Unit II Electrochemistry.

8/18/2019 Unit II Electrochemistry.

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NARASARAOPETA NGINEERING COLLEGE - NARASARAOPE ENGG.CHEM

Dr. P . V. Narayana Page 1

UNIT - II

El ec tr ochemistr y

Intr oduc tion

Electrochemistry is the branch of chemistry, which deals with chemical applications of

electricity. The passage of electricity through a substance is called electrical conductance.

Electrical conductance involves movement of electrons or ions. A substance which allowselectric current to pass through it, is called a conductor. Eg: all metals fused salts, acids and

alkalis.

The electrical conductors are of two types.

1. Metallic or Electronic conductors

2. Electrolytic conductors

1.Metallic or Electronic conductors:

Metallic conductors conduct electricity due to the movement of electrons from one end to

another end. In a solid, the electrical conduction involves the free movement of electrons in the

metallic lattice, without any movement of the lattice atom; this type pf conduction is called

metallic conduction. In metallic conductors, the electricity is carried by the electrons, the atomic

nuclei remaining stationary. These conductors are further sub classified in to three types.

A. Good conductor

B. Semi- conductor

C. Non- conductor or Insulator.

Good conductor: It is a substance, which conducts electricity fully and freely. Eg: Metals likeCopper, Aluminum, and Iron.

Semi-conductor: It is a substance, which partially conducts electricity. Eg: Silicon, Germanium.

Non-conductor or Insulators: It is a substance, which does not conduct electric current i.e.,

which does not allow the passage of current through it. Eg: Wood, Graphite.


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2.Electrolytic conductor: It is a substance, which in aqueous solution of in molten state

liberates ions and allows electric current to pass through. Electrolytic conductors are further sub

classified in to three types, depending upon the extent of dissociation at ordinary dilutions.

Strong Electrolytes: Strong electrolytes are completely dissociated into ions at all

concentrations. Eg. NaCl, HCl, NaOH.

Weak Electrolytes: Weak electrolytes dissociate only to a small extent even at very high

dilutions. Eg: CH3COOH, NH4OH.

Non-Electrolytes: Non-Electrolytes do not dissociate into ions even at low dilutions.

Eg: Glucose, Sugar.

S.No Metallic conductors Electrolytic conductors

1. It involves the flow of electrons in a

conductor.

It involves the movement of ions in a

solution.

2. It does not involve any transfer of matter. It involves transfer of electrolyte in the

forms of ions.

3. Generally metallic conduction shows an

increase in resistance as the temperature

is raised.

But the resistance of an electrolytic

solution decreases as the temperature is

raised.

4. No net chemical change takes place Chemical reactions take place at the two

electrodes.

Conductivity of electrolytes: Electrolytic conduction involves the transfer of electrons through

the migration of positive and negative ions towards the electrodes. In an ionic solution, the

cations and the anions are free to move and both can transport charge. When a current is passed

through the solution, the ions carry a current. The ability of the ions in solution to carry current is

conductivity. The conductivity of a solution depends on the number of anions and cations present

in it and also on how readily these ions can move. Like electronic conductors, electrolytes also

obey ohm’s law.

Ohm’s law: The resistance (R) of a conductor is directly proportional to its Length (l) and

inversely proportional to its cross sectional area (a)

R l


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R a

R

R =

= specific resistance

Thus, if length of the conductor is 1 cm and its cross sectional area is 1 cm2, so

R =

=

=

R =

Units: Resistance = R = Ohm

= specific resistance = Ohm cm

Conductance and Specific conductance:

Resistance and specific resistance are commonly used for metallic conductors, where the

atoms are static. In case of electrolytic solutions the electricity is virtually conducted by

constantly moving ions. So it was thought more meaningful to define another quantity called

conductance. The conductance of an electrolytic solution is defined as the reciprocal of its

resistance.

Conductance (C) =

=Ohm-1 or mho or Ω-1 or Siemens (S)

Specific Conductance: reciprocal of specific resistance is known as Specific conductance,

this type of conductivity is called specific conductivity.

κ =

(R =

= R

)

κ =

Specific conductivity is defined as the conductance of a one centimeter cube of substance or

solution. It is represented by the symbol (κ ) κ = kappa.

In CGS system the conductivity is expressed in S cm-1 or mho cm-1

Specific Conductivity κ =

. = mho. cm-1 = S cm-1

In SI system, meter being the fundamental unit of distance, the conductivity is expressed

as S m-1. The conversion factor of both CGS & SI system is,

1 S cm-1 = 100 S m-1 or 1 S m-1 = 0.01 S cm-1


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Equivalent Conductance:

Equivalence conductance is defined as the conductance of all ions present in one gram

equivalent of the substance or electrolyte in the solution at the given concentration.

Equivalent conductance is represented by Λ (Greek : Capital Lambda).

Λ =×

Units of Equivalent conductivity: Ohm-1 cm2 eq -1

Molar Conductance:

Molar conductance is defined as the conductance of all the ions produced by dissolving

mole of the electrolyte in the given solution.

µ =×

The units of molar conductivity: Ohm-1 cm2 mol-1 or S cm2 mol-1

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 decreased. 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∞.

Ionic Conduc tivity and Ionic Mobil it ies

Equivalent conductivity of any electrolyte at any dilution is directly proportional to the

charge carried by the ions and their velocities. The conductivity is thus given by the products of

charge and velocity of individual ions. At infinite dilution the ionization is complete and the

solution containing one equivalent of various electrolytes contains equivalent number of ions.Hence at infinite dilution total charge carried by all ions is same in every case. Because the total

charge is constant at infinite dilution, the Λ must depend exclusively on ionic velocities.

Defining the ionic velocity or mobility as the speed with which a charged a particle at infinite

dilution moves under a potential gradient of one volt per cm, we have,


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u+ u-

Or

= ku+ = ku-

Where k is a proportionality constant and u+ and u- represent the ionic velocities at

infinite dilution. Since 1 equivalent of an ion under unit potential gradient carries a charge of 1

Faraday per sec., the proportionality constant k = 96500 coulombs.

Therefore, u+ = /k and u- = /k

Ionic conductivity is expressed in S.cm2, while ionic mobility is expressed in cm.s-1.

Kohl r ausch Law :

In a state of infinite dilution, all electrolytes ionize or dissociate cent percentage and the

ions can be carried by the ions independent of each other. During our discussion on the variation

of equivalent conductivity with square root of concentration we noticed that equivalent

conductivity increased with dilution to reach a limiting value characteristic of the electrolyte and

named it as equivalent conductivity at infinite dilution (Λ∞). In 1875 Kohlrausch made a series

of measurements, involving electrolytes with common cations, common anions and calculated

their differences. He obtain a constant difference in Λ∞ values of an ion pair, irrespective of the

nature of the common counter ion employed, as exemplified by following results obtained at

250C.

Basing on similar observations, kohlrausch advised a hypothesis called law of

independent migration of ions. This law states that at infinite dilution each ion makes a definite

contribution towards the equivalent conductivity of an electrolyte, irrespective of nature of the

co-ion with which it is associated in the solution.

Therefore equivalent conductivity at infinite dilution is made up of two independent terms,

called ion conductivities, characteristic of each ionic constituent of an electrolyte in solution. We

can express the same conclusion mathematically as

Λ∞ = +∞ + -∞

Where ∞+and -∞ are the contributions of cation and anion (or) ionic conductance of anion

respectively.


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Appl ications of Kohl r ausch l aw

1. Determination of degree of dissociation.

2.Determination of the solubility of sparingly soluble salt.

3.Calculate the ionic product of water.

1. Determination of degree of dissociation constant:

Degree of dissociation constant () is the fraction of the total number of molecules ionized

into ions.

=No.of molecules ionized into ions

Total number of molecules taken

= = .

Λc, can be obtained from conductivity measurements, Λ∞ can calculated using

Kohlrausch’s law.

Thus degree of dissociation (), is the ratio of the equivalent conductivity at particular

dilution to the equivalent conductivity at infinite dilution.

2. Determination of the solubility of sparingly soluble salt:

The solubility of the sparingly soluble salts like Gal, BaSO4 can calculate by using the

following relation.

Λ =×

C represents the solubility in g. eq. L –1

The equivalent conductance at infinite dilution is

Λ∞= κ v×V

The solubility ‘S’ can be calculated as

S= ×

g/ L

‘E’ is the equivalent weight of the substance

S =××

(

Thus, knowing κ v, the solubility of sparingly soluble salts can be calculated


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3.Ionic Product of Water

Water is a weak electrolyte. It dissociated slightly only

H2O H++ OH-

The dissociation equilibrium of water can be considered as,

According to law of mass action,

K eq =[H +][OH- ]/[H2O]

Water dissociates slightly only so the undissociated water can be taken as constant

K =[H +][OH- ]

Since water as a solvent is always in excess and change in concentration due its

dissociation is negligible. Hence water concentration is assumed to be constant.

K eq = [H+] [OH – ] = K w

The ionic product of water (K w) is defined as the product of concentrations of hydrogen

(H+) and hydroxide (OH – ) ions at 298 K.

The specific conductance of purest form of water at 250C is found to be 5.54×10-8 Ω-1

cm-1. The conductance of 1 litre water will be 5.54×10-8×1000.

Thus, [H+] =.×

. = 1.01×10-7

[OH-] =.×

. = 1.01×10-7

K w= [H+][OH-] = (1.01×10-7)(1.01×10-7)

= 1.02×10-14

Conduc to metr ic Titr ations

“A titration in which electrical conductance of a solution is measured during the course

of the titration”.

Conductometric titration is a type of titration in which the electrolytic conductivity of the

reaction mixture is continuously monitored as one reactant is added. The equivalence point is the

point at which the conductivity undergoes a sudden change.

The principle of conductometric titration is based on the fact that during the titration, one

of the ions is replaced by the other and invariably these two ions differ in the ionic conductivity

with the result that conductivity of the solution varies during the course of titration. The


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equivalence point may be located graphically by plotting the change in conductance as a function

of the volume of titrant added.

Acid-Base Titrations

Strong Acid Vs Strong Base: If HCl is considered as strong

acid and NaOH is considered as strong base, at the beginning of

the titration, the conductance is high due to the presence of

highly mobile hydrogen ions. When the base is added, the

conductance falls due to the replacement of hydrogen ions by

the added OH- ions to form undissociated water. The decrease in

the conductance continues till the equivalence point. At the

equivalence point, the solution contains only NaCl. After the equivalence point, the conductance

gradually increases due to the high mobile OH- ions.

Strong Acid with a Weak Base : If strong acid ( HCl) is titrated against a weak base (NH4OH)

the conductance at the beginning starts falling (due to the removal of

H+ ions) to form practically unionized water and slow moving NH4+

ions.

HCl + NH4OH NH4Cl + H2O

However, when the entire acid is neutralized, further addition of

poorly ionized ammonium hydroxide does not cause any appreciablechange in the conductance. The shape of the curve thus obtained and

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

Weak Acid with a Strong Base: Consider CH3COOH as weak acid and NaOH as strong base

Initially the conductance is low due to the feeble ionization of acetic acid. On the addition of

base, there is decrease in conductance not only due to the

replacement of H+ by Na+ but also suppresses the dissociation of

acetic acid due to common ion acetate. But very soon, the

conductance increases on adding NaOH as NaOH neutralizes the

un-dissociated CH3COOH to CH3COONa which is the strong

electrolyte. This increase in conductance continues raise up to the


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equivalence point. The graph near the equivalence point is curved due the hydrolysis of salt

CH3COONa. Beyond the equivalence point, conductance increases more rapidly with the

addition of NaOH due to the highly conducting OH- ions.

Weak Acid with a Weak Base: 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.

Advantages of Conductometric Titrations

1. No special care is necessary near the end point as the end point is ascertained graphically.

2. Colored solution which cannot be titrated by ordinary volumetric methods with the help

of indicators can be titrated.

3. The titrations of weak acid against weak bases can be performed; they cannot be

measured in volumetric titration because they do not produce sharp color change of

indicators.

4. Very dilute solutions can be titrated.

5. The method is very accurate in dilute solutions.

6.

It can even be used in colored or turbid solutions.


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Dr. P . V. Narayana P a ge 1 0

El ec tr o-c hemic al Cel l s

Electro-chemical cells are entirely different from electrolytic cells. In electrolytic cell

electrical energy is converted to chemical energy. But in electro chemical cells or voltaic cells or

galvanic cells chemical energy is converted to electrical energy. Daniel cell is an example for the

galvanic cell.

It consists of Zinc electrode dipped in ZnSO4 solution acts as oxidation half cell where

oxidation takes place. While copper electrode dipped in CuSO4 solution acts as reduction half

cell where the reduction takes place. These two solutions are separated by a salt bridge (U-tube

containing concentrated solution of KCl or NH4 NO3 in an agar-agar gel). Or porous pot, it

provides an electrical contact between two solutions.

The following chemical reactions take placeAt Anode: Zn Zn2+ + 2e- (Oxidation half Reaction)

At Cathode: Cu+2 + 2e- Cu (Reduction half reaction)

Net cell reaction Zn + Cu2+ Zn2+ +Cu

Representation of Galvanic cells or cell diagram:

It consists of two electrodes anode & cathode. It is represented by keeping in view the following

points.

1. Anode must be written on the left hand side, cathode on the right hand side.

2. Left hand side electrode is written by written by writing the metal first and then

electrolyte. These two are separated by vertical line or a semicolon. The electrolyte may be

written by the formulae of the compounds or by ionic species.

Zn/Zn2+

(or) Zn; Zn2+

(or) Zn; ZnSO4


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3. Cathode is written on the right hand side. In this first case, electrolyte is represented then

only electrode; these two are separated by a vertical line or a semicolon.

Cu2+

/ Cu (or) Cu2+

; Cu (or) CuSO4 / Cu

4. A salt bridge is denoted by two vertical lines, which separates the two half cells. Eg:

Daniel cells

Zn; Zn2+

|| Cu2+

; Cu

EMF of Electrochemical cell:

The difference of potential, which causes flow of electrons from the electrode of higher

potential to the electrode of lower potential, is called electromotive force (EMF). The emf of

galvanic cell is calculated from the reduction half-cell potentials using the following relation.

E(cell) = E(right) – E(left)

Where E(cell) = e.m.f of cell

E(right) = reduction potential of right hand side electrode

E(left) = reduction potential of left hand side electrode

Applications of EMF measurement:

5. Potentiometer titrations can be carried out.

6. Transport number of the ions can be determined.

7. Measurement of pH using hydrogen, quinhydrone and glass electrode.

8. From the EMF data the free energy changes, equilibrium constant of a reaction can be

found out.

9. Hydrolysis constant can be determined.

10. Solubility of a sparingly soluble salt can be found out.

El ec tr ode potent ial

Potential difference is developed between the metal ions from metal to the solution (or)

from solution to the metal. At equilibrium the potential difference remains constant, this is

known as electrode potential of metal. It is represented as E.

or

It is a measure of the tendency of the metal electrode to lose or gain electrons, when it is

in contact with of its own ionic solution.


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The electrode potential of a metal is defined as the direct measure of its tendency to get

reduced is called reduction potential, its value is +x Volts. Similarly the tendency of an electrode

to lose electrons is a measure of its tendency to get oxidized is called oxidation potential, its

value is –x Volts.

Standard electrode potential (Eº); it is the potential measured under standard conditions i.e.,

When,

1. Concentration of the metal ion solution is 1M.

2. Temperature is 298 K.

3. Pressure is 1 atm.

The standard electrode potential is represented as “Eº”

E E E anodecathodecell000

Ner nst Equation

In electrochemistry, the Nernst equation is an equation that can be used to determine the

equilibrium reduction potential of a half-cell in an electrochemical cell. It can also be used to

determine the total voltage (electromotive force) for a full electrochemical cell. It is named after

the German physical chemist who first formulated it, Walther Nernst.Consider the following redox reaction

Mn+ + ne- ↔ M

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

(K) are related as;

G = -RT ln k + RT ln[]

[]

G0 + RT ln[]

[]……………Eq- (1)

Where G0 = standard free energy change

The above equation is known as Van’t Hoff Isotherm.

The decrease in free energy in the reversible reaction will produce electrical energy i.e.

-G = nEF and

G0 = -nE0F ……………Eq-(2)


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Where E = Electrode potential

E0 = Standard electrode potential

F = Faraday (96,500 coulombs)

Comparing equation 1 & 2

-nEF = -nE0F + RT ln

[]

= -nE0F +

ln

[]

Where, concentration of the metal is unity or

-nEF = -nE0F -

ln [Mn+]

Dividing the equation by –nF

E= E0 +

ln [Mn+]

E= E0 +.

log [Mn+]

E= E0 +.

log [Mn+] ……………Eq-(3)

This equation-3 is known as “ Nernst Equation” for electrode potential.

El ec tr oc hemic al Ser ies

The electrochemical series is the arrangement of various electrode systems in the

increasing order of their standard reduction potentials. The electrochemical series consisting ofsome electrode systems along with their half cell reactions is give below.

Important features of electrochemical series

1. The electrode systems having negative values of standard reduction potentials act as

anode when connected to a standard hydrogen electrode whiles those having positive

values as cathode.

2. The metal placed at the bottom (Li) of the series has the minimum value of standard

reduction potential i.e. it has the minimum tendency to get reduced or maximum

tendency to get oxidized.

3. The substance placed at the top of the series (F) possesses the highest value of standard

reduction potential i.e. it has the strongest tendency to get reduced. Consequently, it acts

as the strongest oxidizing agent. The oxidizing power of the substances decreases in

going from bottom to the top of the series.


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Electrochemical Series

Conc entr ation Cel l s

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 higherconcentration (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.


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Types of concentration cells:

The concentration differences could be affected in the electrode material or in the

electrolyte. Further they could be sub divide into cells with transference or without transference.

Concentration cells mainly divided in to two types, they are

1. Concentration cells without transference

2. Concentration cells with transference

1.Concentration cells without transference: In these cells, the potential difference is developed

between two electrodes at different concentrations dipped in the same solution of the electrolyte.

For example, two hydrogen electrodes at different gas pressure in the same solution of hydrogen

ions constitute a cell of this type.

Pt, H2 ( P1) | H+ (a) | H2 (P2), Pt

If oxidation occurs at L.H.S electrode and reduction occurs at R.H.S electrode.

The Nernst equation can be used to derive an expression for the potential of this electrode

concentration cell.

H2 (P1) = 2H+ + 2e-

2H+ + 2e- = H2 (P2)

------------------------

H2 (P1) = H2 (P2)

------------------------

Ecell = .

log

at 250C

The standard cell potential is zero because a cell cannot derive a current through a circuit with

identical electrodes. Similarly we can construct different gas concentration cells, when coupled

with their counter ions.

2.Concentration cells with transference: In these cells, electrodes are identical but these are

immersed in solutions of the same electrolyte of different concentrations. The source of electrical

energy in the cell is the tendency of the electrolyte of diffuse from a solution of higher

concentration to that or lower concentration. With the expiry of time, the two concentrations tend

to become equal. Thus, at the start the emf of the cell is maximum and it gradually falls to zero.

Such a cells are represented in the following manner.


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M | Mn+ (c1) || Mn+ (c2) | M (C2 is greater than C1)

Zn | Zn+2 (c1) || Zn+2 (c2) | Zn (C2 is greater than C1)

The emf of the cell is given by the following expression:

At Anode : MMn+(C1)+ne-

At Cathode: Mn+(C2)+ne- M

Over all 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).

The general equation for emf of such cell is given

Ecell = .

log

at 250C

The above examples are typical examples of concentration cell with transference.

Transference indicates the presence of salt bridge or liquid-liquid contact of electrolytes.

Applications of concentration cells:

The concentration cells are used to determine the solubility of sparingly soluble salts, valency of

the cation of the electrolyte and transition point of the two allotropic forms of a metal used as

electrodes, etc.

Potent iometr ic Titr ations

The titrations in which end point or equivalent point is determined with the measurement

of electrode potentials of the reaction mixture.

Potentiometric titration is a very interesting application of electrode potentials. They

involve a study of variation of emf with the volume of titrant added.

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


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

curve b gives the end point.

(a) (b)Fig. (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.

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 100 mL beaker and platinum electrode is 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 and the volume of titrant 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.

Redox 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.


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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.

Solutions containing more than one halide can be analyzed in a single titration

against silver nitrate.

Standar d el ec tr odes (or ) Refer enc e El ec tr odes

The electrode of standard potential, with which we can compare the potentials of another

electrode, is called a “Reference Electrode”. The best “Primary reference electrode” used is

standard hydrogen electrode, whose electrode potential is taken as zero. But it is not always

convenient to use this gas electrode in day to day potentiometric measurements. The main

difficulty with this primary reference electrode isa) Inconvenience in handling gases

b) Maintenance of accurate pressures throughout the measurements

c) Availability of sufficiently pure hydrogen gas

d) Necessity to platinize the platinum electrode with a solution of chloro platonic acid quite

often.

For these reasons some secondary reference electrodes like calomel electrode and silver

chloride electrode whose standard potentials are very accurately determined against hydrogen

electrode are widely used.

Standard Hydrogen Electrode (SHE)

For determination of half-reaction current flows and voltages, we use the standard hydrogen

electrode. The figure below illustrates the standard hydrogen electrode. A platinum wire


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conducts the electricity through the circuit. The wire is immersed in a 1.0 M strong acid solution

and H2 gas is bubbled in at a pressure of one atmosphere and a temperature of 25°C.

The half-reaction at this electrode is

.

Under these conditions, the potential for the hydrogen reduction is defined as exactly zero.

We call this , the standard reduction potential. This electrode is used to measure the

potentials of other electrodes in the half-cell. A metal and one of its salts (sulfate is often used) is

in the second half-cell.

Standard Calomel Electrode – SCE

Calomel electrode is particularly very simple to construct, free from surface sensitivity

and accurate to use even in a very normal laboratory.

The calomel electrode consists of an inner glass tube and an outer jacket. In the inner

glass tube a platinum wire is dipped into mercury which rests on a paste of mercurous chloride,

Hg2Cl2 (commercially known as calomel) and mercury. This paste is in contact with KCl present


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in the outer jacket, through the glass frit plug fixed at the bottom of inner glass tube. The

calomel electrode comes in contact with the experimental solution through a frit arranged to the

outer jacket. The potential of this electrode depends on the concentration of KCl taken in the

outer jacket. Some of the most popularly used concentrations of KCl and corresponding single

electrode (reduction) potentials on the hydrogen scale at 250C are given below.

0.1M KCl | Hg2Cl2 (s) | Hg, pt 0.3338 V

1.0M KCl | Hg2Cl2 (s) | Hg, pt 0.2800 V

Saturated KCl | Hg2Cl2 (s) | Hg, pt 0.2415V

Ion Sel ec tive el ec tr odes:

These are the electrodes, which responds to specific ions only and develops a potential

against that ions while ignoring the other ions present in the solution.

Principle:

ISE consist of an ion selective membrane in contact with an analyte solution on one side

and an internal reference electrode on the other side. An internal reference electrode is

constituted in contact with reference solution. In ion selective electrodes there must be a

membrane.

The types (classification) of ion-sensitive membranes:

a)

Glass membrane

b) Solid state membrane

c) Heterogeneous membrane

d) Liquid membrane

Glass Electrode: Most often used pH electrodes are

called glass electrodes and belong to the family of

ISEs. They are sensitive only to H+ ions. Typical

glass electrode is made of glass tube engaged with

small glass bubble sensitive to protons. Inside of the

electrode is usually filled with buffered solution of

chlorides in which silver wire covered with silver

chloride is immersed. pH of internal solution varies- for example it can be 1.0 (0.1M HCl) or 7.0


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Active part of the electrode is the glass bubble. While tube has strong and thick walls,

bubble is made to be as thin as possible. Surface of the glass is protonated by both internal and

external solution till equilibrium is achieved. Both sides of the glass are charged by the adsorbed

protons, this charge is responsible for potential difference. This potential in turn is described by

the Nernst equation and is directly proportional to the pH difference between solutions on both

sides of the glass.

The majority of pH electrodes available now a day are combination electrodes that have

both glass H+ ion sensitive electrode and reference electrode compartments, conveniently placed

in one housing.

Applications of Ion-sensitive electrode:

1) To determine ions like H+, K+, Li+, etc.

2)

To determine hardness of water (Ca+2 and Mg+2 ions)

3) To determine concentration of F-, NO3-, CN- etc.

4) To determine concentration of a gas using gas-sensing electrodes.

5) To determine pH of a solution using H+ ion sensitive electrode.

Batter ies and fuel c el l s

Battery

“Battery is a device consisting of two or more galvanic cells arranged in series or

parallel or both, that can generate electrical energy”.

When two or more electrochemical cells are electrically interconnected, each of which

containing two electrodes and an electrolyte is called a Battery. In everyday usage, “battery” is

also used to refer to a single cell. No one battery design is perfect for every application. There

are many parameters like cost, voltage, duty cycle, dimension, stability with time and

temperature, shelf life, etc., on which a battery is selected for a particular operation.

Batteries are classified into two categories depending on their recharging capabilities.1. Primary batteries

2. Secondary Batteries


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Primary Batteries:

These are non-rechargeable and are meant for a single use and meant to be discarded

after use. Primary batteries are non-rechargeable and are less expensive and are often used in

ordinary gadgets like torch lights, watches and toys. Commercially many kinds of primary

batteries are available.

Eg : Leclanche cell, Alkaline cell and Lithium cell

The most familiar example of this type is the dry cell

(known as Leclanche cell after its discoverer) which is used

commonly in our transistors and clocks. The cell consists of

a zinc container that also acts as anode and the cathode is a

carbon (graphite) rod surrounded by powdered manganese

dioxide and carbon. The space between the electrodes is

filled by a moist paste of ammonium chloride (NH4Cl) and

zinc chloride (ZnCl2). The electrode reactions are complex,

but they can be written approximately as follows :

Anode: Zn(s) Zn2+ + 2e –

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

NH4++OH- → NH3+ H2O

Ammonia produced in the reaction forms a complex with Zn

2+

to give [Zn(NH3)2Cl2].The cell has a potential of nearly 1.5 V.

Secondary Batteries:

Secondary batteries are rechargeable and are meant for a multi cycle use. After every use

the electrochemical reaction could be reversed by external application of voltage. The cycle is

reversed till the capacity fades or lost due to leakage or internal short circuit.

These cells are rechargeable and reusable. A combination of all reversible

electrochemical cells gives secondary batteries. Many kinds of secondary batteries are available

in the market.

Eg : Lead-acid cell, Ni/Cd cells, Ni-Metal hydride cell and Lithium ion cells.


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Lead-acid cells:

Anode: Sponge metallic Lead

Cathode: Lead-dioxide (PbO2)

Electrolyte: Dilute mixture of aqueous sulfuric acid.

Applications: Automobiles and construction equipment, standby/backup systems

Used mainly for engine batteries, these cells represent over half of all battery sales. Some

advantages are their low cost, long life cycle, and ability to withstand mistreatment. They also

perform well in high and low temperatures and in high-drain applications.

The half-cell reactions are:

Pb + SO42- PbSO4 + 2e- 0.356 V

PbO2 + SO42- + 4H+ +2e- PbSO4 + 2H2O 1.685 V

There are a few problems with this

design, if the cell voltages exceeds 2.39V, the

water breaks down into hydrogen and oxygen

and may lead to explosion. Another problem

arising from this system is that fumes from the

acid solution may have a corrosive effect on

the area surrounding the battery.

These cells have a low cycle life, a

quick self discharge, and low energy densities.

However, with a nominal voltage of 2V and power densities of up 600 W/kg, the lead-acid cell is

an adequate, if not perfect, design for car batteries.

Nickel-Cadmium Cells:

Anode: Cadmium

Cathode: Nickel oxy-hydroxide (NiO(OH))

Electrolyte: Aqueous potassium hydroxide (KOH)

Applications: Calculators, digital cameras, pagers, laptops, tape recorders, flashlights, medical

devices (defibrillators), electrical vehicles, space applications.


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The cathode is Nickel-plated, woven mesh, and the

anode is a Cadmium-plated net. The electrolyte, KOH, acts

only as an ion conductor and does not contribute

significantly to the cell’s reaction. That’s why not many

electrolytes are needed, so this keeps the weight down.

The cell reactions are as follows:

Cd + 2OH- Cd(OH)2 + 2e- 0.81 V

NiO(OH) + 2H2O + 2e- Ni(OH)2.H2O + OH

- 0.45 V

Cd + NiOOH + 2H2O +OH- Cd(OH)2 + Ni(OH)2.H2O 1.26 V

Advantages include good performance in high-discharge and low-temperature

applications. They also have long shelf and use life. Disadvantages are that they cost more than

the lead –acid battery and have lower power densities. Possibly it’s most well-known limitation

is a memory effect, where the cell retains the characteristics of the previous cycle. This term

refers to a temporary loss of cell capacity, which occurs when a cell is recharged without being

fully discharged. This can cause cadmium hydroxide to passive the electrode, or the battery to

wear out. In the former case, a few cycles of discharging and charging the cell help correct the

problem, but may shorten the lifetime of the battery.

Fuel c el l s:

“Fuel cells are the galvanic cells which convert chemical energy of a fuel- oxidant

system directly into electrical energy by oxidation of fuel at anode and reduction of oxidant at

cathode”.

or

“A fuel cell is a device that converts the chemical energy of a fuel (hydrogen, natural

gas, methanol, gasoline, etc.) and an oxidant (air or oxygen) into electricity”.

Fuel + oxygen Oxidation products + electrical energy


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A cell in which one or both of the reactants are not permanently contained in the cell, but

are continuously supplied from a source external to the cell and the reaction products

continuously removed is called a fuel cell. Unlike the metal anodes typically used in batteries,

the fuels in a fuel cell are usually gas or liquid, with oxygen as the oxidant.

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 1.0 V. 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 musti) 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 outcome

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


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Direct methanol-Oxygen fuel cell: It consists of two electrodes separated by a proton

exchange membrane (PEM) and connected via an external circuit that allows the conversion of

free energy from the chemical reaction of methanol with air or oxygen to be directly converted

into electrical energy.

Anode: CH3OH + H2O → CO2 + 6H+ + 6e-

Cathode: 3/2O2 + 6H+ + 6e- → 3H2O

Overall: CH3OH(l) +3/2O2(g) → CO2(g) + 2H2O(l) Eocell = 1.21 V

The standard electromotive force (emf) of an ideal direct methanol-oxygen fuel cell is

1.21 V. However, because methanol may be incompletely oxidized which lead to the

formation of methanal or methanoic acid, the emf produced is lower than the theoretical one.

In this experiment, direct methanol-oxygen fuel cell that

uses platinum (Pt) as catalyst and dilute sulfuric acid as

the electrolyte. Methanol is the fuel at anode

compartment and the oxygen gas dissolved in the

solution will undergo reduction reaction at the cathode.

Advantages of Fuel Cells:

No emission of toxic gases and chemical wastes

are in safe limits. The reactants and products are environment friendly and only we have

to bother about disposal of cell material.

High efficiency (75-85%) of energy conversion from chemical energy to electrical

energy. So offer an excellent use of our renewable energy resources.

No noise pollution like in generators and low thermal pollution.

Low maintenance costs, fuel transportation costs, cell parts are modular and

exchangeable.

The fuels and electrolyte materials are available in plenty and inexhaustible unlike fossil

fuels.

Unlike solar cells, fuels cells are compact and transportable.

Unlike acid cells used in automotives the fuel cells are far less corrosive.

Unlike nuclear energy, fuel cell energy is economical and safe.

ΩΩΩΩΩ

Unit II Electrochemistry.

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Transcript of Unit II Electrochemistry.