NATIONAL 5 CHEMISTRY UNIT 3 CHEMISTRY IN SOCIETY

1. NATIONAL 5 CHEMISTRY UNIT 3 CHEMISTRY IN SOCIETY

2. CONTENTS • Metals • Properties of plastics • Fertilisers • Nuclear Chemistry • Chemical Analysis

3. Metallic bonding and properties (they conduct electricity) • Reactivity of metals with oxygen, acid and water (balanced ionic equations can be written). • Metal ores and percentage (%) composition • Extracting metals by heat alone, heating with carbon and electrolysis. [depends on reactivity of metal and all metals are reduced in the process]. (balanced ionic equations can be written and reduction equations for the metals can be written.) • Electrochemical cells and REDOX equations • Fuel Cells and Rechargeable batteries METALS

4. Metallic bonding • Positive metal ions surrounded by de-localised electrons (electrons that are free to move). This is why metals have the properties they have

5. Properties of Metals • Density – this is the mass of a substance in a given volume. A high density material is much heavier than the same volume of a low density material e.g. aluminium (low density) – used to build aircraft. Lead (high density) – is used as weights for fishing nets/lines. • Thermal Conductivity - metals all conduct heat well • because of the close contact of the atoms. • E.g. pots/pans. • Electrical Conductivity - metals all conduct electricity when solid and when molten because electrons can travel easily through the structure. • E.g. cables

6. Malleability - metals can be beaten into different shapes. • E.g. jewellery. • Strength - most metals are strong because of the metallic bond which holds the atoms together. • E.g. bridges, cars, buildings etc. • Recycling Metals - Metals need to be recycled because they will not last forever (they are finite resources).

7. Alloys • The properties of metals can be extended or altered by mixing them with other metals or with non-metals. • Iron can be changed into stainless steel by mixing it with small amounts of chromium. This stops the metal rusting.

10. Reactions of Metals

11. METAL REACTIVITY The reactions of metals that we will cover are; • reaction with oxygen metal + oxygen  metal oxide • reaction with water metal + water  metal hydroxide + hydrogen • reaction with dilute acid Metal + acid  salt + hydrogen

12. Reactivity Series • Metals have similar chemical properties. However, some metals are more reactive than others. • Based on their reactivity, chemists produced a ‘league table’ of metals as shown below.

13. most reactive least reactive

14. Metals Reacting with Oxygen • All metals above silver in the reactivity series react with oxygen when heated to form a metal oxide. • The higher the metal in the reactivity series the more vigorous the reaction with oxygen.

15. METAL + OXYGEN METAL OXIDE e.g. magnesium + oxygen  magnesium oxide 2Mg(s) + O2(g)  2MgO(s) • Potassium, sodium and lithium are so reactive they are stored under oil to prevent them from reacting with the oxygen and water in the air.

16. Oxygen can be made by heating potassium permanganate in a test tube and allowing the gas to pass through the preheated metal. Metal + Oxygen  Metal oxide  E.g. Magnesium + Oxygen  Magnesium oxide Mg + O2 MgO

17. METAL + WATER METAL HYDROXIDE + HYDROGEN Metals Reacting with Water • All metals above aluminium in the reactivity series react with water to produce the metal hydroxide and hydrogen gas: e.g. sodium + water  sodium hydroxide + hydrogen Na(s) + H2O(l)  NaOH(aq) + H2(g)

18. METAL + ACID SALT +HYDROGEN Metals Reacting with Acids • All metals above copper in the reactivity series react with dilute acids such as hydrochloric and sulphuric acid to produce a salt and hydrogen gas:

19. e.g. zinc + hydrochloric acid zinc chloride + hydrogen Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)

21. When a metal reacts with an acid it produces bubbles of hydrogen gas. • Generally, the faster the bubbles are produced, the more reactive the metal. • Aluminium is the exception to this. It reacts very slowly for the first 20 mins, after which it reacts quickly. • The reason for this is that the metal is protected by a thin layer of aluminium oxide, which must first be removed by the acid.

22. Summary

23. Metal Ores • Ores are naturally-occuring compounds of metals from which metals can be extracted. • The three main types of ore are metal carbonates, metal oxide and metal sulphides.

24. Common Ores

25. Percentage Composition

26. Extracting Metals • Metals such as gold and silver occur uncombined on earth because they are unreactive and because of this these elements were among the first to be discovered. • Other metals, such as those in the table are found in compounds and have to be extracted (which is an example of reduction).

27. Extraction of Metals from Ores • The method used to extract a metal depends on the reactivity of the metal. • The more reactive the metal, the more difficult it is to extract. • The less reactive the metal, the easier it is to extract.

28. Methods of extractiona) Heating metal oxides Silver oxide  Silver + Oxygen Ag2O  Ag + O2 • Few metals can be obtained in this way.

29. b) Heating Metal Oxides with Carbon Metal oxide + Carbon  Metal + Carbon dioxide E.g. Iron oxide + Carbon  Iron + Carbon dioxide Fe2O3 + C  Fe + CO2 • This method is used to extract metals below aluminium in the reactivity series.

30. c) Using Electricity • Electricity can be used to split ionic compounds into their elements in a process called electrolysis. • The method is used to extract reactive metals above zinc in the reactivity series. • A large electric current is passed through the molten compound, and metal appears at the negative electrode.

31. Electrolysis • http://www.youtube.com/watch?v=i9xS9t-KMpc – electrolysis explained

32. Electrochemical Series Potassium Sodium Calcium Magnesium Aluminium Zinc Iron Nickel Tin Lead Copper Mercury Silver Gold Must be electrolysed to release metal from ore Separated from ore by heating with CHARCOAL, thus releasing CARBON DIOXIDE Can be broken by heat alone

33. Batteries and Cells • We generate electricity from burning fossil fuels, harnessing the power of water (hydroelectric), or nuclear energy. • But, we also need electricity for personal stereos, mobile phones etc. • We use batteries. Chemical reactions in a battery produce electricity

34. Electricity is a flow of electrons. • When electricity is produced in a battery, electrons flow from the battery, through the wires, to the device to which it is connected. • In most batteries the electrons come from a layer of zinc metal.

35. Dry-cell Battery • Zinc cup forms the negative terminal of the battery and the carbon rod is the positive terminal. • Between the two terminals is a paste of ammonium chloride. This completes the circuit by allowing ions to flow through it –acts as an electrolyte. An electrolyte is a substance that will conduct electricity when dissolved in water or melted. This is due to the movement of ions.

36. Some batteries are re-chargable. The chemicals can be restored by giving the battery a supply of electrons. • e.g. a car battery contains lead metal. When the battery is being used the lead metal atoms turn into ions. During recharging, the ions are turned back into lead atoms.

37. Simple Cells • Electricity can be produced by connecting different metals together, with an electrolyte, to form a simple cell. • In the cell shown above, electrons flow from the zinc to the copper. The sodium chloride solution acts as an electrolyte and completes the circuit.

38. A voltmeter measures the voltage produced and it is seen that different voltages are obtained when different metals are used. • The voltage between different pairs of metals varies and this leads to the Electrochemical Series (ECS) (page 7 of Data Booklet).

39. metal A V filter paper soaked in a sodium chloride solution metal B • When two different metals are joined together, electrons flow through the wire from the metal higher in the ECS series to the metal lower in the series e.g. from lithium to silver. • The further apart the metals are in the ECS, the higher the voltage produced. • The closer together the metals are in the ECS, the lower the voltage produced.

40. Displacement Reactions • Displacement reactions occur when a metal is added to a solution containing ions of a metal lower in the electrochemical series. Example • If zinc metal is added to a solution of copper(II) sulphate, the zinc slowly becomes smaller and a brown solid covers it. At the same time the blue copper(II) sulphate solution loses its colour.

41. Why does this happen? • The zinc atoms have LOST electrons and turned into zinc ions, which go into solution. • The copper ions GAIN the electrons lost by the zinc and turns into copper metal atoms.

42. This is called a DISPLACEMENT REACTION and the overall reaction can be represented by; DISPLACEMENT REACTION: Formation of a metal from a solution containing its own ions when a metal higher than itself in the electrochemical series is added to it.

43. As a general rule, a metal will displace a metal lower than itself in the ECS. • e.g. • - iron would displace silver ions from a solution of silver nitrate as iron is above silver in the ECS. • - lead would not displace tin ions from a solution of tin chloride as lead is lower than tin in the ECS.

44. Hydrogen in the ECS • Hydrogen and other non-metals are also in the ECS. • Hydrogen can be placed in the ECS by considering the reactions of metals with dilute acids.

45. Metals down to lead in the ECS react with dilute acids to produce hydrogen gas, i.e. they displace hydrogen ions from acids. • Copper, silver, gold and platinum do not react with dilute acids. • So hydrogen can be placed below lead but above copper in the ECS.

46. Half-Cells • Cells can also be set up by connecting two-half cells together. • A half-cell consists of a metal in contact with a solution of its own ions, such as a strip of copper metal in a beaker of copper(II) sulphate solution.

47. Electricity is produced when two half-cells containing different metals are connected as shown:

48. The metals are joined by wires (electrons flow) and the two solutions are connected using an ion bridge (ions flow). Filter paper soaked in sodium chloride solution is often used. • The ion bridge completes the circuit and allows ions to move across it. If it is removed, the circuit will be broken and no electricity will be produced.

49. In the cell shown the zinc atoms lose electronsand form zinc ions, while the copper(II) ions gain electrons to form atoms of copper metal.

50. The ion-electron equations to represent these reactions are; • Zinc metal would turn into zinc ions and the copper(II) ions would decrease until the cell would stop producing electricity.