1. Redox Chemistry and Corrosion

2. Oxidation and Reduction
So far we have looked at precipitation reactions and acid-base reactions.
Now we shall look at a third group of chemical reactions.
They are called oxidation-reduction reactions.
These reactions are commonly referred to as redox reactions.

3. Redox Reactions
Many of the chemical reactions that play a significant role in maintaining our environment are redox reactions.
Corrosion and the deterioration of metals are redox reactions.
Iron which is used as a structural base for buildings and bridges is particularly prone to corrosion.
Australia spends about 3 billion dollars a year in an effort to prevent corrosion and replacing structures that have corroded.

4. Redox Reactions
These reactions are also used in the processing of mineral ores to extract from then the metals our society requires.
One of Australia’s biggest exports is the mining of these mineral ores.

5. Redox Reactions Other redox reactions include:
The respiration reaction that is the source of energy in almost all living things.
Photosynthesis in green plants
Burning of fuels to propel our cars.
Combustion of coal in electricity power stations.
Use of chemicals such as chlorine to disinfect swimming pools.
Manufacture and use of explosives.
Use of electrolysis to produce many chemicals.
Production and use of fertilisers.

6. Redox Reactions Many chemicals react with oxygen.
Reactions such as these were described as oxidation reactions.
In air, the combustion of carbon, sulfur, iron or even octane always produced at least one oxide:
C(s) + O2(g) ―› CO2(g)
S(s) + O2(g) ―› SO2(g)
4Fe(s) + 3O2(g) ―› 2Fe2O3(s)
2C8H18(l) + 25O2(g) ―›16CO2(g) + 18H2O(l)

7. Oxidation Oxidation means the addition of oxygen.
When oxygen reacts with an element, the element is said to be oxidised.
Because elemental iron reacts with oxygen, there are no large deposits of elemental iron found on earth.
Iron is generally found as a compound of mineral oxide ores (haematite (Fe2O3) and magnetite (Fe3O4))

8. Reduction
Iron used in modern society has been extracted from iron ores.
This extraction process involves reduction of the iron oxide to iron.
It involves the removal or oxygen.
When oxygen is removed from a substance, that material has been reduced.

9. Reduction
The production of iron from haematite can be represented by the reduction equation:
Fe2O3(s) + 3CO(g) ―› 2Fe(s) + 3CO2(g)
The iron(III) oxide has lost an oxygen – it has been reduced.
Reduction cannot occur without oxidation occurring at the same time.
In this reaction the carbon monoxide has gained an oxygen – it has been oxidised.
Reduction – loss of oxygen
Oxidation – gain of oxygen

10. A Better Definition
There are many oxidation and reduction reactions that don’t involve oxygen.
Instead we define oxidation as the loss of electrons.
Similarly, reduction is the gain of electrons rather than the loss of oxygen.

11. OIL RIG Oxidation is the loss of electrons
Reduction is the gain of electrons.

12. Magnesium Oxide
You have used magnesium in class before, remember how it has a coating on it that sometimes we have had to scrape off.
That is magnesium oxide which results in corrosion of magnesium in air.

13. Magnesium Oxide
The magnesium has reacted with atmospheric oxygen to form magnesium oxide.
The magnesium has been oxidised.
2Mg(s) + O2(g) ―> 2MgO(s)

14. 2Mg(s) + O2(g) ―> 2MgO(s)
Magnesium oxide is an ionic compound and consists of Mg2+ ions and O2- ions.
Each magnesium ion, therefore must have lost two electrons to form an Mg2+ ion. Each oxygen atom in the oxygen molecule must have gained two electrons to form an oxide ion O2-.
The reaction can now be represented by two half equations.

15. 2Mg(s) + O2(g) ―> 2MgO(s)
The first half equation show the gain of two electrons by each oxygen atom in the oxygen molecule:
Mg(s) ―> Mg2+(s) + 2e-
The second show the gain of two electrons by each oxygen atom in the oxygen molecule.
O2(g) + 4e- ―> 2O2-(s)

16. Mg(s) ―> Mg2+(s) + 2e- O2(g) + 4e- ―> 2O2-(s)
So the oxidation of magnesium involves the transfer of electrons from magnesium atoms to oxygen atoms.
Note that there is no real ‘loss of electrons’ but rather a transfer of electrons from the magnesium to the oxygen.
If an atom loses electrons, there must be another atom that can gain electrons.
Therefore oxidation and reduction occur simultaneously.

17. Writing Redox Half Equations
Worked Example 16.2a page 275
16.2b

18. Your Turn
Page 278
Question 1
Question 2

19. Writing an Overall Redox Equation
When we write equation for redox reactions, we normally write the two half equations first.
We then follow this with the overall equation.
In the overall equation we do not show any electrons transferred as:
The electrons lost in the oxidation reaction are gained in the reduction reaction.

20. Copper and the solution of silver ions
In the previous example:
Each copper atom that is oxidised loses two electrons
Each Ag+ ion that is reduced gains one electron.
When writing full equations we must balance the electrons first.
Therefore two Ag+ ion must be reduced to take up the electrons lost by each copper atom that is oxidised.

21. Copper and the solution of silver ions
Cu(s) ―> Cu2+(aq) + 2e-
Ag+(aq) + e- ―> Ag(s)
So we need to times the silver ions by 2
The overall equation is:
Cu(s) + 2Ag+(aq) ―> Cu2+(aq) + 2Ag(s) (
) x 2

22. Remember In both half and overall equations.
The number of atoms of each element present in the products is equal to the number present in the reactants.
Atoms are conserved in all chemical equations.
The total charge on the product side of the equation is equal to the total charge on the reactant side of the equation.
Charge is conserved in chemical reactions.

23. Worked Example 16.2c
When sodium is oxidised by atmospheric oxygen, the reaction can be represented by the following half equations:
Na(s) ―> Na+(s) + e-
O2(g) + 4e- ―> 2O2-(s)
Identify the half equation representing the oxidation reaction and write the balanced overall equation.

24. Oxidants and Reductants
An oxidant (or oxidising agent) is a species that causes another to be oxidised.
A reductant (or reducing agent) is a species that causes another to be reduced.
The oxidant itself is reduced.
The reductant is oxidised.

25. Your Turn
Page 278
Question 3 and 4

26. Predicting electron transfer
Read pages 283 – 285
What is a galvanic cell?

27. Galvanic Cell All galvanic cells are composed of two half cells.
Oxidation occurs in one half cell.
Reduction occurs in the other.
A half cell must contain an electrode and an electrolyte.
An electrode is an electronic conductor – a material that has delocalised electrons that can move through the circuit.

28. Galvanic Cells
The electrode at which oxidation takes place is called the anode.
The electrode at which reduction takes place is called the cathode.

29. Galvanic Cells Zinc is the anode. Copper is the cathode.
In galvanic cells the anode is negatively charged and the cathode is positively charged.

30. Galvanic Cells Cu2+ ions are reduced to Cu atoms at the cathode.
Cations will migrate from the salt bridge into the beaker containing that cathode to compensate for the loss of the Cu2+ ions.
At the anode, zinc metal is oxidised and so more Zn2+ ions are added to the solution in that beaker.

31. The salt bridge
To avoid the build up of a positive charge, anions (negatively charged ions) will migrate from the salt bridge into the beaker and so maintain electrical neutrality.

32. Electrolyte
An electrolyte contains ions that are free to move through the solution.
In the example the electrolyte in beaker A was the zinc chloride.
The electrolyte in beaker B was the copper sulfate solution.

33. Galvanic Cells Comprise Of:
Two half cells, which are separate and do not mix.
A length of wire connecting the electrodes of the half cells. This is the external current.
A salt bridge to connect the solutions in the half cells. This is the electrical conductor.
The salt bridge balances the overall charge during the circuit.

34. The electrochemical series
Sodium, magnesium and iron are all metals that corrode easily because they are easily oxidised.
Sodium is oxidised so easily that it is stored under paraffin oil.
Other metals, however, do not corrode readily. Platinum and gold are sufficiently inert to be found free in nature.

35. The Electrochemical Series
Table 16.2 on page 287 represents the electrochemical series.
What can you tell me about the electrochemical series?

37. The electrochemical series
Each half equation represents the reduction reactions.
The top equation is the strongest oxidant so it is most easily reduced.
The strongest reductants are at the bottom and are oxidised quite easily. What kind of metals do these mainly consist of?
In general the smaller amount of energy required to remove a valance electron the more readily the metal will act as a reductant and itself be oxidised.

38. Electrochemical Series
Non-metals tend to gain electrons and therefore act as oxidants.
Reactive metals tend to be stronger reductants.
Transition metals are less readily oxidised.

39. Predicting Redox Reactions
We use the electrochemical series to predict redox reactions.
More reactive metals tend to be found on the lower right of the electrochemical series.
A more reactive metal will be oxidised by, and donate its electrons to the cation of a less reactive metal.
The cation receives the electrons and is reduced.

40. Predicting Redox Reactions
A spontaneous redox reaction can be expected to occur when a relatively strong oxidant is mixed with a relatively strong reductant.
The oxidant is reduced and the half equation occurs in the forward direction.
The reductant is reduced and the half equation occurs in the reverse direction of the that on the electrochemical series.

41. Predicting Redox Reactions
We can predict that zinc metal with react with Cu2+ ions because zinc is more reactive than copper.
Is reduced
What is the overall Equation????
Cu2+(aq)
+ 2e- ―> Cu(s)
Reacts with
Is oxidised
Zn(s)
Zn2+ + 2e- <―

42. AN OIL RIG CAT Anode + Oxidation is loss of electrons:
Reduction is gain of electrons + Cathode
A way to remember oxidation occurs at the anode. Reduction occurs at the cathode.

43. Predicting Reactions
For reactions to occur spontaneously, the aqueous cation in the solution must be a stronger oxidant than the cation of the metal added.
Your Turn
Try Question 13 on page 291
Try Question 15 as well

44. Your Turn
Finish reading this chapter yourself about corrosion.