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This is a summary of the topic "metals" in the GCE O levels subject: Chemistry. Students taking either the combined science (chemistry/physics) or pure chemistry will find this useful. These slides are prepared according to the learning outcomes required by the examinations board.

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  1. 1. METALS
  3. 3. PROPERTIES OF METALS Metals Generally solids at room temperature Good conductors of heat and electricity Malleable and ductile High melting and boiling points Shiny in appearance Malleable: Able to be hammered or pressed into shape without breaking or cracking. Ductile: Able to be drawn out into a thin wire.
  4. 4. PROPERTIES OF METALS Metals exist as a lattice of positive ions in a ‘sea of electrons’. This ‘sea of electrons’ are mobile and able to carry electrical charges thus giving metals the ability to conduct electricity.
  5. 5. ALLOYS
  6. 6. ALLOYS An alloy is a mixture of a metal with another element. Brass is a mixture of copper and zinc. Stainless steel is a mixture of carbon and chromium. Alloys are stronger than pure metals. This is because the different sizes of atoms in an alloy disrupt the orderly arrangement of atoms which means it harder for the atoms to slide past each other.
  7. 7. ALLOYS
  9. 9. REACTIVITY SERIES Metals are arranged from the most reactive at the top to the least reactive at the bottom. They are determined by their reaction with cold water or steam, as well as their reaction with hydrochloric acid.
  10. 10. REACTIVITY SERIES Metal Symbol Potassium K Sodium Na Calcium Ca Magnesium Mg Aluminium Al Zinc Zn Iron Fe Tin Sn Lead Pb *Hydrogen H Copper Cu Silver Ag Gold Au Most reactive Least reactive Mnemonics Popular Scientist Can Make A Zebra In The Lab He Cannot See Girls
  11. 11. Metal(s) Observations and equation for reaction with cold water Observations and equation for reaction with steam Potassium Reacts very violently, forms potassium hydroxide and hydrogen gas. 2K(s) + 2H2O(l)  2KOH(aq) + H2(g) React explosively. Never do this in the lab! Sodium Reacts violently, forms sodium hydroxide and hydrogen gas. 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g) Calcium Reacts readily, forms calcium hydroxide and hydrogen gas. Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g) Magnesium Reacts very slowly with cold water to form magnesium hydroxide and hydrogen gas. Mg(s) + 2H2O(l)  Mg(OH)2(aq) + H2(g) Hot Magnesium reacts violently with steam to form magnesium oxide (white solid) and hydrogen gas. A bright white glow is produced during the reaction. Mg(s) + H2O(g)  MgO(s) + H2(g) Zinc No reaction occurs. Hot Zinc reacts readily with steam to form zinc oxide and hydrogen gas. Zinc oxide is yellow when hot and white when cold. Zn(s) + H2O(g)  ZnO(s) + H2(g) Iron No reaction occurs. Hot Iron reacts slowly with steam to form iron oxide and hydrogen gas. The iron must be heated constantly for the reaction to proceed. 3Fe(s) + 4H2O(g)  Fe3O4(s) + 4H2(g) Lead Copper Silver No reaction occurs. No reaction occurs.
  12. 12. REACTIVITY SERIES If the metals do react with cold water, General formula: Metal + water  metal hydroxide + hydrogen If the metals do not react with cold water but with steam, General formula: Metal + steam  metal oxide + hydrogen
  13. 13. Metal(s) Observation for reaction with dilute hydrochloric acid Equation Potassium React explosively. Never do this in the lab. 2K(s) + 2HCl(aq)  2KCl(aq) + H2(g) Sodium 2Na(s) + 2HCl(aq)  2NaCl(aq) + H2(g) Calcium Reacts violently to give hydrogen gas. Ca(s) + HCl(aq)  CaCl2(aq) + H2(g) Magnesium Reacts rapidly to give hydrogen gas. Mg(s) + HCl(aq)  MgCl2(aq) + H2(g) Zinc Reacts moderately fast to give hydrogen gas. Zn(s) + HCl(aq)  ZnCl2(aq) + H2(g) Iron Reacts slowly to give hydrogen gas. Fe(s) + HCl(aq)  FeCl2(aq) + H2(g) Lead Copper Silver No reaction occurs. NA
  15. 15. EXTRACTION OF METALS Most metals occur as compounds called ores found in the ground. These ores are rarely in the pure state. We are only interested in the metal elements and will obtain them using a process called extraction. The method used for extracting a metal from its ores depends on the position of the metal in the reactivity series (Reactivity decreases down the series).
  16. 16. EXTRACTION OF METALS Metal Method of Extraction Potassium Sodium Calcium Magnesium Aluminium Electrolysis of their molten compounds. Zinc Iron Tin Lead (Hydrogen) Copper Using coke (carbon) to reduce the metal oxide to obtain the metal. Silver Gold Found as free metallic elements
  17. 17. EXTRACTION OF METALS Metals at the top of the series are very reactive and form very stable compounds. Electrolysis is used to obtain the metal. Electrolysis is a costly method to extract metals as large amounts of electricity is needed.
  18. 18. EXTRACTION OF METALS Metals in the middle of the reactivity series usually exist as oxides or sulfates. The are extracted using reduction with carbon. The carbon will remove the oxygen from the metal oxides to give the pure metal. Carbon (Coke) is used cause it is cheap.
  19. 19. EXTRACTION OF METALS Metals at the bottom of the reactivity series are least reactive. They are found “native” in nature. For example Gold.
  21. 21. EXTRACTION OF IRON Iron is extracted from its ore called haematite or Iron (III) Oxide, Fe2O3 by reduction with coke in a blast furnace. This process removes oxygen from the haematite to give us the pure iron. We need limestone, coke and air for this process.
  23. 23. EXTRACTION OF IRON Coke: Carbon supply for reduction. Hot air: Oxygen supply. Limestone (Calcium Carbonate, CaCO3): To remove acidic impurities (mainly sand) in the iron ore.
  24. 24. EXTRACTION OF IRON Haematite, limestone and coke are added from the top. Hot air blasted in from the bottom. Haematite is reduced while all other impurities removed by the Calcium oxide (which comes from limestone).
  25. 25. EXTRACTION OF IRON Carbon reacts with oxygen from hot air to form carbon dioxide. C(s) + O2(g)  CO2(g) Carbon dioxide reacts with more carbon to form carbon monoxide. CO2(g) + C(s)  2CO(g)
  26. 26. EXTRACTION OF IRON Carbon monoxide reduces haematite to iron. Fe2O3(s) + 3CO(g)  2Fe(l) + 3CO2(g) Molten iron runs to the bottom of the furnace.
  27. 27. EXTRACTION OF IRON Limestone decomposes to produce carbon dioxide and calcium oxide. CaCO3(s)  CaO(s) + CO2(g) Calcium oxide is a basic oxide which reacts with acidic silicon dioxide to form calcium silicate. CaO(s) + SiO2(s)  CaSiO3(l) Calcium silicate floats on molten iron.
  28. 28. EXTRACTION OF IRON Calcium silicate, which is also known as slag, is used to make roads. The molten iron obtained is further purified and processed.
  29. 29. RUSTING
  30. 30. RUSTING Rusting or corrosion of iron is the process that produces rust. Iron + oxygen + water  hydrated iron (III) oxide 4Fe(s) + 3O2(g) + 2xH20(l)  2Fe2O3.xH2O(s)
  31. 31. RUSTING Conditions for rusting: Both oxygen (air) and water are required.
  32. 32. RUST PREVENTION Electroplating: Coating iron with an layer of unreactive metal. For example, platinum plating. Painting or greasing: Applying paint or oil on the iron’s surface. Sacrificial protection: A more reactive metal is attached to iron, and will be corroded first before iron.
  34. 34. RECYCLING METALS Why we need to recycle metals? Metals are finite resources, which means they will run out some day. How are metals recycled? Metals are recovered from scrap metal. Lead is recovered from car batteries. Aluminum is recovered from food and drinks cans.
  35. 35. RECYCLING METALS Advantages of recycling metals. Conservation: Recycling helps to conserve natural resources. Recycling reduces environmental problems. We can minimize mining and reduce destruction of huge amounts of land. Economic: Recycling might be cheaper than metal extraction from ores.
  36. 36. RECYCLING METALS However, there are problems related to recycling. Economic: It can be expensive to separate metals from waste and incur transportation costs, making recycling more expensive than extracting metals. Social issues: Time and resources are spent to educate the public and a long time before people will adopt a recycling lifestyle.
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