1. Chapter 22 REDOX

2. Chapter 22 REDOX The Meaning of Oxidation and Reduction
Oxidation Numbers
Balancing Redox Equations

3. Ch 22.1 The Meaning of Oxidation and Reduction
Oxygen in Redox Reactions
Electron Transfer in Redox Reactions
Corrosion

4. Oxygen in Redox Reactions
Oxidation – the combination of an element with oxygen to produce oxides

5. Oxygen in Redox Reactions
Burning

6. Oxygen in Redox Reactions
Bleaching

7. Oxygen in Redox Reactions
Rusting

8. Reduction
Reduction – the loss of oxygen from a compound

9. Redox Reactions Reduction and Oxidation always occur together
2Fe2O3(s) + 3C(s)  4Fe(s) + 3CO2(g)
reduction oxidation

10. Electron Transfer in Redox Reactions
Oxidation – loss of electrons, gain oxygen
Reduction – gain of electrons, loss of oxygen
“LEO the lion goes GER”
LEO – Lose electrons oxidation
GER – Gain electrons reduction

11. Electron Transfer in Redox Reactions
Oxidation: Mg  Mg2+ + 2e-
Loss of electrons
Reduction: S + 2e-  S2-
Gain of electrons

12. Corrosion

13. Corrosion 2Fe(s) + O2(g) + 2H2O(l)  2Fe(OH)2(s)
4Fe(OH)2(s) + O2(g) + 2H2O(l)  4Fe(OH)3(s)
Corrosion of iron

14. Corrosion Some metals completely corrode
Iron
Some metals form a protective coating
Aluminum
Some metals do not corrode at all
Gold

18. Chapter 22.3 Balancing Redox Reactions
Identifying Redox Reactions
Using Oxidation Number Changes
Using Half Reactions

19. Identifying Redox Reactions
Two types of reactions:
REDOX – electrons are transferred
Everything else: single replacement, double replacement, combustion, ….
NO transfer of electrons

20. Identifying Redox Reactions
REDOX – the oxidation number of an element changes
N2(g) + O2(g)  2NO(g)

21. Using Oxidation Number Changes
Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
Step 1 – Assign oxidation numbers to all atoms in the equation –

22. Using Oxidation Number Changes
Step 2 – Identify which atoms are oxidized and reduced
Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
Iron – reduced, Carbon - oxidized

23. Using Oxidation Number Changes
Step 3 – Use a bracket line to connect the atoms undergoing oxidation and one to connect the lines undergoing reduction
Fe2O3(s) + CO(g)  Fe(s) + CO2(g) +2 -3

24. Using Oxidation Number Changes
Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients
Fe2O3(s) + CO(g)  Fe(s) + CO2(g)
3 x (+2) = 6
2 x (-3) = - 6

25. Using Oxidation Number Changes
Step 5 – Finally make sure the equation is balanced for both atoms and charge
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)

31. FRH – Flameless Ration Heater
Mg + H2O  Mg(OH)2 + H2 + Heat
Problem: Mg forms a coating from corrosion – MgO – which is not water soluble, prevents the above reaction from happening
Solution: Add NaCl and Fe to the mix, breaks down the MgO and allows the reaction to happen

32. Chapter 23 Electrochemistry
Electrochemical Cells
Half Cells and Cell Potentials
Electrolytic Cells

33. Electrochemical Cells
The Nature of Electrochemical Cells
Voltaic Cells
Dry Cells
Lead Storage Batteries
Fuel Cells

34. The Nature of Electrochemical Cells
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

35. The Nature of Electrochemical Cells
The zinc bar becomes copper plated
Zinc loses electrons and dissolves slowly
Copper gains electrons and becomes a solid
Oxidation: Zn(s) Zn2+(aq) + 2e-
Reduction: Cu2+(aq) + 2e- Cu(s)

36. The Nature of Electrochemical Cells
Reference Table J
Look at any two metals, the metal that is higher on the table is the one that is more readily oxidized

37. Electrochemical Cell
Any device that converts chemical energy into electrical energy or electrical energy into chemical energy
REDOX reactions must occur
If an electrochemical cell is to be used for electrical energy, the two half reactions must physically be separated

38. Voltaic Cell
Alessandro Volta (1745 – 1827)
First electrochemical cell

39. Voltaic Cell Convert chemical energy into electrical energy
Half Cell – part of a voltaic cell, consists of a metal rod in a solution of ions
Salt Bridge – Separates half cells, tube containing a strong electrolyte (can also use a porous plate)

41. Voltaic Cell Anode – the anode where oxidation occurs
Cathode – the cathode where reduction occurs

42. Dry Cell
A voltaic cell in which the electrolyte is a paste

43. Alkaline Battery
Improved dry cell, the Zinc electrode doesn’t corrode as fast

44. Lead Storage Batteries
                                                                                           
A group of cells connected together

45. Fuel Cell Voltaic cell in which a fuel substance undergoes oxidation
Do not have to be recharged
Oxidation: 2H2(g) + 4OH-(aq)  4H2O(l) + 4e-
Reduction: O2(g) + 2H2O(l) + 4e-  4OH-(aq)

48. Hydrogen Refueling Stations

52. The diagram shows a voltaic cell with copper and aluminum electrodes immediately after the external circuit is completed.
1 Balance the redox equation using the smallest whole-number coefficients. [1]
2 As this voltaic cell operates, the mass of the Al(s) electrode decreases. Explain, in terms of particles, why this decrease in mass occurs. [1]
3 Explain the function of the salt bridge. [1]

53. Answers 3 Cu2+ (aq) + 2 Al(s)  3 Cu(s) + 2 Al3+ (aq)
Aluminum particles are losing electrons and becoming aluminum ions that are entering the solution.
It allows migration of ions, maintains neutrality, prevents polarization

54. Electrolytic Cells
An electrochemical cell used to cause chemical change through the application of electrical energy (electrical energy is added)

56. Differences Flow of electrons is spontaneous
Voltaic (Galvanic) Cells
Electrolytic Cells
Flow of electrons is spontaneous
Flow of electrons is pushed by an outside power source
Anode negative
Cathode positive
Anode positive
Cathode negative

57. Similarities Voltaic (Galvanic) Cells and Electrolytic Cells
Electrons flow from anode to cathode
Reduction – cathode
Oxidation – anode

58. Electroplating is an electrolytic process used to coat metal objects with a more expensive and less reactive metal. The diagram below shows an electroplating cell that includes a battery connected to a silver bar and a metal spoon. The bar and spoon are submerged in AgNO3(aq).

59. Explain why AgNO3 is a better choice than AgCl for use in this electrolytic process. [1]
Explain the purpose of the battery in this cell. [1]

60. Acceptable responses include, but are not limited to:
Silver nitrate produces more ions than silver chloride in water.
AgNO3 readily dissolves in H2O; AgCl dissolves only slightly in H2O.
The battery provides the electrical energy necessary for the reaction to occur.

61. The apparatus shown in the diagram consists of two inert platinum electrodes immersed in water. A small amount of an electrolyte, H2SO4, must be added to the water for the reaction to take place. The electrodes are connected to a source that supplies electricity.

62. What type of electrochemical cell is shown? [1]
What particles are provided by the electrolyte that allow an electric current to flow? [1]

63. Electrolytic or electrolysis.
Acceptable responses include, but are not limited to:
Ions, charged particles, H3O+, SO42–

64. Because tap water is slightly acidic, water pipes made of iron corrode over time, as shown by the balanced ionic equation below:
2Fe + 6H+  2Fe3+ + 3H2
Explain, in terms of chemical reactivity, why copper pipes are less likely to corrode than iron pipes. [1]

65. Acceptable responses include, but are not limited to:
Copper is less reactive than iron.
Cu below H2 on Table J