IB Chemistry on Voltaic Cell, Standard Electrode Potential and Standard Hydrogen Electrode

1. Types voltaic cell
Conversion electrical energy
to chemical energy
Electrochemistry
Electrolytic cellVoltaic cell
NH4CI and ZnCI2
Chemical and electrical energy
Redox rxn
(Oxidation/reduction)
Movement electron
Produce electricity
Conversion chemical energy
to electrical energy
Electrodes– different metal (Half cell) Electrodes – same metal (Half cell)
Chemical
rxn
Electric current
Daniell cell Alkaline cellDry cell Nickel cadmium cell
Primary cell (Non rechargeable)
MnO2 and KOH
Secondary cell (Rechargeable)

2. Conversion electrical to chemical energy
Electrochemistry
ElectrolyticcellVoltaic cell
Conversion chemicalto electricalenergy
Cathode (+ve) - Reduction Cathode (-ve) - Reduction
Vs
Electron flow from anode (-ve) to cathode (+ve) electrode Electron flow from anode (+ve) to cathode (-ve) electrode
Anode
(-ve)
Spontaneous rxn Non Spontaneous rxn
Anode (-ve) – Oxidation Anode (+ve) – Oxidation
++
О
О
О
О
- -
Zn → Zn 2+ + 2e
(oxidized)
Cu2+ + 2e → Cu
(reduced)
Zn2+
Zn2+
Zn2+
Zn2+-
-
-
-
→ +
+
+
Cu2+
Cu2+
Cu2+
-e
-e
+
+
+ -
-
-
X-→ X + -e
(oxidized)
X
-
X
-
X
-
Anode
(+ve)
Cathode
(-ve)
Cathode
(+ve)
-e
-e
Y+ + e- → Y
(reduced)
Y+
Y+
Y+
-e
-e
-e
-e
Anode Cathode
Voltaic Cell Electrolytic Cell
Anode Oxidation Negative (-ve) Oxidation Positive (+ve)
Cathode Reduction Positive (+ve) Reduction Negative (-ve)
Cation (+ve ion) to cathode (-ve)Anion (-ve ion) to anode (+ve)

3. Current– measured in Amperes or Coulombs per second
1A = 1 Coulomb charge pass througha point in 1 second = 1C/s
1 Coulomb charge (electron)= 6.28 x 10 18 electronspassing in 1 second
1 electron/protoncarry charge of – 1.6 x 10 -19 C ( very small)
6.28 x 10 18 electron carry charge of - 1 C
Electric current
Flow electric charges (electron, -ve)
From High electric potential – low potential
Due to Potential Difference– measure with ammeter
ond
electron
ond
Coulomb
A
sec.1
.1028.6
sec1
1
1
18


Click here current/voltage
Current Electric Current – movingcharges in solid wire or solution
Flow of
charges
-
-
-
Solid/WireSolution/Electrolyte
Electron move in random
No current flow cause
No potential difference
Electrons & Protons
-
-
+
+
1A = 6.28 x 1018 e
1 second
Video on current/voltage
Potential Difference across wire
Electron move in one direction
Current flow
+ve ions -ve ions
(cations) (anions)
Potential Difference applied/Battery used
ItQ  t = Time/ s
Find amt charges pass through a sol if
Current is 2.ooA, time is 15 mins
ItQ 
Current flow
Q = Amt Charges/ C I = Current/ A
CQ 1800601500.2 

4. Electric Potential
C
J
Volt
1
1 
Potential Diff/Voltage
-Measured in Volt with Voltmeter
- 1 V = 1 Joule energy released when 1 Coulomb
charge pass through 1 point
- 1 V = 1 J/C
6V battery - 6J energy for every Coulomb
moved bet its terminals.
V = Potential Diff
I = Current
R = Resistance
Potential diff bet 2 points is 1 V
↓
1 J energy released when 1 C charge passes through
Voltmeter across
1Volt
1 V
Potential Diff/Voltage/PotentialEnergy
+ -
1 Ω 2 Ω
Charges (-ve)
flow down
A
R
V
I
RIV
2
3
6


VV
RIV
212 

-
+
-
+
VV
RIV
422 

Total current
Potential Diff(PD)vs Current
PD = Water Pressure
PD = 1.5V – 1.5J energy released 1C charge flow down
PD – cause charge flow- charge flow = Called CURRENT
Potential Diff(PD)vs Current
1.5V = 1.5J/C
A
DElectric potential/PD/Voltage = Electric Pressure = Volt
Electric Current/Current = Charge flow = Amp
Electric Potential Energy = Work done to bring a charge to a point = Joule
Voltage NOT same as energy, Voltage = energy/charge
Battery lift charges, Q to higher potential
Potential Energy bet 2 terminals in battery stored as chemical energy
2A 2A

5. EMF vs PD
V = Potential Diff
I = Current
R = Resistance
Max potential diff bet two
electrodes of battery source.
+ -
1 Ω 2 Ω
A
R
V
I
RIV
2
3
6


VV
RIV
212 

VV
RIV
422 

Total current
Current flow Circuit complete
Circuit complete
↓
Current flow
↓
Internal resistance
(battery - 1Ω)
↓
Terminal PD = 8V
(Voltage drop)
Potential Diff/Voltage in Volt
Symbol for EMF = E or ℰ
Click here voltage drop
internal resistance
No Current flow in circuit
EMF (ElectromotiveForce) in Volt
Battery = EMF = 9V
9 Volt
).(9 currentnoVEMFV
IRV


EMF Internal resistance Ir
Place voltmeter across – EMF= 9V
No current flow.
A
rR
E
I
rRIE
IrIREMFE
1
9
9
)18(
9
)(
)(
)(







VV
RIV
881 

VV
RIV
111 

EMF = 8V+1V
8 Volt
1 Volt
Voltage measured across terminal = 8V Click here EMF notes Click here PD, PE and I
EMF (6V) = 2V + 4V
4 Volt2 Volt
Charges passing through wire
Current flow Circuit complete

6. Series vs Parallel Circuit
3 Ω
A
R
V
I
RIV
5.0
18
9


VV
V
RIV
5.2
55.0



VV
V
RIV
5
105.0



Total Resistance
3 + 10 + 5 = 18Ω
Parallel CircuitSeries Circuit
EMF (9V) = 2.5V + 5V + 1.5V
R = 0.625Ω
Voltage same across all component. VT = V1 = V2 = V3
Total current = sum of current in each branch. IT = I1 + I2 + I3
Total equi resistance < value of any individual resistor
Total Current
Current same in circuit
Total voltage = sum of component in series
5 Ω
10 Ω0.5 A
VV
V
RIV
5.1
35.0



Voltage across all component
Same = 9V
Total Current = sum of all current in each branch
Total Resistance
10Ω 2Ω 1Ω
AI
R
V
I
9.0
10
9


AI
R
V
I
5.4
2
9


AI
R
V
I
9
1
9


Total current (14.4A) = 0.9A + 4.5A + 9A
Ohm’s Law
Sum voltage drop equal to voltage source (EMF). VT = V1 + V2 + V3
Current same in all components in series.
Total resistance = sum of individual resistances. RT = R1 + R2 + R3
Click here voltage drop
Series Circuit Parallel Circuit
• Voltage do not flow
• Charges/Current flow
• Voltage cause current to flow
• Voltage ≠ Energy
• Battery do not supply electron
• Wire contain electron, they flow
Electron in wire repel by -ve terminal move in circuit
Electron move slowly, drift velocity, electric field move at speed of light
Electric signal travel speed of light, bulb light up instantaneously,
Electric signal/field travel faster than movement of electron
Movement e – cause electric field – travel speed of light – bulb light up
Voltage Diff – Pressure diff across
Voltage Diff – Cause water/current to flow
14.4A 4.5A0.9A 9A

7. Potential Diff bet Zn/Zn2+
Electrode potential Zn/Zn2+ = -ve
-
Electrode Potential
Redox Equilibrium
Zn2+
Zn → Zn 2+ + 2e
(Oxidation)
Zn 2+ + 2e → Zn
(Reduction)
Zn 2+ + 2e ↔ Zn
(At equilibrium)
Metal Zn placed in its sol Zn2+ ion
Equilibrium bet Zn/Zn2+
Zn metal reactive lose e form Zn2+
Equilibrium shift to right ←
Potential Diff form bet Zn/Zn2+
Potential Diff
Electrode potential = -ve
Zn2+
Zn2+
Zn
Zn2+
Zn
Zn2+
Zn2+ Zn2+
Zn 2+ + 2e ↔ Zn
Equi shift to ←
-
--
Zn
-
-
-
-
+
+
+
+
+ +
+ +
+
Voltage of Zn/Zn2+ can’t be measured.
Abs electrodepotentialcan’t measured.
Only Diff in electrode potentialcan be measured.
Cannot measure
Abs Potential
Metal Cu placed in its sol Cu2+ ion
Equilibrium bet Cu/Cu2+
Cu2+ ion gain -2e form Cu
Equilibrium shift to left ←
Potential Diff form bet Cu/Cu2+
Potential Diff
Electrode potential = +ve
Cu
Cu2+
Cu2+
Cu2+
Cu2+
Cu → Cu2+ + 2e
(Oxidation)
Cu2+ + 2e → Cu
(Reduction)
Cu2+ + 2e ↔ Cu
(At equilibrium)
Cu
-e
-e
-e
Cu2+
Cu2+
Cu2+
Cu2+ + 2e ↔ Cu
Equi shift to →
Zn Half Cell
+
+
+
Cu
+
+
+
---
-
--- ----
--
Potential Diff bet Cu/Cu2+
Electrode potential Cu/Cu2+ = +ve
Cannot measure
Abs Potential
Voltage of Cu/Cu2+can’t be measured.
Abs electrodepotentialcan’t measured.
Only Diff in electrode potentialcan be measured.
Click here chem database
(std electrode potential)
Click here chem database
(std electrode potential)
Click here interactive ECS Click here pdf version ECS
Cu Half Cell

8. PotentialDiff Cu/Cu2+
Electrode potential
Cu/Cu2+ = +ve
PotentialDiff Zn/Zn2+
Electrode potential
Zn/Zn2+ = -ve
Zn2+
Zn → Zn 2+ + 2e
(Oxidation)
Zn 2+ + 2e → Zn
(Reduction)
Zn 2+ + 2e ↔ Zn
(At equilibrium)
Zn2+
Zn2+
Zn
Zn2+
Zn
Zn2+
Zn2+ Zn2+
Zn 2+ + 2e ↔ Zn
Equi shift to ←
-
-
-
Zn
-
--
-
+
++
+
+ +
+
+
+
Can’t measure
Abs Potential
Cu
Cu2+
Cu2+
Cu2+
Cu2+
Cu → Cu2+ + 2e
(Oxidation)
Cu2+ + 2e → Cu
(Reduction)
Cu2+ + 2e ↔ Cu
(At equilibrium)
Cu
-e
-e
-e
Cu2+
Cu2+
Cu2+
Cu2+ + 2e ↔ Cu
Equi shift to →
Zn Half Cell
+
+
+
Cu
+
+
+
-
Cu Half Cell
Zn/Cu Voltaic Cell
External circuit – flow of electrons
Complete circuit
-
--
--
-
-
----
-- -
Connect 2 Half Cell with wire/ salt bridge
Zn half cell (-ve)
Oxidation
Cu half cell (+ve)
Reduction
Salt Bridge – flow of ions
Complete the circuit
Cu2+ + 2e → CuZn → Zn 2+ + 2e
Zn + Cu2+ → Zn2+ + Cu
Anode Cathode
Maintain electrical
neutrality
Salt bridge – saturated KNO3
Zn2+ increase ↑
NO3
- flow in to balance excess Zn2+
Cu2+ decrease ↓, excess –ve ion ↑
K+ flow in to balance loss of Cu2+
Zn Cu
--
-
-
Zn2+
Zn2+
Zn2+
Excess of Zn2+ ion
+
+
++
-
-
-
-
---
-
-
-
-
-
Excess of –ve ion
+
+
+
+
++
+
Without Salt Bridge
-+
+
+
+
With Salt Bridge
(electron unable to flow due to ESF)
NO3
-
NO3
-
NO3
-
NO3
-
+
+
+ K
+
K
+
K
+
-
-
-
K+ flow in to balance
excess of – ion
NO3
-
flow in to balance
excess of + ion
2 Half Cell to make a Voltaic Cell
-e -e
-
-
-
-
+
+
+
+

9. PotentialDiff Cu/Cu2+
Electrode potential
Cu/Cu2+ = +ve
PotentialDiff Zn/Zn2+
Electrode potential
Zn/Zn2+ = -ve
Zn2+
Zn → Zn 2+ + 2e
(Oxidation)
Zn 2+ + 2e → Zn
(Reduction)
Zn 2+ + 2e ↔ Zn
(At equilibrium)
Zn2+
Zn2+
Zn
Zn2+
Zn
Zn2+
Zn2+ Zn2+
Zn 2+ + 2e ↔ Zn
Equi shift to ←
-
-
-
Zn
-
--
-
+
++
+
+ +
+
+
+
Can’t measure
Abs Potential
Cu
Cu2+
Cu2+
Cu2+
Cu2+
Cu → Cu2+ + 2e
(Oxidation)
Cu2+ + 2e → Cu
(Reduction)
Cu2+ + 2e ↔ Cu
(At equilibrium)
Cu
-e
-e
-e
Cu2+
Cu2+
Cu2+
Cu2+ + 2e ↔ Cu
Equi shift to →
+
+
+
Cu
+
+
+
-
External circuit – flow of electrons
Complete circuit
-
--
--
-
-
----
-- -
Connect 2 Half Cell with wire/ salt bridge
Zn half cell (-ve)
Oxidation
Cu half cell (+ve)
Reduction
Voltmeter – High resistance
(No current flow) Salt Bridge – flow of ions
Complete the circuit
Cu2+ + 2e → CuZn → Zn 2+ + 2e
1.10Volt
Potential diff can be measured.
Voltmeter across – EMF
1.10 Volt
Zn + Cu2+ → Zn2+ + Cu
Anode Cathode
Zn(s) | Zn2+
(aq) || Cu2+
(aq)| Cu (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Phase boundarySalt Bridge Flow
electrons
Maintain electrical
neutrality
Salt bridge – saturated KNO3
Zn2+ increase ↑
NO3
- flow in to balance excess Zn2+
Cu2+ decrease ↓
K+ flow in to balance loss of Cu2+
Zn/Cu Voltaic Cell 2 Half Cell to make a Voltaic Cell
Zn Half Cell Cu Half Cell
-e -e
-
-
-
-
+
+
+
+

10. PotentialDiff Ag/Ag2+
Electrode potential
Ag/Ag2+ = +ve
PotentialDiff Zn/Zn2+
Electrode potential
Zn/Zn2+ = -ve
Zn2+
Zn → Zn 2+ + 2e
(Oxidation)
Zn 2+ + 2e → Zn
(Reduction)
Zn 2+ + 2e ↔ Zn
(At equilibrium)
Zn2+
Zn2+
Zn
Zn2+
Zn
Zn2+
Zn2+ Zn2+
Zn 2+ + 2e ↔ Zn
Equi shift to ←
-
-
-
Zn
-
--
-
+
++
+
+ +
+
+
+
Can’t measure
Abs Potential
Ag
Ag+
Ag+
Ag+
Ag+
Ag → Ag+ + e
(Oxidation)
Ag+ + e → Ag
(Reduction)
Ag+ + e ↔ Ag
(At equilibrium)
Ag
-e
-e
-e
Ag+
Ag+
Ag+
Ag+ + e ↔ Ag
Equi shift to →
+
+
+
Ag
+
+
+
-
External circuit – flow of electrons
Complete circuit
-
--
--
-
-
----
-- -
Connect 2 Half Cell with wire/ salt bridge
Zn half cell (-ve)
Oxidation
Ag half cell (+ve)
Reduction
Voltmeter – High resistance
(No current flow) Salt Bridge – flow of ions
Complete the circuit
Ag+ + e → AgZn → Zn 2+ + 2e
1.56Volt
Potential diff can be measured.
Voltmeter across – EMF
1.56 Volt
Zn + 2Ag+ → Zn2+ + 2Ag
Anode Cathode
Zn(s) | Zn2+
(aq) || Ag+
(aq)| Ag (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Phase boundarySalt Bridge Flow
electrons
Maintain electrical
neutrality
Salt bridge – saturated KNO3
Zn2+ increase ↑
NO3
- flow in to balance excess Zn2+
Ag+ decrease ↓
K+ flow in to balance loss of Ag+
Zn/Ag Voltaic Cell 2 Half Cell to make a Voltaic Cell
Zn Half Cell Ag Half Cell
Ag
Ag+
-e -e
-
-
-
-
+
+
+
+

11. PotentialDiff Ag/Ag2+
Electrode potential
Ag/Ag2+ = +ve
PotentialDiff Cu/Cu2+
Electrode potential
Cu/Cu2+ = -ve
Cu2+
Cu → Cu 2+ + 2e
(Oxidation)
Cu 2+ + 2e → Cu
(Reduction)
Cu 2+ + 2e ↔ Cu
(At equilibrium)
Cu2+
Cu2+
Cu
Cu2+
Cu
Cu2+
Cu2+ Cu2+
Cu 2+ + 2e ↔ Cu
Equi shift to ←
-
-
-
Cu
-
--
-
+
++
+
+ +
+
+
+
Can’t measure
Abs Potential
Ag
Ag+
Ag+
Ag+
Ag+
Ag → Ag+ + e
(Oxidation)
Ag+ + e → Ag
(Reduction)
Ag+ + e ↔ Ag
(At equilibrium)
Ag
-e
-e
-e
Ag+
Ag+
Ag+
Ag+ + e ↔ Ag
Equi shift to →
+
+
+
Ag
+
+
+
-
External circuit – flow of electrons
Complete circuit
-
--
--
-
-
----
-- -
Connect 2 Half Cell with wire/ salt bridge
Cu half cell (-ve)
Oxidation
Ag half cell (+ve)
Reduction
Voltmeter – High resistance
(No current flow) Salt Bridge – flow of ions
Complete the circuit
Ag+ + e → AgCu → Cu 2+ + 2e
0.46Volt
Potential diff can be measured.
Voltmeter across – EMF
0.46 Volt
Cu + 2Ag+ → Cu2+ + 2Ag
Anode Cathode
Cu(s) | Cu2+
(aq) || Ag+
(aq)| Ag (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Phase boundarySalt Bridge Flow
electrons
Maintain electrical
neutrality
Salt bridge – saturated KNO3
Cu2+ increase ↑
NO3
- flow in to balance excess Cu2+
Ag+ decrease ↓
K+ flow in to balance loss of Ag+
Cu/Ag Voltaic Cell 2 Half Cell to make a Voltaic Cell
Cu Half Cell Ag Half Cell
Ag
Ag+
Cu
Cu2+
-e -e
-
-
-
-
+
+
+
+

12. Standard Electrode Potential
Standard HydrogenElectrode (SHE)
Platinum coat with Platinum oxide/black
– increase surface area for adsorption H2
- catalyze equilibrium bet H2 /H+
- H2 ↔ 2H+ + 2e-
Eθ
Standard Reference electrode
All Cell Potential are measured against
• Conc ( 1M)
• Pressure (1 atm)
• Temp (298K)
• Platinum-inert electrode
(sys without metal)
Standard
condition
H2 at 1 atm
Platinum
H2 gas
Pt wire
Platinum
2H+ + 2e ↔ H2
Eθ
= 0V
Types of Half Cells
Metal/ Ion (M/M+)
Gas/ Ion (M/M-)
Ion/ Ion (Fe3+/Fe2+)
• Pure Zn metal
• Conc (1M Zn2+)
• Pressure (1 atm)
• Temp(298K)
Condition Std Zn/Zn2+
Condition Std CI2/CI-
• CI2 gas
• Platinum electrode
• Conc (1M CI-)
• Pressure (1 atm)
• Temp(298K)
• Platinum electrode
• Conc (1M Fe3+/Fe2+)
• Pressure (1 atm)
• Temp(298K)
Condition Std Fe3+/ Fe2+
Zn2+
Zn
Fe3+/Fe2+
CI-
Condition for Standard
C
A
N
T
M
E
A
S
U
R
E
A
B
S
P
O
T
E
N
T
I
A
L
1
2
3
How to measure
electrode
potential ?
Pt
1M H+
Measure
Difference?

13. Standard Electrode Potential
Std HydrogenElectrode (SHE)
Eθ
= 0V
Types of Half Cells
Metal/ Ion (M/M+)
Gas/ Ion (M/M+)
Ion/ Ion (Fe3+/Fe2+)
• Pure Zn metal
• Conc (1M Zn2+)
• Pressure (1 atm)
• Temp(298K)
Condition Std Zn/Zn2+
Condition Std CI2/CI-
• CI2 gas
• Platinum electrode
• Conc (1M CI-)
• Pressure (1 atm)
• Temp(298K)
• Platinum electrode
• Conc (1M Fe3+/Fe2+)
• Pressure (1 atm)
• Temp(298K)
Condition Std Fe3+/ Fe2+
Zn2+
Zn
Fe3+/Fe2+
1
2
3
Connect to SHE
Connect to SHE
Connect to SHE
Eθ
= 0V
Eθ
= 0V
Eθ
= -0.76V
Standard electrode potential Zn/Zn2+ = -0.76V
Eθ
cell = -0.76V
Eθ
= +0.77V
Eθ
= +1.35V
Standard electrode potential Fe3+/Fe2+ = +0.77V
Eθ
cell = +0.77V
Standard electrode potential CI2 /CI- = +1.35V
Eθ
cell = +1.35V
Eθ
= -0.76V
Eθ
= +0.77V
Eθ
= +1.35V
2 Half Cellwith SHE as referenceelectrode
CI-
Pt
+
+
+
Pt

14. Standard Electrode Potential
Std Electrode Potential diff systems
Eθ
= 0V
Eθ
= 0V
Eθ
= 0V
Eθ
= -0.76V
Standard electrode potential Zn/Zn2+ = -0.76V
Eθ
cell = -0.76V
Eθ
= +0.77V
Eθ
= +1.35V
Standard electrode potential Fe3+/Fe2+ = +0.77V
Eθ
cell = +0.77V
Standard electrode potential CI2 /CI- = +1.35V
Eθ
cell = +1.35V
Eθ
= -0.76V
Eθ
= +0.77V
Eθ
= +1.35V
STANDARD Reduction potential – Hydrogen as std
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2-
+ 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ +0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.35
MnO4
-
+ 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
-ve
reduction
potential
+ve
reduction
potential
Click here std analogy video
Click here std analogy
Click here chem database
(std electrode potential)
Compared to
H2 as std
Eθ
cell/CellPotential= EMF in volt
EMF prod when half cell connect to SHE at std condition
Std electrode potential written as std reduction potential

15. Zn half cell (-ve)
Oxidation
H2 half cell (+ve)
Reduction
Anode Cathode
Zn(s) | Zn2+
(aq) || H+
(aq) , H2(g) | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = 0.00 – ( Eθ Zn )
0.76 = 0.00 - Eθ Zn
Eθ Zn = -0.76V
Zn2+ + 2e ↔ Zn Eθ = ?
2H+ + 2e ↔ H2 Eθ = 0.00V
Std electrode potential as std reductionpotential
Find Eθ
cell (use formula)
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn ????
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2-
+ 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.35
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
-0.76V
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
Zn/H2 = 0.76V
Zn/H2
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)
Zn
Zn2+ H+
Pt
H2
-
-
- +
-e
Zn/H2 Cell Determine Eθ cell Zn/Zn2+
Zn2+ + 2e →Zn Eθ
= -0.76V

16. H2 half cell (-ve)
Oxidation
Fe3+/2+ half cell (+ve)
Reduction
Anode Cathode
Pt(s) | H2, H+
(aq) || Fe3+ Fe2+ | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Std electrode potential as std reductionpotential
Find Eθ
cell (use formula)
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ ?????
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.35
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+0.77V
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Pt
Fe3+
H+
Pt
H2
+
+
+--
-e
H2 /Fe3+,Fe2+ Cell
H2 /Fe3+,Fe2+
2H+ + 2e ↔ H2 Eθ = 0.00V
Fe3+ + e ↔ Fe2+ Eθ = ????
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = Eθ Fe3+ – (-0.00)
0.77 = Eθ Fe3+
Determine Eθ cell Fe 3+/Fe2+
Eθ
H2 /Fe3+ = 0.77V
Fe3+ + e →Fe2+ Eθ
= +0.77V
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

17. H2 half cell (-ve)
Oxidation
CI2 half cell (+ve)
Reduction
Anode
Pt(s) | H2, H+
(aq) || CI2 ,CI- | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Std electrode potential as std reductionpotential
Find Eθ
cell (use formula)
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- ?????
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+1.35V
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
H+
Pt
H2 --
-e
H2 /CI2 Cell
2H+ + 2e ↔ H2 Eθ = 0.00V
CI + e ↔ CI- Eθ = ?????
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = Eθ CI2 – (-0.00)
1.35 = Eθ CI2
H2 /CI2 Cell
+
Pt
CI - CI2
Determine Eθ cell H2 /CI2
Eθ
H2 /CI2 = 1.35V
1/2CI- + e →CI- Eθ
= +1.35V
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

18. Zn half cell (-ve)
Oxidation
Cu half cell (+ve)
Reduction
Anode Cathode
Zn(s) | Zn2+
(aq) || Cu2+
(aq) | Cu (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Zn/Cu Voltaic Cell
-e -e
Zn/Cu half cell
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +0.34 – (-0.76) = +1.10V
Zn 2+ + 2e ↔ Zn (anode) Eθ = -0.76V
Cu2+ + 2e ↔ Cu (cathode) Eθ = +0.34V
Std electrode potential as std reduction potential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
Zn + Cu2+ → Zn2+ + Cu Eθ = ?????
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Zn 2+ + 2e ↔ Zn Eθ = -0.76V
Cu2+ + 2e ↔ Cu Eθ = +0.34V
Zn ↔ Zn2+ + 2e Eθ = +0.76V
Cu2+ + 2e ↔ Cu Eθ = +0.34V
Zn + Cu2+ → Zn 2+ + Cu Eθ = +1.10V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn - 0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu + 0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ +0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.35
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+1.10 V
Eθ
Zn/Cu = 1.10V
Cu2+
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
-
-
-
-
Zn Cu
+
+
+
+
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

19. Zn half cell (-ve)
Oxidation
Ag half cell (+ve)
Reduction
Anode Cathode
Zn(s) | Zn2+
(aq) || Ag+
(aq) | Ag (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Zn/Ag Voltaic Cell
-e -e
Zn/Ag half cell
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +0.80 – (-0.76) = +1.56V
Zn 2+ + 2e ↔ Zn (anode) Eθ = -0.76V
Ag+ + e ↔ Ag(cathode) Eθ = +0.80V
Std electrode potential as std reduction potential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
Zn + Ag+ → Zn2+ + Ag Eθ = ?????
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Zn 2+ + 2e ↔ Zn Eθ = -0.76V
Ag+ + e ↔ Ag Eθ = +0.80V
Zn ↔ Zn2+ + 2e Eθ = +0.76V
Ag+ + e ↔ Ag Eθ = +0.80V
Zn + Ag+ → Zn 2+ + Ag Eθ = +1.56V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn - 0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ +0.77
Ag+ + e- ↔ Ag + 0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+1.56 V
Ag
Eθ
Zn/Ag = 1.56V
Ag+
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
-
-
-
-
+
+
+
+
Zn
Eθ value DO NOT dependon stoichiometric coefficient
(Independentof stoichiometric eqn)

20. Cu half cell (-ve)
Oxidation
Ag half cell (+ve)
Reduction
Anode Cathode
Cu(s) | Cu2+
(aq) || Ag+
(aq) | Ag (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Cu/Ag Voltaic Cell
-e -e
Cu/Ag half cell
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +0.80 – (+0.34) = +0.46V
Cu 2+ + 2e ↔ Cu (anode) Eθ = +0.34V
Ag+ + e ↔ Ag(cathode) Eθ = +0.80V
Std electrode potential as std reductionpotential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
Cu + 2Ag+ → Cu2+ + 2Ag Eθ = ?????
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Cu 2+ + 2e ↔ Cu Eθ = +0.34V
Ag+ + e ↔ Ag Eθ = +0.80V
Cu ↔ Cu2+ + 2e Eθ = -0.34V
2Ag+ + e ↔ 2Ag Eθ = +0.80V
Cu + 2Ag+→ Cu 2+ + 2Ag Eθ = +0.46V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ +0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+0.46V
AgCu
Cu2+
Half cell- high electrodepotential is cathode(+)
Half cell - low electrodepotential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
Cu/Ag = 0.46V
Ag+
-
-
-
-
+
+
+
+
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

21. Mn half cell (-ve)
Oxidation
Ni half cell (+ve)
Reduction
Anode Cathode
Mn(s) | Mn2+
(aq) || Ni2+
(aq) | Ni (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Mn/Ni Voltaic Cell
-e -e
Mn/Ni half cells
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = -0.26 – (-1.19) = +0.93V
Mn 2+ + 2e ↔ Mn (anode) Eθ = -1.19V
Ni2+ + 2e ↔ Ni (cathode) Eθ = -0.26V
Std electrode potential as std reduction potential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
Mn + Ni2+ → Mn2+ + Ni Eθ = ?????
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Mn 2+ + 2e ↔ Mn Eθ = -1.19V
Ni2+ + 2e ↔ Ni Eθ = -0.26V
Mn ↔ Mn2+ + 2e Eθ = +1.19V
Ni2+ + 2e ↔ Ni Eθ = -0.26V
Mn + Ni2+ → Mn2+ + Ni Eθ = +0.93V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni - 0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ +0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
-
+ 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+0.93 V
Eθ
Mn/Ni = 0.93V
Ni2+
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
-
-
-
-
NiMn
+
+
+
+Mn2+
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

22. Fe half cell (-ve)
Oxidation
MnO4- half cell (+ve)
Reduction
Anode Cathode
Fe(s) | Fe2+
(aq) || MnO4
- ,H+, Mn2+ | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Fe/MnO4
- Voltaic Cell
-e -e
Fe/MnO4
- half cells
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +1.51 – (-0.45) = +1.96V
Fe2+ + 2e ↔ Fe Eθ = -0.45V
MnO4
- + 5e ↔ Mn2+ + 4H2O E θ = +1.51V
Std electrode potential asstd reduction potential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
5Fe + 2MnO4
- + 16H+→5Fe2+ +2Mn2+ + 8H2O Eθ = ?
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Fe 2+ + 2e ↔ Fe Eθ = -0.45V
MnO4
- + 5e ↔ Mn2+ + 4H2O Eθ = +1.51V
Fe ↔ Fe2+ + 2e Eθ = +0.45V
MnO4
- +5e ↔ Mn2+ + 4H2O Eθ = +1.51V
Fe + MnO4
- → Mn2+ + Fe2+ Eθ = +1.96V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ +0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+1.96V
PtFe
Fe2+
Eθ
Fe/MnO4
- = 1.96V
MnO4
-
Mn2+
Using platinumelectrode
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
-
-
-
-
+
+
+
+
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

23. Zn half cell (-ve)
Oxidation
Fe3+/2+ half cell (+ve)
Reduction
Anode Cathode
Zn(s) | Zn2+
(aq) || Fe3+ , Fe2+
(aq) | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Zn/Fe3+,Fe2+ Cell
-e -e
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +0.77 – (-0.76) = +1.53V
Zn2+ + 2e ↔ Zn Eθ = -0.76V
Fe3+ + e ↔ Fe2+ Eθ = +0.77V
Std electrode potential asstd reduction potential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
Zn + 2Fe3+→ Zn2+ +2Fe2+ Eθ = ?
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Zn 2+ + 2e ↔ Zn Eθ = -0.76V
Fe3+ + e ↔ Fe2+ Eθ = +0.77V
Zn ↔ Zn2+ + 2e Eθ = +0.76V
Fe3+ +e ↔ Fe2+ Eθ = +0.77V
Zn + 2Fe3+ → Zn2+ + 2Fe2+ Eθ = +1.53V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+1.53V
PtZn
Zn2+
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
Zn/Fe3+ = 1.53V
Fe3+-
Fe2+
Using platinumelectrode
Zn/Fe3+,Fe2+
-
-
-
-
+
+
+
+
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

24. Zn half cell (-ve)
Oxidation
I2 half cell (+ve)
Reduction
Anode Cathode
Zn(s) | Zn2+
(aq) || I2 , I-
(aq) | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Zn/I2 , I- Cell
-e -e
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +0.54 – (-0.76) = +1.30V
Zn2+ + 2e ↔ Zn Eθ = -0.76V
I2 + 2e ↔ 2I- Eθ = +0.54V
Std electrode potential asstd reduction potential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
Zn + I2 → Zn2+ +2I- Eθ = ?
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Zn ↔ Zn2+ + 2e Eθ = +0.76V
I2 + 2e ↔ 2I- Eθ = +0.54V
Zn + I2 → Zn2+ + 2I- Eθ = +1.30V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
-
+ 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+1.30V
PtZn
Zn2+
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
Zn/I2 = 1.30V
I--
I2
Using platinumelectrode
-
-
-
-
+
+
+
+
Zn/I2 , I-
Zn2+ + 2e ↔ Zn Eθ = -0.76V
I2 + 2e ↔ 2I- Eθ = +0.54V
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

25. Zn half cell (-ve)
Oxidation
H2 half cell (+ve)
Reduction
Anode Cathode
Zn(s) | Zn2+
(aq) || H+
(aq) , H2(g) | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = 0.00 – (-0.76) = +0.76V
Zn2+ + 2e ↔ Zn Eθ = -0.76V
2H+ + 2e ↔ H2 Eθ = 0.00V
Std electrode potential as std reductionpotential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
Zn + 2H+→ Zn2+ + H2 Eθ = ?
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Zn 2+ + 2e ↔ Zn Eθ = -0.76V
2H+ + 2e ↔ H2 Eθ = 0.00V
Zn ↔ Zn2+ + 2e Eθ = +0.76V
2H+ +2e ↔ H2 Eθ = 0.00V
Zn + 2H+ → Zn2+ + H2 Eθ = +0.76V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2-
+ 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+0.76V
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
Zn/H2 = 0.76V
Using platinumelectrode/H2
Zn/H2
Zn
Zn2+ H+
Pt
H2
-
-
- +
-e
Zn/H2 Cell
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

26. H2 half cell (-ve)
Oxidation
Ag half cell (+ve)
Reduction
Anode Cathode
Pt(s) | H2, H+
(aq) || Ag+
(aq) | Ag (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
H2/Ag Cell
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +0.80 – (-0.00) = +0.80V
2H+ + 2e ↔ H2 Eθ = 0.00V
Ag+ + e ↔ Ag Eθ = +0.80V
Std electrode potential as std reductionpotential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
H2 + 2Ag+ → 2H+ + 2Ag Eθ = ?
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
H2 ↔ 2H+ + 2e Eθ = +0.00V
2Ag+ +2e ↔ 2Ag Eθ = +0.80V
H2 + 2Ag+ → 2H+ + 2Ag Eθ = +0.80V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2-
+ 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+0.80V
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
H2 /Ag = 0.80V
Using platinumelectrode/H2
H2/Ag
Ag
Ag+
H+
Pt
H2
2H+ + 2e ↔ H2 Eθ = 0.00V
Ag+ + e ↔ Ag Eθ = +0.80V
+
+
+--
-e
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

27. H2 half cell (-ve)
Oxidation
Fe3+/2+ half cell (+ve)
Reduction
Anode Cathode
Pt(s) | H2, H+
(aq) || Fe3+ Fe2+ | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Std electrode potential as std reductionpotential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
H2 + 2Fe3+ → 2H+ + 2Fe Eθ = ?
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
H2 ↔ 2H+ + 2e Eθ = +0.00V
2Fe3+ +2e ↔ 2Fe2+ Eθ = +0.77V
H2 + 2Ag+ → 2H+ + 2Ag Eθ = +0.77V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+0.77V
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
H2 /Fe3+ = 0.77V
Using platinum electrode/H2
Pt
Fe3+
H+
Pt
H2
+
+
+--
-e
H2 /Fe3+,Fe2+ Cell
H2 /Fe3+,Fe2+
2H+ + 2e ↔ H2 Eθ = 0.00V
Fe3+ + e ↔ Fe2+ Eθ = +0.77V
2H+ + 2e ↔ H2 Eθ = 0.00V
Fe3+ + e ↔ Fe2+ Eθ = +0.77V
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +0.77– (-0.00) = +0.77V
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

28. H2 half cell (-ve)
Oxidation
CI2 half cell (+ve)
Reduction
Anode Cathode
Pt(s) | H2, H+
(aq) || CI2 ,CI- | Pt (s)
Cell diagram
Anode Cathode
Half Cell Half Cell
(Oxidation) (Reduction)
Salt Bridge Flow
electrons
Std electrode potential as std reductionpotential
Find Eθ
cell (use reduction potential)Find Eθ
cell (use formula)
CI2 + H2 → 2CI- + 2H+ Eθ = ?
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
H2 ↔ 2H+ + 2e Eθ = +0.00V
CI2 +2e ↔ 2CI- Eθ = +1.35V
H2 + CI2 → 2H+ + 2CI- Eθ = +1.35V
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ 1/2H2 + OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2- + 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ + 0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ + 7H2O +1.33
1/2CI2 + e- ↔ CI- +1.35
MnO4
- + 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
+
+1.35V
+ve/high electrode potentialis cathode(+)
-ve/ low electrode potential is anode (-)
Electronsflow from anode (- ) to cathode (+ )
Eθ
H2 /CI2 = 1.35V
Using platinum electrode/H2
H+
Pt
H2 --
-e
H2 /CI2 Cell
2H+ + 2e ↔ H2 Eθ = 0.00V
CI + e ↔ CI- Eθ = +1.35V
Eθ
cell = Eθ
(cathode) – Eθ
(anode)
Eθ
cell = +1.35 – (-0.00) = +1.35V
H2 /CI2 Cell
2H+ + 2e ↔ H2 Eθ = 0.00V
CI + e ↔ CI- Eθ = +1.35V
+
Pt
CI - CI2
Eθ value DO NOT depend on stoichiometric coefficient
(Independentof stoichiometric eqn)

29. Standard Electrode Potential
STANDARD Reduction potential – H2 as std
Oxidized sp ↔ Reduced sp Eθ/V
Li+ + e- ↔ Li -3.04
K+ + e- ↔ K -2.93
Ca2+ + 2e- ↔ Ca -2.87
Na+ + e- ↔ Na -2.71
Mg 2+ + 2e- ↔ Mg -2.37
Al3+ + 3e- ↔ AI -1.66
Mn2+ + 2e- ↔ Mn -1.19
H2O + e- ↔ H2+OH- -0.83
Zn2+ + 2e- ↔ Zn -0.76
Fe2+ + 2e- ↔ Fe -0.45
Ni2+ + 2e- ↔ Ni -0.26
Sn2+ + 2e- ↔ Sn -0.14
Pb2+ + 2e- ↔ Pb -0.13
H+ + e- ↔ 1/2H2 0.00
Cu2+ + e- ↔ Cu+ +0.15
SO4
2-
+ 4H+ + 2e- ↔ H2SO3 + H2O +0.17
Cu2+ + 2e- ↔ Cu +0.34
1/2O2 + H2O +2e- ↔ 2OH- +0.40
Cu+ + e- ↔ Cu +0.52
1/2I2 + e- ↔ I- +0.54
Fe3+ + e- ↔ Fe2+ +0.77
Ag+ + e- ↔ Ag +0.80
1/2Br2 + e- ↔ Br- +1.07
1/2O2 + 2H+ +2e- ↔ H2O +1.23
Cr2O7
2-+14H+ +6e- ↔ 2Cr3+ +7H2O +1.33
1/2CI2 + e- ↔ CI- +1.36
MnO4
-
+ 8H+ + 5e- ↔ Mn2+ + 4H2O +1.51
1/2F2 + e- ↔ F +2.87
-ve
reduction
potential
+ve
reduction
potential
Compared to
H2 as std
Eθ
cell/CellPotential= EMF in volt
EMF when half cell connect to SHE std condition
Std potentialwritten as std reduction potential
TOP right
• High ↑ tendency lose e
• Li → Li +
+ e
• Eθ
Li = +3.04V
• STRONG reducingAgent
•Oxi favourable (Eθ =+ve)
STRONG
Reducing Agent
WEAK
Reducing Agent
BOTTOM right
• Low ↓ tendency lose e
• F - → 1/2F2 + e
• Eθ
F2 = - 2.87V
• WEAK reducingAgent
•Oxi NOT favourable (Eθ =-ve)
WEAK
Oxidizing Agent
Strong
Oxidizing Agent
TOP left
• Low ↓ tendency gain e
• Li+
+ e → Li
• Eθ
Li= - 3.04V
• WEAK oxidizingAgent
• Red NOT favourable
(Eθ =-ve)
BOTTOM left
• High ↑ tendency gain e
• F2 + 2e → 2F-
• Eθ
F2= +2.87V
• STRONG oxidizing Agent
•Red favourable
(Eθ =+ve)

30. Acknowledgements
Thanks to source of pictures and video used in this presentation
Thanks to Creative Commons for excellent contribution on licenses
http://creativecommons.org/licenses/
http://spmchemistry.onlinetuition.com.my/2013/10/electrolytic-cell.html
http://www.chemguide.co.uk/physical/redoxeqia/introduction.html
http://educationia.tk/reduction-potential-table
http://2012books.lardbucket.org/books/principles-of-general-chemistry-v1.0/s23-
electrochemistry.html
Prepared by Lawrence Kok
Check out more video tutorials from my site and hope you enjoy this tutorial
http://lawrencekok.blogspot.com