1. Electrochemistry is the branch of chemistry that deals with the use of chemical reaction to generate a potential or voltage.

2. Electrochemical Reactions
In electrochemical reactions, electrons are transferred from one species to another.
Mg (s) + HCl (aq) →

3. Oxidation Numbers
In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

4. Oxidation and Reduction
A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.

5. Oxidation and Reduction
A species is reduced when it gains electrons.
Here, each of the H+ gains an electron, and they combine to form H2.

6. Oxidation and Reduction
What is reduced is the oxidizing agent.
H+ oxidizes Zn by taking electrons from it.
What is oxidized is the reducing agent.
Zn reduces H+ by giving it electrons.

7. Assigning Oxidation Numbers, ON
Elements in their elemental form have an oxidation number of 0.
Mg Cl2 Pb
The oxidation number of a monatomic ion is the same as its charge.
Zn2+ F- Na+

8. HF NaH Nonmetals tend to have negative oxidation numbers
In combination with other elements:
Nonmetals tend to have negative oxidation numbers
Oxygen has an oxidation number of −2, except in the peroxide ion, which has an oxidation number of −1.
MgO Na2O
Hydrogen is −1 when bonded to a metal and +1 when bonded to a nonmetal.
HF NaH

9. Assigning Oxidation Numbers
Fluorine always has an oxidation number of −1.
HF NaF
The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.
ZnBr CuCl2

10. Assigning Oxidation Numbers
The sum of the oxidation numbers in a neutral compound is 0.
HCl H2O K2O
5. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.
CO ClO4-

11. Examples
Determine the oxidation number of sulfur in the following compounds:
H2S
SO42- S8
S2O32-

12. Some Redox reactions:
Redox reactions involve transfer of electrons between reactants
When a substance loses electrons, it undergoes oxidation;
The substance oxidized is called reducing agent.
When a substance gains electrons, it undergoes reduction; the substance reduced is called an oxidizing agent.
Example:
The reaction of a metal with either an acid or a metal salt, displacement reaction; the ion is displaced or replaced through oxidation of an element.
A + BX → AX + B
Mg (s) + HCl →
Zn (s) + CuSO4 → ZnSO4 (s) + Cu (s)

13. Electrochemical Reactions

14. In a spontaneous Oxidation-Reduction (redox) Reaction electrons are transferred and energy is released
as the reaction proceeds:
Zn strip dissolves…
The blue color due to Cu2+ fades.
Copper metal is deposited.

15. 20-1 Electrode Potentials and Their Measurement
Cu(s) + 2Ag+(aq)
Cu2+(aq) + 2 Ag(s)
Cu(s) + Zn2+(aq)
No reaction
FIGURE 20-1
Behaviour of Ag+(aq) and Zn+(aq) in the presence of copper
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General Chemistry: Chapter 20

16. Voltaic Cells
We can use that energy to do electrical work.
If we make the electrons flow through an external pathway rather than directly between reactants.
We call such a setup a voltaic cell.

17. …the reaction is the same: Zn (s) + Cu2+ (aq)
The significant difference is that the reactants, Zn (s) and Cu2+ (aq), are not in direct contact in the voltaic cell.
We physically separate the reduction half of the reaction from the oxidation half of the reaction.
Zn metal is in contact with a solution of Zn2+ (aq).
Cu metal is in contact with a solution of Cu2+ (aq).

18. An electrochemical half-cell consists of a metal electrode, M, partially immersed in an aqueous solution of its ions, Mn+.

19. Voltaic Cells
An electrochemical cell consists of two half-cells with electrodes joined by a wire and solutions joined by a salt bridge.
The oxidation occurs at the anode.
The reduction occurs at the cathode.

20. Voltaic Cells
Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop.

21. Voltaic Cells
Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced.
Cations move toward the cathode.
Anions move toward the anode.

22. Voltaic Cells
In the cell, then, electrons leave the anode and flow through the wire to the cathode.
As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

23. Voltaic Cells
As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode.
The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

24. Electromotive Force (emf)
Water only spontaneously flows one way in a waterfall.
Likewise, electrons only spontaneously flow one way in a redox reaction—from higher to lower potential energy.

25. Cell Diagrams and Terminology
Chemistry 140 Fall 2002
Cell Diagrams and Terminology
Electromotive force, Ecell
The cell voltage or cell potential.
Cell diagram
Shows the components of the cell in a symbolic way.
Anode (where oxidation occurs) on the left.
Cathode (where reduction occurs) on the right.
Boundary between phases shown by |.
Boundary between half cells (commonly a salt bridge) shown by ||.
Several memory devices have been proposed for the oxidation/anode and reduction/cathode relationships. Perhaps the simplest is that in the oxidation/anode relationship, both terms begin with a vowel: o/a; in the reduction/cathode relationship, both begin with a consonant: r/c. Another method is to remember that Cations go to the Cathode (both begin with C), that cations are formed at the anode is then obvious.
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General Chemistry: Chapter 20

26. Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Chemistry 140 Fall 2002
This slide shows the correct symbolic description of the voltaic, or galvanic, cell in figure 20-4.
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Ecell = V
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
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27. Galvanic (or voltaic) cells
Chemistry 140 Fall 2002
Galvanic (or voltaic) cells
Produce electricity as a result of spontaneous reactions.
Electrolytic cells
Non-spontaneous chemical change driven by electricity.
The electrochemical cells of Figures 20-3 and 20-4 produce electricity as a result of spontaneous chemical reactions; as such, they are called voltaic, or galvanic, cells. In Section 20-7 we will consider electrolytic cells—electrochemical cells in which electricity is used to accomplish a nonspontaneous chemical change.
General Chemistry: Chapter 20

28. 20-2 Standard Electrode Potentials
Cell voltages, the potential differences between electrodes, are among the most precise scientific measurements.
The potential of an individual electrode is difficult to establish.
Arbitrary zero is chosen.
The Standard Hydrogen Electrode (SHE)
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29. The standard hydrogen electrode (SHE)
Chemistry 140 Fall 2002
2 H+(a = 1) + 2 e H2(g, 1 bar) E° = 0 V
Pt|H2(g, 1 bar)|H+(a = 1)
Because hydrogen is a gas at room temperature, electrodes cannot be constructed from it. The standard hydrogen electrode consists of a piece of platinum dipped into a solution containing 1 M H+(aq) with a stream of hydrogen passing over its surface. The platinum does not react but provides a surface for the reduction of H3O+(aq) to H2(g) as well as the reverse oxidation half-reaction.
The standard hydrogen electrode (SHE)
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General Chemistry: Chapter 20

30. Standard Electrode Potential, E° is the voltage of a cell formed from two standard electrodes
E°cell = E°cathode (right) – E°anode,(left)
The tendency for a reduction process to occur at an electrode.
All ionic species present at a=1 (approximately 1 M).
All gases are at 1 bar (approximately 1 atm).

31. E°cell = E°cathode - E°anode
Cu2+(1M) + 2 e- → Cu(s) E°Cu2+/Cu = ?
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = V
anode
cathode
Standard cell potential: the potential difference of a cell formed from two standard electrodes.
E°cell = E°cathode - E°anode
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General Chemistry: Chapter 20

32. Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = 0.340 V
E°cell = E°cathode - E°anode
E°cell = E°Cu2+/Cu - E°H+/H2
0.340 V = E°Cu2+/Cu - 0 V
E°Cu2+/Cu = V
H2(g, 1 atm) + Cu2+(1 M) → H+(1 M) + Cu(s) E°cell = V
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General Chemistry: Chapter 20

33. Measuring standard electrode potentials
Figure: 20-06
Title:
Measuring standard electrode potentials
Caption:
(a) A SHE is the anode, and copper is the cathode. Contact between the half-cells occurs through a porous plate that acts like a salt bridge. (b) This cell has the same connections as that in part (a), but with zinc substituting for copper. However, the electron flow here is opposite that in (a), as noted by the negative voltage. This means zinc is the anode in this case.
A SHE is the anode, and copper is the cathode.
Contact between the half-cells occurs through a porous plate that acts like a salt bridge.
(b) This cell has the same connections as that in part (a), but with zinc substituting for copper. However, the electron flow here is opposite that in (a), as noted by the negative voltage. This means zinc is the anode in this case.

34. Voltaic Cells
Two different metals (plus an electrolyte) produce a voltage when circuit is completed.
Different metal pairs produce different voltages.
Electrons flow from the metal higher in the activity series to the metal lower in the series through an external circuit.

35. Electromotive Force (emf) (the voltage produced by a voltaic cell)
The potential difference between the anode and cathode in a cell is called the electromotive force (emf).
It is also called the cell potential and is designated Ecell.
Electromotive force (emf) is the force required to push electrons through the external circuit

36. Standard Half Cell potentials, Reduction Potentials is a measure of the driving force for reduction to occur under standard conditions:1 M concentrations, P= 1atm, at 25 °C. . .
Standard reduction potentials, E°red, are measured relative to a standard.

37. Oxidizing and Reducing Agents
The strongest oxidizers have the most positive reduction potentials.
The strongest reducers have the most negative reduction potentials.

38. General Chemistry: Chapter 20
TABLE Some Selected Standard Electrode (Reduction) Potentials at 25°C
Reduction Half-Reaction E°, V
Acidic Solution
Slide 38 of 53
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General Chemistry: Chapter 20

39. General Chemistry: Chapter 20
TABLE Some Selected Standard Electrode (Reduction) Potentials at 25°C (continued)
Reduction Half-Reaction E°, V
Acidic Solution
Slide 39 of 53
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General Chemistry: Chapter 20

40. General Chemistry: Chapter 20
TABLE Some Selected Standard Electrode (Reduction) Potentials at 25°C (continued)
Reduction Half-Reaction E°, V
Acidic Solution
Basic Solution
Slide 40 of 53
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General Chemistry: Chapter 20

41. Corrosion: Unwanted Voltaic Cells

42. Figure: 20-20
Title:
Protection of iron against electrolytic corrosion
Caption:
In the anodic reaction, the metal that is more easily oxidized loses electrons to produce metal ions: in (a), this is iron; in (b), it is zinc. In the cathodic reaction, oxygen gas, which is dissolved in a thin film of water on the metal, is reduced to OH-. Rusting of iron occurs in (a), but not in (b). When iron corrodes, Fe2+ and OH- ions from the half-reactions initiate these further reactions:
Fe2+ + 2OH--> Fe(OH)2(s)
4Fe(OH)2(s) + O2 + 2H2O --> 4Fe(OH)3(s)
2Fe(OH)3(s) --> Fe2O3 • H2O(s) + 2H2O(l)

43. Figure: UN
Title:
Magnesium sacrificial anodes
Caption:
The small cylindrical bars of magnesium attached to the steel ship provide cathodic protection against corrosion. Galvanization is another type of cathodic protection, the use of magnesium here is more cost effective.

44. Cell Potentials  Ered = −0.76 V  Ered = +0.34 V
For the oxidation in this cell,
For the reduction,
Ered = −0.76 V 
Ered = V 

45. Cell Potentials Ecell  = Ered (cathode) − (anode)
= V − (−0.76 V)
= V

46. Oxidizing and Reducing Agents
Cathode is always the electrode with more positive reduction potential.
The greater the difference between the two reduction potential, the greater the voltage of the cell.

47. A voltaic cell is based on the following two standard half-reactions:
Cd2+ (aq) + 2e Cd (s)
Sn2+ (aq) + 2e Sn(s)
By using the data from Standard Reduction Potentials table, determine:
the half reactions that occur at the cathode and the anode
the standard cell potential.
Cd2+(aq) + 2 e > Cd(s) V
Answer: V.

48. Ecell as a function of Concentration
Cell potential depends on concentration of reactants, products and temperature. The Nernst Equation allows us to calculate the voltage produced by any electrochemical cell given Eo values for its electrodes and the concentrations of reactants and products.

49. At 25oC the Nernst equation is simplified to
E = Eo - (0.0592/n) × logQ
Eo = standard electrochemical cell potential n = number of electrons passed from anode to cathode. Q = the equilibrium expression of the electrochemical reaction.

50. Example: If we have the reaction:
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
If [Cu2+] = 5.0 M and [Zn2+] = 0.050M:

51. Calculate the voltage produced by the cell Sn(s)|Sn2+||Ag+|Ag(s) at 25oC given: [Sn2+] = 0.15 M [Ag+] = 1.7 M Answer: 0.70 Volt.

52. Concentration Cells
Notice that the Nernst equation implies that a cell could be created that has the same substance at both electrodes.
For such a cell,
be 0, but Q would not.
Ecell 
Therefore, as long as the concentrations are different, E will not be 0.

53. Applications of Oxidation-Reduction Reactions

54. +4 Pb + PbO2 + 2H2SO4 → 2PbSO4 + H2O 0 +2 +2 Figure: 20-15 Title:
A lead-acid (storage) cell
Caption:
The composition of the electrodes is described in the text. The cell reaction that occurs as the cell is discharged in given in equation (20.24). The voltage of the cell is 2.02 V. In this depiction, two anode plates are connected in parallel fashion, as are two cathode plates. Car batteries may have up to 6 cells connected in parallel, as opposed to the 2 shown here. +4 +2
Pb + PbO2 + 2H2SO4 → 2PbSO4 + H2O

55. Hydrogen Fuel Cells

56. General Chemistry: Chapter 20
Chemistry 140 Fall 2002
Fuel Cells
O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq)
2{H2(g) + 2 OH-(aq) → 2 H2O(l) + 2 e-}
2H2(g) + O2(g) → 2 H2O(l)
E°cell = E°O2/OH- - E°H2O/H2
= V – ( V) = V
 is the efficiency value used to evaluate fuel cell performance.
FIGURE 20-18
 = ΔG°/ ΔH° = 0.83
A hydrogen-oxygen fuel cell
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General Chemistry: Chapter 20

57. The flow of electrons from anode to cathode is spontaneous
The flow of electrons from anode to cathode is spontaneous. What is the “driving force”? • Electrons flow from anode to cathode because the cathode has a lower electrical potential energy than the anode. • Potential difference is the difference in electrical potential. • The potential difference is measured in volts.

58. Standard Cell Potentials
The cell potential at standard conditions can be found through this equation:
Ecell 
= Ered (cathode) − Ered (anode)
The cell potential for all spontaneous reactions
at Standard conditions is a positive value.
The standard conditions include 1 M concentrations,
and P= 1 atm (g), for reactants and products at 25 °C.