UNIT II ELECTROCHEMISTRY Rev.Ed. 2013 -14 · PDF fileArrhenius theory & Debye- Huckel’s...

UNIT II ELECTROCHEMISTRY Rev.Ed. 2013-14

Engineering Chemistry Page 26

ELECTROCHEMISTRY

Syllabus: Concept of ionic mobilities – Applications of Kohlrausch’s law – Conductometric

titrations – Galvanic cells – Electrode potentials – Nernst equation – Electrochemical series –

Potentiometric titrations – Concentration cells – Ion selective electrode: Batteries and Fuel cells

Objectives: Knowledge of galvanic cells, electrode potentials, and concentration cells is necessary

for engineers to understand corrosion problem and its control; also this knowledge helps in

understanding modern bio-sensors, fuel cells and improve them.

OUTLINES

Introduction

Arrhenius theory & Debye- Huckel’s theory

Kohlrausch’s law of independent migration of ions

Applications of Kohlrausch’s law

Conductometric titrations

Galvanic cells – electrode potentials

Electrochemical series

Electromotive force (Nernst equation)

Potentiometric titrations

Concentration cells

Ion selective electrodes

Batteries

Fuel cells


Engineering Chemistry 7

1. INTRODUCTION

Electric current is the flow of electrons generated by a battery, when the circuit is completed. A

substance which allows electric current to pass through it is called a conductor, e.g., all metals,

graphite, fused salts, aqueous solutions of acids, bases and salts; while insulator or non-conductor is a

substance which does not conduct electric current i.e., which does not allow the passage of electric

current through it.

The conductors are of two types

a. Metallic conductors are the substances which conduct electricity, but are not decomposed by

it, e.g., all metals, graphite etc.

b. Electrolyte is a substance which in aqueous solution or in molten state liberates ions and

allows electric current to pass through, thereby resulting in its chemical decomposition, e.g.,

acids, bases and some salts.

1.1 IONIC MOBILITIES- Arrhenius theory

The main postulates of the Arrhenius theory are..

1. In solutions, all electrolytes are spontaneously dissociated, to some extent, into charged particles,

called ions. The ions carrying positive charge are called cations; while those carrying negative

charge are called anions.

2. Cations are generally, metallic radicals obtained by loss of electrons from the metal atoms; while

anions are non-metallic atoms or radicals (a group of atoms of two or more elements) obtained by

gain of electrons:

M Mn+

+ n e-

Metal cation electrons

A + n e- A

n-

Non-metal anion

3. The total positive charge on the cations present in a solution is equal, but opposite to the total

negative charge present on the anions. Thus, the solution of an electrolyte is neutral as a whole.

4. The cations and anions present in a solution are constantly reuniting to form un dissociated,

electrically neutral parent molecules and a state of dynamic equilibrium exists between the

ionized and the unionized molecules

AB

Unionized molecule

A+ + B-

cation anion

Thus, the process of electrolytic dissociation is a reversible process.

5. The ions are free to move under the influence of electric current, they are directed towards

oppositely charged electrodes. Cations move towards cathode, while anions move toward the

anode, when electricity is passed through the solution.

6. The properties of electrolytes in solution are the properties of the ions produced.



7. The electrolyte at a given dilution may not be completely ionized and the fraction of the total

molecules ionized, is termed as degree of ionization, i.e.

No. of molecules dissociated into ionsDegree of ionization = -------------------------------------------------- Total number molecules taken

Electrolytes like mineral acids, alkalis and many salts, which give solutions with high conductivities,

are called strong electrolytes. This is because strong electrolytes are more or less completely ionized,

when in solution, e.g.,

HCl↔ H+ + Cl

-

NaCl ↔ Na+ + Cl

-

KOH ↔ K+ + OH

-

In other words, their degree of ionization is 1 or 100% .

8. Electrolytes like many organic acids, bases, water and ammonium hydroxide, which are poor

conductors in solutions, are weak electrolytes. This is due to the fact that only a small fraction of

their molecules dissociate into ions in solutions. Their solutions contain smaller number of ions

in equilibrium with unionized molecules, e.g.,

H2O ↔ H+ + OH

-

CH3COOH ↔ H+ + CH3COO

–

9. The conductivity of an electrolytic solution generally increases with the increase in temperature,

because the average kinetic energy of the ions increases as temperature increases.

1.2 DEBYE-HUCKEL THEORY

The Arrhenius theory is found to be valid only in case of weak electrolytes. The failure of the

theory, in case of strong electrolytes, has been satisfactorily explained by Debye-Huckel and

Onsager (1923). According to them:

1. All strong electrolytes are completely ionized even in solid state, i.e., their degree of

dissociation is 1 or 100%. However, the ions are not free to move in solid state and, therefore,

cannot conduct electricity.

2. On melting or dissolution, the ions become mobile and mobility of these ions depends upon:

viscosity of the medium, the extent of solvation (the number of solvent molecules attached to

each ion).

3. The ratio of ʌeqc / ʌeq

∞ for strong electrolytes does not represent the degree of ionization, but it

is only a conductivity ratio.

4. The increase in conductivity (ʌeqc) of strong electrolyte solution on dilution is due to increase

in the mobility of ions.

5. The lower values of mobility of ions in concentrated solutions, in spite of complete

dissociation is due to the following ionic interferences:

ʌeqc = ʌeq

∞− (A + B ʌeq

∞) √



1.4 CONDUCTIVITY OF ELECTROLYTES

According to ohm’s law, the resistance of a conductor(R) is directly proportional to its length (L) and

inversely proportional to its cross sectional area(A)

R = ρ (L/A), where ρ is a constant of proportionality, termed as specific resistance

Thus, if L = 1cm and A = 1 cm2 R = ρ, i.e., resistivity or specific resistance of a conductor is the

resistance between two opposite faces of centimeter cube of that substance. The units of resistivity are

ohm-cm

a. Conductance : It is the tendency of a material to allow the flow of current through it. It is the

reciprocal of resistance.

Conductance (C) = 1/R

b. Specific conductivity (κ) is the reciprocal of specific resistance of an electrolyte solution i.e.,

κ(ka-pa) = (1/ρ) = (L/AR)

The usual unit of specific conductivity is ohm-1

cm-1

or S cm-1

If, L = 1cm and A = 1 cm2, then κ = 1/R = Specific conductance

Hence, specific conductivity (κ) is the conductivity of 1 cm3 solution

While observing conductance of solutions, electrodes have to be used and they may not have exact

surface area of 1cm2 and distance of 1 cm. Hence the reciprocal of resistance does not give specific

conductivity but a value proportional to it. Such measurements in solutions are difficult. Hence

indirect method is employed to calculate specific conductance from the observed conductance. The

ratio of length(L) and area(A) of electrodes in a conductivity cell is defined as cell constant(x)

Then x becomes the ratio of specific conductance and observed conductance.

Specific conductance κv = x X observed conductance

The units of cell constant are cm-1

. For the determination of cell constant, the specific conductance of

N/50 KCl solution at 250C is taken as 0.002765 ohm

-1.cm

-1. Such a solution is prepared by dissolving

0.372g of pure KCl in 250ml of conductivity water and its conductivity is observed. Then

Cell constant x = 0.002765/ observed conductance

Once cell constant is determined it can be used for all experiments provided care is taken to see that

the distances are not altered.

Problems on cell constant and conductance:

1. A conductance cell has two parallel electrodes of 1.25 cm2 area placed 10.5 cm apart; filled

with an electrolyte solution the resistance was 1995 Ω. Calculate the cell constant and specific

conductance.

Solution:



2. Equivalent conductance of 0.05 N NaOH is 240 mho. cm2. The electrode of the cell are 1 cm

apart with area 1 sq. cm. calculate the specific conductance.

Solution:

3. A cell whose resistance when filled with 0.1 N KCl is 192.3 ohm. It has 6306 ohm when

filled with 0.01N NaCl at 25 °C. Calculate the cell constant, specific conductance of NaCl

solution, equivalent conductance of NaCl solution and degree of ionization of NaCl.

Solution: Given Specific conductivity of 0.1 N KCl is 0.01289 at 25 °C and

equivalent conductance at infinite dilution of NaCl at 25 °C is 126.45

(i)

(ii) ( ( ))

( )

Ohm

-1cm

-1

( ) ( )

( )

Ohm

-1cm

2equiv

-1

(iii) Degree of ionization

λα(NaCl) = 126.45 Ohm-1

cm2equiv

-1

α(NaCl) =

7

Degree of ionization of NaCl = 0.31

c. Equivalent conductance of an electrolytic solution is defined as “the conductance of all the ions

present in one equivalent of the electrolyte in the solution at a given dilution”. If one equivalent of the

electrolyte is contained in V mL, then:

ʌeq = V x specific conductance of 1 cm3 solution = Vx κ

V = (1/N) L or (1000/N) ml (dilution)

ʌeq = (1000 κ /N)

The unit of equivalent conductance is ohm-1

cm2 eq

-1

d. Molar conductivity of a solution is defined as the conductance of all the ions present in one

mole of electrolyte in the solution. If M is the molar concentration in mol/L then:

ʌm = (1000 κ /M)

The unit of molar conductivity is ohm-1

cm2 mol

-1



e. Effect of dilution on conductance: As dilution increases more and more (i.e addition of water),

the ionization of the electrolyte increases and number of ions in any cubic volume decrease. As

specific conductance is conductance of the ions present in one cc of solution, its value decreases with

dilution. Since equivalent conductance and molecular conductance are products specific conductance

and volume of solution, they also tend to increase with dilution. Ionization increases with dilution at

attains a limiting value and further addition of water will not produce any further ionization.

Consequently this limiting value of equivalence conductance is at infinite dilution Λ∞.

2. KOHLRAUSCH’S LAW OF INDEPENDENT MIGRATION OF IONS

According to the law, at infinite dilution, when dissociation is altogether assumed to be complete

and all interionic effects vanish, each ion moves independently irrespective of the nature and presence

of the co-ion and contributes a definite value of its share to the total molar conductance of an

electrolyte, which depends only on its nature and not at all on the ion with which it is linked.

Thus, molar conductivity at infinite dilution of any electrolyte is equal to the sum of the molar

conductances of the cation and anion present, since each ion contributes a definite amount to

the total conductance of the electrolyte, i.e.,

ʌm∞

= V+ λ+∞ + V− λ−

∞

Where V+ and V

- are the number of cations and anions per formula of electrolyte;

λ+∞ and λ-

∞ are molar conductivities at infinite dilution of cation and anion respectively. So is the case

with equivalent conductance.

For example, for acetic acid (CH3COOH), V+

= V− = 1 so

λm∞ (CH3COOH) = λ

∞ (H

+) + λ

∞ (CH3COO

−)

And for MgCl2 V+ =1, V

− = 2

λm∞ (MgCl2) = λ+

∞ (Mg

+2) + 2λ−

∞ (Cl

−)

2.1 APPLICATIONS OF KOHLRAUSCH’S LAW:

2.1.1 Determination of molar conductivity of a weak electrolyte: As mentioned earlier, it is

not possible to determine the values of ʌm∞

of a weak electrolyte by extrapolation of molar

conductivity values at infinite dilution, because (A) even at very high dilutions, these

electrolytes are not completely ionized (B) the curve does not follow a straight line path.

Kolrausch’s law has provided a method for finding the ʌm∞

for weak electrolytes from ʌm∞

measurements of strong electrolytes. Suppose we want to compute the ʌm∞

value for acetic acid,

this value can be obtained from the ʌm∞

values of HCl, NaCl and CH3COONa, by Kolrausch’s

law we have

ʌm∞ (HCl) + ʌm

∞ (CH3COONa) − ʌm

∞ (NaCl)

= [λ∞

(H+) + λ

∞ (Cl

−) ] + [λ

∞ (CH3COO

-) + λ

∞ (Na

+)] − [λ

∞ (Na

+) + λ

∞ (Cl

−)]

= λ∞

(H+) + λ

∞ (CH3COO

−)



= λ∞

(CH3COOH)

Thus the values of ʌ∞

m for strong electrolytes can be obtained by extrapolation of ʌmc – √ graphs.

Thus ʌ∞

m value for any weak electrolyte can be obtained from the ʌm∞

values of the selected strong

electrolytes.

2.1.2 Determination of degree of dissociation: The term degree of dissociation is “the fraction

of the total number of molecules dissociated into ions” i.e.,

αc =

We also know that the conductivity of a solution is due to the presence of ions in solution,

greater the number of ions, greater is the conductivity.

ʌmc, the molar conductivity at a particular dilution

No. Of molecules dissociated into ions at this dilution ………… (i)

But at infinite dilution all molecules are in ionic form

ʌm∞

Total number of molecules taken ………… (ii)

Dividing (i) by (ii), we get

ʌmc /ʌm

∞

Thus, degree of dissociation at any dilution is the ratio of the molar conductivity at that

dilution to the molar conductivity at infinite dilution.

The value of ʌ cm can be obtained by direct measurement; while ʌm

∞ can be obtained with the

help of Kolrausch’s law, i.e.,

ʌm ∞

= V+ λ+∞ + V− λ−

∞

Hence, the values of αc can be calculated using the law.

2.1.3 Determination of solubility of sparingly soluble salts: For ordinary purpose, ionic salts

like AgCl, PbS, BaSO4, CaCO3, Fe(OH)3, etc., are regarded as insoluble, but they do have a very

small, but definite solubilities in water. The solubility of such a sparingly soluble salt is obtained by

determining the specific conductivity (k) of a saturated salt solution. Since only a very small amount

of salt is present in solution, the equivalent conductivity at such dilution can practically be taken as

ʌeq∞.

For every sparingly soluble salts: ʌeqc = ʌeq

∞ + kυ−

The value of ʌ∞

can be calculated with the help of Kolrausch’s law.

υ = ʌeqc = ʌeq

∞ = λ+

∞ + λ−

∞

k k k

But υ cm3 of saturated solution contains = 1 eq

1000 cm3 or 1 L solution contains = (1000/υ) eq



2.1.4. Calculation of ionic product of water (Kw)

Pure water is a poor conductor of electricity. But the conductivity and other measurements

support that water ionizes to a very small extent producing equal number of H+ and OH

- ions. The

ionization equilibrium is represented as

H2O H+ + OH

- Since H

+ is associated to water molecule,

2H2O H3O+ + OH

-

The equilibrium constant K is given by

2

2

-

3

O]H[

]OH[]O[H K

K. [H2O]2 = [H3O

+] [OH

-]

Since, the ionization of water is negligible, [H2O] can be taken as constant

K[H2O]2 = Kw = constant

Kw = [H3O+] [H2O] The constant Kw is known as ionic product of water at given

temperature.

“The ionic product of water, Kw at a given temperature, is defined as the product of the

concentrations of H+ and OH

- ions in water or in aqueous solution”.

At 25 oC (room temperature), Kw is 1.0 x 10

-14 moles

2/Lit

2 .

As the temperature increases, ionization of water increases and the value of Kw increases.

Kw = [H+][OH

-] = 1.008 x 10

-14 moles

2/Lit

2 at 22

oC.

Each water molecule dissociates to give equal number of opposite ions,

[H+] = [OH

-]

Kw = [H+][OH

-] = [H

+]

2 or [OH

-]

2 = 1.0 x 10

-14 M

2/Lit

2

[H+] = [OH

-] =

22-14 /LitMoles 10 x 1 = 1.0 x 10-7

moles/lit

Problems on Kohlrausch law:

1. Calculate the equivalent conductivity at infinite dilution for chloro acetic acid from the following

data. The λα of HCl, NaCl and ClCH2COONa are respectively 426,126 and 90mho cm-1 equiv-1.

Sol: λα of ClCH2COOH is to be calculated.

ClCH2COOH ClCH2COO- + H

+

λα → λa + λc

Given the values of λα of the three electrolytes which contains ClCH2COO- and H

+ are



H+Cl

-, Na

+Cl

-, ClCH2COO

-Na

+ = I + III - II = H

+Cl

- + ClCH2COO

- Na

+ - Na

+ Cl

- +

I II III = ClCH2COO-H

+

λa and λc are present in electrolyte I and II which is to be added and the excess Na+

and Cl- are to

subtracted.

Add the λα values of HCl and ClCH2COONa and subtract the ClCH2COOH value of NaCl. The λα of

ClCH2COONa is obtained.

= 426 + 90 – 126 = 390 mho cm-1

equiv-1

2. The specific conductance of saturated solution of silver chloride at 18oC is 1.24x10

-6

mho. The mobilities of Ag+ and Cl

- ions at this temperature are 53.8 and 65.3

respectively. Calculate the solubility of silver chloride in grams per liter.

Sol. Kv = 1.2x10-6

λα = λa + λc

λα = 53.8 + 65.3 = 119.1

λv = Kv + V

V = λα/Kv

10

6 ml

10

6 ml of contains 1 gm equivalent of the AgCl = (143.5 gms of AgCl)

1000ml of water should contain AgCl =

10

-6

Solubility of AgCl = 1.494 x 10-3

gms/ 1000cc.

3. Conductometric titrations

Conductometric titration is instrumental method of volumetric analysis based on the change in

conductance of the solution, at the equivalence point (or end point) during titration. This method is

based on the fact that conductance of an aqueous solution, containing an electrolyte, depends upon: (i)

the number of free ions in the solution and their nature; (ii) the charge on the free ions, (iii) mobility

(or speed) of the ions.

During the course of titration (i.e., addition of one electrolytic solution to that of another), the

number of free ions in the solution changes. Not only that, even the identity of the ions also changes.

As a result of this, conductance of the solution (contained in cell) also undergoes a change. Let us take

a few specific cases for illustration.

Acid- base titrations

(a) Titration of a strong acid ( HCl) with a strong base ( NaOH)

Before base is added, the conductivity of acid solution is high (mainly due to the presence of highly

mobile H+ ions). This is represented by point A on the curve. The acid is taken in the conductivity

cell and the alkali in the burette. On gradual addition of NaOH from the burette, highly mobile H+



ions of the acid are removed by the OH- ions added from the burette to form nearly non-

conducting water molecules.

H+ Cl- + Na+ OH- Na+ Cl- + H2O

Acid Base Water (non- conducting)Salt

Hence, the conductivity of the solution decreases progressively, till the equivalence point B is

reached. On further addition of NaOH, the conductivity of the solution will raise along the curve BC

(due to the addition and presence of highly mobile OH- ions to the solution since they are no more

neutralized). Thus, the descending branch of the curve (i.e., AB) gives the conductance of a mixture

of acid and salt , and the ascending branch of the curve no excess of either acid or base and hence,

the intersection corresponds to the equivalence point.

H+Cl + Na+ Cl-

A

B

C

Na+Cl- + Na+OH-

Equivalence point

Volume of strong base solution

Fig. 1. Coductometric titration of a strong acid and strong base

(b) Titration of a weak acid ( acetic acid) with a strong base (sodium hydroxide):

The titration curve has the form of A, B, C. Acetic acid (a weak acid) has low conductivity, as

represented by A. As NaOH is added, the poorly conducting acid is converted into highly ionized

salt, CH3COONa.

CH3COOH + NaOH CH3COO-Na+ + H2O (unionized)

Poorly conducting acid Highly ionized salt

And consequently, the conductivity

goes up along AB. When the acid is

neutralized, further addition of the

alkali causes a sharp rise in

conductance along BC (due to

addition of more conducting OH-

ions). The intersection of AB and

BC, therefore, represents the

equivalence point.

HAc + Na+Ac-

Na+Ac- + Na+ OH-

Equivalence point

Volume of strong base solution

AB

C

Fig. 2. Coductometric titration of a weak acid and strong base



(c) If strong acid ( HCl) is titrated against a weak base ( NH4OH): the conductance in the beginning

starts falling (due to the removal of H+ ions) to form practically unionized water plus slow moving

NH4+ ions

H+Cl- + NH4+OH-

Strong acid weak base NH4

+Cl- + H2O

Salt practically unionized

However, when the entire acid is neutralized, further addition of poorly ionized ammonium

hydroxide does not cause any appreciable change in the conductance. The shape of the curve thus

obtained and the intersection corresponding to the equivalence point are shown in the figure.

Volume of NH4OH solution

Equivalence point

H+Cl + NH4+ Cl-

NH4OH + NH4Cl-

Fig. 3. Coductometric titration of a strong acid and weak base

(d) If it is required to titrate a weak acid ( CH3COOH) against a weak base (NH4OH), the addition of

ammonium hydroxide may even cause a decrease in conductance in the beginning, because the

common ions formed depresses the dissociation of their respective electrolyte (i.e., CH3COOH,

NH4OH).

CH3COOH + NH4OH NH4+ + CH3COO

- + H2O

Weak acid weak base highly ionized salt practically unionized

However, on further addition of ammonium hydroxide, an increase in the conductivity of the

solution results, the conductance of the highly ionized salt (e.g., ammonium acetate) exceeds the

conductance of the weak acid (acetic acid) it replaces. After the neutralization of the acid, further

addition of poorly ionized ammonium hydroxide does not cause any appreciable change in the

conductance.

Volume of NH4OH solution

NH4+Ac- + NH4OH

Fig. 4. Coductometric titration of a weak acid against a weak base

Equivalence point



In a similar way redox titrations, precipitation titrations can also be performed and results are obtained

in a better way.

ADVANTAGES OF CONDUCTOMETRIC TITRATIONS

1. They give accurate end-points than the conventional visual titrations, where end point is

detected only after adding excess of the titrant. One titration is enough.

2. There is no need to use indicator here, since end point is determined graphically by

intersection of lines and many errors are also minimized due to graphical method.

3. They are useful in the titrations involving colored solutions, where the color change of

indicator is not clear for detection.

4. They are very useful in the titrations of very dilute solutions

5. They are useful for titrating weak acids against weak alkalis, which otherwise do not give

sharp endpoints in visual titrations

6. In this method, no keen observation may not be necessary near the end point, since it is

obtained graphically and hence possible errors are minimized.

4. Electrochemical Cell (Galvanic cell)

An electrochemical cell is a device in which a redox chemical reaction is utilized to get electrical

energy. An electrochemical cell is generally referred to as voltaic cell or galvanic cell. The

electrode where the oxidation occurs is called anode and the electrode where reduction occurs, is

called cathode. Example: Daniels Cell, Leclanche cell

The Daniel cell (Figure above) consists of zinc electrode dipped in ZnSO4 solution and copper

electrode, dipped in CuSO4 solution. The two solutions are separated by salt bridge so as to avoid

direct contact or mixing with each other. The electrode reactions in Daniel cell are

At anode: Zn Zn2+

+ 2e- (Oxidation)

At Cathode: Cu2+

+ 2e- Cu (Reduction)

Cell Reaction: Zn + Cu2+

Zn2+

+ Cu

Zn has more tendency to form Zn2+

and hence Zn metal acquires a negative charge; and Cu2+

has more

tendency to get deposited as Cu. Hence, copper electrode becomes positively charged. As a result, the



electrons via the external circuit constitutes the electric current in the opposite direction. The emf of

the cell is 1.1 volts.

4.1 Representation of a Galvanic Cell

A galvanic cell can be represented as follows;

a) Anode is written on the left hand side; while the cathode is written on the right hand side.

b) The electrode on the left (anode) is represented by writing the metal or solid phase first and the

electrolyte separated by a vertical line or semicolon

Zn(s) | Zn2+

(aq) or Zn(s); Zn2+

c) The cathode of the cell is written on the right hand side. In this case, the electrolyte is

represented first and the metal or solid phase, thereafter separated by a vertical line or semicolon.

Cu2+

| Cu(s) or Cu2+

; Cu(s)

d) A salt bridge is indicated by two vertical lines, separating the two half cells.

Thus, applying above considerations to Daniel Cell, we may represent it as

Zn(s) | Zn2+

(1M) || Cu2+

(1M) | Cu(s)

5. Electrochemical series: It is the series of the elements arranged in increasing order

of their standard electrode reduction potentials.

Metal ion Standard reduction potential in volts

Li + e- Li - 3.05

K+ + e

- K - 2.93

Ca2+

+ 2e- Ca

- 2.90

Na+ + e

- Na - 2.71

Mg2+

+ 2e- Mg - 2.37

Al3+

+ 3e- Al - 1.66

Zn2+

+ 2e- Zn - 0.76

Cr3+

+ 3e- Cr - 0.74

Fe2+

+ 2e- Fe - 0.44

Ni2+

+ 2e- Ni - 0.23

Sn2+

+ 2e- Sn - 0.14

Pb2+

+ 2e- Pb - 0.13

Fe3+

+ 3e- Fe - 0.04

Direction of electron (e-) flow

Anode Zn ZnSO4 (aq) || CuSO4 (aq) Cu Cathode (-)

Direction of electric current flow

(+)



H++ e

- ½ H 0.00

Cu2+

+ 2e- Cu + 0.34

Ag+ + e

- Ag + 0.80

Pt4+

+ 4e- Pt + 0.86

Au+ + e

- Au + 1.69

½ F2 + e- F

- + 2.87

The electro chemical series provide valuable information regarding:

5.1.1 Relative ease of oxidation or reduction

A system with high reduction potential has a great tendency to undergo reduction. For example, the

standard reduction potentials of F2/F- System and Li

+/Li System is + 2.87V and -3.05V respectively.

The former one can easily gain electrons than the later one. So F2 can easily be reduced to F- and Li is

easily oxidized to Li+.

5.1.2 Replacement Tendency

Metal with greater oxidation potential can displace metals with lower oxidation potentials from their

salt solution. For Example, Cu2+

has more tendency to replace Zn. Zinc will displace copper from

the solution of CuSO4.

5.1.3 Predicting spontaneity of redox reactions

Positive value of E of a cell reaction indicates that the reaction is spontaneous. If the value of E is

negative, the reaction is not feasible.

5.1.4 Calculation of Equilibrium Constant

The standard electrode potential

eqeq

0 KlognF

2.303RTKln

nF

RT E

Hence, 0.0592V

nE

2.303RT

E x nF K log

00

eq

6. ELECTROMOTIVE FORCE (EMF)

Electromotive force is the difference of potential produced by sources of electrical energy, which

can be used to drive current through electrical circuit. It is not a force but a scalar quantity expressed

in volts.

6.1 Electrode Potential

A metal (M) consists of metal ions (Mn+

), with the valence electrons that bind the metal atoms

together. If a metal is in contact with a solution of its own salt, the following two chemical reactions

will take place.



a) Positive metallic ions passing into solution

M Mn+

+ ne- (Oxidation)

b) Positive ions get deposited on the metal electrode

Mn+

+ ne- M (Reduction)

Zn2+

Zn2+

ZnRod

Zn2+ ions

+

+

+

+

+

+

+

+

+

++

+

Cu2+

Cu2+

CuRod

Cu2+ ions

+

+

+

+

+

+

+

+

+

++

+

CuSO4 SolutionZnSO4 Soution

Zn2+ ions

moving from

Zn electrode

to solution

leaving free

electrons

Cu2+ ions entering

the copper metal

leaving behind free

negativly charged

ions in solution

Fig. 6. Electrode potential

In the first case, ‘n’ electrons are left behind on the metal and it acquires negative charge. In

the second case, the metal acquires a positive charge. A dynamic equilibrium is ultimately

established in the reaction. At equilibrium, the potential difference between the metal and solution

attains a constant value, which is called the “Electrode potential”.

Thus, “Electrode potential of a metal is the measure of tendency of a metal to lose or gain electrons,

when it is in contact with a solution of its own salt of unit molar concentration at 25 oC”.

As a convention, the tendency of an electrode to lose electrons is termed as oxidation

potential (+X) and the tendency of an electrode to gain electrons as reduction potential (-X).

6.2 Redox Reaction

Oxidation involves loss of electrons and reduction involves gain of electrons those electrons.

In other words, oxidation and reduction must always go side by side and are not independent in an

electro chemical cell.

If we place zinc metal in a solution of copper sulphate, immediate deposition of Cu takes

place and metal zinc goes into solution as Zn+2

Zn(s) + Cu2+

(aq) Zn2+

(aq) + Cu(s)

In this change, the zinc atom (Zn) is oxidized to Zinc ion (Zn2+

), losing electrons; while the

copper ion (Cu2+

is reduced to copper atom, gaining those electrons.

)Oxidation(2e(aq) Zn Zn(S) 2

Cu(S)(aq) Zn (aq)Cu Zn(S)

Cu(S)2e(aq)Cu22

2 (Reduction)

(Redox)

The overall reaction is called redox or oxidation –reduction reaction. Each of these reactions

is known as a half-reaction. The reaction, in which loss of electrons takes place, is called oxidation-



half reaction; while the other reaction, in which gain of electrons takes place, is called reduction half

reaction.

Thus each electro chemical cell is made up of two electrodes which are dipped in their salt solutions

and connected through a salt bridge and give emf on connection in an external circuit. Each of these

electrodes will have its potential. An electrode dipped in its salt solution is a half cell and its potential

called single electrode potential, which cannot be measured directly.

6.3 EMF of an Electrochemical Cell

An electrochemical or galvanic cell is obtained by coupling two half cells. Mathematically,

the emf of an electrochemical cell is the algebraic sum of the single electrode potentials; provided

proper signs are given according to the actual reaction taking place on the electrodes.

Standard EMF of the cell = [standard oxidation potential at oxidation half cell reaction + Standard

reduction potential of reduction half cell reaction]

= [standard reduction potential at reduction half cell reaction - Standard

reduction potential of oxidation half cell reaction]

[oxidation potential = - reduction potential]

Ecell = Ecathod – Eanode

Ecell = Eright – Eleft ;

Where Ecell = e.m.f. of the cell

Eright = reduction potential of right hand side electrode

Eleft = reduction potential of left hand side electrode

The positive value of Ecell indicates that the cell reaction is feasible.

Problems on EMF

1. Calculate the e.m.f the following reaction at 25 oC

Cu++ + Zn Zn++ + Cu

Eo Zn (Ox) = 0.736 V

Eo Cu (Ox) = -0.337 V

E (Cell) = E(R) - E(L)

= 0.337-(-0.763)

= 0.337 + 0.763

= 1.1volt

Zn Zn2+aq (1M) Cu2+

aq (0.1M) CuThe potential of the cell at 25 oC is 1. volt. calulate the

potentilal of the cell Zn Zn2+aq (0.5M) Cu2+

aq (0.01M) Cu at the same temperature

2.



E(cell) = Eo(cell) +

0.0591

2log10

Cu 2+

Zn 2+

E(cell) = 1.1 v

= 1.1 +0.0591

2log10

0.1

0.5

= 1.1 + log100.0296 X 0.2

= 1.1 + 0.698970.0296 X

= 1.1 + 0.0297

= 1.1207 V

6.4 Measurement of Electrode Potential

It is not possible to know the

absolute value of a single electrode

potential. We can only measure the

difference in potential between two

electrodes potentiometrically, if we fix

arbitrarily the potential of any one of

electrode. For this purpose, the

potential of a standard hydrogen

electrode (SHE) or normal hydrogen

electrode (NHE) has been arbitrarily

fixed as zero.

The SHE or NHE contains a platinum foil electrode coated with platinum black (for better

adsorption of hydrogen gas) in contact with 1M HCl solution through which hydrogen gas is bubbled

at a constant rate of one atmosphere pressure.

This electrode is represented as..

Pt(s), H2 (g) (1 atm); H+(1 M)

The following reactions take place at this electrode

H+(aq) + e

- → ½ H2 (g) Reduction

½ H2 (g) → H+ (aq) + e

- Oxidation

Thus when both of these reactions can take place at the electrode, it is capable of acting as anode or

cathode depending on the other electrode coupled, to form the cell. The EMF of the cell is measured

with a potentiometer and since the single electrode potential value of the hydrogen electrode is

UNIT II

ELECTROCHEMISTRY Rev.Ed. 2013-14


arbitrarily taken as zero, the single electrode potential of the other electrode can be calculated, since

the EMF value determined is the algebraic sum of the values of single electrode potentials.

In order to measure the electrode potential ( for eg; Zn dipped in ZnSO4), this electrode is coupled

with SHE through a salt bridge. Zinc electrode acts as anode, so the electrode potential is oxidation

potential with a value + 0.76V and consequently, its reduction potential is -0.76V. The electrode

potential of copper electrode is +0.34V.

6.5 Expression for single electrode potential

Consider a general redox reaction

Mn+

(aq) + ne- M(S)

For a reversible reaction, the free energy change (G) and its equilibrium constant (K) are

inter-related.

G [Reactant]

[Product]ln RT G0

Where G0 is known as the standard free energy change.

In a reversible reaction, the electrical energy is produced at the expense of the free energy i.e.

- G = nFE and G0 = - nFE

0

Where E is the electrode potential; E0 is the standard electrode potential; F = Faraday (or

96,500 coulombs). Consequently,

n

0

M

Mln RT nFE - nFE -

unity)isMofionConcentrat(M

1ln RT nFE -

n

0

n0 Mln RT nFE - nFE -

n0 Mln

nF

RT E E

n0 Mlog

nF

RT 2.303 E

Electrical work done = Electrical energy produced

= Quantity of electricity flow x EMF

For every one mole electrons transferred in the cell reaction the

quantity of electricity that flows through the cell is one faraday

(1F= 96500) if n moles of electrons are transferred.

Then, Electrical work done = nF. E cell

UNIT II



This expression is known as Nernst’s equation for electrode Potential.

From the above equation, it is clear that

i) If concentration of solution [Mn+

] is increased, the electrode potential increases and vice versa.

ii) If temperature is increased, electrode potential is increased and vice-versa.

When the elements are arranged in increasing order of reduction potential, a series called

electro chemical series is obtained.

6.6 Nernst equation for a cell reaction

Consider the Daniel Cell Zn | ZnSO4 (aq) || CuSO4 (aq) | Cu

Reaction is Zn(S) + Cu2+

(aq) Zn2+

(aq) + Cu(S)

Ecell = E0cell

2

2

Cu

Znlog

nF

RT303.2

At 298K, the Nernst equation can be written as

2

2

0

Anode

0

CathodecellCu

Znlog

0592.0)E(E E

n

2

2

0

|Zn

0

Cu|Cu Zn

Culog

0592.0)E(E 22

nZn

7. Potentiometric titrations

Potentiometric titration is an important application of emf measurement. In this method, a cell

is constructed, in which at least one of the electrodes is reversible with respect to one of the ions

taking part in the titration reaction.

Theory: The potential of an electrode dipping in a solution of an electrolyte depends on upon

the concentration of active ions (i.e., which changes the electrode potential)

E = E0 + (RT/nF) log C

UNIT II



A small change in the active ion concentration in the solution changes the electrode potential

to the corresponding level. During the course of titration, the concentration of the active ion

decreases, thereby electrode potential of the indicator electrode decreases. Thus, measurement of

indicator electrode potential can serve as a means of detection of equivalence point of the titration

reaction. The potential of the indicator electrode is, usually measured potentiometrically by

connecting it to a reference electrode like saturated calomel electrode.

Detection of end point: The emf of a cell changes by the addition of small volumes of titrant, so the

concentration of reversible ion in contact with indicator electrode changes. The change in emf with

every small addition is recorded. The change of potential will be slow at first, but at equivalence the

point change will be sharp or quite sudden with a jump or rise i n potential. The values of potentials

are plotted against corresponding volume of titrant added. A curve (a) like the one shown in figure is

obtained. The end point corresponds to the point of inflexion, i.e., point where the slope of the curve

is maximum as shown. If the inflexion is not sharp alternatively, the change in emf with every small

addition of titrant is plotted against volume V to obtain a curve shown in (b). The maximum of the

curve b gives the end point.

V (Volume of titrant)

End point

V (Volume of titrant)

End point

(a) (b)

Fig. 11. (a) In potentiometeric titration, the point of inflation is the end-point,(b) plot of E/ V against volume (V) Maxima gives the more acurare end-point.

E/

V

Em

f of

c ell

(E

)

Maxima

There are three important types of potentiometric titrations, which are described below.

a. Acid base reactions: In acid base titrations, quinhydrone electrode is employed as the indicator

electrode. The reference electrode is, generally the saturated calomel electrode.

A definite volume of the given acid solution is taken in a 100ml beaker. To it a pinch of

quinhydrone is added and a stirrer and platinum electrode are placed in it. This electrode is then

connected to saturated calomel electrode through a potentiometer. On adding standard alkali solution

from the burette, the emf of the cell increases at first slowly, but at the end point the rate of change of

potential will be suddenly quite large with a possible jump or drop in potential. The end point of the

titration is then located by plotting ΔE/ΔV versus V as shown in figure (11b) and the volume of titrant

UNIT II



corresponding to the peak in the curve gives end point. The advantage of such a titration lies in the

fact that this method can be used for titration of coloured solutions.

b. Oxidation-reduction reactions: Titrations involving oxidizing agents such as potassium

dichromate, potassium permanganate etc and reducing agents like ferrous ammonium sulphate can be

followed potentiometrically by using platinum indicator electrode.. On adding potassium dichromate

from the burette, emf of the cell will increase first slowly, but at the equivalence point, there will be

sudden jump or drop in the potential since the change in the ratio of Fe+2

/Fe+3

ions concentration, is

quite rapid at the equivalence point.

E = E0 + (RT/nF) log (Fe

+2)/ (Fe

+3)

Advantages of potentiometric titrations:

1. Coloured solutions, where the use of indicator is not possible can be estimated by

potentiometric titrations.

2. Since no indicator is necessary, there is no problem with regard to the choice of indicators

based on pH value of the solutions.

3. Since end point is determined graphically, many errors in titration are minimized and single

titration is enough.

4. Polybasic acids can be titrated in steps corresponding to different steps of neutralization.

5. Dependence on colour and external indicators are avoided for redox titration by Potentiometric

titrations.

6. Solutions containing more than one halide can be analyzed in a single titration against silver

nitrate.

8. Concentration Cell

In a galvanic cell, electrical energy arises from the decrease in free energy

(-G) of the chemical reactions taking place in the cell.

In a concentration cell, there is no net chemical reaction. The electrical energy in a

concentration cell arises from the transfer of a substance from the solution of a higher concentration

(around one electrode) to solution of lower concentration (around the other electrode).

A concentration cell is made up of two half cells having identical electrodes, identical

electrolyte, except that the concentrations of the reactive ions at the two electrodes are different. The

two half cells may be joined by a salt bridge.

UNIT II



For example

Ag Ag +AgNO3 (C2)

(Con)

Salt bridge

of stand

NH4NO3

AgNO3 (C1)

Ag-Ag +or

(C2 C1)V

AgNO3 (C1M) AgNO3 (C2M)

C. 25at . C

Clog

n

0592.0 E 0

1

2Cell

The general equation for emf of such cell is given by 1

2Cell

C

Clog

nF

RT303.2 E

In this cell, the following reactions occur

At anode : M Mn+

(C1) + ne-

At cathode : Mn+

(C2) + ne- M

Cell reaction : Mn+

(C2) Mn+

(C1)

The emf so developed is due to the mere transference due to concentration gradient of metal

ions from the solution of higher concentration (C2) to the solution of lower concentration (C1).

9. Ion Selective Electrode: An electrode that responds to the activity of a particular ion is

called ion-selective (or) ion-sensitive electrode (ISE). ISE are prepared basing on the principle that

membrane potentials are developed due to concentration gradient.

Principle: Whenever two solutions of different concentrations are separated by a membrane, a

potential difference arises across the membrane due to the unequal distribution of ions in the

solutions. This potential difference is known as membrane potential. This membrane potential

difference will be measured by ion-selectometers with the help of ion selective electrodes.

There are six types of ISE based on the membranes used in the electrode:

1. Glass membrane electrode

2. Liquid membrane electrode

3. Solid membrane electrode

4. Gas – sensing electrode

5. Enzyme – based electrode

6. Bio catalytic membrane electrode

9.1 Glass membrane electrode:

The electrode has a thin specially made glass membrane, which is selective to various

univalent cations. The internal reference electrode enables electrical contact between the inner

surface of the membrane via the reference solution and ion electrometer.

UNIT II



Glass Electrode: It contains a bulb containing 0.1M HCl

and a silver wire coated with silver chloride which acts as

internal reference electrode. It is immersed in a solution

whose hydrogen ion concentration is to be determined.

The upper end of the glass electrode is sealed. A

potential develops between two surfaces of the membrane

and PD developed is proportional to the difference in pH

value. Hence, the potential difference between the outer

surface of the glass bulb the solution into which the glass

bulb immersed is varied and so, the overall potential

is governed by the hydrogen ion concentration of the test solution. To determine the pH of the

unknown solution a glass electrode is combined with SCE and placed in a solution under test the emf

of the total cell is measured and the pH is calculated.

Ag/AgCl

0.1M KCl

0.1M HCl

Glass memberane

Fig. 12. Glass memberane electrode

The glass used for construction of the glass electrode is a special glass of approximate composition of

SiO2 (72%), Na2O (22%) and CaO (6%). Such electrodes satisfactorily work over the pH range 0-10.

Glass electrode are easy for operation, not easily poisoned. The equilibrium is rapidly reached with

accurate results. However glass electrodes can be used in pH range 0-10 and at higher pH the glass

attacked by alkalis.

,

The active membrane is a single crystal of LaF3 doped with europium (II) to lower its

electrical resistance and facilitate ionic charge transport. The LaF3 crystal, seated into the end of a

rigid plastic tube, is in contact with the internal and external solutions. Typically the internal solution

is 0.1M each NaF and NaCl; the fluoride ion activity controls the potential of the inner surface of the

LaF3 membrane, and the chloride ion activity fixes the potential of the internal Ag/AgCl reference

electrode the electro chemical cell for this electrode.

Ag/AgCl(s), Cl-(0.1M), F

-(0.1M)/LaF3 crystal/test solubility //Reference electrode

9.2 Solid Membrane Electrode

In solid membrane electrodes, the glass

membrane is replaced by an ion conducting

membrane. Fluoride electrode is example which

is used to measure fluoride ions accurately

quickly and economically. It consists of single

crystal of LaF3 as membrane. It contains an

internal reference electrode and internal

fluoridestandard electrode and LaF3 ion

exchange.

Initial filling solution

Reference electrode

Solid state ionic conductor

Fig. 13. Solid Membrane Electrode

UNIT II



It obeys Nernst relation of the form E = constant + F

RT ln [F

-]int /[F

-]ext

Since [F]int is constant, the equation simplifies to E = constant + 0.0591 pH.

At 25 oC.

The fluoride electrode also responds to hydroxide ion concentration. Hence, the hydroxide ion

concentration is kept constant with buffer solutions. It contains of 0.25M acetic acid, 0.75M sodium

acetate, 1M sodium chloride and 1mg sodium citrate. Sodium citrate masks Al3+

and Fe3+

which

interfere by complexing Fluoride. The buffer controls the overall ionic strength as well as the pH.

This electrode is highly useful in environmental studies, fluoride and other ions.

9.3 Liquid Membrane Electrode

Liquid membrane or ion exchange electrodes are prepared using an organic liquid ion-

exchanger which is immiscible with water or with ion sensing material is dissolved in an organic

solvent which is immiscible with water. The solvent is placed in a tube sealed at the lower end by a

thin hydrophobic membrane such as cellulose acetate paper; aqueous solutions will not penetrate this

film.

Liquid membrane electrode consists of a double concentric tube arrangement in which the

inner tube contains the aqueous reference solution and internal reference electrode. The outer

compartment contains organic liquid ion exchanger reservoir which occupies the pores of a

hydrophobic membrane.

Calcium responsive electrode is

an example for the liquid

membrane electrode. It contains

calcium salts, bis (2-ethyl hexy1)

phosphoric acid (d2EHP) and

dissolved in straight chain

alcohols (or) di decyl hydrogen

phosphate dissolved in di-n-

octyle phenyl/phosphate.

Ag/AgCl ref. electrode

Internal aq. solution

Ion exchange reservoir

Porous memberane

Liquid internal exchange layer with in porous memberane

Fig. 14. Liquid Membrane Electrode ,

9.4 Enzyme based electrode

These electrodes make use of an enzyme to convert the substance to be determined into an ionic

product, which can itself be detected by a known ion selective electrode. A typical example is the

urea electrode in which the enzyme urease is employed to hydrolyse urea and the progress of the

reaction can be followed by means of a glass electrode, which is sensitive to ammonium ions. The

final concentration of ammonium ions determined can be related to the urea present.

UNIT II



NH2CONH2 + H2O + H+ 2NH4+ + CO2

urease

The enzyme is incorporated in a poly acrylamide gel, which is allowed to set on the bulb of

the glass electrode and may be held in position by a nylon gauze. Then the electrode is inserted in to

a solution containing urea. Ammonium ions produced, diffuse through the gel and cause a response by

the ammonium ion probe.

9.5 Gas Sensing Electrode

The construction of ammonia sensing glass electrode

is shown. Dissolved ammonia from the sample

diffuses through a gas permeable fluoro carbon

membrane until a reversible equilibrium is

established between the ammonia level of the

sample and internal filling solution. Hydroxide ions

are formed in the internal filling solution by the

reaction of ammonia with water.

Memberane

Sensing element

Internal filling solution

Ref.electrode

Fig. 15. Gas Sensing Electrode

NH3 + H2O NH4+ +OH

-

The hydroxide ion concentration level of the internal filling solution is measured by the internal

sensing element, which is directly proportional to the level of ammonia in the sample.

9.6. Bio-catalytic Membrane Electrode

In biocatalytic membrane electrodes, a bio-catalyst is immobilized at the surface of an

electro chemical sensor (membrane electrode). The membrane electrode may be ion-selective (glass,

solid (or) polymer) or gas sensing (NH3, CO2 (or) H2S) electrode. The biocatalysts may be an

enzyme, tissue, or bacteria.

Bio-catalytic electrode life time is

dependent on the stability of the bio-catalyst,

which in turn is dictated by (1) methods of the

bio-catalyst immobilization (2) pH of the

solution, (3) storage conditions and (4) presence

of activators (or) inhibitors. The average

reported life time for gas sensing based bio-

catalytic membrane electrodes is 28 days. The

range of life time is 0.4 to 240 days. The

electrodes have excellent selectivity

characteristics owing to bio-catalytic specificity.

Memberan

Bulk solution

Biocatalytic layer

Fig. 16. Biocatalytic Membrane Electrode

UNIT II



10. Batteries

A device that stores chemical energy for a release at a desired later stage as electricity is

known as cell and a series of such cells in a group is called a battery. They are commercial portable

source of electro-chemical cells which supply DC at a constant voltage. They convert chemical

energy of a selected reaction into electrical energy. Batteries can be any of the three types.

10.1. Classification

Primary Battery (or Primary cells): These are the cells in which the cell reaction is not

reversible. When the reactants have for the most part been converted into products, no more electric

current is produced and the battery becomes ‘dead’. They can be used as long as the reactants are

active and can not be recharged. Eg: Leclanche cell

Secondary Battery (or Secondary Cells): Here the cell reaction can be reversed by passing the

direct electric current in opposite direction. A secondary battery may be used through a large number

of cycles of discharging and charging. Eg; Lead storage cell

Flow battery/ Fuel cells: These are the cells in which materials (reactants, products, electrolytes)

pass through the battery, which is simply an electrochemical cell that converts chemical to electrical

energy. Energy can be withdrawn indefinitely as long as fuel supply is maintained. They do not store

energy and cannot be reused. Eg; Hydrogen –oxygen fuel cell

10.2 Dry or Leclanche Cell

This is a cell is without fluid component and hence is conveniently portable. The anode of the cell is

a zinc can, containing an electrolyte consisting of solid NH4Cl, ZnCl2 and MnO2, to which starch is

added to make it thick paste to prevent leakage on transport. A graphite rod acts as a cathode and is

immersed in the electrolyte in the centre of the cell as displayed in the figure below:

Metal cap (+ ve)

Insulating washer

collar to keep rod in postion

Zinc Cup (negative)

Mixture of mangnese (iv)oxide, graphite, ammonium chloride and zinc chloride

Graphite rod

Metal cover (negative)

Fig. 17. Oxidation- reduction potentiometric titration

UNIT II



The cathode reaction is quite complex.

At Cathode: 2MnO2 (s) + H2 + 2e- Mn2O3(s) + 2OH (aq)

NH4 (aq) + OH (aq) NH3 (g) + H2O (l)

ZnCl2(s) + 2NH3 (g) [Zn (NH3)2]Cl2(s)

2MnO2 + 2NH4Cl (aq) + 2e- [Zn(NH3)2]Cl2 (s) + Mn2O3 + H2O

At Anode: Zn(s) Zn2+

(aq) + 2e-

Net reaction: Zn(s) + 2NH4Cl (aq) + 2MnO2 (s) Mn2O3(s) + [Zn(NH3)2]Cl2 (s) + H2O

Since various reactions involved cannot be reversed by passing electricity back through the

cell, this cell is irreversible and for one time use only. It gives a voltage of about 1.5V, which can be

increased further to desired level by connecting such cells in series. Dry cell finds wide variety of

applications in flash-lights, transistor radios, calculators.

Disadvantages

1. When current is drawn rapidly from it, products build up on the electrodes, thereby causing a

drop in the voltage.

2. Since the electrolytic medium is acidic, zinc metal dissolves slowly, thereby cell rundown is

slow.

3. Since the reactants are in acidic medium, zinc undergoes corrosion at a faster rate and decreases

the life of the cell.

10.3 Alkaline Battery: This is an improved form over the dry cell, in which NH4Cl is replaced by

KOH as an electrolyte. Zinc in powdered form is mixed with KOH to get a Gel. Graphite anode rod

is surrounded by a paste containing MnO2. The outside body is made with zinc. The cell reactions

are

Anode: Zn(s) + 2OH-(aq) Zn(OH)2(S) + 2e

-

Cathode: 2MnO2(s) + H2O(l) + 2e- Mn2O3(s) + 2OH

-(aq)

Net reaction: Zn(S)+2MnO2(S) H2O(l) Zn(OH)2(S) + Mn2O3(S)

The main advantages of alkaline cell over dry cell are

i) Zinc does not dissolve as readily in a basic medium and hence is not corroded.

ii) Alkaline battery maintains a better voltage as the current is drawn from it

iii) The life of alkaline battery is longer than dry cell, since there is no corrosion of Zn.

Uses: These are used in camera exposure controls, calculators, watches etc.

10.4 Nickel – Cadmium Battery:

This consists of a cadmium anode and cathode composed of a paste of NiO (OH) (s).

The cell reaction is

UNIT II



Anode: Cd(s) + 2OH-(aq) Cd(OH)2+ 2e

-

Cathode: 2NiO(OH)(s) + 2H2O + 2e- 2Ni(OH)2(s) + 2OH

-(aq)

Net reaction: 2NiO(OH)(s) + Cd(s) + 2H2O(l) Cd(OH)2(s) +2Ni(OH)2(s)

The reaction can be readily reversed, because the reaction products, Ni(OH)2(S) and Cd(OH)2, adhere

to the electrode surface.

It is portable, rechargeable cell with a fairly long life >20 years with a high amperage and maintains a

constant voltage of 1.4 to 1.45 volts.

It can be left for long periods of time without any appreciable deterioration, since no gases are

produced during discharging or charging.

Uses: They are used in electronic calculators, electronic flash units, cordless electronic shavers,

transistors and other battery powered small tools, emergency power supply and military tools.

10.5 Mercury Battery

This is a tiny cell used for special medical applications and in advanced electronics. The anode is zinc

amalgam and strong alkaline paste containing KOH, Zn(OH)2 and HgO serves as a cathode. These are

enclosed in a steel case, where cathode and anode are separated by a paper driver, which allows the

migration of ions.

The cell reaction is

Anode : Zn amalgam (s) + 2OH-(aq) ZnO (s) + H2O (l) + 2e

-

Cathode: HgO (s) + H2O (l) + 2e- Hg(l) + 2OH

-(aq)

Net reaction: Zn amalgam (s) + HgO (s) ZnO (s) + Hg (l)

Since there is no change in the electrolyte composition during its operation and overall

reaction involves only solid substances, it provides nearly constant voltage of 1.34V for 95% of its

life time. It is more expensive than ordinary dry cell but has specific applications.

Uses: It gives excellent performance in heart pacemakers, hearing aids, light meters, digital watches

etc.

10.6 Fuel Cells: A fuel cell is an electro-chemical cell which converts the chemical energy of an

easily available fuel -oxidant system directly into electricity. The chemical energy is provided by the

fuel and the oxidant is stored outside the cell and maintains a continuous supply. The essential

process in a fuel cell is

Fuel + oxygen Oxidation products + electrical energy

Characteristics of a fuel cell

They do not store chemical energy. Reactants are to be supplied constantly, while products are

removed constantly. A fuel cell resembles an engine than a battery. The efficiency of the fuel

cell is about twice that of a conventional power plant for generating electricity. Fuel cell generators

UNIT II



are free from noise, vibration, heat transfer thermal pollution and other problems normally associated

with conventional power plants.

Despite these characteristics, the fuel cells are not yet commercialized in operation. A major

problem lies in the choice and availability of suitable catalysts (for electrodes) able to function

efficiently for long periods to time without deterioration and contamination.

10.7 H2-O2 Fuel cell: One of the simplest and most successful fuel cell is hydrogen–oxygen fuel cell,

which consists essentially of an electrolytic solution such as 25% KOH solution and two inert porous

electrodes. Hydrogen and oxygen gases are bubbled, through the anode and the cathode

compartment.

The following reactions take place.

Anode: 2H2 (g) + 4OH-

(aq) 4H2O(l) + 4e-

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

- (aq)

Net reaction: 2H2 (g) + O2 (g) 2H2O (l)

The standard emf of the cell is

E0 = E

0ox+E

0Red = 0.83 + 0.40 = 1.23V.

In actual practice, the emf of the cell is 0.8 to 1.0V.

The product discharged by the cell is water.

Usually, a large number of these cells are stacked

together in series to make a battery, called the fuel

cell battery or fuel battery.

The electrodes must meet the stringent requirements. They must

i) be good conductors

ii) be good electron sources or sinks

iii) not to be consumed or deteriorated by the electrolyte heat or electrode reactions.

iv) be excellent catalysts for the reactions that take place on their surface.

When hydrogen is used as the fuel, the electrodes are made of either graphite impregnated with finely

divided platinum, or a 75:25 alloy of Palladium and Silver or nickel. Electrolytes used are aqueous

KOH or H2SO4 or ion-exchange resin saturated with water. For low temperature operating fuel battery

(-54 oC to 72

oC), potassium thiocyanate dissolved in liquid ammonia is employed.

Applications: Hydrogen – Oxygen fuel cells are used as auxiliary energy source in space vehicles,

submarines and other military vehicles. It has light weight and the product out come is useful water.

The initial cost is high but the maintenance cost is low.

10.8 Other fuel cells

In addition to H2/O2 system, a number of other fuel cells are also developed.

i) Propane – oxygen fuel Cell

UNIT II



The half reactions are

Anode: C3H8 (g) + 6H2O (l) 3CO2 (g) + 2OH+ + 2Oe

-

Cathode: 5O2 (g) + 2OH+

(aq) + 2Oe- 10H2O (l)

Net reaction: C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (l)

ii) Methyl alcohol – oxygen fuel Cell

The half reactions are

Anode: CH3OH (l) + H2O (l) CO2 (g) + 6H+

(aq) + 6e-

Cathode: 2

3 O2 (g) + 6H+ + 6e

- 3H2O (l)

Net reaction: CH3OH (l) + 2

3 O2 (g) CO2 (g) + 2H2O (l)

* * *

UNIT II



Assignment Questions

1. (a) Define transport number and ionic mobilities

(b) State and explain Kohlrausch law

(c) Discuss the applications of Kohlaursch law.

2. (a) Define and explain Specific conductance, equivalent conductance and molar conductance.

(b) Show that the degree of dissociation of a weak acid is reciprocal of the square root its

concentration.

3. Discuss in detail about various Conductometric titrations.

4. What is a galvanic cell? Explain its construction, working and applications

5. Define electrode potential. Derive an expression for the determination of electrode potential.

6. What is electrochemical series? Derive an expression for Nernst Equation.

7. Discuss in detail about Potentiometric titrations.

8. What are ion selective electrodes? Explain the construction working and applications of (a) NHE

(b) calomel electrode (c) glass membrane electrode (d) Ag-AgCl electrode.

9. (a) What is a battery? Explain the classification of batteries

(b) Explain the construction and working of Leclanche cell and Ni-Cd battery

(c) Explain in detail about concentration cells

10. What is a fuel cell? Explain the construction and working of Hydrogen-Oxygen fuel cell and

Methanol – Oxygen fuel cell.

UNIT II ELECTROCHEMISTRY Rev.Ed. 2013 -14 · PDF fileArrhenius theory & Debye- Huckel’s...

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