Redox Chemistry and Corrosion

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) + 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.