1. ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 14
DR. AUGUSTINE OFORI AGYEMAN
Assistant professor of chemistry
Department of natural sciences
Clayton state university

2. CHAPTER 14
ELECTRODE POTENTIALS

3. REDOX CHEMISTRY - Electron transfer occurs in redox reactions
Oxidation
- Loss of electrons
Reduction
- Gain of electrons
Oxidizing agent (oxidant) is the species reduced
Reducing agent (reductant) is the species oxidized

4. REDOX CHEMISTRY Oxidizing Agent - The species that gains electrons
- The species that is reduced
- Causes oxidation
Aox + ne- ↔ Ared
Cu2+(aq) + 2e- ↔ Cu(s)

5. REDOX CHEMISTRY Reducing Agent - The species that loses electrons
- The species that is oxidized
- Causes reduction
Bred ↔ Box + ne-
Fe(s) ↔ Fe2+(aq) + 2e-

6. REDOX CHEMISTRY The Overall Reaction
- Both an oxidation and a reduction must occur in a redox reaction
- The oxidizing agent accepts electrons from the reducing agent
Aox + Bred ↔ Ared + Box
Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq)
- Reducing agent
- Oxidized species
- Electron loss
- Oxidizing agent
- Reduced species
- Electron gain

7. REDOX CHEMISTRY Half Reactions - Oxidation half reaction
Bred ↔ Box + ne-
Fe(s) ↔ Fe2+(aq) + 2e-
- Reduction half reaction
Aox + ne- ↔ Ared
Cu2+(aq) + 2e- ↔ Cu(s)

8. ELECTRODE - Conducts electrons into or out of a redox reaction system
Examples
platinum wire
carbon (glassy or graphite)
indium tin oxide (ITO)
Electroactive Species
- Donate or accept electrons at an electrode

9. REDOX CHEMISTRY Charge (q) of an electron = - 1.602 x 10-19 C
Charge (q) of a proton = x C
C = coulombs
Charge of one mole of electrons
= (1.602 x C)(6.022 x 1023/mol) = x 104 C/mol
= Faraday constant (F)
q = n x F

10. CURRENT - The quantity of charge flowing past a point in an
electric circuit per second
Units
Ampere (A) = coulomb per second (C/s)

11. VOLTAGE Potential Difference (E)
- Work done by or on electrons when they
move from one point to another
Units: volts (V or J/C)
Work (J) = E (V) x q (C)

12. CHEMICAL CHANGE Spontaneous Process
- Takes place with no apparent cause
Nonspontaneous Process
- Requires something to be applied in order for it to occur
(usually in the form of energy)

13. ELECTROLYSIS - Voltage is applied to drive a redox reaction that
would not otherwise occur
Examples
- Production of aluminum metal from Al3+
- Production of Cl2 from Cl-

14. ELECTROLYSIS CELL - Nonspontaneous reaction
- Requires electrical energy to occur

15. GALVANIC CELL - Spontaneous reaction - Produces electrical energy
- Can be reversed electrolytically for reversible cells
Example
Rechargeable batteries
Conditions for Non-reversibility
- If one or more of the species decomposes
- If a gas is produced and escapes

16. GALVANIC CELL - A spontaneous redox reaction generates electricity
- One reagent is oxidized and the other is reduced
- The two reagents must be separated (cannot be in contact)
- Electrons flow through a wire (external circuit)

17. VOLTAIC (GALVANIC) CELL
Oxidation Half reaction
- Loss of electrons
- Occurs at anode (negative electrode)
- The left half-cell by convention
Reduction Half Reaction
- Gain of electrons
- Occurs at cathode (positive electrode)
- The right half-cell by convention

18. GALVANIC CELL Salt Bridge
- Connects the two half-cells (anode and cathode)
- Filled with gel containing saturated aqueous salt solution (KCl)
- Ions migrate through to maintain electroneutrality
- Prevents charge buildup that may cease the reaction process
Preparation
- Heat 3 g of agar and 30 g of KCl in 100 mL H2O
- Heat until a clear solution is obtained
- Pour into a U-tube and allow to gel
- Store in a saturated aqueous KCl

19. VOLTAIC (GALVANIC) CELL
For the overall reaction
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) e-
Voltmeter e- - +
Cu electrode
Zn electrode
Cl- K+
Zn2+
Salt bridge (KCl)
Cu2+
Anode
Oxidation
Zn(s) → Zn2+(aq) + 2e-
Cathode
Reduction
Cu2+(aq) + 2e- → Cu(s)

20. GALVANIC CELL Voltage or Potential Difference (E)
- Is the voltage measured
- Measured by a voltmeter (potentiometer) connected to electrodes
Greater Voltage
- More favorable net cell reaction
- More work done by flowing electrons

21. GALVANIC CELL Line Notation
Phase boundary: represented by one vertical line
Salt bridge: represented by two vertical lines
Fe(s) FeCl2(aq) CuSO4(aq) Cu(s)

22. STANDARD POTENTIALS Electrode Potentials
- A measure of how willing a species is to gain or lose electrons
Positive Voltage (spontaneous process)
- Electrons flow into the negative terminal of voltmeter
(flow from negative electrode to positive electrode)
Negative Voltage (nonspontaneous process)
- Electrons flow into the positive terminal of voltmeter
(flow from positive electrode to negative electrode)
Conventionally
- Negative terminal is on the left of galvanic cells

23. STANDARD POTENTIALS Standard Reduction Potential (Eo)
- Used to predict the voltage when different cells are connected
- Potential of a cell as cathode compared to
standard hydrogen electrode
- Species are solids or liquids
- Activities = 1
- We will use concentrations for simplicity
Concentrations = 1 M
Pressures = 1 bar

24. STANDARD POTENTIALS Standard Hydrogen Electrode (SHE)
- Used to measure Eo for half-reactions
- Connected to negative terminal
- Pt electrode
- Acidic solution in which [H+] = 1 M
- H2 gas (1 bar) is bubbled past the electrode
H+(aq, 1 M) + e- ↔ 1/2H2 (g, 1 bar)
Conventionally, Eo = 0 for SHE

25. STANDARD POTENTIALS The Eo for Ag+ + e- ↔ Ag(s) is 0.799 V
Implies that if a sample of silver metal is placed
in a 1 M Ag+ solution, a value of V will be
measured with S. H. E. as reference
Pt(s) H2(g, 1 bar) H+(aq, 1 M ) Ag+ (aq, 1 M) Ag(s)
SHE Ag+ (aq, 1 M) Ag(s)

26. STANDARD POTENTIALS Silver does not react spontaneously with hydrogen
2H+(aq) + 2e- → H2(g) Eo = V
Ag+(aq) + e- → Ag(s) Eo = V
Reverse the second equation (sign changes)
Ag(s) → Ag+(aq) + e- Eo = V
Multiply the second equation by 2 (Eo is intensive so remains)
2Ag(s) → 2Ag+(aq) + 2e- Eo = V
Combine (electrons cancel)
2Ag(s) + 2H+(aq) → 2Ag+(aq) + H2(g) Eo = V

27. STANDARD POTENTIALS Consider Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Cu2+(aq) + 2e- → Cu(s) Eo = V
Zn2+(aq) + 2e- → Zn(s) Eo = V
Reverse the second equation (sign changes)
Zn(s) → Zn2+(aq) + 2e- Eo = V
Combine (electrons cancel)
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Eo = V
Eo is positive so reaction is spontaneous
Reverse reaction is nonspontaneous

28. STANDARD POTENTIALS
- Half-reaction is more favorable for more positive Eo
Formal Potential
- The potential for a cell containing a specified concentration
of reagent other than 1 M

29. STANDARD POTENTIALS Eo (V) Half Reaction 2.890 F2 + 2e- ↔ 2F- 1.507
1.280
1.229
0.799
0.339
0.000
-0.402
-0.440
-0.763
-1.659
-2.936
-3.040
Half Reaction
F e- ↔ 2F-
MnO e- ↔ Mn2+
Ce e- ↔ Ce3+ (in HCl)
O H e- ↔ 2H2O
Ag+ + e- ↔ Ag(s)
Cu e- ↔ Cu(s)
2H+ + 2e- ↔ H2(g)
Cd e- ↔ Cd(s)
Fe e- ↔ Fe(s)
Zn e- ↔ Zn(s)
Al e- ↔ Al(s)
K+ + e- ↔ K(s)
Li+ + e- ↔ Li(s)
Oxidizing
agents
Reducing
agents
Increasing oxidizing power
Increasing reducing power

30. The half-cell potential (at 25 oC), E, is given by
NERNST EQUATION
For the half reaction
aA + ne- ↔ bB
The half-cell potential (at 25 oC), E, is given by

31. NERNST EQUATION Eo = standard electrode potential
R = gas constant = J/K-mol
T = absolute temperature
F = Faraday’s constant = x 104 C/mol
n = number of electrons

32. NERNST EQUATION - The standard reduction potential (Eo)
when [A] = [B] = 1M
- [B]b/[A]a = Q = reaction quotient
- Concentration for gases are expressed as pressures in bars
- Q = 1 for [ ] = 1 M and P = 1 bar
logQ = 0 and E = Eo
- Q is omitted for pure solids, liquids, and solvents

33. NERNST EQUATION - When a half reaction is multiplied by a factor
Eo remains the same
- For a complete reaction
Ecell = E+ - E-
and
Eo = E+o - E-o
E+ = potential at positive terminal
E- = potential at negative terminal

34. NERNST EQUATION For the Cu – Fe cell at standard conditions
Cu e- ↔ Cu(s) V
Fe e- ↔ Fe(s) V
Ecell = V
Galvanic Reaction
Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq)
Fe Fe2+ (1M) Cu2+ (1 M) Cu

35. NERNST EQUATION - Positive E implies spontaneous forward cell reaction
- Negative E implies spontaneous reverse cell reaction
If cell runs for long
- Reactants are consumed
- Products are formed
- Equilibrium is reached
- E becomes 0
- Reason why batteries run down

36. NERNST EQUATION At cell equilibrium at 25 oC
E = 0 and Q = K (the equilibrium constant) Or
Positive Eo implies K > 1
Negative Eo implies K < 1

37. REFERENCE ELECTRODES Indicator Electrode
- Responds directly to the analyte
Reference Electrode
- Provides known and constant potential
Examples
Silver-silver chloride electrode (Ag/AgCl)
Saturated Calomel electrode (SCE)

38. REFERENCE ELECTRODES Saturated Calomel electrode (SCE)
- Saturated with KCl
1/2Hg2Cl2(s) + e- ↔ Hg(l) + Cl-
E = V
In this case, the reference is not V (SHE)
but V (SCE)

39. REFERENCE ELECTRODES Saturated Calomel electrode (SCE)
- Different KCl concentrations can be used
- 0.1 M KCl is least temperature sensitive
- Saturated KCl solution is easier to make and maintain

40. REFERENCE ELECTRODES Silver-Silver Chloride Electrode (Ag/AgCl)
- Saturated with KCl
AgCl(s) + e- ↔ Ag(s) + Cl-
E = V

41. REFERENCE ELECTRODES Emeasured = Eo - 0.241 (SCE)
Emeasured = Eo (Ag/AgCl)
Eo(SHE) E(SCE) E(Ag/AgCl)
Cu e- ↔ Cu(s) V V V
Fe e- ↔ Fe(s) V V V