Electrochemical Electrochemical SystemSystem

Week Topic Topic Outcomes

8-10 Electrochemical systems

•Electrochemical system and its thermodynamics

•Chemical reactions in electrochemical cells and the Nernstequation

•The relationship between the cell emf and the equilibrium constant, Gibbs energies and reaction entropies

•Thermodynamics of Galvanic cells and Fuel cell

It is expected that students are able to:

•Define the electrochemical system.

•Explain about electrochemical cell and differentiate galvanic cell and electrolyte cell

•Apply Nernst equation of thermodynamic principles to electrochemical cells.

•Determine ∆Go, ∆Ko, ∆So, ∆Ho, ∆Cpo

activity coefficient and pH of cell’s chemical reaction.

•Understand the cell emf in IUPAC convention.

•Estimate the liquid junction potential from emf measurement..

Topic Outcomes

Redox: Reduction-Oxidation, an electron transfer process

Oxidation: loss one or more e-

Reduction: gain one or more e-

Oxidation number: the charge of the atom or molecule would have

Pure element and neutral compound 0

Mono- and poly-atomic ions sum of oxidation number is ionic charge

Quick Review

CaBr2; Ca = +2, Br = -1,Oxidation number (pure element), =(+2)+2(-1)=0

Check your understanding

K2Cr2O7

O = -2 K = +1

Cr = ?

HCO3-

O = -2 H = +1

C = ?

Electrochemistry: study of the interchange between chemical change and electrical work

Electrochemical cells: systems utilizing a redoxreaction to produce or use electrical energy

Quick Review (Cont’d)

Becomes smaller

Becomes larger

Znanode

Cucathode

Cell and Cell Notation

Anode: Electrode site of oxidation

Cathode: Electrode site of reduction

Electrons flow from anode to cathode

Salt bridge: Maintains electrical neutrality, allows current to flow, prevent solutions mixing

Positive ion migrates to cathode

Negative ion migrates to anode

Cell notation: Anode | Salt Bridge | Cathode

The difference in electrical potential between the anode and cathode is called:

• cell voltage

• electromotive force (emf)

• cell potential

000reductionoxidationCell EEE += (Unit: Volts)

Note: Volt (V) = Joule (J) Coulomb (C)

Cell Potential

Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.

E0 = 0 VReference electrode

Standard hydrogen electrode (SHE)

2e- + 2H+ (1 M) H2 (1 atm)

Reduction Reaction

Standard Electrode Potentials

(Anode)

(Cathode)

Note: Copper reduction potential, Eo = +0.34(If copper is reduced (at cathode), voltmeter can give a reading- positive)

Note: Zn reduction potential is negative.Change Zn as anode

E0 is for the reaction as written

The half-cell reactions are reversible; the sign of E0

changes when the reaction is reversed

Changing the stoichiometriccoefficients of a half-cell reaction does not change the value of E0

The more positive E0 the greater the tendency for the substance to be reduced

+ E°CELL : spontaneous reaction

E°CELL = 0 :equilibrium

- E°CELL: non-spontaneous reaction

More positive E°CELL means stronger oxidizing agent or more

likely to be reduced

Spontaneity of Redox Reaction000reductionoxidationCell EEE +=

Eoreduction=-Eooxidation

Write separate equations (half-reactions) for oxidation and reduction

For each half-reaction

Balance elements involved in e- transfer

Balance number e- lost and gained

To balance e- multiply each half-reaction by whole numbers

Add half-reactions/cancel like terms (e-)

Check that all atoms and charges balance

Balance the Half-Cell Reactions

e–

e – e–

e–

Zn 2+

SO4 2–

Zn(s)1.0 M Zn 2+

solution

Anode

1.0 M Cu 2+

solution

Cathode

Cu 2+

SO4 2–

Cu(s)

Check your understanding

??

?? ??

Check your understanding

A half-cell reduction reaction is

Fe(OH)2 (s) + 2e- Fe (s) + 2OH- (aq) Eo=-0.877 V

Is to combine with another half-cell, which the reduction reaction is given as:

a) Al3+ (aq) + 3e- Al (s) Eo= -1.66 V

b) AgBr (s) + e- Ag (s) + Br- (aq) Eo=+0.071 V

Determine Eocell, and whether Fe is oxidized or reduced.

∆G =∆G0 + RTlnQ ∆G0 = -nFE0cell

-nFEcell= -nFE0cell + RTln Q

Ecell= E0cell - RTln QnF

Ecell= E0cell - 0.0257ln Qn

Ecell= E0cell – 0.0592log Qn

Nernst Equation

Note: At equilibrium, ∆G=0, thus Ecell=0

the emf for an electrochemical cell to be calculated if the activity and E°are known.

Faraday, F: charge on 1 mole e-F = 96485 C/mole

∆Go = -nFEocell

ΔG0 = -RT ln K

ΔG0 = -nFE0cell-RT ln K = -nFE0cell

KnFRTECell ln

0 =( )

0257.096485

29831.8=

⎟⎠⎞

⎜⎝⎛

=

moleC

KmolK

J

FRT

Kn

Kn

ECell log0592.0ln0257.00 ==

n= moles of balanced e-

Entropy Change

The reaction entropy related to ΔG°R

pp

RR T

EnFTGS ⎟⎟

⎠

⎞⎜⎜⎝

⎛∂∂

=⎟⎟⎠

⎞⎜⎜⎝

⎛∂Δ∂

−=Δoo

o

Measurement of the T dependence of E° can be used to determine ΔS°R

To obtain the entropy change for the cell reaction

The reaction entropy related to ΔG°R

PTEnFTnFEH ⎟⎟

⎠

⎞⎜⎜⎝

⎛∂∂

+−=Δo

oo

Enthalpy Change

∆Ho = ∆Go + T∆So

Example (from text book)

The half-cell reduction reactions,

Fe3+ (aq) + e- Fe2+ (aq) Eo=+0.771 V

Fe2+ (aq) + 2e- Fe (s) Eo=-0.447 V

Calculate Eo for half-cell reduction reaction

Fe3+ (aq) + 3e- Fe (s)

Example

Calculate ΔGor and the equilibrium constant at 298.15K for two half-cell reduction reactions:

Cr2O72-(aq) + 14H+(aq) + 6e- → 2Cr3+(aq) + 7H2O(l)

Eº = +1.232 V

2H+(aq) + 2e- → H2(g)

Eº = 0.00 V

The standard potential E° for a given cell in which a two

electron transfer takes place is 1.10 V at 298.15 K and

(∂E°/∂T) = –6.50 × 10–5 V K–1. Calculate ΔG°rxn, ΔS°rxn,

ΔH°rxn. Assume that n = 2.

Example

Check your understanding

Calculate the equilibrium constant at

298.15 K for the following half-cell reactions

a) Cathode: NO3- + 4H+ + 3e- → NO + 2H2O Eº = +0.957 V

Anode: 2H2O → O2 + 4H+ + 4e- Eº = –1.229 V

b) Anode: Cd(OH)2 + 2e- → Cd + 2OH- Eº = –0.809V

Cathode: O2 + 2H2O + 4e- → 4OH- Eº = +0.401V

Calculate the equilibrium constant at 25°C for the

reaction occurring in the Daniell cell, if the standard

EMF is 1.10 V.

Calculate the equilibrium constant for the reaction

H2 + 2Fe3+ ⇔ 2H+ + 2Fe2+

Example

Galvanic cell vs. Galvanic cell vs. Electrolytic cellElectrolytic cell

Galvanic cell vs. Electrolytic cell

The external battery supplies the electrons. They enter through the cathode and come out through the anode.

The electrons are supplied by the species getting oxidized. They move from anode to the cathode in the external circuit.

Here, the anode is positive and cathode is the negative electrode. The reaction at the anode is oxidation and that at the cathode is reduction.

Anode is negative and cathode is the positive electrode. The reaction at the anode is oxidation and that at the cathode is reduction.

Both the electrodes are placed in a same container in the solution of molten electrolyte

The two half-cells are set up in different containers, being connected through the salt bridge or porous partition

Redox reaction is not spontaneous and electrical energy has to be supplied to initiate the reaction

Redox reaction is spontaneous and is responsible for the production of electrical energy

Converts electrical energy into chemical energy

Converts chemical energy into electrical energy

Electrolytic cellElectrochemical cell (Galvanic Cell)