Electrochemistry B Sc III-SEM-VI

Dr. Y. S. THAKARE
M.Sc. (CHE) Ph D, NET, SET
Assistant Professor in Chemistry,
Shri Shivaji Science College, Amravati
Email: yogitathakare_2007@rediffmail.com
UNIT- VI
PART-I
Electrochemistry

Electrochemistry: The Branch of physical chemistry which
deals with the study of interconversion of electrical energy
in to chemical energy or vice-versa
Electrolytic Cell: Conversion of electrical energy in to
chemical energy
Electrochemical Cell: Conversion of chemical energy into
electrical energy
Electrode – The metal rod dipped in its salt solution

Electrochemical Series: Standard electrodes are arranged
according to decreasing order of its standard reduction
potential with respect to standard reduction potential of
hydrogen (0.00V).

Oxidation half cell Reduction half cell
Left hand side electrode Right hand side electrode
Nernst Equation
𝐸 = 𝐸0
−
2.303𝑅𝑇
𝑛𝐹
𝑙𝑜𝑔
[𝑂𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛 ℎ𝑎𝑙𝑓 𝑐𝑒𝑙𝑙]
[𝑅𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛 ℎ𝑎𝑙𝑓 𝑐𝑒𝑙𝑙]
𝐸 = 𝐸0
−
0.0591
𝑛
𝑙𝑜𝑔
[𝑃𝑟𝑜𝑑𝑢𝑐𝑡]
[𝑅𝑒𝑎𝑐𝑡𝑎𝑛𝑡]
Ecu= +0.34
EZn= -0.76
Electrode| Electrolyte || Electrolyte | Electrode
Ecell = ER - EL

Electrode potential: Due to difference in rate of dissolution and
deposition of ions the potential develops across the metal solution interface
is known as electrode potential. It depends upon the concentration
(activity) of the ions in the solution.
Cathode Anode
Denoted by a positive sign since
electrons are consumed here
Denoted by a negative sign since
electrons are liberated here
A reduction reaction occurs in the
cathode of an electrochemical cell
An oxidation reaction occurs here
Electrons move into the cathode Electrons move out of the anode
The electrode whose potential is known or arbitrarily fixed is known as
reference electrode.
E.g. Hydrogen electrode, Sat Calomel Electrode
The electrode whose potential is to be determining by combined with
another electrode of known potential is known as indicator electrode.
E.g. Glass electrode, quinhydrone electrode.

Emf of Hydrogen electrode
Consider Nernst equation
𝑬 = 𝑬𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝒏
𝒍𝒐𝒈
[𝑷𝒓𝒐𝒅𝒖𝒄𝒕]
[𝑹𝒆𝒂𝒄𝒕𝒂𝒏𝒕]
𝑬𝑯 = 𝑬𝑯
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝒍𝒐𝒈
[𝑷𝑯𝟐
]
[𝑯+]𝟐
{Since 𝑷𝑯𝟐
=1 and 𝑬𝑯
𝟎
=0}, we get
𝑬𝑯 = 𝟎 −
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝒍𝒐𝒈
𝟏
[𝑯+]𝟐
𝑬𝑯 = −
𝟎.𝟎𝟓𝟗𝟏
𝟐
{−𝟐 𝒍𝒐𝒈𝑯+
}
𝑬𝑯 = −
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝟐 {− 𝒍𝒐𝒈𝑯+
}
𝑬𝑯 = −𝟎. 𝟎𝟓𝟗𝟏pH
Representation Electrode reaction Emf of Cell pH
Pt(s) | H2(g),(1atm.) H+
(aq) (a1)
2𝐻+
+ 2𝑒−
→ 𝐻2 𝑔 (1𝑎𝑡𝑚) 𝐸𝐻 = −0.0591pH
𝑝𝐻 =
𝐸𝑐𝑒𝑙𝑙
0.0591
Hydrogen electrode is the primary standard reference electrode
Standard Hydrogen Electrode

Advantages of Standard Hydrogen Electrode (SHE)
 It can be used over the entire pH range.
 The electrode can be used as reference electrode for measuring the potential of
other electrodes.
 It does not exhibit salt error.
 It is highly accurate.
 It has low internal resistance and hence ordinary potentiometer can be used for
emf measurement.
Disadvantages of Standard Hydrogen Electrode (SHE)
 It is not possible to maintain the unit activity of H+ ions in the solutions
 It can not be used in a solution containing ions of metals that are below
hydrogen in the electrochemical series. Interaction with the hydrogen will occur
and the metal will be deposited on the electrode surface.
 It can not be used in presence of oxidizing and reducing agents.
 Platinum can be easily poisoned by the adsorbed impurities from the solution.
 Black coating deteriorates and hence it should be renewed from time to time.
 Adsorbed impurities reduces the life of SHE.
 It is not easy to get pure and dry hydrogen gas.
 It is difficult to maintain the pressure of hydrogen gas at a fixed value
(1 atm).
 It is very difficult to construct SHE.

Calomel Electrode
Emf of Hydrogen electrode
𝑬 = 𝑬𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝒏
𝒍𝒐𝒈
𝑬𝒄𝒂𝒍 = 𝑬𝒄𝒂𝒍
𝟎
−
𝟎. 𝟎𝟓𝟗𝟏
𝟏
𝒍𝒐𝒈[𝑪𝒍−
]
The emf of Calomel electrode is depend
upon concentration of Cl- ions.
From emf of cell it is observed that with
decreasing concentration Cl- ions the emf
of cell increases
Representation Electrode reaction Emf of Cell (Volt)
Pt, Hg(l) |Hg2Cl2(s) |KCl(aq)
1
2
𝐻𝑔2𝐶𝑙2 + 𝑒− → 𝐻𝑔(𝑙) + 𝐶𝑙−
SCE=0.2415
NCE=0.2800
DNCE=0.3338
Calomel electrode is the secondary reference electrode
SCE= Saturated Calomel electrode
NCE= Normal Calomel electrode
DNCE=Deci normal Calomel electrode

Glass Electrode
Ag | AgCl | 0.1M HCl| Glass | unknown sol (pH=?) 2𝐻+
+ 2𝑒−
→ 𝐻2 𝑔 (1𝑎𝑡𝑚) 𝐸𝐺 = 𝐸𝐺
0
− 0.0591 𝑝𝐻 𝑝𝐻 =
𝐸𝑐𝑒𝑙𝑙 − 0.2415 + 𝐸𝐺
0
0.0591
Emf of Glass electrode
𝑬 = 𝑬𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝒏
𝒍𝒐𝒈
𝑬𝑮 = 𝑬𝑮
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝒍𝒐𝒈
[𝑷𝑯𝟐
]
[𝑯+]𝟐
{Since 𝑷𝑯𝟐
=1}, we get
𝑬𝑮 = 𝑬𝑮
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝒍𝒐𝒈
𝟏
[𝑯+]𝟐
𝑬𝑮 = 𝑬𝑮
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
{−𝟐 𝒍𝒐𝒈𝑯+}
𝑬𝑮 = 𝑬𝑮
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝟐 {− 𝒍𝒐𝒈𝑯+}
But − 𝒍𝒐𝒈𝑯+=pH
∴ 𝑬𝑮= 𝑬𝑮
𝟎
− 𝟎. 𝟎𝟓𝟗𝟏pH
Glass electrode is the indicator electrode generally use to determine pH of solution
Introduced by F. Haber and K. Klemensiewicz

Advantages of glass electrode
1. It may be used in the presence of strong oxidizing,
reducing and alkaline solution.
2. It can be used for solution having pH values 2 to 10 with
some special glass values can be extended up to 12.
3. It is simple to operate
4. It is immune to poisoning
5. It has no salt or protein error and the equilibrium is
reached quickly.
6. It can be used in colored turbid and colloidal solution
also.
Disadvantages of glass electrode
1. In glass electrode the bulb is very fragile and therefore
has to be used with great care.
2. As the glass membrane has a very high electrical
resistance hence it cannot be used with ordinary
potentiometer .
3. It cannot be employed in pure ethyl alcohol, acetic acid
and gelatin.

Pt | H2Q, Q,
H+(unknown)
𝑸 + 𝟐𝑯+
+ 𝟐𝒆−
→ 𝑯𝟐𝑸 𝐸𝑄 = 0.6996 − 0.0591 𝑝𝐻
𝑝𝐻 =
0.4581 − 𝐸𝑐𝑒𝑙𝑙
0.0591
Quinhydrone Electrode Emf of Glass electrode
𝑬 = 𝑬𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝒏
𝒍𝒐𝒈
𝑬Q = 𝑬Q
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝒍𝒐𝒈
[𝐻2𝑄]
[𝑸][𝑯+]𝟐
{Since [𝐻2𝑄]=[Q]=1}, we get
𝑬Q = 𝑬Q
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝒍𝒐𝒈
𝟏
[𝑯+]𝟐
𝑬Q = 𝑬Q
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
{−𝟐 𝒍𝒐𝒈𝑯+
}
𝑬Q = 𝑬Q
𝟎
−
𝟎.𝟎𝟓𝟗𝟏
𝟐
𝟐 {− 𝒍𝒐𝒈𝑯+
}
But − 𝒍𝒐𝒈𝑯+
=pH and 𝑬Q
𝟎
=0.6996
∴ 𝑬𝑸= 𝟎. 𝟔𝟗𝟗𝟔 − 𝟎. 𝟎𝟓𝟗𝟏pH
Quinhydrone electrode is the indicator electrode
Introduced by Bill mann in 1921
Quinhydrone is the equimolar mixture of
Quinone (Q) and hydroquinone (𝐇𝟐𝐐)

Advantages of Quinhydrone electrode
 It is very simple to setup and gives accurate values of pH because
equilibrium is attained quickly.
 Small quantity of solution is sufficient for pH measurement.
 It gives accurate results even in presence of oxidizing ions below pH8 only.
 It can be used for measuring pH of solution containing Zn2+, Pb2+, Cu2+ etc.
where hydrogen electrode is unsuitable.
 Ordinary potentiometer can be used for emf measurement.
 It is stable for longer time as well as at higher temperature i.e. above
300C.
Disadvantages of Quinhydrone electrode
 It can not be used in alkaline solutions above pH = 8, because H2Q acts as a
weak dibasic acid in more alkaline solution.
 The potential of this electrode is affected by even small concentration of
neutral salts. This is known as salt error.
 The potential of electrode is affected by small oxidizing and reducing
agent. Hence it can not be used in solutions containing such substance.

DETERMINATION OF pH OF SOLUTION USING HYDROGEN ELECTRODE
Here
Indicator electrode –-Hydrogen electrode (acts as anode)
Reference electrode– Saturated calomel electrode (acts as cathode)
The cell is represented as
−(anode) + (cathode)
Pt | H2(1atm) | H+(c= unknown) || KCl sat. solution |Hg2Cl2(s) |Hg(l)
EL ER
EMF of the cell is given by
Ecell = ER – EL
= E calomel - E hydrogen
= 0.2415 – (-0.0591 pH)
= 0.2415+ 0.0591 pH
0.0591 pH = Ecell - 0.2415
OR 𝒑𝑯 =
𝑬𝒄𝒆𝒍𝒍−𝟎.𝟐𝟒𝟏𝟓
𝟎.𝟎𝟓𝟗𝟏

DETERMINATION OF pH OF SOLUTION USING GLASS ELECTRODE
Here
Indicator electrode –-Glass electrode (acts as anode)
Reference electrode– Saturated calomel electrode (acts as cathode)
− (anode) +(cathode)
Ag | AgCl | 0.1M HCl| Glass | unknown sol (pH=?)|| KCl sat. solution |Hg2Cl2(s) |Hg(l)
EL ER
Ecell = ER – EL
= E calomel - E Glass
= 0.2415 – (𝑬𝑮
𝟎
− 𝟎. 𝟎𝟓𝟗𝟏pH)
= 0.2415- 𝑬𝑮
𝟎
+ 𝟎. 𝟎𝟓𝟗𝟏pH
0.0591pH = 𝑬𝑮
𝟎
- 0.2415 + Ecell
OR 𝒑𝑯 =
𝑬𝒄𝒆𝒍𝒍−𝟎.𝟐𝟒𝟏𝟓+ 𝑬𝑮
𝟎
𝟎.𝟎𝟓𝟗𝟏

DETERMINATION OF pH OF SOLUTION USING QUINHYDRONE ELECTRODE
Here
Indicator electrode –-Glass electrode (acts as cathode)
Reference electrode– Saturated calomel electrode (acts as anode)
−(anode) +(cathode)
Hg(l) | Hg2Cl2(s) | KCl(sat.) || H+, Q, H2Q | Pt
EL ER
Ecell = EQ - Ecalomel
= 0.6996- 0.0591pH – 0.2415
Ecell = 0.4581- 0.0591pH
0.0591pH=0.4581- Ecell
OR 𝒑𝑯 =
𝟎.𝟒𝟓𝟖𝟏−𝑬𝒄𝒆𝒍𝒍
𝟎.𝟎𝟓𝟗𝟏

POTENTIOMETRIC TITRATION
The titration in which end point is determined by measuring potential of indicator
electrode is called as Potentiometric titration.
There is a sudden change in potential (emf) of indicator electrode near the equivalence
point (end point) of the titration.
 POTENTIAL OF A GALVANIC CELL IS DETERMINED AT VARIOUS POINTS DURING
THE TITRATION PROCESS
 POTENTIAL OF AN INDICATOR ELECTRODE DEPENDS ON THE CONCENTRATION OF
IONS TO WHICH IT IS REVERSIBLE
 In potentiometric titration the indicator electrode must be reversible w.r.to the
ions whose concentration alters during the course of reaction
 CONCENTRATION OF IONS CHANGES DURING A TITRATION THE ELECTRODE
POTENTIAL CHANGES
 EQUIVALENCE POINT IS DETERMINED GRAPHICALLY BY PLOTTING EMF VS
VOLUME OF TITRANT ADDED.
Types of potentiometric titration
• Acid-Base Titration
• Redox Titration
• Precipitation Titration

ACID BASE TITRATION
Potentiometric acid base titration can be performed by coupling an indicator
electrode which is reversible with respect to H+ ions with a reference electrode.
indicator electrode : Hydrogen electrode or Quinhydrone electrode or Glass
electrode
reference electrode : Saturated Calomel electrode
For example: HCL + NaOH → NaCl + H2O
a) Using a Hydrogen electrode-
− +
Pt | H2(1atm) | H+(c= unknown) || KCl sat. solution |Hg2Cl2(s) |Hg(l)
EL ER
Ecell = ER – EL
= E calomel - E hydrogen
= 0.2415 – (-0.0591 pH)
= 0.2415+ 0.0591 pH

b) Using a Quinhydrone electrode-
− +
Hg│Hg2Cl2(s), KCl (sat) ║ H+ (unknown conc.), Q, QH2│Pt
EL ER
Ecell = Eright – Eleft
Ecell = EQ - Ecalomel
= 0.6996- 0.0591pH – 0.2415
Ecell = 0.4581- 0.0591pH

c) Using Glass Electrode-
− +
Ag | AgCl | 0.1M HCl| Glass | unknown sol (pH=?)|| KCl sat. solution |Hg2Cl2(s) |Hg(l)
EL ER
Ecell = ER – EL
= E calomel - E Glass
= 0.2415 – (𝑬𝑮
𝟎
− 𝟎. 𝟎𝟓𝟗𝟏pH)
= 0.2415- 𝑬𝑮
𝟎
+ 𝟎. 𝟎𝟓𝟗𝟏pH

OXIDATION REDUCTION TITRATION
E.g. Titration of ferrous ammonium sulphate (FAS)solution against potassium
dichromate in acidic medium. Fe3+ / Fe2+ and Cr6+ / Cr3+
6 Fe2+ + Cr2O7
2- + 14 H+ → 6 Fe3+ + 2 Cr3+ + 7 H2O
Indicator electrode --- Redox electrode, Here Fe3+/Fe2+ (oxidation) and Cr6+ /Cr3+(reduction)
Reference electrode--- Calomel electrode (Ecal= + 0.2415)
The cell may be represented as
Ecalomel ║ EFe
3+
/ Fe
2+
− +
Hg(l) | Hg2Cl2(s) | KCl(sat.) ║ Fe2+, Fe3+| Pt
The EMF of cell is given by
E(cell) = ER - EL
𝐸𝑐𝑒𝑙𝑙 = 𝐸Fe
3+
/ Fe
2+ − 𝐸𝑐𝑎𝑙
= (𝐸Fe
3+
/ Fe
2+
0
−
0.0591
1
𝑙𝑜𝑔
[𝐹𝑒2+]
[𝐹𝑒3+]
) − 0.2415
𝐸𝑐𝑒𝑙𝑙 = 0.77 +
0.0591
1
𝑙𝑜𝑔
[𝐹𝑒3+]
[𝐹𝑒2+]
) − 0.2415
Fe3+ + e- → Fe2+
𝑬Fe
3+
/ Fe
2+ = (𝑬Fe
3+
/ Fe
2+
𝟎
−
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏 𝑭
log10
[𝑷𝒓𝒐𝒅𝒖𝒄
[𝑹𝒆𝒂𝒄𝒕𝒂𝒏
𝑬Fe
3+
/ Fe
2+ = (𝑬Fe
3+
/ Fe
2+
𝟎
−
𝟐.𝟑𝟎𝟑 𝑹 𝑻
𝒏 𝑭
log10
[𝑭𝒆𝟐+]
[𝑭𝒆𝟑+]
𝑬𝑭𝒆𝟑+
/𝑭𝒆𝟐+ = − 𝑬𝑭𝒆𝟐+
/𝑭𝒆𝟑+
(𝑬Fe
3+
/ Fe
2+
𝟎
= + 0.77 V)

PRECIPITATION TITRATIONS
Potentiometric precipitation titration may be explained by taking the example of
titration of silver nitrate solution against KCl solution.
The cell may be represented as.
− +
Hg(l) | Hg2Cl2(s) | KCl(sat.) ║ AgNO3 (aq) | Ag
EMF of this cell is given by
𝐸𝑐𝑒𝑙𝑙 = 𝐸𝐴𝑔+/𝐴𝑔 − 𝐸𝑐𝑎𝑙𝑜𝑚𝑒𝑙
= (𝐸𝐴𝑔+/𝐴𝑔
0
+ 0.0591log[𝐴𝑔+
] − 0.2415)
= (0.799 + 0.0591log[𝐴𝑔+
] − 0.2415)
𝐴𝑔+
+ 𝑒−
= 𝐴𝑔(𝑆)
𝑬Ag+/Ag
= 𝐸Ag+/Ag
0
−
𝟎.𝟎𝟓𝟗𝟏
𝒏
𝒍𝒐𝒈
= 𝐸Ag+/Ag
0
−
𝟎.𝟎𝟓𝟗𝟏
𝒏
𝒍𝒐𝒈
[𝐴𝑔(𝑆)]
[𝐴𝑔+]
𝐸Ag+/Ag
0
= +0.7990𝑉 and [𝐴𝑔(𝑆)] = 1
𝑬Ag+/Ag
= 𝟎. 𝟕𝟗𝟗𝟎 −
𝟎. 𝟎𝟓𝟗𝟏
𝒏
𝒍𝒐𝒈
𝟏
[𝐴𝑔+]
KCl Solution
Silver
Electrode
AgNO3
Solution

Advantages of potentiometric titration
 The method can be used for colored solutions and also for solutions
where precipitation occurs during the titrations.
 This method can be used for the titration of a mixture of two acids with
the same NaOH solution. In such a case, two different points of
inflection are obtained .
 This method can be used for titrating weak acid against weak base.
 The potentiometric titrations have been used extensively due to the
fact that apparatus used is not very expensive and freely available .
 This method can be used for the analysis of dilute solutions with high
degree of accuracy.
 This method is applicable for the estimation of a mixture of chloride,
bromide and iodide ions by titration with AgNO3 solution.
 Dissociation constants (pKa) of weak acids may be obtained by titrating
the weak acid with a strong base.
 It is used to determine hydrolysis constant
 The potentiometric titration is applicable to determine Hammet
constant

CONCENTRATION CELLS-
A cell in which emf is generated as a result of difference in concentration is called as
concentration cell. In such a cell, substance is transferred from a system of high
concentration to the one at low concentration to produce the electric current.
Types of concentration cells-
Concentration cells may be classified into two main types.
a) Amalgam cells or electrode concentration cells (first type)-
These are the cells in which the two electrodes of different concentrations are dipped
in the same solution of their salt.
b) Electrolyte concentration cells (second type)-
These are the cells in which the two electrodes are of the same material which are
dipped into two solutions of the same electrolyte but of different concentrations.
These cells are further sub classified into two types-
i) Concentration cells without transference (or transport)-
In these types of cells, there is no direct transfer of electrolyte from one solution to
another. The transfer occurs as a result of a chemical reaction
− +
Zn(s) | Zn SO4(aq) (C1) || Zn SO4(aq) (C2) | Zn(s)
ii) Concentration cells with transference (or transport)-
In these types of cells, the two electrolyte solutions are in direct contact with each
other and hence there is direct transfer of the electrolyte from a more concentrated
solution to less concentrated solution.
Pt(s)| H2(g) (1atm) | HCl (C1) ¦ HCl (C1) |H2(g) (1atm) | Pt(s)

CONCENTRATION CELLS
A cell in which emf is generated as a result of difference in
concentration is called as concentration cell.
− +
Zn(s) | Zn SO4(aq) (C1) || Zn SO4(aq) (C2) | Zn(s) 𝑪𝟐 > 𝑪𝟏
At left hand electrode,
Zn(s) → Zn2+
(aq) (C1) + 2e- … (i)
At right hand electrode,
Zn2+
(aq) (C2) + 2e- → Zn(s) … (ii)
The Net cell reaction
Zn2+
(aq) (C2) → Zn2+
(aq) (C1) … (iii)
Applying Nernst equation, emf of the cell is given by
𝐸𝑊𝑂𝑇 = 𝐸𝑐𝑒𝑙𝑙
0
−
0.0591
𝑛
𝑙𝑜𝑔
For concentration cell without transference 𝐸𝑐𝑒𝑙𝑙
0
= 0
𝐸𝑊𝑂𝑇 = −
0.0591
𝑛
𝑙𝑜𝑔
= −
0.0591
𝑛
𝑙𝑜𝑔
𝐶1
𝐶2
… (iv)
This equation gives emf of concentration cell without transference. In the above cell
substituting n=2 in equation (iv), we get,
𝐸𝑊𝑂𝑇 = −
0.0591
2
𝑙𝑜𝑔
𝐶1
𝐶2
∴ 𝐸𝑊𝑂𝑇 = 0.0295 𝑙𝑜𝑔
𝐶2
𝐶1

• Determination of Pka of weak
acid
Henderson’s equation
𝒑𝑯 = 𝒑𝑲𝒂 +
𝒍𝒐𝒈
[𝑺𝒂𝒍𝒕]
[𝑨𝒄𝒊𝒅]
At half the neutralization point,
[salt] =[acid] and hence pKa = pH.

Electrochemistry B Sc III-SEM-VI

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Electrochemistry B Sc III-SEM-VI