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9.6 Nitrogen and Sulphur Learning outcomes: Candidates should be able to: (a)Explain the lack of reactivity of nitrogen (b)Describe: (i)The formation and.

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Presentation on theme: "9.6 Nitrogen and Sulphur Learning outcomes: Candidates should be able to: (a)Explain the lack of reactivity of nitrogen (b)Describe: (i)The formation and."— Presentation transcript:

1 9.6 Nitrogen and Sulphur Learning outcomes: Candidates should be able to: (a)Explain the lack of reactivity of nitrogen (b)Describe: (i)The formation and structure of the ammonium ion (ii)The displacement of ammonia from its salts (c) describe the Haber process for the manufacture of ammonia from its elements, giving essential operating conditions, and interpret these conditions (qualitatively) in terms of the principles of kinetics and equilibria (d) Understand the industrial importance of ammonia and nitrogen compounds derived from ammonia (e) Understand the environmental consequences of the controlled use of nitrate fertilisers (f) Understand and explain the occurrence and catalytic removal of oxides of nitrogen

2 (g) Explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide (h) Describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels (i)State the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain (j) State the main details of the Contact process for sulphuric acid production (k) Understand the industrial importance of sulphuric acid (l) Describe the use of sulphur dioxide in food preservation

3 (a)Explain the lack of reactivity of nitrogen: Group V N::N Nitrogen is a gas making up approximately 78% of the atmosphere with a triple bond energy of 945 KJmol -1 hence explaining the lack of reactivity compared to other non-metals It does react with the s-block metals, e.g. 3Mg(s) + N 2 (g) Mg 3 N 2 (s) Magnesium nitride This results in a large activation energy, E a

4 (b) Describe: The formation and structure of the ammonium ion The displacement of ammonia from its salts Ammonia, is produced commercially by the Haber Process. N 2 (g) + 3H 2 (g) 2NH 3 (g)  H = -92 KJmol -1 It can be acidified to give the ammonium ion: NH 3 (aq) + H + (aq) NH 4 + (aq) Ammonia is alkaline (a weak base) in aqueous solution and a typical reaction would be using dilute hydrochloric acid: NH 3 (aq) + HCl(aq) NH 4 + (aq) + Cl - (aq) Ammonia is clearly a base as it is a proton acceptor according to the Bronsted- Lowry definition of acids and bases.

5 Structure of the ammonium ion Ammonia dissociates in water to a small extent (1%) NH 3 (aq) + H 2 0(l) NH 4 + (aq) + OH - (aq) Can you remember the conjugate pairs (of acids and bases)? It has a tetrahedral shape similar to methane (CH 4 )

6 In the gas phase, ammonia gas can react with hydrogen chloride: NH 3 (g) + HCl(g) NH 4 Cl(s) NH 3 (g) HCl(g) cotton wool soaked in conc. ammonia in conc. HCl(aq)

7 Displacement of ammonia from its salts If we add sodium hydroxide to ammonium chloride, the characteristic smell of ammonia is produced: NaOH(aq) + NH 4 Cl(aq) NH 3 (aq) + NaCl(aq) + H 2 O(l) The same is also true of organic amine salts: NaOH(aq) + CH 3 CH 2 NH 3 Cl(aq) CH 3 CH 2 NH 2 (aq) + NaCl(aq) + H 2 O(l) The group I metal hydroxide is clearly the stronger base!

8 (c) Describe the Haber process for the manufacture of ammonia from its elements, giving essential operating conditions, and interpret these conditions (qualitatively) in terms of the principles of kinetics and equilibria Ammonia, NH 3, is produced commercially by the Haber Process. N 2 (g) + 3H 2 (g) 2NH 3 (g)  H = -92 KJmol -1 Fe 460 o C 200 atm

9 Haber Process Iron catalyst

10 Operating Conditions N 2 and H 2 are pumped into a chamber. The pre-heated gases are passed through a heating coil to the catalyst bed. The catalyst bed is kept at  C under high pressure. The product gas stream (containing N 2, H 2 and NH 3 ) is passed over a cooler to a refrigeration unit. In the refrigeration unit, ammonia liquefies, but not N 2 or H 2. Haber Process (contd.)

11 1. Kinetics A catalyst (iron) is used to speed up the rate of reaction and to lower the high activation energy in breaking the N 2 triple bond. However, if the temperature is too high, it begins to get destroyed and must be changed more regularly. Haber Process (contd.)

12 In the Haber Process for the production of ammonia, based on the reversible reaction: it is observed that: As the total pressure increases, the amount of ammonia present at equilibrium increases. As the temperature decreases, the amount of ammonia at equilibrium increases. Effect of temperature and pressure on Haber Process An example:

13 Haber Process

14 (d) Understand the industrial importance of ammonia and nitrogen compounds derived from ammonia Most of the ammonia produced is used to make nitric acid and fertilizers. 1.Ammonia and oxygen are passed over a hot platinum catalyst: About 95% of the ammonia is converted into nitrogen monoxide 2. The gases are cooled and mixed with more air. Then under pressure, they are passed up a tower that has water trickling down it and nitric acid is produced: 4NO(g) + 3O 2 (g) + 2H 2 O(l) 4HNO 3 (aq) Overall: NH 3 (g) + 2O 2 (g) HNO 3 (aq) + H 2 O(l) 4NH 3 (g) + 5O 2 (g) 4NO(g) + 6H 2 O(l)  H = -909 KJmol -1 Conditions: 1100K and 4-10 atm. pressure Pt

15 The conversion of NO into nitric acid actually occurs in three stages: a.2NO(g) + O 2 (g) 2NO 2 (g)  H = -115 KJmol -1 b. 2NO 2 (g) N 2 O 4 (g)  H = -58 KJmol -1 c. 3N 2 O 4 (g) + 2H 2 O(l) 4HNO 3 (aq) + 2NO(g) Nitric acid is used to make fertilizers, explosives and the polyamide, Nylon Commercial uses of nitric acid: Fertilizers 化肥 Ammonium salts are used as fertilizers because they contain nitrogen in a form that plants can use. The fertilizer Nitram is ammonium nitrate and is made from a solution of ammonia in water and nitric acid in an acid-base reaction: NH 3 (aq) + HNO 3 (aq) NH 4 NO 3 (aq) Similar neutralization reactions with phosphoric(V) and sulphuric acids: 2NH 3 (aq) + H 2 SO 4 (aq) (NH 4 ) 2 SO 4 (aq)

16 (e) Understand the environmental consequences of the controlled use of nitrate fertilisers Ammonium nitrate is a good fertilizer: it contains 35% by mass of nitrogen all of which is available to plants as it dissolves easily. However, some nitrate ions are washed into rivers and streams which encourages rapid growth of algae. However, when algae die they are decomposed by bacteria and this uses up oxygen dissolved in water. The fish begin to suffocate and die from a lack of oxygen. Nitrates also get into drinking water as they are not purified and at high levels can cause problems for small babies by affecting haemoglobin (used in respiration and transport of O 2 in the blood).

17 In today’s lesson….. (d) Understand the industrial importance nitrogen compounds derived from ammonia (revision from yesterday) (e) Understand the environmental consequences of the controlled use of nitrate fertilisers (revision from yesterday) (f) Understand and explain the occurrence and catalytic removal of oxides of nitrogen (new) (g) Explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide (new) We can test some of our knowledge in a past-paper style question on catalytic removal of NO in car engines We can test some of our knowledge in a past-paper style question on manufacture of nitric acid

18 (d) Understand the industrial importance of ammonia and nitrogen compounds derived from ammonia Most of the ammonia produced is used to make nitric acid and fertilizers. 1.Ammonia and oxygen are passed over a hot platinum catalyst: About 95% of the ammonia is converted into nitrogen monoxide 2. The gases are cooled and mixed with more air. Then under pressure, they are passed up a tower that has water trickling down it and nitric acid is produced: 4NO(g) + 3O 2 (g) + 2H 2 O(l) 4HNO 3 (aq) Overall: NH 3 (g) + 2O 2 (g) HNO 3 (aq) + H 2 O(l) 4NH 3 (g) + 5O 2 (g) 4NO(g) + 6H 2 O(l)  H = -909 KJmol -1 Conditions: 1100K and 4-10 atm. pressure Pt

19 The conversion of NO into nitric acid actually occurs in three stages: a.2NO(g) + O 2 (g) 2NO 2 (g)  H = -115 KJmol -1 b. 2NO 2 (g) N 2 O 4 (g)  H = -58 KJmol -1 c. 3N 2 O 4 (g) + 2H 2 O(l) 4HNO 3 (aq) + 2NO(g) Concentrated nitric acid (66%) is a strong acid: HNO 3 (aq) H + (aq) + NO 3 - (aq) It is also an oxidising agent and reacts with most metals to form nitrates and nitrogen oxides: Cu(s) + 2HNO 3 (aq) Cu(NO 3 ) 2 (aq) + H 2 O(l) + NO 2 (g) Properties of nitric(v) acid

20 Nitric acid is used to make fertilizers, explosives and the polyamide, Nylon Commercial uses of nitric acid: Fertilisers 化肥 Ammonium salts are used as fertilisers because they contain nitrogen in a form that plants can use. The fertiliser Nitram is ammonium nitrate and is made from a solution of ammonia in water and nitric acid in an acid-base reaction: NH 3 (aq) + HNO 3 (aq) NH 4 NO 3 (aq) Similar neutralization reactions with phosphoric(V) and sulphuric acids: 2NH 3 (aq) + H 2 SO 4 (aq) (NH 4 ) 2 SO 4 (aq)

21 (e) Understand the environmental consequences of the controlled use of nitrate fertilisers Ammonium nitrate is a good fertiliser: it contains 35% by mass of nitrogen all of which is available to plants as it dissolves easily. However, some nitrate ions are washed into rivers and streams which encourages rapid growth of algae 藻类. However, when algae die they are decomposed by bacteria and this uses up oxygen dissolved in water. The fish begin to suffocate and die from a lack of oxygen. Nitrates also get into drinking water as they are not purified and at high levels can cause problems for small babies by affecting haemoglobin 血红蛋白 (used in respiration and transport of O 2 in the blood).

22 (f) Understand and explain the occurrence and catalytic removal of oxides of nitrogen  The combustion engines in cars produce a number of polluting gases, including a small amount of NO x :  Carbon monoxide, which is poisonous.  Nitrogen oxides (NO, NO 2 and N 2 O 4 ) which cause acid rain and help to destroy the ozone (O 3 ) layer in the upper atmosphere.  Unburnt hydrocarbons may cause cancer.  Unburnt hydrocarbons react with nitrogen oxides in sunlight to form smogs (containing ozone) at low levels in the atmosphere, especially in large hot cities. N 2 (g)+ O 2 (g) 2NO(g) : the excess heat in an engine produces NO

23 Catalytic removal of nitrogen oxides These poisonous gases may be removed by the use of a catalytic converter In car exhaust systems. It is honeycomb in shape to provide a large surface area and is coated with platinum and rhodium metal catalysts: Carbon monoxide + nitrogen oxide nitrogen + carbon dioxide Hydrocarbons + nitrogen oxide nitrogen + carbon dioxide + water Pt or Rh 1100K NO-CON2N2 CO 2 ON CO 2 CO Inside a catalytic converter

24 (g) Explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide Nitrogen oxides (NO, NO 2 and N 2 O 4 ) which cause acid rain and help to destroy the ozone (O 3 ) layer in the upper atmosphere. In more detail: Nitric oxide (NO) reacts with oxygen to produce nitrogen dioxide: 2NO(g) + O 2 (g) 2NO 2 (g) NO 2 dissolves in rain to form a mixture of nitric(v) and nitric(III) acids: 2NO 2 (g) + H 2 O(l) HNO 3 (aq) + HNO 2 (aq) What type of redox reaction is this? disproportionation These acids are help to form acid rain; they are corrosive and damage construction materials that are attacked by acids especially marble ( 大理石 ), so some ancient monuments lose their features.

25 NO and NO 2 have a catalytic role in the conversion of atmospheric sulphur dioxide to sulphur trioxide: 2NO(g) + O 2 (g) 2NO 2 (g) SO 2 (g) + NO 2 (g) SO 3 (g) + NO(g) Did you know? The explosive TNT uses fuming nitric acid (98%) in its manufacture Trinitrotoluene Sulphur trioxide dissolves in rain water to form sulphuric acid and aquatic life may be destroyed in lakes and it also destroys limestone and marble.

26 To summarize….. (d) Understand the industrial importance nitrogen compounds derived from ammonia (e) Understand the environmental consequences of the controlled use of nitrate fertilisers (f) Understand and explain the occurrence and catalytic removal of oxides of nitrogen (g) Explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide In tomorrow’s lesson we will study sulphur chemistry, in particular the contact process for formation of SO 3 and H 2 SO 4

27 In today’s lesson….. (f) Understand and explain the occurrence and catalytic removal of oxides of nitrogen (revision) (g) Explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide (rev.) (h) Describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels (new) (i) State the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain (new) (j) State the main details of the contact process for sulphuric acid production (also see equilibria hand-out) (new) We will test our knowledge using an exam style question to check our revision of nitrogen chemistry, nitric oxides and catalytic role with SO 2

28 (f) Understand and explain the occurrence and catalytic removal of oxides of nitrogen  The combustion engines in cars produce a number of polluting gases, including a small amount of NO x :  Carbon monoxide, which is poisonous.  Nitrogen oxides (NO, NO 2 and N 2 O 4 ) which cause acid rain and help to destroy the ozone (O 3 ) layer in the upper atmosphere.  Unburnt hydrocarbons may cause cancer.  Unburnt hydrocarbons react with nitrogen oxides in sunlight to form smogs (containing ozone) at low levels in the atmosphere, especially in large hot cities. N 2 (g)+ O 2 (g) 2NO(g) : the excess heat in an engine produces NO

29 Catalytic removal of nitrogen oxides These poisonous gases may be removed by the use of a catalytic converter In car exhaust systems. It is honeycomb in shape to provide a large surface area and is coated with platinum and rhodium metal catalysts: Carbon monoxide + nitrogen oxide nitrogen + carbon dioxide Hydrocarbons + nitrogen oxide nitrogen + carbon dioxide + water Pt or Rh 1100K NO-CON2N2 CO 2 ON CO 2 CO Inside a catalytic converter

30 (g) Explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide Nitrogen oxides (NO, NO 2 and N 2 O 4 ) which cause acid rain and help to destroy the ozone (O 3 ) layer in the upper atmosphere. In more detail: Nitric oxide (NO) reacts with oxygen to produce nitrogen dioxide: 2NO(g) + O 2 (g) 2NO 2 (g) NO 2 dissolves in rain to form a mixture of nitric(v) and nitric(III) acids: 2NO 2 (g) + H 2 O(l) HNO 3 (aq) + HNO 2 (aq) What type of redox reaction is this? disproportionation These acids are help to form acid rain; they are corrosive and damage construction materials that are attacked by acids especially marble ( 大理石 ), so some ancient monuments lose their features.

31 NO and NO 2 have a catalytic role in the conversion of atmospheric sulphur dioxide to sulphur trioxide: 2NO(g) + O 2 (g) 2NO 2 (g) SO 2 (g) + NO 2 (g) SO 3 (g) + NO(g) Did you know? The explosive TNT uses fuming nitric acid (98%) in its manufacture Trinitrotoluene Sulphur trioxide dissolves in rain water to form sulphuric acid and aquatic life may be destroyed in lakes and it also destroys limestone and marble.

32 Sulphur (Group VI) (h) Describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels Coal, an impure form of graphite is used as a fossil fuel and is still used in boilers to heat water. It was more popular in England more than one hundred years ago in the Victorian Days. However, small amounts of sulphur are oxidized to corrosive and foul smelling sulphur dioxide, which contributes to pollution: S(s) + O 2 (g) SO 2 (g) structure of SO 2

33 SO 2 (g) + H 2 O(l) H 2 SO 3 (aq) (i) State the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain Sulphur dioxide released into the atmosphere dissolves in rain and forms sulphuric (IV) acid more commonly known as sulphurous acid; some SO 2 comes from natural sources such as volcanoes, where hot sulphur reacts with O 2 : This is corrosive and damages construction materials that are attacked by acids especially marble ( 大理石 ), so some ancient monuments lose their features. Acid-rain also effects forests and kills some plants, algae and fish in natural lakes, and the more acid-rain produced the worse the effect. formation of acid-rain

34 (j) State the main details of the contact process for sulphuric acid production (see equilibria hand-out) Sulphuric acid, H 2 SO 4 is produced from sulphur, water and oxygen by the Contact process. 1. Sulphur is burnt to sulphur dioxide: S(s) + O 2 (g) SO 2 (g) 2. Sulphur dioxide is converted to suphur trioxide by reaction with more oxygen using a catalyst of vanadium(v) oxide: 2SO 2 (g) + O 2 (g) 2SO 3 (g)  H = -197 KJmol -1 V 2 O 5 710K

35 Effect of temperature and pressure on the contact process Le Chateliers principle predicts that for the exothermic forward reaction cooling will produce the maximum yield (95%): the equilibrium is shifted to the right. The SO 3 is removed by dissolving it in conc. H 2 SO 4. However, the rate of reaction to reach equilibrium increases with temperature increase. The yield can be increased to 99.5% by passing the remaining reactant gases over the hot catalyst (on a fourth catalyst bed). Although increasing the pressure will shift the equilbrium to the right as fewer gas molecules are produced, there are more operational energy costs and in practice the process is run at near atmospheric pressure (only an extra 0.5% conversion achieved!) 2SO 2 (g) + O 2 (g) 2SO 3 (g)  H = 197 KJmol -1 V 2 O 5 710K  Temperature  Pressure

36 A flow diagram Purifier and drier Air Heat exchanger Catalytic chamber Absorption tower Purifier and drier Sulphur burner Air H 2 SO 4 store water 98% H 2 SO 4 c.H 2 SO 4 oleum drySO 2 +air SO 3

37 3. Formation of H 2 SO 4 The absorbed SO 3 reacts with the water in sulphuric acid to form more sulphuric acid so that the acid becomes even more concentrated: SO 3 (g) + H 2 O(l) H 2 SO 4 (l) SO 3 (g) + H 2 SO 4 (l) H 2 S 2 O 7 (l) Properties of H 2 SO 4 : Viscous liquid with hydrogen bonding 1. Strong acid: H 2 SO 4 (l) H + (aq) + HSO 4 - (aq) HSO 4 - (aq) H + (aq) + SO 4 2- (aq) 2. Dehydrating agent: CuSO 4.5H 2 O(s) CuSO 4 (s) + 5H 2 O(l) 3. Oxidizing agent 8HI(g) + H 2 SO 4 (l) 4H 2 O(l) + H 2 S(g) + 4I 2 (s) H 2 SO 4 H2OH2O H 2 S 2 O 7 (l) + H 2 O(l) 2H 2 SO 4 (l) oleum

38 Practice questions Sulphuric acid is manufactured from sulphur in four stages S(s) + O 2 (g) SO 2 (g) 2SO 2 (g) + O 2 (g) 2SO 3 (g) SO 3 (g) + H 2 SO 4 (l) H 2 S 2 O 7 (l) H 2 S 2 O 7 (l) + H 2 O(l) 2H 2 SO 4 (l) (a)Give the oxidation number of sulphur in each species. (b)Which two steps are redox reactions? 0+IV +VI √ √

39 To summarize what we have learnt today….. (h) Describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels (i) State the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain (j) State the main details of the Contact process for sulphuric acid production Tomorrow we will understand: (k) Understand the industrial importance of sulphuric acid (l) Describe the use of sulphur dioxide in food preservation Revise nitrogen and sulphur

40 In today’s lesson….. (h) Describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels (rev.) (i) State the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain (rev.) (j) State the main details of the contact process for sulphuric acid production (also see equilibria hand-out) (rev.) (k) Understand the industrial importance of sulphuric acid (new) (l) Describe the use of sulphur dioxide in food preservation (new) Keywords: SO 2, acid-rain, SO 3, Contact Process, H 2 SO 4, SO 2 -food preservation

41 Sulphur (Group VI) (h) Describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels Coal, an impure form of graphite is used as a fossil fuel and is still used in boilers to heat water. It was more popular in England more than one hundred years ago in the Victorian Days. However, small amounts of sulphur are oxidized to corrosive and foul smelling sulphur dioxide, which contributes to pollution: S(s) + O 2 (g) SO 2 (g) structure of SO 2

42 SO 2 (g) + H 2 O(l) H 2 SO 3 (aq) (i) State the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain Sulphur dioxide released into the atmosphere dissolves in rain and forms sulphuric (IV) acid more commonly known as sulphurous acid; some SO 2 comes from natural sources such as volcanoes, where hot sulphur reacts with O 2 : This is corrosive and damages construction materials that are attacked by acids especially marble ( 大理石 ), so some ancient monuments lose their features. Acid-rain also effects forests and kills some plants, algae and fish in natural lakes, and the more acid-rain produced the worse the effect. formation of acid-rain

43 (j) State the main details of the contact process for sulphuric acid production (see equilibria hand-out) Sulphuric acid, H 2 SO 4 is produced from sulphur, water and oxygen by the Contact process. 1. Sulphur is burnt to sulphur dioxide: S(s) + O 2 (g) SO 2 (g) 2. Sulphur dioxide is converted to suphur trioxide by reaction with more oxygen using a catalyst of vanadium(v) oxide: 2SO 2 (g) + O 2 (g) 2SO 3 (g)  H = -197 KJmol -1 V 2 O 5 710K

44 Effect of temperature and pressure on the contact process Le Chateliers principle predicts that for the exothermic forward reaction cooling will produce the maximum yield (95%): the equilibrium is shifted to the right. The SO 3 is removed by dissolving it in conc. H 2 SO 4. However, the rate of reaction to reach equilibrium increases with temperature increase. The yield can be increased to 99.5% by passing the remaining reactant gases over the hot catalyst (on a fourth catalyst bed). Although increasing the pressure will shift the equilbrium to the right as fewer gas molecules are produced, there are more operational energy costs and in practice the process is run at near atmospheric pressure (only an extra 0.5% conversion achieved!) 2SO 2 (g) + O 2 (g) 2SO 3 (g)  H = 197 KJmol -1 V 2 O 5 710K  Temperature  Pressure

45 A flow diagram Purifier and drier Air Heat exchanger Catalytic chamber Absorption tower Purifier and drier Sulphur burner Air H 2 SO 4 store water 98% H 2 SO 4 c.H 2 SO 4 oleum drySO 2 +air SO 3

46 3. Formation of H 2 SO 4 The absorbed SO 3 reacts with the water in sulphuric acid to form more sulphuric acid so that the acid becomes even more concentrated: SO 3 (g) + H 2 O(l) H 2 SO 4 (l) SO 3 (g) + H 2 SO 4 (l) H 2 S 2 O 7 (l) Properties of H 2 SO 4 : Viscous liquid with hydrogen bonding 1. Strong acid: H 2 SO 4 (l) H + (aq) + HSO 4 - (aq) HSO 4 - (aq) H + (aq) + SO 4 2- (aq) 2. Dehydrating agent: CuSO 4.5H 2 O(s) CuSO 4 (s) + 5H 2 O(l) 3. Oxidizing agent 8HI(g) + H 2 SO 4 (l) 4H 2 O(l) + H 2 S(g) + 4I 2 (s) H 2 SO 4 H2OH2O H 2 S 2 O 7 (l) + H 2 O(l) 2H 2 SO 4 (l) oleum

47 (k) Understand the industrial importance of sulphuric acid Sulphuric acid is an important industrially produced chemical. Its main uses are:-  Paints and pigments 染料 涂料  Fertilisers 化肥  Fibres and dyes 县委 染料  Soaps and detergents 香皂 洗涤剂  Cleaning metals 除锈剂  Chemicals and plastics 化学品 塑料  Tanning leather 皮革

48 (l) Describe the use of sulphur dioxide in food preservation (new) Sulphur dioxide is poisonous to bacteria and is used as a food preservative (E220 on food labels in the U.K.) especially in the preparation of fruit juices and soft drinks and to kill bacteria in wine-making as it is slightly soluble in water. SO 2 (g) + H 2 O(l) H 2 SO 3 (aq) H + (aq) + HSO 3 - (aq) Sodium metabisulphite (Na 2 S 2 O 5 ) is another preservative (E222) and liberates SO 2 (g) upon acidification. In today’s lesson we learnt: (k) Understand the industrial importance of sulphuric acid (l) Describe the use of sulphur dioxide in food preservation


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