1. IGCSE CHEMISTRY SECTION 2 LESSON 3

2. Content The iGCSE Chemistry course Section 1 Principles of Chemistry Section 2 Chemistry of the Elements Section 3 Organic Chemistry Section 4 Physical Chemistry Section 5 Chemistry in Society

3. Content Section 2 Chemistry of the Elements a)The Periodic Table b)Group 1 Elements c)Group 7 Elements d)Oxygen and Oxides e)Hydrogen and Water f)Reactivity Series g)Tests for ions and gases

4. Lesson 3 d)Oxygen and oxides e)Hydrogen and water d) Oxygen and oxides 2.16 recall the gases present in air and their approximate percentage by volume 2.17 explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air 2.18 describe the laboratory preparation of oxygen from hydrogen peroxide,using manganese(IV) oxide as a catalyst 2.19 describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced 2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid 2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate 2.22 describe the properties of carbon dioxide, limited to its solubility and density 2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density 2.24 understand that carbon dioxide is a greenhouse gas and may contribute to climate change. e) Hydrogen and water 2.25 describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron 2.26 describe the combustion of hydrogen 2.27 describe the use of anhydrous copper(II) sulfate in the chemical test for water 2.28 describe a physical test to show whether water is pure.

5. Magnesium + Oxygen

6. Magnesium + Oxygen Carbon + Oxygen

7. Magnesium + Oxygen Carbon + Oxygen Sulphur + Oxygen

8. Magnesium + Oxygen

9. What happens? When heated in a Bunsen burner flame magnesium burns very brightly to form magnesium oxide.

10. What happens? When heated in a Bunsen burner flame magnesium burns very brightly to form magnesium oxide. Magnesium + Oxygen  Magnesium oxide 2Mg (s) + O 2(g)  2MgO (s)

11. What happens? When heated in a Bunsen burner flame magnesium burns very brightly to form magnesium oxide. Magnesium + Oxygen  Magnesium oxide 2Mg (s) + O 2(g)  2MgO (s) Magnesium oxide is a basic oxide. MgO is almost insoluble, so when added to water not many hydroxide ions are formed, but it is slightly alkaline (pH~9)

12. Carbon + Oxygen

13. What happens? Carbon, as graphite (solid), burns to form gaseous carbon dioxide.

14. What happens? Carbon, as graphite (solid), burns to form gaseous carbon dioxide. Carbon + Oxygen  Carbon dioxide C (s) + O 2(g)  CO 2(g)

15. What happens? Carbon, as graphite (solid), burns to form gaseous carbon dioxide. Carbon + Oxygen  Carbon dioxide C (s) + O 2(g)  CO 2(g) Carbon dioxide is an acidic oxide and reacts with water to give carbonic acid CO 2(g) + H 2 O (l)  H 2 CO 3(aq)

16. Sulphur + Oxygen

17. What happens? Solid sulphur burns in air to form sulphur dioxide, a suffocating gas.

18. What happens? Solid sulphur burns in air to form sulphur dioxide, a suffocating gas. Sulphur + Oxygen  Sulphur dioxide S (s) + O 2(g)  SO 2(g)

19. What happens? Solid sulphur burns in air to form sulphur dioxide, a suffocating gas. Sulphur + Oxygen  Sulphur dioxide S (s) + O 2(g)  SO 2(g) Sulphur dioxide is an acidic oxide and reacts with water to give sulphurous acid SO 2(g) + H 2 O (l)  H 2 SO 3(aq)

20. Carbon dioxide Laboratory preparation:

21. Carbon dioxide Laboratory preparation: Carbon dioxide is usually prepared in the laboratory by the reaction of dilute hydrochloric acid and calcium carbonate (marble chips). The gas can be collected over water. Hydrochloric acid Carbon dioxide Calcium carbonate

22. Carbon dioxide Laboratory preparation: Carbon dioxide is usually prepared in the laboratory by the reaction of dilute hydrochloric acid and calcium carbonate (marble chips). The gas can be collected over water. Hydrochloric acid Carbon dioxide Calcium carbonate CaCO 3(s) + 2HCl (aq)  CaCl 2(aq) + H 2 O (l) + CO 2(g)

23. Carbon dioxide Laboratory preparation: Carbon dioxide can also be prepared by the thermal decomposition of a metal carbonate such as copper carbonate.

24. Carbon dioxide Laboratory preparation: Carbon dioxide can also be prepared by the thermal decomposition of a metal carbonate such as copper carbonate. Thermal decomposition = when a compound is split up by heating to form products which do not recombine on cooling

25. Carbon dioxide Laboratory preparation: Carbon dioxide can also be prepared by the thermal decomposition of a metal carbonate such as copper carbonate. http://www.bbc.co.uk/schools/gcsebitesize/scien ce/aqa_pre_2011/rocks/limestonerev1.shtml

26. Carbon dioxide Laboratory preparation: Carbon dioxide can also be prepared by the thermal decomposition of a metal carbonate such as copper carbonate. http://www.bbc.co.uk/schools/gcsebitesize/scien ce/aqa_pre_2011/rocks/limestonerev1.shtml CuCO 3(s)  CuO (s) + CO 2(g)

27. Carbon dioxide Properties: PHYSICALCHEMICAL

28. Carbon dioxide Properties: PHYSICALCHEMICAL Carbon dioxide is a colourless gas It is denser than air and can be ‘poured’ over a burning splint to put it out Under pressure it is quite soluble in water – fizzy drinks contain carbon dioxide Solid carbon dioxide is known as ‘dry ice’. When heated: Solid  Gas (sublimation)

29. Carbon dioxide Properties: PHYSICALCHEMICAL Carbon dioxide is a colourless gas When bubbled through water it dissolves to form the weakly acidic carbonic acid. It is denser than air and can be ‘poured’ over a burning splint to put it out It does not support combustion Under pressure it is quite soluble in water – fizzy drinks contain carbon dioxide It turns calcium hydroxide solution (limewater) milky or cloudy. This is the classic test for carbon dioxide gas. Solid carbon dioxide is known as ‘dry ice’. When heated: Solid  Gas (sublimation)

30. Carbon dioxide Uses : FIZZY DRINKS A large amount of carbon dioxide can be dissolved in water. The carbon dioxide is released when the bottle is opened and the pressure reduced. The fizzing is carbon dioxide leaving the solution.

31. Carbon dioxide Uses : FIRE EXTINGUISHERS Carbon dioxide fire extinguishers are primarily used for fighting electrical fires. CO 2 is denser than air, so when sprayed onto a fire it smothers it and prevents oxygen from getting in.

32. Carbon dioxide As a greenhouse gas:

33. Carbon dioxide As a greenhouse gas: When fossil fuels burn, carbon dioxide gas is produced. The build-up of gases such as carbon dioxide could increase the greenhouse effect of the Sun’s radiation, leading to a heating up of the Earth’s atmosphere.

34. Carbon dioxide As a greenhouse gas: EARTH Atmosphere

35. Carbon dioxide As a greenhouse gas: EARTH Heat rays from the Sun Atmosphere

36. Carbon dioxide As a greenhouse gas: EARTH Heat rays from the Sun Atmosphere Heat rays trapped in the Earth’s atmosphere lead to a greenhouse effect.

37. Lesson 3 d)Oxygen and oxides e)Hydrogen and water d) Oxygen and oxides 2.16 recall the gases present in air and their approximate percentage by volume 2.17 explain how experiments involving the reactions of elements such as copper, iron and phosphorus with air can be used to investigate the percentage by volume of oxygen in air 2.18 describe the laboratory preparation of oxygen from hydrogen peroxide,using manganese(IV) oxide as a catalyst 2.19 describe the reactions of magnesium, carbon and sulfur with oxygen in air, and the acid-base character of the oxides produced 2.20 describe the laboratory preparation of carbon dioxide from calcium carbonate and dilute hydrochloric acid 2.21 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates such as copper(II) carbonate 2.22 describe the properties of carbon dioxide, limited to its solubility and density 2.23 explain the use of carbon dioxide in carbonating drinks and in fire extinguishers, in terms of its solubility and density 2.24 understand that carbon dioxide is a greenhouse gas and may contribute to climate change. e) Hydrogen and water 2.25 describe the reactions of dilute hydrochloric and dilute sulfuric acids with magnesium, aluminium, zinc and iron 2.26 describe the combustion of hydrogen 2.27 describe the use of anhydrous copper(II) sulfate in the chemical test for water 2.28 describe a physical test to show whether water is pure.

38. Magnesium Dilute hydrochloric and dilute sulphuric acids AluminiumZincIron

39. What is an acid?

40. An acid is a substance which contains hydrogen which may be replaced by a metal to form a salt. An acid is a substance which forms hydrogen ions in solution. ACID CORNER

41. What is an acid? An acid is a substance which contains hydrogen which may be replaced by a metal to form a salt. An acid is a substance which forms hydrogen ions in solution. ACID CORNER Acids change moist litmus paper from blue to red. Acids are soluble in water. Acids are electrolytes. Acids can have a sour or sharp taste.

42. What is an acid? An acid is a substance which contains hydrogen which may be replaced by a metal to form a salt. An acid is a substance which forms hydrogen ions in solution. ACID CORNER Acids change moist litmus paper from blue to red. Acids are soluble in water. Acids are electrolytes. Acids can have a sour or sharp taste. Hydrochloric acid HCl Sulphuric acid H 2 SO 4 Nitric acid HNO 3 Carbonic acid H 2 CO 3

43. What is an acid? An acid is a substance which contains hydrogen which may be replaced by a metal to form a salt. An acid is a substance which forms hydrogen ions in solution. ACID CORNER Acids change moist litmus paper from blue to red. Acids are soluble in water. Acids are electrolytes. Acids can have a sour or sharp taste. Hydrochloric acid HCl Sulphuric acid H 2 SO 4 Nitric acid HNO 3 Carbonic acid H 2 CO 3 Acid + Alkali  Salt + Water

44. Acids + Metals acid + reactive metal  salt + hydrogen A ‘reactive’ metal is any metal higher than copper in the reactivity series.

45. Acids + Metals Hydrochloric acid + Magnesium  Magnesium chloride + Hydrogen 2HCl(aq) + Mg(s) MgCl 2 (aq) + H 2 (g) Sulphuric acid + Magnesium  Magnesium sulphate + Hydrogen H 2 SO 4 (aq) + Mg(s) MgSO 4 (aq) + H 2 (g)

46. Acids + Metals Hydrochloric acid + Aluminium  Aluminium chloride + Hydrogen 6HCl(aq) + 2Al(s) 2AlCl 3 (aq) + 3H 2 (g) Sulphuric acid + Aluminium  Aluminium sulphate + Hydrogen 3H 2 SO 4 (aq) + 2Al(s) Al 2 (SO 4 ) 3 (aq) + 3H 2 (g)

47. Acids + Metals Hydrochloric acid + Zinc  Zinc chloride + Hydrogen 2HCl(aq) + Zn(s) ZnCl 2 (aq) + H 2 (g) Sulphuric acid + Zinc  Zinc sulphate + Hydrogen H 2 SO 4 (aq) + Zn(s) ZnSO 4 (aq) + H 2 (g)

48. Acids + Metals Hydrochloric acid + Iron  Iron chloride + Hydrogen 6HCl(aq) + 2Fe(s) 2FeCl 3 (aq) + 3H 2 (g) Sulphuric acid + Iron  Iron sulphate + Hydrogen 3H 2 SO 4 (aq) + 2Fe(s) Fe 2 (SO 4 ) 3 (aq) + 3H 2 (g)

49. The combustion of hydrogen

50. Properties of hydrogen colourless and odourless gas lightest of all gases and diffuses very rapidly not very soluble in water not poisonous, but does not support life boiling point of 20K (-253 o C) has three isotopes, hydrogen, deuterium and tritium

51. The combustion of hydrogen Hydrogen burns very rapidly in oxygen: 2H 2(g) + O 2(g)  2H 2 O (g)

52. The combustion of hydrogen Hydrogen burns very rapidly in oxygen: 2H 2(g) + O 2(g)  2H 2 O (g) This is a commonly used reaction in rocket engines, and because the product is water vapour, hydrogen cells are being developed for cars

53. Describe the use of anhydrous copper (II) sulphate in the chemical test for water

54. A salt which has lost its water of crystallisation is called an anhydrous salt. When water is added to an anhydrous salt, the salt becomes hydrated

55. Testing for water + water Anhydrous copper sulphate Hydrated copper sulphate CuSO 4(s) + 5H 2 O (l)  CuSO 4.5H 2 O (aq) + heat

56. Testing for water + water Anhydrous copper sulphate Hydrated copper sulphate CuSO 4(s) + 5H 2 O (l)  CuSO 4.5H 2 O (aq) + heat exothermic

57. How can we test to show whether water is pure?

58. 110 o C 100 90 80 70 60 50 40 30 20 10 0 Only pure water boils at exactly 100 o C.

59. How can we test to show whether water is pure? 110 o C 100 90 80 70 60 50 40 30 20 10 0 Only pure water boils at exactly 100 o C. If it boils above or below this then impurities must be present.

60. How can we test to show whether water is pure? 110 o C 100 90 80 70 60 50 40 30 20 10 0 Only pure water boils at exactly 100 o C. If it boils above or below this then impurities must be present. Don’t assume that pH7 is a test for pure water!

61. End of Section 2 Lesson 3 In this lesson we have covered: The reactions of oxygen with metals Carbon dioxide – preparation and properties Hydrogen – reactions and properties