The document provides an overview of oxidation and reduction concepts including: - Oxidation involves loss of electrons while reduction involves gain of electrons. - Examples of oxidation and reduction reactions are given for sodium-chlorine, magnesium-oxygen, and zinc-copper sulfate. - Oxidizing and reducing agents are defined as substances that cause oxidation or reduction in other substances. - The electrochemical series orders metals by their tendency to be oxidized. - Electrolysis and examples like copper plating and extracting copper from scrap iron using electrolysis are summarized.

1. • 1.5 Oxidation and
• Reduction
• (Time needed: 7 class periods)

2. Learning Outcomes
• Introduction to oxidation and reduction: simple examples only, e.g. Na
with Cl2, Mg with O2, Zn with Cu2+. Oxidation and reduction in terms
of loss and gain of electrons.
• Oxidising and reducing agents.
• The electrochemical series as a series of metals arranged in order of
their ability to be oxidised (reactions, other than displacement
reactions, not required).
• Electrolysis of (i) copper sulfate solution with copper electrodes and
(ii) acidified water with inert electrodes. (Half equations only
required.)

3. Oxidation and reduction
• Oxidation = addition of oxygen to a
substance
• C + O2 CO2
• Reduction is loss of oxygen or addition of
hydrogen
• CuO + H2  Cu + H2O

4. examples
• Sodium + chlorine  sodium chloride
• Na + Cl  Na+ + Cl• Na loses an electron [oxidised]
• Cl gains an electron [reduced]

7. Example 2
• Magnesium +oxygen  magnesium oxide
• Mg + O  MgO
• Mg  Mg+2 loses 2 electrons [oxidation]
• O  O-2 gains 2 electrons [reduction]

8. Example 3
• Zinc +copper sulphate  Zinc sulphate+ Copper
• Zn + Cu+2  Zn+2 + Cu
• Zinc loses electrons (oxidised)
• Copper gains electrons (reduced)

9. Oxidising agent
• A substance that causes oxidation in another
substance

10. Reducing agent
• A substance that causes reduction in another
substance.

11. •
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Oxidation is loss of electrons;
Reduction is gain of electrons
CuO + H2  Cu + H2O
CuO  Cu+2 and O-2
Cu+2  Cu [gains 2 electrons] reduced
H2  H2+2[loses 2 electrons] oxidised
O-2  O-2 [ no change]

12. Oxidation numbers
• The charge that an atom has or appears to
have assuming that the compound is ionic.
• Electrons always go the the most
electronegative element

13. Oxidation number rules 1
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Elements on their own = 0
H2 = 0
Zn = 0
Cl2 = 0

14. Oxidation number rules 2
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Ions = same as charge
Cu +2 = +2
O-2 = -2
Cl-1 = -1

15. Oxidation number rules 3
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Charges of all elements in a compound = 0
CuSO4
Cu = +2
S = +6
O4 = -8 [O = -2]
Total = +2 +6 –8 = 0

16. Oxidation number rules 4
• Oxygen = -2
• Exceptions are
• peroxides O = -1 [H2O2, Na2O2 ]
• OF2 O = +2, F = -1

17. Oxidation number rules 5
• Hydrogen = +1
• Exceptions are the metal hydrides
• NaH
Na = +1, H = -1

18. Oxidation number rules 6
• Halogens [ Cl, F, I, Br] are always –1
except when joined to more electropositice
element
• Cl2O
• Cl = +1, O = -2

19. Oxidation number rules 7
• In a complex ion the sum of all the charges
= the chartge on the ion.
• SO4-2
• S = +6, O4 = -8 [O = -2]
• +6 –8 = -2

20. redox
• Oxidation is an increase on oxidation
number
• Reduction is a decrease in oxidation
number.

21. Electrochemical Series
• Electrochemical Series – Elements listed
in order of ability to be oxidised

22. Metals
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King [K]
Neptune [Na]
Caught [Ca]
Many [Mg]
Angry [Al]
Zulus [Zn]
Fighting [Fe]
Police [Pb]
Constables [Cu]
Having [Hg]
Asthma [Ag]
Attacks [Au]

23. Zinc
Metals above
hydrogen in the
Reactivity Series
react with acids to
produce hydrogen
gas.
Potassium
Sodium

24. Displacement of metals
• Displacement reactions occur when a metal
from the electrochemical series is mixed
with the ions of a metal lower down in the
series. The atoms of the more reactive
metal push their electrons on to ions of the
less reactive metal.

25. Displacement
• More reactive metal displaces less
reactive from a solution
• Mg + CuSO4 = MgSO4 + Cu
• Mg + Cu+2  Mg+2 + Cu
• Mg loses electrons (Oxidised)
• Cu+2 gains electrons (reduced)

26. Learning Outcomes
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Rusting of iron.
Swimming-pool water treatment.
Use of scrap iron to extract copper.
Electroplating. Purification of copper.
Chrome and nickel plating. Cutlery.

27. Rust
• Rust is the formation of iron oxides
(usually red oxides), formed by the reaction
of iron and oxygen in the presence of water
or air moisture.
• Oxidation

28. Swimming pools
• The water in swimming pools is kept sterile by the
addition of oxidizing agents, chlorine or chlorine
compounds, which kill microorganisms by oxidation.
The active agent is usually chloric(1) acid (HOCl). It
may be formed in two ways
• 1. Direct chlorination of the water:
• Cl 2(aq) +H2O (l)  HOCl (aq) + Cl− (aq) + H+ (aq)
• Note that when the Cl2 reacts with the water it is both
oxidized and reduced

29. Swimming pools
• 2. The addition of sodium chlorate(I) [sodium
hypochlorite]:
• NaOCl (s) + H2O (l) Na+ (aq) + OH− (aq) + HOCl (aq)
• Nowadays chlorine is not used, mainly on grounds of
safety. Pools are sterilized with chlorine compounds,
which produce chloric(I) acid when they dissolve in
water. These compounds act in essentially the same
way as chlorine. Sodium chlorate(I) is one such
compound.

30. Use of scrap iron to extract copper.
•
(Dissolved CuSO4) + (Metallic Fe) ==> (Dissolved FeSO4) + (Metallic Cu)

31. Electrolysis
• Chemical reaction
caused by the passage
of an electric current
through a liquid
known as the
electrolyte

32. Definitions
•
Electrolyte - liquid in which electrolysis takes place. Usually an ionic solution but it can also be a
fused [melted] ionic compound
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Anode - positive electrode. Positive because the battery sucks electrons out of it
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Cation - positive ion. Called cation because it is attracted to the opposite charge of the cathode.
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Inert Electrodes - do not react with the electrolyte Graphite and Pt
•
Active electrodes - react with electrolyte e.g. Copper and iron
Cathode. Negative electrode. Negative because the battery pumps electrons into it.
Anion - negative ion. Called anion because it is attracted to the opposite charge of the anode

33. Electrolysis

34. Electroplating
• Electroplating
• Covering cathode in metal e.g. Cu by
making it cathode in copper sulphate
solution

35. Copper plating

36. Copper Plating
Anode reaction
• Cu(s) = Cu2+(aq) + 2e• Anode loses mass as copper dissolves off
• Impurities [Au, Ag, Pt etc.] fall to bottom
Cathode reaction
• Cu2+(aq) + 2e- = Cu(s)
• Cathode gains mass as Cu is deposited on it
• Cu is 99.9% pure

37. Learning Outcomes
• Mandatory experiment 1.2 (half equations only required,
e.g. 2Br– – 2e– → Br2).
• Demonstration of ionic movement.
• Demonstration of electrolysis of aqueous sodium sulfate
(using universal indicator)
• and of aqueous potassium iodide (using phenolphthalein
indicator) with inert electrodes. (Half equations only
required.)

38. Ionic Movement
• During electrolysis of a solution of Copper Chromate in
dil. Hydrochloric acid, positive ions (cations) are attracted
to the negative electrode (cathode) and negative ions
(anions) are attracted to the positive electrode (anode). If
these ions are coloured, their movement may be observed
visually.
• Examples of coloured ions include;
• copper(II) [Cu2+] - blue
• chromate(VI) [CrO42- ] – yellow

39. Q & A to Ionic Movement Expt
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(1) What colour is the copper(II) chromate solution?
Copper(II) chromate solution is an olive green colour.
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(2) What colour is observed at the positive electrode after the power supply has been turned on for some time?
A yellow colour is observed at the positive electrode.
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(3) What colour is observed at the negative electrode after the power supply has been turned on for some time?
A blue colour is observed at the negative electrode.
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(4) Explain in terms of the movement of ions why different colours are formed at each electrode.
When the circuit is completed, positive copper ions (Cu2+) are attracted to the negative electrode. These ions have a
blue colour. Similarly negative chromate(VI) ions (CrO42-) are attracted to the positive electrode. These ions are
coloured orange.
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(5) What is the function of the dilute hydrochloric acid?
The dilute hydrochloric acid is required to complete the circuit.

40. Electrolysis of Sodium sulphate
• Solution of Na2SO4 + universal indicator
• H+ ions are produced at the positive
electrode (oxidation of O2- in water) while
OH- ions are produced at the negative
electrode as the H+ in water is reduced to
H2(g).

41. Sodium Sulphate and Universal
Indicator

42. Electrolysis of Sodium Sulphate
• Red is acid at the positive electrode
•
2H2O(l) O2(g) + 4H+(aq) + 4 e•
lose electrons = oxidation = anode
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• Purple is base at the negative electrode
•
H2O(l) + 2 e- H2(g) + 2OH-(aq)
•
gain electrons = reduction = cathode

43. Electrolysis of Potassium Iodide
• Solution of KI + phenolphthalein
• Brown I2(s) forms at the positive electrode
and some yellow/orange I3- forms in
solution. At the negative electrode, H+ is
again reduced to H2(g) and the
phenolphthalein turns pink due to the OHions.

44. Electrolysis of Potassium Iodide

45. Electrolysis of Potassium Iodide
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KI  K+ + IIodide loses electrons  Brown iodine
2I-  I2 + 2e- Anode, Oxidation
H2O  H+ + OH-
• OH- is basic , Phenolphthalein Purple.