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Nitrogen and its Compound
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  1. Chapter 43 Nitrogen and its Compound 43.1Introduction 43.2Unreactive Nature of Nitrogen 43.3Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides 43.4Ammonia 43.5Nitric(V) Acid 43.6Nitrates(V)

  2. 43.1 Introduction (SB p.110) • Nitrogen (first member of Group VA): • Electronic configuration: 1s22s22p3 • Complete octet by forming diatomic molecules N  N • Non-metal, colourless and odourless gas • Very low melting and boiling points • Slightly soluble in water and does not support combustion Some information about nitrogen

  3. 43.1 Introduction (SB p.110) • Nitrogen • Mainly as free N2 molecules in the atmosphere (78% by volume) • Combine with other elements in the form of proteins in all living things • Liquid N2 is used as coolant • Raw material for Haber process(manufacture of ammonia) • Ammonia is the major component of nitrogenous fertilizers

  4. 43.2 Unreactive Nature of Nitrogen (SB p.111) • Nitrogen in gaseous state • As diatomic molecules (N2) which are held by weak van der Waals’ forces • 2 atoms are joined by extremely strong triple covalent bonds • Bond enthalpy of the triple bond = +944 kJ mol–1 • Due to extremely strong covalent bonds and absence of bond polarity •  Nitrogen molecule is very unreactive

  5. 43.2 Unreactive Nature of Nitrogen (SB p.111) Bond enthalpies of some common covalent bonds

  6. Reactions involving nitrogen usually have high activation energies and unfavourable equilibrium constants • e.g. At 25°C • N2(g) + O2(g) 2NO(g) Kc = 4.5  10–31 • The presence of catalyst and high temperatureand pressure may be required for nitrogen to react • N2(g) + 3H2(g) 2NH3(g) 400 – 500°C, 300 – 1000 atm Fe as catalyst 43.2 Unreactive Nature of Nitrogen (SB p.111)

  7. N2will not react at room temperature due to high bond enthalpy • At high temperature, N2 shows some reactions with other elements • ∵ sufficient energy to break N  N triple bond • At high temperature, • N2(g) + O2(g) 2NO(g)  43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.112)

  8. The electric discharge in lightning provides sufficient energy to break the N  N triple bond and then react with O2 • N2(g) + O2(g)  2NO(g) • 2NO(g) + O2(g)  2NO2(g) lightning colourless Reddish brown (poisonous) colourless 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.112)

  9. 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.112) • The above reactions are very important in nature • The NO2 formed dissolves in rainwater to produce nitric(V) acid and nitric(III) acid • 2NO2(g) + H2O(l)  HNO3(aq) + HNO2(aq)

  10. The NO2 absorbs sunlight and breaks down into NO and O atom • NO2(g)  NO(g) + O(g) • These leads to formation of photochemical smog sunlight 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.112) • At high temperatures in car engines, N2 & O2 react to form NO(g) which emitted into air with exhausted gas • The NO formed will be oxidized to NO2

  11. 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.113) • In laboratory, we use the apparatus shown on the right to convert N2 into NO2 • When current is switched on, electric discharges occur in the gap between the electrodes • NO is formed and followed by NO2

  12. Other than NO and NO2, N2 can form other oxides • e.g. 2 NO2 molecules (brown) can combine to form a N2O4 molecule (yellow) • NO2 & N2O4 exist in equilibrium in gas phase • 2NO2(g) N2O4(g) H = –58 kJ mol–1 yellow brown 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.113)

  13. NO2 (left) and N2O4 (right) predominate in hot water and ice water respectively 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.113) • The formation of N2O4 is exothermic •  N2O4 predominantes at low temperatures • NO2 predominantes at high temperatures •  the colour of mixture fades on cooling, darkens on warming

  14. (a) (i) Dinitrogen monoxide (ii) Nitrogen monoxide 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.114) Check Point 43-1 (a) Draw the structures of the following compounds. (i) Dinitrogen monoxide (ii) Nitrogen monoxide (iii) Dinitrogen trioxide (iv) Nitrogen dioxide (v) Dinitrogen tetraoxide (vi) Dinitrogen pentaoxide Answer

  15. (iii) Dinitrogen trioxide (iv) Nitrogen dioxide 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.114)

  16. (v) Dinitrogen tetraoxide (vi) Dinitrogen pentaoxide 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.114)

  17. (b) The ascending order of reactivity is: N2 < O2 < F2. The reactivity of diatomic molecules depends on the bond enthalpy of covalent bonds. The bond enthalpy of N  N is greater than that of O = O, which in turn is greater than that of F – F. Therefore, the breakage of N  N bond requires the greatest amount of energy, whereas the breakage of F – F bond requires the least amount of energy. 43.3 Direct Combination of Nitrogen and Oxygen leading to Formation of Nitrogen Oxides (SB p.114) Check Point 43-1 (b) Arrange N2, O2 and F2 in an ascending order of reactivity. Explain the order briefly. Answer

  18. Ammonia • colourless, pungent gas • polar molecules • trigonal pyramidal shape with a lone pair of electrons on nitrogen • extremely soluble in water and easy to condense to liquid due to hydrogen bonds • good solvent for ionic compounds • weakly alkaline • NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) • Kb = 1.8  10–5 mol dm–3 43.4 Ammonia (SB p.114)

  19. 43.4 Ammonia (SB p.114) • Ammonia • one of the most fundamental raw materials for modern industries • important source of fertilizers and 85% of ammonia is used to make nitrogenous fertilizers (e.g. (NH4)2SO4, NH4NO3) • making fibres and plastics (rayon, nylon) • making nitric(V) acid (used to make fertilizers, dyes) • making household cleaners • making detergents

  20. 43.4 Ammonia (SB p.115) Percentages of ammonia used in different industries

  21. NH3 is manufactured industrially by the Haber Process, named after the German chemist Fritz Haber • The process involves direct combination of N2 and H2 under special conditions • N2(g) + H2(g) 2NH3(g) H = –92 kJ mol–1 Fritz Haber (1868 – 1934) 43.4 Ammonia (SB p.115) Manufacture of Ammonia by the Haber Process

  22. 43.4 Ammonia (SB p.116) Flow diagram for the Haber process

  23. Process of Haber Process: • N2 is obtained from fractional distillation of liquid air • H2 is obtained from methane, naphtha or mixture by steam reforming • CH4(g) + H2O(g)  CO(g) + 3H2(g) • CH4(g) + air  CO(g) + 2H2(g) + N2(g) Ni 900°C Ni 900°C 43.4 Ammonia (SB p.116)

  24. 43.4 Ammonia (SB p.116) • Mixture of CO & H2O is mixed with steam and passed over a heated catalyst • CO(g) + H2O(l)  CO2(g) + H2(g) • The CO2 formed is dissolved in water under pressure • The gases (N2 & H2) are purified before proceeding to the next stage • ∵ Compounds of oxygen and sulphur will poison the catalyst

  25. 43.4 Ammonia (SB p.116) • Purified N2 & H2 are mixed in ratio of 3 : 1 by volume •  Compressed to 200 – 1000 atm and heated in the heat exchanger •  Hot gaseous mixture is passed over iron in the catalytic chamber •  Gases contain 10 – 15% of NH3 and unreacted N2 and H2 when leaving the chamber •  The gases are cooled after passing through the heat exchanger •  NH3 is liquefied under pressure and unreacted gases are recycled

  26. Synthesis of ammonia is an exothermic and reversible reaction • N2(g) + H2(g) 2NH3(g) H = –92 kJ mol–1 43.4 Ammonia (SB p.117) Physico-chemical principles: • According to Le Chatelier’s principle, • (1) high pressure will increase the yield • (2) low temperature will increase the yield

  27. 43.4 Ammonia (SB p.117) • Apart from increasing yield, the reaction rate should be fast •  Low temperatures would lower the rate of reaction • ∴ optimum temperature is around 500°C which is high enough for reaction to proceed quickly but low enough to give satisfactory yield • Catalyst is used to increase the reaction rate •  poisoned by CO, CO2, H2S •  Gases entering the catalytic chamber should have high purity!!

  28. NH3partly ionizes in water to give NH4+ and OH– ions • ∴ NH3(aq) is alkaline • NH3(g) + H2O(l) NH4+(aq) + OH–(aq) • Kb = 1.8  10–5 mol dm–3 • ∴ NH3(aq) is a weak base 43.4 Ammonia (SB p.118) Chemical Properties of Ammonia As a base

  29. NH3 neutralizes acids to give ammonium salts • e.g. 2NH3(aq) + H2SO4(aq)  (NH4)2SO4(aq) • NH3(aq) + HNO3(aq)  NH4NO3(aq) ammonium sulphate(VI) ammonium nitrate(V) Filter paper soaked with NH3 Filter paper soaked with HCl 43.4 Ammonia (SB p.118) Reaction with Acids • Formation of NH4Cl by reacting NH3 with HCl • NH3(aq) + HCl(aq)  NH4Cl(s)

  30. NH3 precipitates the hydroxide of many metals from solutions of their salts • CaSO4(aq) + 2NH3(aq) +2H2O(l)  Ca(OH)2(s) + (NH4)2SO4(aq) • ZnSO4(aq) + 2NH3(aq) +2H2O(l)  Zn(OH)2(s) + (NH4)2SO4(aq) • Pb(NO3)2(aq) + 2NH3(aq) +2H2O(l)  Pb(OH)2(s) + 2NH4NO3(aq) • CuSO4(aq) + 2NH3(aq) +2H2O(l)  Cu(OH)2(s) + (NH4)2SO4(aq) • FeSO4(aq) + 2NH3(aq) +2H2O(l)  Fe(OH)2(s) + (NH4)2SO4(aq) • Fe2(SO4)3(aq) + 6NH3(aq) +6H2O(l)  2Fe(OH)3(s) + 3(NH4)2SO4(aq) white white white blue dirty green reddish brown 43.4 Ammonia (SB p.118) Reaction with Metal Salts

  31. Pb(OH)2(s) Cu(OH)2(s) Fe(OH)2(s) Fe(OH)3(s) 43.4 Ammonia (SB p.119)

  32. Some metal hydroxides (e.g. Zn(OH)2 & Cu(OH)2) redissolve in excess NH3 solution and form complex compounds • Zn(OH)2(s) + 4NH3(aq)  [Zn(NH3)4]2+(aq) + 2OH–(aq) • Cu(OH)2(s) + 4NH3(aq)  [Cu(NH3)4]2+(aq) + 2OH–(aq) colourless deep blue Cu(OH)2(s) [Cu(NH3)4]2+(aq) A solution containing Cu2+(aq) 43.4 Ammonia (SB p.119)

  33. Silver(I) ions also form a complex with ammonia • AgCl is insoluble in water and acids, but dissolves in excess NH3 forming soluble complex ion [Ag(NH3)2]+(aq) • AgCl(s) Ag+(aq) + Cl–(aq) • Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) Addition of excess NH3(aq) water AgCl(s) dissolves AgCl(s) 43.4 Ammonia (SB p.120)

  34. 4NH3(g) + 3O2(g)  2N2(g) + 6H2O(g)  Laboratory set-up for oxidation of ammonia 43.4 Ammonia (SB p.120) As a Reducing Agent Reaction with Oxygen • NH3does not burn in air or support combustion • It burns in O2 with a yellow flame, forming N2 and water vapour

  35. 4NH3(g) + 5O2(g) •  4NO(g) + 6H2O(g) • This is called catalytic oxidation of ammonia • Key reaction in the preparation of HNO3 Pt  Laboratory set-up for catalytic oxidation of ammonia 43.4 Ammonia (SB p.120) • In the presence of catalyst (red hot spiral coil of platinum at 800 – 900°C), NH3 is oxidized to NO by O2

  36. When dry NH3 is passed over heated black CuO, NH3 is oxidized to N2 and H2O • The CuO turns from black to reddish brown as it is reduced to Cu • 2NH3(g) + 3CuO(s)  3Cu(s) + N2(g) + 3H2O(g)  Laboratory set-up for oxidation of ammonia by copper(II) oxide 43.4 Ammonia (SB p.121) Reaction with Copper(II) Oxide

  37. (a) (i) CH4(g) + H2O(g)  CO(g) + 3H2(g) (ii) C(s) + H2O(l)  CO(g) + H2(g)  43.4 Ammonia (SB p.121) Check Point 43-2 (a) Write chemical equations to show how hydrogen is produced from (i) the reaction of natural gas (mainly methane) with water; (ii) the reaction of coal (mainly carbon) with water. Answer

  38. Check Point 43-2 (cont’d) (b) Consider the following reversible reaction: N2(g) + 3H2(g) 2NH3(g) H = –92 kJ mol–1 Discuss how each of the following factors affects the above equilibrium: (i) increase in temperature (ii) decrease in pressure (iii) addition of a suitable catalyst 43.4 Ammonia (SB p.121) Answer

  39. 43.4 Ammonia (SB p.121) (b) (i) The forward reaction is exothermic. According to Le Chatelier’s principle, exothermic reactions are favoured at low temperatures. Therefore, an increase in temperature will favour the backward reaction, and thus decrease the yield of ammonia. (ii) According to Le Chatelier’s principle, a high pressure will increase the yield of ammonia as the forward reaction is accompanied by a decrease of volume from four to two volumes of the gas. Therefore, a decrease in pressure will decrease the yield of ammonia. (iii) Addition of a suitable catalyst will increase the rate of both forward and backward reactions to the same extent. As it does not change the position of the equilibrium, the yield of ammonia remains constant.

  40. 43.4 Ammonia (SB p.121) Check Point 43-2 (cont’d) (c) Ammonia reacts with oxygen in two different ways. Give equations for both of these reactions and explain how one of them is used industrially to produce nitric(V) acid. Answer

  41. 43.4 Ammonia (SB p.121) (c) In the absence of catalyst, ammonia burns to give molecular nitrogen and water vapour. 4NH3(g) + 3O2(g)  2N2(g) + 6H2O(g) Industrially, in the presence of red hot platinum-rhodium at about 850°C, ammonia is catalytically oxidized to nitrogen monoxide. 4NH3(g) + 5O2(s) 4NO(g) + 6H2O(g) The nitrogen monoxide formed then reacts with oxygen from the air to give nitrogen dioxide. 2NO(g) + O2(g)  2NO2(g) The nitrogen dioxide reacts with excess air and water to produce aqueous nitric(V) acid. 4NO2(g) + O2(g) + 2H2O(l)  4HNO3(aq) H2O(l) + 3NO2(g)  2HNO3(aq) + NO(g) The NO(g) is recycled and subsequently combines with more oxygen and water to give more nitric(V) acid. Finally, the product is distilled to give concentrated nitric(V) acid (containing 68% HNO3). Pt – Rh 850°C

  42. 43.5 Nitric(V) Acid (SB p.121) • Nitric(V) acid • a very strong acid • turns yellow on storage as the formation of dissolved NO2 from decomposition of some acid • 4HNO3(l)  4NO2(aq) + 2H2O(l) +O2(g) • keep in brown bottles as light will speed up decomposition • used to make explosives, nylon, fertilizers and dyestuff synthesis

  43. Most of the ammonia formed is converted to nitric(V) acid by Ostward process • Ostward process is divided into 3 stages: • 1. Mixture of ammonia and excess air is passed over Pt-Rh catalyst at around 700-800°C under low pressure • 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g) Pt-Rh 850°C 43.5 Nitric(V) Acid (SB p.122) Manufacture of Nitric(V) Acid from the Catalytic Oxidation of Ammonia

  44. 43.5 Nitric(V) Acid (SB p.122) 2. The NO formed then reacts with O2 to form NO2 2NO(g) + O2(g)  NO2(g) 3. The NO2 reacts with excess air and water to give aqueous HNO3 4NO2(g) + O2(g) + 2H2O(l)  4HNO3(aq)

  45. 43.5 Nitric(V) Acid (SB p.122) Nitric(V) Acid as an Oxidizing Agent • HNO3 is a strong oxidizing agent, especially when concentrated • NO3– acts as an electron acceptor when H+ ions are present • HNO3 can be reduced to different nitrogen compounds with different oxidation states, depending on • 1. the conc. of HNO3 • 2. nature of substance being oxidized

  46. 43.5 Nitric(V) Acid (SB p.122) • If dilute or moderately concentrated HNO3 is reduced, NO will be formed • 4HNO3(aq) + 3e–  3NO3 –(aq) + 2H2O(l) + NO(g)or NO3–(aq) + 4H+(aq) + 3e–  NO(g) + 2H2O(l) • If concentrated HNO3 is reduced, NO2will be formed • 2HNO3(aq) + e–  NO3 –(aq) + NO2(g) + H2O(l)or NO3–(aq) + 2H+(aq) + e–  NO2(g) + H2O(l) • The electrons are supplied by the reducing agent in the reaction

  47. 43.5 Nitric(V) Acid (SB p.123) Reaction with Copper • Cu reacts with warm dilute HNO3 to give NO3Cu(s) + 8HNO3(aq)  • 3Cu(NO3)2(aq) + 4H2O(l) + 2NO(g) • The NO formed reacts with atmospheric O2 to give NO22NO(g) + O2(g)  2NO2(g)

  48. 43.5 Nitric(V) Acid (SB p.123) • Conc. HNO3 (~14 M) reacts with Cu to give NO2and a blue solution of Cu(NO3)2 • Cu(s) + 4HNO3(aq)  Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g)

  49. 43.5 Nitric(V) Acid (SB p.123) Reaction with Iron(II) Ion • Conc. HNO3 oxdizes green Fe2+ ions to brown Fe3+ ions while itself reduced to NO • 3Fe2+(aq) + NO3–(aq) + 4H+(aq)  3Fe3+(aq) + NO(g) + 2H2O(l) • The NO formed reacts with atmospheric O2 to form NO2 • 2NO(g) + O2(g)  NO2(g)

  50. 43.5 Nitric(V) Acid (SB p.123) Reaction with Sulphur • Hot concentrated HNO3 oxidizes sulphur to give sulphuric(VI) acid and brown fumes of NO2 • S(s) + 6HNO3(aq)  • H2SO4(aq) + 6NO2(g) + 2H2O(l)