acid containing potassium dichromate. It can also be obtained by the thermal decomposition of ammonium dichromate. (NH4)2Cr2O7  Heat → N2 + 4H2O + Cr2O3 Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide. Ba(N3)2 → Ba + 3N2 Chemistry 174 2022-23  Properties Dinitrogen is a colourless, odourless, tasteless and non-toxic gas. 14 15 Nitrogen atom has two stable isotopes: N and N. It has a very low 3 solubility in water (23.2 cm per litre of water at 273 K and 1 bar pressure) and low freezing and boiling points (Table 7.1). Dinitrogen is rather inert at room temperature because of the high bond enthalpy of N ≡ N bond. Reactivity, however, increases rapidly with rise in temperature. At higher temperatures, it directly combines with some metals to form predominantly ionic nitrides and with non-metals, covalent nitrides. A few typical reactions are: 6Li + N2  Heat → 2Li3N 3Mg + N2  Heat → Mg3N2 It combines with hydrogen at about 773 K in the presence of a catalyst (Haber’s Process) to form ammonia: y –1 N2(g) + 3H2(g) ‡ 773 k ˆˆˆˆˆ ˆˆˆˆ ˆ 2NH3(g); † ∆f H = –46.1 kJmol Dinitrogen combines with dioxygen only at very high temperature (at about 2000 K) to form nitric oxide, NO. Heat N2 + O2(g) ‡ ˆˆˆˆˆ ˆˆˆˆ ˆ 2NO(g) † Uses Uses: The main use of dinitrogen is in the manufacture of ammonia and other industrial chemicals containing nitrogen, (e.g., calcium cyanamide). It also finds use where an inert atmosphere is required (e.g., in iron and steel industry, inert diluent for reactive chemicals). Liquid dinitrogen is used as a refrigerant to preserve biological materials, food items and in cryosurgery. Example 7.3 Write the reaction of thermal decomposition of sodium azide. Solution Thermal decomposition of sodium azide gives dinitrogen gas. 2NaN 3 → 2Na + 3N 2 Intext Question 7.3 Why is N2 less reactive at room temperature? 7.3 Ammonia Preparation Ammonia is present in small quantities in air and soil where it is formed by the decay of nitrogenous organic matter e.g., urea. NH2 CONH2 + 2H2O → ( NH4 )2CO3 ƒ 2NH3 + H2O + CO2 On a small scale ammonia is obtained from ammonium salts which decompose when treated with caustic soda or calcium hydroxide. 2NH4Cl + Ca(OH)2 → 2NH3 + 2H2O + CaCl2 (NH4)2 SO4 + 2NaOH → 2NH3 + 2H2O + Na2SO4 175 The p-Block Elements 2022-23  On a large scale, ammonia is manufactured by Haber’s process. 0 –1 N2(g) + 3H2(g) Ö 2NH3(g); ∆f H = – 46.1 kJ mol In accordance with Le Chatelier’s principle, high pressure would favour the formation of ammonia. The optimum conditions for the 5 production of ammonia are a pressure of 200 × 10 Pa (about 200 atm), a temperature of ~ 700 K and the use of a catalyst such as iron oxide with small amounts of K2O and Al2O3 to increase the rate of attainment of equilibrium. The flow chart for the production of ammonia is shown in Fig. 7.1. Earlier, iron was used as a catalyst with molybdenum as a promoter. Fig. 7.1 Flow chart for the manufacture of ammonia Properties Ammonia is a colourless gas with a pungent odour. Its freezing and boiling points are 198.4 and 239.7 K respectively. In the solid and N liquid states, it is associated through hydrogen bonds as in the case H H of water and that accounts for its higher melting and boiling points H than expected on the basis of its molecular mass. The ammonia molecule is trigonal pyramidal with the nitrogen atom at the apex. It has three bond pairs and one lone pair of electrons as shown in the structure. Ammonia gas is highly soluble in water. Its aqueous solution is – weakly basic due to the formation of OH ions. + – NH3(g) + H2O(l) l NH4 (aq) + OH (aq) It forms ammonium salts with acids, e.g., NH4Cl, (NH4)2 SO4, etc. As a weak base, it precipitates the hydroxides (hydrated oxides in case of some metals) of many metals from their salt solutions. For example, ZnSO4 ( aq ) + 2NH4 OH ( aq ) → Zn ( OH )2 ( s ) + ( NH4 )2 SO4 ( aq ) ( white ppt ) FeCl3 (aq ) + NH 4OH (aq ) → Fe 2O3.x H2 O (s ) + NH4 Cl (aq ) ( brown ppt ) Chemistry 176 2022-23  The presence of a lone pair of electrons on the nitrogen atom of the ammonia molecule makes it a Lewis base. It donates the electron pair and forms linkage with metal ions and the formation of such complex compounds finds applications in detection of metal ions 2+ + such as Cu , Ag : 2+ 2+ Cu (aq) + 4 NH3(aq) Ö [Cu(NH3)4] (aq) (blue) (deep blue) Ag + (aq ) + Cl − ( aq ) → AgCl ( s ) (colourless) (white ppt) AgCl (s ) + 2NH3 (aq ) →  Ag ( NH3 )2  Cl (aq ) (white ppt) (colourless) Uses Uses: Ammonia is used to produce various nitrogenous fertilisers (ammonium nitrate, urea, ammonium phosphate and ammonium sulphate) and in the manufacture of some inorganic nitrogen compounds, the most important one being nitric acid. Liquid ammonia is also used as a refrigerant. Example 7.4 Why does NH3 act as a Lewis base ? Solution Nitrogen atom in NH3 has one lone pair of electrons which is available for donation. Therefore, it acts as a Lewis base. Intext Questions 7.4 Mention the conditions required to maximise the yield of ammonia. 7.5 How does ammonia react with a solution of Cu2+? 7.4 Oxides of Nitrogen forms a number of oxides in different oxidation states. The names, formulas, preparation and physical appearance of these oxides Nitrogen are given in Table 7.3. Table 7.3: Oxides of Nitrogen Name Formula Oxidation Common Physical state of methods of appearance and nitrogen preparation chemical nature NH4 NO3 → Heat Dinitrogen oxide N2O + 1 colourless gas, N 2O + 2H 2O [Nitrogen(I) oxide] neutral Nitrogen monoxide NO + 2 2NaNO2 + 2FeSO 4 + 3H2SO 4 colourless gas, [Nitrogen(II) oxide] → Fe 2 (SO 4 )3 + 2NaHSO 4 neutral + 2H2O + 2NO 177 The p-Block Elements 2022-23  250 K Dinitrogen trioxide N2O3 + 3 2NO + N 2O 4 → 2N 2O 3 blue solid, [Nitrogen(III) oxide] acidic 2Pb(NO3 )2  673K → Nitrogen dioxide NO2 + 4 brown gas, 4NO2 + 2PbO + O2 [Nitrogen(IV) oxide] acidic Cool Dinitrogen tetroxide N2O4 + 4  2NO2    N 2O 4 colourless solid/ Heat [Nitrogen(IV) oxide] liquid, acidic Dinitrogen pentoxide N2O5 +5 4HNO3 + P4O10 colourless solid, [Nitrogen(V) oxide] → 4HPO3 + 2N 2O5 acidic Lewis dot main resonance structures and bond parameters of oxides are given in Table 7.4. Table 7.4: Structures of Oxides of Nitrogen Chemistry 178 2022-23  Why does NO2 dimerise ? Example 7.5 NO2 contains odd number of valence electrons. It behaves as a typical Solution odd molecule. On dimerisation, it is converted to stable N2O4 molecule with even number of electrons. Intext Question 7.6 What is the covalence of nitrogen in N2O5 ? 7.5 Nitric Acid Nitrogen forms oxoacids such as H2N 2O 2 (hyponitrous acid), HNO 2 (nitrous acid) and HNO 3 (nitric acid). Amongst them HNO3 is the most important. Preparation In the laboratory, nitric acid is prepared by heating KNO3 or NaNO3 and concentrated H2SO4 in a glass retort. NaNO3 + H2 SO4 → NaHSO4 + HNO3 On a large scale it is prepared mainly by Ostwald’s process. This method is based upon catalytic oxidation of NH3 by atmospheric oxygen. 4NH3 ( g ) + 5O2 ( g )  Pt / Rh gauge catalyst 500 K , 9 bar → 4NO ( g ) + 6H2 O ( g ) (from air) Nitric oxide thus formed combines with oxygen giving NO2. 2NO ( g ) + O2 ( g )
2NO2 ( g ) Nitrogen dioxide so formed, dissolves in water to give HNO3. 3NO2 ( g ) + H2 O ( l ) → 2HNO3 ( aq ) + NO ( g ) NO thus formed is recycled and the aqueous HNO 3 can be concentrated by distillation upto ~ 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated H 2SO4. Properties It is a colourless liquid (f.p. 231.4 K and b.p. 355.6 K). Laboratory grade nitric acid contains ~ 68% of the HNO3 by mass and has a specific gravity of 1.504. In the gaseous state, HNO3 exists as a planar molecule with the structure as shown. In aqueous solution, nitric acid behaves as a strong acid giving hydronium and nitrate ions. + – HNO3(aq) + H2O(l) → H3O (aq) + NO3 (aq) Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals such as gold and platinum. The 179 The p-Block Elements 2022-23  products of oxidation depend upon the concentration of the acid, temperature and the nature of the material undergoing oxidation. 3Cu + 8 HNO3(dilute) → 3Cu(NO3)2 + 2NO + 4H2O Cu + 4HNO3(conc.) → Cu(NO3)2 + 2NO2 + 2H2O Zinc reacts with dilute nitric acid to give N2O and with concentrated acid to give NO2. 4Zn + 10HNO3(dilute) → 4 Zn (NO3)2 + 5H2O + N2O Zn + 4HNO3(conc.) → Zn (NO3)2 + 2H2O + 2NO2 Some metals (e.g., Cr, Al) do not dissolve in concentrated nitric acid because of the formation of a passive film of oxide on the surface. Concentrated nitric acid also oxidises non–metals and their compounds. Iodine is oxidised to iodic acid, carbon to carbon dioxide, sulphur to H2SO4, and phosphorus to phosphoric acid. I2 + 10HNO3 → 2HIO3 + 10NO2 + 4H2O C + 4HNO3 → CO2 + 2H2O + 4NO2 S8 + 48HNO3 → 8H2SO4 + 48NO2 + 16H2O P4 + 20HNO3 → 4H3PO4 + 20NO2 + 4H2O Brown Ring Test: The familiar brown ring test for nitrates depends 2+ on the ability of Fe to reduce nitrates to nitric oxide, which reacts 2+ with Fe to form a brown coloured complex. The test is usually carried out by adding dilute ferrous sulphate solution to an aqueous solution containing nitrate ion, and then carefully adding concentrated sulphuric acid along the sides of the test tube. A brown ring at the interface between the solution and sulphuric acid layers indicates the presence of nitrate ion in solution. - 2+ + 3+ NO3 + 3Fe + 4H → NO + 3Fe + 2H2O [Fe (H2O)6 ]2+ + NO → [Fe (H2O)5 (NO)] + H2O 2+ (brown) Uses Uses: The major use of nitric acid is in the manufacture of ammonium nitrate for fertilisers and other nitrates for use in explosives and pyrotechnics. It is also used for the preparation of nitroglycerin, trinitrotoluene and other organic nitro compounds. Other major uses are in the pickling of stainless steel, etching of metals and as an oxidiser in rocket fuels. 7.6 Phosphorus — Phosphorus is found in many allotropic forms, the important ones Allotropic being white, red and black. White phosphorus is a translucent white waxy solid. It is poisonous, Forms insoluble in water but soluble in carbon disulphide and glows in dark (chemiluminescence). It dissolves in boiling NaOH solution in an inert atmosphere giving PH3. P4 + 3NaOH + 3H2O → PH3 + 3NaH2 PO2 (sodium hypophosphite ) Chemistry 180 2022-23  P White phosphorus is less stable and therefore, more reactive than the other solid phases under normal conditions because of angular 60° strain in the P4 molecule where the angles are only 60°. It readily P P catches fire in air to give dense white fumes of P4O10. P4 + 5O2 → P4 O10 P It consists of discrete tetrahedral P4 molecule as shown in Fig. 7.2. Fig. 7.2 Red phosphorus is obtained by heating white phosphorus at 573K White phosphorus in an inert atmosphere for several days. When red phosphorus is heated under high pressure, a series of phases of black phosphorus is formed. Red phosphorus possesses iron grey lustre. It is odourless, non- poisonous and insoluble in water as well as in carbon disulphide. Chemically, red phosphorus is much less reactive than white phosphorus. It does not glow in the dark. P P P It is polymeric, consisting of chains of P 4 tetrahedra linked together in the manner as shown in Fig. 7.3. P P P P P P Black phosphorus has two forms α-black P P phosphorus and β-black phosphorus. α-Black P phosphorus is formed when red phosphorus is Fig.7.3: Red phosphorus heated in a sealed tube at 803K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystals. It does not oxidise in air. β-Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It does not burn in air upto 673 K. 7.7 Phosphine Preparation Phosphine is prepared by the reaction of calcium phosphide with water or dilute HCl. Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3 Ca3P2 + 6HCl → 3CaCl2 + 2PH3 In the laboratory, it is prepared by heating white phosphorus with concentrated NaOH solution in an inert atmosphere of CO2. P4 + 3NaOH + 3H2O → PH3 + 3NaH2 PO2 ( sodium hypophosphite ) When pure, it is non inflammable but becomes inflammable owing to the presence of P2H4 or P4 vapours. To purify it from the impurities, it is absorbed in HI to form phosphonium iodide (PH4I) which on treating with KOH gives off phosphine. PH 4 I + KOH → KI + H 2 O + PH 3 Properties It is a colourless gas with rotten fish smell and is highly poisonous. It explodes in contact with traces of oxidising agents like HNO3, Cl2 and Br2 vapours. It is slightly soluble in water. The solution of PH3 in water decomposes in presence of light giving red phosphorus and H2. When absorbed in 181 The p-Block Elements 2022-23 copper sulphate or mercuric chloride solution, the corresponding phosphides are obtained. 3CuSO 4 + 2PH3 → Cu3 P2 + 3H 2 SO 4 3HgCl2 + 2PH 3 → Hg 3 P2 + 6HCl Phosphine is weakly basic and like ammonia, gives phosphonium compounds with acids e.g., PH3 + HBr → PH 4 Br Uses Uses: The spontaneous combustion of phosphine is technically used in Holme’s signals. Containers containing calcium carbide and calcium phosphide are pierced and thrown in the sea when the gases evolved burn and serve as a signal. It is also used in smoke screens. In what way can it be proved that PH3 is basic in nature? Example 7.6 PH3 reacts with acids like HI to form PH4I which shows that it is Solution basic in nature. PH3 + HI → PH 4 I Due to lone pair on phosphorus atom, PH3 is acting as a Lewis base in the above reaction. Intext Questions + 7.7 (a) Bond angle in PH4 is higher than that in PH3. Why? (b) What is formed when PH3 reacts with an acid? 7.8 What happens when white phosphorus is heated with concentrated NaOH solution in an inert atmosphere of CO2 ? 7.8 Phosphorus Halides Phosphorus forms two types of halides, PX3 (X = F, Cl, Br, I) and PX5 (X = F, Cl, Br). 7.8.1 Phosphorus Preparation Trichloride It is obtained by passing dry chlorine over heated white phosphorus. P4 + 6Cl2 → 4PCl3 It is also obtained by the action of thionyl chloride with white phosphorus. P4 + 8SOCl 2 → 4PCl3 + 4SO2 + 2S2 Cl2 Properties It is a colourless oily liquid and hydrolyses in the presence of moisture. PCl3 + 3H2O →H3 PO3 + 3HCl P It reacts with organic compounds containing –OH group such as CH3COOH, C2H5OH. 3CH3COOH + PCl 3 → 3CH3 COCl + H3 PO3 Cl Cl Cl 3C2 H5 OH + PCl 3 → 3C2 H 5Cl + H 3 PO 3 3 7.8.2 Phosphorus It has a pyramidal shape as shown, in which phosphorus is sp Pentachloride hybridised. Chemistry 182 2022-23  Preparation Phosphorus pentachloride is prepared by the reaction of white phosphorus with excess of dry chlorine. P4 + 10Cl2 → 4PCl5 It can also be prepared by the action of SO2Cl2 on phosphorus. P4 + 10SO2 Cl 2 → 4PCl5 + 10SO 2 Properties PCl5 is a yellowish white powder and in moist air, it hydrolyses to POCl3 and finally gets converted to phosphoric acid. PCl5 + H2O → POCl3 + 2HCl POCl 3 + 3H2 O → H3 PO4 + 3HCl When heated, it sublimes but decomposes on stronger heating. Heat PCl 5  → PCl 3 + Cl 2 It reacts with organic compounds containing –OH group converting them to chloro derivatives. C2 H5OH + PCl5 → C2 H5Cl + POCl 3 + HCl CH3COOH + PCl5 → CH3COCl + POCl3 +HCl Finely divided metals on heating with PCl5 give corresponding Cl chlorides. Cl 2Ag + PCl5 → 2AgCl + PCl3 240 pm Sn + 2PCl5 → SnCl 4 + 2PCl 3 It is used in the synthesis of some organic 202 P pm compounds, e.g., C2H5Cl, CH3COCl. In gaseous and liquid phases, it has a trigonal Cl Cl bipyramidal structure as shown. The three equatorial P–Cl bonds are equivalent, while the two axial bonds are longer than equatorial bonds. This is due to the fact that Cl the axial bond pairs suffer more repulsion as compared to equatorial bond pairs. Example 7.7 Why does PCl3 fume in moisture ? Solution PCl3 hydrolyses in the presence of moisture giving fumes of HCl. PCl3 + 3H2O →H3 PO3 + 3HCl Example 7.8 Are all the five bonds in PCl5 molecule equivalent? Justify your answer. Solution PCl5 has a trigonal bipyramidal structure and the three equatorial P-Cl bonds are equivalent, while the two axial bonds are different and longer than equatorial bonds. 183 The p-Block Elements 2022-23  Intext Questions 7.9 What happens when PCl5 is heated? 7.10 Write a balanced equation for the reaction of PCl5 with water. 7.9 Oxoacids of Phosphorus forms a number of oxoacids. The important oxoacids of Phosphorus phosphorus with their formulas, methods of preparation and the presence of some characteristic bonds in their structures are given in Table 7.5. Table 7.5: Oxoacids of Phosphorus Name Formula Oxidation Characteristic Preparation state of bonds and their phosphorus number Hypophosphorous H3PO2 +1 One P – OH white P4 + alkali (Phosphinic) Two P – H One P = O Orthophosphorous H3PO3 +3 Two P – OH P2O3 + H2O (Phosphonic) One P – H One P = O Pyrophosphorous H4P2O5 +3 Four P – OH PCl3 + H3PO3 One P – O – P Hypophosphoric H4P2O6 +4 Four P – OH red P4 + alkali Two P = O One P – P Orthophosphoric H3PO4 +5 Three P – OH P4O10+H2O One P = O Pyrophosphoric H4P2O7 +5 Four P – OH heat phosphoric Two P = O acid One P – O – P Metaphosphoric* (HPO3)n +5 Three P – OH phosphorus acid Three P = O + Br2, heat in a Three P – O – P sealed tube * Exists in polymeric forms only. Characteristic bonds of (HPO3)3 have been given in the Table. The compositions of the oxoacids are interrelated in terms of loss or gain of H2O molecule or O-atom. The structures of some important oxoacids are given next. In oxoacids phosphorus is tetrahedrally surrounded by other atoms. All these acids contain at least one P=O bond and one P–OH bond. The oxoacids in which phosphorus has lower oxidation state (less than +5) contain, in addition to P=O and P–OH bonds, either P–P (e.g., in H4P2O6) or P–H (e.g., in H3PO2) bonds but not both. These acids in +3 oxidation state of phosphorus tend to disproportionate to higher and lower oxidation states. For example, orthophophorous acid (or phosphorous acid) on heating disproportionates to give orthophosphoric acid (or phosphoric acid) and phosphine. 4H3 PO3 → 3H3 PO4 + PH3 Chemistry 184 2022-23  O O O O O P P P P P HO OH HO O H OH H OH OH OH OH OH OH H H3PO4 H4P2O7 H3PO3 H3PO2 Orthophosphoric acid Pyrophosphoric acid Orthophosphorous acid Hypophosphorous acid O O O P P O OH O P HO P OH HO P P P O OH O O O O O O OH OH P H4P2O5 OH O OH Pyrophosphorous O OH Cyclotrimetaphosphoric acid, (HPO3)3 Polymetaphosphoric acid, (HPO3)n Fig. 7.4 Structures of some important oxoacids of phosphorus The acids which contain P–H bond have strong reducing properties. Thus, hypophosphorous acid is a good reducing agent as it contains two P–H bonds and reduces, for example, AgNO3 to metallic silver. 4 AgNO3 + 2H2O + H3PO2 → 4Ag + 4HNO3 + H3PO4 + These P–H bonds are not ionisable to give H and do not play any role in basicity. Only those H atoms which are attached with oxygen in P–OH form are ionisable and cause the basicity. Thus, H3PO3 and H3PO4 are dibasic and tribasic, respectively as the structure of H3PO3 has two P–OH bonds and H3PO4 three. Example 7.9 How do you account for the reducing behaviour of H3PO2 on the basis of its structure ? Solution In H3PO2, two H atoms are bonded directly to P atom which imparts reducing character to the acid. Intext Questions 7.11 What is the basicity of H3PO4? 7.12 What happens when H3PO3 is heated? 7.10 Group 16 Oxygen, sulphur, selenium, tellurium, polonium and livermorium Elements 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. 185 The p-Block Elements 2022-23 7.10.1 Occurrence Oxygen is the most abundant of all the elements on 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, baryte BaSO4 and sulphides such as galena PbS, zinc blende ZnS, copper pyrites CuFeS2. Traces of sulphur occur as hydrogen sulphide in volcanoes. Organic materials such as eggs, proteins, garlic, onion, mustard, hair and wool contain sulphur. 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. Livermorium is a synthetic radioactive element. Its symbol is Lv, atomic number 116, atomic mass 14 10 2 4 292 and electronic configuration [Rn] 5f 6d 7s 7p. It has been produced only in a very small amount and has very short half-life (only a small fraction of one second). This limits the study of properlies of Lv. Here, except for livermorium, important atomic and physical properties of other elements of Group16 along with their electronic configurations are given in Table 7.6. Some of the atomic, physical and chemical properties and their trends are discussed below. Table 7.6: Some Physical Properties of Group 16 Elements Property O S Se Te Po Atomic number 8 16 34 52 84 –1 Atomic mass/g mol 16.00 32.06 78.96 127.60 210.00 2 4 2 4 10 2 4 10 2 4 14 10 2 4 Electronic configuration [He]2s 2p [Ne]3s 3p [Ar]3d 4s 4p [Kr]4d 5s 5p [Xe]4f 5d 6s 6p a Covalent radius/(pm) 66 104 117 137 146 2– b Ionic radius, E /pm 140 184 198 221 230 Electron gain enthalpy, –141 –200 –195 –190 –174 –1 /∆egH kJ mol Ionisation enthalpy (∆iH1) 1314 1000 941 869 813 –1 /kJ mol Electronegativity 3.50 2.58 2.55 2.01 1.76 –3 c d e Density /g cm (298 K) 1.32 2.06 4.19 6.25 – f Melting point/K 55 393 490 725 520 Boiling point/K 90 718 958 1260 1235 Oxidation states* –2,–1,1,2 –2,2,4,6 –2,2,4,6 –2,2,4,6 2,4 a b c d e f Single bond; Approximate value; At the melting point; Rhombic sulphur; Hexagonal grey; Monoclinic form, 673 K. * Oxygen shows oxidation states of +2 and +1 in oxygen fluorides OF2 and O2F2 respectively. 7.10.2 Electronic The elements of Group16 have six electrons in the outermost shell and 2 4 Configuration have ns np general electronic configuration. 7.10.3 Atomic Due to increase in the number of shells, atomic and ionic radii increase and Ionic from top to bottom in the group. The size of oxygen atom is, however, Radii exceptionally small. Chemistry 186 2022-23 7.10.4 Ionisation Ionisation enthalpy decreases down the group. It is due to increase in Enthalpy size. However, the elements of this group have lower ionisation enthalpy values compared to those of Group15 in the corresponding periods. This is due to the fact that Group 15 elements have extra stable half- filled p orbitals electronic configurations. 7.10.5 Electron Because of the compact nature of oxygen atom, it has less negative Gain electron gain enthalpy than sulphur. However, from sulphur onwards Enthalpy the value again becomes less negative upto polonium. 7.10.6 Next to fluorine, oxygen has the highest electronegativity value amongst Electronegativity the elements. Within the group, electronegativity decreases with an increase in atomic number. This implies that the metallic character increases from oxygen to polonium. Elements of Group 16 generally show lower value of first ionisation Example 7.10 enthalpy compared to the corresponding periods of group 15. Why? Due to extra stable half-filled p orbitals electronic configurations of Solution Group 15 elements, larger amount of energy is required to remove electrons compared to Group 16 elements. 7.10.7 Physical Some of the physical properties of Group 16 elements are given in Properties Table 7.6. Oxygen and sulphur are non-metals, selenium and tellurium metalloids, whereas polonium is a metal. Polonium is radioactive and is short lived (Half-life 13.8 days). All these elements exhibit allotropy. The melting and boiling points increase with an increase in atomic number down the group. The large difference between the melting and boiling points of oxygen and sulphur may be explained on the basis of their atomicity; oxygen exists as diatomic molecule (O2) whereas sulphur exists as polyatomic molecule (S8). 7.10.8 Chemical Oxidation states and trends in chemical reactivity Properties The elements of Group 16 exhibit a number of oxidation states (Table 7.6). The stability of -2 oxidation state decreases down the group. Polonium hardly shows –2 oxidation state. Since electronegativity of oxygen is very high, it shows only negative oxidation state as –2 except in the case of OF2 where its oxidation state 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 state in their compounds with oxygen and + 6 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 is primarily covalent. Anomalous behaviour of oxygen The anomalous behaviour of oxygen, like other members 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. 187 The p-Block Elements 2022-23  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 shells can be expanded and covalence exceeds four. (i) Reactivity with hydrogen: All the elements of Group 16 form hydrides of the type H2E (E = O, S, Se, Te, Po). Some properties of hydrides are given in Table 7.7. Their acidic character increases from H2O to H2Te. The increase in acidic character can be explained in terms of decrease in bond enthalpy for the dissociation of H–E bond down the group. Owing to the decrease in enthalpy for the dissociation of H–E bond 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 7.7: Properties of Hydrides of Group 16 Elements Property H 2O H 2S H 2Se H 2 Te m.p/K 273 188 208 222 b.p/K 373 213 232 269 H–E distance/pm 96 134 146 169 HEH angle (°) 104 92 91 90 –1 ∆f H/kJ mol –286 –20 73 100 –1 ∆diss H (H–E)/kJ mol 463 347 276 238 a –16 –7 –4 –3 Dissociation constant 1.8×10 1.3×10 1.3×10 2.3×10 a Aqueous solution, 298 K (ii) Reactivity with oxygen: All these elements form oxides of the EO2 and EO3 types where E = S, Se, Te or Po. Ozone (O3) and sulphur dioxide (SO2) are gases while selenium dioxide (SeO2) is solid. Reducing property of dioxide decreases from SO2 to TeO2; SO2 is reducing while TeO2 is an oxidising agent. Besides EO2 type, sulphur, selenium and tellurium also form EO3 type oxides (SO3, SeO3, TeO3). Both types of oxides are acidic in nature. (iii) Reactivity towards the halogens: Elements of Group 16 form a large number of halides of the type, EX6, EX4 and EX2 where E is an element of the group and X is a halogen. The stability of the – – – – halides decreases in the order F > Cl > Br > I. Amongst hexahalides, hexafluorides are the only stable halides. All hexafluorides are gaseous in nature. They have octahedral structure. Sulphur hexafluoride, SF6 is exceptionally stable for steric reasons. Amongst tetrafluorides, SF4 is a gas, SeF4 a liquid and TeF4 a solid. 3 These fluorides have sp d hybridisation and thus, have trigonal bipyramidal structures in which one of the equatorial positions is occupied by a lone pair of electrons. This geometry is also regarded as see-saw geometry. All elements except oxygen form dichlorides and dibromides. These 3 dihalides are formed by sp hybridisation and thus, have tetrahedral structure. The well known monohalides are dimeric in nature. Examples Chemistry 188 2022-23  are S2F2, S2Cl2, S2Br2, Se2Cl2 and Se2Br2. These dimeric halides undergo disproportionation as given below: 2Se2Cl2 → SeCl4 + 3Se H2S is less acidic than H2Te. Why? Example 7.11 Due to the decrease in bond (E–H) dissociation Solution enthalpy down the group, acidic character increases. Intext Questions 7.13 List the important sources of sulphur. 7.14 Write the order of thermal stability of the hydrides of Group 16 elements. 7.15 Why is H2O a liquid and H2S a gas ? 7.11 Dioxygen Preparation Dioxygen can be obtained in the laboratory by the following ways: (i) By heating oxygen containing salts such as chlorates, nitrates and permanganates. 2KClO3  Heat MnO → 2KCl + 3O2 2 (ii) By the thermal decomposition of the oxides of metals low in the electrochemical series and higher oxides of some metals. 2Ag2O(s) → 4Ag(s) + O2(g); 2Pb3O4(s) → 6PbO(s) + O2(g) 2HgO(s) → 2Hg(l) + O2(g) ; 2PbO2(s) → 2PbO(s) + O2(g) (iii) Hydrogen peroxide is readily decomposed into water and dioxygen by catalysts such as finely divided metals and manganese dioxide. 2H2O2(aq) → 2H2O(1) + O2(g) On large scale it can be prepared from water or air. Electrolysis of water leads to the release of hydrogen at the cathode and oxygen at the anode. Industrially, dioxygen is obtained from air by first removing carbon dioxide and water vapour and then, the remaining gases are liquefied and fractionally distilled to give dinitrogen and dioxygen. Properties Dioxygen is a colourless and odourless gas. Its solubility in water is to 3 3 the extent of 3.08 cm in 100 cm water at 293 K which is just sufficient for the vital support of marine and aquatic life. It liquefies at 90 K and 16 17 freezes at 55 K. Oxygen atom has three stable isotopes: O, O and 18 O. Molecular oxygen, O2 is unique in being paramagnetic inspite of having even number of electrons (see Class XI Chemistry Book, Unit 4). Dioxygen directly reacts with nearly all metals and non-metals except some metals ( e.g., Au, Pt) and some noble gases. Its combination with other elements is often strongly exothermic which helps in sustaining the reaction. However, to initiate the reaction, some external 189 The p-Block Elements 2022-23  heating is required as bond dissociation enthalpy of oxgyen-oxygen –1 double bond is high (493.4 kJ mol ). Some of the reactions of dioxygen with metals, non-metals and other compounds are given below: 2Ca + O 2 → 2CaO 4Al + 3O 2 → 2Al 2O 3 P4 + 5O2 → P4 O10 C + O 2 → CO 2 2ZnS + 3O2 → 2ZnO + 2SO2 CH 4 + 2O2 → CO2 + 2H2O Some compounds are catalytically oxidised. For example, V2 O5 2SO2 + O2   → 2SO3 CuCl2 4HCl + O2  → 2Cl 2 + 2H2O Uses Uses: In addition to its importance in normal respiration and combustion processes, oxygen is used in oxyacetylene welding, in the manufacture of many metals, particularly steel. Oxygen cylinders are widely used in hospitals, high altitude flying and in mountaineering. The combustion of fuels, e.g., hydrazines in liquid oxygen, provides tremendous thrust in rockets. Intext Questions 7.16 Which of the following does not react with oxygen directly? Zn, Ti, Pt, Fe 7.17 Complete the following reactions: (i) C2H4 + O2 → (ii) 4Al + 3 O2 → 7.12 Simple A binary compound of oxygen with another element is called oxide. As already stated, oxygen reacts with most of the elements of the periodic Oxides table to form oxides. In many cases one element forms two or more oxides. The oxides vary widely in their nature and properties. Oxides can be simple (e.g., MgO, Al2O3 ) or mixed (Pb3O4, Fe3O4). Simple oxides can be classified on the basis of their acidic, basic or amphoteric character. An oxide that combines with water to give an acid is termed acidic oxide (e.g., SO2, Cl2O7, CO2, N2O5 ). For example, SO2 combines with water to give H2SO3, an acid. SO2 + H2 O → H2 SO3 As a general rule, only non-metal oxides are acidic but oxides of some metals in high oxidation state also have acidic character (e.g., Mn2O7, CrO3, V2O5). The oxides which give a base with water are known as basic oxides (e.g., Na2O, CaO, BaO). For example, CaO combines with water to give Ca(OH)2, a base. CaO + H2 O → Ca (OH )2 Chemistry 190 2022-23  In general, metallic oxides are basic. Some metallic oxides exhibit a dual behaviour. They show characteristics of both acidic as well as basic oxides. Such oxides are known as amphoteric oxides. They react with acids as well as alkalies. For example, Al2O3 reacts with acids as well as alkalies. Al 2O3 (s ) + 6HCl (aq ) + 9H2 O ( l ) → 2 [ Al(H2O)6 ] 3+ (aq ) + 6Cl − ( aq ) Al 2O3 (s ) + 6NaOH ( aq ) + 3H2O ( l ) → 2Na 3 [ Al (OH )6 ] (aq ) There are some oxides which are neither acidic nor basic. Such oxides are known as neutral oxides. Examples of neutral oxides are CO, NO and N2O. 7.13 Ozone Ozone is an allotropic form of oxygen. It is too reactive to remain for long in the atmosphere at sea level. At a height of about 20 kilometres, it is formed from atmospheric oxygen in the presence of sunlight. This ozone layer protects the earth’s surface from an excessive concentration of ultraviolet (UV) radiations. Preparation When a slow dry stream of oxygen is passed through a silent electrical discharge, conversion of oxygen to ozone (10%) occurs. The product is known as ozonised oxygen. V –1 3O2 → 2O3 ∆H (298 K) = +142 kJ mol Since the formation of ozone from oxygen is an endothermic process, it is necessary to use a silent electrical discharge in its preparation to prevent its decomposition. If concentrations of ozone greater than 10 per cent are required, a battery of ozonisers can be used, and pure ozone (b.p. 101.1K) can be condensed in a vessel surrounded by liquid oxygen. Properties Pure ozone is a pale blue gas, dark blue liquid and violet-black solid. Ozone has a characteristic smell and in small concentrations it is harmless. However, if the concentration rises above about 100 parts per million, breathing becomes uncomfortable resulting in headache and nausea. Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen results in the liberation of heat (∆H is negative) and an increase in entropy (∆S is positive). These two effects reinforce each other, resulting in large negative Gibbs energy change (∆G) for its conversion into oxygen. It is not really surprising, therefore, high concentrations of ozone can be dangerously explosive. Due to the ease with which it liberates atoms of nascent oxygen (O3 → O2 + O), it acts as a powerful oxidising agent. For example, it oxidises lead sulphide to lead sulphate and iodide ions to iodine. PbS(s) + 4O3(g) → PbSO4(s) + 4O2(g) 2I–(aq) + H2O(l) + O3(g) →2OH–(aq) + I2(s) + O2(g) When ozone reacts with an excess of potassium iodide solution buffered with a borate buffer (pH 9.2), iodine is liberated which can be titrated against a standard solution of sodium thiosulphate. This is a quantitative method for estimating O3 gas. 191 The p-Block Elements 2022-23  Experiments have shown that nitrogen oxides (particularly nitrogen monoxide) combine very rapidly with ozone and there is, thus, the possibility that nitrogen oxides emitted from the exhaust systems of supersonic jet aeroplanes might be slowly depleting the concentration of the ozone layer in the upper atmosphere. NO ( g ) + O3 ( g ) → NO2 ( g ) + O2 ( g ) Another threat to this ozone layer is probably posed by the use of freons which are used in aerosol sprays and as refrigerants. The two oxygen-oxygen bond lengths in the ozone molecule are identical (128 pm) and the molecule is angular o as expected with a bond angle of about 117. It is a resonance hybrid of two main forms: Uses Uses: It is used as a germicide, disinfectant and for sterilising water. It is also used for bleaching oils, ivory, flour, starch, etc. It acts as an oxidising agent in the manufacture of potassium permanganate. Intext Questions 7.18 Why does O3 act as a powerful oxidising agent? 7.19 How is O3 estimated quantitatively? 7.14 Sulphur — Sulphur forms numerous allotropes of which the yellow rhombic Allotropic (α-sulphur) and monoclinic (β -sulphur) forms are the most important. The stable form at room temperature is rhombic sulphur, which Forms transforms to monoclinic sulphur when heated above 369 K. α-sulphur) Rhombic sulphur (α 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 dissolves to some extent in benzene, alcohol and ether. It is readily soluble in CS2. β-sulphur) Monoclinic sulphur (β 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 β-sulphur are formed. It is stable above 369 K and transforms into α-sulphur below it. Conversely, α-sulphur is stable below 369 K and transforms into β-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 Fig. 7.5(a). Chemistry 192 2022-23  Several other modifications of sulphur containing 6-20 sulphur atoms per ring have been synthesised in the last two decades. In cyclo-S6, the ring adopts the chair form and the molecular dimensions are as shown in Fig. 7.5 (b). (a) (b) At elevated temperatures (~1000 K), S2 is the dominant Fig. 7.5: The structures of (a) S8 ring in species and is paramagnetic rhombic sulphur and (b) S6 form like O2. Which form of sulphur shows paramagnetic behaviour ? Example 7.12 In vapour state sulphur partly exists as S2 molecule which has two Solution unpaired electrons in the antibonding π * orbitals like O2 and, hence, exhibits paramagnetism. 7.15 Sulphur Preparation Dioxide Sulphur dioxide is formed together with a little (6-8%) sulphur trioxide when sulphur is burnt in air or oxygen: S(s) + O2(g) → SO2 (g) In the laboratory it is readily generated by treating a sulphite with dilute sulphuric acid. SO32-(aq) + 2H+ (aq) → H2O(l) + SO2 (g) Industrially, it is produced as a by-product of the roasting of sulphide ores. 4FeS2 (s ) + 11O2 ( g ) → 2Fe2 O3 ( s ) + 8SO2 ( g ) The gas after drying is liquefied under pressure and stored in steel cylinders. Properties Sulphur dioxide is a colourless gas with pungent smell and is highly soluble in water. It liquefies at room temperature under a pressure of two atmospheres and boils at 263 K. Sulphur dioxide, when passed through water, forms a solution of sulphurous acid. SO2 ( g ) + H2 O ( l ) ƒ H2 SO3 ( aq ) It reacts readily with sodium hydroxide solution, forming sodium sulphite, which then reacts with more sulphur dioxide to form sodium hydrogen sulphite. 2NaOH + SO2 → Na2SO3 + H2O Na2SO3 + H2O + SO2 → 2NaHSO3 In its reaction with water and alkalies, the behaviour of sulphur dioxide is very similar to that of carbon dioxide. 193 The p-Block Elements 2022-23  Sulphur dioxide reacts with chlorine in the presence of charcoal (which acts as a catalyst) to give sulphuryl chloride, SO2Cl2. It is oxidised to sulphur trioxide by oxygen in the presence of vanadium(V) oxide catalyst. SO2(g) + Cl2 (g) → SO2Cl2(l) 2SO2 ( g ) + O2 ( g )  V2O5 → 2SO3 ( g )  When moist, sulphur dioxide behaves as a reducing agent. For example, it converts iron(III) ions to iron(II) ions and decolourises acidified potassium permanganate(VII) solution; the latter reaction is a convenient test for the gas. 2Fe3+ + SO2 + 2H2 O → 2Fe 2+ + SO24− + 4H + 5SO2 + 2MnO−4 + 2H2 O → 5SO42− + 4H+ + 2Mn2 + The molecule of SO2 is angular. It is a resonance hybrid of the two canonical forms: Uses Uses: Sulphur dioxide is used (i) in refining petroleum and sugar (ii) in bleaching wool and silk and (iii) as an anti-chlor, disinfectant and preservative. Sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite (industrial chemicals) are manufactured from sulphur dioxide. Liquid SO2 is used as a solvent to dissolve a number of organic and inorganic chemicals. Intext Questions 7.20 What happens when sulphur dioxide is passed through an aqueous solution of Fe(III) salt? 7.21 Comment on the nature of two S–O bonds formed in SO2 molecule. Are the two S–O bonds in this molecule equal ? 7.22 How is the presence of SO2 detected ? 7.16 Oxoacids of Sulphur forms a number of oxoacids such as H2SO3, H2S2O3, H2S2O4, Sulphur H2S2O5, H2SxO6 (x = 2 to 5), H2SO4, H2S2O7, H2SO5, H2S2O8. Some of these acids are unstable and cannot be isolated. They are known in aqueous solution or in the form of their salts. Structures of some important oxoacids are shown in Fig. 7.6. Fig. 7.6: Structures of some important oxoacids of sulphur Chemistry 194 2022-23 7.17 Sulphuric Manufacture Sulphuric acid is one of the most important industrial chemicals Acid worldwide. Sulphuric acid is manufactured by the Contact Process which involves three steps: (i) burning of sulphur or sulphide ores in air to generate SO2. (ii) conversion of SO2 to SO3 by the reaction with oxygen in the presence of a catalyst (V2O5), and (iii) absorption of SO3 in H2SO4 to give Oleum (H2S2O7). A flow diagram for the manufacture of sulphuric acid is shown in (Fig. 7.7). The SO2 produced is purified by removing dust and other impurities such as arsenic compounds. The key step in the manufacture of H2SO4 is the catalytic oxidation of SO2 with O2 to give SO3 in the presence of V2O5 (catalyst). 2SO2 ( g ) + O2 ( g )  V2O5 → 2SO3 ( g ) ∆ r H 0 = −196.6 kJmol −1  The reaction is exothermic, reversible and the forward reaction leads to a decrease in volume. Therefore, low temperature and high pressure are the favourable conditions for maximum yield. But the temperature should not be very low otherwise rate of reaction will become slow. In practice, the plant is operated at a pressure of 2 bar and a temperature of 720 K. The SO3 gas from the catalytic converter is absorbed in concentrated H2SO4 to produce oleum. Dilution of oleum with water gives H2SO4 of the desired concentration. In the industry two steps are carried out simultaneously to make the process a continuous one and also to reduce the cost. SO3 + H2SO4 → H2S2O7 (Oleum) The sulphuric acid obtained by Contact process is 96-98% pure. Water Conc. H2SO4 Impure spray spray Conc. H2SO4 SO2+O2 Dry SO2+O2 SO3 V2O5 Quartz Sulphur Preheater Air Waste Waste Catalytic Sulphur converter Oleum burner water acid (H2S2O7) Arsenic purifier Washing and Drying containing Dust cooling tower tower gelatinous hydrated precipitator ferric oxide Fig. 7.7: Flow diagram for the manufacture of sulphuric acid 195 The p-Block Elements 2022-23  Properties Sulphuric acid is a colourless, dense, oily liquid with a specific gravity of 1.84 at 298 K. The acid freezes at 283 K and boils at 611 K. It dissolves in water with the evolution of a large quantity of heat. Hence, care must be taken while preparing sulphuric acid solution from concentrated sulphuric acid. The concentrated acid must be added slowly into water with constant stirring. The chemical reactions of sulphuric acid are as a result of the following characteristics: (a) low volatility (b) strong acidic character (c) strong affinity for water and (d) ability to act as an oxidising agent. In aqueous solution, sulphuric acid ionises in two steps. + – H2SO4(aq) + H2O(l) → H3O (aq) + HSO4 (aq); K a1 = very large ( K a1 >10) – + 2- –2 HSO4 (aq) + H2O(l) → H3O (aq) + SO4 (aq) ; K a2 = 1.2 × 10 The larger value of K a1 ( K a1 >10) means that H2SO4 is largely + – dissociated into H and HSO4. Greater the value of dissociation constant (Ka), the stronger is the acid. The acid forms two series of salts: normal sulphates (such as sodium sulphate and copper sulphate) and acid sulphates (e.g., sodium hydrogen sulphate). Sulphuric acid, because of its low volatility can be used to manufacture more volatile acids from their corresponding salts. 2 MX + H2SO4 → 2 HX + M2SO4 (X = F, Cl, NO3) (M = Metal) Concentrated sulphuric acid is a strong dehydrating agent. Many wet gases can be dried by passing them through sulphuric acid, provided the gases do not react with the acid. Sulphuric acid removes water from organic compounds; it is evident by its charring action on carbohydrates. C12H22O11  H2 SO4 → 12C + 11H2O Hot concentrated sulphuric acid is a moderately strong oxidising agent. In this respect, it is intermediate between phosphoric and nitric acids. Both metals and non-metals are oxidised by concentrated sulphuric acid, which is reduced to SO2. Cu + 2 H2SO4(conc.) → CuSO4 + SO2 + 2H2O S + 2H2SO4(conc.) → 3SO2 + 2H2O C + 2H2SO4(conc.) → CO2 + 2 SO2 + 2 H2O Uses Uses: Sulphuric acid is a very important industrial chemical. A nation’s industrial strength can be judged by the quantity of sulphuric acid it produces and consumes. It is needed for the manufacture of hundreds of other compounds and also in many industrial processes. The bulk of sulphuric acid produced is used in the manufacture of fertilisers (e.g., ammonium sulphate, superphosphate). Other uses are in: (a) petroleum refining (b) manufacture of pigments, paints and dyestuff intermediates (c) detergent industry (d) metallurgical applications (e.g., cleansing metals before enameling, electroplating and galvanising (e) storage batteries (f) in the manufacture of nitrocellulose products and (g) as a laboratory reagent. Chemistry 196 2022-23  What happens when Example 7.13 (i) Concentrated H2SO4 is added to calcium fluoride (ii) SO3 is passed through water? (i) It forms hydrogen fluoride Solution CaF2 + H 2 SO 4 → CaSO 4 + 2HF (ii) It dissolves SO3 to give H2SO4. SO3 + H2O → H2 SO4 Intext Questions 7.23 Mention three areas in which H2SO4 plays an important role. 7.24 Write the conditions to maximise the yield of H2SO4 by Contact process. 7.25 Why is Ka Br – Br > I – I. A reason for this anomaly is the relatively large electron-electron repulsion among the lone pairs in F2 molecule where they are much closer to each other than in case of Cl2. Example 7.15 Although electron gain enthalpy of fluorine is less negative as compared to chlorine, fluorine is a stronger oxidising agent than chlorine. Why? Solution It is due to (i) low enthalpy of dissociation of F-F bond (Table 7.8). – (ii) high hydration enthalpy of F (Table 7.8). 7.18.8 Chemical Oxidation states and trends in chemical reactivity Properties All the halogens exhibit –1 oxidation state. However, chlorine, bromine and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states also as explained below: Halogen atom ns np nd in ground state 1 unpaired electron accounts (other than fluorine) for –1 or +1 oxidation states 3 unpaired electrons account First excited state for +3 oxidation states