1. Redox Chemistry and Corrosion Chapter 16

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) + O 2 (g) ―› CO 2 (g)  S(s) + O 2 (g) ―› SO 2 (g)  4Fe(s) + 3O 2 (g) ―› 2Fe 2 O 3 (s)  2C 8 H 18 (l) + 25O 2 (g) ―›16CO 2 (g) + 18H 2 O(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 (Fe 2 O 3 ) and magnetite (Fe 3 O 4 ))

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: Fe 2 O 3 (s) + 3CO(g) ―› 2Fe(s) + 3CO 2 (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) + O 2 (g) ―> 2MgO(s)

14. ► Magnesium oxide is an ionic compound and consists of Mg 2+ ions and O 2- ions. ► Each magnesium ion, therefore must have lost two electrons to form an Mg 2+ ion. Each oxygen atom in the oxygen molecule must have gained two electrons to form an oxide ion O 2-. ► The reaction can now be represented by two half equations.

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

16. Mg(s) ―> Mg 2+ (s) + 2e - O 2 (g) + 4e - ―> 2O 2- (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. Cu(s) ―> Cu 2+ (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) ―> Cu 2+ (aq) + 2Ag(s) Copper and the solution of silver ions ( ) 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 - O 2 (g) + 4e - ―> 2O 2- (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 ► Cu 2+ 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 Cu 2+ ions. ► At the anode, zinc metal is oxidised and so more Zn 2+ 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 Cu 2+ ions because zinc is more reactive than copper. Cu 2+ (aq) Is reduced + 2e- ―> Cu(s) Zn(s) Is oxidised Zn 2+ + 2e- <― Reacts with What is the overall Equation????

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.