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DEPARTMENT OF CHEMISTRY
WELCOMES YOUALL
22CYT12 & Chemistry for Computer Systems
2022R
Unit-I-Electrochemistry
Prepared By
Krishnaveni K
Assistant Professor
Department of Chemistry
Kongu Engineering
College, Perundurai,
Erode
Course Outcome: Apply the principle of
electrochemistry for various applications
APPLIED CHEMISTRY
• The development of science and technology has
been giving us a lot of benefits. The advanced
technology has often required the basic research.
• Applied Chemistry is the scientific field for
understanding basic chemical properties of
materials and for producing new materials with
well-controlled functions.
• Applied chemistry is increasingly important in
solving environmental problems and contributing
to the development of new materials, both of which
are key issues in the 21st century.
ELECTROCHEMISTRY
22-Feb-24
Introduction – cells – types - representation of galvanic cell - electrode
potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from
single electrode potential - reference electrode: construction, working and
applications of standard hydrogen electrode, standard calomel electrode - glass
electrode – EMF series and its applications - potentiometric titrations (redox) -
conductometric titrations - mixture of weak and strong acid vs strong base.
UNIT-II
ELECTROCHEMISTRY
History of Electrochemistry
 16 th Century - William Gilbert –Father of Magnetism
 18 th Century – William Nicholson & Wilhelm Ritter – Decomposition of water – Electrolysis
 Svante Arrhenius - Dissociation of electrolytes
 Walther Hermann Nernst – Theory of Electromotive Force
 Conductance? Ability to conduct current , mho
ELECTROCHEMISTRY
INTRODUCTION
 It is a branch of chemistry
 The study of process involved in the interconversion of
chemical and electrical energy.
KEY TERMS IN ELECTROCHEMISTRY
 Conductor: Material which conduct electric current
 Non conductor: Material which do not conduct electric current
 Current: The flow of electrons through a wire or any conductor
 Oxidation: Loss of electrons
 Reduction: Gain of electrons
 Redox reaction: oxidation and reduction reactions occur simultaneously
 Reducing agent: A reactant in which donates an electron to the reduced species. (The reducing agent
is oxidized)
22-Feb-24
 Oxidizing agent: A reactant in which accepts an electron from the oxidized species. (The oxidizing agent
is reduced)
 Anode: The electrode at which oxidation occurs
 Cathode: The electrode at which reduction occurs
 Electrolyte: A water soluble substance and conduct an electric current
 Half cell: A single electrode immersed in an electrolytic solution and developing a definite potential
difference.
 Cell: Two half cells are connected through one wire
 Oxidation Potential : It is the tendency of an electrode to loss electrons
 Reduction potential: It is the tendency of an electrode to gain electrons
 Electrode Potential: It is the tendency of an electrode to loss or gain electrons
 Single Electrode Potential: It is the tendency of an electrode to loss or gain electrons when it is dipped in
its own salt solution. (Standard- 1M concentration at 250C).
22-Feb-24
22-Feb-24
LEOGER BOARD
 NaCl  Na+ + Cl-
 Anode : Cl-  Cl2 + 2e-
 Cathode  2Na+ + 2e-  2Na
 Over all reaction : 2Cl- + 2Na+  Cl2 + 2Na  2NaCl
ELECTROCHEMICAL CELL
Introduction
An electrochemical cell is a device in
which a redox reaction is utilized to get
electrical energy.
An electrochemical cell is also commonly
referred to as voltaic or galvanic cell.
The electrode where reduction occurs is
called cathode.
The electrode where oxidation occurs is
called anode.
22-Feb-24
Construction
 Electrochemical Cells are made up of two half-cells, each consisting of an electrode
which is dipped in an electrolyte. The same electrolyte can be used for both half cells.
These half cells are connected by a salt bridge which provides the platform for ionic
contact between them. A salt bridge minimizes or eliminates the liquid junction
potential.
 The practical application of an electrochemical or galvanic cell is the Daniel cell.
 It consists of a Zn electrode dipping in ZnSO4 solution and a Cu electrode dipping in
CuSO4 solution.
EMF= Eoxi + E Red
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Cell reaction
 Anode : Zn → Zn2+ + 2e- (Oxidation) {0.76V}
 Cathode : Cu2+ + 2e- → Cu (reduction) {0.34V}
 Overall : Zn + Cu2+ → Zn2+ + Cu (Redox)
 Representation of Daniel cell : Zn / Zn2+ || Cu2+ / Cu
 Zn / ZnSO4 (1M) // CuSO4 (1M) / Cu
 Cell EMF : 1.1 V
 EMF= Eoxi + E Red
= EZn + Ecu = 0.76+0.34
 CuSO4 - Cu2+ + SO4
2-
Anode : Zn Cathode : Cu
Zn Zn2+ ZnSO4 CuSO4 Cu2+ Cu
, / ;
Zn / Zn2+ Cu2+ / Cu
Zn / ZnSO4 (1M) CuSO4 (1M) / Cu
Zn , Zn2+ Cu2+ , Cu
Zn / Zn2+ || or // Cu2+ / Cu
Electrolytic cells
 Electrical Energy -- Chemical Energy
 Anode  positive Charge - oxidation --- 2Cl-  Cl2 + 2e-
 Cathode  negative charge  reduction --- 2Na+ + 2e-  Na
 Overall reaction --- 2Na+ + 2Cl-  2NaCl
Electrochemical Series
 The standard electrode potentials of a number of electrodes are arranged in the
increasing order of reduction potential at 25°C is referred to as emf or electrochemical
series.
Characteristics of electrochemical series:
 Lithium is the first member of the series.
 Highly reactive metal systems are at the top of the series.
 In other words, good reducing agents are at the top of the series, having the negative sign and act as
anode.
 All good oxidizing agents are at the bottom of the series , having the positive sign and act as cathode.
 Hydrogen system is at the middle of the series. All the elements which displace hydrogen from dilute
acids are placed above it.
Applications of Electrochemical Series
 To Find Reactivity of Metals
 As we move down in the electrochemical series reactivity of metal
decreases
 Alkali metals and alkaline earth metals at the top are highly reactive.
They can react with cold water and evolve hydrogen. They dissolve in
acids forming salts.
 Metals like Fe, Pb, Sn, Ni and Co which lie a little down in the series,
do not react with cold water but react with steam and evolve hydrogen.
 Metals like Cu, Ag and Au which lie below the hydrogen are less
reactive and do not evolve hydrogen from water.
Calculation of standard EMF of the cell
 EMF= Eoxi + E Red
 Zn & Cu Couple
 EMF= Eoxi + E Red
= EZn + E Cu
= 0.76+ + 0.34
= 1.1V
 Fe & H2
 EMF= Eoxi + E Red
 EMF= EFe + E H2
 = 0.441+ 0
0.441V
 Ni & Hg Couple
 Ni – Anode
 Hg - Cathode
 EMF= Eoxi + E Red
 = ENi + E Hg
= 0.236 + 0.61= 0.846V
 EMF = Standard reduction potential of R.H.S electrode- Standard reduction potential of L.H.S
electrode
 E0 = E0
RHS - E0
LHS
 = E0
Hg- E0
Ni
 = 0.61 – (-0.236)
 = 0.61+0.236 = 0.846V
= ENi + E Hg
= 0.236+0.61
= 0.846V
Cr & Sn Couple
Cr – Anode
Sn - Cathode
EMF= Eoxi + E Red
EMF= ECr + E Sn
= -0.74+(-0.14)
= 0.60V
EMF = Standard reduction potential of R.H.S electrode-
Standard reduction potential of L.H.S electrode
E0 = E0
RHS - E0
LHS
= E0
Sn - E0
Cr
= – 0.14 -(-0.74)
= -0.14+0.74 = 0.60V
 Zn + CuSO4  ZnSO4 + Cu
 Cu + ZnSO4  No reaction
 Zn + H2SO4  ZnSO4 + H2
 Ag + H2SO4  no reaction
For Studying displacement reaction
 Elements having higher reduction potential will gain electrons and that having lower
reduction potential will lose electrons. Hence element higher in electrochemical series
can displace an element placed lower in electrochemical series from its salt solution.
Example
Can zinc displaces copper from its salt solution?
Zn displaces Cu from CuSO4, because, zinc is placed higher in electrochemical series
while Cu is placed lower in electrochemical series. Hence zinc can easily displace
copper from CuSO4.
Zn+CuSO4 --------> ZnSO4 + Cu
For choosing elements as Oxidizing Agents
 The elements which have more electron-accepting tendency are oxidizing agents. The
strength of an oxidizing agent increases as the value of reduction potential becomes more
and more positive. Elements at the bottom of the electrochemical series have higher (+ve)
reduction potential. So they are good oxidizing agents. Thus, oxidizing power increases
from top to bottom in the series.
Example- F2 is a stronger oxidant than Cl2, Br2 and I2.
Cl2 is a stronger oxidant than Br2 and I2.
For choosing elements as Reducing Agents
The elements which have more electron losing tendency are reducing agents. The
power of reducing agent increases as the value of reduction potential becomes more and
more negative. Elements at the top of the electrochemical series have higher (-ve)
reduction potential. So they are good reducing agents. Thus, reducing power decreases
from top to bottom in the series.
Example-
The element like Zn, K, Na, Fe, etc. are good reducing agent.
Displacement of hydrogen from dilute acids by metals
 The metal which can provide electrons to H+ ions present in dilute acids for reduction evolve hydrogen
from dilute acids. The metal having negative values of reduction potential possesses the property of
losing an electron or electrons.
 Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from
dilute acids and on descending in the series, tendency to liberate hydrogen gas from dilute acids
decreases.
 The metals which are below hydrogen in the electrochemical series like Cu, Hg, Au and Pt do not evolve
hydrogen from dilute acids.
Example
Zinc reacts with dil.H2SO4 to give H2 but Ag does not. Why?
Zn+H2SO4 --------> ZnSO4 + H2 ; E0
Zn = -0.76 volts
Ag+H2SO4 --------> No reaction; E0
Ag = +0.80 volts
The metal with a positive reduction potential will not displace hydrogen from an acid solution.
 Displacement of hydrogen from water
 Iron and the metals above iron are capable of liberating hydrogen from water. The tendency
decreases from top to bottom in the electrochemical series.
 Alkali metals and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe
liberate hydrogen from hot water or steam.
 For Calculation of Standard emf of the cell
Standard reduction potential values are given in emf series. From the values E0
cell is calculated
using formula
E0
cell or standard emf of a cell = E0
oxi(cathode) - E0
red(anode)
For predicting spontaneity of the cell reaction
E0
cell > 0 cell reaction is spontaneous
E0
cell < 0 cell reaction is non-spontaneous
E0
cell = 0 cell reaction is in equilibrium
For determination of equilibrium constant for a reaction
We know that
-∆G0 = RTlnK
= 2.303RT logK
log K =
log K = (-∆G0 = nFE0)
Thus, from the value of E0 for a cell reaction, its equilibrium constant can be calculated.
CrCr3+ + 3e- (0.74 v)
 -∆G0 = RTlnK
 -∆G0 = nFE0
 n=3, F=96500C, E0 = 0.74V, R=8.314J/K/mol, T=298K
 Log K = -∆G0 /2.303RT
 = nFE0/2.303*8.314*298
=3*96500*0.74/5705.8
= 36.5
ZnZn2+ + 2e- (0.76V)
 N=2, F=96500C, E0 = 0.76V, R=8.314J/K/mol, T=298K
 Log K = -∆G0 /2.303RT
 = nFE0/2.303*8.314*298
=2*96500*0.76/5705.8
= 25.71
Reference electrode
 The electrode of standard potential with which we can compare the potentials of other
electrode is called a reference electrode.
 It can acts both as anode or cathode depending upon the nature of other electrode.
Classification:
i) Primary reference electrodes Ex : Standard Hydrogen Electrode (SHE)
ii) Secondary reference electrodes Ex: Calomel, Ag/AgCl electrodes and
Quinehydrone electrodes
v
Reference
Electrode
Working or Indicator
Electrode
The part of
the cell that
is kept
constant
The part of
the cell that
contains the
unknown
solution
Construction and Working of Standard Calomel Electrode (SCE)
 A common reference electrode.
 It consists of a wide glass tube.
 Mercury is placed at the bottom of the
glass tube.
 A paste of mercury and mercurous
chloride(Calomel) is placed above the
mercury. The remaining portion above the
paste is filled with a KCl solution of
known concentration (0.1N, 1.0N and
saturated) .
 A platinum wire is immersed into the
mercury to obtain electrical contact.
 The side arm is provided for making
electrical contact through a salt bridge.
Pt wire
Electrode representation:
Hg, Hg2Cl2(s)// KCl(satd. solution)
Working of the electrode:
If it acts as Cathode :
Hg2Cl2 Hg2
2+ +2Cl-
Hg2
2+ +2e- 2Hg
Hg2Cl2+2e- 2Hg+2Cl-
if it acts as anode :
2Hg Hg2
2+ +2e-
Hg2
2+ +2Cl-
2Hg+2Cl- Hg2Cl2+2e-
Hg2Cl2
KCl E in Volts
saturated 0.2422V
1.0N 0.2800 V
0.1N 0.3338V
Electrode potential
Measurement of pH using Calomel electrode
Measurement of pH using Calomel electrode
Hydrogen electrode containing a solution of unknown pH combine with the
calomel electrode to set up a complete cell.
We use a saturated calomel electrode as the reference and the complete cell can be represented as :
Pt, H2 (1atm)/ H+ (C=?)/ /KCl(satd. solution)/Hg2Cl2(s), Hg
Merits
 It is easy to construct and easy to transport.
 It provides almost a constant potential value with varying temperature and finds
application in laboratories for measuring potential of an electrode.
 It is used in corrosion studies.
Ion selective electrode- Glass electrode
 In Ion selective electrode, a membrane is in contact with a solution, with which it
can exchange ions. This ISE is responsive towards H+ and extensively used to
measure the pH of solution.
 In 1906, cremer found that a thin bulb of glass conducted electricity when he put
two solutions of different acid strengths inside and outside the bulb.
 The potential developed at the glass was in accordance with Nernst equation.
 Ion selective electrode is one which selectively responds to a specific ion in a
mixture and the potential developed at the electrode is a function of the
concentration of that ion in solution
Example- Glass electrode
Ion selective electrode- Glass electrode
Construction:
It consists of a thick walled glass tube with a very thin glass bulb placed at the bottom.
The thickness of the bulb is 0.01-0.03mm. Glass have high electrical resistivity.
 In glass electrode potential depends upon the pH of the medium
 The glass electrode consists of a glass bulb made up of special type of glass called Corning-
015 contains Na2O(22%), CaO(6%) and SiO2(72%) with high electrical conductance and
high hygroscopic in nature..
 The mixtures of the oxides is melted and cooled to form the glass. By altering the
composition of the glass, it is possible to make the electrode selective for different ions.
 The glass bulb is filled with a solution of constant pH(0.1M HCl) and insert with a Ag-
AgCl electrode, which is the Internal reference electrode and also serves for the external
electrical contact.
Thin
walled
glass bulb
0.1M HCl
AgCl
coated Ag
wire
Glass electrode
Electrode representation
Ag-AgCl /(0.1M) HCl/ Glass
Working:
The glass electrode works on the principle that when a thin glass membrane is placed between two
different concentration of a solution, a potential is developed at layers of the glass membrane. This
potential arises due to difference in the concentration of H+ ion inside and outside the membrane.
The potential developed is known as glass electrode potential EG and can be expressed as
H+ + e-  1/2H2
EG = E0
G --- 0.0591 Log[ H+]
--------
n
EG = E0
G - 0.0591 pH
potentiometer
 The glass electrode is placed in the solution under test and coupled with a saturated calomel electrode.
Cell representation
Ag-AgCl /(0.1M) HCl/ Glass/ solution of unknown pH//saturated calomel electrode
 The EMF of the cell is determined experimentally. From the emf, pH of the solution is calculated as
follows.
E cell = Ecalomel- Eglass
E cell = 0.2422 - (E0
G + 0.0591 pH)
E cell = 0.2422- E0
G -0.0591 pH
pH =
The value of E0
G can be determined by using a solution of known pH.
0.2422- Ecell -E0
G
0.0591
Advantages of Glass electrode
 It is very easy to construct and simple to operate.
 The potential developed remains constant for long time.
 This electrode can be used with very small amount of the test solution.
 This electrode can be used even in the presence of oxidized impurities, reducing impurities ,poison
molecules etc.,
 It can be used in turbid coloured and colloidal solutions.
Limitations of Glass electrode
 Since the glass membrane offers very high resistance, ordinary potentiometer cannot be used. It is
necessary to use electronic potentiometers.
 This electrode cannot be used to determine the pH above 12.
Potentiometric Titration
22-Feb-24
 It is similar to direct volumetric titration.
 Instead of indicator, potential is measured across the analyte
 Two electrodes are used – an indicator electrode and reference electrode
 Since the potential of reference electrode is constant and with the potential of indicator electrode, the
concentration of ion in the analyte can be measured.
 Ecell is recorded at intervals as the titrant is added.
 A graph of potential against volume added can be drawn and the end point of the reaction is halfway
between the jump in voltage.
 Ecell depends on the concentration of the interested ions with which the indicator electrode is in contact.
 For example, the electrode reaction may be
 Mn++ ne−-----> M
 As the concentration of Mn+ changes, the Ecell changes correspondingly. Thus the potentiometric
titration involve measurement of Ecell with the addition of titrant.
Theory
 Potential of an Electrode dipping in solution of eletrolyte depends upon the concentration of active ions.
E= E⁰ + (RT/nF) log C
 Small Change in active ion concentration in the solution changes the electrode potential correspondingly
 Concentration of Active ion decreases electrode potential of indicator electrode decreases
 The potential of Indicator electrode is measured potentiometrically by connecting with a reference electrode
(Saturated Calomel Electrode)
22-Feb-24
Determination of End point
 The emf of a cell changes by the addition of a small amount of titrant. So concentration
of reversible ion in contact with indicator electrode changes.
 Record the change in emf with every small addition
 The changes of potential will be slow at first, but at equivalence, the point change will
be sharp
 The values are plotted against corresponding volume changes.
 Change in emf with addition of titrant (⧍E/⧍V) is plotted against volume (V)
 The maximum of the curve gives the end point.
22-Feb-24
 Fig (a)– Volume of Titrant Vs Emf
 Fig(b) Volume of titrant Vs (⧍E/⧍V)
22-Feb-24
Classification of Potentiometric Titration
 Acid –Base titration
 Redox titration(Reduction- Oxidation)
 Precipitation titration
22-Feb-24
When it is titrated against K2Cr2O7 the following redox reaction takes place
Fe2+ is converted to Fe3+ and its concentration increases.
The potential is determined by the ratio of [Fe2+] / [Fe3+]
Till the end point, there is variation in potential with respect to the ratio of [Fe2+] / [Fe3+]
and after end point there’s no much variation in potential.
(Oxidised)
(reduced)
22-Feb-24
Advantages of Potentiometric Titrations
Potentiometric titrations can be carried out in colored solutions, where indicators
cannot be used
There is no need of prior information about the relative strength of titrant before the
titration
CONDUCTOMETRIC TITRATION
 Volumetric method based on the measurement of conductance of the solution during the titration
22-Feb-24
Conductance
Number and Charge on the free
ions
Mobility of the ions
Measurement of conductance using cells
 A conductance of the solution is measured using conductance cell.
 Which is made of a glass tube in which the platinized thin foils of
platinum electrodes are firmly fixed by sealing on a glass base.
 Polarization is removed when the electrodes are coated with finely
divided platinum black( chloroplatinic acid + lead acetate )
 The electrode are then washed repeatedly with distilled water and
finally with conductivity water.
 After usage the electrodes should be kept in conductivity water.
 The conductance of solution may be determined by measuring the
resistance of solutions into which a conductance cell is dipped.
 Conductance measurements are used extensively in chemistry and
in chemical industries. The use of the method is based on the
information from the behavior of electrolytes.
Process
 Taking a solution to be titrated in a beaker kept in a water
bath at a constant temperature.
 Conductivity cell is dipped and connected to a conductivity
bridge.
 The titrant is added from the burette(Fig)
 Conductance is measured each addition of solution.
 Recorded value is plotted the value of conductance against
the volume of the titrant.
 Since the conductance of solution is proportional to the
concentration of ions present, the conductance first
decreases with increase in volume of titrant, it reach the
saturation point it increase with the addition of titrant.
 From the graph end point is noted.
22-Feb-24
Procedure
 Calibrate the instrument by releasing the calibration knob
 Standard Sodium Hydroxide is taken in the burette
 The given acids is made upto 100ml in the standard measuring flask (SMF)
 20 ml of made up acids + 20 ml of conductivity water are added in 100 ml beaker
 Conductance is noted for addition of every addition of 1ml of Standard Sodium
Hydroxide
 Plot a graph between Volume of Standard Sodium Hydroxide Vs Conductance
 End Points are noted from the graph
 Equivalent Weight of Hydrochloric Acid = 36.5
 Equivalent Weight of Acetic Acid = 60
22-Feb-24
Types of Conductometric Titrations
 Acid –Base titration
 Strong Acid Vs Strong Base
Weak Acid Vs Strong Base
Mixture of Weak and Strong Acid Vs Strong Base
 Precipitation titration
 Replacement titration
 Redox titration
 Complexometric titration
22-Feb-24
Strong Acid Vs Strong Base(HCl Vs NaOH)
 Solution of electrolytes conducts electricity due to the presence of ions. The specific conductance of
solution is proportional to the concentration of ions in it. The reaction between HCl and NaOH may be
represented as
• H+ + OH------ H2O
 When a solution of hydrochloric acid is titrated with NaOH, the fast moving hydrogen ions are
progressively replaced by slow moving sodium ions. As a result conductance of the solution decreases.
This decrease in conductance will take place until the end point is reached. Further addition of alkali
raises the conductance sharply as there is an excess of hydroxide ions.
 A graph is drawn between volume of NaOH added and the conductance of solution. The exact end
point is the point of intersection of the two straight lines.
22-Feb-24
HCl + NaOH NaCl + H2O
22-Feb-24
HCl + NaOH NaCl + H2O
Weak Acid Vs Strong Base
( CH3COOH Vs NaOH)
22-Feb-24
CH3COOH+NaOH 3COO‾ + Na+ +H2O
H+ +OH---> H20
CH3COO- + Na-- CH3COONa
Mixture of Weak and Strong Acid Vs Strong Base
(HCl and CH3COOH Vs NaOH)
22-Feb-24
HCl + NaOH NaCl +H2O
CH3COOH+NaOH CH3COO
-Na+ +H2O
Advantages
 This method can be used with very diluted Solution
 This method can be used with Coloured and Turbid solution in which the end point
cannot be seen clearly
 This method can be used in which there is no suitable indicator
 Used for acid-base, redox, precipitation titration etc.,
22-Feb-24
REFERENCES:
 1.Palanisamy P.N., Manikandan P., Geetha A.& Manjula Rani K, “Applied
Chemistry”, 6th Edition, Tata McGraw Hill Education Private Limited, New
Delhi, 2019.
 2 .Paya Payal B.Joshi, Shashank Deep., “Engineering Chemistry”, Oxford
University Press, New Delhi, 2019.
 3.Palanna O., “Engineering Chemistry”, McGraw Hill Education, New
Delhi, 2017.
22-Feb-24
22-Feb-24
THANK
YOU
THANKYOU

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  • 1. DEPARTMENT OF CHEMISTRY WELCOMES YOUALL 22CYT12 & Chemistry for Computer Systems 2022R Unit-I-Electrochemistry Prepared By Krishnaveni K Assistant Professor Department of Chemistry Kongu Engineering College, Perundurai, Erode Course Outcome: Apply the principle of electrochemistry for various applications
  • 2. APPLIED CHEMISTRY • The development of science and technology has been giving us a lot of benefits. The advanced technology has often required the basic research. • Applied Chemistry is the scientific field for understanding basic chemical properties of materials and for producing new materials with well-controlled functions. • Applied chemistry is increasingly important in solving environmental problems and contributing to the development of new materials, both of which are key issues in the 21st century.
  • 3.
  • 5. Introduction – cells – types - representation of galvanic cell - electrode potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from single electrode potential - reference electrode: construction, working and applications of standard hydrogen electrode, standard calomel electrode - glass electrode – EMF series and its applications - potentiometric titrations (redox) - conductometric titrations - mixture of weak and strong acid vs strong base. UNIT-II ELECTROCHEMISTRY
  • 6. History of Electrochemistry  16 th Century - William Gilbert –Father of Magnetism  18 th Century – William Nicholson & Wilhelm Ritter – Decomposition of water – Electrolysis  Svante Arrhenius - Dissociation of electrolytes  Walther Hermann Nernst – Theory of Electromotive Force  Conductance? Ability to conduct current , mho
  • 7. ELECTROCHEMISTRY INTRODUCTION  It is a branch of chemistry  The study of process involved in the interconversion of chemical and electrical energy. KEY TERMS IN ELECTROCHEMISTRY  Conductor: Material which conduct electric current  Non conductor: Material which do not conduct electric current  Current: The flow of electrons through a wire or any conductor  Oxidation: Loss of electrons  Reduction: Gain of electrons  Redox reaction: oxidation and reduction reactions occur simultaneously  Reducing agent: A reactant in which donates an electron to the reduced species. (The reducing agent is oxidized) 22-Feb-24
  • 8.  Oxidizing agent: A reactant in which accepts an electron from the oxidized species. (The oxidizing agent is reduced)  Anode: The electrode at which oxidation occurs  Cathode: The electrode at which reduction occurs  Electrolyte: A water soluble substance and conduct an electric current  Half cell: A single electrode immersed in an electrolytic solution and developing a definite potential difference.  Cell: Two half cells are connected through one wire  Oxidation Potential : It is the tendency of an electrode to loss electrons  Reduction potential: It is the tendency of an electrode to gain electrons  Electrode Potential: It is the tendency of an electrode to loss or gain electrons  Single Electrode Potential: It is the tendency of an electrode to loss or gain electrons when it is dipped in its own salt solution. (Standard- 1M concentration at 250C). 22-Feb-24
  • 10.
  • 11.  NaCl  Na+ + Cl-  Anode : Cl-  Cl2 + 2e-  Cathode  2Na+ + 2e-  2Na  Over all reaction : 2Cl- + 2Na+  Cl2 + 2Na  2NaCl
  • 12. ELECTROCHEMICAL CELL Introduction An electrochemical cell is a device in which a redox reaction is utilized to get electrical energy. An electrochemical cell is also commonly referred to as voltaic or galvanic cell. The electrode where reduction occurs is called cathode. The electrode where oxidation occurs is called anode. 22-Feb-24
  • 13. Construction  Electrochemical Cells are made up of two half-cells, each consisting of an electrode which is dipped in an electrolyte. The same electrolyte can be used for both half cells. These half cells are connected by a salt bridge which provides the platform for ionic contact between them. A salt bridge minimizes or eliminates the liquid junction potential.  The practical application of an electrochemical or galvanic cell is the Daniel cell.  It consists of a Zn electrode dipping in ZnSO4 solution and a Cu electrode dipping in CuSO4 solution. EMF= Eoxi + E Red 22-Feb-24
  • 14. Cell reaction  Anode : Zn → Zn2+ + 2e- (Oxidation) {0.76V}  Cathode : Cu2+ + 2e- → Cu (reduction) {0.34V}  Overall : Zn + Cu2+ → Zn2+ + Cu (Redox)  Representation of Daniel cell : Zn / Zn2+ || Cu2+ / Cu  Zn / ZnSO4 (1M) // CuSO4 (1M) / Cu  Cell EMF : 1.1 V  EMF= Eoxi + E Red = EZn + Ecu = 0.76+0.34  CuSO4 - Cu2+ + SO4 2-
  • 15. Anode : Zn Cathode : Cu Zn Zn2+ ZnSO4 CuSO4 Cu2+ Cu , / ; Zn / Zn2+ Cu2+ / Cu Zn / ZnSO4 (1M) CuSO4 (1M) / Cu Zn , Zn2+ Cu2+ , Cu Zn / Zn2+ || or // Cu2+ / Cu
  • 16. Electrolytic cells  Electrical Energy -- Chemical Energy  Anode  positive Charge - oxidation --- 2Cl-  Cl2 + 2e-  Cathode  negative charge  reduction --- 2Na+ + 2e-  Na  Overall reaction --- 2Na+ + 2Cl-  2NaCl
  • 17. Electrochemical Series  The standard electrode potentials of a number of electrodes are arranged in the increasing order of reduction potential at 25°C is referred to as emf or electrochemical series. Characteristics of electrochemical series:  Lithium is the first member of the series.  Highly reactive metal systems are at the top of the series.  In other words, good reducing agents are at the top of the series, having the negative sign and act as anode.  All good oxidizing agents are at the bottom of the series , having the positive sign and act as cathode.  Hydrogen system is at the middle of the series. All the elements which displace hydrogen from dilute acids are placed above it.
  • 18.
  • 19.
  • 20. Applications of Electrochemical Series  To Find Reactivity of Metals  As we move down in the electrochemical series reactivity of metal decreases  Alkali metals and alkaline earth metals at the top are highly reactive. They can react with cold water and evolve hydrogen. They dissolve in acids forming salts.  Metals like Fe, Pb, Sn, Ni and Co which lie a little down in the series, do not react with cold water but react with steam and evolve hydrogen.  Metals like Cu, Ag and Au which lie below the hydrogen are less reactive and do not evolve hydrogen from water.
  • 21. Calculation of standard EMF of the cell  EMF= Eoxi + E Red  Zn & Cu Couple  EMF= Eoxi + E Red = EZn + E Cu = 0.76+ + 0.34 = 1.1V  Fe & H2  EMF= Eoxi + E Red  EMF= EFe + E H2  = 0.441+ 0 0.441V
  • 22.  Ni & Hg Couple  Ni – Anode  Hg - Cathode  EMF= Eoxi + E Red  = ENi + E Hg = 0.236 + 0.61= 0.846V  EMF = Standard reduction potential of R.H.S electrode- Standard reduction potential of L.H.S electrode  E0 = E0 RHS - E0 LHS  = E0 Hg- E0 Ni  = 0.61 – (-0.236)  = 0.61+0.236 = 0.846V = ENi + E Hg = 0.236+0.61 = 0.846V
  • 23. Cr & Sn Couple Cr – Anode Sn - Cathode EMF= Eoxi + E Red EMF= ECr + E Sn = -0.74+(-0.14) = 0.60V EMF = Standard reduction potential of R.H.S electrode- Standard reduction potential of L.H.S electrode E0 = E0 RHS - E0 LHS = E0 Sn - E0 Cr = – 0.14 -(-0.74) = -0.14+0.74 = 0.60V
  • 24.  Zn + CuSO4  ZnSO4 + Cu  Cu + ZnSO4  No reaction  Zn + H2SO4  ZnSO4 + H2  Ag + H2SO4  no reaction
  • 25. For Studying displacement reaction  Elements having higher reduction potential will gain electrons and that having lower reduction potential will lose electrons. Hence element higher in electrochemical series can displace an element placed lower in electrochemical series from its salt solution. Example Can zinc displaces copper from its salt solution? Zn displaces Cu from CuSO4, because, zinc is placed higher in electrochemical series while Cu is placed lower in electrochemical series. Hence zinc can easily displace copper from CuSO4. Zn+CuSO4 --------> ZnSO4 + Cu
  • 26. For choosing elements as Oxidizing Agents  The elements which have more electron-accepting tendency are oxidizing agents. The strength of an oxidizing agent increases as the value of reduction potential becomes more and more positive. Elements at the bottom of the electrochemical series have higher (+ve) reduction potential. So they are good oxidizing agents. Thus, oxidizing power increases from top to bottom in the series. Example- F2 is a stronger oxidant than Cl2, Br2 and I2. Cl2 is a stronger oxidant than Br2 and I2.
  • 27. For choosing elements as Reducing Agents The elements which have more electron losing tendency are reducing agents. The power of reducing agent increases as the value of reduction potential becomes more and more negative. Elements at the top of the electrochemical series have higher (-ve) reduction potential. So they are good reducing agents. Thus, reducing power decreases from top to bottom in the series. Example- The element like Zn, K, Na, Fe, etc. are good reducing agent.
  • 28. Displacement of hydrogen from dilute acids by metals  The metal which can provide electrons to H+ ions present in dilute acids for reduction evolve hydrogen from dilute acids. The metal having negative values of reduction potential possesses the property of losing an electron or electrons.  Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from dilute acids and on descending in the series, tendency to liberate hydrogen gas from dilute acids decreases.  The metals which are below hydrogen in the electrochemical series like Cu, Hg, Au and Pt do not evolve hydrogen from dilute acids. Example Zinc reacts with dil.H2SO4 to give H2 but Ag does not. Why? Zn+H2SO4 --------> ZnSO4 + H2 ; E0 Zn = -0.76 volts Ag+H2SO4 --------> No reaction; E0 Ag = +0.80 volts The metal with a positive reduction potential will not displace hydrogen from an acid solution.
  • 29.  Displacement of hydrogen from water  Iron and the metals above iron are capable of liberating hydrogen from water. The tendency decreases from top to bottom in the electrochemical series.  Alkali metals and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe liberate hydrogen from hot water or steam.  For Calculation of Standard emf of the cell Standard reduction potential values are given in emf series. From the values E0 cell is calculated using formula E0 cell or standard emf of a cell = E0 oxi(cathode) - E0 red(anode)
  • 30. For predicting spontaneity of the cell reaction E0 cell > 0 cell reaction is spontaneous E0 cell < 0 cell reaction is non-spontaneous E0 cell = 0 cell reaction is in equilibrium For determination of equilibrium constant for a reaction We know that -∆G0 = RTlnK = 2.303RT logK log K = log K = (-∆G0 = nFE0) Thus, from the value of E0 for a cell reaction, its equilibrium constant can be calculated.
  • 31. CrCr3+ + 3e- (0.74 v)  -∆G0 = RTlnK  -∆G0 = nFE0  n=3, F=96500C, E0 = 0.74V, R=8.314J/K/mol, T=298K  Log K = -∆G0 /2.303RT  = nFE0/2.303*8.314*298 =3*96500*0.74/5705.8 = 36.5 ZnZn2+ + 2e- (0.76V)  N=2, F=96500C, E0 = 0.76V, R=8.314J/K/mol, T=298K  Log K = -∆G0 /2.303RT  = nFE0/2.303*8.314*298 =2*96500*0.76/5705.8 = 25.71
  • 32. Reference electrode  The electrode of standard potential with which we can compare the potentials of other electrode is called a reference electrode.  It can acts both as anode or cathode depending upon the nature of other electrode. Classification: i) Primary reference electrodes Ex : Standard Hydrogen Electrode (SHE) ii) Secondary reference electrodes Ex: Calomel, Ag/AgCl electrodes and Quinehydrone electrodes v Reference Electrode Working or Indicator Electrode The part of the cell that is kept constant The part of the cell that contains the unknown solution
  • 33. Construction and Working of Standard Calomel Electrode (SCE)  A common reference electrode.  It consists of a wide glass tube.  Mercury is placed at the bottom of the glass tube.  A paste of mercury and mercurous chloride(Calomel) is placed above the mercury. The remaining portion above the paste is filled with a KCl solution of known concentration (0.1N, 1.0N and saturated) .  A platinum wire is immersed into the mercury to obtain electrical contact.  The side arm is provided for making electrical contact through a salt bridge. Pt wire
  • 34. Electrode representation: Hg, Hg2Cl2(s)// KCl(satd. solution) Working of the electrode: If it acts as Cathode : Hg2Cl2 Hg2 2+ +2Cl- Hg2 2+ +2e- 2Hg Hg2Cl2+2e- 2Hg+2Cl- if it acts as anode : 2Hg Hg2 2+ +2e- Hg2 2+ +2Cl- 2Hg+2Cl- Hg2Cl2+2e- Hg2Cl2 KCl E in Volts saturated 0.2422V 1.0N 0.2800 V 0.1N 0.3338V Electrode potential
  • 35. Measurement of pH using Calomel electrode
  • 36. Measurement of pH using Calomel electrode Hydrogen electrode containing a solution of unknown pH combine with the calomel electrode to set up a complete cell. We use a saturated calomel electrode as the reference and the complete cell can be represented as : Pt, H2 (1atm)/ H+ (C=?)/ /KCl(satd. solution)/Hg2Cl2(s), Hg
  • 37.
  • 38. Merits  It is easy to construct and easy to transport.  It provides almost a constant potential value with varying temperature and finds application in laboratories for measuring potential of an electrode.  It is used in corrosion studies.
  • 39. Ion selective electrode- Glass electrode  In Ion selective electrode, a membrane is in contact with a solution, with which it can exchange ions. This ISE is responsive towards H+ and extensively used to measure the pH of solution.  In 1906, cremer found that a thin bulb of glass conducted electricity when he put two solutions of different acid strengths inside and outside the bulb.  The potential developed at the glass was in accordance with Nernst equation.  Ion selective electrode is one which selectively responds to a specific ion in a mixture and the potential developed at the electrode is a function of the concentration of that ion in solution Example- Glass electrode
  • 40. Ion selective electrode- Glass electrode Construction: It consists of a thick walled glass tube with a very thin glass bulb placed at the bottom. The thickness of the bulb is 0.01-0.03mm. Glass have high electrical resistivity.  In glass electrode potential depends upon the pH of the medium  The glass electrode consists of a glass bulb made up of special type of glass called Corning- 015 contains Na2O(22%), CaO(6%) and SiO2(72%) with high electrical conductance and high hygroscopic in nature..  The mixtures of the oxides is melted and cooled to form the glass. By altering the composition of the glass, it is possible to make the electrode selective for different ions.  The glass bulb is filled with a solution of constant pH(0.1M HCl) and insert with a Ag- AgCl electrode, which is the Internal reference electrode and also serves for the external electrical contact. Thin walled glass bulb 0.1M HCl AgCl coated Ag wire Glass electrode
  • 41.
  • 42. Electrode representation Ag-AgCl /(0.1M) HCl/ Glass Working: The glass electrode works on the principle that when a thin glass membrane is placed between two different concentration of a solution, a potential is developed at layers of the glass membrane. This potential arises due to difference in the concentration of H+ ion inside and outside the membrane. The potential developed is known as glass electrode potential EG and can be expressed as H+ + e-  1/2H2 EG = E0 G --- 0.0591 Log[ H+] -------- n EG = E0 G - 0.0591 pH potentiometer
  • 43.
  • 44.
  • 45.  The glass electrode is placed in the solution under test and coupled with a saturated calomel electrode. Cell representation Ag-AgCl /(0.1M) HCl/ Glass/ solution of unknown pH//saturated calomel electrode  The EMF of the cell is determined experimentally. From the emf, pH of the solution is calculated as follows. E cell = Ecalomel- Eglass E cell = 0.2422 - (E0 G + 0.0591 pH) E cell = 0.2422- E0 G -0.0591 pH pH = The value of E0 G can be determined by using a solution of known pH. 0.2422- Ecell -E0 G 0.0591
  • 46. Advantages of Glass electrode  It is very easy to construct and simple to operate.  The potential developed remains constant for long time.  This electrode can be used with very small amount of the test solution.  This electrode can be used even in the presence of oxidized impurities, reducing impurities ,poison molecules etc.,  It can be used in turbid coloured and colloidal solutions. Limitations of Glass electrode  Since the glass membrane offers very high resistance, ordinary potentiometer cannot be used. It is necessary to use electronic potentiometers.  This electrode cannot be used to determine the pH above 12.
  • 48.  It is similar to direct volumetric titration.  Instead of indicator, potential is measured across the analyte  Two electrodes are used – an indicator electrode and reference electrode  Since the potential of reference electrode is constant and with the potential of indicator electrode, the concentration of ion in the analyte can be measured.  Ecell is recorded at intervals as the titrant is added.  A graph of potential against volume added can be drawn and the end point of the reaction is halfway between the jump in voltage.  Ecell depends on the concentration of the interested ions with which the indicator electrode is in contact.  For example, the electrode reaction may be  Mn++ ne−-----> M  As the concentration of Mn+ changes, the Ecell changes correspondingly. Thus the potentiometric titration involve measurement of Ecell with the addition of titrant.
  • 49. Theory  Potential of an Electrode dipping in solution of eletrolyte depends upon the concentration of active ions. E= E⁰ + (RT/nF) log C  Small Change in active ion concentration in the solution changes the electrode potential correspondingly  Concentration of Active ion decreases electrode potential of indicator electrode decreases  The potential of Indicator electrode is measured potentiometrically by connecting with a reference electrode (Saturated Calomel Electrode) 22-Feb-24
  • 50. Determination of End point  The emf of a cell changes by the addition of a small amount of titrant. So concentration of reversible ion in contact with indicator electrode changes.  Record the change in emf with every small addition  The changes of potential will be slow at first, but at equivalence, the point change will be sharp  The values are plotted against corresponding volume changes.  Change in emf with addition of titrant (⧍E/⧍V) is plotted against volume (V)  The maximum of the curve gives the end point. 22-Feb-24
  • 51.  Fig (a)– Volume of Titrant Vs Emf  Fig(b) Volume of titrant Vs (⧍E/⧍V) 22-Feb-24
  • 52. Classification of Potentiometric Titration  Acid –Base titration  Redox titration(Reduction- Oxidation)  Precipitation titration 22-Feb-24
  • 53. When it is titrated against K2Cr2O7 the following redox reaction takes place Fe2+ is converted to Fe3+ and its concentration increases. The potential is determined by the ratio of [Fe2+] / [Fe3+] Till the end point, there is variation in potential with respect to the ratio of [Fe2+] / [Fe3+] and after end point there’s no much variation in potential. (Oxidised) (reduced)
  • 54.
  • 55. 22-Feb-24 Advantages of Potentiometric Titrations Potentiometric titrations can be carried out in colored solutions, where indicators cannot be used There is no need of prior information about the relative strength of titrant before the titration
  • 56. CONDUCTOMETRIC TITRATION  Volumetric method based on the measurement of conductance of the solution during the titration 22-Feb-24 Conductance Number and Charge on the free ions Mobility of the ions
  • 57. Measurement of conductance using cells  A conductance of the solution is measured using conductance cell.  Which is made of a glass tube in which the platinized thin foils of platinum electrodes are firmly fixed by sealing on a glass base.  Polarization is removed when the electrodes are coated with finely divided platinum black( chloroplatinic acid + lead acetate )  The electrode are then washed repeatedly with distilled water and finally with conductivity water.  After usage the electrodes should be kept in conductivity water.  The conductance of solution may be determined by measuring the resistance of solutions into which a conductance cell is dipped.  Conductance measurements are used extensively in chemistry and in chemical industries. The use of the method is based on the information from the behavior of electrolytes.
  • 58. Process  Taking a solution to be titrated in a beaker kept in a water bath at a constant temperature.  Conductivity cell is dipped and connected to a conductivity bridge.  The titrant is added from the burette(Fig)  Conductance is measured each addition of solution.  Recorded value is plotted the value of conductance against the volume of the titrant.  Since the conductance of solution is proportional to the concentration of ions present, the conductance first decreases with increase in volume of titrant, it reach the saturation point it increase with the addition of titrant.  From the graph end point is noted. 22-Feb-24
  • 59. Procedure  Calibrate the instrument by releasing the calibration knob  Standard Sodium Hydroxide is taken in the burette  The given acids is made upto 100ml in the standard measuring flask (SMF)  20 ml of made up acids + 20 ml of conductivity water are added in 100 ml beaker  Conductance is noted for addition of every addition of 1ml of Standard Sodium Hydroxide  Plot a graph between Volume of Standard Sodium Hydroxide Vs Conductance  End Points are noted from the graph  Equivalent Weight of Hydrochloric Acid = 36.5  Equivalent Weight of Acetic Acid = 60 22-Feb-24
  • 60. Types of Conductometric Titrations  Acid –Base titration  Strong Acid Vs Strong Base Weak Acid Vs Strong Base Mixture of Weak and Strong Acid Vs Strong Base  Precipitation titration  Replacement titration  Redox titration  Complexometric titration 22-Feb-24
  • 61. Strong Acid Vs Strong Base(HCl Vs NaOH)  Solution of electrolytes conducts electricity due to the presence of ions. The specific conductance of solution is proportional to the concentration of ions in it. The reaction between HCl and NaOH may be represented as • H+ + OH------ H2O  When a solution of hydrochloric acid is titrated with NaOH, the fast moving hydrogen ions are progressively replaced by slow moving sodium ions. As a result conductance of the solution decreases. This decrease in conductance will take place until the end point is reached. Further addition of alkali raises the conductance sharply as there is an excess of hydroxide ions.  A graph is drawn between volume of NaOH added and the conductance of solution. The exact end point is the point of intersection of the two straight lines. 22-Feb-24 HCl + NaOH NaCl + H2O
  • 62. 22-Feb-24 HCl + NaOH NaCl + H2O
  • 63. Weak Acid Vs Strong Base ( CH3COOH Vs NaOH) 22-Feb-24 CH3COOH+NaOH 3COO‾ + Na+ +H2O H+ +OH---> H20 CH3COO- + Na-- CH3COONa
  • 64. Mixture of Weak and Strong Acid Vs Strong Base (HCl and CH3COOH Vs NaOH) 22-Feb-24 HCl + NaOH NaCl +H2O CH3COOH+NaOH CH3COO -Na+ +H2O
  • 65. Advantages  This method can be used with very diluted Solution  This method can be used with Coloured and Turbid solution in which the end point cannot be seen clearly  This method can be used in which there is no suitable indicator  Used for acid-base, redox, precipitation titration etc., 22-Feb-24
  • 66. REFERENCES:  1.Palanisamy P.N., Manikandan P., Geetha A.& Manjula Rani K, “Applied Chemistry”, 6th Edition, Tata McGraw Hill Education Private Limited, New Delhi, 2019.  2 .Paya Payal B.Joshi, Shashank Deep., “Engineering Chemistry”, Oxford University Press, New Delhi, 2019.  3.Palanna O., “Engineering Chemistry”, McGraw Hill Education, New Delhi, 2017. 22-Feb-24