Electrochemistry

Electrochemistry deals with the links between chemical reactions and electricity.

This includes the study of chemical changes caused by the passage of an electric current

across a medium, as well as the production of electric energy by chemical reactions.

Electrochemistry also embraces the study of electrolyte solutions and the chemical

equilibria that occur in them.

The devices used for the inter-conversion between chemical and electrical forms

of energy are called electrochemical cells.

Electrochemical cells which generate an electric current are called voltaic cells or

galvanic cells, and common batteries consist of one or more such cells. In other

electrochemical cells an externally supplied electric current is used to drive a chemical

reaction which would not occur spontaneously. Such cells are called electrolytic cells.

Electrolytic cell

An electrolytic cell decomposes chemical compounds by means of electrical

energy, in a process called electrolysis. The result is that the chemical energy is

increased.

An electrolytic cell has three component parts: an electrolyte and two electrodes

(a cathode and an anode). The electrolyte is usually a solution of water or other solvents

in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes.

When driven by an external voltage applied to the electrodes, the electrolyte provides

ions that flow to and from the electrodes, where charge-transferring or redox, reactions

can take place.

The electrolytic cell is the industrial chloralkali cell in which brine (an aqueous sodium

chloride solution) is electrolytically converted to chlorine and caustic soda (sodium

hydroxide, NaOH). The external power source supplies electric energy to drive the

overall reaction.

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2Cl− + 2H2O → Cl2 + H2 + 2OH−          

Chloride ion is oxidized to chlorine gas at the carbon electrode, and water is reduced to

hydrogen gas (H2) and hydroxide ion (OH−) at the iron electrode. The electrolytes are

maintained as electrically neutral by a flow of sodium ions through the separator (such as

an ion exchange membrane).

For the electrolytic cell, the external markings of anode and cathode are opposite the

chemical definition. That is, the electrode marked as anode for discharge acts as the

cathode while charging and the electrode marked as cathode acts as the anode while

charging.

Voltaic Cells

When zinc metal is placed in a solution of copper ions as described by the net

ionic equation shown below.

Cu+2 (aq) + Zn (s) -------> Cu(s) + Zn+2 (aq)

The zinc metal slowly "dissolves" as its oxidation produces zinc ions which enter into

solution. At the same time, the copper ions gain electrons and are converted into copper

atoms which coats the zinc metal or sediments to the bottom of the container. The energy

produced in this reaction is quickly dissipated as heat, but it can be made to do useful

work by a device called, a voltaic cell.

A voltaic cell is composed to two compartments or half-cells, each composed of

an electrode dipped in a solution of electrolyte. These half-cells are designed to contain

the oxidation half-reaction and reduction half-reaction separately as shown below.

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The half-cell, called the anode, is the site at which the oxidation of zinc occurs as shown

below.

Zn (s) ----------> Zn+2 (aq) + 2e-

During the oxidation of zinc, the zinc electrode will slowly dissolve to produce zinc ions

(Zn+2), which enter into the solution containing Zn+2 (aq) and SO4-2 (aq) ions.

The half-cell, called the cathode, is the site at which reduction of copper occurs as

shown below.

Cu+2 (aq) + 2e- -------> Cu (s)

When the reduction of copper ions (Cu+2) occurs, copper atoms accumulate on the surface

of the solid copper electrode.

The reaction in each half-cell does not occur unless the two half cells are connected to

each other.

In order for oxidation to occur, there must be a corresponding reduction reaction that is

linked or "coupled" with it. Moreover, in an isolated oxidation or reduction half-cell, an

imbalance of electrical charge would occur, the anode would become more positive as

zinc cations are produced, and the cathode would become more negative as copper

cations are removed from solution. This problem can be solved by using a "salt bridge"

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connecting the two cells as shown in the diagram below. A "salt bridge" is a porous

barrier which prevents the spontaneous mixing of the aqueous solutions in each

compartment, but allows the migration of ions in both directions to maintain electrical

neutrality. As the oxidation-reduction reaction occurs, cations ( Zn+2) from the anode

migrate via the salt bridge to the cathode, while the anion, (SO4-2), migrates in the

opposite direction to maintain electrical neutrality.

The two half-cells are also connected externally. In this arrangement, electrons

provided by the oxidation reaction are forced to travel via an external circuit to the site of

the reduction reaction.

The differences between galvanic and electrolytic cells can be summarised in a table.

Galvanic/Voltaic Cells Electrolytic Cells

chemical energy electrical energy electrical energy  chemical energy

Two half-cells with separate electrolytes and a

salt bridge (or porous barrier).

Electrodes usually in the same electrolyte

chemical reaction is spontaneous chemical reaction is forced by applying a voltage -

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Eo total is positive it is not spontaneous

Eo total is negative

The electrode on which oxidation takes place

is called the anode and the electrode on which

reduction takes place is called cathode.

anode - negative terminal

cathode - positive terminal

The electrode which is connected to the negative

terminal of the battery is called cathode / the cations

migrate to it which gains electrons and hence

reduction takes place.

anode - positive electrode

Cathode - negative electrode

Reversible and Irreversible cells:

Reversible cells

Electrochemical cells can also be said as two types namely reversible and

irreversible cell. A cell works reversibly in the thermodynamic conditions i.e., during the

measurement of EMF no current flows through the cell and no chemical reaction takes

place. Such cells are called as reversible cells.

If the external EMF is infinitely greater than that of the cell emf, an extremely

small amount of current flows through the cell in the opposite direction and small amount

of the chemical reaction takes place in the reverse direction. E.g. Daniel cell is a

reversible cell. Its cell potential is 1.1 V. Thus in Daniel cell (a galvanic cell), zinc

undergoes dissolution and copper undergoes deposition to realize an emf of 1.1V, as per

the following reaction sequence:

Zn + Cu2+ === Zn2+ + Cu …….. 1.1 V

If an emf of –1.101 V is impressed on Daniel cell, copper undergoes dissolution and zinc

undergoes deposition.

Irreversible cells

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Cells which do not obey the conditions of thermodynamic reversibility are called

irreversible cells. Irreversible is a cell where the cell reaction cannot be reversed even on

applying infinitesimally small but excess applied emf i.e. the products produced during

the cell reaction are not available for recombination on reversal of voltage. Example of an

irreversible cell is the cell used for the electrolysis of brine or the dry cell used in pen-

torches. In the electrolysis of brine (aqueous NaCl solution) for example, on applying

voltage, Na+ ions move towards cathode, gain one electrode and become elemental

sodium atoms. But the sodium atoms immediately react with water to form sodium

hydroxide. Similarly, chloride ions move towards anode, loose one electron to form

chlorine atoms. These chlorine atoms recombine forming molecular chlorine, which is

evolved as a gas. The reaction sequence is given below:

Na+ + e → Na ; 2 Na + 2 H2O → 2 NaOH + H2

Cl- → Cl + e ; Cl + Cl → Cl2 ↑

Cell diagram or representation of a cell  

In general, the electrode at which reduction takes place is written on the RHS of the salt

bridge and the electrode at which oxidation takes place is written on the LHS of the salt

bridge. The salt bridge linking the aqueous solutions is represented by two vertical

parallel lines having ions on both sides.  

 

For Zn – Cu cell, Zn electrode is written on the LHS while the Cu electrode on the RHS

of the salt bridge.  

Zn Zn2+ Cu2+ Cu

The symbol for an inert electrode, like the platinum electrode is often enclosed in a slash.

For example,

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Mg Mg2+ H+ H2 Pt

The value of emf of a cell is written on the right of the cell diagram. Thus a zinc-copper cell has emf 1.1 V and is represented as

Zn Zn2+ Cu2+ Cu E = +1.1 V

Electrode Potential

Consider a zinc rod being dipped in ZnSO4 solution. The zinc atoms on the surface of the

metal with in the solution have tendency to release Zn+2 into solution retaining the

electrons on the surface of the metal. This process is called dissolution or solution

pressure of the metal and it is oxidative in nature.  

 

The zinc ions of the solution have a tendency to accept the electrons on the surface of

zinc rod to form neutral zinc atoms and get deposited on the zinc rod. This process is

called deposition or the osmotic pressure of solution is reductive in nature.  

 

These two processes will be taking place simultaneously at different rates. In this case,

the rate of dissolution is found to be greater than the rate of deposition. Consequently, by

the time equilibrium is reached, more of dissolution would have occurred and the solution

becomes negatively charged. Due to the attractive electrostatic forces, the Zn ions

accumulate around the Zn rod and an electrical double layer of opposite charges is

formed and between the Zn rod and the solution the potential developed is called single

electrode potential or Electrode Potential.

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Tendency of an electrode to lose or gain electrons when it is in contact with its

own ions in solution is called electrode potential. Tendency to gain electrons means

reduction potential and tendency to lose electrons means the oxidation potential.

Standard Electrode Potential (Eo)

It is the potential developed when the pure metal is in contact with its own ions at

one molar concentration at a temperature of 25oC or 298 K.

Cell potential or Electro Motive Force (EMF)

The emf (electro motive force) of a cell is the algebraic sum of the potentials of

the two constituent single electrode systems. It is obvious that cell is made of two half-

cells / single electrode systems. A cell is generally represented with the negative

electrode / anode written first at the left and then the cathode / positive electrode at the

right. Thus the emf of a galvanic cell is calculated from the half-cell potentials using the

relation

Ecell = Eright - Eleft = Ecathode - Eanode

Here it is to be noted the values of std potentials (reduction potentials) of cathodes

are more positive and those of anodes are more negative, so that the cell potential is

positive.

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e.g. for Daniel cell, Ecell = Eright - Eleft = Ecathode - Eanode =

ECu2+

/ Cu – EZn2+

/ Zn = 0.34 – (-0.76) = 1.10 Volt

Measurement of EMF

The potential difference, which causes current flow from the electrode of

higher potential to the electrode of lower potential, is called electromotive force (emf) of

the cell. The emf of a cell cannot be directly determined by connecting across a

voltmeter, as some part of the cell current is drawn by the voltmeter during its

measurement. This results in the formation of reaction products at the electrodes and

hence a change in the electrolyte concentration around the electrodes. This difficulty can

be overcome by the measurement of excess applied opposite emf that just nullifies the

cell emf (Poggendorff’s external compensation method). Care is to be taken during the

measurement such that the current taken from the cell is negligibly small and the ionic

concentrations are not appreciably altered. The emf of the cell thus remains constant and

its value can be determined with high degree of accuracy / precision.

The emf of a cell can be determined by Poggendorff’s external compensation

method. In this potentiometric measurement of cell emf, a standard cell is used whose

emf is known and does not vary with time. Weston cadmium cell is the conventionally

used standard cell. Fig. below is the schematic of a simple potentiometer.

It consists of a uniform wire AB of high resistance. A storage battery C of constant but

large emf is connected to the ends A and B of the wire through a variable resistance

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G

X (Cell unknown)

S (Cell - std)

A

C

B

R

K

D D’

(rheostat) R. The cell X whose emf is to be determined is included in the circuit through a

galvanometer G and a sliding contact D. The circuit is closed using a plug key K. The

position of the sliding contact D is slowly changed along the wire AB till a point is

reached at which there is no net current flowing through the galvanometer (its deflector

points to zero). This null point position D is noted.

The standard cell S (whose emf is known) is then introduced into the circuit and

the circuit closed using the plug key K. The position of the sliding contact D’ is again

slowly changed along the wire AB till a point is reached at which there is no net current

flowing through the galvanometer (its deflector points to zero). This null point position

D’ is noted. The emf of the cell X (unknown) is determined using the relation mentioned

below:

Emf of cell X (Ex) α balancing length AD

Emf of standard cell S (Es) α balancing length AD’s

Thus Es / Ex = (length AD’s) / (length ADx), from which Ex can be determined, as the

value of Es is known. Nowadays the digital potentiometers are used which have the in-

built circuitry of potentiometer set up and standard cell with switching arrangement for

standard cell so that the unknown cell is connected externally to directly read the cell emf

as digital display.

Relation between Gibb's Free Energy and Cell Potential (EMF) – Nernst Equation.

When a cell reaction takes place electrical energy is produced which results in decrease

in the free energy of the system.

Electrical work = Decrease in free energy

 In an electro chemical cell,

 Electric work done = Quantity of current produced x E.M.F.

 For one mole of electrons quantity of current is 1F (96500 coulomb)

 Therefore for n moles it is nF.

 Electric work done = nFEcell

  and For a standard cell

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For a general cell reaction

 

 Van’t-Hoff isotherm says

G = Go + RT

Hence G = -nFE

-nFE = -nFEo + RT

Divide by nF and reverse the sign

 

Substituting R = 8.314 J/K/mole; F = 96500 coulombs; T = 298 K and multiplying 2.303

for conversion of ln to log, the equation becomes

 

Applications:

1) Determination of potential of cell and electrode

2) Free energy change can be determined using the equation

3) Equilibrium constant can be calculated using the relation

-∆G0 = RTlnK

Reference Electrodes:

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The potential of an electrode system (electrode of interest or working electrode)

can be measured by coupling with other electrode with a voltmeter introduced between

them. The coupled electrode should not possess any charge transfer reaction (electrode

reaction) in the electrolyte used or it should not be polarized. Such ideally non-

polarizable electrodes used for the measurement of working electrodes are called

reference electrodes. Reference electrodes are two types namely primary and secondary

reference electrodes. Primary reference electrode is one that is universally used such as

standard hydrogen electrode (SHE) and its potential is arbitrarily taken as zero. But

SHE involves tedious and cumbersome construction. This difficulty is overcome by the

use of ‘secondary reference electrodes, which can be constructed easily and their

potentials can be determined with SHE as reference.

It consists of a platinum electrode immersed in a 1 M solution of H+ ions

maintained at 25C. Hydrogen gas at one atmosphere enters the glass hood and bubbles

over the platinum electrode. The hydrogen gas at the platinum electrode passes into

solution, forming H+ ions and electrons. It is represented as;

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1 M HCl

Pt. foil

H2 (1 atm.)

Pt, H2 (1atm) H+ (1M) Eo = 0 V

Limitations

i) It requires hydrogen gas and is difficult to set up and transport

ii) It requires large volume of test solution

iii) The solution may become poison, if any impurities presents on the surface of

Pt electrode.

iv) The potential of the electrode is dependent on atmospheric pressure.

Secondary Reference Electrode

Saturated calomel electrode (SCE)

Examples of secondary reference electrodes are calomel electrode, silver-silver

electrode, glass electrode, quinhydrone electrode etc. Calomel electrode is set up with

mercurous chloride, mercury and potassium chloride electrolyte and represented as Hg |

Hg2Cl2 | KCl.

Depending on the concentration of KCl, calomel electrode is of three types namely

saturated normal and decinormal calomel electrodes. The electrode reaction for calomel

electrode is Hg2Cl2 (s) + 2e- → 2Hg(s) + 2Cl-

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Saturated KCl

Hg + Hg2Cl2

Hg

Pt wire

The potentials of the different calomel electrodes against SHE reference is given below:

[ KCl ] Saturated 1.0 N 0.1 N

Potential , V 0.2422 0.2810 0.3335

Merits and Demerits of Calomel electrode: 1. Gives relative pH values compared to

hydrogen electrode, which gives absolute values of pH.

Ag/AgCl electrode is prepared by depositing a thin layer of AgCl electrolytically on a Ag

or Pt wire and immersing in a solution containing the chloride ions. It is represented as

Ag | AgCl | Cl-( M)

The electrode reaction of Ag/AgCl electrode is : AgCl + e → Ag + Cl-

Ion – Selective Electrode (ISE)

An Ion-selective electrode (ISE) is a transducer (sensor) which converts the

activity of a specific ion dissolved in a solution into an electrical potential which can be

measured by a voltameter or pH meter. The voltage if theoretically dependent on the

logarithm of the ionic activity, according to the Nernst equation.

The sensitive part of the electrode is usually made as an ion-specific membrane

along with a reference electrode. Hence ISE is also known as a specific ion electrode

(SIE).

Ion-selective electrodes are used in biochemical and biophysical research, where

measurements of ionic concentration in an aqueous solution are required, usually on a

real time basis.

Principle

At equilibrium, the membrane potential is mainly dependent on the concentration

of the target ion outside the membrane and is described by the Nernst equation. Briefly,

the measured voltage is proportional to the Logarithm of the concentration, and the

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sensitivity of the electrode is expressed as the electrode Slope - in millivolts per decade

of concentration. Unknown samples can then be determined by measuring the voltage

and plotting the result on the calibration graph.

The use of metals directly as ISE has the following disadvantages: (i) slow electrode

response (ii) Nernst equation not followed (iii) change of electrode potential due to the

availability of electrons on the electrode surface (iv) no well defined electron change.

Hence various membranes are used in ISEs and such electrodes are called Ion selective

membrane electrodes (ISMEs). ISMEs show some degree of specificity and selectivity,

These electrodes utilize some membrane to confine an inner solution and the reference

electrode. Membranes in the ISE and reference electrode (RE) sides function by ion

exchange (IE) mechanism.

There are three main types of membranes used in the Ion selective electrodes; they are

(i) Glass membrane

(ii) Solid-state crystal membrane and

(iii) Liquid ion-exchange membrane

Based on the ion selective membranes used, the ISE are classified into three groups

a) Glass membrane electrode

b) Crystalline membrane electrode

c) Liquid membrane electrode

1. Glass Membrane Electrodes

Glass membranes are made from an ion-exchange type of glass (silicate of

chalcogenide). This type of ISE has good selectivity, but only for several single-charged

cations; mainly H+, Na+, and Ag+. The glass membrane has excellent chemical durability

and can work in very aggressive media. A very common example of this type of electrode

is the pH glass electrode.

The selectivity of glass membranes depends depends on the composition of glass.

Generally they are based on Na2O-Al2O3-SiO2 mixtures.

Two typical composition of glass membranes are:

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i) Na2O (22%), CaO (6%) and SiO2 (72%) - responds to H+ ions

ii) Na2O (11%), Al2O3 (18%) and SiO2 (71%) – responds to alkali metal ions

Glass membrane electrodes are further subdivided into three types based on the

selectivity characteristics

i) pH type glass electrode

ii) Cation sensitive type

iii) Sodium sensitive type

Glass Electrode:

The glass elctrode assembly consists of a thin glass bulb filled with 0.1 N HCl and a

silver wire coated with silver chloride immersed in it. The Ag/AgCl electrode here acts as

the internal reference electrode. The glass electrode is represented as

Ag | AgCl(s) | 0.1 M HCl | glass.

Determination of pH of an aqueous solution by glass electrode:

The determination pf pH of a solution is one of the important applications of the

EMF measurements. When glass electrode is immersed in the solution whose pH is to be

determined, a potential difference is set up between the two surfaces of the glass

membrane. The potential value developed is proportional to the pH of the test solution

(sample).

The magnitude of this difference of potential (single electrode potential) is given by

EG = EoG - 0.0591 log [H+]

EG = EoG + 0.0591 pH

Where EoG is constant for the given glass electrode and it depends upon the nature and

composition of glass membrane.

Actually, the glass membrane of the glass electrode undergoes an ion-exchange reaction

in which the sodium ions of the glass membrane are exchanged with protons of the

sample solution. The electrode reaction of the glass electrode immersed in the test

solution can be represented as

glass ---- Na + H+ = glass ---- H + Na+

To carry out the determination of pH of a solution, the glass electrode is connected with a

saturated calomel electrode.

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The cell is therefore represented as;

Ag, AgCl / HCl (0.1 N) / Glass / unknown solution / SCE

Ecell = Eright - Eleft

= ESCE - EG

= ESCEl – (EoG + 0.0591 pH)

= ESCE – EoG - 0.0591 pH

pH = ESCE – EoG - Ecell

0.0591

Advantages of Glass electrode

i) To determine pH of any solution

ii) Small quantity of solution is sufficient for determination

iii) Used even in the presence of metallic ions and poisons

iv) Equilibrium is easily reached

Disadvantages:

i) Glass membrane used in the bulb is very thin and its resistance is very high,

electronic potentiometers are needed for the measurements

ii) Cannot be used in strongly alkaline solution (pH > 10), in such cases special

type of glass membranes has to be used.

Electrochemical series:

Electrochemical series is the arrangement of elements in the ascending (or descending)

order of their electrode (reduction) potential values with hydrogen at the centre.

Electrode systems appearing earlier in the series have the oxidation reaction spontaneous

and are termed as anodes and those appearing later in the series have the reduction

reaction spontaneous and are termed as cathodes.

Half-reaction   E° (V)  

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Li+ + e− ⇄ Li(s)  −3.0401

Cs+ + e− ⇄ Cs(s)  −3.026

K+ + e− ⇄ K(s)  −2.931

Ca2+ + 2e− ⇄ Ca(s)  −2.868

Na+ + e− ⇄ Na(s)  −2.71

Mg2+ + 2e− ⇄ Mg(s)  −2.372

Ti2+ + 2e− ⇄ Ti(s)  −1.63

Mn2+ + 2e− ⇄ Mn(s)  −1.185

Zn2+ + 2e− ⇄ Zn(s)  −0.7618

Fe2+ + 2e− ⇄ Fe(s)  −0.44

Cd2+ + 2e− ⇄ Cd(s)  −0.40

Sn2+ + 2e− ⇄ Sn(s)  −0.13

2H+ + 2e− ⇄ H2(g)    0.0000

Cu2+ + 2e− ⇄ Cu(s)  +0.340

Bi+ + e− ⇄ Bi(s)  +0.50

Ag+ + e− ⇄ Ag(s)  +0.7996

Ce4+ + e− ⇄ Ce3+  +1.44

F2(g) + 2e− ⇄ 2F−  +2.87

The potential of a redox electrode system is given merely as a number (modulus value,

irrespective sign). Thus if the reduction potential is positive, the oxidation potential is

negative for the same system and vice-versa. Std. electrode potential conventionally

represents the reduction potential or potential of the reduction reaction.

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Importance / Significance / Applications of electrochemical series: Electrochemical series

is useful / significant in the prediction of

1. Spontaneity of reactions: The feasibility / spontaneity of reactions of the reactions

be predicted from the knowledge of the electrode potential values. Such processes are

spontaneous whose standard potential is positive.

2. Corrosion behaviour of metals and alloys: Metals / alloys with negative values of

std. electrode potentials are prone / susceptible to corrosion and those with positive

values of std. electrode potentials are resistant to corrosion (an undesirable

phenomenon).

e.g. zinc is more easily corroded than copper (or) copper is more resistant to corrosion

than zinc.

3. Redox behaviour of materials: Materials with more negative values of

std.electrode potentials are used in reduction reactions – addition of electrons (as

they can donate electrons) and Materials with more posittive values of std.

electrode potentials are used in oxidation reactions – removal of electrons (as they

can accept electrons). E.g. zinc, tin etc. are used as reducing agents whereas

oxides of copper etc are used as oxidizing agents.

4. Displacement characteristics of metals: Metals with more negative std. electrode

potentials will displace metals with more positive std. electrode potential values.

E.g. zinc will displace copper from its salt solution and not vice-versa.

5. Determination of equilibrium constant of the reaction from the knowledge of

electrode potential values, using the relation ΔG = - nFE = - RT ln K.

Potentiometric Titration:

Potentiometric titration is used for following the course of reactions involving

electrolytes, where there is no proper indicator available. In a potentiometric titration, a

suitable electrode immersed in the solution to be titrated acts as the ‘indicator’. The

indicator electrode is paired with a reference electrode e.g., platinum electrode acts as

indicator whereas calomel acts as reference electrode.

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The potential of the electrode system is determined using the platinum and

calomel electrodes immersed in the reaction mixture after regular additions of the titrant.

The potential observed is plotted against the volume of titrant added. The endpoint is

determined graphically from the change in trends before and after the completion of the

reaction.

The potentiometric titrations may be of three types:

(a) Acid-base titrations

(b) Oxidation-reduction titrations

(c) Precipitation titration

Oxidation – Reduction titrations

The titration of ferrous (ammonium) sulphate (FAS) and potassium dichromate is

considered to be an example of oxidation-reduction titration. Fe2+ is oxidized to Fe3+,

while Cr6+ is reduced to Cr3+

6 Fe2+ + Cr2O72-- + 14 H+ 6 Fe3+ + 2Cr3+ + 7H2O

The redox reaction between ferrous (ammonium) sulphate (FAS) and potassium

dichromate is followed by determining the potential using the platinum and calomel

electrodes immersed in the reaction mixture after adding regular volumes of dichromate.

The potential observed is plotted against the volume of titrant added (E Vs V plot). The

endpoint is determined graphically from the change in trends before and after the

completion of the reaction. A derivative plot can also de made between change in

potential to change in volume (ΔE/ΔV) and the average volume of the titrant [(V1+V2)/2].

The volume corresponding to the peak in the derivative plot directly gives the end point

condition whereas the involved electrode systems and their potentials can be got from the

E Vs V plot. The model graphs for the potentiometric titrations are given below:

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The S-shaped curve for the E Vs V plot is due to the fact that the electrode system

itself is changed after the completion of reaction i.e. initially when a small amount of

dichromate (Cr2O72- with chromium in its Cr6+ state) is added to the reaction mixture

containing FAS and dilute sulphuric acid in a beaker, corresponding amount of ferrous

ions (Fe2+) is oxidized to ferric state (Fe3+) and Cr6+ is reduced to Cr3+ state. Thus the

beaker contents are two ionic (redox) species of iron (i.e.Fe2+ and Fe3+) with platinum

electrode (inert electrode, contributing only electronic conductivity) immersed. This

constitutes the iron electrode system, whose potential varies gradually (as given by

Nernst equation) with regular additions of titrant. The After the end point (reaction

completion), all the ferrous ions are completely oxidized to ferric ions and the excess

added dichromate ions (Cr6+) exist as redox couple with Chromium (III) ions, leaving

only one type of species for iron (Fe3+) in the beaker. Now the electrode system is

chromium electrode system, whose potential varies gradually (as given by Nernst

equation) with regular addition. As the electrode system itself is changed during the

reaction, there is a shoot up in potential in the

E Vs V plot and hence it is S-shaped.

Precipitation titration:In precipitation titration, an ion from the solution is precipitated out by the

addition of a titrant. The changes in the concentration of the ion during the course of the

titration can be followed by measuring the potential of an electrode reversible to the ion.

A typical precipitation is that of sodium chloride solution against silver nitrate solution.

In the titration of silver nitrate solution by sodium chloride, a silver wire is dipped in

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Volume of titrant (ml)V1 + V2

2

∆E

/ ∆

V

mV

/ml

Em

f (V

)

silver nitrate solution. When sodium chloride solution is added, AgCl is formed and we

have a silver-silver chloride electrode. It is combined with a reference electrode like SCE

and the emf of the cell.

Ag AgCl Cl- SCE

is given by

E = ESCE – E0Cl

- AgCl Ag - RT/F ln acl-

But the activity of chloride ions is controlled by the solubility product of silver

chloride through the relation

KSp = aAg+aCl-

Hence the cell emf can be written as

E = ESCE – E0Cl

- AgCl Ag - RT/F ln KSp + RT/F ln aAg+

At 298 K

E = Const + 0.0591 log aAg+

Thus the activity of silver ions decreases the potential of the cell decreases. The

decrease is, however, small in the initial stages of the titration. At the equivalence point

silver ion concentration is very small due to slight solubility of silver chloride, the

potential therefore changes rapidly. Excess of chloride ions, the potential therefore

changes rapidly. Excess of chloride ions after the equivalence point do not affect

significantly the concentration of silver ions and hence the potential of the electrode

changes rather slowly.

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Em

f (V

)

Volume of titrant (ml)

Conductometric titration is a method of volumetric analysis based on the

change in conductance of the solution, at the equivalence point (end point) during

titration. The principle of conductometric titration is that the conductance of a reaction

follows a specific trend before the completion of the reaction and it follows a different

trend after the completion of the reaction. From the change in trends, the end point of the

reaction can be determined graphically. The conductance of the solution depends on

i) Number of free ions in the solution

ii) Charge of the free ions

iii) Mobility of the ions

When adding one electrolyte solution to other solution, the conductance of the

solution changes due to the number of free ions in the solution changes during titration.

Conductometric titration of strong acid against a strong base:

The course of neutralization of a strong acid by a strong base can be determined by

condutometric method (without the use of indicator).

In the case of neutralization of strong acid (say HCl) by a strong base (say NaOH), the

conductance of HCl (taken in a beaker) is determined at regular additions of NaOH, from

the burette.

H+Cl- + Na+OH- → Na+Cl- + H2O

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The conductance of the reaction mixture (HCl) decreases till the end point because lighter

protons (H+) are replaced by heavier Na+ ions, whose mobility is lower than that of

protons. The conductance of the reaction mixture increases after the end point because

heavier chloride ions (Cl- whose atomic mass is 35.5) are replaced by lighter hydroxyl

(OH-) ions (with mass 17), whose mobility is higher than that of chloride ions. A plot of

the conductance of the reaction mixture against the volume of the titrant gives two

straight lines of opposite slopes. The point of intersection of the straight lines is the end

point.

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