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

Types voltaic cell

Conversion electrical energy to chemical energy

Electrochemistry

Electrolytic cell Voltaic 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 cell Dry cell Nickel cadmium cell

Primary cell (Non rechargeable)

MnO2 and KOH

Secondary cell (Rechargeable)

Conversion electrical to chemical energy

Electrochemistry

Electrolytic cell Voltaic cell

Conversion chemical to electrical energy

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)

Current – measured in Amperes or Coulombs per second 1A = 1 Coulomb charge pass through a point in 1 second = 1C/s 1 Coulomb charge (electron) = 6.28 x 10 18 electrons passing in 1 second 1 electron/proton carry 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

CoulombA

sec.1

.1028.6

sec1

11

18

Click here current/voltage

Current Electric Current – moving charges in solid wire or solution

Flow of

charges

-

-

-

Solid/Wire Solution/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

https://www.youtube.com/watch?v=1xPjES-sHwg

https://www.youtube.com/watch?v=1xPjES-sHwg

http://www.howequipmentworks.com/electricity_basics/

Electric Potential

C

JVolt

11

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/Potential Energy

+ -

1 Ω 2 Ω

Charges (-ve)

flow down

AR

VI

RIV

23

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

D Electric 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

EMF vs PD

V = Potential Diff

I = Current

R = Resistance

Max potential diff bet two electrodes of battery source.

+ -

1 Ω 2 Ω

AR

VI

RIV

23

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 (Electromotive Force) in Volt Battery = EMF = 9V

9 Volt

).(9 currentnoVEMFV

IRV

EMF Internal resistance Ir

Place voltmeter across – EMF= 9V No current flow.

ArR

EI

rRIE

IrIREMFE

19

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 Volt 2 Volt

Charges passing through wire

Current flow Circuit complete

https://www.youtube.com/watch?v=Che8SRxp92I

https://www.youtube.com/watch?v=Che8SRxp92I

http://www.bbc.co.uk/bitesize/higher/physics/elect/energy_volts/revision/2/

http://www.bbc.co.uk/bitesize/higher/physics/elect/energy_volts/revision/2/

http://cnx.org/contents/650c4bb3-d190-44fd-9f2f-445a783abdec@4/Electric_Potential_Energy:_Pot

http://cnx.org/contents/650c4bb3-d190-44fd-9f2f-445a783abdec@4/Electric_Potential_Energy:_Pot

Series vs Parallel Circuit

3 Ω

AR

VI

RIV

5.018

9

VV

V

RIV

5.2

55.0

VV

V

RIV

5

105.0

Total Resistance

3 + 10 + 5 = 18Ω

Parallel Circuit Series 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 Ω 0.5 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

VI

9.010

9

AI

R

VI

5.42

9

AI

R

VI

91

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 4.5A 0.9A 9A

https://www.youtube.com/watch?v=XoHxAAnIOaw

https://www.youtube.com/watch?v=XoHxAAnIOaw

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 electrode potential can’t measured. Only Diff in electrode potential can 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/Cu 2+ can’t be measured. Abs electrode potential can’t measured. Only Diff in electrode potential can be measured.

Click here chem database (std electrode potential)


Click here interactive ECS Click here pdf version ECS

Cu Half Cell

http://www.fptl.ru/biblioteka/spravo4niki/handbook-of-Chemistry-and-Physics.pdf

http://www.hbcpnetbase.com/


http://www.hbcpnetbase.com/

http://www2.hkedcity.net/sch_files/a/lsc/lsc-chem/public_html/nss/redox/ecs.html

http://downloads.gphysics.net/UACh/Medicina/TablasIones.pdf

http://www2.hkedcity.net/sch_files/a/lsc/lsc-chem/public_html/nss/redox/ecs.html

http://downloads.gphysics.net/UACh/Medicina/TablasIones.pdf

Potential Diff Cu/Cu2+


Potential Diff Zn/Zn2+


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 → Cu Zn → 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

-

-

-

-

+

+

+

+





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

+ + +

-


- - -

- - -

- - - - - - - -




Voltmeter – High resistance (No current flow) Salt Bridge – flow of ions

Complete the circuit

Cu2+ + 2e → Cu Zn → 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 boundary Salt Bridge Flow

electrons



Zn2+ increase ↑


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

-

-

-

-

+

+

+

+

Potential Diff Ag/Ag2+

Electrode potential Ag/Ag2+ = +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

+ + +

-


- - -

- - -

- - - - - - - -



Ag half cell (+ve) Reduction



Ag+ + e → Ag Zn → Zn 2+ + 2e



1.56 Volt

Zn + 2Ag+ → Zn2+ + 2Ag

Anode Cathode

Zn(s) | Zn2+(aq) || Ag+

(aq)| Ag (s)

Cell diagram

Anode Cathode



electrons



Zn2+ increase ↑


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

-

-

-

-

+

+

+

+

Potential Diff Ag/Ag2+

Electrode potential Ag/Ag2+ = +ve


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

+ + +

-


- - -

- - -

- - - - - - - -


Cu half cell (-ve) Oxidation




Ag+ + e → Ag Cu → Cu 2+ + 2e



0.46 Volt

Cu + 2Ag+ → Cu2+ + 2Ag

Anode Cathode

Cu(s) | Cu2+(aq) || Ag+

(aq)| Ag (s)

Cell diagram

Anode Cathode



electrons



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

-

-

-

-

+

+

+

+

Standard Electrode Potential

Standard Hydrogen Electrode (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?

http://2012books.lardbucket.org/books/principles-of-general-chemistry-v1.0/s23-electrochemistry.html


Std Hydrogen Electrode (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 Cell with SHE as reference electrode

CI-

Pt

+

+

+

Pt


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


Compared to

H2 as std

Eθ cell/Cell Potential = EMF in volt EMF prod when half cell connect to SHE at std condition Std electrode potential written as std reduction potential

https://www.youtube.com/watch?v=0dLBwyTh9Dg

https://www.youtube.com/watch?v=0dLBwyTh9Dg

http://www.chemguide.co.uk/physical/redoxeqia/introduction.html





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

Eθcell = Eθ

(cathode) – Eθ (anode)

Eθcell = 0.00 – ( Eθ Zn )

0.76 = 0.00 - Eθ Zn Eθ Zn = -0.76V

Zn2+ + 2e ↔ Zn Eθ = ? 2H+ + 2e ↔ H2 E

θ = 0.00V

Std electrode potential as std reduction potential

Find Eθcell (use formula)

Eθcell = Eθ

(cathode) – Eθ(anode)


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 potential is cathode (+) -ve/ low electrode potential is anode (-) Electrons flow from anode (- ) to cathode (+ )

Eθ Zn/H2 = 0.76V

Zn/H2

Eθ value DO NOT depend on stoichiometric coefficient (Independent of stoichiometric eqn)

Zn

Zn2+

H+

Pt

H2

-

-

- +

-e

Zn/H2 Cell Determine Eθ cell Zn/Zn2+

Zn2+ + 2e →Zn Eθ = -0.76V

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


Salt Bridge Flow

electrons



Eθcell = Eθ



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


Pt

Fe3+

H+

Pt

H2

+

+

+ - -

-e

H2 /Fe3+,Fe2+ Cell

H2 /Fe3+,Fe2+

2H+ + 2e ↔ H2 Eθ = 0.00V

Fe3+ + e ↔ Fe2+ Eθ = ????

Eθcell = Eθ


Eθcell = 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



CI2 half cell (+ve) Reduction

Anode

Pt(s) | H2, H+

(aq) || CI2 ,CI- | Pt (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons



Eθcell = Eθ




H+ + e- ↔ 1/2H2 0.00 Cu2+ + e- ↔ Cu+ +0.15 SO4


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


H+

Pt

H2 - -

-e

H2 /CI2 Cell

2H+ + 2e ↔ H2 Eθ = 0.00V

CI + e ↔ CI- Eθ = ?????

Eθcell = Eθ


Eθcell = 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




Anode Cathode

Zn(s) | Zn2+(aq) || Cu2+

(aq) | Cu (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons

Zn/Cu Voltaic Cell

-e -e

Zn/Cu half cell

Eθcell = Eθ


Eθcell = +0.34 – (-0.76) = +1.10V

Zn 2+ + 2e ↔ Zn (anode) Eθ = -0.76V Cu2+ + 2e ↔ Cu (cathode) Eθ = +0.34V


Find Eθcell (use reduction potential) Find Eθ

cell (use formula)

Zn + Cu2+ → Zn2+ + Cu Eθ = ?????

Eθcell = Eθ


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



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+


-

-

-

-

Zn Cu

+

+

+

+




Anode Cathode

Zn(s) | Zn2+(aq) || Ag+

(aq) | Ag (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons

Zn/Ag Voltaic Cell

-e -e

Zn/Ag half cell

Eθcell = Eθ


Eθcell = +0.80 – (-0.76) = +1.56V

Zn 2+ + 2e ↔ Zn (anode) Eθ = -0.76V Ag + + e ↔ Ag(cathode) Eθ = +0.80V



cell (use formula)

Zn + Ag+ → Zn2+ + Ag Eθ = ?????

Eθcell = Eθ


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



Zn2+ + 2e- ↔ Zn - 0.76


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+


-

-

-

-

+

+

+

+

Zn


Cu half cell (-ve) Oxidation


Anode Cathode

Cu(s) | Cu2+(aq) || Ag+

(aq) | Ag (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons

Cu/Ag Voltaic Cell

-e -e

Cu/Ag half cell

Eθcell = Eθ


Eθcell = +0.80 – (+0.34) = +0.46V

Cu 2+ + 2e ↔ Cu (anode) Eθ = +0.34V Ag + + e ↔ Ag(cathode) Eθ = +0.80V



cell (use formula)

Cu + 2Ag+ → Cu2+ + 2Ag Eθ = ?????

Eθcell = Eθ


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


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

Ag Cu

Cu2+

Half cell- high electrode potential is cathode (+) Half cell - low electrode potential is anode (-) Electrons flow from anode (- ) to cathode (+ )

Eθ Cu/Ag = 0.46V

Ag+

-

-

-

-

+

+

+

+


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


Salt Bridge Flow

electrons

Mn/Ni Voltaic Cell

-e -e

Mn/Ni half cells

Eθcell = Eθ


Eθcell = -0.26 – (-1.19) = +0.93V

Mn 2+ + 2e ↔ Mn (anode) Eθ = -1.19V Ni2+ + 2e ↔ Ni (cathode) Eθ = -0.26V



cell (use formula)

Mn + Ni2+ → Mn2+ + Ni Eθ = ?????

Eθcell = Eθ


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


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-+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+


-

-

-

-

Ni Mn

+

+

+

+ Mn2+


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


Salt Bridge Flow

electrons

Fe/MnO4- Voltaic Cell

-e -e

Fe/MnO4- half cells

Eθcell = Eθ


Eθcell = +1.51 – (-0.45) = +1.96V

Fe2+ + 2e ↔ Fe Eθ = -0.45V MnO4

- + 5e ↔ Mn2+ + 4H2O E θ = +1.51V



cell (use formula)

5Fe + 2MnO4- + 16H+→5Fe2+ +2Mn2+ + 8H2O Eθ = ?

Eθcell = Eθ


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


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-+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

Pt Fe

Fe2+

Eθ Fe/MnO4- = 1.96V

MnO4-

Mn2+

Using platinum electrode


-

-

-

-

+

+

+

+




Anode Cathode

Zn(s) | Zn2+(aq) || Fe3+ , Fe2+

(aq) | Pt (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons

Zn/Fe3+,Fe2+ Cell

-e -e

Eθcell = Eθ


Eθcell = +0.77 – (-0.76) = +1.53V

Zn2+ + 2e ↔ Zn Eθ = -0.76V Fe3+ + e ↔ Fe2+ Eθ = +0.77V



cell (use formula)

Zn + 2Fe3+→ Zn2+ +2Fe2+ Eθ = ?

Eθcell = Eθ


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



Zn2+ + 2e- ↔ Zn -0.76


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

Pt Zn

Zn2+


Eθ Zn/Fe3+ = 1.53V

Fe3+-

Fe2+


Zn/Fe3+,Fe2+

-

-

-

-

+

+

+

+



I2 half cell (+ve) Reduction

Anode Cathode

Zn(s) | Zn2+(aq) || I2 , I

-(aq) | Pt (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons

Zn/I2 , I- Cell

-e -e

Eθcell = Eθ


Eθcell = +0.54 – (-0.76) = +1.30V

Zn2+ + 2e ↔ Zn Eθ = -0.76V I2

+ 2e ↔ 2I- Eθ = +0.54V



cell (use formula)

Zn + I2 → Zn2+ +2I- Eθ = ?

Eθcell = Eθ


Zn ↔ Zn2+ + 2e Eθ = +0.76V I2

+ 2e ↔ 2I- Eθ = +0.54V Zn + I2

→ Zn2+ + 2I- Eθ = +1.30V



Zn2+ + 2e- ↔ Zn -0.76


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


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

Pt Zn

Zn2+


Eθ Zn/I2 = 1.30V

I--

I2


-

-

-

-

+

+

+

+

Zn/I2 , I-

Zn2+ + 2e ↔ Zn Eθ = -0.76V I2

+ 2e ↔ 2I- Eθ = +0.54V



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


Salt Bridge Flow

electrons

Eθcell = Eθ


Eθcell = 0.00 – (-0.76) = +0.76V

Zn2+ + 2e ↔ Zn Eθ = -0.76V 2H+ + 2e ↔ H2 E

θ = 0.00V



cell (use formula)

Zn + 2H+→ Zn2+ + H2 Eθ = ?

Eθcell = Eθ


Zn 2+ + 2e ↔ Zn Eθ = -0.76V 2H+ + 2e ↔ H2 E

θ = 0.00V

Zn ↔ Zn2+ + 2e Eθ = +0.76V 2H+ +2e ↔ H2

Eθ = 0.00V Zn + 2H+ → Zn2+ + H2 Eθ = +0.76V



Zn2+ + 2e- ↔ Zn -0.76


H+ + e- ↔ 1/2H2 0.00 Cu2+ + e- ↔ Cu+ +0.15 SO4


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


Eθ Zn/H2 = 0.76V

Using platinum electrode/H2

Zn/H2

Zn

Zn2+

H+

Pt

H2

-

-

- +

-e

Zn/H2 Cell




Anode Cathode

Pt(s) | H2, H+

(aq) || Ag+(aq) | Ag (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons

H2/Ag Cell

Eθcell = Eθ


Eθcell = +0.80 – (-0.00) = +0.80V

2H+ + 2e ↔ H2 Eθ = 0.00V

Ag+ + e ↔ Ag Eθ = +0.80V



cell (use formula)

H2 + 2Ag+ → 2H+ + 2Ag Eθ = ?

Eθcell = Eθ


H2 ↔ 2H+ + 2e Eθ = +0.00V 2Ag+ +2e ↔ 2Ag Eθ = +0.80V H2 + 2Ag+ → 2H+ + 2Ag Eθ = +0.80V



Zn2+ + 2e- ↔ Zn -0.76


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


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


Eθ H2 /Ag = 0.80V


H2/Ag

Ag

Ag+

H+

Pt

H2

2H+ + 2e ↔ H2 Eθ = 0.00V

Ag+ + e ↔ Ag Eθ = +0.80V

+

+

+ - -

-e




Anode Cathode

Pt(s) | H2, H+

(aq) || Fe3+ Fe2+

| Pt (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons



cell (use formula)

H2 + 2Fe3+ → 2H+ + 2Fe Eθ = ?

Eθcell = Eθ


H2 ↔ 2H+ + 2e Eθ = +0.00V 2Fe3+ +2e ↔ 2Fe2+ Eθ = +0.77V H2 + 2Ag+ → 2H+ + 2Ag Eθ = +0.77V



H+ + e- ↔ 1/2H2 0.00 Cu2+ + e- ↔ Cu+ +0.15 SO4


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


Eθ H2 /Fe3+ = 0.77V


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 E

θ = 0.00V Fe3+ + e ↔ Fe2+ Eθ = +0.77V

Eθcell = Eθ


Eθcell = +0.77– (-0.00) = +0.77V



CI2 half cell (+ve) Reduction

Anode Cathode

Pt(s) | H2, H+

(aq) || CI2 ,CI- | Pt (s)

Cell diagram

Anode Cathode

Half Cell Half Cell


Salt Bridge Flow

electrons



cell (use formula)

CI2 + H2 → 2CI- + 2H+ Eθ = ?

Eθcell = Eθ


H2 ↔ 2H+ + 2e Eθ = +0.00V CI2

+2e ↔ 2CI- Eθ = +1.35V H2 + CI2

→ 2H+ + 2CI- Eθ = +1.35V



H+ + e- ↔ 1/2H2 0.00 Cu2+ + e- ↔ Cu+ +0.15 SO4


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


Eθ H2 /CI2

= 1.35V


H+

Pt

H2 - -

-e

H2 /CI2 Cell

2H+ + 2e ↔ H2 Eθ = 0.00V

CI + e ↔ CI- Eθ = +1.35V

Eθcell = Eθ


Eθcell = +1.35 – (-0.00) = +1.35V

H2 /CI2 Cell

2H+ + 2e ↔ H2 Eθ = 0.00V

CI + e ↔ CI- Eθ = +1.35V

+ Pt

CI - CI2



STANDARD Reduction potential – H2 as std


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-+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/Cell Potential = EMF in volt EMF when half cell connect to SHE std condition Std potential written as std reduction potential

TOP right • High ↑ tendency lose e • Li → Li + + e

• Eθ Li = +3.04V • STRONG reducing Agent •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 reducing Agent •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 oxidizing Agent • 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)

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

http://creativecommons.org/licenses/

http://spmchemistry.onlinetuition.com.my/2013/10/electrolytic-cell.html




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Date post:	18-Jul-2015
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IB Chemistry on Voltaic Cell, Standard Electrode Potential and Standard Hydrogen Electrode

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