1. Chemsheets AS006 (Electron arrangement)
14/04/2017
ELECTROCHEMISTRY
© A Jul-12

2. Zn2+(aq) e–  Zn(s)

4. Zn  Zn2+ + 2 e- oxidation Cu2+ + 2 e-  Cu reduction
- electrode
anode
oxidation
+ electrode
cathode
reduction
electron flow
At this electrode the metal loses electrons and so is oxidised to metal ions.
These electrons make the electrode negative.
At this electrode the metal ions gain electrons and so is reduced to metal atoms.
As electrons are used up, this makes the electrode positive. Zn Cu
Zn  Zn e-
oxidation
Cu e-  Cu
reduction

7. Standard Conditions Concentration
1.0 mol dm-3 (ions involved in ½ equation)
Temperature
298 K
Pressure
100 kPa (if gases involved in ½ equation)
Current
Zero (use high resistance voltmeter)

8. S tandard
H ydrogen
E lectrode

9. Emf = E = Eright - Eleft

12. Pt(s) | H2(g) | H+(aq) || Cu2+(aq) | Cu(s)

13. ROOR Ni(s) | Ni2+(aq) || Sn4+(aq), Sn2+(aq) | Pt(s)
K(s) | K+(aq) || Mg2+(aq) | Mg(s)

14. ELECTRODE POTENTIALS – Q1
Emf = Eright - Eleft
= Eright - 0
Eright = V

15. ELECTRODE POTENTIALS – Q2
Emf = Eright - Eleft
Emf =
Emf = V

16. ELECTRODE POTENTIALS – Q3
Emf = Eright - Eleft
Emf = (-0.76)
Emf = V

17. ELECTRODE POTENTIALS – Q4
Emf = Eright - Eleft
= Eleft
Eleft = = V

18. ELECTRODE POTENTIALS – Q5
Emf = Eright - Eleft
a) Emf = (-0.25) = V
b) Emf = = V
c) Emf = = V

19. ELECTRODE POTENTIALS – Q6
Emf = Eright - Eleft
a) Eright = = V
Ti3+(aq) + e-  Ti2+(aq)
b) Eleft = = V
K+(aq) + e-  K(aq)
c) Eright = = V
Ti3+(aq) + e-  Ti2+(aq)

20. ELECTRODE POTENTIALS – Q7
a) Cr(s) | Cr2+(aq) || Zn2+(aq) | Zn(s)
Emf = (-0.91) = V
b) Cu(s) |Cu2+(aq)|| Fe3+(aq),Fe2+(aq)| Pt(s)
Emf = = V
c) Pt(s) | Cl-(aq)| Cl2(g) || MnO4-(aq),H+(aq),Mn2+(aq)| Pt(s)
Emf = – 1.36 = V

22. The more +ve electrode gains electrons (+ charge attracts electrons)

23. – + e– + 1.10 V + 0.34 V – 0.76 V Cu2+ + Zn → Cu + Zn2++
–ve electrode
+ve electrode e– V V
Cu e-  Cu
– 0.76 V
Zn e-  Zn
Cu2+ + Zn → Cu + Zn2+

24. PREDICTING REDOX REACTIONS – Q1+
–ve electrode
+ve electrode e– V
– 0.25 V
Ni e-  Ni
– 0.76 V
Zn e-  Zn
Ni2+ + Zn → Ni + Zn2+

25. PREDICTING REDOX REACTIONS – Q2+
–ve electrode
+ve electrode e– V V
Ag+ + e-  Ag V
Cu e-  Cu
2 Ag+ + Cu → 2 Ag + Cu2+

26. PREDICTING REDOX REACTIONS – Q3 a
Mg(s)|Mg2+(aq)||V3+(aq),V2+(aq)|Pt(s) –
–ve electrode
+ve electrode e– V
– 0.26 V
V3+ + e-  V2+
– 2.36 V
Mg e-  Mg
YES: Mg reduces V3+ to V2+

27. PREDICTING REDOX REACTIONS – Q3 b+
–ve electrode
+ve electrode e– V V
Cl2 + 2 e-  2 Cl- V
Fe3+ + e-  Fe2+
NO: Cl- won’t reduce Fe3+ to Fe2+

28. PREDICTING REDOX REACTIONS – Q3 c
Pt(s)|Br-(aq),Br2(aq)||Cl2(g)|Cl-(aq)|Pt(s) +
–ve electrode
+ve electrode e– V V
Cl2 + 2 e-  2 Cl- V
Br2 + 2 e-  2 Br-
YES: Cl2 oxidises Br- to Br2

29. PREDICTING REDOX REACTIONS – Q3 d
Sn(s)|Sn2+(aq)||Fe3+(aq),Fe2+(aq)|Pt(s) – +
–ve electrode
+ve electrode e– V V
Fe3+ + e-  Fe2+
– 0.14 V
Sn e-  Sn
YES: Sn reduces Fe3+ to Fe2+

30. PREDICTING REDOX REACTIONS – Q3 e+
–ve electrode
+ve electrode e– V V V
Cl2 + 2 e-  2 Cl-
Cr2O H+ + 6 e-  2 Cr H2O
NO: H+/Cr2O72- won’t oxidise Cl- to Cl2

31. PREDICTING REDOX REACTIONS – Q3 f
Pt(s)|Cl-(aq)|Cl2(g)||MnO4- (aq),H+(aq),Mn2+(aq)|Pt(s) +
–ve electrode
+ve electrode e– V V V
MnO H+ + 5 e-  Mn H2O
Cl2 + 2 e-  2 Cl-
YES: H+/MnO4- oxidises Cl- to Cl2

32. PREDICTING REDOX REACTIONS – Q3 g
Fe(s)|Fe2+(aq)||H+(aq)|H2(g)|Pt(s) –
–ve electrode
+ve electrode e– V
0.00 V
2 H+ + 2 e-  H2
– 0.44 V
Fe e-  Fe
YES: H+ oxidises Fe to Fe2+

33. PREDICTING REDOX REACTIONS – Q3 h+
–ve electrode
+ve electrode e– V V
Cu e-  Cu
0.00 V
2 H+ + 2 e-  H2
NO: H+ won’t oxidise Cu to Cu2+

34. PREDICTING REDOX REACTIONS – Q4+ V
Cl2 + 2 e-  2 Cl- V V
Fe3+ + e-  Fe2+
MnO H+ + 5 e-  Mn H2O NO V
YES
Cr2O H+ + 6 e-  2 Cr H2O NO

35. PREDICTING REDOX REACTIONS – Q5a+
–ve electrode
2.19 = Eleft
+ve electrode
Eleft = 0.34 – 2.19 = – 1.85 V e– V V
Cu e-  Cu
? V
Be e-  Be
Be + Cu2+ → Be2+ + Cu

36. PREDICTING REDOX REACTIONS – Q5b
–ve electrode
When using SHE
+ve electrode
E = cell emf = – 1.90 V e–
1.90 V V
2 H+ + 2 e-  H2
? V
Th e-  Th
4 H+ + Th → 2 H2 + Th4+

37. PREDICTING REDOX REACTIONS – Q6a
Pt(s)|H2(g)|H+(aq)||Br2(aq),Br-(aq)|Pt(s) +
–ve electrode
+ve electrode e– V V
Br2 + 2 e-  2 Br-
0.00 V
2 H+ + 2 e-  H2
H2 + Br2 → 2 H+ + 2 Br-

38. PPT - Electrode potentials
PREDICTING REDOX REACTIONS – Q6b
14/04/2017
Cu(s)|Cu2+(aq)||Fe3+(aq),Fe2+(aq)|Pt(s) +
–ve electrode
+ve electrode e– V V
Fe3+ + e-  Fe2+ V
Cu e-  Cu
2 Fe3+ + Cu → 2 Fe2+ + Cu2+

39. Electrochemical cells
i appreciate that electrochemical cells can be used as a commercial source of electrical energy
j appreciate that cells can be non-rechargeable (irreversible), rechargeable and fuel cells
k be able to use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells and to deduce the e.m.f. of a cell
l understand the electrode reactions of a hydrogen-oxygen fuel cell and appreciate that a fuel cell does not need to be electrically recharged
m appreciate the benefits and risks to society associated with the use of these cells

40. Non-rechargeable (primary) cells – Zinc-carbon
-0.80 V Zn(NH3) e-  Zn NH3
+0.70 V 2 MnO H e-  Mn2O3 + H2O
Standard cell
Short life
Determine: a) cell emf
b) overall reaction during discharge

41. Non-rechargeable (primary) cells – alkaline
-0.76 V Zn e-  Zn
+0.84 V MnO2 + H2O + e-  MnO(OH) + OH-
Determine: a) cell emf
b) overall reaction during discharge
Longer life

42. Non-rechargeable (primary) cells – lithium
Very long life
High voltage
Determine: a) cell emf
b) overall reaction during discharge

43. Rechargeable (secondary) cells
In non-rechargeable (primary) cells, the chemicals are used up so the voltage drops
In rechargeable (secondary) cells the reactions are reversible – they are reversed by applying an external current.
It is important that the products from the forward reaction stick to the electrodes and are not dispersed into the electrolyte.

44. Rechargeable (secondary) cells – Li ion
+0.60 V Li+ + CoO2 + e-  LiCoO2
-3.00 V Li+ + e-  Li
Rechargeable
Most common rechargeable cell
Determine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge

45. Rechargeable (secondary) cells – lead-acid
+1.68 V PbO2 + 3 H+ + HSO e-  PbSO4 + 2 H2O
-0.36 V PbSO4 + H e-  Pb + HSO4-
Determine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge
Used in sealed car batteries (6 cells giving about 12 V overall)

46. Rechargeable (secondary) cells – nickel-cadmium
+0.52 V NiO(OH) + 2 H2O + 2 e-  Ni(OH)2 + 2 OH-
-0.88 V Cd(OH) e-  Cd OH-
Determine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge

47. FUEL CELLS High efficiency (more efficient than burning hydrogen)
How is H2 made?
Input of H2/O2 to replenish so no need to recharge
+0.40 V O2 + 2 H2O + 4 e-  4 OH-
-0.83 V 2 H2O e-  H OH-
Determine: a) cell emf b) overall reaction

49. From wikipedia (public domain)

50. Pros & cons of cells
+ portable source of electricity
Pros & cons of non-rechargeable cells
+ cheap, small
– waste issues
Pros & cons of rechargeable cells
+ less waste, cheaper in long run
– still some waste issues
Pros & cons of fuel cells
+ water is only product
– most H2 is made using fossil fuels, fuels cells expensive