electro chemistry
TRANSCRIPT
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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
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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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