Prepared by:Mrs Faraziehan Senusi

PA-A11-7C

Electrochemical Cells

Corrosion & Prevention

Chapter 3Oxidation and Reduction

Oxidation-Reduction Concepts

Voltaic Cell

Electrolytic Cell

Reference: Chemistry: the Molecular Nature of Matter and Change, 6th ed, 2011, Martin S. Silberberg, McGraw-Hill

Voltaic (Galvanic) Cells• Use spontaneous reaction (G < 0) to generate electrical energy• Difference in Chemical Potential energy between higher energy reactants

and lower energy products is converted to electrical energy to power electrical devices

Electrolytic Cells• Uses electrical energy to drive nonspontaneous reaction (G > 0)• Electrical energy from an external power supply converts lower energy

reactants to higher energy products

Free energy and electrical work

DGorxn = S mDGo

f (products) - S nDGof (reactants)

• A spontaneous reaction has a negative free energy change (ΔG < 0), and a spontaneous electrochemical reaction has a positive cell potential (Ecell > 0)

• These two indications of spontaneity are proportional to each other:

ΔG α –Ecell

Free energy and electrical work

ocell

o

cell

-4

--

--

-

cell

nFEΔG

states, standard their in are components theall When

nFEΔGe V.mol

J1065.9

e mol

96485CconstantFaraday F, e of mole 1 of Charge

e mol

chargex e of moles Chargeor nF Charge

:(F) e of mole 1 of charge the timesdtransferre

(n) electrons of moles ofnumber theequals cell the throughflows charge The

chargeEΔG

• We also can relate the standard cell potential to the equilibrium constant of redox reaction.

KnF

RT

KRT

KRT

lnE

lnnFE

lnΔG

ocell

ocell

o

Free energy and electrical work

At equilibrium,

• The interconnections among the standard free energy change, the equilibrium constant, and the standard cell potential.

Free energy and electrical work

(Product favoured)

(Reactant favoured)

FO

RW

AR

D R

EA

CT

ION

RE

VE

RS

E R

EA

CT

ION

Table 20.2 The relationship between DGo and K at 25 oC

DGo (kJ) K significance

200

100

50

10

1

0

-1

-10

-50

-100

-200

9 x 10-36

3 x 10-18

2 x 10-9

2 x 10-2

7 x 10-1

1

1.5

5 x 101

6 x 108

3 x 1017

1 x 1035

Essentially no forward reaction; reverse reaction goes to completion

Forward and reverse reactions proceed to the same extent

Forward reaction goes to completion; essentially no reverse reaction

Example:

• The relation between cell potential and concentration based on the relation between free energy and concentration: ΔG equals ΔG° (the free energy change when the system moves from standard-state concentrations to equilibrium) plus RT In Q (the free energy change when the system moves from nonstandard-state to standard-state concentrations).

ΔG = ΔG° + RT ln Q Where,ΔG = - nFEcell ΔG° = - nFE°cell

So, - nFEcell = - nFE°cell + RT ln Q

Therefore, the Nerst Equation:Q

nF

RTEcell lnEo

cell

Free energy and electrical work

* More discussion will be futher discuss on Chapter 4

Corrosion & Prevention

Metals corrode because they oxidize easily. standard reduction potentials less positive than that of oxygen gas. When any one of these half-reactions is reversed (to show oxidation of the

metal) and combined with the reduction half-reaction for oxygen, the result is a positive E value. Thus the oxidation of most metals by oxygen is spontaneous.

CorrosionOrdinary corrosion is the redox process by which metals are

oxidized by oxygen, O2, in the presence of moisture.

A point of strain (metals are most “active”) in a steel object acts as an anode where the iron is oxidized to iron(II) ions, and pits are formed

Fe Fe2+ + 2e- (oxidation,anode)

The electrons produced then flow through the nail to areas exposed to O2. These act as cathodes where oxygen is reduced to hydroxide ions, OH–.

O2 + 2H2O + 4e- 4OH– (reduction,cathode)

At the same time, the Fe2+ ions migrate through the moisture on the surface.

The overall reaction is obtained by balancing the electron transfer and adding the two half-reactions.

The Fe2+ ions can migrate from the anode through the solution toward the cathode region, where they combine with OH– ions to form iron(II) hydroxide.

Iron is further oxidized by O2 to the 3+ oxidation state. The material we call rust is a complex hydrated form of iron(III) oxides and hydroxides with variable water composition; it can be represented as Fe2O3. xH2O.

The overall reaction for the rusting of iron is

Corrosion

• There are several methods for protecting metals against corrosion.

Applying a protective coating, such as paint Iron objects are frequently painted to keep out O2 and moisture, but if the

paint layer chips, rusting proceeds.

Connecting the metal directly to a “sacrificial anode,” a piece of another metal that is more active and therefore preferentially oxidized

CORROSION PROTECTION

• Another popular coating method is galvanizing.• Is a steel coating method which uses zinc.• Since zinc is a more active metal than iron, as the potentials for

the oxidation half-reactions show, any oxidation that occurs dissolves zinc rather than iron.

Fe Fe2+ + 2e- E = 0.44 V

Zn Zn2+ +2e- E = 0.76 V

• Thus, zinc acts as a “sacrificial” coating on steel, instead of the iron.

CORROSION PROTECTION

Galvanized objects are steel coated with zinc

• Cathodic protection is a method most often used to protect steel in buried fuel tanks and pipelines.

• An active metal, such as magnesium, is connected by a wire to the pipeline or tank to be protected.

• Because magnesium is a better reducing agent than iron, electrons are furnished by the magnesium, keeping the iron from being oxidized.

• Then, the active metal is sacrified instaed of the iron.

CORROSION PROTECTION