Topic 2 Electrolysis. In a nutshell!

Transcription

1 Topic 2 Electrolysis In a nutshell!

2 Indicators: the ph scale Guess the ph

3 Your stomach contains hydrochloric acid. This aids digestion and allows the stomach to kill any harmful bacteria. The acid has a ph of ~2. If to much acidic foods are eaten it can result in indigestion, heartburn and acid reflux. This causes pain in the stomach and in the oesophagus.

4 Alkali s and Bases Metal oxides The Facts! Ammonia An alkali is a base that dissolves in water.

5 Neutralisation acid + base salt + water Sodium hydroxide Hydrochloric acid Na OH H Cl The sodium replaces the hydrogen from HCl Na Cl H 2 O Sodium chloride Water

6 Particle model of Neutralisation Sodium chloride Hydrochloric acid Sodium hydroxide Water

7 What is Titration? The technique of titration is used to find out accurately how much of a chemical substance is dissolved in a given volume of a solution. The technique uses a particular set of apparatus with which volumes of solutions can be measured to an accuracy of greater than 0.1cm 3

8 Titration demonstration B Use the funnel to help you to pour some hydrochloric acid into the burette. Put the beaker under the burette and open the tap to let some acid into the tip of the burette. Close the tap again and pour the acid in the beaker down the sink. C Fill the burette up to just below the zero mark with hydrochloric acid this does not have to be exact. D Use the measuring cylinder to pour exactly 10 cm 3 sodium hydroxide solution into to the conical flask. Add a few drops of indicator and swirl it so that the liquids mix. Stand the flask under the burette. E Read the volume of acid in the burette and write it down. F Add the acid from the burette to the flask a little at a time about 1 cm 3 each time is about right. Swirl the flask gently each time you add some acid. Stop when the indicator turns red and stays red. G Record the volume of acid in the burette. Work out how much acid has been added and write it down. H Wash out the conical flask with water and repeat steps D to G. This time, add acid about 1 cm 3 at a time until you have nearly added the same amount as before. Then add acid in much smaller amounts drop by drop.

9

10 Applying ideas. Copper (II) oxide is a, which reacts with sulfuric acid to make a and. A soluble base is called an. This is called a reaction.

11 The Rules Hydrochloric Acid Chloride Sulfuric Acid Sulfate Nitric Acid Nitrate

12 Must remember. METAL + ACID SALT + HYDROGEN METAL OXIDE + ACID SALT + WATER METAL CARBONATE + ACID SALT + WATER + CARBON DIOXIDE

13 Activity Use the general equation to finish of the word equations below metal + acid a salt + hydrogen magnesium + sulphuric acid magnesium + sulphate hydrogen iron + nitric acid + iron nitrate hydrogen calcium + sulphuric acid + calcium sulphate hydrogen Zinc + hydrochloric + acid zinc chloride hydrogen

14 Metal oxides and acids Metal oxide + acid a salt + water Most metal oxides are not soluble. i.e. They are bases but not alkalis. This means they are often slower and may need heating to make them react Acid Oxide Copper oxide + sulphuric acid copper sulphate + water

15 Reactions of metal oxides with acid A metal oxide is a compound containing a metal and oxide. They are sometimes called BASES. For example: Mg O Na O Magnesium oxide Sodium oxide Aluminium oxide Na Al Al O O O METAL OXIDE + ACID SALT + WATER Mg O H H Cl Cl Cl Mg Cl H O H Copy and complete the following reactions: 1) Magnesium oxide + hydrochloric acid 2) Calcium oxide + hydrochloric acid 3) Sodium oxide + sulphuric acid

16 Base + acid a salt + water Sulfuric acid + copper carbonate When your base is a carbonate, carbon dioxide Gas is also made!!!

17 Task - write out and complete these half equations 1. Hydrochloric acid + sodium hydroxide 2. Hydrochloric acid + calcium carbonate 3. Sulfuric acid + copper oxide 4. Sulfuric acid + iron carbonate 5. Nitric acid + ammonia 6. Phosphoric acid + iron oxide C

18 1. Sodium chloride + water 2. Calcium chloride + water + carbon dioxide 3. Copper sulfate + water 4. Iron sulfate + water + carbon dioxide 5. Ammonium nitrate + water 6. Iron phosphate + water

19 Must remember. METAL + ACID SALT + HYDROGEN METAL OXIDE + SALT SALT + WATER METAL CARBONATE SALT + WATER + CARBON DIOXIDE

20 Complete the worksheet Neutralisation reactions

21 Extracting metals from ores Potassium Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Tin Lead Copper Silver Gold Platinum Metals ABOVE CARBON, because of their high reactivity, are extracted by ELECTROLYSIS and this needs a lot of energy Metals BELOW CARBON are extracted by heating them with carbon in a BLAST FURNACE These LOW REACTIVITY metals blatantly won t need to be extracted because they are SO unreactive you ll find them on their own, not in a metal oxide

22 Fill in the gaps A solid compound cannot conduct. This is because the particle are not free to move. When we dissolve or melt the compound the are free to move so can conduct.

23 Copper sulphate Copper Sulphate solid Copper ion Sulphate on

24 Electrolysis of copper sulphate Positive electrode anode + - Negative electrode cathode electrolyte + - Cation, copper Anion, oxygen

25 Electrolysis Electrolysis is used to extract a HIGHLY REACTIVE metal. When we electrolysed copper sulphate the negative sulphate ions moved to the positive electrode and the positive copper ions moved to the negative electrode OPPOSITES ATTRACT!!! = sulphate ion = copper ion

26 Electrolysis is splitting up substances using electricity (Lysis is latin for splitting)

27 What substances can we split up? We can separate compounds like: Zinc chloride Lead bromide Aluminium oxide Nickel sulphate Write a sentence to explain what electrolysis is and write what will be made during electrolysis of these chemicals. They must be liquids

28 Positive Anode Don t get stressed in the exam: Remember PANIC Negative Is Cathode

29 What happens at the electrodes + e- - At the positive electrode the negative anions are attracted to it. They give electrons to the electrode to lose their charge

30 What happens at the electrodes - e- + At the negative electrode the positve cations are attracted to it. They take electrons from the electrode to lose their charge

31 Electrolysis of HCl ANODE + - CATHODE + - Cl - H + Cl - Cl - H + H + Cl - H +

32 What is Test For Hydrogen? POP! Hydrochloric Acid Magnesium Ribbon

33 What is the test for Chlorine? Damp blue litmus paper Turns Then Pink White

34 Hydrogen chloride and hydrochloric acid have the formula HCl Hydrogen chloride hydrogen + chloride HCl H + Cl HCl H 2 + Cl 2 2HCl H 2 + Cl 2

35 Chlorine gas was used as a weapon in WW1. The gas dissolves the moisture in the eyes and lungs. The cholrine then forms an acid. The victim usually died very painfully within a few days. Chlorine gas in the First World War

36 Can we use chlorine for anything? Paper making PVC

37 Electrolysis of brine When NaCl dissolves in water, it s ions become free to move. So the solution can be electrolysed. In water some of the molecules of water will naturally split apart.

38 Watch the demonstration of the electrolysis of brine Product Test for substance Hydrogen Chlorine Sodium hydroxide Blows out a lighted splint and makes a squeaky pop. Damp litmus paper goes from blue to red then bleaches white Goes dark blue or purple when universal indicator is added.

39 Brine, Sodium Chloride in water Sodium chloride solid sodium ion chlorine on

40 Sodium chloride Sodium chloride solid sodium ion chloride on

41 Electrolysis of NaCl solution

42 Electrolysis of salt 1 Negative ions Salt consists of sodium ions (Na + ) and chloride ions (Cl - ). Chloride ions go to the anode where they lose an electron The neutral chlorine atoms produced join up into pairs Chlorine 2Cl - - gas + 2e- is formed Cl 2

43 Electrolysis of salt 3 Na + ions move to the cathode but do not accept electrons. It is the hydrogen ions that gain electrons As a result hydrogen gas is formed at the cathode. 2H + + 2e- H 2

44 How does the sodium hydroxide form? Sodium chloride solution has four types of ions: Na + and Cl - ions from the sodium chloride H + and OH - ions from the water. The Cl - ions form chlorine at the positive electrode and the H + ions form hydrogen at the negative electrode. So, what s left? Na + and OH - ions are left behind and so a solution of sodium hydroxide (NaOH) is formed. What is the overall equation for the electrolysis of a sodium chloride solution? 2NaCl (aq) + 2H 2 O (l) H 2 (g) + Cl 2 (g) + 2NaOH (aq)

45 Draw what you did Testtube Beaker Power pack

46 Sodium + water sodium + hydrogen + Chlorine Chloride hydroxide NaCl + H 2 O NaOH + H 2 + Cl 2 2NaCl + 2H 2 O 2NaOH + H 2 + Cl 2

47 What would we get if we electrolyse water? How much of each gas would we expect to get? H 2 O

48 Electrolysis of dilute sulfuric acid

49 Uses of hydrogen Rocket fuel

50 Uses of oxygen Welding equipment

51 Compound Anode Cathode Sodium Bromide Potassium Iodide Calcium Fluoride Magnesium Oxide Lithium Chloride

52 O.I.L.R.I.G. Oxidation is loss of electrons Reduction is gain of electrons

53 Half Equations The half equation shows what happens at each electrode e.g. 3Al e - 3Al The aluminium ions collect 3 electrons at the cathode to form aluminium atoms. 2Cl - Cl e - The chlorine ions drop off electrons at the anode to form chlorine atoms (which in turn react to form a chlorine molecule (covalent bonding)

54 Balance the following half equations. 1. Cl 2 + e - Cl - 2. Al 3+ + e - Al 3. H + + e - H 2 4. Pb 2+ + e - 2Pb 5. O e - O 2

55 Watch the animation on page 119 of the active book

56

57

58 The Reactivity Series Potassium Sodium Calcium Magnesium The Reactivity Series lists metals in order of reactivity: Make a pneumonic to help you remember this Aluminium Carbon Zinc Iron Lead Copper Silver Gold

59 The premier league of reactivity Think about it like a football league. If the team near the bottom played the team near the top, who would you expect to win?? It s the same for the reactivity series, if an element near the top of the reactivity series is mixed with an ore from the bottom, the element displaces the less reactive metal from it s ore. V Objective: Explain why we need different methods of extraction depending on the reactivity of the metal

60 Displacement reactions A displacement reaction is one where a MORE REACTIVE metal will DISPLACE a LESS REACTIVE metal from a compound. Magnesium Copper sulphate Mg Cu SO 4 The magnesium DISPLACES the copper from copper sulphate Mg SO 4 Cu Magnesium sulphate Magnesium + copper sulphate Copper magnesium sulphate + copper

61 Watch the dating demo

62 Some example reactions Reaction Prediction Observations Zinc + copper sulphate Zinc + lead nitrate Copper + lead nitrate Copper + silver nitrate Extension work write down the equations for these reactions

63 Zinc Zinc sulphate X Copper sulphate Iron sulphate Magnisium sulphate Copper X Iron X magnesium X

64 4. Aisha placed small samples off four different metals on a spotting tile. She added drops of copper sulphate solution to each metal. copper iron magnesium zinc Aisha repeated the experiment with fresh samples of the four metals and solutions of different salts. She recorded some of her results in a table. shows that a reaction took place X shows that no reaction took place. spotting tile

65 solutions metals copper sulphate copper iron magnesium zinc iron sulphate magnesium sulphate zinc sulphate (a)the four metals have different reactivities. (i) Use the information in the table to put the four metals in a reactivity series. 1 mark (ii) Use the reactivity series to complete the table by writing in or X in the three empty boxes. 2 marks (b) Copper reacts with silver nitrate solution. (i)complete the word equation for the reaction: Copper + silver nitrate marks (ii) Platinum does not react with silver nitrate. Put the metals platinum, copper and silver in the correct order according to their reactivity. (iii) 1 mark (c) In many houses the hot water pipes are made from copper and the boiler is made from iron. Which of these metals will corrode first? Explain your answer. 1 mark Maximum 7 marks

66 Rocks and Minerals Rocks in the earth s crust are very rarely pure substances If the minerals in rocks contain metals then they are called ores. Unreactive metals are found in their pure state and not as an ore. Can you think of any unreactive metals? Metals vary in cost depending on the availability of the ore, how much metal is contained in the ore and the method of extraction. You are going to try one method of extraction today. Objective: Recall that most metals are extracted as ores from the earth s crust

67 Match the metal ore to the picture haematite(iron) cinnabar(mercury) malachite(copper) bauxite(aluminium) Galena(lead) Starter Activity

68 Match the metal ore to the picture cinnabar(mercury) haematite(iron) bauxite(aluminium) malachite(copper) Galena(lead) Answers

69 Which method of extraction? There are two main methods for the extraction of metals from their ores: Chemical reduction (like you ve carried out in class) Electroysis The method used for extraction depends on the reactivity of the metal you are trying to extract. Scientists use something called the reactivity series to help them decide which method to use, the higher the metal, the more reactive it is. Objective: Explain why we need different methods of extraction depending on the reactivity of the metal

70 Extracting metals from ores Potassium Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Tin Lead Copper Silver Gold Platinum Metals ABOVE CARBON, because of their high reactivity, are extracted by ELECTROLYSIS and this needs a lot of energy Metals BELOW CARBON are extracted by heating them with carbon in a BLAST FURNACE These LOW REACTIVITY metals blatantly won t need to be extracted because they are SO unreactive you ll find them on their own, not in a metal oxide

71

72

73 Oxidation Burning A reaction where the products contain more oxygen than the original fuel. Adding more oxygen to a substance is called oxidation. Gloss paint Iron nails

74 Oxidation Fuels like carbon when burnt are oxidised. C CO 2 S SO 2 It has had oxygen added to it. Iron has been oxidised to make iron oxide (rust).

75 Reduction Removing oxygen from a chemical is called reduction. It reduces the amount of oxygen in the substance.

76 Reduction Reduction can occur when burning iron oxide. 3C + 2Fe 2 O 3 4Fe + 3Co 2 The iron oxide gives its oxygen to the carbon. What can we say has happened to the Carbon?

77 Redox reactions Although some fuels gain oxygen some will also lose oxygen. Thus oxidation and reduction are happening at the same time. These are known as REDOX reactions. Write a word equation for methane (CH 4 ) reacting with oxygen. State what where redox is occurring.

78 NaOCl NaCl CH 3 CH 2 OH CH 3 COOH Oxidation or Reduction? How do we know?

79

80

81 What are the properties of metals

82 What are the properties of nonmetals

83 Comparing the electrical conductivity of metals: Are all metals good conductors? Are some better than others? How could we investigate this? Monday, 16 March 2015

84 What are the advantages of recycling metals? Make metals last longer Less energy needed to recycle compared to extracting ore Reduces the need to mine ores less damage to the environment Less pollution produced Less waste materials put into landfill Monday, 16 March 2015

85 What are the disadvantages of recycling Costs and energy involved in: metals? Collecting materials Sorting materials Transporting materials Can be more expensive to recycle Monday, 16 March 2015

86 Monday, 16 March 2015 The graph

87

88

89 What is an alloy? An alloy is a mixture of a metal with at least one other element. Steel is a common example of an alloy. It contains iron mixed with carbon and other elements. Adding other elements to a metal changes its structure and so changes its properties. The final alloy may have very different properties to the original metal.

90 What types of alloys are there? Other well-known alloys include: brass: an alloy of copper and zinc. solder: an alloy of tin and lead amalgam: an alloy of mercury and silver or tin.

91 What is steel? Steel is an alloy of iron and other elements, including carbon, nickel and chromium. Steel is stronger than pure iron and can be used for everything from sauce pans to suspension bridges!

92 Why is steel stronger than iron? The atoms in pure iron are arranged in densely-packed layers. These layers can slide over each other. This makes pure iron a very soft material. When other elements are added to iron, their atoms distort the regular structure of the iron atoms. It is more difficult for the layers of iron atoms in steel to slide over each other and so this alloy is stronger than pure iron.

93

94