CHAPTER 1

ELECTROCHEMICAL NATURE OF AQUEOUS

CORROSION

CMT555

Electrochemical & Corrosion Science

ELECTROCHEMISTRY What is electrochemistry? Electrochemistry is a branch of chemistry that

deals with the interconversion between electrical energy and chemical energy.

Electrical Chemical The conversion takes place in an electrochemical

cell.

REDOX REACTIONS

In an electrochemical cell, there are two half-reactions (redox processes) occur at two different electrodes:

Reaction Electrode Reduction (Gain of electron)

Oxidation (Loss of electron)

Anode

Cathode

* Remember: OIL RIG

Example 1:

Half-reaction Anode:

Cathode:

Overall:

)(2)(22

gaqHeH

)(2

2

)()()( 2 gaqaqs HZnHZn

eZnZn aqs 22

)()(

(reduction)

(oxidation)

A summary of redox terminology Process Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)

OXIDATION One reactant loses electrons. Reducing agent is oxidized. Oxidation number increases.

Zinc loses electrons. Zinc is the reducing agent and become oxidized. The oxidation no. of Zn increases from 0 to +2.

REDUCTION Other reactant gains electrons. Oxidising agent is reduced. Oxidation number decreases.

Hydrogen ion gains electrons. Hydrogen ion is the oxidizing agent and becomes reduced. The oxidation no. of H decreases from +1 to 0.

Oxidising agent: ?

Reducing agent:?

When a piece of zinc is placed in an aqueous solution of copper(II) sulphate, brown solids (copper metal) will form and collect at the bottom of the container. The blue colour of the solution slowly fades, and the piece of zinc gets smaller.

Electron acceptor

Electron donor

In the reaction, the zinc metal gives up electrons and slowly dissolves:

Zn Zn2+ + 2e-

The electrons are taken up by the Cu2+ ions in the solution, and get deposited as copper:

Cu2+ + 2e- Cu

In the process, zinc gets oxidised while the Cu2+ ions get reduced.

Zn + Cu2+ Zn2+ + Cu

In the reaction, zinc supplies electrons to Cu2+ ions. Hence, zinc is the reducing agent. Cu2+ ions accept electrons from zinc. Hence, the Cu2+ ions is the oxidising agent.

In a redox reaction, the reducing agent undergoes oxidation, while the oxidising agent undergoes reduction.

Zn + Cu2+ Zn2+ + Cu

oxidation

reduction

reducing

agent oxidising

agent

Balancing Redox Equations using Half-reactions

This method divides the overall redox reaction into oxidation and reduction half-reactions.

Since neither oxidation nor reduction can actually occur without the other, we refer to the separate equations as half-reactions.

The general rule involves the following: 1. Divide the reaction into two half-reactions.

2. Balance each half reaction separately.

a) Balance atoms other than O & H in each half-reaction separately.

b) Balance O by adding H2O to the opposite side.

c) Balance H by adding H+ as appropriate.

d) Balance the charge by adding e-. For example, if the reactant side of the equation has a total charge of +3, the product side must also equal +3.

e) Balance the charges of the two half-reactions by multiplying appropriately.

3. Add two half-reactions together and balance the final equation by inspection.

4. Cancel the electrons on both sides.

Example 2:

Consider the following reaction:

2

2

4 ClMnClMnO

Reduction Oxidation

1)

2)

3)

4) 5)

1)

2) ..........

3) ..........

4)

5)

24 MnMnO 2ClCl

24 MnMnO

OHMnMnO 22

4 4

OHMnMnOH 22

4 48

OHMnMnOHe 22

4 485

OHMnMnOHe 22

4 8221610

22 ClCl

eClCl 22 2 eClCl 10510 2

Finally, combine half-reactions and cancel terms:

eClOHMn

ClMnOHe

10582

1021610

22

2

4

22

2

4 58210216 ClOHMnClMnOH

Exercise 1

Balance the following equations:

332272 FeCrFeOCr

2422 SOClSOCl

2242 SOCOSOHC

(a)

(b)

(c)

ELECTROCHEMICAL CELL

Oxidation-reduction or redox reactions take place in electrochemical cells.

There are several types of electrochemical cells:

1. Galvanic/voltaic/chemical cell/daniel cell

2. Electrolytic cell

3. Concentration cell

There are 4 essential components in a cell:

1. Anode

2. Cathode

3. Ionic conductor (electrolyte)

4. Metallic conductor (electrical connection)

Galvanic Cells

Spontaneous chemical reaction which generates electrical energy

Chemical energy Electrical energy

In the cell reaction the difference in chemical potential energy between higher energy reactants and lower energy products is converted into electrical energy.

A common galvanic cell is shown below:

Electrolytic Cells Electrical energy is used to bring about a non-

spontaneous reaction.

Electrical energy Chemical energy

This type of cell is formed when an external source of electrical energy is introduced into the system.

In the cell reaction electrical energy from an external power supply converts lower energy reactants into higher energy products.

A common electrolytic cell is shown below:

Similarities between Galvanic Cell and Electrolytic Cell

1. Both involve redox reactions.

2. Anode is the site of oxidation.

3. Cathode is the site of reduction.

4. Involve the flow of electrons from the

Anode to the Cathode.

5. Both can have a salt bridge (for the passage of ions so as to maintain electrical neutrality).

Differences between Electrolytic Cell and Galvanic Cell

Galvanic cell Electrolytic cell

1. Chemical energy Electrical energy

1. Electrical energy Chemical energy

2. Spontaneous reaction 2. Non-spontaneous reaction

3. Positive terminal of a cell is cathode

Negative terminal of a cell is anode

3. Positive terminal of a cell is anode

Negative terminal of a cell is cathode

4. G < 0 4. G > 0

Cell Notation

Cell notation in chemistry is a shorthand way of expressing a certain reaction in an electrochemical cell.

Components of

anode compartment (oxidation half-cell)

Components of cathode compartment

(reduction half-cell)

LEFT SIDE RIGHT SIDE

Cell Notation

CuCuZnZn 22

Electrode (s)

Electrolyte (aq)

Electrolyte (aq)

Electrode (s)

Reduction Oxidation

Phase boundary (aqueous/solid)

Salt bridge or porous partition

Concentration Cells

A galvanic cell in which both compartments have the same material but at different concenration.

The electrons flow in the direction that tends to equalize concentrations.

Concentration Cell

Diluted solution Concentrated solution

[Zn2+] in the two compartments are different, 1.0 M and 0.1 M, respectively.

The cell will try to equalize the [Zn2+] in the two compartments by transfering electrons from:

Compartment containing 0.1 M Zn2+ to the one containing 1.0 M Zn2+ (left to right).

The electron transfer will produce more Zn2+ in the left compartment and consume Zn2+ (to form Zn) in the right compartment.

CELL POTENTIAL

Electron flows from the anode to the cathode because there is a difference in electrical potential energy between the electrodes.

The difference in electrical potential between the anode and the cathode is measured by a voltmeter.

The voltage across the electrodes of a cell is called the cell voltage or cell potential or electromotive force (emf) (E)

STANDARD ELECTRODE POTENTIAL

According to the IUPAC convection, all half-cell reactions are written as REDUCTION.

The standard electrode potential is sometimes called as standard reduction potential or standard redox potential.

The electrode potential measured under standard-state conditions is called standard electrode potential, E measured in volts.

Conditions for the measurement of the electrode potential have to be standardised.

The standard-state conditions are:

Temperature (T) fixed at 25C or 298 K

Pressure (P) fixed at 1 atm or 101 kPa

Concentration of aqueous ions fixed at 1.0 M

The absolute value of the standard electrode potential of an electrode system cannot be measured.

The value can be determined if a particular electrode is chosen as the standard or reference electrode and all other electrode systems are measured against this standard electrode.

The standard hydrogen electrode (SHE) was chosen by IUPAC as the reference electrode for all measurements of E.

The standard reference half-cell is a standard hygrogen electrode, which consists of a specially prepared platinum electrode immersed in a 1 M aqueous solution of a strong acid, H+ (aq) [or H3O

+ (aq) ], through which H2 gas at 1 atm is bubbled.

)1;(2)1;(2 2 atmgHeMaqH Ereference = 0.00 V

E for the hydrogen electrode is arbitrarily fixed at 0.00 V.

Tables of Standard Electrode Potentials for Half-Reactions allow us to determine the voltage of electrochemical cells.

These tables compare the ability of different half-reactions to compete for electrons (become reduced).

Since the values are given in their ability to be reduced, the bigger the E, the easier they are to be reduced.

Selected Standard Electrode Potentials (298K)

Half-Reaction Eo(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

streng

th o

f redu

cing

ag

ent

stre

ng

th o

f o

xid

izin

g a

gen

t

Example 3:

If copper and hydrogen half-cells are joined together, we find that the copper half-cell will gain electrons from the hydrogen half-cell:

saq CueCu

22 VE 34.0

aqaq HeH 222 VE 00.0

Since both half-reactions cannot undergo reduction, we must reverse the equation of the reaction that will undergo OXIDATION.

saq CueCu 22

eHH aqg 222

VE 34.0

VE 00.0

Reduction

Oxidation

Calculating Voltages of Electrochemical Cells

Before calculating the voltage of a cell you must first determine which half-cell will be oxidized and which one will be reduced.

anodecathodecell EEE

Example 4 Zinc copper electrochemical cell

From Table of Standard Electrode Potential:

In the Table, all reactions are written as reduction reactions.

The E values indicate which half-reaction is better at competing for electrons.

saq ZneZn 22

saq CueCu 22 VE 34.0

VE 76.0

Since the copper half-reaction has a larger value for E than the zinc half-reaction, copper will be reduced, forcing zinc to be oxidized.

So we reverse the zinc equation.

We can then add the two equations together to get the full redox reaction and determine the voltage of the cell.

eZnZn s 2

2

saq CueCu 22 V34.0

V76.0

ssaq CuZnZnCu aq 22 V10.1

A positive value of E indicates a

spontaneous reaction A negative value of E indicates a

non-spontaneous reaction

Exercise 2

Predict the spontaneity of the following reactions:

a) Reduction of Sn2+ to Sn by Mg

b) Oxidation of Cl- to Cl2 by acidified Cr2O72-

SPONTANEITY OF REDOX REACTIONS

In an electrochemical cell, chemical energy is converted to electrical energy.

Electrical energy in this case is the product of the emf of the cell and the total electrical charge (in coulombs) that passes through cell:

W = Q E

unit: W = Joules (J) Q = Coulombs (C) E = Volts (V)

The total charge is determined by the number of moles of electrons (n) that pass through the circuit.

By definition: total charge = nF (Q = nF) 1F = 96500 C/mol The measured emf is the max. voltage that the cell can

achieve. This value is used to calculate the maximum amount of electrical energy that can be obtained from the chemical reaction.

This energy is used to do electrical work (wele), so

wmax = wele = -nFEcell wmax = is the max. amount of work that can be done.

Gibbs Free Energy (G) The Gibbs free energy (G) represent the max. amount of useful work that can be obtained from a reaction. G is the negative value of the max. electrical work: G = Wmax = -nFEcell For reactions in which reactants and products are in their standard states,

G = -nFEcell

Spontaneous reaction: G = negative

Ecell = positive

Non-spontaneous reaction: G = positive

Ecell = negative

Applications of Thermodynamics to Corrosion

A definite relation between the free energy change and the cell potential of an electrochemical reaction.

G is used to predict the spontaneous direction of any electrochemical reaction.

The sign of G is the most important factor to indicate whether or not the reaction is spontaneous.

G = -nFEcell

In any electrochemical reaction, the most negative (active) half-cell tends to be oxidized and the most positive (noble) half-cell tends to be reduced.

From the rule:

All metals with potentials more negative (active) than hydrogen will tend to be corroded by acid solutions.

However Cu, Ag, Hg, Pt and Au are not corroded in acid solutions.

But if dissolved oxygen is present, there is a possibility of corrosion to occur for Cu and Ag due to oxygen reduction.

Cu and Ag tend to corrode spontaneously in the presence of oxygen.

Example 5

i) Cu + H2SO4 No reaction

(Cu/Cu2+ > positive than H2/H+)

ii) 2Cu + 2H2SO4 + O2 2CuSO4 + 2H2O

(O2/H2O > positive than Cu/Cu2+)

Exercise 3

1. Would it be possible to store a silver spoon in a zinc nitrate solution?

2. Would it be possible to store a silver nitrate solution in a copper container?

3. Calculate the standard free-energy change for the following reaction at 25C:

Is these reaction spontaneous?

)(

32

)( 3)1(2)1(32 ss CaMAuMCaAu

)(22)()(2

gaqsHZnClHClZn

Nernst Equation

Used when all or some of the components are under non-standard-state conditions.

Shows how the cell potential depends on the concentrations of the cell components.

QnF

RTEE cellcell ln

QnF

RTEE cellcell log

3.2

Qn

EE cellcell log0592.0

Where R = universal gas constant

(8.314472 J K-1 mol-1)

T = temperature (K)

F = Faraday constant (96500 C mol-1)

n = no. of electron

Q = reaction quotient

treac

productQ

tan

Cell Potential and the Nernst Equation

The Nernst equation can be used to calculate the cell potential of electrochemical cells which are not under standard conditions.

Consider the galvanic cell operate at 25C:

Zn (s) Zn2+ (1.8 M) Cu2+ (0.2 M) Cu (s)

Exercise 4

Predict whether the following reaction would proceed spontaneously as written at 298 K:

Co(s) + Fe2+

(aq) Co2+

(aq) + Fe(s)

given that

[Co2+] = 0.15 M and [Fe2+] = 0.68 M.

Equilibrium Constants

For a general cell reaction:

where a, b, c, and d are the stoichiometric coefficients for the reacting species A, B, C, and D.

For the reaction at a particular temperature:

where K is the equilibrium constant.

dDcCbBaA

ba

dc

BA

DCK

][][

][][

At equilibrium, there is no net transfer of electrons, so E = 0 and Q = K.

Applying these conditions to the Nernst equation (valid at 25C)

Therefore,

Qn

EE cellcell log0592.0

Kn

E cell log0592.0

0

0592.0

)(log

cellEnK

?K

Example 6

Calculate the equilibrium constant, K, for the following reaction:

AgSnAgSn 22 42

Relationship of Gibbs Free Energy and Nernst Equation

Consider a redox reaction of type:

The free energy change, G, is given by thermodynamic equation:

Where Q = reaction quotient

dDcCbBaA

QRTGG ln

ba

dc

BA

DC = [Product]

[Reactant]

Because G = -nFE and G = -nFE, the equation can be expressed as

Dividing the equation through by nF, we get

Using the base-10 logarithm of Q

QRTnFEnFE ln

QnF

RTEE ln (Nernst Equation)

Qn

EE log0592.0

Go

Eocell K

Go K Reaction at standard-state

conditions

Eocell

The interrelationship of Go, Eo, and K.

< 0 spontaneous

at equilibrium

nonspontaneous

0

> 0

> 0

0

< 0

> 1

1

< 1

Go = -RT lnK Go = -nFEocell

Eocell = RT lnK

nF

Kn

Ecell log0592.0

(at 298 K)

HOMEWORK

1. Calculate E and Ecell for the following cell reactions.

a.

b.

2. What is the emf of a cell consisting of a Pb2+/Pb half-cell and a Pt/H+/H2 half-cell if [Pb

2+] = 0.10 M, [H+] = 0.050 M and PH2 = 1.0 atm?

MSnMMgSnMgSnMg saqaqs

035.0,045.0 22

2)(

2

MZnMCrCrZnCrZn saqaqs

0085.0,010.0

2323

23

2)(

3

3. Calculate the emf of the following concentration cell:

READ on batteries topic

-dry cell battery

-mercury battery

-lead storage battery

-solid state lithium battery

ss MgMMgMMgMg 53.024.022

Solubility Product Constant When a sparingly soluble salt (a salt which has

very low solubility in water) is added a little at a time to water, a saturated solution is eventually formed.

That is, the solution is in equilibrium with excess undissolved solid.

When a saturated solution of silver chloride is in contact with solid silver chloride, the following equilibrium is established.

)()()( aqClaqAgsAgCl

The product of the concentration of silver ions and chloride ions is known as the solubility product (Ksp) of silver chloride.

The solubility product of a sparingly soluble salt is defined as the product of the concentration of the ions in a saturated solution of the salt, raised to the power of the stoichiometric coefficient of the respective ions.

]][[ ClAgKsp

Example 7

1.

2.

3.

)()()( 22 aqSaqNisNiS ]][[22 SNiKsp

)(2)()( 22 aqClaqPbsPbCl

22 ]][[ ClPbKsp

343

243 2)(3)()( POaqCasPOCa23

4

32 ][][ POCaKsp

For a sparingly soluble salt AxBy, the Ksp is given by:

The value of Ksp changes when there is a change in temperature.

Generally, the higher the value of Ksp, the higher the solubility.

)()()( aqyBaqxAsBA xyyx

yxxy

sp BAK ][][

Example 8

Write down the expression for the solubility product for the following salts.

a) Pb(OH)2

b) Ag2CrO4

c) Al(OH)3

d) MgF2

Example 9

Calculate the solubility products of the salts from the data below:

a) The solubility of silver chloride in water is 1.453 x 10-3 g dm-3.

b) The solubility of silver sulphate in water is 1.5 x 10-2 mol dm-3.

Solubility Product and Precipitation

When an aqueous solution of calcium nitrate is added to an aqueous solution of sodium sulphate, one of the products is calcium sulphate, a sparingly soluble salt.

Whether a precipitate of CaSO4 will be formed or not depends on the concentration of the Ca2+

ions and SO4- ions in the mixture.

)(2)()()()( 344223 aqNaNOsCaSOaqSONaaqNOCa

If,

[Ca2+][SO42-] > Ksp : precipitation will occur

[Ca2+][SO42-] < Ksp : no precipitation takes place

Example 10

25.0 cm3 of 1.0 x 10-5 mol dm-3 silver nitrate is added to 25.0 cm3 of 1.0 x 10-5 mol dm-3 sodium chloride.

Common Ion Effect Consider a saturated solution of silver chloride:

If a little sodium chloride is added to the above solution, according to Le Chateliers Principle, the equilibrium will shift to the left to get rid of some of the Cl- ions added.

Hence, the amount of solid silver chloride will increase.

In other words, the solubility of silver chloride is decreased by the addition of sodium chloride (which contains the Cl- ions).

This phenomenon is known as the common ion effect.

)()()( aqClaqAgsAgCl