1. Electricity from Chemical Reactions

2. Electrochemistry
The production of electrical energy from chemical reactions
Redox reactions involve the transfer of electrons
Redox means that reduction and oxidation are occurring simultaneously

3. Reduction
Occurs when there is a decrease in oxidation number Zn2+  Zn
Gains electrons
Loses Oxygen
Converting a complex substance into a simpler form i.e. smelting iron to produce the pure metal iron

4. Oxidation
Occurs where there is an increase in oxidation number Zn  Zn2+
Loses electrons
Gains oxygen
The reaction used to describe the reaction of any substance with oxygen

5. Determining Oxidation Numbers
The atoms in elements have an Oxidation Number of zero eg Fe, C, Cl2
For a neutral molecule, the sum of the oxidation numbers are zero eg CO2
For a monatomic ion, the oxidation number is the same as it’s charge Cl – , Na+

6. Determining Oxidation Numbers
Oxygen usually takes – 2 in compounds. In peroxides (H2O2 & BaO2) it is – 1
Hydrogen takes + 1 in compounds, except in hydrides (NaH, CaH2) where it takes – 1

7. Determining Oxidation Numbers
For a polyatomic ion, the sum of the oxidation numbers of its component atoms is the same as its charges
For polyatomic molecules or ions, the, most electronegative element has a negative oxidation number and the least electronegative element has a positive oxidation number

8. Redox Half Reactions
Consider the reaction when a strip of zinc is dropped in a solution of Copper Sulphate
Zn(s) + Cu 2+(aq)  Zn2+(aq) + Cu(s)
Electrons are transferred from zinc atoms to copper ions
Reaction occurs spontaneously, that is with no external force or energy being applied

9. Redox Half Reactions Redox reactions consist of two half reactions
Oxidation Zn(s)  Zn2+(aq) + 2e–1
Reduction Cu 2+(aq) + 2e–1  Cu(s)
It is possible to use redox reactions to produce electricity

10. Galvanic Cells Also called Electrochemical Cells
Achieved by separating the half equations into half cells
Transferred electrons are forced to pass through an external circuit
Such an apparatus is called a Galvanic Cell

11. Galvanic Cells – + Zn2+ Cu2+ Flow of electrons zinc copper
Salt
bridge
Zn2+
Cu2+
Negative Electrode
(ANODE)
Positive Electrode
(CATHODE)

12. Standard Electrode Potentials
The electrical potential of a galvanic cell is the ability of the cell to produce an electric current.
Electrical potential is measured in volts
Cannot measure the electrode potential of an isolated half cell
Can measure the difference in in potential between two connected half cells

13. Standard Electrode Potentials
Electrical potential of a cell results from competition between 2 half cells for electrons
Half cell with the greatest tendency to attract electrons will undergo REDUCTION
Other half cell will lose electrons and undergo OXIDATION

14. Standard Electrode Potentials
The Reduction Potential of a half cell is a measure of the tendency of the oxidant to accept electrons and so undergo reduction
The difference between the reduction potentials of the two half cells is called the Cell Potential Difference

15. Standard Electrode Potentials
The Standard Cell Potential Difference (E0 cell) is the measured cell potential difference when the concentration of each species = 1M, pressure = 1 atm and Temp = 25 C
E0 cell = E0 oxidant – E0 reductant

16. Standard Electrode Potentials
A Standard Hydrogen Half cell is used as a comparative measure the reduction potentials of other cells
The SHE is given a value of 0.00 V
All other half cells are given a reduction potential value in comparison to this SHE by being connected to it

17. Standard Hydrogen Electrode
Glass
sleeve
Platinum wire
H2 gas
(1 Atm)
Salt Bridge to
Other half-cell
1.00M
Acid solution
Platinum electrode

18. Standard Hydrogen Electrode
SHE is used to measure reduction potential of other cells
If a species accepts electrons more readily than hydrogen, its electrode potential is positive
If a species accepts electrons less readily than hydrogen, its electrode potential is negative

19. Electrochemical Series
The reaction that is higher on the electrochemical series will occur as it appears and will reverse the direction of the reaction that occurs lower on the table

20. Potential Difference Is measured by a volt meter
Can be estimated by using electrochemical series
Connect Mg2+/Mg and Cl2/Cl– half cells get a potential difference of 3.7V
Looking at the electrochemical series

21. Potential Difference Cl2 + 2e–  Cl– has an E0 of 1.36V
Mg2+ + 2e–  Mg has an E0 of – 2.38V
The potential difference can be calculated
1.36 – (– 2.38) = 3.74V

22. Galvanic Cells Primary Cells Secondary Cells
Produce energy until one component is used up, then discarded
Secondary Cells
Store energy and may be recharged

23. Primary Cells
Dry Cells
Alkaline Cells
Button Cells

24. Dry Cells The ordinary zinc – carbon cell Anode oxidation (–)
Zn (s)  Zn 2+ (aq) + 2e –
Cathode oxidation (+)
2MnO2 (s) + NH4+ (aq) + 2e–  Mn2O2 (s) + 2NH3 (aq) + H2O (l)

25. Dry Cells The new cell produces about 1.5V
Once reaction reaches equilibrium its “flat”

26. Dry Cell Metal Cap (+) Mixture of Carbon & Manganese Dioxide
Cathode Carbon Rod
Ammonium Chloride & Zinc Chloride Electrolyte
Anode Zinc Case (–)

27. Alkaline Cells The ordinary zinc – carbon cell Anode oxidation (–)
Zn (s)  Zn 2+ (aq) + 2e –
Immediately reacts with OH – ions in the electrolyte to form zinc hydroxide
Zn (s) + 2OH –(aq)  Zn(OH)2 (s) + 2e –

28. Alkaline Cells Cathode reduction (+)
2MnO2 (s) + H2O(l) + 2e–  MnO2 (s) + OH –(aq) + H2O (l)
Five times the life of the dry cell

29. Alkaline Cell Metal Cap (+) Cathode outer steel case
Potassium Hydroxide Electrolyte
Powdered Zinc
Anode Steel or Brass
Mixture of Carbon & Manganese Dioxide
Metal Base (–)

30. Button Cells
Used in very small applications like watches, cameras etc.
Two main types
Mercury zinc and silver zinc
Anode Oxidation (–)
Zn (s) + 2OH –(aq)  Zn(OH)2 (s) + 2e –

31. Button Cells Cathode Reduction (+) depends on the type of battery
HgO(s) + H2O (l) + 2e –  Hg (l) + 2OH –(aq)
Ag2O(s)+H2O (l) + 2e –  2Ag (s) + 2OH (aq)
Produce an almost constant 1.35V

32. Button Cell Metal Cap (–) Zinc Powder
Cathode outer container of nickel or steel (+)
Electrolyte
Mercury Oxide

33. Secondary Cells Lead – Acid (Car Battery) Nickel cadmium Cells
Fuel Cells

34. Lead Acid Battery Car Batteries p 211-2
Also called storage batteries or accumulators
Each cell produces 2 volts so typical 12 volt car battery contains 6 cells
Both electrodes are lead plates separated by some porous material like cardboard

35. Lead Acid Battery
Positive electrode is coated with PbO2 Lead (IV) Oxide
The electrolyte is a solution of 4M sulfuric acid

36. Lead Acid Battery Anode Oxidation (–) Cathode Reduction (+)
Pb(s) + SO4 2-  PbSO4 (s) + 2e –
Cathode Reduction (+)
PbO2(s) + SO H+ + 2e –  PbSO4 (s) + 2H2O (l)
Overall Reaction
Pb(s) + PbO2(s) + 2H2SO4  2PbSO4 (s) +2H2O (l)

37. Nickel Cadmium Cells Often called Nicads
Electrodes are Nickel and Cadmium
Electrolyte is Potassium Hydroxide
Reactions involve the hydroxides of the two metals

38. Nickel Cadmium Cells Anode (Oxidation) (– ) Cathode (Reduction) (+)
Cd (s) + 2OH– (aq)  Cd(OH)2 (s) + 2 e–
Cathode (Reduction) (+)
NiO-OH (s) + H2O (l) + e–  Ni(OH)2 (s) + OH– (aq)
Overall Reaction
Cd (s) +NiO-OH(s) + H2O(l)  Cd(OH)2 (s)+ Ni(OH)2 (s)

39. Fuel Cells
Limitation of dry cells looked at so far is that they contain reactants in small amounts and when they reach equilibrium.
Primary Cells are then discarded, secondary cells are then recharged
A cell that can be continually fed reactants would overcome this and allow for a continual supply of electricity

40. Fuel Cells
Fuel cells transform chemical energy directly into electrical energy
60% efficiency
Space Program uses hydrogen and oxygen with an electrolyte of Potassium Hydroxide

41. Fuel Cells Anode Oxidation (–) Cathode Reduction (+) Overall Equation
H2(g) + 2OH –(aq)  2H2O (l) + 2e–
Cathode Reduction (+)
O2(g) + 2H2O(l) + 4e–  4OH–(aq)
Overall Equation
H2(g) + O2(g)  2H2O (l)

42. Hydrogen Oxygen Fuel Cell– +
Electrolyte
HydrogenGas Inlet
Oxygen Gas
Inlet
Porous
Anode
Porous
Cathode
Water outlet