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ANALYTICAL CHEMISTRY
CHEM 3811
CHAPTER 14
DR. AUGUSTINE OFORI AGYEMAN
Assistant professor of chemistry
Department of natural sciences
Clayton state university
CHAPTER 14
ELECTRODE POTENTIALS
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
REDOX CHEMISTRY
Oxidizing Agent
- The species that gains electrons
- The species that is reduced
- Causes oxidation
Aox + ne- ↔ Ared
Cu2+(aq) + 2e- ↔ Cu(s)
REDOX CHEMISTRY
Reducing Agent
- The species that loses electrons
- The species that is oxidized
- Causes reduction
Bred ↔ Box + ne-
Fe(s) ↔ Fe2+(aq) + 2e-
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)
- Oxidizing agent
- Reduced species
- Electron gain
- Reducing agent
- Oxidized species
- Electron loss
REDOX CHEMISTRY
Half Reactions
- Oxidation half reaction
Bred ↔ Box + neFe(s) ↔ Fe2+(aq) + 2e-
- Reduction half reaction
Aox + ne- ↔ Ared
Cu2+(aq) + 2e- ↔ Cu(s)
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
REDOX CHEMISTRY
Charge (q) of an electron = - 1.602 x 10-19 C
Charge (q) of a proton = + 1.602 x 10-19 C
C = coulombs
Charge of one mole of electrons
= (1.602 x 10-19 C)(6.022 x 1023/mol) = 9.649 x 104 C/mol
= Faraday constant (F)
q=nxF
CURRENT
- The quantity of charge flowing past a point in an
electric circuit per second
Units
Ampere (A) = coulomb per second (C/s)
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)
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)
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-
ELECTROLYSIS CELL
- Nonspontaneous reaction
- Requires electrical energy to occur
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
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)
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
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
VOLTAIC (GALVANIC) CELL
For the overall reaction
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
e-
Voltmeter
-
Zn electrode
e-
+
Cu electrode
ClK+
Zn2+
Salt bridge (KCl)
Anode
Oxidation
Zn(s) → Zn2+(aq) + 2e-
Cu2+
Cathode
Reduction
Cu2+(aq) + 2e- → Cu(s)
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
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)
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
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
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
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 0.799 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)
STANDARD POTENTIALS
Silver does not react spontaneously with hydrogen
2H+(aq) + 2e- → H2(g)
Ag+(aq) + e- → Ag(s)
Eo = 0.000 V
Eo = +0.799 V
Reverse the second equation (sign changes)
Ag(s) → Ag+(aq) + eEo = -0.799 V
Multiply the second equation by 2 (Eo is intensive so remains)
2Ag(s) → 2Ag+(aq) + 2e- Eo = -0.799 V
Combine (electrons cancel)
2Ag(s) + 2H+(aq) → 2Ag+(aq) + H2(g) Eo = -0.799 V
STANDARD POTENTIALS
Consider Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Cu2+(aq) + 2e- → Cu(s)
Zn2+(aq) + 2e- → Zn(s)
Eo = +0.339 V
Eo = -0.762 V
Reverse the second equation (sign changes)
Zn(s) → Zn2+(aq) + 2eEo = +0.762 V
Combine (electrons cancel)
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Eo = +1.101 V
Eo is positive so reaction is spontaneous
Reverse reaction is nonspontaneous
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
STANDARD POTENTIALS
Reducing
agents
Increasing reducing power
Increasing oxidizing power
Oxidizing
agents
Half Reaction
F2 + 2e- ↔ 2FMnO4- + 5e- ↔ Mn2+
Ce4+ + e- ↔ Ce3+ (in HCl)
O2 + 4H+ + 4e- ↔ 2H2O
Ag+ + e- ↔ Ag(s)
Cu2+ + 2e- ↔ Cu(s)
2H+ + 2e- ↔ H2(g)
Cd2+ + 2e- ↔ Cd(s)
Fe2+ + 2e- ↔ Fe(s)
Zn2+ + 2e- ↔ Zn(s)
Al3+ + 3e- ↔ Al(s)
K+ + e- ↔ K(s)
Li+ + e- ↔ Li(s)
Eo (V)
2.890
1.507
1.280
1.229
0.799
0.339
0.000
-0.402
-0.440
-0.763
-1.659
-2.936
-3.040
NERNST EQUATION
For the half reaction
aA + ne- ↔ bB
The half-cell potential (at 25 oC), E, is given by
RT  B b 

EE 
ln 
a 
nF  A 
O
b



B
0.05916
E  EO 
log 
a
n
 A



NERNST EQUATION
Eo = standard electrode potential
R = gas constant = 8.314 J/K-mol
T = absolute temperature
F = Faraday’s constant = 9.649 x 104 C/mol
n = number of electrons
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
NERNST EQUATION
- When a half reaction is multiplied by a factor
Eo remains the same
- For a complete reaction
Ecell = E+ - Eand
Eo = E+o - E-o
E+ = potential at positive terminal
E- = potential at negative terminal
NERNST EQUATION
For the Cu – Fe cell at standard conditions
Cu2+ + 2e- ↔ Cu(s)
0.339 V
Fe2+ + 2e- ↔ Fe(s)
-0.440 V
Ecell = 0.779 V
Galvanic Reaction
Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq)
Fe Fe2+ (1M)
Cu2+ (1 M) Cu
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
NERNST EQUATION
At cell equilibrium at 25 oC
E = 0 and Q = K (the equilibrium constant)
E
o
0.05916

logK
n
Or
K  10
nEo /0.05916
Positive Eo implies K > 1
Negative Eo implies K < 1
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)
REFERENCE ELECTRODES
Saturated Calomel electrode (SCE)
- Saturated with KCl
1/2Hg2Cl2(s) + e- ↔ Hg(l) + ClE = + 0.241 V
In this case, the reference is not 0.000 V (SHE)
but 0.241 V (SCE)
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
REFERENCE ELECTRODES
Silver-Silver Chloride Electrode (Ag/AgCl)
- Saturated with KCl
AgCl(s) + e- ↔ Ag(s) + ClE = + 0.197 V
REFERENCE ELECTRODES
Emeasured = Eo - 0.241 (SCE)
Emeasured = Eo - 0.197 (Ag/AgCl)
Eo(SHE)
E(SCE)
E(Ag/AgCl)
Cu2+ + 2e- ↔ Cu(s)
0.339 V
0.098 V
0.142 V
Fe2+ + 2e- ↔ Fe(s)
-0.440 V
-0.681 V
-0.637 V