Lecture #15

Chapter 18 - Electrochemistry

Chapter 18 - Electrochemistry

the branch of chemistry that examines the transformations between chemical and

electrical energy

Redox Chemistry Revisited

Znº(s) + Cu2+(aq) → Zn2+(aq) + Cuº(s)

Sum of two half-reactions: One species gains e– (reduction) while another species loses e– (oxidation) Oxidizing agents vs. reducing agents

Znº(s) → Zn2+(aq) Cu2+(aq) → Cuº(s)

Zn = red. agent; Cu2+ = oxid. agent

A Spontaneous Redox Reaction

Cu2+NO3-Zn2+

K+

?

NO3-

Zn0 + Cu2+ Zn2+ + Cu0

electrons

A Voltaic Cell

Electrochemical Cells

The anode is for oxidation !!

Reduction takes place at the cathode !!

Voltaic Cell

Spontaneous Reaction

Chemical energy is transformed into electrical energy.

Electrolytic Cell

External source of electrical energy

required

Electrical energy is transformed into chemical energy.

Cell Components

Anode = electrode at which oxidation half-reaction (loss of electrons) takes place.

Cathode = electrode at which reduction half-reaction (gain of electrons) takes place.

A bridge connects the two solutions of the cell; balances flow of electrons, eliminates

accumulation of charge in either compartment.

Writing Cell Diagrams

Write chemical symbol of anode at the far left, symbol of cathode at the far right, and a double vertical for connecting

bridge halfway between them.

Work inward from electrodes toward the bridge, using vertical lines to indicate phase changes and symbols of ions

or compounds to represent electrolytes surrounding the electrode that are changed by the cell reaction.

Indicate concentrations of dissolved species and partial pressures of any gases (if known).

Cu(s) + 2 Ag+(aq) Cu2+(aq) + 2 Ag(s)

Cell Diagram Example

1. Cu(s) . . . || . . . Ag(s)

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

3. Cu(s) | Cu2+(1.00 M) || Ag+(1.00 M) | Ag(s)

anode half-cell cathode half-cell

} }

Standard Potentials

NCDPI Reference Tables for Chemistry (October 2006 form A-v1) Page 7

ACTIVITY SERIES of Halogens:

2F

2Cl

2Br

2I

ACTIVITY SERIES of Metals

Li Rb K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb [ 2H ] Sb Bi Cu Hg Ag Pt Au

Polyatomic Ions +4NH Ammonium

3BrO− Bromate

CN− Cyanide

2 3 2C H O−

3(CH COO )− Acetate

4ClO− Perchlorate

3ClO− Chlorate

2ClO − Chlorite

ClO− Hypochlorite

3IO − Iodate

4MnO− Permanganate

3NO− Nitrate

2NO− Nitrite

OH− Hydroxide

3HCO− Hydrogen carbonate

4HSO− Hydrogen sulfate

SCN− Thiocyanate 23CO − Carbonate

22 7Cr O − Dichromate

24CrO − Chromate

24SO − Sulfate 23SO − Sulfite 34PO − Phosphate

React with oxygen to form oxides

Replace hydrogen from acids

Replace hydrogen from steam

Replace hydrogen from cold water

Recollection: The Activity Series of Metals

Atoms are Strong Reducing Agents

Cations are Strong Oxidizing Agents

NCDPI Reference Tables for Chemistry (October 2006 form A-v1) Page 7

ACTIVITY SERIES of Halogens:

2F

2Cl

2Br

2I

ACTIVITY SERIES of Metals

Li Rb K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb [ 2H ] Sb Bi Cu Hg Ag Pt Au

Polyatomic Ions +4NH Ammonium

3BrO− Bromate

CN− Cyanide

2 3 2C H O−

3(CH COO )− Acetate

4ClO− Perchlorate

3ClO− Chlorate

2ClO − Chlorite

ClO− Hypochlorite

3IO − Iodate

4MnO− Permanganate

3NO− Nitrate

2NO− Nitrite

OH− Hydroxide

3HCO− Hydrogen carbonate

4HSO− Hydrogen sulfate

SCN− Thiocyanate 23CO − Carbonate

22 7Cr O − Dichromate

24CrO − Chromate

24SO − Sulfate 23SO − Sulfite 34PO − Phosphate

React with oxygen to form oxides

Replace hydrogen from acids

Replace hydrogen from steam

Replace hydrogen from cold water

Recollection: The Activity Series of Metals

Atoms are Strong Reducing Agents

Cations are Strong Oxidizing Agents

2 Naº + 2 H2O(l) → 2 Na+ + 2 OH- + H2

Snº + H2O + H+ → Sn2+ + OH- + 2 H2

2 Cuº + O2 + → 2 CuO

2 Znº + 4 H2O(g) → 2 Zn2+ + 4 OH- + 2 H2Δ

Standard Potentials

Standard reduction potential (Eº) – the potential of a half-reaction in which all reactants and products are their

standard states at 25ºC.

Standard cell potential (Eºcell) – a measure of how forcefully an electrochemical cell (in standard state) can pump

electrons through an external circuit.

Eºcell = Eºcathode – Eºanode

19.3 Standard Reduction Potentials

Eºcell = Eºcathode – Eºanode

Eºcathode (Cu2+ → Cu(s)) = 0.342 V

Eºanode (Zn2+ → Zn(s)) = –0.762 V

Eocell = (0.342 V) – (–0.762 V) = 1.104 V

Standard Cell Potential (Ecell)

Znº(s) → Zn2+(aq) Cu2+(aq) → Cuº(s)

1) Calculate the standard cell potential for the reaction:

2 Fe3+(aq) + 2 I– (aq) → 2 Fe2+(aq) + I2(s)

Eºcell = Eºcathode – Eºanode

Eºcathode (Fe3+(aq) → Fe2+(aq) ) = + 0.77 V

Eºanode ( I2º(s)→2I- (aq)) = + 0.54 V

Eocell = (+ 0.77 V) – (+ 0.54 V) = + 0.23 V

Fe3+(aq) → Fe2+(aq) 2 I- (aq) → I2º (s)

2) Calculate the standard cell potential for the reaction:

2 NiO(OH)(s) + 2 H2O (l) + Cd (s)→ 2 Ni(OH)2 (s) + Cd(OH)2 (s)

2 NiO(OH) + 2 H2O → 2 Ni(OH)2 + 2 OH-

Cd + 2 OH- → Cd(OH)2

Eºcell = Eºoxidation + Eºreduction

Eocell = (+ 0.81 V) + (+ 0.49 V) = + 1.30 V

Eºoxidation = + 0.81 V

Eºreductiion = + 0.49 V

Eºcell = Eºcathode – Eºanode

(Cd → Cd2+)

(Ni3+ → Ni2+)

2 Fe3+(aq) + 2 I– (aq) → 2 Fe2+(aq) + I2(s)

Cd + 2 OH- → Cd(OH)2

2 NiO(OH)(s) + 2 H2O (l) + Cd (s)→ 2 Ni(OH)2 (s) + Cd(OH)2 (s)

2 NiO(OH) + 2 H2O → 2 Ni(OH)2 + 2 OH-

Znº(s) + Cu2+(aq) → Zn2+(aq) + Cuº(s)

Cu2+(aq) → Cuº(s)

Fe3+(aq) → Fe2+(aq)

(Ni3+ → Ni2+)

Znº(s) → Zn2+(aq)

2 I- (aq) → I2º (s)

(Cd → Cd2+)

Cu2+(aq) → Cuº(s) +0.34

Zn2+(aq) → Znº(s) -0.76

Fe3+(aq) → Fe2+(aq) +0.77

I2º(s) → 2I- (aq) +0.54

Cd2+ → Cdº -0.81

Ni3+ → Ni2+ +0.49

Standard Reduction PotentialsStrongerOxidizingAgents

WeakerOxidizingAgents

StrongerReducingAgents

WeakerReducingAgents

Alessandro Volta Luigi Galvani

Dry Cells

Lead Storage Battery

Fuel Cell

Chemical Energy and Electrical Work

Current and Voltage

Current - the number of electrons that flow through the system per second

1 A = 1 Ampere = 1 Coulomb of charge/second = 6.242 x 1018 electrons/second

Potential difference - the difference in potential energy between reactants and products

1 V of force = 1 J of energy/Coulomb of charge

(The potential difference can also be thought of as the voltage needed to drive electrons through the external circuit.)

Electromotive force (emf) - the amount of force pushing the electrons through the wire

ΔGºcell = Welec = –CEºcell

Welec = work done by the cell C = charge (coulombs) Eºcell = electromotive force (emf); cell voltage,Volts = J/C

ΔGºcell = -nFEºcell Faraday constant (F) is 9.65 × 104 C/(mol e–) n = number of moles of electrons

Voltage and Electrical Work

1.602 x 10-19 C/e 6.022 x 1023 e mol ×

3) Calculate the value of ΔGº and the work done on the circuit for the reaction:

Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s) taking place in a voltaic cell that produces 2.71V.

ΔGºcell = Welec = -nFEºcell = (-2)(9.65 x 104 C/mol e)(2.71 J/C) = -523 J

Mg(s) → Mg2+(aq)

Cu2+(aq) → Cu(s)

[oxidation]

[reduction]

Eocell = (+ 0.34 V) – (– 2.37 V) = + 2.71 VEºcell = Eºcathode – Eºanode

4) Calculate the value of Eºcell for the following reaction by calculating ΔGº:

Cu(s) + 2 Fe3+(aq) → Cu2+(aq) + 2 Fe2+ (aq)

ΔGºcell =[65.5 + 2(-78.9)] - [0.0 + 2(-4.7)]ΔGºcell =[-92.3] - [0.0 - 9.4] = -82.9 kJ

Eºcell = -82.9 x 103 J/mol -(2)(9.65 x 104 C) = 0.430 J/C = 0.430 V

A Reference Point:

The Standard Hydrogen Electrode

A Reference Point: The Standard Hydrogen Electrode

2 H+(aq) + 2 e– →H2(g)

ESHE = 0.00 V

Pt ❘ H2(g), 1.0 atm ❘ H+(1.0 M) ❘ ❘

(Can serve as anode or cathode)

A Reference Point: The Standard Hydrogen Electrode

The Standard Hydrogen Electrode (SHE) reduction potential is defined to be exactly 0.00 V.

Half reactions with a stronger tendency toward reduction that the SHE have a positive value for Eºreduction.

Half reactions with a strong tendency toward oxidation than the SHE have a negative value for Eºreduction.

Eºcell = Eºcathode – Eºanode Eºcell = Eºoxidation + Eºreduction

Eºoxid = -Eºred

When adding Eº values for the half-cells, do not multiply the half-cell Eº values.

!!H2"!!(1!atm)!

0.762 V = ESHE – EZn 0.762 V = 0.00 V – EZn

0.342 V = ECu – ESHE 0.342 V = ECu – 0.00 V

Determination of Eo

Selected Standard Electrode Potentials (298K)

Half-Reaction E0(V)

2H+(aq) + 2e- H2(g)

F2(g) + 2e- 2F-(aq)

Cl2(g) + 2e- 2Cl-(aq)

MnO2(g) + 4H+(aq) + 2e- Mn2+(aq) + 2H2O(l)

NO3-(aq) + 4H+(aq) + 3e- NO(g) + 2H2O(l)

Ag+(aq) + e- Ag(s)

Fe3+(g) + e- Fe2+(aq)

O2(g) + 2H2O(l) + 4e- 4OH-(aq)

Cu2+(aq) + 2e- Cu(s)

N2(g) + 5H+(aq) + 4e- N2H5+(aq)

Fe2+(aq) + 2e- Fe(s)

2H2O(l) + 2e- H2(g) + 2OH-(aq)

Na+(aq) + e- Na(s)

Li+(aq) + e- Li(s)

+2.87

-3.05

+1.36

+1.23

+0.96

+0.80

+0.77

+0.40

+0.34

0.00

-0.23

-0.44

-0.83

-2.71

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Overall Cell Potential

Eºcell = Eºcathode - Eºanode

Eºcell = + 0.34 - (-0.76) = 1.10 V

Eºcell = Eºoxid + Eºred

Eºcell = - (-0.76) + 0.34 =1.10 V

Cu2+ (aq) + 2e ➞ Cu (s)

Zn2+ (aq) + 2e ➞ Zn (s)

Eº(V)0.34

-0.76