1. IB topic 9 Oxidation-reduction
Define oxidation and reduction in terms of electron loss and gain.

2. Taking notes
Do not copy, write in your own words or draw diagrams or pictures
10 % better performance if you copy
35% better performance if you write in your own words

3. 2Mg + O2  2MgO
Reduction-charge goes down
OILRIG
Redox always occurs together
Question 1

4. A:  

5. B

6. C

7. D

8. 9.2Redox equations
Deduce simple oxidation and reduction half-equations given the species involved in a redox reaction.
2Fe + 3Cl2 2FeCl3

9. Deduce the oxidation number of an element in a compound.
+ means loss, – gain of e-
Rules
Elements Na, O2, S8 = 0
Group 1 = +1 H=+1
O = -2 halides -1
Many exceptions
Ox. # add up to the charge on the species
In covalent compounds more electronegative is – ie NH3, CCl4

10. Give Ox. numbers to each element
H2SO4, SO32-
NH4+, Fe2O3, K2Cr2O7, CuCl2,
Question 2
Question 3

11. State the names of compounds using oxidation numbers.
MnO2, FeO, CuCl, Na2O
Manganese (IV) oxide, iron (II) oxide, Copper (I) chloride, sodium oxide
[Cu(H2O)6]2+ [CuCl4]2-
Hexaaquacopper(II) ion
Tetrachlorocopper (II) ion

12. Deduce whether an element undergoes oxidation or reduction in reactions using oxidation numbers.
Ca + Sn2+  Ca2+ + Sn
4NH3 + 5O2 4NO + 6H2O
Disproportionation Cl2 + H2OHCl + HClO
Question 4

13. Deduce redox equations using half-equations.
Steps
Assign O numbers and write half reactions
Balance atoms other than H and O
Balance O by adding H2O as needed
Balance H by adding H+ as needed
Balance Charges by adding e- to the + side
Equalize the e- by multiplying
Add the half reactions together

14. Try NO3- + Cu  NO + Cu2+
+5, -2, 0 on left +2,-2,+2 on right
Cu  Cu2+ ox NO3- NO red
4H+ + NO3-  NO H2O
Cu  Cu2+ + 2e-
4H+ + NO3- + 3e- NO H2O
8H+ + 2NO3- + 6e- + 3Cu  NO + 4H2O + 3Cu2+ + 6e-

15. Fe+2 + MnO4-  Fe+3 + Mn+2

16. SO32- + Cr2O72- SO42- + Cr3+

17. Internet example
Question 5

18. Define the terms oxidizing agent and reducing agent.
A substance that gets reduced causes oxidation so it is an oxidizing agent

19. Identify the oxidizing and reducing agents in redox equations.
Fe2O3 + 3C  2Fe + 3CO2
Fe oxidizing C reducing
IO3- + 5I- + 6H+  3I2 +3H2O
IO3- oxidizing I- reducing
question 6

20. 9.3Reactivity
Deduce a reactivity series based on the chemical behavior of a group of oxidizing and reducing agents.
More reactive metals lose their e- more readily becoming a strong reducing agent
Zn + CuSO4
Stronger Mg, AL, Zn, Fe, Pb, Cu, Ag
simulation

21. Non metals F2 strongest oxidizing agent, most readily becomes reduced Cl2, Br2, I2
Question 7

22. Deduce the feasibility of a redox reaction from a given reactivity series.
Yes or no
ZnCl2 + Ag
2FeCl3 + 3 Mg
Cl2 + 2KI
question 8

23. 9:4 Voltaic Cells (battery)
Explain how a redox reaction is used to produce electricity in a voltaic cell
Zn(s) → Zn2+(aq) + 2e- red. Agent
Other half cell
Cu2+ + 2e- → Cu(s) reduced
This combination is a voltaic cell

24. State that oxidation occurs at the negative electrode (anode) and reduction occurs at the positive electrode (cathode)
Which is the anode (where e- leave) /cathode?
Draw this setup.
Animation
Draw Zn/Zn2+ and Ag/Ag+ and give the potential, show the flow of e-
Where is oxidation and reduction

25. Connect these half cells with a salt bridge
this is a spontaneous reaction
Animation
Question 9

26. 9:5 Electrolytic cells
Describe, using a diagram, the essential components of an electrolytic cell.
Opposite of a voltaic cell
Requires electrical energy
│‌ means + then – in diagrams
Animation

27. The power source pushes e- to the – electrode
-electrode attracts + ions
-electrode is the cathode cations gain e- so are reduced
Show the electrolysis of MgF2

28. Describe how current is conducted in an electrolytic cell
Do questions 10

29. Deduce the products of the electrolysis of a molten salt
Diagram the electrolysis of molten(melted) NaCl
Tell where oxidation and reduction occurs
Do question 11

30. 19.1 Standard electrode potentials
Describe the standard hydrogen electrode.

31. Standard H cell Conditions - Pt electrode H2 gas at 1 atm pressure
1 mol dm-3 H+
298 K or 25oC
0.00 V
Attach a half cell if e- flows to H2 it is –
Like Zn which is V

32. conventions Zn(s)/Zn+2││H+(aq)/1/2 H2(g) (Pt) Oxidation on left side
More – value of electrode potentials give off e-

33. Define the term standard electrode potential (E Ö ) .
relative electrode potential compared under standard conditions with the standard hydrogen electrode
Look at your data booklet

34. 19.1.3 Calculate cell potentials using standard electrode potentials.
Try Cr2O72- and Br2
Answer 0.26 V

35. Predict whether a reaction will be spontaneous using standard electrode potential values.
Can a solution of tin II ions reduce a solution of iron III ions?
Yes 0.91V

36. Can a solution of Sn4+ ions reduce Fe3+ to Fe
Sn e- → Sn2+ Eo = +1.33
No what does work
Do question 12

37. 19.2 Electrolysis
Predict and explain the products of electrolysis of aqueous solutions.
For water need DC in a dilute solution of H2SO4
H+ to H2 given off at the – electrode
OH- to O2 at the + electrode
2H2O  4H+ + O2 + 4 e-

38. Electrolysis of NaCl(aq)
- electrode H2
+ electrode dilute OH- to O2 conc Cl- to Cl2
Write half reactions
Do question 13

39. Determine the relative amounts of the products formed during electrolysis.
Position in the electrochemical series
+ ions lower in the series will gain e- at the – electrode (cathode) in preference to those higher
Hydroxide ions release e- to form oxygen and H2O in preference to other anions at the positive electrode

40. In some cases concentrations ( more concentrated may be discharged)
Nature of the electrode C and Pt are inert
List all the cations and anions
Cations lower in the series gain e- more readily

41. Describe the use of electrolysis in electroplating.
CuSO4(aq) with copper electrodes
Cu2+ goes to – electrode and plates Cu
+ electrode Cu goes to Cu2+(use impure ore)

42. Electroplating
Object to be electroplated is put at the negative electrode and is placed in a solution of ions of the metal used to plate it.

43. Factors affecting relative amounts
Charge on the ions
Na+ Cu2+ Al3+ Al takes more energy to make
Quantity of e- (amperage and time) charge = current x time
Do question 14

44. Do questions 1-14 on chapter 10 in your IB Study Guide and turn in