Electrochemistry Lecture Notes

LESSON PLAN

LP- Engg.Chem II

LP Rev.No:00

Date:

Page 1of 6

Sub Name & Name :CY 2161 ENGINEERING CHEMISTRY- II

Unit: 1 Branch: Common to all I Yr B.E / B.Tech

(except Marine Engineering)

Unit syllabus: ELECTROCHEMISTRY 9 hrs.

Electrochemical cells – reversible and irreversible cells – EMF – measurement of emf – Single electrode potential – Nernst equation (problem) – reference electrodes –Standard Hydrogen electrode -Calomel electrode – Ion selective electrode – glass electrode and measurement of pH – electrochemical series – significance – potentiometer titrations (redox - Fe²+ vs dichromate and precipitation – Ag+ vs CI- titrations) and conduct metric titrations (acid-base – HCI vs, NaOH) titrations,

Objective: To impart knowledge on Electro chemistry detailing electrodes, cells, conventional representation of cells and measurement of potentials.

SessionNo Topics to be covered

Timemin Ref

Teaching Aids

1 Electrochemical cells –Galvanic & Electrolytic cells reversible and irreversible cells, examples

50 1,2 BB & Chalk

2 EMF – Definition, measurement of emf (Poggendorff’s method ) – Single electrode potential

-do- 1,2 -do-

3 Nernst equation – Derivation from Thermodynamic considerations – reference electrodes –Standard Hydrogen electrode -Calomel electrode

-do- 1,2 -do-

4 Ion selective electrode – glass electrode and measurement of pH

-do- 1,2 -do-

5 electrochemical series – significance. -do- 1,2 -do-

6 potentiometer titrations (redox - Fe²+ vs dichromate and precipitation – Ag+ vs CI- titrations)

-do- 1,2 -do-

7 conductometric titrations (acid-base – HCI vs, NaOH) titrations,

-do- 1,2 -do-

8 Problems on measurement of electrode and cell potentials using Nernst equation

-do- 1,2 -do-

9. Revision

Lecturer Notes –Electrochemistry

1

Lecturer Session No: 01 – Topic: Electrochemical cells –Galvanic & Electrolytic cells reversible and irreversible cells, examples

Electrochemical cells:

The devices used for the inter-conversion between chemical and electrical forms

of energy are called electrochemical cells. They are of two types namely galvanic cells

and electrolytic cells.

Electrolytic cell is one, which uses electric energy for the chemical reactions

(decomposition) to take place. Redox processes are effected by electrolytic cells.

Ex : Chlor-alkali generation by the passage of electric current through brine (NaCl)

solution.

At Cathode:

Na+ + e– Na (Primary process)

Na + H2O NaOH+ 1/2 H2 (Secondary process)

At anode:

Cl - Cl + e- (Primary process)

Cl + Cl Cl2 (Secondary process)

Galvanic cell / Electrochemical cell is a device, which produces electric energy at the

expense of chemical energy.

Ex: Daniel cell made of copper and zinc electrodes possesses an emf of 1.1 V with the

cell reaction mentioned below:

Zn + Cu2+ Zn2+ + Cu

At Anode: Zn Zn2+ + 2 e ;

At Cathode: Cu2+ + 2 e Cu

Net cell reaction Zn + CuSO4 ZnSO4 + Cu

Electrolytic cells consume electric energy to effect chemical reactions whereas

electrochemical cells produce electric energy at the expense of chemical energy.

Reversible and Irreversible Cells:

2

Electrochemical cells are of two types namely reversible and irreversible cell.

Reversible cell is an electrochemical with the reversibility of the electrode realizable i.e.

if an infinitesimally small but excess applied emf (electro motive force) could

change the direction of the electrode reaction, the cell is termed as reversible cell. E.g.

Daniel cell is a reversible cell. Its cell potential is 1.1 V. Thus in Daniel cell (a galvanic

cell), zinc undergoes dissolution and copper undergoes deposition to realize an emf of

1.1V, as per the following reaction sequence:

Zn + Cu2+ === Zn2+ + Cu …….. 1.1 V

If an emf of –1.101 V is impressed on Daniel cell (by interchanging the polarity /

terminals – Remember that voltage is a vector quantity, characterized by direction),

copper undergoes dissolution and zinc undergoes deposition. These changes can be

visually observed by the fading / deepening of the blue colour of the cupric (sulphate)

solution on interchanging the terminals. Also, the solid deposit becomes brown if copper

and grey / black if zinc.

Irreversible is a cell where the cell reaction cannot be reversed even on applying

infinitesimally small but excess applied emf i.e. the products produced during the cell

reaction are not available for recombination on reversal of voltage. Example of an

irreversible cell is the cell used for the electrolysis of brine or the dry cell used in pen-

torches. In the electrolysis of brine (aqueous NaCl solution) for example, on applying

voltage, Na+ ions move towards cathode, gain one electrode and become elemental

sodium atoms. But the sodium atoms immediately react with water to form sodium

hydroxide. Similarly, chloride ions move towards anode, loose one electron to form

chlorine atoms. These chlorine atoms recombine forming molecular chlorine, which is

evolved as a gas. The reaction sequence is given below:

Na+ + e → Na ; 2 Na + 2 H2O → 2 NaOH + H2

Cl- → Cl + e ; Cl + Cl → Cl2 ↑

Thus the products are not available for recombination, even on reversal of the

voltage. A similar type of reactions takes place with dry cell also. Thus these cells are

irreversible cells.

3

Another standard example of a reversible cell is a cell made of zinc and silver

electrodes immersed in dil. H2SO4. When these electrodes are connected externally, zinc

dissolves with the liberation of hydrogen gas.

Zn + H2SO4 → ZnSO4 + H2 ↑

The reaction at the silver electrode, on the other hand, is

2 Ag+ + 2 e → 2 Ag

The thermodynamic concept of reversibility applied to reversible cells involves the

following conditions:

1. The driving force is infinitesimally greater than the opposing force.

2. The process can be reversed if the external force, infinitesimally greater than the opposing force is applied to the system in opposite direction.

Thus on applying slight excess emf in the opposite direction, the reaction cannot be reversed as hydrogen gas has already evolved. Such a cell, which does not obey the conditions of thermodynamic reversibility, is termed as “irreversible cell”.


Lecturer Session No: 02 – Topic: EMF – Definition, measurement of emf (Poggendorff’s method ) – Single electrode potential

4

Electromotive Force

The emf (electro motive force) of a cell is the algebraic sum of the potentials of

the two constituent single electrode systems. It is obvious that cell is made of two half-

cells / single electrode systems. A cell is generally represented with the negative

electrode / anode written first at the left and then the cathode / positive electrode at the

right. Thus the emf of a galvanic cell is calculated from the half-cell potentials using the

relation

Ecell = Eright - Eleft = Ecathode - Eanode

Here it is to be noted the values of std potentials (reduction potentials) of cathodes

are more positive and those of anodes are more negative, so that the cell potential is

positive.

e.g. for Daniel cell, Ecell = Eright - Eleft = Ecathode - Eanode =

ECu2+

/ Cu – EZn2+

/ Zn = 0.34 – (-0.76) = 1.10 Volt

Measurement of EMF

The potential difference, which causes current flow from the electrode of

higher potential to the electrode of lower potential, is called electromotive force (emf) of

the cell. The emf of a cell cannot be directly determined by connecting across a

voltmeter, as some part of the cell current is drawn by the voltmeter during its

measurement. This results in the formation of reaction products at the electrodes and

hence a change in the electrolyte concentration around the electrodes. This difficulty can

be overcome by the measurement of excess applied opposite emf that just nullifies the

cell emf (Pogendorff’s external compensation method). Care is to be taken during the

measurement such that the current taken from the cell is negligibly small and the ionic

concentrations are not appreciably altered. The emf of the cell thus remains constant and

its value can be determined with high degree of accuracy / precision.

The emf of acell can be determined by Pogendorff’s external compensation method. In

this potentiometric measurement of cell emf, a standard cell is used whose emf is known

and does not vary with time. Weston cadmium cell is the conventionally used standard

cell. Fig. below is the schematic of a simple potentiometer.

C R

5

E E’

A B

Std. Cell S

K

Cell X

It consists of a uniform wire AB of high resistance. A storage battery C of constant but

large emf is connected to the ends A and B of the wire through a variable resistance

(rheostat) R. The cell X whose emf is to be determined is included in the circuit through a

galvanometer G and a sliding contact E. The circuit is closed using a plug key K. The

position of the sliding contact D is slowly changed along the wire AB till a point is

reached at which there is no net current flowing through the galvanometer (its deflector

points to zero). This null point position Dx is noted.

The standard cell S (whose emf is known) is then introduced into the circuit and

the circuit closed using the plug key K. The position of the sliding contact D is again

slowly changed along the wire AB till a point is reached at which there is no net current

flowing through the galvanometer (its deflector points to zero). This null point position

Ds is noted. The emf of the cell X (unknown) is determined using the relation mentioned

below:

Emf of cell E α balancing length AD

Emf of standard cell S Es α balancing length ADs

Emf of cell X Ex α balancing length ADx

Thus Es / Ex = (length ADs) / (length ADx), from which Ex can be determined, as the

value of Es is known. Nowadays the digital potentiometers are used which have the in-

built circuitry of potentiometer set up and standard cell with switching arrangement for

standard cell so that the unknown cell is connected externally to directly read the cell emf

as digital display.

Single Electrode Potential

When a zinc metal strip is immersed in copper sulphate solution for a short while

and taken out, a brown deposit (of copper) can be observed on the bright colored strip,

6

which turns black (due to the formation of cupric oxide) with time. On the other hand,

zinc strip when immersed in dilute hydrochloric acid, produces copious gas (hydrogen)

evolution. These phenomena indicate that metals are in equilibrium with their cations in

solution and the equilibrium is shifted towards either metal dissolution or metal

deposition, depending on the nature of metal. Metal dissolution (M Mn+ + n e) process

can be viewed as oxidation and metal deposition process (Mn+ + n e →M) as reduction.

The metal dissolution (oxidation) process is spontaneous for metals such as zinc whereas

the metal deposition (reduction) process is spontaneous for metals such as copper. The

former types of metals are called anodes and the latter type as cathodes. Thus the anodic

or cathodic behaviour of metal explains the phenomena such as the spontaneous

immersion deposit of copper and gas evolution by zinc with HCl. Similarly non-metallic

elements will be in equilibrium with their anions (ex. Cl with Cl -). The tendency and the

extent of the equilibrium between the element and its ion varies with element and this

tendency termed as electrode potential.

Electrode potential of any system can be defined as the tendency of any element to

be in equilibrium with its either cation or anion in solution (aqueous). The system

possessing this tendency is termed as electrode system. Generally metals will be in

equilibrium with their cations in solutions whereas non-metallic elements will be in

equilibrium with their anions in solutions. Thus the condition for an electrode system is

the establishment of equilibrium between the element and its ion and not the electronic

conduction.

Ex : Copper electrode system consists of a copper strip immersed in any copper salt

solution as electrolyte. Electrode reaction for this system is Cu 2+ + 2 e Cu. The

electrode system is represented as Cu | Cu2+. Since some electrode systems lack electronic

conductivity, platinum electrode (which does not take part in electrode reaction) is

provided and the electrode system set up.

Ex. Hydrogen electrode is set up with Pt and represented as H2 | H+ | Pt.

Electrode reaction: H2 2 H+ + 2 e-

7

Platinum does not involve in electrode reaction. Thus, hydrogen electrode consists of

platinum metal immersed in hydrochloric acid (of concentration 1 molar) into which

hydrogen gas is bubbled at a particular pressure (1 atmosphere), called standard

hydrogen electrode (SHE).Thus an electrode system consists of an element, its cation or

anion in solution with or without platinum support for electronic conductivity. The

equilibrium exists between the element and the ion but the extent varies from system to

system. The associated reversible reaction is termed as electrode reaction.

1. 2H+ + 2e H2 is the electrode reaction in hydrogen electrode.

2. Cu2+ + 2e Cu is the electrode reaction in copper electrode.

Thus the tendency of element to be in equilibrium with its cation or anion or the extent of

reversibility of the associated electrode reaction is spoken of as electrode potential. It

takes the unit of volt and is a vector quantity. If the reaction is spontaneous, potential is

positive and vice-versa. For a given electrode system, either oxidation (loss of electrons)

or reduction (gain of electrons) is spontaneous and accordingly the potential will be of

positive or negative sign. For ease of convention, reduction potential is taken as positive

i.e. a system having positive potential has the reduction reaction spontaneous.

Ex: electrode potential is + 0.34 V means that the reaction Cu2+ + 2e Cu is spontaneous

for copper electrode. Zinc electrode potential is - 0.76 V meaning that Zn Zn2+ + 2 e is

spontaneous reaction for zinc electrode.

Anodes are represented with the element symbol first, followed by a vertical line

(which shows electrode-electrolyte interface), which is turn is followed by the ionic form

of the electrolyte. Ex. Zinc electrode is represented as Zn | Zn2+; cathodes are represented

in similar manner but in the reverse order. Ex: Copper electrode is represented as Cu2+ |

Cu.

The potential of an electrode system depends on (i) the nature of the metal /

element (electron transfer number), (ii) temperature of the electrolyte and (iii)

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concentration of the electrolyte or pressure of the gas for gaseous systems. As the

potential of the same electrode system varies with these conditions, potential under a set

of standard conditions is considered as standard electrode potential. Thus standard

electrode potential is defined as the (reduction) potential of the electrode system when all

the components are at their unit active levels i.e. tendency of element to be in equilibrium

with its cation or anion or the extent of reversibility of the associated electrode reaction

when the electrolyte concentration is unity (1 molar) and the gas pressure is 1

atmosphere.


Lecturer Session No: 03 – Topic: Nernst equation – Derivation from

Thermodynamic considerations – reference electrodes –Standard Hydrogen

electrode -Calomel electrode

Nernst Equation:

9

The potential of any electrode system (E) depends on (i) nature of the metal / element

(ii) temperature and concentration of the electrolyte. The functional dependence of

potential of any electrode system (E) on these factors is given by Nernst equation

E = E 0 + (RT / nF) ( log Mn+) where Eo , n, T, R and F denote respectively. The standard

electrode potential, charge transfer number, absolute / Kelvin temperature, Universal Gas

Constant and Faraday’s constant. The derivation / explanation of the terms are as follows:

It is obvious that the concept of electrode potential evolves from the inter-conversion

between chemical and electrical forms of energy. The equation pertaining to these two

forms of energy is

G = G0 + RT ln K ----- (1) where G, G0, K and R are respectively

the free energy change for a process under given conditions of Kelvin Temperature,

standard free energy change for the process, equilibrium constant for the process and

universal gas constant.

The free energy change for any chemical to electrical energy conversion process

is given by the equation

G = - nFE ----- (2) where E is the potential of the electrode system,

F – Faraday’s constant = 96496 or 96500 coulombs and ‘n’ is the number of electrons

transferred between the element and the ion which are in equilibrium.

Combining equations (1) and (2), we get

- nFE = - nFE0 + RT ln K ----- (1) where E0 is the standard electrode potential of the electrode system.

Consider the reduction reaction

Mn+ + n e → M

Equation (3) becomes

E = E0 - (RT / nF) ln (aproducts / areactants ) ----- (4)

10

where ‘a’ is the activity of the species. The activity of a (homogenous / uniform) solid is taken as unity, that of the electrolyte expressed in terms of the concentration and that of a gas (or gaseous mixture) expressed in terms of pressure (or partial pressure) of the gas.

Equation (4) can be written as

E = E0 - (RT / nF) ln (aM / aMn+ )

(or) E = E0 + (RT / nF) ln [Mn+] ----- (4)

where [Mn+] is the concentration of the electrolyte / metal ion in solution.

i.e. E = E0 + (RT / nF) ln [Mn+] ----- (5)

substituting the values of R= 8.314 Joules, F = 96500 Coulombs and introducing the factor 2.303 to convert natural logarithms to common logarithms, equation (5) becomes

E = E0 + (0.059 / n) log [Mn+] ----- (6)

Eequations (5) and (6) are the two forms of Nernst equation which gives the dependence of electrode potential on the factors mentioned.

Problems:Calculate the std emf of the cell Zn | ZnSO4 || CuSO4 | Cu, if the std

electrode potentials of copper and zinc are respectively 0.337 V and –0.763 V.

Reference Electrodes:

The potential of an electrode system (electrode of interest or working electrode)

can be measured by coupling with other electrode with a voltmeter introduced between

them. The coupled electrode should not possess any charge transfer reaction (electrode

reaction) in the electrolyte used or it should not be polarized. Such ideally non-

polarizable electrodes used for the measurement of working electrodes are called

11

reference electrodes. Reference electrodes are two types namely primary and secondary

reference electrodes. Primary reference electrode is one that is universally used such as

standard hydrogen electrode (SHE) and its potential is arbitrarily taken as zero. But

SHE involves tedious and cumbersome construction. This difficulty is overcome by the

use of ‘secondary reference electrodes, which can be constructed easily and their

potentials can be determined with SHE as reference.

Examples of secondary reference electrodes are calomel electrode, silver-silver electrode,

glass electrode, quinhydrone electrode etc. Calomel electrode is set up with mercurous

chloride, mercury and potassium chloride electrolyte and represented as Hg | Hg2Cl2 |

KCl. Depending on the concentration of KCl, calomel electrode is of three types namely

saturated normal and decinormal calomel electrodes.

12

1 N HCl

Pt. foil

H2 (1 atm.)

Saturated KCl

Hg + Hg2Cl2

Hg

Pt wire

The electrode reaction for calomel electrode is ½ Hg2Cl2 (s) + e → Hg(s) + Cl-

The potentials of the different calomel electrodes against SHE reference is given below:

[ KCl ]Saturated 1.0 N 0.1 N

Potential , V 0.2422 0.2810 0.3335

Merits and Demerits of Calomel electrode: 1. Gives relative pH values compared to

hydrogen electrode, which gives absolute values of pH.

Ag/AgCl electrode is prepared by depositing a thin layer of AgCl electrolytically on a Ag

or Pt wire and immersing in a solution containing the chloride ions. It is represented as

Ag | AgCl | Cl-( M)

The electrode reaction of Ag/AgCl electrode is : AgCl + e → Ag + Cl-

13


Lecturer Session No: 04 – Topic: Ion selective electrode – glass electrode and

measurement of pH

Glass electrode is made of special glass of relatively low melting point and electrical

conductivity with composition Na2O (22%), CaO (6%) and SiO2 (72%). The glass

elctrode assembly consists of a thin glass bulb filled with 0.1 N HCl and a silver wire

coated with silver chloride immersed in it. The Ag/AgCl electrode here acts as the

internal reference electrode. The glass electrode is represented as

Ag | AgCl(s) | 0.1 M HCl | glass.

14

Working: when glass electrode is immersed in the solution whose pHis to be determined,

a potential difference is set up between the two surfaces of the glass membrane. The

potential value developed is proportional to the pH of the test solution (sample). Actually,

the glass membrane of the glass electrode undergoes an ion-exchange reaction in which

the sodium ions of the glass membrane are exchanged with protons of the sample

solution. The electrode reaction of the glass electrode immersed in the test solution can

be represented as

glass ---- Na+ + H+ = glass ---- H+ + Na+

Ion selective electrodes (ISEs): ISEs are the electrodes useful in the qualitative and

quantitative analysis of sample ion only, in a mixture of variety of species. As the

potential of an electrode system varies with the ionic concentration of the electrolyte, the

electrode system can be used as an ISE by coupling with a suitable reference electrode.

ISEs are normally used with saturated calomel electrode (SCE) reference and the

developed potentials are measured using potentiometers or pH meters. The use of metals

directly as ISE has the following disadvantages: (i) slow electrode response (ii) Nernst

equation not followed (iii) chage of electrode potential due to the availability of electrons

on theelectrode surface (iv) no well defined electron change.

Hence various membranes are used in ISEs and such electrodes are called Ion selective

membrane electrodes (ISMEs). ISMEs show some degree of specificity and selectivity,

These electrodes utilize some membrane to confine an inner solution and the reference

electrode. Membranes in the ISE and reference electrode (RE) sides function by ion

exchange (IE) mechanism. Various types of ISEs are as follows: 1. Glass membrane

electrodes 2. Liquid membrane electrodes 3. Double membrane electrodes 4. Solid-state

membrane electrodes 5. Precipitate membrane electrodes.

15


Lecturer Session No: 05 – Topic: Electrochemical series – significance.

Electrochemical series:

Electrochemical series is the arrangement of elements in the ascending (or descending)

order of their electrode (reduction) potential values with hydrogen at the centre.

Electrode systems appearing earlier in the series have the oxidation reaction spontaneous

and are termed as anodes and those appearing later in the series have the reduction

reaction spontaneous and are termed as cathodes.

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Electrode system Standard Potential (E0), volt

Li | Li+

Mg | Mg2+

Zn | Zn2+

Sn | Sn2+

H2 | 2 H+

Cu2+ | Cu

Ag+ | Ag

F2 | 2 F-

-3.05

-2.40

-0.76

-0.44

0.00

+0.34

+0.80

+2.87

The potential of a redox electrode system is given merely as a number (modulus value,

irrespective sign). Thus if the reduction potential is positive, the oxidation potential is

negative for the same system and vice-versa. Std. electrode potential conventionally

represents the reduction potential or potential of the reduction reaction.

Importance / Significance / Applications of electrochemical series: Electrochemical series

is useful / significant in the prediction of

1. Spontaneity of reactions: The feasibility / spontaneity of reactions of the reactions

be predicted from the knowledge of the electrode potential values. Such processes

are spontaneous whose standard potential is positive.

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2. Corrosion behaviour of metals and alloys: Metals / alloys with negative values of

std. electrode potentials are prone / susceptible to corrosion and those with positive

values of std. electrode potentials are resistant to corrosion (an undesirable phenomenon).

e.g. zinc is more easily corroded than copper (or) copper is more resistant to corrosion

than zinc.

3. Redox behaviour of materials: Materials with more negative values of std.

electrode potentials are used in reduction reactions – addition of electrons (as they

can donate electrons) and Materials with more posittive values of std. electrode

potentials are used in oxidation reactions – removal of electrons (as they can

accept electrons). E.g. zinc, tin etc. are used as reducing agents whereas oxides of

copper etc are used as oxidizing agents.

4. Displacement characteristics of metals: Metals with more negative std. electrode

potentials will displace metals with more positive std. electrode potential values.

E.g. zinc will displace copper from its salt solution and not vice-versa.

5. Determination of equilibrium constant of the reaction from the knowledge of

electrode potential values, using the relation ΔG = - nFE = - RT ln K.

Galvanic series is the arrangement of industrial metals or alloys in the ascending or

descending order of their electrode potential values in the seawater electrolyte medium. It

predicts the corrosion tendencies of metals and alloys in common corroding media. It

differs from the electrochemical series in that it is more practical than electrochemical

series and the potentials are expressed with reference to Saturated Calomel Electrode

(SCE) whereas in electrochemical series, the potentials are expressed with reference to

Standard Hydrogen Electrode (SHE). It is to be noted that the potential of pure metal is

different from that of the alloy. Galvanic series is more useful as it predicts the corrosion

behaviour of different metals and alloys in various corroding media.

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Lecturer Session No: 06 – Topic: Potentiometer titrations (redox - Fe²+ vs dichromate

and precipitation – Ag+ vs CI- titrations)

Potentiometric Titration:

Potentiometric titration is used for following the course of reactions involving

electrolytes, where there is no proper indicator available. The potential of the electrode

system is determined using the platinum and calomel electrodes immersed in the reaction

mixture after regular additions of the titrant. The potential observed is plotted against the

volume of titrant added. The endpoint is determined graphically from the change in

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trends before and after the completion of the reaction. E.g. the redox reaction between

ferrous (ammonium) sulphate (FAS) and potassium permanganate or dichromate is

followed by determining the potential using the platinum and calomel electrodes

immersed in the reaction mixture after adding regular volumes of permanganate or

dichromate. The potential observed is plotted against the volume of titrant added (E Vs V

plot). The endpoint is determined graphically from the change in trends before and after

the completion of the reaction. A derivative plot can also de made between change in

potential to change in volume (ΔE/ΔV) and the average volume of the titrant [(V1+V2)/2].

The volume corresponding to the peak in the derivative plot directly gives the end point

condition whereas the involved electrode systems and their potentials can be got from the

E Vs V plot. The model graphs for the potentiometric titrations are given below:

The S-shaped curve for the E Vs V plot is due to the fact that the electrode system

itself is changed after the completion of reaction i.e. initially when a small amount of

permanganate (MnO4- with manganese in its Mn7+ state) is added to the reaction

mixture containing FAS and dilute sulphuric acid in a beaker, corresponding amount

of ferrous ions (Fe2+) is oxidized to ferric state (Fe3+) and permanganate is reduced to

Mn2+ state. Thus the beaker contents are two ionic (redox) species of iron (i.e.Fe2+ and

Fe3+) with platinum electrode (inert electrode, contributing only electronic

conductivity) immersed. This constitutes the iron electrode system, whose potential

varies gradually (as given by Nernst equation) with regular additions of titrant. The

Volume of titrant (ml)V1 + V2

2

∆E

/ ∆

V

mV

/ml

20

Em

f (V

)

After the end point (reaction completion), all the ferrous ions are completely oxidized

to ferric ions and the excess added permanganate ions (Mn7+) exist as redox couple

with manganous ions (Mn2+), leaving only one type of species for iron (Fe3+) in the

beaker. Now the electrode system is manganese electrode system, whose potential

varies gradually (as given by Nernst equation) with regular addition. As the electrode

system itself is changed during the reaction, there is a shoot up in potential in the

E Vs V plot and hence it is S-shaped.

pH determination: A cell is constructed with saturated calomel electrode (SCE) or and

a glass electrode or quinhydrone electrode immersed in the solution, whose pH is to

be determined and its potential measured. By applying Nernst equation for the cell

constructed, the pH of the sample can be determined. where the potentials are in

volts. The glass electrode potential can be determined by previously dipping in

solutions of known pH (buffers) and noting the potential observed with the similar

cells construted. The actual cell representation is

Glass | solution of unknown pH || SCE

pH of the buffer = pHb = (ESCE- EG0- Ecell) / 0.0591 ------- (i)

where Ebcell is the cell potential observed with the buffer.

pH of the sample = pHs = (ESCE- EG0- Ecell) / 0.0591 ------ (ii)

where Escell is the cell potential observed with the sample.

From Equations (i) and (ii) pH of the sample can be calculated from the observed cell

potentials. Similarly, pH determination can be made with other electrodes such as

quinhydrone electrode.

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Pt. wire

0.1M HCl

Thin walled glass bulb

Saturated KCl

Hg + Hg2Cl2

Hg

Pt wire

Potentiometer


Lecturer Session No: 07 – Topic: Conductometric titrations (acid-base – HCI vs,

NaOH) titration).

Conductometric titration of strong acid against a strong base:

The course of neutralization of a strong acid by a strong base can be determined by

condutometric method (without the use of indicator). The principle of conductometric

titration is that the conductance of a reaction follows a specific trend before the

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completion of the reaction and it follows a different trend after the completion of the

reaction. From the change in trends, the end point of the reaction can be determined

graphically. The model graph for this titration is shown in fig.

In the case of neutralization of strong acid (say HCl) by a strong base (say NaOH), the

conductance of HCl (taken in a beaker) is determined at regular additions of NaOH, from

the burette.

H+Cl- + Na+OH- → Na+Cl- + H2O

The conductance of the reaction mixture (HCl) decreases till the end point because lighter

protons (H+) are replaced by heavier Na+ ions, whose mobility is lower than that of

protons. The conductance of the reaction mixture increases after the end point because

heavier chloride ions (Cl- whose atomic mass is 35.5) are replaced by lighter hydroxyl

(OH-) ions (with mass 17), whose mobility is higher than that of chloride ions. A plot of

the conductance of the reaction mixture against the volume of the titrant gives two

straight lines of opposite slopes. The point of intersection of the straight lines is the end

point.

Vol. of titrant (ml) Vol. of titrant (ml)

Con

duct

ance

(m

ho)

Con

duct

ance

(m

ho)

Neutralisation titration Precipitation titration

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Conductometric titration of weak acid against a strong base:

The course of neutralization of a weak acid by a strong base can be determined by





graphically. In the case of neutralization of weak acid (say CH3COOH) by a strong base

(say NaOH), the conductance of CH3COOH (taken in a beaker) is determined at regular

additions of NaOH, from a burette.

CH3COOH + Na+OH- → CH3COO-Na+ + H2O(weak) (feebly ionized)

The conductance of the reaction mixture (CH3COOH) is initially low because of the poor

dissociation of the weak electrolyte. On the addition of strong alkali, the strong

electrolyte, sodium acetate is formed. It tends to suppress the ionoization of acetic acid,

due to common ion effect. Further there is increase in conductance of the reaction

mixture, because of the larger proportion of the strong electrolyte, CH3COONa.

Immediately after the end point, the conductance increases sharply because of the fast

moving hydroxyl ions compared to the acetate ions. A plot of the conductance of the

reaction mixture against the volume of the titrant gives two straight lines of positive, but

varying slopes. The point of intersection of the straight lines is the end point.

Conductometric titration of strong acid against a weak base:

The course of neutralization of a strong acid by a weak base can be determined by





graphically. In the case of neutralization of strong acid (say HCl) by a weak base (say

NH4OH), the conductance of HCl (taken in a beaker) is determined at regular additions of

NH4OH, from a burette.

H+Cl- + NH4OH → NH4+Cl- + H2O

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(weak)

The conductance of the reaction mixture (HCl) decreases sharply, with the addition of the

weak base, ammonium hydroxide, because of the formation of the salt ammonim

chloride, which is a weaker electrolyte than HCl. Immediately after the end point, the

conductance decreases further but slowly because ammonium hydroxide is a still weaker

electrolyte than ammonium chloride.A plot of the conductance of the reaction mixture

against the volume of the titrant gives two straight lines of negative, but varying slopes.

The point of intersection of the straight lines is the end point.

Conductometric titration of weak acid against a weak base:

The course of neutralization of a weak acid by a weak base can be determined by





graphically. In the case of neutralization of weak acid (say CH3COOH) by a weak base

(say NH4OH), the conductance of CH3COOH (taken in a beaker) is determined at regular

additions of NH4OH, from a burette.

CH3COOH + NH4OH → CH3COONH4 + H2O (weak)

The conductance of the reaction mixture (CH3COOH) increases gradually with the

addition of the weak base, ammonium hydroxide, because of the formation of the salt

ammonium acetate. Immediately after the end point, the conductance remains almost

constant, because the titrant, ammonium hydroxide is a weake electrolyte. A plot of the

conductance of the reaction mixture against the volume of the titrant gives first a sloping

line and a line parallel to one of the co-ordinate axes. The point of intersection of the

straight lines is the end point. This method is quite suitable when there is no proper

indicator available, particularly this case.

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Lecturer Session No: 08 – Topic: Problems on measurement of electrode and cell

potentials using Nernst equation

26


Lecturer Session No: 09 – Topic: Revision

TEXT BOOKS FOR THIS UNIT. 1. P.C.Jain and Monica Jain, “Engineering Chemistry” Dhanpat Rai Pub, Co., New Delhi

(2002). 2. S.S.Dara “A text book of Engineering Chemistry” S.Chand & Co.Ltd., New Delhi

(2006).

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REFERENCES BOOKS FOR THIS UNIT.

:

1. B.Sivasankar “Engineering Chemistry” Tata McGraw-Hill Pub.Co.Ltd, New Delhi (2008). 2. B.K.Sharma “Engineering Chemistry” Krishna Prakasan Media (P) Ltd., Meerut (2001).

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Electrochemistry Lecture Notes