Chapter 11. Reactivity of metals

Transcription

1 Chapter 11 Reactivity of metals 11.1 Comparing reactivities of common metals 11.2 The metal reactivity series 11.3 Chemical equations 11.4 Metal reactivity series and the tendency of metals to form positive ions 11.5 Displacement reactions of metals P. 1 / 72

2 11.6 Ionic equations 11.7 Relation between the extraction method and the metal reactivity series Key terms Progress check Summary Concept map P. 2 / 72

3 11.1 Comparing reactivities of common metals Different reactivities of metals Some metals are more reactive than others. Figure 11.1 Potassium reacts vigorously with water. Figure 11.2 Gold is a very unreactive metal. It does not react with air and water. P. 3 / 72

4 Learning tip Potassium metal catches fire easily when placed in oxygen or when added to water. Hence, it is flammable. Reactivity of a metal refers to the readiness of it to react with other substances. Three factors are considered when comparing the reactivities of metals: 1.The lowest temperature at which the reaction starts 2.The rate of reaction 3.The amount of heat given out during reaction 11.1 Comparing reactivities of common metals P. 4 / 72

5 Metal Conditions for reaction Observation Word equation Potassium gentle heating It burns vigorously with a lilac (pale purple) flame to produce a white powder. potassium + oxygen potassium oxide Sodium BURN gentle heating It burns vigorously with a golden yellow flame to produce a white powder. sodium + oxygen sodium oxide Calcium strong heating It burns quite vigorously with a brick-red flame to produce a white powder. Table 11.1 Reactions of some common metals with air Comparing reactivities of common metals calcium + oxygen calcium oxide P. 5 / 72

6 Metal Magnesium Aluminium Zinc Conditions for reaction BURN strong heating strong heating strong heating Observation It burns with a very bright white flame to produce a white powder. Aluminium powder burns to give out much heat; a white powder forms. Zinc powder burns to give out some heat; a powder (yellow when hot, white when cold) forms. Table 11.1 Reactions of some common metals with air Comparing reactivities of common metals P. 6 / 72 Word equation magnesium + oxygen magnesium oxide aluminium + oxygen aluminium oxide zinc + oxygen zinc oxide

7 Metal Iron Lead Copper Mercury Conditions for reaction BURN DO NOT BURN strong heating strong heating very strong heating very strong heating Observation Iron powder burns with yellow sparks to produce a black solid. It melts; a powder (orange when hot, yellow when cold) is seen on the surface. Its surface turns black. A red powder forms on the surface. Table 11.1 Reactions of some common metals with air Comparing reactivities of common metals P. 7 / 72 Word equation iron + oxygen iron(ii, III) oxide lead + oxygen lead(ii) oxide copper + oxygen copper(ii) oxide mercury + oxygen mercury(ii) oxide

8 Metal Conditions for reaction Observation Word equation Silver Platinum Gold NO REACTION No observable change even on very strong heating. No observable change even on very strong heating. No observable change even on very strong heating. Table 11.1 Reactions of some common metals with air Comparing reactivities of common metals P. 8 / 72

9 (a) Potassium burns with a lilac flame. (b) Sodium burns with a golden yellow flame. (c) Calcium burns with a brick-red flame. (d) Magnesium burns with a very bright white flame. (e) Iron burns with yellow sparks Figure 11.3 Different metals burn in air to give different flame colours Comparing reactivities of common metals P. 9 / 72

10 Reactivity of metals towards air decreases most reactive least reactive Metal Reaction with air Word equation Potassium, K Sodium, Na Calcium, Ca Magnesium, Mg Aluminium, Al Zinc, Zn Iron, Fe Lead, Pb Copper, Cu Mercury, Hg Silver, Ag Platinum, Pt Gold, Au react and burn in air (oxygen) react but do not burn in air (oxygen) no reaction with air (oxygen) P. 10 / 72 metal + oxygen metal oxide Table 11.2 The reactivity of common metals towards air (oxygen) Comparing reactivities of common metals

11 Appearance of metals and storage methods All metals look shiny when they are freshly cut. But the shiny surface of very reactive metals soon becomes dull when exposed to air. Metals react with oxygen in air, forming an oxide layer on the metal surface. Very reactive metals (such as potassium and sodium) are stored under paraffin oil Comparing reactivities of common metals P. 11 / 72

12 Calcium, which is a quite reactive metal, is kept in an airtight container. Figure 11.4 Sodium is stored under paraffin oil while calcium is stored in an airtight container. Gold is the least reactive of all metals. It does not react with oxygen and is always shiny and attractive in appearance Comparing reactivities of common metals Class practice 11.1 P. 12 / 72

13 Action of water on potassium Melts to form a silvery ball Moves about very quickly on the water surface with a hissing sound Burns with a lilac flame potassium + water potassium hydroxide + hydrogen The resultant solution is alkaline. potassium hydroxide is produced which turns red litmus paper blue Comparing reactivities of common metals P. 13 / 72

14 Learning tip A large amount of heat is produced when potassium reacts with water. The heat causes potassium to melt. Action of water on sodium Less vigorously than potassium Melts to form a silvery ball. Moves about quickly on the water surface Burns with a golden yellow flame 11.1 Comparing reactivities of common metals P. 14 / 72

15 sodium + water sodium hydroxide + hydrogen The resultant solution is alkaline. sodium hydroxide is produced. Figure 11.5 Potassium reacts vigorously with water. Figure 11.6 Sodium reacts with water less vigorously than potassium Comparing reactivities of common metals P. 15 / 72

16 Action of water on calcium When we add a calcium granule to cold water, it sinks to the bottom. Colourless gas bubbles form at a moderate rate. The gas can be collected using an inverted funnel. When the gas is tested with a burning splint, it burns with a pop sound. calcium + water calcium hydroxide + hydrogen 11.1 Comparing reactivities of common metals P. 16 / 72

17 hydrogen water inverted funnel calcium granule Figure 11.7 Calcium reacts with water at a moderate rate. Calcium granule gradually decreases in size and eventually disappears. A milky suspension (calcium hydroxide solid) is produced, which is only slightly soluble in water Comparing reactivities of common metals Skill corner 11.1 P. 17 / 72

18 Action of steam on magnesium Magnesium has almost no reaction with cold water. It reacts slowly with hot water to give magnesium hydroxide (only slightly soluble in water) and hydrogen. magnesium + water magnesium hydroxide + hydrogen Magnesium reacts more vigorously with steam Comparing reactivities of common metals P. 18 / 72

19 wet sand magnesium ribbon delivery tube hydrogen water heat trough Figure 11.8 The reaction of magnesium with steam Comparing reactivities of common metals P. 19 / 72

20 With strong heating, the water in the wet sand changes into steam. Steam then reacts with magnesium to give an intense white light. A white solid, magnesium oxide, forms. magnesium + steam magnesium oxide + hydrogen 11.1 Comparing reactivities of common metals P. 20 / 72

21 Action of steam on aluminium, zinc and iron Zinc and iron do not react with cold or hot water. They react with steam. The reaction is less vigorous for zinc, and even less for iron. zinc + steam zinc oxide + hydrogen iron + steam iron(ii, III) oxide + hydrogen 11.1 Comparing reactivities of common metals P. 21 / 72

22 Learning tip Iron(II, III) oxide refers to a mixture of iron(ii) oxide and iron(iii) oxide. Aluminium metal is usually covered with a very thin layer of aluminium oxide which protects the metal from reaction. If the oxide layer is removed, the aluminium obtained would be more reactive than zinc. aluminium + steam aluminium oxide + hydrogen (after removing the oxide layer) 11.1 Comparing reactivities of common metals P. 22 / 72

23 Lead, copper, mercury, silver, platinum and gold, even if heated strongly, do not react with steam. Learning tip Remember that magnesium can react with hot water and steam Comparing reactivities of common metals P. 23 / 72

24 Reactivity of metals towards water or steam decreases most reactive least reactive Metal Potassium, K Sodium, Na Calcium, Ca Magnesium, Mg Aluminium, Al Zinc, Zn Iron, Fe Lead, Pb Copper, Cu Mercury, Hg Silver, Ag Platinum, Pt Gold, Au Table 11.3 The reactivity of common metals towards water or steam Comparing reactivities of common metals Reaction with water or steam metals react with cold water heated metals react with steam heated metals do not react with water or steam Class practice 11.2 P. 24 / 72 Equation metal + water metal hydroxide + hydrogen metal + steam metal oxide + hydrogen

25 Reactions of metals with dilute hydrochloric acid and dilute sulphuric acid When we add a magnesium ribbon to dilute hydrochloric acid, magnesium dissolves and many colourless gas bubbles are given out. The test tube quickly becomes warm as heat is given out. magnesium + dilute hydrochloric acid magnesium chloride + hydrogen 11.1 Comparing reactivities of common metals P. 25 / 72

26 Similar observations can be made when we add a magnesium ribbon to dilute sulphuric acid. magnesium + dilute sulphuric acid magnesium sulphate + hydrogen iron magnesium dilute dilute hydrochloric hydrochloric acid acid (a) Magnesium reacts vigorously (b) Iron reacts slowly with with dilute hydrochloric acid. dilute hydrochloric acid. Figure Comparing reactivities of common metals P. 26 / 72

27 copper dilute hydrochloric acid (c) There is no observable change when copper is added to dilute hydrochloric acid. Figure 11.9 (a) Magnesium, (b) iron and (c) copper react differently with dilute hydrochloric acid Comparing reactivities of common metals P. 27 / 72

28 Metal Reaction with dilute acid Equation Reactivity of metals towards dilute hydrochloric/ sulphuric acid decreases most reactive Potassium, K Sodium, Na Calcium, Ca Magnesium, Mg Aluminium, Al Zinc, Zn Iron, Fe Lead, Pb Copper, Cu Mercury, Hg explosive reaction reacts with dilute acid, more slowly down the series very slow reaction metal + dilute hydrochloric acid metal chloride + hydrogen or metal + dilute sulphuric acid metal sulphate + hydrogen Silver, Ag no reaction Platinum, Pt least reactive Gold, Au Table 11.4 The reactivity of common metals towards dilute hydrochloric acid or dilute sulphuric acid Comparing reactivities of common metals Experiment 11.1 Class practice 11.3 P. 28 / 72 Experiment 11.1

29 11.2 The metal reactivity series Comparing their reactions with air, water and dilute acids arrange common metals in order of reactivity Metal reactivity series Metals at the top of the series are the most reactive. Metals at the bottom are the least reactive. P. 29 / 72

30 Potassium, K Sodium, Na Calcium, Ca Platinum, Pt Figure Metal reactivity series for Gold, Au least reactive common metals. Example 11.1 Class practice The metal reactivity series Magnesium, Mg Aluminium, Al Zinc, Zn Iron, Fe Lead, Pb Copper, Cu Mercury, Hg Silver, Ag most reactive P. 30 / 72 decreasing reactivity

31 11.3 Chemical equations What is a chemical equation? When magnesium burns in air (or oxygen), magnesium oxide is produced. magnesium + oxygen magnesium oxide magnesium oxide magnesium Figure Magnesium burns in air to form magnesium oxide. P. 31 / 72

32 In the reaction between magnesium and oxygen, two magnesium atoms react with one oxygen molecule to form two formula units of magnesium oxide. OO 1 oxygen molecule Mg Mg 2 magnesium atoms Mg 2+ O 2 Mg 2+ O 2 2 formula units of magnesium oxide Figure A simple diagram showing what happens to the particles during the reaction between magnesium and oxygen Chemical equations P. 32 / 72

33 Learning tip One formula unit of magnesium oxide consists of one magnesium ion (Mg 2+ ) and one oxide ion (O 2 ). Chemists usually use chemical equations to represent reactions instead. Chemical equation for this reaction is written as: 2Mg(s) + O 2 (g) 2MgO(s) 11.3 Chemical equations Think about P. 33 / 72

34 What does a chemical equation tell us? Example 2Mg(s) + O 2 (g) 2MgO(s) 1. The reactants involved 2. The products formed 3. Physical states of the substances involved 4. The relative number of particles (i.e. atoms, molecules, ions or formula units) These are magnesium (Mg) and oxygen (O 2 ), written on the left-hand side of the arrow. This is magnesium oxide (MgO), written on the right-hand side of the arrow. Mg and MgO are solids, represented by a state symbol (s); O 2 is a gas (g). Other state symbols are: liquid (l) and aqueous solution (aq). 2 atoms of Mg would react with 1 molecule of O 2 to produce 2 formula units of MgO. Table 11.5 The information that a chemical equation contains Chemical equations P. 34 / 72

35 Single arrow between the two sides of an equation indicates that the reaction goes one way only. Double arrow is also used, e.g. N 2 (g) + 3H 2 (g) 2NH 3 (g) The means that the reaction is reversible i.e. both forward and backward reactions occur at the same time Chemical equations Class practice 11.5 P. 35 / 72

36 Balancing a chemical equation Example 2Mg(s) + O 2 (g) 2MgO(s) 2Mg(s) + O 2 (g) 2MgO(s) left-hand side 2 Mg atoms 2 O atoms right-hand side 2 Mg atoms 2 O atoms The numbers of magnesium atoms and oxygen atoms are the same on both sides of the equation. The equation is balanced Chemical equations P. 36 / 72

37 Learning tip Atoms cannot be created or destroyed in a reaction. They just rearrange to give new substances. The numbers in front of the formulae of reactants and products in a balanced chemical equation are called stoichiometric coefficients Chemical equations P. 37 / 72

38 Steps in writing a chemical equation 1. Stoichiometric coefficients must be placed in front of the formulae where necessary. The formulae themselves must not be changed. 2. The stoichiometric coefficient in front of a chemical formula is different from the subscript in a chemical formula Chemical equations Problem-solving strategy 11.1 P. 38 / 72

39 stoichiometric coefficient (affects both H and O) 4H 2 O subscript (affects only H) The coefficient 4 means that there are 4 water molecules. There are totally 8 hydrogen atoms and 4 oxygen atoms. The subscript 2 means that there are 2 hydrogen atoms in a water molecule Chemical equations P. 39 / 72

40 3. Some equations involve polyatomic ions (e.g. SO 4 2, NO 3, OH ). When balancing such equations, we should consider the polyatomic ions as a single unit. For example, if there are two SO 4 2 ions on the reactant side of the equation, there should be two SO 4 2 ions on the product side. Example Chemical equations Class practice 11.6 P. 40 / 72

41 11.4 Metal reactivity series and the tendency of metals to form positive ions Metals react by losing electrons to form positive ions Each magnesium atom loses two outermost shell electrons magnesium ion (Mg 2+ ) Electrons lost from magnesium atoms are gained by oxygen atoms oxide ions (O 2 ) Figure Electron diagrams showing the formation of magnesium oxide from the reaction between magnesium and oxygen. P. 41 / 72

42 Learning tip A magnesium atom loses two outermost shell electrons in order to get the electronic arrangement of a noble gas atom. Key point Metals react by losing electrons to form. positive ions 11.4 Metal reactivity series and the tendency of metals to form positive ions P. 42 / 72

43 Reactivity and readiness to lose electrons The more readily the metal loses electrons, the more reactive is the metal. The readiness of elements to lose electrons decreases as we go across a period. E.g. Across Period 3, reactivity of metals: Na > Mg > Al The readiness of elements to lose electrons increases down a group. E.g. Reactivity of metals: Li < Na < K (Group I) Be < Mg < Ca (Group II) 11.4 Metal reactivity series and the tendency of metals to form positive ions P. 43 / 72

44 increasing readiness to lose electrons increasing reactivity of metals increasing readiness to lose electrons increasing reactivity of metals Figure Readiness of metals to lose electrons (and hence reactivity of metals) decreases across a period and increases down a group Metal reactivity series and the tendency of metals to form positive ions P. 44 / 72

45 From the above reasoning, the order of reactivity of some metals can be explained: K > Na, Ca > Mg > Al... Key point A metal higher in the reactivity series has a higher reactivity, and its atoms would lose electrons to form positive ions more easily Metal reactivity series and the tendency of metals to form positive ions P. 45 / 72

46 11.5 Displacement reactions of metals Copper in silver nitrate solution When we place copper in silver nitrate solution, the copper slowly dissolves. Some shiny silver deposits form on the copper surface. The solution gradually turns pale blue. Cu(s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag(s) colourless pale blue P. 46 / 72

47 copper wire silver deposits Figure Copper displaces silver from silver nitrate solution. Note the silver deposits formed on the copper wire and the pale blue colour of the resultant solution. Key point Displacement reaction is a reaction in which one element displaces another element from its compound Displacement reactions of metals P. 47 / 72

48 Zinc in copper(ii) sulphate solution A similar displacement reaction occurs when we place zinc into copper(ii) sulphate solution. Zn(s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu(s) silvery blue colourless reddish brown When putting copper into zinc sulphate solution no reaction occurs Displacement reactions of metals P. 48 / 72

49 zinc reddish brown copper coated on zinc copper(ii) sulphate solution Figure The reaction between zinc and copper(ii) sulphate solution. The part of the zinc metal dipped into the solution is coated with reddish brown copper. The blue colour of copper(ii) sulphate solution becomes paler. Experiment 11.2 Experiment Displacement reactions of metals P. 49 / 72

50 Key point A metal (M 1 ) higher in the reactivity series will displace any metal (M 2 ) lower in the series from the solution of a compound of M 2. A metal higher in the reactivity series is more reactive. its atoms lose electrons more readily to form positive ions. The positive ions of the less reactive metal would accept these electrons. forming the atoms of the less reactive metal Displacement reactions of metals Example 11.3 P. 50 / 72

51 11.6 Ionic equations Example Displacement reaction between copper and silver nitrate solution Cu Cu Cu Cu Cu Cu Cu copper Ag Ag Ag Ag Ag Ag Ag copper Ag + NO 3 NO 3 silver nitrate solution Ag + Ag + NO 3 Ag + NO 3 NO 3 silver deposits coated on copper Cu 2+ Cu 2+ NO 3 NO 3 NO 3 P. 51 / 72

52 Chemical equation: Cu(s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag(s) Silver nitrate and copper(ii) nitrate are ionic compounds which are soluble in water. In aqueous solutions, they appear as mobile ions i.e. Ag +, Cu 2+ and NO 3. Rewrite the chemical equation: Cu(s) + 2Ag + (aq) + 2NO 3 (aq) Cu 2+ (aq) + 2NO 3 (aq) + 2Ag(s) silver nitrate solution copper(ii) nitrate solution 11.6 Ionic equations P. 52 / 72

53 Nitrate ions (NO 3 ) remain unchanged in the displacement reaction. Ions which do not actually take part in the reaction spectator ions By cancelling out the spectator ions from the equation, we get: Cu(s) + 2Ag + (aq) Cu 2+ (aq) + 2Ag(s) This is called an ionic equation Ionic equations P. 53 / 72

54 Learning tip An ionic equation must be balanced with respect to the net ionic charges. Key point An ionic equation is an equation which includes only those ions that are formed or changed during the reaction. Example 11.4 Problem-solving strategy 11.2 Example 11.5 Class practice Ionic equations P. 54 / 72

55 11.7 Relation between the extraction method and the metal reactivity series Ease of extraction most difficult easiest most reactive Potassium, K Sodium, Na Calcium, Ca Magnesium, Mg Aluminium, Al Zinc, Zn Reactivity Iron, Fe Lead, Pb Copper, Cu Mercury, Hg Silver, Ag Platinum, Pt Gold, Au least reactive P. 55 / 72

56 Metals at the top of the reactivity series (e.g. potassium, sodium) give stable metal ores and are more difficult to be extracted. Electrolysis is used for extraction of these metals from their molten ores. Metals in the middle of the series (e.g. zinc, iron) are easier to be extracted. These metals are often extracted by heating their metal ores with carbon (i.e. carbon reduction) 11.7 Relation between the extraction method and the metal reactivity series P. 56 / 72

57 Metals near the bottom of the series (e.g. copper, mercury) are the easiest to be extracted. These metals are often extracted by heating the ore alone (or in air) or displacement from solutions Relation between the extraction method and the metal reactivity series P. 57 / 72

58 Metals at the bottom of the series (e.g. gold) are extracted by mechanical separation. Key point The lower the position of the metal in the reactivity series, the more easily it can be extracted from its ore. Example Relation between the extraction method and the metal reactivity series Class practice 11.8 P. 58 / 72

59 Key terms 1. balanced chemical equation 平衡化學方程式 2. chemical equation 化學方程式 3. displace 置換 4. displacement reaction 置換反應 5. ionic equation 離子方程式 6. metal reactivity series 金屬活性序 7. reactivity 活性 8. spectator ion 旁觀離子 9. stoichiometric coefficient 計量系數 P. 59 / 72

60 Progress check 1. How can we compare the reactivity of metals? 2. How do some common metals react with oxygen? 3. How do some common metals react with water? 4. How do some common metals react with dilute acids? 5. How can we construct a metal reactivity series? 6. How can we write the chemical equations for the reactions of common metals with oxygen, water and dilute acids? P. 60 / 72

61 7. How can we transcribe word equations into chemical equations? 8. How can we write balanced chemical equations to describe various reactions? 9. What is the relation between the reactivity of a metal and its tendency to form a positive ion? 10. What is a displacement reaction? 11. How can we write balanced ionic equations? 12. What is the relation between the extraction methods and the reactivity of metals? Progress check P. 61 / 72

62 Summary 11.1 Comparing reactivities of common metals 1. Reactivity of a metal is the readiness of it to react with other substances. 2. The reactivity of metals can be found by comparing their reactions with air, water and dilute acids. Refer to p.4 11 for the results of the reactions. P. 62 / 72

63 11.2 The metal reactivity series 3. The metal reactivity series is a series of common metals arranged in a decreasing order of reactivity. 4. The following table summarizes the appearances and reactions of metals in the reactivity series. Summary P. 63 / 72

64 Metal K Na Appearance of metal dull (stored under paraffin oil) air (oxygen) burns vigorously, forming metal oxide (Example 1) Reaction of metal with water/steam metal + water metal hydroxide + hydrogen (Example 4) Ca reacts with Mg decreasing metal + Al vigour, generally dull forming metal Zn oxide Fe (Example 2) Pb Cu Hg Ag Au generally shiny a layer of metal oxide formed on the surface (Example 3) no reaction steam metal oxide + hydrogen (Example 5) no reaction dilute hydrochloric acid reacts explosively, forming metal chloride and hydrogen (Example 6) reacts with decreasing vigour: metal + hydrochloric acid metal chloride + hydrogen (Example 7) no reaction Displacement reaction not applicable these three metals react with water to give hydrogen a metal displaces any other metal lower in the series from a solution of its compound (Example 8) Summary P. 64 / 72

65 Example 1: 4Na(s) + O 2 (g) 2Na 2 O(s) Example 2: 2Ca(s) + O 2 (g) 2CaO(s) Example 3: 2Cu(s) + O 2 (g) 2CuO(s) Example 4: 2Na(s) + 2H 2 O(l) 2NaOH(aq) + H 2 (g) Example 5: Zn(s) + H 2 O(g) ZnO(s) + H 2 (g) Example 6: 2K(s) + 2HCl(aq) 2KCl(aq) + H 2 (g) (NEVER attempt this experiment!) Example 7: Fe(s) + 2HCl(aq) FeCl 2 (aq) + H 2 (g) Example 8: Mg(s) + 2AgNO 3 (aq) Mg(NO 3 ) 2 (aq) + 2Ag(s) or Ionic equation: Mg(s) + 2Ag + (aq) Mg 2+ (aq) + 2Ag(s) Summary P. 65 / 72

66 11.3 Chemical equations 5. A chemical equation shows the physical states and relative numbers of particles of the reactants and products in a chemical reaction. 6. A reversible reaction is represented by a double arrow. 7. The steps in writing a chemical equation are shown in Problem-solving strategy 11.1 on p.16. Summary P. 66 / 72

67 11.4 Metal reactivity series and the tendency of metals to form positive ions 8. Metals react by losing electrons to form positive ions. Different metals have different reactivities because they have different tendencies to lose electrons. Atoms of a reactive metal lose electrons readily. Summary P. 67 / 72

68 11.5 Displacement reactions of metals 9. A metal (M 1 ) higher in the reactivity series will displace any metal (M 2 ) lower in the series from the solution of a compound of M 2. This is because a more reactive metal loses electrons more readily Ionic equations 10. An ionic equation is an equation which includes only those ions that are formed or changed during the reaction. Summary P. 68 / 72

69 11. An ionic equation must be balanced with respect to the ionic charges as well as the number of atoms. (Refer to Problem-solving strategy 11.2 on p.23.) 11.7 Relation between the extraction method and the metal reactivity series 12. The method used to extract a metal from its ores depends on the position of the metal in the reactivity series. 13. The lower the position of the metal in the reactivity series, the more easily it can be extracted from its ore. Summary P. 69 / 72

70 Concept map Extraction method of metal depends on the position of metal in Metal reactivity series Readiness of metals to lose electrons and form metal ions arranged in decreasing order REACTIVITY OF METALS P. 70 / 72

71 Metal reactivity series can be used to predict Displacement reactions arranged in decreasing order REACTIVITY OF METALS can be compared by The lowest temperature at which the reaction starts The rate of reaction The amount of heat given out Concept map P. 71 / 72

72 REACTIVITY OF METALS reaction of metals with Oxygen Water if no reaction, allow metal to react with Steam Dilute acid gives gives gives gives Metal oxide Metal hydroxide and hydrogen Metal oxide and hydrogen Salt and hydrogen Concept map P. 72 / 72