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iGCSE chemistry Section 2 lesson 3

iGCSE chemistry Section 2 lesson 3. 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. Content. Section 2

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iGCSE chemistry Section 2 lesson 3

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  1. iGCSE chemistrySection 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 • The Periodic Table • Group 1 Elements • Group 7 Elements • Oxygen and Oxides • Hydrogen and Water • Reactivity Series • Tests for ions and gases

  4. 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. • Lesson 3 • Oxygen and oxides • Hydrogen and water

  5. Magnesium + Oxygen

  6. Magnesium + Oxygen Carbon + Oxygen

  7. Magnesium + Oxygen Sulphur + Oxygen Carbon + 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) + O2(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) + O2(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) + O2(g)  CO2(g)

  15. What happens? Carbon, as graphite (solid), burns to form gaseous carbon dioxide. Carbon + Oxygen  Carbon dioxide C(s) + O2(g)  CO2(g) Carbon dioxide is an acidic oxide and reacts with water to give carbonic acid CO2(g) + H2O(l) H2CO3(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) + O2(g)  SO2(g)

  19. What happens? Solid sulphur burns in air to form sulphur dioxide, a suffocating gas. Sulphur + Oxygen  Sulphur dioxide S(s) + O2(g)  SO2(g) Sulphur dioxide is an acidic oxide and reacts with water to give sulphurous acid SO2(g) + H2O(l) H2SO3(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 CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(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/science/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/science/aqa_pre_2011/rocks/limestonerev1.shtml CuCO3(s)  CuO(s) + CO2(g)

  27. Carbon dioxide Properties:

  28. Carbon dioxide Properties:

  29. Carbon dioxide Properties:

  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. CO2 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: Atmosphere EARTH

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

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

  37. 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. • Lesson 3 • Oxygen and oxides • Hydrogen and water

  38. Dilute hydrochloric and dilute sulphuric acids Magnesium Aluminium Zinc Iron

  39. What is an acid?

  40. 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

  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. 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. ACID CORNER

  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. 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 H2SO4 Nitric acid HNO3 Carbonic acid H2CO3 ACID CORNER

  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. 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 H2SO4 Nitric acid HNO3 Carbonic acid H2CO3 ACID CORNER 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) MgCl2(aq) + H2(g) Sulphuric acid + Magnesium  Magnesium sulphate + Hydrogen H2SO4(aq) + Mg(s) MgSO4(aq) + H2(g)

  46. Acids + Metals Hydrochloric acid + Aluminium  Aluminium chloride + Hydrogen 6HCl(aq) + 2Al(s) 2AlCl3(aq) + 3H2(g) Sulphuric acid + Aluminium  Aluminium sulphate + Hydrogen 3H2SO4(aq) + 2Al(s) Al2(SO4)3(aq) + 3H2(g)

  47. Acids + Metals Hydrochloric acid + Zinc  Zinc chloride + Hydrogen 2HCl(aq) + Zn(s) ZnCl2(aq) + H2(g) Sulphuric acid + Zinc  Zinc sulphate + Hydrogen H2SO4(aq) + Zn(s) ZnSO4(aq) + H2(g)

  48. Acids + Metals Hydrochloric acid + Iron  Iron chloride + Hydrogen 6HCl(aq) + 2Fe(s) 2FeCl3(aq) + 3H2(g) Sulphuric acid + Iron  Iron sulphate + Hydrogen 3H2SO4(aq) + 2Fe(s) Fe2(SO4)3(aq) + 3H2(g)

  49. The combustion of hydrogen

  50. The combustion of hydrogen • 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 (-253oC) • has three isotopes, hydrogen, deuterium and tritium

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