1. ELECTROCHEMISTRY Chapter 18

2. Introduction
* Electrochemistry: the study of the relationship between electron flow and redox reactions.
Significance and Applications?
What is a redox reaction? An oxidation – reduction reaction in which electrons are transferred (same # e- lost and gained).
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

3. I. Oxidation and Reduction Reactions A. Introduction
1. Reactions involving the transfer of electrons.
2. One species has a desire to lose electrons an another species has a desire to gain electrons. With the correct conditions, energy in the form of current flow is produced.
3. Driving force is tendency for electron flow.
4. Examples: batteries, combustion, rusting of iron, metabolism of food in the body.

4. B. Definitions Fe  Fe2+ + 2 e-
1. Oxidation - The loss of electron(s).
Oxidation Reaction
M  M n n e -
While M is oxidized, it acts as an reducing agent because it causes something else to be reduced (i.e.,it gives electrons to something else).
Fe  Fe e-

5. 2. Reduction - The gain of electron(s). Reduction Reaction
X n e -  X n-
While X is reduced, it acts as an oxidizing agent because it causes something else to be oxidized (i.e., it takes electrons from something else).
Cl e-  2 Cl-
Ag+ + e-  Ag

6. E electrons (products) O oxidation the lion says G gain of
L loss of
E electrons (products)
O oxidation
the lion says
G gain of
E electrons (reactants)
R!!! reduction

7. C. Properties of Oxidation-Reduction Rxns
1. “Redox reaction” always consists of the two processes (oxidation and reduction) which occurs simultaneously. One species must lose electrons so that another species can gain those same # of electrons.
2. A redox reaction may be recognized if a metal is being oxidized or reduced.
Metal gains or loses electrons as it goes from reactants to products.
How can we recognize this in a reaction?

8. Cu2+(aq) + Mg(s)  Mg2+(aq) + Cu(s)
Example Problems
For the redox reaction given, identify which reactant is being oxidized, which reactant is being reduced, which is the oxidizing agent, and which is the reducing agent?
Cu2+(aq) + Mg(s)  Mg2+(aq) + Cu(s)
MnO4-(aq) + Fe2+(aq)  Mn2+(aq) + Fe3+(aq)

9. D. Related Terminology For Redox Rxns
* Recall complete redox reaction always includes both oxidation and reduction.
* Half-reaction: One of two parts of an oxidation - reduction reaction, one part of which involves a loss of electrons and the other a gain of electrons.
For the reaction: Fe + Cu2+  Fe2+ + Cu
1. Write the reduction half reaction.
2. Write the oxidation half reaction.

10. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) d. What is the reducing agent?
For the reaction:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
d. What is the reducing agent?
Reducing agent: The reactant that is oxidized.
It loses electrons so that some other reactant can be reduced.
e. What is the oxidizing agent?
Oxidizing Agent: The reactant that is reduced.
It gains electrons so that some other reactant can be oxidized.

11. E. Balancing Redox Reactions
1. Introduction
Oxidation Number Method
Arbitrary book-keeping system for electrons which can be used to identify what is being oxidized or reduced.
Half-Reaction Method
*** You will be responsible for a simplified version of this.

12. 2. Simplified Rules for Balancing Redox
A. Split the reaction into half-reactions. (by inspection) B. Balance number of atoms of each element on both sides of each equation by changing coefficients.
C. Balance charges on both sides of each equation
by adding the appropriate electrons. (Electrons added to more positive side.)
D. Same number of electrons must be transferred in
redox reaction. Multiply one or both half reactions by a whole number to get same number of electrons transferred.
E. Add half-reactions together for balanced redox reaction. Recheck to see that electrons cancel out and atoms of each element balance.

13. Example Problems Balance the following redox reactions:
Sn2+(aq) + Fe3+(aq) → Sn4+(aq) + Fe2+(aq)
Sb3+(aq) + Sr(s) → Sb(s) + Sr2+(aq)
How many mol es of electrons are being transferred in the following balanced
redox reaction?
2 Al(s) + 3 Cl2(g) → 2 AlCl3(aq)

14. II. Voltaic (Galvanic) Cells
A. Introduction
1. An electrochemical cell in which a product-favored (spontaneous) redox reaction generates an electric current.
2. The reaction produces an electron flow through an outside conductor (wire).
3. Examples: batteries

15. B. Example of a Voltaic Cell 1. Reaction
Zn(s)  Zn2+(aq) e-
__Cu2+(aq) + 2 e-  Cu (s)________
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
If you just ran the redox reaction in a test tube, the reaction would occur, but no current (energy) would be captured. See slide #15
However in a galvanic (voltaic) cell, energy is captured in the form of current flow. Slide # 16.
The 2 half-reactions occur in separate half-cells.

16. Redox Reaction With No Current Flow

17. Redox Reaction With Current Flow
(Voltaic Cell)

18. C. Requirements of Galvanic Cell
1. Anode: an electrode (conductor such as metal strip or graphite) where oxidation occurs.
2. Cathode: an electrode (conductor such as metal strip or graphite) where reduction occurs.
3. Salt Bridge: A tube of an electrolyte (sometimes in a gel) that is connected to the two half-cells of a voltaic cell: the salt bridge allows the flow of ions but prevents the mixing of the different solutions that would allow direct reaction of the cell reactants. Charge does not build up in half cells. Electrical neutrality must be maintained.

19. Cathode A
Reduction
Anode U T
Oxidation

20. Cell Diagrams
A cell diagram is “shorthand” for an electrochemical cell.
The anode is placed on the left side of the diagram.
The cathode is placed on the right side.
A single vertical line ( | ) represents a boundary between phases, such as between an electrode and a solution.
A double vertical line ( || ) represents a salt bridge or porous barrier separating two half-cells.

21. D. Galvanic Cells and Electrical Potential
1. Electron flow in galvanic cell can do work /
produce energy.
2. Electrical potential energy measured in volts.
1 volt = (1 joule) / (1 coulomb)
work unit of charge due to
6.24 x 1018 electrons
3. Coulombs = amperes x seconds
C = A x s or
A = C / s

22. E. Standard Cell Voltages
1. Cell voltages can be measured under standard conditions (1 atm pressure, 250 C, and 1.0 M concentrations). E0cell.
2. The standard cell potential is determined by the equation
E0cell = E0red + E0ox
3. If E0cell is positive, the net cell reaction is product- favored (spontaneous).
4. If E0cell is negative, the net cell reaction is reactant- favored (nonspontaneous).

23. F. Standard Electrode Potentials
1. Standard Electrode Potentials are measured for half-reactions, relative to a standard hydrogen electrode potential (which is assigned 0 volts).
2. See Table 18.1 for Standard Electrode Potentials
a. Each half reaction is written as a reduction.
b. Each half reaction could occur in either direction.
c. The more “+” the standard electrode potential, the greater the tendency to undergo reduction, meaning a good oxidizing agent.

24. d. The more “-” the standard electrode potential,
d. The more “-” the standard electrode potential, the greater the tendency to undergo oxidation, meaning a good reducing agent.
e. If a half-reaction is written in the reverse direction, the sign of the corresponding standard electrode potential must change.
(Same magnitude, opposite in sign)
f. If a half-reaction is multiplied by a factor (coefficient), the standard electrode potential is not multiplied by that factor.

25. Selected Electrode (Reduction) Potentials
F2(g) + 2e-  2F V F2/F-
Au3+ + 3e-  Au(s) V Au3+/Au
Hg2+ + 2e-  Hg(l) V Hg2+/Hg
2H+ + 2e-  H2(g) V H+/H2 Fe2+ + 2e-  Fe(s) V Fe2+/Fe
Li+ + e-  Li(s) V Li+/Li

26. G. Problem Solving and Voltaic Cells
1. A voltaic (galvanic) cell is constructed using a piece
of Sn in 1.00 M Sn2+ for one half-cell and a piece of Au in 1.00 M Au3+ for the other half-cell. The reaction is run under standard conditions of pressure and temperature.
a. Determine the net cell potential (voltage).
b. Determine the net cell reaction.
c. What is the oxidizing agent for the net reaction? d. What is the reaction occurring at the anode?
e. Sketch the cell information in the following diagram.

28. H. Cell Potential and Gibbs Free Energy
A positive E0cell indicates a spontaneous or product-favored reaction. There must be a relationship between E0cell and free energy (ΔG0).
ΔG0 = -n F E0cell
n = # moles electrons transferred
F = Faraday constant = 9.65 x 104 C / mol e-
(also recall that volt = Joules / Coulomb = J / C)
** A positive E0cell would produce a negative ΔG0

29. 1. Calculate the standard free energy change (ΔG0)
for the reaction below under standard conditions.
NiO2 + 2Cl- + 4 H+  Cl2 + Ni H2O
E0cell = v

30. I. Equilibrium Constant and Cell Potential We know ΔG0 = -RTlnKeq
and: ΔG0 = -n F E0cell
Therefore: n F E0cell = -RTlnKeq
E0cell = RTlnKeq / nF
Rearranging: ln Keq = n F E0cell / RT

31. 1. Calculate Kc for the following reaction under standard conditions:
Cu2+(aq) Ag(s)  Cu(s) Ag+(aq)
Is the reaction product-favored?
Do you expect a large amount of product formed at equilibrium?

32. Thermodynamics, Equilibrium, and Electrochemistry: A Summary
From any one of the three quantities Keq, ΔG°, E°cell, we can determine the others.

34. J. Concentration Effects on Cell Potential
1. When all the concentrations in a voltaic
cell are 1.0 M, under standard conditions,
the cell potential equals the standard cell
potential.
2. As concentrations of reactants or products
change, the cell potential (voltage changes).
3. When equilibrium is reached, the cell
potential will drop to zero.

36. Mathematical Relationship Nernst Equation
R = gas constant = J / mol K
T = Kelvin temperature
n = # moles of electrons transferred
Q = reaction quotient. What is it??
or (easiest to use equation) RT
Ecell = E°cell – –––– ln Q nF

37. Cr3+(aq) + 3 e-  Cr(s) E0 = -.76v
Problem:
A galvanic cell contains Ni2+ (aq) in contact with Ni (s) and Cr3+ (aq) in contact with Cr (s). When [Ni2+] = 1.0 x 10-4 and [Cr3+] = 2.0 x 10-3, determine the value of Ecell . Given:
Ni2+(aq) e-  Ni(s) E0 = -.25v
Cr3+(aq) e-  Cr(s) E0 = -.76v

38. K. Applications of Voltaic Cells
1. The Lead Storage Battery (Secondary Battery)
Anode Reaction
Pb(s) + HSO41-(aq)  PbSO4(s) + H+ + 2e-
Cathode Reaction
PbO2(s) + 3H+ + HSO41-(aq) + 2e-  PbSO4(s) + 2H2O (l)

39. Lead Storage Battery

40. 2. Dry Cell Batteries 3. Fuel Cells What are they?
Where are they used?
How do they differ from batteries?

41. Corrosion of an Iron Piling
One way to minimize rusting is to provide a different anode reaction.

42. IV. Electrolytic Cells A. Electrolytic vs. Voltaic (Galvanic) Cells
Redox reaction which proceeds spontaneously in a product favored direction, generating electricity.
Electrolytic Cell
Redox reaction in which an electrical current is supplied to drive a nonspontaneous, reaction- favored reaction.

43. B. Properties of Electrolytic Cell
1. Energy requiring (in form of electric current).
2. No physical separation needed for the two electrode reactions.
3. Usually no salt bridge required.
4. Conducting medium is molten salt or aqueous solution.
5. For electrolytic redox reaction:
E0cell is negative.
ΔG0 is positive.
Kc is small (<1).

44. C. Electrolysis of Molten Sodium Chloride 1. Redox Reaction:
2 Na+ + 2 e-  2 Na(l) cathode (reduction)
2 Cl  Cl2(g) + 2 e anode (oxidation)
2 Na+ + 2 Cl-  2 Na(l) + Cl2(g) net cell rxn
2. Electrolytic Cell:
See next slide

47. E. Stoichiometry of Electrolysis
1. Electrons Treated as Stoichiometric Factor
a. Given the half-reactions:
Na+ + e- Na(s) and Cu e-  Cu(s)
How many mol e- required to form 1 mol Na(s)?
How many mol e- required to form 4 mol Cu(s)?
***
b. Chemical change is directly proportional to
*** electron flow (charge).
2. Coulomb Relationship
*** Charge = current x time
2 mol e- = 1 mol Cu(s)
1 mol e- = 9.65 x 104 C
C = Amps x seconds

48. Steps for Electrolysis Calculations

49. a. How many moles of electrons are required in an
3. Problem Solving
a. How many moles of electrons are required in an
electrolytic cell to deposit 2.00 grams of chromium,
Cr(s), from a solution of CrCl3?

50. b. What mass of aluminum metal can be produced per
hour in the electrolysis of a molten aluminum salt
by a current of 26 A?

51. F. Applications of Electrolysis Electroplating
Refining of Metals
Copper
Aluminum

52. Electroplating
Electrolysis can be used to coat one metal onto another, a process called electroplating.
Usually, the object to be electroplated, such as a spoon, is cast of an inexpensive metal. It is then coated with a thin layer of a more attractive, corrosion-resistant, and expensive metal, such as silver or gold.

53. Refining of Aluminum