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 Notes –Electrochemistry
Lecturer Session No: 02 – Topic: EMF – Definition, measurement of emf (Poggendorff’s method ) – Single electrode potential
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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,
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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-
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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 Notes –Electrochemistry
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)
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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.
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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 Notes –Electrochemistry
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 Notes –Electrochemistry
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 Notes –Electrochemistry
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 Notes –Electrochemistry
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
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
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. 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
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
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. 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
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
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. 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 Notes –Electrochemistry
Lecturer Session No: 08 – Topic: Problems on measurement of electrode and cell
potentials using Nernst equation
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Lecturer Notes –Electrochemistry
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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