1. Chapter 16 – Redox Chemistry & Corrosion
Week 8, Lesson 1
Chapter 16 – Redox Chemistry & Corrosion

2. Oxidation & Reduction
Redox reactions are reactions that are called oxidation-reduction reactions.
Many of the chemical reactions that play a significant role in maintaining our environment are redox reactions.
For example, the reactions of water and the gases of the atmosphere that cause corrosion and so the deterioration of metals are redox reactions.

3. Redox Reactions 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 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 electrolytes to produce many chemicals
Production and us of fertilisers.

4. Redox Reactions cont…
Once the reactive component of air was first produced by Joseph Priestley identified by Antoine Lavoisier as the element 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.
This type of reaction was classified as oxidation, which here means the addition of oxygen.

5. Redox Reactions cont…
When oxygen reacts with an element, the element is said to be oxidised.
The iron that society uses has been extracted from iron ores.
This process involves reduction of the iron oxide to iron.
When oxygen has been removed from a substance the material has been reduced.

6. Redox Reactions Example…
The iron (III) oxide has lost oxygen – it has been reduced.
Reduction can not occur without oxidation occurring at the same time.
In this reaction it is the carbon monoxide that has gained oxygen – and it is said to have been oxidised.
Oxidation and reduction always occur simultaneously.

7. An Improved Definition – Electron Transfer
While the definition of oxidation as the addition of oxygen was adequate for combustion reactions occuring in air, many similar reactions not involving oxygen as a reactant have since been investigated.
Magnesium ribbon is often dull and covered with a coating that needs to be scraped of. This is magnesium oxide which results from the corrosion of Mg in air.
The magnesium reacted with atmospheric oxygen to form magnesium oxide – the magnesium has been oxidised.
This reaction is represented by the equation:
2Mg(s) + O2  2MgO(s)

8. Electron Transfer cont…
Magnesium oxide is an ionic compound and consists of Mg2+ and O2- ions.
Each magnesium atom has lost 2 electrons to form this ion.
Each oxygen molecule has gained 2 electrons to form this ion.
The reaction can be represented by two half equations.
The first shows the loss of electrons:
Mg(s)  Mg2+(s) + 2e-
The second shows the gain of electrons:
O2(g) + 4e-  2O2-(s)

9. Electron Transfer cont…
Therefore oxidation is defined as the loss of electrons.
Reduction is defined as the gain of electrons.
Reactions that involve oxidation and reduction are called redox reactions.

10. Writing Half Equations example…
When a strip of copper wire is suspended in a solution of silver nitrate, long crystals of silver metal can be observed. The solution changes to a pale blue colour, indicating the presence of Cu2+ ions. Write half equations for this reaction and identify the substance oxidised and reduced. STEP 1: The first equation is the copper metal forming copper ions. Cu(s) Cu2+(aq) + 2e- STEP 2: The second equation involves the silver ions form silver metal. Ag+(aq) + e-  Ag(s) Since the copper has lost electrons is has undergone oxidation and the silver has undergone reduction.

11. Writing an Overall Redox Reaction
When writing equations for redox reactions, normally the two half equation are written first and then they are added together to produce an overall equation.
In an overall equation, no electron transfer is shown.
That is, the electrons that are lost in oxidation, are gained in the reduction reaction.

12. Writing an Overall Redox Reaction Example…
In the example of copper and the solution of silver ions:
Each copper atom that is oxidised loses two electrons.
Each silver ion that is reduced gains on electron.
Two silver ions must be reduced to take up the electrons lost by each copper atom oxidised.
To write an overall equation, the half equation involving the reduction of silver is multiplied by 2, before it is combined with the half equation for the oxidation of copper.
Cu(s)Cu2+(aq) + 2e-
(Ag+(aq) + e-  Ag(s)) x 2
The overall equation is
Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)

13. Overall Redox Reactions cont…
Remember in both half and full redox 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 reactions.
The total charge on the product side of the equation is equal to the total charge on the reactants side of the equation. Charge is conserved in all reactions.

14. Oxidants and Reductants
An oxidant (or oxidising agent) is a species that causes another to be oxidised.
A reductant is a species that causes another to be reduced.
An oxidant causes oxidation and is reduced.
A reductant causes reduction and is oxidised.

15. Oxidants and Reductants cont…

16. Week 8, Lesson 2
Oxidation Numbers
To determine whether or not an equation is a redox reaction and to help us identify oxidants and reactants, oxidation numbers are assigned to elements involved in a reaction.
These numbers are determined by applying a set of rules that consider unequal sharing of electrons between two atoms.

17. Oxidation Number Rules
Oxidation numbers are determine according to the following rules in ascending order:
The oxidation number of a free element is 0.
The oxidation number of a simple ions is equal to the charge on that ion.
In compounds, some elements have oxidation numbers that are regarded as fixed, except in a few exceptional circumstances.
Main group elements have an oxidation number equal to their valency.
Fluorine has an oxidation number of -1.
Hydrogen has an oxidation number of +1 when is forms compounds with non metals and -1 in metal hydrides.
Oxygen almost always has an oxidation number of -2. In F2O it has an oxidation number of +2 and in peroxides such as H2O2 and BaO2 its oxidation number is -1.
4. The sum of oxidation numbers in a neutral compound is zero.
5. The sum of oxidation number in a polyatomic ion is equal to the charge on the ion.

18. Using Oxidation Numbers
A change in oxidation numbers is an indication that a redox reaction have taken place.
Oxidation involves an increase in oxidation number.
Reduction involves a decrease in oxidation number.

19. Half Equations for Complex Redox Reactions
Week 8, Lesson 3
Half Equations for Complex Redox Reactions
Look at the half equation for the oxidation of iron (II) ions to iron (III) ions:
Fe2+(aq)  Fe3+(aq) + e-
The conversion of MnO4- to Mn2+ involves reduction.
The following steps are used to balance this half equation:
Balance all atoms in the half equation except hydrogen and oxygen.
Balance the oxygen by adding water.
Balance the hydrogen atoms by adding H+
Balance the charges on both sides by adding electrons to the more positive side.
Then add the oxidation half equation to the reduction half equation to get the overall equation.
Make sure the number of electrons used in reduction equals the number of electrons released during oxidation.

20. Evidence for Electron Transfer
Week 9, Lesson 1
Evidence for Electron Transfer
When zinc metal is placed in a solution of copper (II) sulfate, the zinc is oxidised and a brown deposit of copper metal forms.
Electrons flow from the zinc atoms to the copper(II) ions as these collide.
Evidence for this flow can be obtained by setting up an electrochemical or galvanic cell.
A galvanic cell is an energy converter that separates the oxidant from the reductant.

21. Galvanic Cells
The zinc metal and solution containing copper(II) ions are in separate beakers.
A copper strip, dipping in the solution of copper (II) ions is connected by a piece of wire to a galvanometer.
A galvanometer is a meter for detecting a flow of electrons.
A zinc strip, dipping into a solution of zinc chloride is also connected to this by a piece of wire.
The solutions in the two beakers are connected by a salt bridge.
The salt bridge contains an ionic compound.
The ions flow through the salt bridge between the solutions to complete the circuit.
The salt bridge can be a piece of filter paper soaked in potassium nitrate.

22. Galvanic Cells cont…
When this particular galvanic cell operates, the galvanometer indicates that electrons are flowing from the zinc strip in beaker B to the copper strip in beaker A.
This observation provides evidence that the oxidation and reduction reactions between zinc atoms and copper ions also occur what reactants are separated in a galvanic cell.
It also supports the idea that there is a transfer of electrons.

23. Zinc and Copper Galvanic Cell
At the copper strip, in beaker A, Cu2+ ions accept electrons to form copper atoms:
Cu2+(aq) + 2e-  Cu(s)
At the zinc strip in beaker B, zinc atoms lose electrons to form zinc ions:
Zn(s)  Zn2+(aq) + 2e-
Because of these reactions, the electrons flow through the zinc strip and connecting wire towards the copper strip.
The zinc is the source of negatively charged electrons so it is assigned a negative charge.
The electrons are attracted to the copper strip so this is assigned a positive charge.

24. Salt Bridge
To maintain the balance of positive and negative charge in each beaker, ions flow from the salt bridge into the two solutions.
In the beaker containing the copper sulfate solution, Cu2+ ions are reduced to Cu atoms, leaving SO42- ions in the solution.
The K+ ions flow from the salt bridge into this beaker to replace the Cu2+ ions.
At the same time, NO3- ions flow into the beaker containing ZnCl2 solution to balance the positive charge of the Zn2+ ions being produced at the zinc electrode.

25. Galvanic Cells All galvanic cells are composed of two half cells.
Oxidation occurs in one and reduction in the other.
A half cell must contain al electrode and electrolyte.
An electrode is an electronic conductor – a material that can move through the circuit.
The electrode at which oxidation takes place is called the anode.
The electrode at which reduction takes place is call the cathode.

26. Galvanic Cells cont…
In galvanic cells the anode is the negatively charged electrode.
The cathode is the positively charged electrode.
In the previous example, the copper ions are reduced to copper atoms at the cathode.
Cations will migrate from the salt bridge in the beaker containing the cathode to compensate for the loss of copper ions.
At the anode, zinc metal is oxidised and so more zinc ions are added to the solution in the beaker.
To avoid a build up of positive charge, anions will migrate from the salt bridge into the beaker and so maintain electrical neutrality.
An electrolyte contains ions that are free to move through the solution.

27. Galvanic Cell

28. Electrochemical Series
Sodium, magnesium and iron are metals that easily corrode because they are easily oxidised.
Chemists have combined such information with experimental data from constructing galvanic cells from various combinations of half cells to produce the electrochemical series.
The electrochemical series lists half equations in order of their tendency to occur as reduction reactions.

29. Electrochemical Series cont…

30. Electrochemical Series cont…
The half cell represented by the equation at the top of the list attracts electrons most strongly.
It will form the positive half cell when it is connected to any other half cell represented in the series.
An electrochemical series shows the relative reactivities of a number of metals and their corresponding.
The most reactive metals are those that are oxidised most easily; the most reactive metal cations are those that are reduced most readily.

31. Electrochemical Series cont…
As you go down the series, the:
Metals, which are on the right hand side, become increasingly easy to oxidise and therefore are more reactive. Metals lower in the series lose electrons more readily and so are stronger reductants.
Metal cations which are on the left hand side, become increasingly hard to reduce and therefore are less reactive. Cations higher in the series have a greater attraction for electrons and so are more readily reduced, and therefore are strong oxidants.

32. Predicting Redox Reactions
The electrochemical series is useful for predicting redox reactions.
More reactive metals tend to be found on the lower right of the 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.
A spontaneous 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 reaction occurs in the forwards direction.
The reductant is oxidised and its half equation is the reverse of that on the electrochemical series.

33. Predicting Redox Reactions cont…
Is reduced
Oxidant
Reacts with
Reductant
Is oxidised

34. Predicting Reactions in Galvanic Cells
An electrochemical series can be used to predict the processes that occur when two half cells are joined in an electrochemical cell.
The galvanic cell on the next slide is made by joining Ag+/Ag half cell with the Cu2+/Cu half cell and linking them with a salt bridge.
In this galvanic cell the two possible oxidants are Ag+ and Cu2+ and the two possible reductants are Cu and Ag.

35. Predicting Reactions in Galvanic Cells cont…
Ag+(aq) is higher on the left-hand side of the electrochemical series than Cu2+ and so is the stronger of the two oxidants.
This means that Ag+ attracts electrons more strongly than Cu2+, so electrons in the wire will flow away from the copper electrode towards the silver electrode.
Since oxidation occurs at the anode and the reduction at the cathode, the half cell reactions are:
At the cathode: Ag+(aq) + e-  Ag(s)
At the anode: Cu(s)  Cu2+(aq) + 2e-

36. AN OIL RIG CAT
Use the following phrase to help you remember the process:
ANode: Oxidation Is Loss of electrons Reduction Is Gain of electrons CAThode.

37. Week 9, Lesson 2
The Corrosion Process
Direct reaction with oxygen in the air to form a metal ocise is known as dry corrosion.
In cases such as aluminium, dry corrosion forms a tough, impervious oxide coating that protects the metal underneath from further contact from oxygen.
When moisture is present, the corrosion of iron quickens.
This process is known as wet corrosion.
The most well known form of wet corrosion is the formation of rust.

38. Wet Corrosion
Wet corrosion can occur in moist air or by direct immersion in water.
Several factors affect the rate at which wet corrosion occurs.
In general, corrosion is accelerated in the presence of:
Water – corrosion in dry air is very slow
Impurities such as salt and acidic pollutants that dissolve in the water
Impurities such as carbon in the iron
Corrosion is reduced however when the iron:
Is alloyed with certain other materials
Has a protective coating

39. The Wet Corrosion Process
Analysis of rust shows that it is a hydrated oxide of iron with the formula Fe2O3.xH2O.
x represents a number between 1 and 3 and xH2O indicates that this is the number of water molecules associated with the iron oxide.

40. The Wet Corrosion Process cont…
Experimental results indicate that the first step in the formation of rust is the oxidation of iron to form Fe2+ ions.
These ions can be detected in solution on the surface of the iron.
An Fe2+ ion has two less electrons than and iron atom.
This step therefore involves the loss of electrons and can be represented by the half equation:
Fe(s)  Fe2+(aq) + 2e-

41. The Wet Corrosion Process cont…
This only represents have the equation. The electrons lost by the iron must be accepted by another reactant.
Oxygen is also involved in corrosion – it gains the electrons lost by the iron.
The oxygen is reduced in the presence of water to hydroxide ions, represented by:
O2(aq) + 2H2O(l) + 4e-  4OH-(aq)

42. The Wet Corrosion Process cont…
The overall equation can be written as:
2Fe(s) + O2(aq) + 2H2O(l)  2Fe2+(aq) + 4OH-(aq)

43. The Wet Corrosion Process cont…
The iron (II) ions then react with the hydroxide ions to form a green precipitate of iron(II) hydroxide.
On exposure to more moisture and oxygen, the iron(II) hydroxide is further oxidised to iron(III) hydroxide, a red brown precipitate.

44. The Wet Corrosion Process cont…
The wet corrosion of iron is a galvanic process.
Galvanic cells are set up on the metal surface where different regions act as anodes and cathodes:
The cathode region, where oxygen gains electrons and is reduced, occurs at sites of high oxygen concentration.
The anode region, where iron loses electrons and is oxidised, oftern occurs in an area of stress in the metal where the metal lattice has been deformed, for example at the tip of a nail.

45. The Wet Corrosion Process cont…
The anode region is one where there is a lower concentration of oxygen.
Corrosion is therefore greatest at the centre of a water drop or under a layer of chipped paint.
The cathode region is at the edge of the water drop or where paint is missing as the oxygen concentration is the greatest.
Wet corrosion is accelerated by the presence of dissolved salts, if the iron contains impurities or if it is in less contact with a less reactive metal.

46. Protecting Against Corrosion
There are several ways to protect iron and steel structures against corrosion.
These depend on the environment in which the structure is located, the design and intended use of the structure, as well as the relative costs of various protection measures.
The simplest of these methods involves creating a barrier between the iron and oxygen and water to provide surface protection.
Other methods are based on an understanding of the electrochemical nature of corrosion.

47. Surface Protection
Involves covering the surface of the iron to prevent contact with oxygen and moisture.
Materials such as paint and plastic can be used for this purpose.
Iron can also be protected by alloying with small quantities of metals such as chromium, nickel, manganese and molybdenum to produce stainless steel.
The atoms of these metals used to make the allot are all bonded into the metallic lattice.
These metals oxidise slightly in air, but the oxide coating produced is continuous and unreactive.
This oxide layer forms a dense, impervious barrier between the metal and oxygen and water, protecting the iron from further oxidation.

48. Surface Protection cont…
Metals such as iron and steel can also be coated with thin layers of other metals in a process known as electroplating.
These other metals, which include chromium and nickel, again form protective oxide coatings.

49. Electrochemical Protection – Impressed Current Cathodic Protection
Impressed current protection involved reversing the direction of electron flow by using a low voltage direct current power supply to give the iron being protected a negative charge.
Because the iron is receiving electrons, oxidation is inhibited (oxidation would produce more electrons).
The iron becomes the site of the reduction reaction, that is the cathode.
This method is used to protect large scale steel structures such as wharves and pipelines.

50. Electrochemical Protection – Sacrificial Protection
This also involves the iron acting as the cathode.
A more easily oxidised metal, such as zinc, forms the anode and is sacrificed.
The more reactive metal loses electrons and forms metal cations in preference to the iron.
The iron is not oxidised in the presence of zinc because the zinc has protected the iron from corrosion.
This type of protection is known as electrochemical or sacrificial protection.
This form of protection is used in roofing, galvanised iron is iron with a zinc coating.