1. NATIONAL 5 CHEMISTRY
UNIT 3
CHEMISTRY IN SOCIETY

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

3. METALS 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

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.

9. Other Elements present
Alloy
Main Metal
Other Elements present
Uses
Reason
Stainless steel
Iron
Chromium, Nickel
Sinks, Cutlery
Non-rusting, strong
Mild steel
Carbon
Girders, Car bodies
Strong, rust resistant
Gold
Copper
Rings, Electrical contacts
Good conductor, unreactive
Solder
Lead (50%)
Tin (50%)
Joining metals, electrical contacts
Low melting point, good conductor
Brass
Zinc
Machine bearings, ornaments
Hard wearing, attractive

10. Reactions of Metals

11. metal + oxygen  metal oxide reaction with water
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 Name Symbol Potassium K Sodium Na Lithium
Calcium Ca
Magnesium Mg
Aluminium Al
Zinc Zn
Iron Fe
Tin Sn
Lead Pb
Copper Cu
Mercury Hg
Silver Au
Gold Ag
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. magnesium + oxygen  magnesium oxide 2Mg(s) + O2(g)  2MgO(s)
e.g.
magnesium + oxygen  magnesium oxide
2Mg(s) O2(g)  MgO(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.
METAL + OXYGEN METAL OXIDE

16. Metal + Oxygen  Metal oxide
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 O  MgO

17. 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)
METAL + WATER METAL HYDROXIDE + HYDROGEN

18. 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:
METAL + ACID SALT +HYDROGEN

19. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
e.g.
zinc + hydrochloric acid zinc chloride + hydrogen
Zn(s) HCl(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 Common name Chemical name Metal present Haematite
Iron oxide
Iron
Bauxite
Aluminium oxide
Aluminium
Galena
Lead sulphide
Lead
Cinnabar
Mercury sulphide
Mercury
Malachite
Copper(II) carbonate
Copper

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 extraction a) 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
Fe2O 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
– 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. In most batteries the electrons come from a layer of zinc metal.
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.
Electricity is a flow of electrons.

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
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. 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.
metal B V
metal A
filter paper soaked in a sodium chloride solution
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 electrons and 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.

51. Cells Involving Non-Metals
In the cell shown, the two half cells are a solution of iodide ions and a solution of iron (III) ions.
2I-  I2 + 2e-
Fe3+ + e-  Fe2+

52. Electrons flow from the iodide ions through the meter to the iron (III) ions. As this happens, the iodide ions turn into iodine molecules.
The iron (III) ions gain electrons and turn into iron (II) ions.

53. Oxidation and Reduction

54. Oxidation
Below are the ionic equations for the reactions of calcium with oxygen, water and dilute acid.
Calcium atoms have lost electrons to become calcium ions as shown below
This is an OXIDATION reaction.

55. An oxidation reaction is one in which there is a loss of electrons.

56. Reduction
Reactions in which there is a gain of electrons are called REDUCTION reactions.
A reduction reaction is one in which there is a gain of electrons.

57. Oxidation Is
Loss of electrons
Reduction Is
Gain of electrons

58. REDOX Reactions
Oxidation and reduction reactions take place at the same time.
In a redox reaction, electrons lost by one substance during oxidation are gained by another substance during reduction.
The formation of a compound by a metal is a redox reaction.

59. e.g. when sodium joins with chlorine to form sodium chloride
All displacement reactions are redox reactions
e.g. the reaction between zinc metal and copper(II) sulphate solution.

60. The oxidation and reduction equations can be combined to show the overall REDOX reaction.

61. The reaction between a metal and a dilute acid can also be considered to be a redox reaction.
e.g. The reaction between magnesium and hydrochloric acid.
The magnesium is oxidised and forms magnesium ions, whilst the hydrogen ions in the acid gain electrons and form hydrogen gas.

62. During electrolysis, oxidation occurs at the positive electrode and reduction occurs at the negative electrode.

63. Hydrogen fuel cell How Fuel Cells work
Advantages and disadvantages of hydrogen as a fuel.

64. What is a Fuel Cell?
Quite simply, a fuel cell is a device that converts chemical energy into
electrical energy, water, and heat through electrochemical reactions.
Fuel and air react when they come
into contact through a porous
membrane (electrolyte) which separates
them.
This reaction results in a transfer of
electrons and ions across the electrolyte
from the anode to the cathode.
 If an external load is attached to this
arrangement, a complete circuit is formed
and a voltage is generated from the flow
of electrical current.
The voltage generated by a single cell is typically rather small (< 1 volt), so many
cells are connected in series to create a useful voltage.

65. Fuel Cell Vs. Battery
Basic operating principles of both are very similar, but there are several
intrinsic differences.
Hydrogen fuel cell
Galvanic cell (battery)
Open system
Anode and cathode are gases in
contact with a platinum catalyst.
Reactants are externally supplied,
no recharging required.
 Closed system
Anode and cathode are metals.
Reactants are internally consumed,
need periodic recharging.

66. Fuel Cell Vs. Internal Combustion Engine
Similarities:
Both use hydrogen-rich fuel.
Both use compressed air as the oxidant.
Both require cooling.
Differences:
Fuel cell:
 Output is electrical work.
 Fuel and oxidant react electrochemically.
 Little to no pollution produced.
I.C. Engine:
Output is mechanical work.
Fuel and oxidant react combustively.
Use of fossil fuels can produce significant pollution.

67. Some History… Fuel cell principle first discovered
by William Grove in 1839.
Grove used four large cells, each
containing hydrogen and oxygen,
to produce electric power which was
then used to split the water in the
smaller upper cell.
Commercial potential first demonstrated by NASA in the 1960’s with the
usage of fuel cells on the Gemini and Apollo space flights. However, these
fuel cells were very expensive.
Fuel cell research and development has been actively taking place since the
1970’s, resulting in many commercial applications ranging from low cost portable
systems for cell phones and laptops to large power systems for buildings.

68. Fuel Cells in Use: Stationary Systems

69. Fuel Cells in Use: Stationary Systems
Fuel cell system for submarine

70. Fuel Cells in Use: Transportation Systems
Buses are most commercially
advanced applications of fuel
cells to date.
Are currently being used by
many American and European
cities.
XCELLSiS fuel cell bus prototypes

71. Fuel Cells in Use: Transportation Systems
Many of the major car companies are developing fuel cell car prototypes
which should come to market during the next decade. The cars use either
pure hydrogen or methanol with an on board reformer.

72. Fuel Cells in Use: Hydrogen Fuel Cell System

73. Fuel Cells in Use: Space Systems
12 kW Space shuttle fuel cell
Weight: 120 kg
Size: 36x38x114 cm
Contains 32 cells in series
1.5 kW Apollo fuel cell
Apollo used two of these
units.

74. Fuel Cells in Use: Portable Systems
A laptop using a fuel cell power source
can operate for up to 20 hours on a
single charge of fuel (Courtesy: Ballard
Power Systems)

75. But Isn’t Hydrogen Explosive?
Many have blamed this disaster on
a hydrogen explosion. However,
hydrogen burns invisibly, and no
evidence of leaks were ever found
(garlic scent added to hydrogen gas).
Using infrared spectrographs, NASA
scientists found that the skin of the
Hindenburg was treated with compounds
which are found in gunpowder and
rocket fuel (nitrates and aluminum
powder). This, combined with a wooden
frame coated with lacquer resulted in
a highly flammable ship.

76. redox

77. Glossary of Terms Used in Describing Fuel Cell Technology
Electrochemical reaction: A reaction involving the transfer of electrons
from one chemical substance to another.
Electrode: An electrical terminal that conducts an electric current into or
out of a fuel cell (where the electrochemical reaction occurs).
Anode: Electrode where oxidation reaction happens (electrons are released).
Cathode: Electrode where reduction reaction occurs (electrons are acquired).
In a fuel cell, hydrogen is oxidized at the anode and oxygen reduction occurs
at the cathode.
Electrolyte: A chemical compound that conducts ions from one electrode to
the other.
Ion: An atom that carries a positive or negative charge due to the loss or gain
of an electron. Anion is a negative ion, cation is a positive ion.
An electrochemical cell consists of 2 electrodes + 1 electrolyte

78. Terminology (cont.)
Catalyst: A substance that participates in a reaction, increasing its rate,
but is not consumed in the reaction.
Polymer: A natural or synthetic compound made of giant molecules which
are composed of repeated links of simple molecules (monomers).
Inverter: A device used to convert direct current electricity produced by a
fuel cell to alternating current.
Reformer: A device that extracts pure hydrogen from hydrocarbons.
Stack: Individual fuel cells connected in series within a generating assembly.