1. Fundamentals of Electrochemistry
Introduction
1.) Electrical Measurements of Chemical Processes
Redox Reaction involves transfer of electrons from one species to another.
Chemicals are separated
Can monitor redox reaction when electrons flow through an electric current
Electric current is proportional to rate of reaction
Cell voltage is proportional to free-energy change
Batteries produce a direct current by converting chemical energy to electrical energy.
Common applications run the gamut from cars to ipods to laptops

2. Fundamentals of Electrochemistry
Basic Concepts
1.) A Redox titration is an analytical technique based on the transfer of electrons between analyte and titrant
Reduction-oxidation reaction
A substance is reduced when it gains electrons from another substance
gain of e- net decrease in charge of species
Oxidizing agent (oxidant)
A substance is oxidized when it loses electrons to another substance
loss of e- net increase in charge of species
Reducing agent (reductant)
(Reduction)
(Oxidation)
Oxidizing
Agent
Reducing
Agent

3. Fundamentals of Electrochemistry
Basic Concepts
2.) The first two reactions are known as “1/2 cell reactions”
Include electrons in their equation
3.) The net reaction is known as the total cell reaction
No free electrons in its equation
4.) In order for a redox reaction to occur, both reduction of one compound and oxidation of another must take place simultaneously
Total number of electrons is constant
½ cell reactions:
Net Reaction:

4. Fundamentals of Electrochemistry
Basic Concepts
5.) Electric Charge (q)
Measured in coulombs (C)
Charge of a single electron is 1.602x10-19C
Faraday constant (F) – 9.649x104C is the charge of a mole of electrons
6.) Electric current
Quantity of charge flowing each second
through a circuit
Ampere: unit of current (C/sec)
Relation between charge and moles:
Coulombs
moles

5. Higher potential difference
Fundamentals of Electrochemistry
Basic Concepts
7.) Electric Potential (E)
Measured in volts (V)
Work (energy) needed when moving an electric charge from one point to another
Measure of force pushing on electrons
Higher potential difference
Relation between free energy, work and voltage:
Volts
Coulombs
Joules
Higher potential difference requires more work to lift water (electrons) to higher trough

6. Fundamentals of Electrochemistry
Basic Concepts
7.) Electric Potential (E)
Combining definition of electrical charge and potential
Relation between free energy difference and electric potential difference:
Describes the voltage that can be generated by a chemical reaction

7. Fundamentals of Electrochemistry
Basic Concepts
8.) Ohm’s Law
Current (I) is directly proportional to the potential difference (voltage) across a circuit and inversely proportional to the resistance (R)
Ohms (W) - units of resistance
9.) Power (P)
Work done per unit time
Units: joules per second J/sec or watts (W)

8. Fundamentals of Electrochemistry
Galvanic Cells
1.) Galvanic or Voltaic cell
Spontaneous chemical reaction to generate electricity
One reagent oxidized the other reduced
two reagents cannot be in contact
Electrons flow from reducing agent to oxidizing agent
Flow through external circuit to go from one reagent to the other
Reduction:
Oxidation:
Net Reaction:
AgCl(s) is reduced to Ag(s)
Ag deposited on electrode and Cl-
goes into solution
Electrons travel from Cd
electrode to Ag electrode
Cd(s) is oxidized to Cd2+
Cd2+ goes into solution

9. Fundamentals of Electrochemistry
Galvanic Cells
1.) Galvanic or Voltaic cell
Example: Calculate the voltage for the following chemical reaction
DG = -150kJ/mol of Cd
Solution:
n – number of moles of electrons

10. Fundamentals of Electrochemistry
Galvanic Cells
2.) Cell Potentials vs. DG
Reaction is spontaneous if it does not require external energy

11. Fundamentals of Electrochemistry
Galvanic Cells
2.) Cell Potentials vs. DG
Reaction is spontaneous if it does not require external energy
Potential of overall cell = measure of the tendency of this reaction to proceed to equilibrium
ˆ Larger the potential, the further the reaction is from equilibrium and the greater the driving force that exists
Similar in concept to balls sitting at different heights along a hill

12. Fundamentals of Electrochemistry
Galvanic Cells
3.) Electrodes
Cathode: electrode where reduction takes place
Anode: electrode where oxidation takes place

13. Fundamentals of Electrochemistry
Galvanic Cells
4.) Salt Bridge
Connects & separates two half-cell reactions
Prevents charge build-up and allows counter-ion migration
Salt Bridge
Contains electrolytes not
involved in redox reaction.
K+ (and Cd2+) moves to cathode with
e- through salt bridge (counter
balances –charge build-up
NO3- moves to anode (counter
balances +charge build-up)
Completes circuit
Two half-cell reactions

14. Fundamentals of Electrochemistry
Galvanic Cells
5.) Short-Hand Notation
Representation of Cells: by convention start with anode on left
Zn|ZnSO4(aZN2+ = )||CuSO4(aCu2+ = )|Cu
Phase boundary
Electrode/solution interface
anode
cathode
2 liquid junctions
due to salt bridge
Solution in contact with
anode & its concentration
Solution in contact with
cathode & its concentration

15. Fundamentals of Electrochemistry
Standard Potentials
1.) Predict voltage observed when two half-cells are connected
Standard reduction potential (Eo) the measured potential of a half-cell reduction reaction relative to a standard oxidation reaction
Potential arbitrary set to 0 for standard electrode
Potential of cell = Potential of ½ reaction
Potentials measured at standard conditions
All concentrations (or activities) = 1M
25oC, 1 atm pressure
Ag+ + e- » Ag(s) Eo = V
Standard Hydrogen Electrode (S.H.E)
Pt(s)|H2(g)(aH2 = 1)|H+(aq)(aH+ = 1)||
Hydrogen gas is bubbled over a Pt electrode

16. Fundamentals of Electrochemistry
Standard Potentials
1.) Predict voltage observed when two half-cells are connected
As Eo increases, the more favorable the reaction and the more easily the compound is reduced (better oxidizing agent).
Reactions always written as reduction
Appendix H contains a more extensive list

17. Fundamentals of Electrochemistry
Standard Potentials
2.) When combining two ½ cell reaction together to get a complete net reaction, the total cell potential (Ecell) is given by:
Where: E+ = the reduction potential for the ½ cell reaction at the positive electrode
E+ = electrode where reduction occurs (cathode)
E- = the reduction potential for the ½ cell reaction at the negative electrode
E- = electrode where oxidation occurs (anode)
Place values on number line to determine the potential difference
Electrons always flow towards more positive potential

18. Fundamentals of Electrochemistry
Standard Potentials
3.) Example: Calculate Eo, and DGo for the following reaction:

19. Fundamentals of Electrochemistry
Nernst Equation
1.) Reduction Potential under Non-standard Conditions
E determined using Nernst Equation
Concentrations not-equal to 1M
For the given reaction:
aA + ne- » bB Eo
The ½ cell reduction potential is given by:
Where: E = actual ½ cell reduction potential
Eo = standard ½ cell reduction potential
n = number of electrons in reaction
T = temperature (K)
R = ideal gas law constant (8.314J/(K-mol)
F = Faraday’s constant (9.649x104 C/mol)
A = activity of A or B
at 25oC

20. Fundamentals of Electrochemistry
Nernst Equation
2.) Example:
Calculate the cell voltage if the concentration of NaF and KCl were each 0.10 M in the following cell:
Pb(s) | PbF2(s) | F- (aq) || Cl- (aq) | AgCl(s) | Ag(s)

21. Fundamentals of Electrochemistry
Eo and the Equilibrium Constant
1.) A Galvanic Cell Produces Electricity because the Cell Reaction is NOT at Equilibrium
Concentration in two cells change with current
Concentration will continue to change until Equilibrium is reached
E = 0V at equilibrium
Battery is “dead”
Consider the following ½ cell reactions:
aA + ne- » cC E+o
dD + ne- » bB E-o
Cell potential in terms of Nernst Equation is:
Simplify:

22. Fundamentals of Electrochemistry
Eo and the Equilibrium Constant
1.) A Galvanic Cell Produces Electricity because the Cell Reaction is NOT at Equilibrium
Since Eo=E+o- E-o:
At equilibrium Ecell =0:
Definition of
equilibrium constant
at 25oC
at 25oC

23. Fundamentals of Electrochemistry
Eo and the Equilibrium Constant
2.) Example:
Calculate the equilibrium constant (K) for the following reaction:

24. Fundamentals of Electrochemistry
Cells as Chemical Probes
1.) Two Types of Equilibrium in Galvanic Cells
Equilibrium between the two half-cells
Equilibrium within each half-cell
If a Galvanic Cell has a nonzero voltage then the net cell reaction is not at equilibrium
Conversely, a chemical reaction within a ½ cell will reach and remain at equilibrium.
For a potential to exist, electrons must flow from one cell to the other which requires the reaction to proceed  not at equilibrium.

25. Fundamentals of Electrochemistry
Cells as Chemical Probes
2.) Example:
If the voltage for the following cell is 0.512V, find Ksp for Cu(IO3)2:
Ni(s)|NiSO4(0.0025M)||KIO3(0.10 M)|Cu(IO3)2(s)|Cu(s)

26. Fundamentals of Electrochemistry
Biochemists Use Eo´
1.) Redox Potentials Containing Acids or Bases are pH Dependent
Standard potential  all concentrations = 1 M
pH=0 for [H+] = 1M
2.) pH Inside of a Plant or Animal Cell is ~ 7
Standard potentials at pH =0 not appropriate for biological systems
Reduction or oxidation strength may be reversed at pH 0 compared to pH 7
Metabolic Pathways

27. Fundamentals of Electrochemistry
Biochemists Use Eo´
3.) Formal Potential
Reduction potential that applies under a specified set of conditions
Formal potential at pH 7 is Eo´
Eo´ (V)
Need to express concentrations as
function of Ka and [H+].
Cannot use formal concentrations!