1. METALS
National 5

2. c Starter Which of the following diagrams could
represent the structure of a metal ? c

3. Lesson 1: Metals
Today we are learning to understand some of the properties of metals.
By the end of today I can:
Complete a table to show the boiling point/melting point of metals and conductivity of metals.
Explain metallic bonding

4. Metals Metals have amazed and intrigued us for thousands of years.
From tungsten, the element with the highest melting point to lithium, a metal that floats on water; from gold, a metal so unreactive it is shiny even after thousands of years, to caesium, a metal so reactive it’s stored under a vacuum.
They can be used to create electricity and convert harmful car exhaust fumes into harmless gases, they can corrode but also protect. They can be strong and rigid or soft and supple, one is even a liquid.
Some are found as pure nuggets lying at the bottom of fast flowing rivers; others are so tightly bound to other elements, huge amounts of energy are required to separate them.
Metals are truly magical!

5. Activity 9.1 Element Symbol M.P B.P State at 25˚C Cadmium Calcium
Gallium
Tungsten
Mercury
Platinum

6. Conductivity

7. Set up either of the following electrical circuits to test the electrical conductivity of metals on the following table

8. (9.2)Conductivity task:

9. Metallic bonding
Most metals are SOLIDS at room temperature and HARD with high melting/boiling points
All elements want to achieve a full outer electron shell.
Metals will give up 1, 2 or 3 electrons to form +ve metal ions

10. Metallic bonding
The greater the number of electrons in the outer shell the stronger the metallic bond.
So the melting point of Al>Mg>Na

11. Metallic elements
Strong electrostatic forces exist between the delocalised electrons (free to move) and the positive metal ions formed due to the loss of outer electrons. These electrostatic attractions are known as metallic bonds.

12. Conducting electricity

13. Plenary flow chart:
Pupils to complete the flow chart in their jotters.
Follow the questions and answer for (a-d) substance z in each scenario.

14. Does it conduct electricity when solid?
Does it conduct electricity when it is liquid?
Does it have a high or low M.P or B.P?
(a)
(b)
LOW
(c)
(d)

15. Lesson 2:
Reactions of Metals

16. Metals Anagram
Conducting electricity is hugely important, but what other properties make metals amazing?
Unscramble these anagrams.
ATHE RONOCTCDU
LLAEMBEAL
HSIYN
LICDUTE
WLO SDNYITE
TTRGHESN

17. Lesson 2: Reaction of Metals
Today we are learning to understand some of the reactions of metals.
By the end of today I can:
List extra properties of metals
Understand how important the reactivity series is.
List some of the alkali metals in order of reactivity due to experiment demonstration.
State some of the reactions a metal can undergo.

18. (9.3) Coloured compounds
Fill in the table to show the colour of these ions.
You may remember doing the flame test and may want to
use your data booklet to help you

19. Reactions of metals Some metals are more reactive than others.
The amount of energy given out when a metal reacts gives a measure of its reactivity.
By observing how vigorous the reaction is between a metal and:
WATER
OXYGEN
DILUTE ACID
We can begin to order the metals in a reactivity series

20. Reactivity Series Based on their reactivity, chemists produced
a table of metals as shown:

21. Reaction with Water: The Alkali Metals (9.5)
Observe the reaction between 3 alkali metals with water.
In your jotter record the reactivity of the three metals, note which gas is produced.
Write a word equation and balanced equation for the reaction of Potassium and water

22. Reaction with Water (9.5)
Half fill three small (100cm3) beakers with water.
Into the first place half a spatula of calcium granules, the second half a spatula of magnesium powder and the third half a spatula of zinc powder.
Observe what happens (you may need to leave them for several minutes).

24. Metals Reacting with Water
ALL metals above aluminium in the reactivity series react with water to produce the metal hydroxide and hydrogen gas.
The metals from groups 1 and 2 react vigorously.
METAL WATER METAL HYDROXIDE +
HYDROGEN

25. Example
Lithium + Water Lithium Hydroxide Hydrogen

26. Lesson 3

27. Starter:
What is the name given to the gas given off when an alkali metal is placed in water ?
Put the 3 metals observed in order of most reactive to least reactive
What colour would the water of gone if universal indicator was dropped into the glass bowl ?

28. (9.6)Metals with acid:
Half fill six test tubes with hydrochloric acid and place them in a test tube rack.
Place a piece of magnesium, aluminium, zinc, iron, tin and copper into each test tube.
Observe what happens.

29. Table of results:

30. Metals Reacting with Dilute Acid
Metals above hydrogen in the reactivity series react with acids to produce a salt and hydrogen gas.
Metals below HYDROGEN do not react with acids.
METAL ACID SALT + HYDROGEN

31. Example
Zinc Hydrochloric
acid
Zinc Chloride
+ Hydrogen

32. (9.7)Metals with oxygen:

33. Table of results:
Metal
Observation
Order of Reactivity

34. Metals Reacting with Oxygen
Some metals react vigorously with oxygen and burn
Fiercely, some react more slowly and others do not react at all.
METALS react with OXYGEN to make a METAL OXIDE.
METAL OXYGEN METAL OXIDE

35. Example
Magnesium Oxygen Magnesium O Oxide

36. Reactions of Metals and the Reactivity Series
The results of experiments with Oxygen, water and dilute acids allow scientists to place metals into an order of reactivity.

37. A Zinc B Copper C Hydrogen D Gold E Carbon dioxide F Calcium GB
Copper C
Hydrogen D
Gold E
Carbon dioxide F
Calcium G
Magnesium H
Potassium I
Oxygen
Which boxes show metals which react with dilute acids?
Which boxes show metals which react with dilute acid but not with cold water?
Which boxes show metals that do not react with dilute acid?
Which boxes show metals which react with oxygen but not with a dilute acid?
Which box shows the gas produced when sodium reacts with water?

38. Lesson 4
Naming salts

39. Naming salts
Salts have both a first and second name like us. The first part comes from the alkali/metal. The second part comes from the acid.

41. Metals and hydrochloric acid- equations
Magnesium + hydrochloric Magnesium chloride + acid hydrogen Aluminium +hydrochloric ________ +___________ acid Zinc + hydrochloric ____________ +_______

42. Metal with sulfuric acid- equation
Magnesium + sulfuric acid Magnesium + hydrogen sulfate Aluminium + _______ ________ + ___________ _______ + __________ _________ + __________

43. Metals and nitric acid- equations
Magnesium +________ ___________ +________ _________+_________ _________ + _________ ________ + _________ __________ + ________

45. Challenge:
Create a word equation using any metal above hydrogen from the previous slide and one of the 3 acids we have been using.

46. Lesson 5
Ionic equations

47. Ionic equations
Today we are learning to write ionic equations for some chemical reactions and about alloys
By the end of today I can…
Write the ionic equation for some chemical reactions
State the definition for an alloy.

48. Ionic equations
Whenever a metal reacts it always turns in to a positive charged ion by losing it’s outer electrons to whatever it is reacting with.
The other reactants must gain these electrons and form an negative charged ion.

49. Example:
Word equation: Sodium + Water Sodium hydroxide + hydrogen Chemical equation: 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) Ionic equation: 2Na(s) + 2H2O(l) 2[Na+](aq) + 2[OH-](aq) + H2(g)

50. Word equation: Chemical equation: Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)
magnesium + hydrochloric acid magnesium chloride +hydrogen
Chemical equation:
Mg(s) HCl(aq) MgCl2(aq) + H2(g)
Ionic equation:
Mg(s)+2H+(aq)+ 2Cl-(aq) Mg2+(aq)+2Cl-(aq) + H2(g)

51. Spectator ion:
If ions are not involved in a reaction, they are called SPECTATOR ions and are not shown in the final ionic equation:
Write ionic equations for the following chemical reactions:
A. Potassium + water potassium hydroxide + hydrogen
B. Zinc + hydrochloric acid zinc(II) chloride + hydrogen
C. Copper + oxygen copper(II) oxide

52. Alloys
When metals combine they do not form compounds, they just mix and form a mixture called an alloy. The physical properties of the metal change however, depending on the proportion of each metal.
For example solder is an alloy of tin and lead. It is used to join wires to components in circuit boards. On their own the tin was too hard and the lead too soft, but by mixing the tin and lead in the right proportion an alloy can be made which is easy to melt but hard enough to withstand impacts on the circuit board.
Some alloys contain one main metal with small quantities of other metals and some non-metals. Stainless steel contains mainly iron with small quantities of carbon, which makes it less brittle, and chromium, to make it shiny and corrosion resistant.

53. Gold Alloys
Pure gold is a relatively soft metal and so it is not usually used to make jewellery. In order to make it stronger it is mixed with other metals such as copper and silver. Alloying also changes the colour of gold. Although alloying gold reduces its value, it is still valuable. The percentage of gold in an alloy is indicated by the number of carats (ct) it has. Pure gold is said to be 24ct. The percentage of gold and other metals in some different alloys of ‘yellow’ gold are shown in the following table.

54. Carats (ct)
Percentage gold (%)
Percentage other metals (%) 22
91.7
Ag: 5.o Cu: 2.0 Zn: 1.3 18
75.0
Ag:15.0 Cu: 10.0 14
58.3
Ag: 30.0 Cu: 11.7 9
37.5
Ag: 42.5 Cu 20.0

55. Heat alone and Heat with Carbon
Lesson 6
Heat alone and Heat with Carbon

56. Starter:
Lithium + water Lithium hydroxide + hydrogen Write the chemical equation and the ionic equation for the above reaction.

57. Learning intentions & success criteria
Today we are learning about extracting metals from their ores and extracting some metals from its ore with the aid of carbon.
By the end of today I can
State what metals are extracted with heat alone
State what metals are extracted with heat and carbon.

58. Extracting metals Metals exist in the earth’s crust as METAL ORES, the
natural compound which they exist as.
Metals are extracted from their ores in different
ways, depending on the REACTIVITY of the metal.
(More reactive metals are harder to separate from
their ores.)

59. Heat alone (9.9) Heat Alone
1. Put a spatulas of silver oxide in one test tube and a spatula of copper oxide in another.
2. Heat strongly using a roaring Bunsen burner flame.
3. Test for the presence of oxygen by placing a glowing splint into each test tube.

60. Heat alone
Some metals can be extracted from their ore (metal oxide) by heat alone. This is only the case for the least reactive metals.
metal oxide metal + oxygen

61. Extracting metals summary

62. (9.10) Heat and Carbon
1. Put two spatulas of copper oxide into a test tube.
2. Hold the test tube with a test tube holder and heat using a roaring blue flame.
3. Take a burning splint and plunge it into the hot copper oxide powder.
4. Continue to heat the test tube.
5. Pour the powder out onto a heat proof mat and look at the inside of the test tube, you should see copper on the inside of the glass.
6. Repeat the experiment with iron(III) oxide.

63. Heat and Carbon
Metals that are slightly more reactive were only discovered when their ore was accidently dropped in a fire. The heat and the carbon caused the metal to be displaced and the carbon to join up with the oxygen. This was eventually known as smelting.
The carbon is called a reducing agent.

64. It was later found that a much higher temperature was required to extract iron. To achieve this they needed to blow hot air into what was called a Blast Furnace; the oxygen reacted with the carbon to form carbon monoxide which then reacted with the iron oxide.
iron(III) + carbon iron carbon dioxide
oxide monoxide  

66. Lesson 7
Electrolysis

67. Starter: What metal did we try to extract with heat and carbon ?
Why did we use a wooden peg to grip the test tubes ?
Why did we aim the test tubes at the wall ?
How would you conduct this experiment on a large scale ?

68. Learning intentions & Success criteria
Today we are learning about electrolysis and the percentage of metal in an ore.
By the end of today I can…
state the definition for electrolysis
Calculate the percentage of metal present in an ore.

69. (9.11)Electrolysis

70. Extracting metals summary:

71. Percentage of Metal in an Ore

72. Work out the %of metal in the following ores:
copper(II) carbonate CuCO3
lead(IV) oxide PbO2
aluminium oxide Al2O3

73. Tenorite Cu2O is an ore of copper. Given that copper has a mass of 63
Tenorite Cu2O is an ore of copper. Given that copper has a mass of 63.5 and oxygen a mass of 16, calculate the percentage by mass of copper in tenorite.
Gibbsite Al(OH)3 is a mineral found in aluminium ore. Given the relative atomic mass of aluminium is 27, oxygen is 16 and hydrogen is 1, calculate the percentage by mass of gibbsite that is aluminium.

74. Electrochemical cells

75. Starter: Chalcopyrite (FeS2) Siderite (FeCO3) Cinnabar (HgS)
Calculate the percentage composition that is a metal in each of these examples.

76. Learning Intentions and Success Criteria
To learn about the different components of an electrochemical cell
To learn about the significance of electrochemical cells within chemistry
Success Criteria
I can recall information about the electrochemical series
I can interpret data within the data booklet
I can differentiate the different components of the electrochemical cell
I can distinguish the accepting metal and receiving metal within the electrochemical cell

78. (9.13) A Lemon Cell

79. (9.14)Experiment

81. Electrochemical series

82. Plenary
On your exit ticket write the answer yes or no to the following question:
Would a voltage be produced in the following experiment set up ?

83. Redox And Displacement

84. Starter Which way do the electrons flow within each cell
Which cell would have the biggest voltage Mg Cu Al Cu

85. Recap
In the last lesson we looked at the lemon cell demo and a simple electrochemical cell where the metals were changed.
In the practical it was found that the metal higher up the electrochemical series donated its electrons to the metal lower down the electrochemical series.

86. Learning Intention and Success Criteria
Learning Intentions
To learn about the processes of oxidation and reduction
To learn about reduction and oxidation half equations
To learn about redox equations
To learn about displacement reactions
Success Criteria
I can state the processes of reduction and oxidation
I can recognise the processes of reduction and oxidation
I can recall the OIL RIG Memory aid
I can manipulate reduction and oxidation half equations

87. Summary of experiment
As the electrons flow from one metal to the other, one metal must donate electrons and the other accept them.
You can see from the table that some metals are better at donating electrons and some metals better at accepting them.
If we place the metals in order of their ability to donate electrons we form what is known as the electrochemical series.
Metals good at donating electrons are at the top and the metals at the bottom are better at accepting electrons.

88. Electrochemical Cells
Electrically conducting solutions containing ions are known as electrolytes.
A simple cell can be made by placing two metals in an electrolyte.
Another type of cell can be made using two half-cells (metals in solutions of their own ions).
An ‘ion bridge’ (salt bridge) can be used to link the half-cells. Ions can move across the bridge to complete an electrical circuit.

89. Electricity can be produced in cells where at least one of the half-cells does not involve metal atoms/ions. A graphite rod can be used as the electrode in such half-cells.
Different pairs of metals produce different voltages. These voltages can be used to arrange the elements into an electrochemical series.
The further apart elements are in the electrochemical series, the greater the voltage produced when they are used to make an electrochemical cell.
Electrons flow in the external circuit from the species higher in the electrochemical series to the one lower in the electrochemical series.

95. Redox

96. Electrolysis = Redox

97. Examples: Copper and bromine ions
(making copper solid and bromine gas)
Magnesium and chlorine ions
(making solid silver and chlorine gas)
Calcium and iodide ions
(making calcium solid and iodine solid)

98. Displacement reaction

100. Redox- Displacement

102. (9.16) Redox Reactions Try these three experiments:
1. Pour out 2cm depth of copper(II) sulfate solution into a test-tube, add a spatula of iron filings. Put a rubber bung in and shake. Observe what happens.
2. Pour out 2cm depth of hydrochloric acid into a test-tube and add a strip of magnesium ribbon. Observe what happens.
3. Pour out 2cm depth of iodine solution into a test-tube and add a spatula of sodium sulfite. Observe what happens.

104. Plenary
On your exit ticket write the memory trick used to remember oxidation and reduction.