Electro Chemistry

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ELECTRO CHEMISTRY

Electrochemistry is the branch of science which deals with the relationship

between chemical reaction and electricity

Relating electricity and chemical reactions

ELECTRODE POTENTIAL

When a metal is placed in its own salt solution it may under go oxidation or

reduction according to its tendency to loose or gain electrons.

i) Oxidation: Loss of electron

M (s) ▬▬► M n+

(aq) + ne-

Metal behaves like an anode

For example:

When Mg electrode is dipped in MgSO4 solution, Mg goes into solution as

Mg2+

ions and Mg electrode attains negative charge due to oxidation.

The negative charge electrode attracts the positive ions from the solution and

forms a sort of layer of positive ions around the metal

Transfer of electrons

Galvanic Cell

In put: Chemical energy

Out put: Electrical energy

Electrolytic Cell

In put: Electrical Energy

Out put : Chemical reaction /energy

www.Vidyarthiplus.com


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ii) Reduction: Gain of electrons

M n+

(aq) + ne- ▬▬► M (s)

Metal behaves like a cathode

For example:

When Cu electrode is dipped in CuSO4 solution, Cu2+

ions from solution

deposits on the metal and Cu electrode attains a positive charge due to

reduction.

The positive charge electrode attracts the negative ions from the solution and

forms a sort of layer of negative ions around the metal.

The layer of positive / negative ions formed on the metal is called

Helmholtz Electrical Double Layer. A difference of potential is set up

between the metal ions and the solution.

At equilibrium, the potential difference becomes a constant value and is

called as electrode potential of the metal.

The tendency of the electrode to lose electrons is called oxidation potential

(EOP) and the tendency of the electrode to gain electrons is called reduction

potential (ERP).



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Single Electrode Potential (E)

It is a measure of the tendency of a metallic electrode to loose or gain

electrons when it is dipped in its own salt solution.

Standard Electrode Potential (Eo) (SEP)

It is a measure of the tendency of the metallic electrode to loose or gain

electrons when it is dipped in its own salt solution of unit concentration

(1M), at 25oC and atmospheric pressure.

Measurement of SEP

SEP cannot be measured directly. The electrode is coupled with a reference

electrode

Examples: Standard Hydrogen electrode (SHE)

Saturated Calomel Electrode (SCE)

REFERENCE ELECTRODES

The electrode of standard potential with which we can compare the

potentials of other electrodes is called a reference electrode.

The potential of the electrode remains constant at all temperatures.

It undergoes specific reduction or oxidation but the potential will be same

only sign will be different.



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Standard Hydrogen Electrode (SHE) /

Normal Hydrogen Electrode (NHE)

Type of electrode: Gas electrode (Primary Reference Electrode)

Components:

Electrode component: Pt-H2

Electrolyte component: HCl (1M)

Electrode representation:

Pt, H2 (1atm) / H+ (1M)

Construction

Hydrogen electrode consists of a Platinum foil connected to a platinum wire

sealed in a glass tube. The electrode is in contact with 1M HCl and hydrogen

gas (1 atmosphere) is constantly bubbled.

Limitations

• It requires pure hydrogen gas and is difficult to set up and to transport

• It requires large volume of test solution

• The potential of the electrode is dependent on atmospheric pressure

Reactions

As anode

H2 ▬▬► 2 H+ + 2e-

As cathode

2 H+ + 2e- ▬▬► H2

Eo = 0 V



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Saturated Calomel Electrode (SCE)

Type/ class: Metal- metal insoluble salt electrode (Secondary Reference

Electrode)

Components:

Electrode component: Pt – Hg

Electrolyte component: Hg2Cl2(s) / KCl

Electrode representation:

Hg, Hg2Cl2(s) - KCl (sat. solution)

Construction:

Calomel electrode consists of a glass tube containing mercury at the bottom

over which mercurous chloride paste (calomel) is placed. The tube is filled

with saturated KCl solution. A platinum wire is fused into the layer of

mercury to provide electrical contact. The electrode potential differs with the

concentration of KCl.

Reactions:

As anode

2Hg (l) ▬▬► Hg22+

+ 2e-

Hg22+

+ 2Cl- ▬▬►Hg2Cl2

As cathode

Hg22+

(2Cl -) + 2e- ▬▬►2Hg (l) +2 Cl

The net reaction can be represented as

2Hg + 2Cl- ↔↔↔↔ Hg2Cl2

E is 0.3335 V for 0.1N KCl (DNCE)

E is 0.2810 V for 1N KCl (NCE)

E is 0.2422V for Sat. KCl (SCE)



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Applying Nernst equation for the net reaction

Ecell = Eo

cell + 0.0591/n . log [Hg]2[Cl−]

2 / [Hg2Cl2]

The [Hg] and [Hg2Cl2] are unity .Therefore the Ecell depends on the

concentration of chloride ions , hence the electrode is said to be reversible

wrt chloride. (Another electrode which is reversible wrt to Chloride ions is

Ag/ AgCl)

The concentration of chloride ions ↓↓↓↓ ses during oxidation and ↑↑↑↑ses during

reduction.

Measurement of single electrode potential

The single / standard electrode potential can be measured by coupling the

electrode with a SHE. The E cell will be the E electrode as the ESHE is zero

For example the potential of Copper electrode can be measured by

constructing the following cell.

Electromotive Force (E.M.F.)

When two electrodes are connected, current starts flowing through the

circuit. The driving force which makes the electrons to flow from a region of

higher potential to a region of lower potential is called the electromotive

force abbreviated as emf. It is measured in Volts (V).



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The emf of the cell cannot be measured accurately using a voltmeter, as (i) a

part of the energy is utilized for its working and (ii) due to the polarization

effects the emf changes

.

Measurement Of EMF Of Cell By Poggendorff’s Compensation Principle

Refer class notes

Nernst Equation For Electrode Potential

Derivation Refer Class Notes

Expression for SEP

Nernst equation for a galvanic cell

Expression for concentration cell

E(elec)= Eo (elec)+ 2.303 RT log [R]/[P] nF

E(con.cell) = 2.303 RT log [a2] nF [a1]

Note a2>a1

E(G.cell) = Eo (cell)+ 2.303 RT log [Cathode] nF [Anode]



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ELECTROCHEMICAL SERIES (e.m.f series)

A series in which elements are arranged in the ascending (increasing ) order

Of their standard reduction potential is called emf series.

Half cell reaction Eo (V)

Li+ + e

- → Li - 3.04

Mg2+

+ 2e- -

→ Mg - 2.37

Al3+

+ 3e-

- →Al - 1.66

Zn2+

+ 2e- - → Zn - 0.76

Fe2+

+ 2e- - → Fe - 0.44

2H+ + 2e

- - → H2 (g) 0.00

Hg 22+

+ 2e▬ → Hg (l) 0.2422

Cu2+

+ 2e- - → Cu 0.34

Cu+ + e

-- → Cu 0.52

Pt,Fe3+

+ e▬ → Fe2+

0.77

Ag+ + e

-- → Ag 0.80

Au+ + e

- → Ag 1.69

F2 + 2e- - → 2F

- 2.87

Application / Significance of electrochemical series

(i) Relative ease of oxidation or reduction

• The metals which lie above hydrogen in the series undergo

spontaneous oxidation and the metals which lie below SHE undergo

reduction spontaneously ( ie. Acts as Anodes and Cathodes

respectively)

• The metals which lie above hydrogen are good reducing agents and

which lies below hydrogen will act as good oxidizing agents

(ii) Replacement tendency



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• The metal lying above in emf series displaces the metal lying below it

from an electrolyte of the later.

Example 1: Ni spatula cannot be used to stir copper sulphate solution

due to the following reaction

Ni(s) + Cu2+

(aq) ▬▬► Ni 2+

(aq) + Cu (s)

Example 2: when zinc is dipped in copper sulphate solution copper gets

deposited (displaced)

Zn (s) + CuSO4 (aq) → Zn SO4 (aq) + Cu (s)

(iii) Liberation of Hydrogen

• The metal with negative reduction potential will displace H2 from an

acid solution

Zn (s) + 2 HCl (aq) → Zn Cl2 (aq) + H2 ↑

Hence acids cannot be stored in galvanized steel containers.

For exactly the same reason galvanized steels are not used to store

food stuffs containing vinegar. (Vinegar is used as food preservative-

vinegar is acetic acid)

(iv) Calculation of equilibrium constant (Keq)

−∆GO

= n F Eo

−∆GO

= 2.303 RT log K(eq)

Therefore log K(eq) = n F Eo

2 .303 RT

log K(eq) = nEo

0.0591



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(v) Calculation of Standard emf of the cell

E cell = E cathode − E anode

( if both reduction potentials are considered)

E cell = E cathode + E anode

( if oxidation potential of anode and the reduction potential of cathode

are considered)

(vi) Corrosion

• The metals higher in the series are anodic and are more prone to

corrosion.

• The metals lower in the series are noble metals (cathodic) and they are

less prone to corrosion.

(vii) Predicting the spontaneity of cell reaction

• Spontaneity of the redox reaction can be predicted from the emf value

of complete cell reaction.

If the value of Ecell is positive, the reaction is feasible. as ∆G will be

negative ( i.e. it is an electrochemical cell)

If the value of Ecell is negative, the reaction is not feasible. as ∆G will be

positive ( i.e. it is an electrolytic cell)

Galvanic Series

In galvanic series, metals and alloys are arranged according to their

tendency to corrode. This series can be used to determine whether galvanic

corrosion is likely to occur and how strong the corrosion reaction will be.



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Importance of galvanic series

i) Provides more practical information on the relative corrosion

tendencies of different metals and alloys

ii) The rate and severity of corrosion depends on the potential

difference existing with in the system

Comparison of emf series with galvanic series

Electrochemical series

Galvanic series

Only Metals are arranged in

increasing order of electrode

potential

Metals and alloys are arranged in

increasing order of electrode

potential

The surface layer is removed before

dipping the metal in the electrolyte

( Ex.: Al2O3 is removed from the

surface of Al)

The metal along with its scale (

example Al2O3 on Al) is dipped in

the electrolyte

Electrolyte : Own salt solution

Electrolyte : Sea Water



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Reference electrode is SHE

Reference electrode is SCE

The position of a given metal in this

series is fixed

Position is not fixed, may shift

For example the alloy of Fe (MS and

SS) occupies different positions.

It gives relative displacement

tendencies

It gives the relative corrosion

tendencies

Cells

Cell is a simple unit comprising of an anode and a cathode dipped in an

electrolyte. Battery is an array of cells.



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Electrochemical Cell or Galvanic cell

It is a device in which a redox reaction is used to derive electrical energy.

During the working of the cell the stored chemical energy decreases and this

decrease is gained as electrical energy.

In the electrochemical cell the electrode at which oxidation occurs is called

anode (− ve) and the electrode at which reduction occurs is called cathode

(+ ve).

Example: Zn acts as anode and Cu acts as cathode in Daniel cell

Cells

Electrochemical cell

Ex.: Lechlanche Cell

Electrolytic cell

Ex. Electroplating

Reversible cells

/ Secondary cells.

Ex. : Daniel cell

Irreversible cells

/ Primary cells Ex.: Leclanche Cell

Functions due to

Potential gradient

Functions due to concentration gradient

Electrode cell

Ex. Amalgam

cell , Gas cells

Electrolyte cell

Ex. Silver ion

conc. cell



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It consists of zinc electrode dipped in 1M zinc sulphate solution and a

copper electrode dipped in 1M copper sulphate solution. Each electrode acts

as a half cell connected by a salt bridge through a voltmeter. The two

solutions can seep through the salt bridge without mixing.

At anode: Oxidation takes place

Zn (s) ▬▬► Zn 2+

+ 2e _

At cathode: Reduction takes place.

Cu 2+

(aq) + 2e _

▬▬► Cu (s)

Net reaction: Zn (s) + Cu

2+ (aq)

▬▬► Zn

2+ (aq) + Cu (s)

Representation of a galvanic cell

(i) Galvanic cell consists of two electrodes, anode and cathode

(ii) The anode is written on the left hand side while the cathode is written on

right side.

(iii) The anode is written with the metal first and then the electrolyte .The

two are separated by a vertical line or semicolon



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Ex: Zn / Zn 2+

(or) Zn / ZnSO4 (or) Zn; Zn 2+

(iv) The cathode is written with electrolyte first and then the metal both are

separated by vertical line or semicolon

Ex: Cu2+

/ Cu (or) CuSO4 / Cu (or) Cu ; Cu2+

(v) The two half cells are connected by a salt bridge which is indicated by

two parallel lines.

▬ +

Zn / ZnSO4 (1M ) ║ CuSO4 ( 1M ) / Cu

Salt bridge: It consists of a U tube filled with a saturated solution of KCl or

(NH4)2NO3 in agar-agar gel. It connects the two half cells and performs the

following functions

• It eliminates the liquid junction potential.

• It provides path for the flow of electrons between two half cells.

• Completes the circuit.

• Maintains electrical neutrality in the two compartments by migration

of ions through the porous material thus ensures the chemical

reactions proceed without hindrance

• Prevents mixing of the electrode solutions.

Reversible cells

A cell works reversibly in the thermodynamic conditions.

Ex. Daniel cell, Secondary batteries, Rechargeable batteries.

The cell is reversible if it satisfies all the following conditions:

(i) If applied emf is equal to derived emf then the net reaction is zero

(ii) If applied emf is infinitesimally smaller than the derived emf then the

cell should act as electrochemical cell (forward reaction)



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(iii) If applied emf is infinitesimally greater than the derived emf then the

cell should act as electrolytic cell (reverse reaction)

Irreversible cells

Cells which do not obey the (above) conditions of thermodynamic

reversibility are called irreversible cells. If one of the products escapes from

the cell then that cell cannot be made reversible by applying an external

current.

Ex. Zinc-Silver cell, Primary cells

Zn –Ag Cell

Zn / H2SO4 (aq) / Ag

Cell reaction:

Anode: Zn + H2SO4 ▬▬► ZnSO4 + H2 ↑

Cathode: 2Ag + + 2e

_ ▬▬► 2Ag

When two electrodes are connected from outside, zinc dissolves liberating

hydrogen gas. Since one of the product hydrogen escapes, the cell reaction

cannot be reversed when connected to an external EMF. The cell does not

obey the conditions of reversibility and is called irreversible cell.

Electrolytic cell

It is a device in which chemical reaction proceed at the expanse of electrical

energy.

Ex. Electro plating and electrolysis

Electrolysis of NaCl

The cell is constituted by dipping two platinum electrode in an appropriate

electrolyte ( NaCl in water ) . The electrodes are connected to the two

terminals of a battery. The electrode connected to positive terminal acts as



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anode (attracts anions) and the other electrode connected to the negative

terminal acts as cathode (attracts cations). Chlorine is liberated at anode and

hydrogen is liberated at cathode

Cell reaction:

At anode: 2 Cl � ▬▬► Cl2 ↑ + 2e

�

At cathode: (i) Na + + H2O

▬▬► NaOH + H+

(ii) 2 H+ + 2e

� ▬▬► H2↑

Net reaction: 2NaCl + 2H2O ▬▬► 2NaOH + H2 + Cl2

Differences between electrolytic and electrochemical cells

Electrolytic cell

Electrochemical cells

/ Galvanic Cell

Conversion of electrical energy into

chemical energy

Chemical energy into electrical

energy

The anode is positive plate and

cathode is negative plate

The anode is negative plate and

cathode is positive plate



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Electrons are supplied to the cell

from the external power supply

Electrons are drawn from the cell.

Not a spontaneous reaction

Spontaneous reaction.

eg. Electroplating

eg. Corrosion

The extent of chemical reaction

occurring at the electrode is

governed by Faraday’s law of

electrolysis.

The e.m.f of the cell depends on the

concentration of the electrolyte

and chemical nature of the electrode

(Nernst Equation)

The amount of electricity passed

during electrolysis is measured by

Coulometer. e.g: Electroplating,

Electrolysis

The e.m.f produced in the cell is

measured by potentiometer.

e.g: Corrosion, Discharging of

battery

Ion Selective / Ion Sensitive Electrodes

Glass electrode Combined glass electrode

Glass electrode is a perfect example of ion selective electrode. It consists of

a glass membrane which is permeable to a specific ion (such as Li+ / Na

+ /

K+

/ NH4+ etc), depending upon its composition.



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The glass membrane has better conductivity than ordinary glass. The

membrane is coated with hydrated silica. Though the ion selective electrode

used for determining pH ,is not permeable to H+

ions, it is sensitive to H+

ions. When the comes in contact with acids of varying concentration at the

inner and outer surface, a potential is developed due to concentration

gradient between the acid within the bulb and outside the bulb.

The magnitude of potential depends on the pH of the solution in which the

glass electrode is immersed. If the pH of the external medium is 7 the

potential becomes zero.

The glass electrode is represented as

Ag/AgCl(s)/HCl (0.1M)/ glass membrane

The glass electrode can also be used as a reference electrode.

As anode

Ag (l) ▬▬► Ag+ + e-

Hg+ + Cl− ▬▬►AgCl

As cathode

Ag+

Cl - + e- ▬▬►Ag (s) +Cl−

The net reaction can be represented as Ag + Cl- ↔↔↔↔ AgCl

Applying Nernst equation for the net reaction

Ecell = Eo

cell + 0.0591/n . log [Ag][Cl−] / [AgCl]

The [Ag] and [AgCl] are unity .Therefore the Ecell depends on the

concentration of chloride ions , hence the electrode is said to be reversible

wrt chloride. (Another electrode which is reversible wrt to Chloride ions is

Calomel)

The concentration of chloride ions ↓↓↓↓ ses during oxidation and ↑↑↑↑ses during

reduction.



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DETERMINATION OF pH USING GLASS ELECTRODE: Refer class notes for cell

construction and derivation of expression for pH determination

APPLICATIONS OF EMF MEASUREMENT

� Calculation of Thermodynamically important parameters

� Calculation of equilibrium constants

� Determination of solubility product of sparingly soluble salts

� Determination of Valency

� Determination of pH of any solution like acids, body fluids, natural

water, waste water, coloured solutions etc.

� Potentiometric titrations

� Conductometric titrations

Calculation of Thermodynamically important parameters

According to first law of thermodynamics, energy can be neither created nor

destroyed, but can be converted from one form to another.

A cell is a device which converts the stored chemical energy to electrical

energy. Hence the loss in free energy is the gain in electrical energy

When the emf of the cell is measured at various temperatures, then ∆G , ∆S

and ∆H can be calculated at any given temperature , using the relationships

(i) −∆G = n F E

(ii) ∆S = nF (∆E/∆T)

(iii) ∆H = ∆G+T∆S

Calculation of equilibrium constants

−∆G = n F Ecell

−∆G = 2.303 RT log K (eq)

Therefore log K(eq) = n F E

2 .303 RT



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Determination of solubility of sparingly soluble salts

The solubility of any sparingly soluble salt can be calculated by constructing

a concentration cell.

For example to measure the concentration of AgCl, a cell is constructed as

follows:

Ag/AgCl(sat)/KCl (0.1M)‖ AgNO3(0.01M)/Ag

If a drop of silver nitrate is added to the anodic compartment, the following

reaction takes place

KCl (aq)+ AgNO3(aq) ▬▬► K NO3 (aq) + AgCl↓

The precipitated silver chloride alters the emf of the cell. From the change in

the emf the solubility product is calculated from the relationship

E = 0.0591/n X log 0.01 / C

E can be measured, n=1 hence C (conc. Of AgCl) can be calculated

Determination of Valency

The “n” ie. The number of electrons involved in the cell reaction (which

indicates the oxidation number / valency of a metal) can be calculated by

measuring the emf of a cell with known concentrations of electrolytes of a

metal for which oxidation state has to be determined , using the expression

E(con.cell) = 2.303 RT log [C2] nF [C1]

Determination of pH of any solution like acids

log K(eq) = nE

0.0591



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Refer Class notes

Potentiometric titrations : 1

Contents

Estimation of ferrous ion

Aim To estimate the amount of Fe2+

Cell Construction

− +

SCE II Fe3+, Fe2+ ,Pt

Type of Reaction

Redox

Reaction

K2Cr2O7+6FeSO4+7H2SO4───►Cr2(SO4)3+3Fe2(SO4)3+7H2O+K2SO4

EMF due to the oxidation of Fe2+ to Fe3+

Measurement Reference electrode is SCE, Indicator electrode is Platinum.

Units

Volts

Reactants

Std. oxidising agent Vs Given Fe2+

Burette Solution

Oxidising Agent (KMnO4, Cerric sulphate,K2Cr2O7 etc.)

Pipette Solution

20 ml of given Fe2+

Additional Solution

20 ml of dil. Sulphuric acid

End Point

Sudden increase in the EMF as the ratio of Fe3+/ Fe2+ attains maxima

Equivalent Wt.

Fe2+ is 55.85



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Calulation

Amt of Fe2+

= Normality of Fe2+

x Eq.Wt.of 55.85(Fe2+

)

Graph 1

Graph 2

X axis

Volume of Oxidising Agent (ml)

Average volume of Oxidising agent (ml)

EMF(Volts)

∆E /∆V (V/ml)

Y axis Shape

Y emf X Vol. of. KMnO4

Y

∆E /∆V

X Av. Vol. of. KMnO4

Result Amount of ferrous in the given solution = _____g/L

Potentiometric titration :2

Contents

Estimation of the amount of Silver chloride / Barium chloride by

precipitating it as a sparingly soluble salt

Aim To estimate the amount of AgCl / BaSO4

Cell Construction

− +

SCE II AgNO3/Ag

Type of Reaction

Precipitation



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Reaction

MCl+ AgNO3 ───► MNO3 + AgCl↓↓↓↓

BaCl2 + M2SO4 ───► 2MCl+ BaSO4↓↓↓↓

EMF due to the change in concentration of AgNO3 / BaCl2

Measurement

Reference electrode is SCE, Indicator electrode is Ag/AgCl.

Units

Volts

Reactants

Std. M Cl agent Vs Given AgNO3 Std M2SO4 Vs Given BaCl2

Burette Solution

M Cl / M2SO4

Pipette Solution

20 ml of given Ag+ / 20 ml of Ba2+

End Point

Sudden increase in the EMF

Calculation

Amt of Ag NO3 = Normality of Ag NO3 x Eq.Wt.169.87 (Ag NO3)

Amt of BaCl2 = Normality of BaCl2 x Eq.Wt.122.14 (BaCl2.2H2O)

Graph 1 Graph 2 X axis

Volume of Precipitating agent (ml)

Average volume of precipitating agent (ml)

EMF(Volts)

∆E /∆V (V/ml)

Y axis Shape

Y emf X Vol. of.K2SO4

Y

∆E /∆V

X Av.Vol. of. MSO4

Result Amount of BaCl2 in the given solution = _____g/L



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Conduct metric titrations

Contents

Estimation of Acid

Aim

To Estimate the amount of given acid

Type of Reaction

Neutralisation

Reaction

H+ + OH- H2O

Measurement

Conductance due to mobility of ions

Units

Observed conductance X 10-3 mho

Reactants

Std Base Vs Given acid

Burette Solution Std Base

Pipette Solution

20 ml of Given Acid

Additional Solution

Water to immerse the Platinum foils

End Point

Increase in Conductance after the initial decrease

Equivalent Wt.

HCl is 36.45 HNO3 is 63 H2SO4 is 49

Calculation

Amt of H+= Nacid x Eq.Wt.of H

+

Model Graph X axis

Volume of Base (ml)

Y axis Conductance (mho)



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Shape

Y Conductance (m.mho) X

Vol.of NaOH ml

Result

Amount of acid present in the given solution = ---------g/L



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Electro Chemistry