ELECTROCHEMISTRY

CHAPTER-1ELECTROCHEMISTRY

1.1 Important Terminologies

Current: Flow of electrons through any conducting material is known as the current.

Oxidation:The tendency to lose electrons

Reduction: The tendency to gain electrons

Electrode:When a metal rod is dipped in its salt solution, it develops a positive or negative potential. This assembly is called as an electrode.

Anode:The electrode at which oxidation occurs is called as anode.

Cathode:The electrode at which reduction occurs is known as cathode.

Electrolytes:Electrolytes are the conductors which are in the form of solution or fused state in which conductance takes place due to the movement of ions present in it.Electrochemical cell: The electrochemical cell consists of two conductors called electrodes that are immersed in an electrolyte and are connected externally by means of a metal conductor.

Electrolysis:The process that occurs in the electrochemical cell is called as electrolysis. It is the breakdown of electrolyte into ions by electricity.

Half cell:A part of a cell containing electrode dipped in an electrolytic solution is called as a half cell.

Oxidation half cell : The electrode where oxidation takes place i.e. where electrons are lost

Reduction half cell : The electrode where reduction takes place i.e. where electrons are gained

Conductors:Material that permits electric current to pass through.

Example: Copper, aluminium, silver, fused salts, bases, aqueous solutions of acids etc.,

Insulators or non-conductors: Materials which do not allow the electric current to pass through it.

Example: Wood, plastics etc.,

Standard electrode potential:It is the potential of the electrode which is in contact with its solution of 1M concentration which is measured at 298K and 1atm pressure.

Single electrode potential:The measure of tendency of a metallic electrode to lose or gain electrons, when it is in contact with a solution of its own salt.

Representation Notation and sign conventions:

1.Molecules, elements, gases and electrode material are represented by the usual chemical symbols. Concentration of ions and molecules and partial pressures of gases are given in brackets.

2.Interface between an electrode and electrolyte or between two electrolytic solutions is represented by a semicolon (;) or by a single vertical line (l). Example Zn ; Zn2+ or Zn l Zn2+

3.Example: Ag, AgCl l Cl-In the example the (,) indicates that the Ag and AgCl should be taken together always to constitute an electrode.

4.The electrode on the left is written in the order the electrode and then the ion (Zn l Zn2+)

5.The electrode on the right is written in the order, ion and then electrode (Zn2+l Zn)

6.A double vertical line (ll) represents a salt bridge which has zero potential difference.Example:Zn l Zn2+ ll Cu2+ l Cu7.The combination of two single electrodes or two half cell constitute an cell

The left hand electrode is anode The right hand electrode is cathodeThe left hand half cell is cathode half cellThe right hand half cell is anode half cell

Example: Zn l Zn2+ ll Cu2+ l Cu(anode)(Cathode)8.At anode oxidation takes place Zn Zn2+ + 2eAt cathode reduction takes place Cu2+ + 2e Cu9.If E0cell of the reaction is positive, the cell reaction takes place and is said to be spontaneous.If E0cell of the reaction is negative, the cell reaction is not feasible and is said to be non-spontaneous.

1.2 Role of Salt bridge

The salt bridges are U-shaped and contain agar-agar gel containing KCl or MNO3. Some times NH4NO3 is also used. The reason for using these salts, is that the cations and anions have almost same speed and same transport number.

The salt bridge provides a passage for the flow of charges in the internal circuit and thus prevents the accumulation of charges. In the presence of salt bridge, the charges present flow towards the oppositely charged electrodes. The negative ions move from the copper electrode passes through the CuSO4 solution, salt bridge and zinc sulphate solution to zinc rod. The positive ions move in the reverse direction i.e., zinc rod to copper rod. Hence flow of electricity is maintained when the salt bridge is used.But when the salt bridge is not used the electrons released by the anode flow to the cathode through the external circuit. During this there occurs an accumulation of charges near the two electrodes, which prevents further flow of current. Hence the electrochemical change stops and current drops to zero.

Introduction

. Electrochemistry is branch of chemistry. This is based on properties of a solution when it is made part of an electrochemical cell. The principle of electrochemical cell deals with the conversion of energy from electrical to chemical and vice versa. This principle is used in the construction of cells and batteries which is a source of electrical energy. Refining of copper involving electrolysis is done to obtain high quality electric cables. The electroplated materials are now used in large scale in automobile industries, in jewelry, etc.. These provide both protection of metals from corrosion and aesthetic appeal. The extraction of the metals from its ores like sodium from common salt also involves electrolysis.

Not only the electrochemical concept is used for protecting the material but also used in explaining the cause of destruction of metal. For example, corrosion of steel beams used in the construction works like bridges, houses etc., occur through this principle. But methods like cathodic protection has been developed involving the same principle, so as to counteract and protect metals from corrosion.

Galvanic cellThese cells convert chemical energy into electrical energy.

These cells generate an electric current due to chemical reaction occurring within them. The redox reaction occurring in a galvanic cell is spontaneous and involves a decrease in the free energy system.

Example:Dry cell, Lechlanche cell, Daniel cell.

Electrolytic cellThese cells convert electrical energy into chemical energy.

In these cells chemical change is brought about when from outer source the current is passed into the electrolyte. This reaction is non-spontaneous but is feasible.

Example: Used in electroplating, electrotyping, electro-refining of metals.

Differences between Electrolytic cells and electrochemical cells.S.NoElectrolytic cellElectrochemical cell

1.In electrolytic cell electrical energy is converted into chemical energyIn electrochemical cell Chemical energy is converted into electrical energy

2.The negative charge is carried by cathodeHere the negative charge is carried by anode.

3.The positive charge is carried by anodeHere the positive charge is carried by cathode

4.The chemical reaction takes place during electrolysis is governed by the Faradays law of electrolysisThe e.m.f produced in the cell depends on the type of the electrode used and on the concentration of the electrolytes

5.Calumniator is used for the measurement of electricity passed during the electrolysis process Potentiometer is used for the measurement of e.m.f produced in the cell

6.Here the external energy is used for the supply of electrons to the cellHere electrons are drawn from the cell.

1.3 Reversible cell

A cell is said to be reversible when

1.there is infinitesimally small difference in the driving and opposing force.

2.it is possible to reverse any change taking place by applying a force infinitesimally greater than one acting on it.Therefore, in a reversible cell the reaction occurs only when the current is drawn from connector.

A reversible cell satisfies the following conditions:

1.If an external EMF equal to that of the cell is applied in opposite direction, no reaction should take place on either of the two electrodes.

2.If the external EMF is slightly greater than the EMF of the cell, the overall cell reaction gets reversed. In this case the current flows in the opposite direction.

ExampleThe Daniel cell is a reversible cell. It maybe represented by

Zn(s) Zn2+(1M) Cu2+(1M) Cu(s)

At standard conditions, in the Daniel cell with emf 1.1. volts, the cell reaction is Zn(s) + Cu2+(1M) Zn2+(1M) + Cu(s)

When an external emf is slightly greater than 1.1 volts the cell reactions gets reversed as follows

Cu(s) + Zn2+(1M) Zn(s) + Cu2+(1M)

The current will now flow in opposite direction.

1.4 Irreversible cell

Cells which do not obey the conditions of thermodynamic reversibility are called as irreversible cells.

An irreversible cell satisfies the following conditions.

1.Even if external EMF applied is equal to emf of the cell the reaction does not stop.2.Even when the external EMF applied is slightly greater than that of the cell, the reaction occurs in the same direction and does not get reversed.

Example:

Zn(S) H2SO4(soln) Cu(s)

The following cell reaction occurs

Zn + H2SO2 ZnSO4 + H2

Liberation of hydrogen occurs which results in irreversibility.

1.5 Electrode Potential

When a metal is in contact with a solution of its own ions at 25C, it may either undergo oxidation reaction or reduction reaction.

Fig 1.1

If it is a oxidation reaction

M(s) Mn+ (aq) + ne

The Mn+ leaves the metal electrode and enters into solution. This leaves the free electrons on metal electrode resulting in a negative charge on the electrode. This negative charge attracts the positively charged ions in the solution leading to the formation of a layer of positive ions around the metal electrode.

Similarly, if it is a reduction reaction Mn+(aq) + ne M(s)

The Mn+ ion from solution enter into the metal electrode resulting in positive charge on the electrode. This attracts the negative charged ions in the solution leading to the formation of a layer of negatively charged ions around the electrode.The formation of a layer of negative or positive ions around the metal electrode is called Helmholtz electrical double layer. Because of this layer a potential difference is set up between the metal and the solution. At equilibrium, this potential difference will become constant and this is known as standard electrode potential of metal.

Single Electrode Potential

Definition:Single electrode potential is the measure of tendency of a metallic electrode to lose or gain electrons, when it is in contact with a solution of its own salt.

The tendency of a metallic electrode to lose electrons is known as oxidation potential and the tendency of an electrode to gain electrons is known as reduction potential.

It is impossible to measure absolute value of single electrode potential because neither oxidation nor reduction takes place independently. It is possible to measure the potential difference between two electrodes with the help of the potentiometer. In other words, we can only measure electrode potential with reference to a reference electrode.

Standard Electrode Potential

DefinitionStandard electrode potential is the tendency of a metallic electrode to lose or gain electrons when in contact with 1M concentration of its own salt solution at 250C.

Electromotive Force and its Measurements

DefinitionElectromotive force or EMF of a galvanic cell is defined as the difference of electrode potential which causes the flow of current from one electrode to another when virtually no current is drawn from the cell.(NOTE: the difference between cell potential and EMF is, cell potential is positive while EMF can have positive as well as negative value.) MeasurementIt is necessary that no current should be drawn from the cell while measuring its EMF. This is possible only when the EMF of the cell under study is balanced by EMF of some other cell. This is done in Poggendorfs compensation method

Poggendorfs compensation methodIn this method the EMF of the given cell is compensated by the EMF of a known standard cell so that no current could flow in the circuit.

ConstructionIt consists of a potentiometer wire AB stretched over a meter scale. The cell C is a source of direct current. Its positive electrode is connected to end A while the negative electrode to the end B of the wire. X is the cell whose EMF is to be measured while the cell S is standard cell of known EMF. The positive electrodes of both the cells X and S are connected to the end A of the wire while their negative electrodes are connected to the two terminals of a double pole double throw switch P. The third terminal of the plug is connected to a sliding contact D through a galvanometer G.

GCell for which emf is to be measuredABDDGalvanometerSliding contactSXStandard cell-+C (Storage battery)RAdjustable resistance Measurement of emf of a cellP

Fig 1.2

WorkingStandard cell S is first brought into the circuit with the help of the double pole double throw switch and the null point D is found when no current passes through the circuit and galvanometer does not record any deflection. At this point, the EMF of the cell C is balanced by that of the cell S. If Ec is the EMF of the cell C and Es is the EMF of the cell S, then

(Ec / Es) = AB/ADNow, the cell S is removed from the circuit and the cell X is connected in the circuit. The bridge is again balanced and the null point D is found for no current flowing in the circuit. At this stage, the EMF of the cell C is balanced by that of the cell X. If EX is the EMF of the cell X, then

(Ec / Ex) = AB/AD

From the above equations

i.e.,

Since the EMF of the standard cell S is known and the length AD and AD can be directly read with the help of meter scale, the EMF of cell X under study can be obtained.

Standard cellPoggendroffs compensation method requires a standard cell whose EMF is standard and is precisely known. For this a saturated Weston cadmium cell is used. The standard cell has a very low temperature coefficient and its EMF does not change with time. These unique properties make it a standard cell.

RepresentationCd(12.5% in Hg) 3CdSO4.8H2O(Satd soln) Hg2SO4(s), Hg(l)

Electrode reactions

At anode :Cd(s) + SO42-(aq) CdSO4(s) + 2e

At cathode :Hg2SO4(s) + 2e 2Hg(l) + SO42- (aq)

Overall cell reaction :

Cd(s) + Hg2SO4(s) CdSO4(s) + 2Hg(l)

Construction

This type of cell consists of a H-shaped glass vessel. The bottom of the glass limbs are sealed with platinum wire. The cathode contains mercury and mercurous sulphate is placed over it. The anode consists of an amalgam of cadmium (Hg-Cd) over which some crystals of CdSO4. (8/3)H2O are placed. The rest of the part of the vessel is filled with saturated solution of CdSO4. The upper end of the arms is closed with corks. The purpose of solid crystals is to keep the electrolyte saturated at all temperatures. The emf of this cell is 1.01807 V at 20C.

+ SaturatedCdSO4 CorkHgSO4 HgCd-HgFig 1.3 The Weston Standard cellCdSO4 8/3 H2O SaturatedCdSO4

Applications of EMF measurements

1.Determination of solubility of sparingly soluble salts and its solubility products2. Determination of the valency of an ion3.Determination of pH4. Determination of standard free energy change and equilibrium constant.5.In potentiometric titrations.

1.Determination of solubility of sparingly soluble salts and its solubility productsThere are salts that are sparingly soluble in water, whose solubility is difficult to be determined by other methods. Example :Silver chloride (AgCl)

The solubility of these sparingly soluble salts can be determined by the EMF measurement by constructing a concentration cell.

Construction of the cell

Ag(s) AgCl(s) KCl (0.01N) AgNO3(0.01 N) Ag(s) (unknown)

In this, one of the Ag electrodes is placed in contact with 0.01N AgNO3 solution and the other electrode is in contact with 0.01N KCl solution. The two solutions are connected through a salt bridge containing saturated solution of ammonium nitrate. A drop of AgNO3 solution is added to the KCl solution. This forms a small amount of AgCl which is sufficient to give a saturated solution.

WorkingOne of the electrodes of this cell thus is in contact with a solution of silver ions of known concentration (0.01N). The other electrode is in contact with a solution of unknown concentration of silver ions furnished by the ionisation of sparingly soluble silver chloride formed.

The EMF of the above cell is given as

i.e.,

where, the only unknown factor is the C, the concentration of Ag+ ions furnished by AgCl in KCl solution. The emf of the cell can be measured and thus the concentration of AgCl C is calculated.

Multiplying the concentration of AgCl with the equivalent weight of AgCl (143.5), the solubility of AgCl is calculated in gm/litre.(OR)

The solubility product of silver chloride is given by

Ksp(AgCl) = C 0.01

The solubility (S) of the silver chloride is then given by

S =

2 Determination of the valency of an ionThe valency of the ions could be determined from the EMF of the concentration cells.

ExampleThe valency of the mercurous ion could be obtained as follows

Construction of the cell

Mercury / Mercurous nitrate (C1) // Mercurous nitrate (C2) / Mercury solution solution

The salt bridge contains the saturated solution of ammonium nitrate.

The EMF of the cell is given by

C2 and C1 are known concentration of mercurous nitrate electrolyte. The emf is obtained experimentally. Substituting these values in the above equation the n the valency of the mercurous ion is calculated.

3.Determination of pHThe pH could be obtained using any one of the indicator electrodes.

Example :Using a standard hydrogen electrode:

ConstructionA cell with a reference electrode and hydrogen electrode is constructed using a salt bridge.

Pt, H2(1atm), H+(c= unknown) // KCl(satd), Hg2Cl2(s), Hg

CalculationThe potential of the reference electrode i.e. calomel electrode is +0.2422 volts is known. The emf of the cell is obtained experimentally. Substituting the values in the following equation Ecell = Eright Eleft(1)

Thus the potential of hydrogen electrode used is calculated as follows.Eleft = Eright Ecell (2)

To obtain the pH, the potential of the hydrogen electrode calculated from eqn (2) is substituted in the following Nernst equation

By convention E0 of hydrogen is zero, therefore

Thus, substituting the value of the hydrogen electrode calculated from equation(2), the pH could be determined from the above equation.

By using glass electrode

A cell with a reference electrode and glass electrode is constructed using a salt bridge as follows.

Construction Pt, HCl(0.1N) / glass / test soln // KCl(satd), Hg2Cl2(s), Hg(l)

The salt bridge consists of KCl solution.

CalculationThe potential of the reference electrode i.e. calomel electrode is 0.2422volts is known. The emf of the cell is obtained experimentally. Thus substituting the values in the following equation

Ecell = Eright Eleft(1)

The potential of a glass electrode used is calculated as follows.

Eleft = Eright Ecell (2)

To obtain the pH, the potential of the glass electrode obtained from eqn (2) is substituted in the equation (3) or (4).

EG = E0G + 0.0591 log[H+]

EG = E0G - 0.0591 pH

. (3)

(OR) ..(4)

[ Where EG = Ecalomel Ecell and glass electrode is anode].

Thus the pH is calculated using the glass electrode.

4.Determination of standard free energy change and equilibrium constantThe standard free energy change (G0) of a reaction can be calculated from the using the standard emf from the following expression ( )G0 = nFE0Where, F = 96,500 coulombsE0 = Standard emf of the cell n = number of electrons involved

The equilibrium constant k of a reaction can be calculated from the following equation:

Where, E0 is the standard EMF of the cell k is the equilibrium constantn is the number of electrons involved in the reaction.

1.6 Electrochemical series

The electrochemical series is the arrangement of various electrode systems in the increasing order of their standard reduction potentials.

Importance of Electrochemical series1.The most active metals are at the top of the series.

2. The electrode system having negative values of standard reduction potentials act as anode when connected to a standard hydrogen electrode, while those having positive values act as cathode.3.On moving from the top to the bottom in the series, tendency to gain electrons i.e. the tendency for reduction to occur increases.Li will act as the strongest reducing agent.

Standard reduction potentials

Electrode reaction

Oxidised formReduced formStandard reductionPotential, E0, V

PotassiumK+ + e- K-2.925

Calcium Ca2+ + 2e- Ca- 2.87

Sodium Na+ +e- Na-2.714

Magnesium Mg2+ +2e- Mg-2.37

Aluminum Al3+ 3e- Al-1.66

Zinc Zn2+ + 2e- Zn-0.763

Chromium Cr3+ + 3e- Cr-0.74

Iron Fe2+ + 2e- Fe-0.440

Lead Pb2+ +2e- Pb-0.126

Hydrogen 2H+ +2e- H20.00

Copper Cu2+ +2e- Cu+0.337

Iodine I2 + 2e- 2I-+0.5355

Mercury Hg22+ + 2e- 2Hg+ 0.789

Silver Ag+ +e- Ag+0.7991

Bromine Br2(l) + 2e- 2Br-+1.0652

PlatinumPt2+ + 2e- Pt+1.2

GoldAu+ + e- Au+1.68

Increasing ease of reduction of reactantFluorine F + 2e- 2F-+ 2.87

Increasing ease of oxidation of reactant

1.92 Engineering Chemistry1.91Electrochemistry

4.On moving from the bottom to the top in the series, the tendency to lose electrons i.e tendency for oxidation increases.F2 will act as the strongest oxidizing agent.5.The systems placed above hydrogen are stronger reducing agents, while the one below the hydrogen are weaker reducing agents.

APPLICATIONS OF ELECTROCHEMICAL SERIES1.To predict the relative oxidizing and reducing powerIn general, Oxidizing agents have positive Eo Values Higher the positive value, stronger will be the oxidizing agent that is placed below the hydrogen.

Reducing gents have negative Eo values Higher the negative value stronger will be the reducing agent that are placed above the hydrogen

2.To predict the spontaneity of any redox reactionFor any spontaneous reaction G = -veSince G = -nFEcell Ecell should be positive for spontaneous reaction.

The Ecell is calculated from the standard redox potentials by using the relationEcell = Ecathode Eanode

3.Displacement of hydrogen by metals from dil. acidsThe displacement of hydrogen from an acid is a reduction process i.e.

2H+ + 2e H2(g) Eo = 0.00 volts

This is brought about by a metal whose reducing power is greater than that of hydrogen. Hence any metal placed above hydrogen in electrochemical series i.e. with negative std. reduction potential can displace hydrogen from a dil. acid.Example :Zn, Mg, Ni.

4.Displacement of metals from salt solutionA metal with lower value of reduction potential can displace another metal with higher value of reduction potential.

Example: Zinc displaces copper from its solution spontaneously, but reverse reaction is not feasible.

Zn + Cu2+ Cu + Zn2+

5.Electropositive character of metals (tendency to form positive ions)The metals at the top of the electrochemical series have the greater tendency to lose electrons. So, these metals at the top, possesses the highest electropositive character. The value of the std. reduction potentials increases in going down the series, which implies that the electropositive character of metals decreases.Example: Alkali metal > Hg, Ag, Au

6.Thermal stability of the metal oxideHigher the electro positive character, greater is the stability of its oxide. Hence the stability of oxides of metals decreases in going from top to bottom of the series. Example: Alkali metal oxides like Na2O is more stable than Ag2O, HgO etc.

7.Calculation of standard cell potentials (EoCell)The EoCell is calculated using

EoCell = EoCathode - EoanodeWhere EoCathode is the standard reduction potentials of cathodeEoanode is the standard reduction potentials of anode.

8.To calculate equlilbrium constants Go = - RTlnKeqm Go = - nFEoEo = (RT/nF) ln keqm ln keqm = nFEo/RTFrom E0 equilibrium constant can be calculated.

9. Prediction of correct metallurgical methodEo values of Cu, H2O and Al are (+)0.34, (-)0.83and (-) 1.66V. Cu gets easily reduced when compared to H2O and H2O get reduced more easily than Al. Thus, copper is produced by electrolysis of aqueous CuSO4 and aluminium cannot be produced form aqueous solution of aluminium

This is because when Al3+(aq) is electrolyzed, only H2O will be electrolysed and not Al3+

1.7 Thermodynamics of a reversible cell

NERNST EQUATION FOR ELECTRODE POTENTIAL

Nernst equation is used to calculate the electrode potential of the electrode assembly under any given condition. It tells us the effect of electrolyte concentration on electrode potential.

Nernst equation is given by

This is applicable to a general redox reaction. Consider a equation

aA + bB cC +dD

For a general reversible chemical system, according to Vant-Hoffs isotherm and Gibbs free energy, the relation between the free energy change G and its equilibrium constant is expressed as

. (1) But , G0 = ( ) RTln K (2)

Substitute equation (2) in (1)

. (3) where, G0 = standard free energy change .

For a reversible reaction, the decrease in free energy (G ) is given by G = () nFE ...(4)

G0 = () nFE0

(or) ()G0 = nFE0 (5)

Where , G0 is the standard free energy change . n is the number of electrons F is Faraday = 96,500 Coulombs of electricity E0 is the standard potential

Substituting the equation (4) and (5) in equation (3), we get

Rearranging the above equation

(6)

Consider a reduction reaction Mn+ + ne M Applying eqn. (6) to an electrode reduction equation

Since [M] = 1 for the solid metal

The above equation may be written as,

Where, R = 8.314 J/K/mole; T = 298C; F = 96500 coulombs.

Substituting the value of R, T and F

The above equation is known as Nernst equation.

Nernst equation for a reduction reaction

Consider an oxidation reaction M Mn+ + ne

Similarly for an oxidation reaction Nernst equation is

Nernst equation for an oxidation reaction

Applications of Nernst equation1. To study the effect of electrolyte concentration on electrode potential2. For calculation of the potential of a cell under non-standard conditions3. Determination of unknown concentration of one of the ionic species in a cell when a concentration of other ionic species is known.4. The pH of a solution can be calculated.5. To find the valency of an ion or the number of electrons involved in the electrode reaction.

1.8 Types of Electrodes

Electrodes are of two types namely:Reference electrode Indicator electrode

We know that absolute value of a single electrode potential cannot be measured, because the oxidation or reduction does not take place independently. The potential difference between two electrodes can be measured only by combining two electrodes to form a complete cell and by using the potentiometer.Thus electrode potential can be measured with reference to a reference electrode.

Reference electrodes:Reference electrode is an electrode which has a constant emf or constant potential, with which we can compare the potentials of other electrodes.

There are two types of reference electrode 1)Primary reference electrode Example : Standard hydrogen electrode (SHE)

2)Secondary reference electrodeExample : Calomel electrode, Ag-AgCl electrode, Glass electrode, Quinhydrone electrode etc.

PRIMARY REFERENCE ELECTRODE Standard Hydrogen Electrode (SHE)

Fig 1.4

Standard hydrogen electrode is an example of primary reference electrode. It consists of platinum foil or wire coated with platinum black dipped into a molar (1M) solution of H+ ions. The platinum wire is surrounded by an outer tube into which the hydrogen enters through a side inlet and escapes at the bottom through the test solution. Hydrogen gas at 1 atmosphere pressure is passed through it continuously at 2980K. Although an open vessel as shown in diagram, in practice the electrode will be used in a stoppered flask with a suitable exit for hydrogen. Its potential is arbitrarily taken as zero at all temperatures. SHE is a reversible electrode i.e, depending on the nature of another electrode to which it is connected, i.e., this electrode may act as anode or cathode.Representation: Pt, H2(g), (1atm) /H+(1M)

Reaction As Anode : H2(g) 2H+(aq) + 2e

As cathode : 2H+(aq) + 2e H2(g) Electrode Potential of SHE : arbitrarily considered as zero

SECONDARY REFERENCE ELECTRODEDue to the difficulties encountered with SHE, secondary reference electrode is developed

Definition:The electrodes whose electrode potentials are found using the primary reference electrode and are always constant at given conditions are used instead of primary reference electrode.

As at first the potential is obtained from the primary reference electrode which is then being used as a reference to another (to a third) electrode in order to obtain its (third electrodes) electrode potential this is called as secondary reference electrode.

Example:Calomel electrode, Ag-AgCl electrode, quinhydrone electrode etc.

Calomel electrode It is a secondary reference electrode

ConstructionIt consists of a glass tube provided with a bent side tube A and another side tube B containing a rubber tubing that could be closed with a screw clip. The glass tube consists of layer of pure mercury in the bottom. The surface of the mercury is covered with a paste of mercurous chloride and mercury in potassium chloride.

Fig 1.5 Calomel ElectrodeABA saturated solution of KCl is filled in the tube. A platinum wire sealed through a glass tube C, is dipped in the mercury layer for electrical contact. The glass tube C contains a little mercury into which an amalgamated copper wire dips. The side tube is used to make electrical contact with other electrode through a salt bridge.

Representation: Hg(l), Hg2Cl2(S) / KCl(satd)Reaction Anode2Hg(l) Hg22+(aq) + eHg22+(aq)+ Cl- Hg2Cl2(S)------------------------------------2Hg(l) + Cl-(aq) Hg2Cl2(S) + e

Cathode

Hg2Cl2(S) + e

2Hg(l)Electrode potential : E = E0Cl-/Hg2Cl2/Hg - (RT/nF) ln [Cl-]

Working

Electrode As anodeHg liberates electrons and sends Hg22+ ions into solution. The Hg22+ would combine with Cl- ions furnished by the KCl forming sparingly soluble Hg2Cl2. The result is a fall in concentration of the chloride ions in the solution with electrons transferred to metal.

Electrode As cathodeIf the electrode involves reduction reaction, the Hg22+ ions furnished by the sparingly soluble Hg2Cl2 would be discharged at the cathode. Hence more and more of calomel would pass into solution. This results in an increase in concentration of the chloride ions. Thus the potential of the calomel electrode depends on the activity of the chloride ions and increases as the activity of the chloride ions decreases. The electrode is reversible with respect to chloride ions.

Potential of saturated calomel electrode (SCE) is (+)0.2422 volts, Normal calomel electrode (NCE) i.e. 1N KCl is (+)0.28 volt and Decinormal calomel electrode (DNCE ) i.e. 0.1N KCl is (+)0.3338 Volts

Advantages of calomel electrode1.It is simple to construct2. Results of cell potential measurements are reproducible and stable over a long period and do not vary with temperature.

Silver Silver chloride electrode

ConstructionIt consists of a silver wire or a silver plated platinum wire coated electronically with a thin layer of silver chloride, dipping into a potassium chloride solution of known concentration. The potentials of the 0.1M and saturated Ag-AgCl electrode at 25C with respect to the normal hydrogen electrode are 0.29. and 0.199 volt respectively.

Representation : Ag(S), AgCl(s) / KCl(satd)Reactions : At anode

Ag(s) Ag+ + e-

Ag+ + Cl-(aq) AgCl(s)

Ag(s) + Cl-(aq) AgCl(s) + e-

Electrode potential = E = E0 +

At cathode

AgCl(s) Ag+ + Cl-(aq)

Ag+ + e- Ag(s)

AgCl(s) + e- Ag(s) + Cl-(aq)

This result in the removal of chloride ions from the solution with electrons transferred to the metal. The overall reaction involves only the concentration of Cl- ions as the variable. Thus the electrode is reversible with respect Cl- ion.

Electrode potential: E = E0Cl + (RT/nF) ln [Cl-]

Working

Electrode as anode

Since the sparingly soluble salt AgCl is in contact wit Cl- ion solution, the solution is saturated. Some Ag passes from the electrode into the solution

Ag Ag+ + e

The Ag+ ions produced combine with Cl- of KCl to form AgCl(s). This result in the removal of chloride ions from solutions with electrons transferred to metal.

Electrode as CathodeIf the electrode involves reduction reaction, the Hg22+ ions furnished by the sparingly soluble Hg2Cl2 would be discharged at the cathode. Hence more and more of calomel would pass into solution. This results in an increase in concentration of the chloride ions.

Thus the potential of the calomel electrode depends on the activity of the chloride ions and increases as the activity of the chloride ions decreases. The electrode is reversible with respect to chloride ions.

UsesThis electrode is used to measure the concentration of the chloride ions.

Quinhydrone electrode :Quinhydrone is a compound of quinine and hydroquinone and in solution is decomposed into equimolar quantities of these substances.C6H4O2.C6H4(OH)2 = C6H4O2 + C6H4(OH)2 Quinhydrone quinine hydroquinoneThis consists of platinum wire dipped in an equimolar solution of hydroquinone and quinone. The electrode is represented as Pt(s) QH2(aq), Q(aq)_, H+(aq)Where QH2 and Q represent hydroquinone C6H4(OH)2 and Quinone C6H4O2 respectively .The electrode reaction involved is

C6H4O2(aq) + 2H+(aq)+ 2e C6H4(OH)2(aq)

Q(aq) + 2H+(aq) + 2e- QH2(aq)

The potential developed on platinum electrode is given by Nernst equation as

In practice the [Q] / [QH2] is maintained constant at unity by saturating the solution with the substance quinhydrone, which is equimolar mixture of quinone and hydroquinone. The middle term reduces to zero.

(since ln[H+]2 = 2ln[H+] )

E = E0 + 0.0591 log [H+] at 250 C

(or) E = E0 0.0591 pH

Where E0 is the standard electrode potential of quinhydrone electrode = 0.6998V.

(or) E = 0.6998 0.0591 pH

Determination of pH of a solution using quinhydrone electrodeTo find the pH of a solution, quinhydrone electrode is connected to a saturated calomel electrode.

This can be represented as SCE // H+ (unknown) / Q(aq) , QH2(aq) / Pt(s) Ecell = Eright - Eleft

Ecell = 0.6998 - 0.0591pH 0.2422

Merits of quinhydrone :1. It attains equilibrium rapidly2. It is not so readily incapacitated as the hydrogen electrode3. It has low internal resistance4. It is free of salt errors and non-reducing errors.5. It can be used for pH measurements in non-aqueous media.

Demerits of quinhydrone1. It cannot be used in solution of pH greater than 8.2. It is not stable for long time at high temperatures.

Indicator electrode or Ion selective electrodeThe electrode in which the potential changes with change in concentration is known as indicator electrode. i.e. the indicator electrode of a cell is one whose potential is dependent upon the activity of a particular ionic species whose concentration is to be determined.Example : glass electrodeGlass electrodeThe glass electrode is the most widely used hydrogen ion responsive electrode

Principle : When a glass membrane is immersed in a solution, a potential is developed between the two surfaces of the membrane which is a linear function of the hydrogen ion concentration of the solution i.e. pH value

ConstructionThe glass tube of the glass electrode is made of Lithium based glasses of composition SiO2 63%, Li2O 28%, Cs2O 2%, BaO 4%, La2O3 3%. Its melting point should be low and electrical conductivity is high.

Fig 1.6 Glass electrode

The electrode consists of a glass bulb. This is filled with a solution of hydrochloric acid (usually 0.1M). A silver- silver chloride electrode is inserted into it. The upper end of the electrode must be sealed. This is done in order to ensure a constant concentration of the inner hydrochloric acid solution. The internal hydrochloric acid is maintained at constant concentration, so that 1. the potential of the silver silver chloride electrode is constant 2. the potential between the hydrochloric acid and the inner surface of the glass bulb is also constant. Hence the only potential which can vary is that existing between the outer surface of the glass bulb and the test solution into which it is immersed.

So the overall potential of the electrode is governed by the hydrogen ion concentration of the test solution.

Representation : Ag(s) , AgCl(s) /HCl (0.1M) /Glass

WorkingTo measure the hydrogen ion concentration of a solution the glass electrode is immersed in the test solution and is combined with a reference electrode a calomel electrode. Thus the cell representation is

Internal std soln/ ref. electrode 2Ref. electrode1 / test soln

External reference electrodeIon selective electrode

Ag(s),AgCl(s) /HCl(0.1M) /Glass/ Test soln // KCl(satd) /Hg2Cl2(S), Hg(l)

Measurement of pH using glass electrode

Fig 1.7

For pH measurements, E0G is first measured by dipping glass electrode in buffer solutions of known pH values. Once E0G for a particular arrangement is determined, the electrode is then placed in the solution of unknown pH, the emf of the cell is measured and using the equation, its pH can be calculated. Ecell = Eright Eleft = 0.2422 (E0G - 0.0591 pH) = 0.2422 E0G +0.0591 pH

(or)

The emf of the cell is expressed as E = K + (RT/F) ln[H+]E = K + 0.0591pH

In these equations K is constant and partly dependent on the nature of the glass used, individual character of each electrode and its value may vary with time. This variation of K with time is related to the existence of an asymmetry potential. The asymmetric potential is due to the differing responses i.e. the strain due to changes of hydrogen ion activity that is developed at the inner and outer surfaces of the glass bulb. So glass electrode has to be standardized frequently placing in a buffer solution of known hydrogen in activity.

Advantages of Glass Electrode 1. It can be used in any solutionsExample: in alkaline, oxidizing, turbid, coloured solutions etc2. It is most convenient and simple to use3. It is not easily poisoned4. Equilibrium is easily achieved5. A small quantity of solution is sufficient for the determination of pH.

Disadvantages1. It gives erroneous results when used in poorly buffered solutions which are nearly neutral.2. It is found to be too sensitive to Na+ ions at very high pH values greater than pH =9 resulting in alkaline error

Types of electrodeType (I) electrodeMetal - Metal ion electrodeMetal amalgam electrodeGas ion electrodeType (II) electrodeMetal Metal insoluble salt electrode reversible to anion Type (III) electrodeMetal Metal insoluble salt electrode reversible to cation Type (IV) electrodeRedox electrode

Type (I) electrode

1.Metal - Metal ion electrode

Pure metal (M) is in contact with a solution of its cation (Mn+)

Representation : Mn+(aq) / M (s)

Reaction : Mn+(aq) + ne M (s)

Electrode potential :

Example Zinc electrode: Zn(s) /ZnSO4(aq) Copper electrode:Cu(s) /CuSO4(aq)

2.Metal amalgam electrodeMetal amalgam is in contact with a solution of metal ion

Representation : Mn+(aq) / M(Hg)(s)

Reaction : Mn+(aq) + ne M(Hg)(s)

Electrode potential :

The active metals which cannot form electrodes when dipped in their salt solution are converted into metal amalgam and used. This is because the activity of the metal is lowered when mixed with mercury.

Example : Sodium electrode : Na(Hg)(s) /Na+(aq)

3.Gas Ion electrodeInert metal is dipped in a solution containing ions to which the gas is reversible. The gas is continuously bubbled through the solution.

Representation : Pt(s) , X2 (p = y atm) / X-(aq)Reaction : X2 (p) + ne X-(aq)

Electrode potential :

Example Std. hydrogen electrode: Pt(s) , H2(g) /H+(aq) Chloride electrode : Pt(s) , Cl2(g) /Cl-(aq)

Metal Metal insoluble salt electrode reversible to anion (type II)Metal (M) is covered by layer of sparingly soluble salt (MX) immersed in a solution containing a common anion.

Representation: X-(satd) // MX(S)/ M(s)Reaction MX(S) + e M(s) + X-(aq)

Electrode potential :

Example : Calomel electrode : Hg(l) /Hg2Cl2(S), KCl(satd)Silver electrode : Ag(s) /AgCl(S), KCl(satd)

Metal Metal insoluble salt electrode reversible to cation (type III)Metal is in contact with one of its sparingly soluble salt together with another insoluble salt having the same anion, in the solution of a salt having a common cation as that of the latter salt.

Representation: M2X2 (soln), M2X1(S), M1X1(S)/ M1(s)

Reaction M1X1(S)+ M2n+(aq) + ne M1(s)+ M2X1(S)

Electrode potential :

E = E0 M2X2 (soln), M2X1(S), M1X1(S)/ M1(s) + RT ln [M2n+(aq)] nF

Example Pb(s) PbSO4(s), SrSO4(s), SrCl2(soln)

Redox electrode (type IV)An inert material is (eg. Pt) is dipped in a solution containing ions in two oxidation states of the substances.

Representation:Pt/Mn1+(C1), Mn2+(c2)Reaction : Mn1+(C1) Mn2+(c2)+ e

E Pt/Mn1+(C1), Mn2+(c2) = E0 Pt/Mn1+(C1), Mn2+(c2) - RT ln [Mn2+] nF [Mn1+]Electrode potential :

Example : Iron electrode : Pt/Fe2+(C1), Fe3+(C2)

GALVANIC CELL (or) VOLTAIC CELLThese cells convert chemical energy into electrical energy on account of some chemical or physical change taking place within the cell. The driving force to cause the physical or chemical change is the decrease in free energy.

Example : Daniel cell can be obtained by joining an oxidation electrode Zn/Zn2+ to a reduction electrode Cu2+/Cu as shown in Fig .

1.1VZn strip (Anode)Cu strip (Cathode) + Zn SO4 solutionCu SO4 solutionPorous plateDaniel cellFig 1.8 ZnSO4 and CuSO4 solution in direct contact through porous plate

1.1VZn strip (Anode)Cu strip (Cathode) + Zn SO4 solutionCu SO4 solutionGalvanic cellFig 1.9 ZnSO4 and CuSO4 solution are in Electrical communication through a salt bridgeSalt bridgeKCl solution

The solutions in which two electrodes are immersed are made in communication with each other by a direct contact through a porous diaphragm or through a salt bridge .The two electrodes are connected at the outer circuit to a device which utilize the electrical energy produced . The cell is represented as

Zn(s) ZnSO4(aq) CuSO4(aq) Cu(s)

Anodic reaction : Zn Zn2+ + 2e- (Oxidation )

Cathodic reaction: Cu2+ + 2e- Cu (Reduction )

Cell reaction: Zn + Cu2+ Zn2++Cu (redox reaction )

The zinc metal acts as anode because it has greater potential (0.76 V ) than copper (+0.34 V ) . The emf of the cell is 1.1 volts .

Galvanic cells are of the following two types

1. Chemical cell2. Concentration cell

CHEMICAL CELL OR FORMATION CELLIn this type, the electrical energy is produced on account of a chemical change within the cell.

Ex.Daniel cell, Lechlanche cell, dry cell

Classification1.Chemical cells with transference2.Chemical cell without transference.

CHEMICAL CELL WITH TRANSFERENCEThe two electrodes are dipped in the electrolyte containing their ions. These electrolytes are in direct contact but are separated by a porous diaphragm. This creates a liquid Junction. Since the ions on the two sides of the liquid junction are of equal velocities, their moving across the junction sets up potential called liquid junction potential. It is represented by Ej. This type of cells involve three types of potential oxidation potential at the anode, reduction potential at the cathode and liquid junction potential. However the Ej could be minimized or even eliminated through the use of a KCl salt bridge.

Representation : M1/M1n+(aq) M2n+(aq) /M2

Reaction Anode : M1(s) M1n+ (aq) + 2e

Cathode :M2n+ (aq) + 2e M2(s)

Cell reaction : M1(s) + M2n+ (aq) M1n+ (aq) + M2(s)

Cell potential Ecell = Eanode + Ecathode+ Eliquid junction

Example : Zn(s) /Zn2+(aq) Cd2+(aq) / Cd(s)

CHEMICAL CELLS WITHOUT TRANSFERENCEThis type of cell consists of two electrodes which are dipped in same solution. One of the electrode is reversible to cation of the electrolyte and the other electrode is reversible to the anion of the electrolyte. Since both electrodes are dipped in same solution the liquid potential is absent. There is no transport of ions so it is called as chemical cell without transference.Representation:Pt(s) , X2(g) / MX(aq) / M(s)

Example : Pt(s) , H2(g) / HCl(aq), AgCl(s) /Ag(s)H2 is reversible to the H+ ions and Ag/AgCl electrode is reversible to the Cl- ions.

ReactionsAnode :(1/2)H2(g) H+(a1) + eCathode :AgCl(s) + e Ag(s) + Cl-(aq)

Cell reaction: (1/2)H2(g) + AgCl(s) H+(aq) + Ag(s) + Cl-(aq)

Cell Potential

ECell = E0Ag / Agcl(s) / Cl- - RT ln aH+ . a Cl- nF (pH2)1/2

1.9 Concentration cell

The electrical energy is produced due to the differences in the concentration of either the electrolyte or the electrode.The physical transfer of matter involves a decrease in the free energy of the system. Thus the process is spontaneous. The electrical energy arises due to the transfer of matter from the solution / electrode of the higher activity to that of lower activity

Types of concentration cell1.Electrode Concentration cell2.Electrolyte concentration cell

ELECTRODE CONCENTRATION CELLImmerse two electrodes of the same material but of different activities (concentrations) in the same electrolyte. The electrical energy is produced due to the transfer of matter from one electrode to other. (i.e. higher activity to that of lower activity electrode)

In these cells there is no liquid junction potential as these cells use only one electrolyte.

Types of electrode concentration cell 1.Amalgam electrode concentration cell2.Gas electrode concentration cell

i.e in electrode concentration cells the electrodes used are either amalgam electrodes or gas electrodes.

Amalgam electrode concentration cellTwo amalgam electrodes of different activities of the same metal is immersed in a single common electrolyte thus forming this cell.

RepresentationM(Hg)(a1) Mn+(aq) M(Hg)(a2)

Cell reactionAnode :M(Hg) (a1) Mn+(aq) + e

Cathode :Mn+(aq) + e M(Hg) (a2)

Cell reaction : M(Hg) (a1) M(Hg) (a2)

Thus a transfer of metal occurs from the anode to cathode giving rise to emf . Cell potential Ecell = Eright Eleft

Thus the emf of the amalgam electrode concentration cell depends on the ratio of the activities of the metal amalgam taken in two electrodes. It does not depend on the concentration of H+ ions in the solution. Thus, the electrode with higher gas pressure should be made anode.

Example : Ag(Hg) (a1) / AgNO3(aq) / Ag(Hg) (a2)

GAS ELECTRODE CONCENTRATION CELL

Two gas electrodes of different activities of the same gas are immersed in a single common electrolyte thus forming this cell.

RepresentationPt(s) , H2 (g) (p1) / HCl(aq)/ H2(g) (p2), Pt(s)

Cell reaction Anode : (1/2)H2(g) (p1) H+(aq) + e

Cathode : H+(aq) + e (1/2)H2(g) (p2)

Cell reaction : (1/2)H2(g) (p1) (1/2)H2(g) (p2)

Cell potential Ecell = Eright Eleft

Thus the emf of the gas electrode concentration cell depends on the ratio of the partial pressure of the gas taken in two electrodes. It does not depend on the concentration of H+ ions in the solution. Thus, the electrode with higher gas pressure should be made anode.

Electrolyte concentration cells It consists of two identical electrodes immersed in into two solutions of same electrolyte at different concentrations. The electrodes of such cells are reversible to one of these ions of the electrolyte. In this there is transfer of matter from electrolyte of higher activity to electrolyte of lower activity. Due to this transfer emf is generated.

Example :

Electrolysis of HCl is carried out by passing electricity in an aqueous solution of HCl through two platinum electrodes. During this electrolysis process, chlorine is evolved at the anode and hydrogen is evolved at the cathode. ( Fig ).

1 N HClPlatinum electrodes Fig 1.10 Electrolysis

When electric current is passed in HCl solution the following reactions take place at the electrodes.

At cathode : The electrode at which reaction occurs is called cathode.

H+ions will move towards the cathode, where H+ ions gets reduced to hydrogen.

2H+ + 2e H2

At anode :The electrode at which oxidation takes place is called the anode.

2Cl Cl2 + 2e

Types of Electrolyte concentration cells

1. Electrolyte concentration cells without transference2. Electrolyte concentration cells with transferenceElectrolyte concentration cells without transferenceTwo chemical cells with no liquid junction potential are combined together in such a way that they oppose each other. The two electrolytes are thus not in direct contact with one another.

Example:Pt, H2(g)(1atm) HCl(a1), AgCl(s) Ag(s)

Pt, H2(g)(1atm) HCl(a2), AgCl(s) Ag(s)

These two cells are combined together so that they oppose one another.

Representation

Pt,H2(g)(1atm) HCl(a1),AgCl(s) Ag(s) - Ag(s) AgCl(s),HCl(a2) H2(g)(1atm),Pt

Half cell reactions

Left half cell reaction (1/2) H2 + AgCl(s) Ag(s) + HCl(a1)

Right half cell reaction Ag(s) + HCl(a2) (1/2) H2 + AgCl(s) ________________________________Overall reaction HCl(a2) HCl(a1)

This indicates there is transfer of HCl from one concentration to other. Since the two half cells oppose each other, the net EMF of the cell is given by

Ecell = Eright Eleft

Thus the emf of the cells depends on the ratios of the concentrations/ activities of the two electrolyte solutions used in the cell.

Electrolyte concentration cells with transferenceIt consists of two identical electrodes immersed in into two solutions of same electrolyte at different concentrations. The electrodes of such cells are reversible to one of these ions of the electrolyte. The two electrolytes are kept in direct contact with one another. Thus a liquid junction potential develops at the interface of two electrolytes.

Electrolyte concentration cells with transference and having cation as the reversible ion.

RepresentationM(s) MX(a1) MX(a2) M(s)Reactions Left half cell reactionM(s) Mn+(a1) + ne

Right half cell reaction Mn+ (a2) + ne M(s) ___________________Overall cell reaction Mn+ (a2) Mn+ (a1)

The overall reaction indicates that only the Mn+ from the electrolyte of concentration C2 is transferred from the electrolyte of concentration C1.

But in actual as the two electrolytes are in direct contact, both the ions present in the electrolyte also migrate. Since the transport numbers of cation and anions are not equal, the relative migration of the both the ions across the liquid junction contribute significantly to emf.

EMF of the electrolyte concentration cells with transference and having cation as the reversible ion

Pt, H2(g)(1atm) HCl(a1) HCl(a2) H2(g)(1atm), Pt

EMF of the electrolyte concentration cells with transference and having anion as the reversible ion Ag (s), AgCl(s), HCl(a1) HCl(a2) AgCl(s), Ag

Hence in both the cases, when the electrodes are reversible to cation or anion, the emf of the cell depends on the transport number of the ion other than to which the electrode is reversible. Therefore, such cells are used in the determination of transport number of the ions.

1.10 Conductometric Titrations

Conductance measurements are frequently employed to find the endpoints of acid-alkali and other titrations.

(NOTE: The titrating solution in the burette is usually five times stronger than the solution taken in the conductivity vessel. This is done inorder to minimize the change in the volume in the conductivity cell due to addition of burette solution)Experimental methodsThe measurement of electrical conductance of a solution amounts to determination of electrical resistance of the solution because conductance is just the reciprocal of resistance. The wheatstone bridge is generally employed for this purpose. This is shown in figure. If the ratio of the arms is equal i.e., R3 = R4 , then the resistance of the cell is equal to that taken from the resistance box R2, when the bridge is balanced. It is then possible to read directly the resistance of the cell.

The reciprocal of this reading is proportional to the specific conductance. This conductance is then plotted against the volume added and by titration curves, the end point may be calculated.

ApparatusThe solution whose conductivity is to be determined, is placed in the conductivity cell. The cell is connected to the resistance box R on one side and to a long wire AB on the other side. I is the induction coil from which the current is led. On moving the sliding contact along the wire AB, a point is reached at which the galvanometer reading is zero.

Resistance of the solution = Length BDResistance of RLength AD

Conductance = (1/Resistance of the solution)

Fig 1.11

In order to use the titrations for high precision work the wheat stone bridge is replaced in the form of visual detector or (telephone earpiece). In this apparatus the alternating current is used. The apparatus shown in fig. is used.

Here the A.C is reduced to about 3 to 5 volts by means of transformer T. The galvanometer G and a rectifier (D) is used. The 800 ohm resistances A and B are used. The rectifier D may be rectifying crystal, a copper oxide rectifier or suitable vacuum tube circuit giving rectification and amplification. The solution to be titrated is placed in the cell. The resistance A and B adjusted and then the current is recorded in G. The titration is now carried out and the galvanometer readings are plotted against the volume of the titrant added.

Fig 1.12

The end point is determined by the point of intersection of two parts of the curves.

ACID - BASE CONDUCTIVITY TITRATIONS

Principle of acid base conductivity titrationElectrical conductance depends upon the number and mobility of ions.

ProcedureThe acid is taken in the conductivity vessel and alkali in the burette. Alkali is added gradually and the conductance is noted. The conductance is plotted against the volume of alkali added. This plot will have two straight lines AB and CD. The point of intersection X of these two lines gives the volume of alkali required for neutralisation .Theory of Strong acid vs Strong base reaction

Increase in conductance due to addition of OH-Decrease in conductance due to removal of H+End pointInitial conductance due to H+Volume of alkali added Titration of a strong acid vs strong baseBCDAConductance

Fig 1.13

The initial conductance is of hydrochloric acid which is due to the presence of hydrogen and chloride ions. As alkali is added gradually, the hydrogen ions are replaced by slow moving sodium ions.

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O

So on continued addition of sodium hydroxide, the conductance will sharply decrease until the acid has been completely neutralised.

Subsequent addition of alkali will result in introducing the fast moving OH- ions. The conductance therefore after reaching a certain minimum value, will begin to sharply increase.

NaOH Na+ + OH-

Theory of Weak acid vs Strong base reaction

Volume of alkali added Titration of a weak acid vs strong baseConductanceBCDA Increase in conductance due to addition of OH-Increase in conductance due to formation of highly ionized saltEnd pointInitial conductance is low due to poor dissociation of weak acids

When a weak acid (acetic acid) is used the initial conductance is low because the weak acids have poor dissociation. As alkali is added gradually, the acid combines with alkali to yield highly ionized salt. (Acetic acid forms sodium acetate). This contributes to the increase in conductance.

CH3COOH + Na+ + OH- Na+ + CH3COO- + H2O

On continued addition of sodium hydroxide, the conductance will slightly increase until the acid has been completely neutralised.

Subsequent addition of alkali will result in introducing the fast moving OH- ions. So there is sharp increase in conductance. NaOH Na+ + OH-

Theory of Strong acid vs Weak base reaction

Titration of a strong acid vs weak baseConductance No change in conductance due to poor dissociation of weak base Decrease in conductance due to removal of H+End pointInitial conductance due to H+Volume of alkali added BCDA

Fig 1.14

The initial conductance is of hydrochloric acid which is due to the presence of hydrogen and chloride ions. As weak alkali is added gradually, the hydrogen ions are replaced by slow moving ammonium ions.

H+ + Cl- + NH4OH NH4+ + Cl- + H2O

On continued addition of ammonium hydroxide, the conductance will sharply decrease until the acid has been completely neutralised.

Subsequent addition of weak alkali will not have appreciable change in the conductance as the weak base has poor dissociation.

Theory of Weak acid vs Weak base reaction

No change in conductance due to poor dissociation of weak baseIncrease in conductance due to formation of highly ionized saltEnd pointInitial conductance is low due to poor dissociation of weak acidsBCDA

Conductance

Volume of alkali added

Fig 1.15Titration of a strong acid vs strong base

When a weak acid (acetic acid) is used the initial conductance is low because the weak acids have poor dissociation. As alkali is added gradually, the acid combines with alkali to yield highly ionized salt. (acetic acid forms ammonium acetate). This contributes to the increase in conductance.

CH3COOH + NH4OH CH3COO- + NH4+ + H2O

On continued addition of ammonium hydroxide, the conductance will increase until the acid has been completely neutralized.

Subsequent addition of weak alkali will not have appreciable change in the conductance as the weak base has poor dissociation.

NaOH Na+ + OH- Theory of Mixture of weak and strong acids vs strong base reactions

BCDA Increase in conductance due to addition of OH-Decrease in conductance due to removal of H+End pointInitial conductance due to H+ Increase in conductance due to formn. of highly ionized saltEnd pointVolume of alkali added Fig 1.16 Titration of a strong acid vs strong baseConductanceEF

The plot resembles the combination of the strong acid against strong base and weak acid against strong base.

The strong acid first gets titrated first. Only after the strong acid is neutralised, the weak acid will be neutralised.

Hence the initial conductance is due to the presence of hydrogen and chloride ions of hydrochloric acid. As alkali is added gradually, the hydrogen ions are replaced by slow moving sodium ions.

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O

On continued addition of sodium hydroxide, first there will be decrease in the conductance due to the removal of fast moving H+ ions. This is followed by slight increase in conductance due to the highly ionised salt formed from weak acid and strong base reaction. This increase is noticed until the weak acid has been completely neutralized. CH3COOH + Na+ + OH- Na+ + CH3COO- + H2O

Subsequent addition of alkali will result in introducing the fast moving OH- ions which results in a sharp increase in conductance.

NaOH Na+ + OH-

The plot of conductance against the volume of alkali added will have three straight lines AB, EF & CD. The point of intersection of AB and EF gives the volume of alkali required for neutralization of strong acid. The point of intersection of lines EF and CD gives the volume of alkali required for neutralisation of weak acid.

Precipitation reactions

Example : Consider the precipitation of silver chloride from potassium chloride (KCl) and silver nitrate (AgNO3)

Principle Electrical conductance depends upon the number and mobility of ions.ProcedureKCl is taken in the conductivity vessel and AgNO3 in the burette. AgNO3 is added gradually and the conductance is noted. The conductance is plotted against the volume of AgNO3 added. The point of intersection of two lines is noted which indicates the complete precipitation of AgCl.

Theory behind the reaction

Volume of AgNO3 added Fig 1.17 Titration of a weak acid vs strong baseConductanceBCDA Increase in conductance due to Ag+NO3-Conductance is constant because Mobility of Cl- and NO3- are same. End point

The Ag+ ions are replaced by the K+ ions. The mobility of K+ ions and ag+ is more or less same. So the conductance will remain more or less constant. After complete precipitation the value increases because of the mobility of K+ ions.

The point of intersection X of these two lines gives the volume of AgNO3 required for precipitation of AgCl.

Replacement titration

ProcedureA strong acid is added to a salt of weak acid. The conductance observed is plotted against the volume of strong acid added.

Theory of Replacement titration

Volume of HCl addedconductanceFig 1.18 Replacement titrationsEnd pointABDC

When a strong acid is added to a salt of weak acid, the weak acid is replaced first.

The chloride ions have higher conducting power than acetate ions. So there will be a slight increase in the conductance upto the end point. But on continued addition of the strong acid, conductance increases more rapidly due to H+ ions

The plot of conductance against the volume of strong acid added will have two straight lines AB and CD. The point of intersection X of these two lines is the end point required for complete replacement of ions.Redox Titration : ( Fe2+ Vs Cr2O72- + H+)

ProcedureIron solution is added to an acidified solution of dichromate. The conductance observed is plotted against the volume of Iron added.

Theory of Redox titration

End pointVolume of Fe2+ addedconductanceFig 1.19 Redox titrationsABCD

With the solution of acidified dichromate, if we add iron solution, due to decrease in H+ ions concentration, the conductance will decrease up to end point.

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

After the equivalent point there will not be any change in the conductance.

The plot of conductance against the volume of iron added will have two straight lines AB and CD. The point of intersection X of these two lines is the end point required for complete oxidation of ions.

1.11 Over voltage

The excess potential that is required in order to operate a cell than the theoretical potential value of the same cell is called overvoltage.

Example:The theoretical value required for Zn2+ to deposit on the cathode as Zn is 0.76 . The experimental value is just the same. So no overvoltage is observed.But in case of evolution of gases say hydrogen or oxygen at anode, the experimental potential required to achieve the cell reaction is much greater than the theoretical potential. This difference in the experimental potential and the theoretical potential is the overvoltage.

ElectrodeHydrogen overvoltage (in volts)

Platinised platinum0.01

Smooth platinum0.26

Nickel0.33

Copper0.54

Mercury1.04

This overvoltage is sometimes referred to as the bubble overvoltage, as it is observed just at the point at which the gas bubbles begin to appear. The value varies largely with nature of the metal.

Hydrogen overvoltageThe hydrogen overvoltage is the difference of potential of the electrode at which hydrogen gas is actually evolved and the potential of the reversible hydrogen electrode. At zero current density and with platinised platinum as electrode, the over voltage of hydrogen is zero.

Determination of hydrogen overvoltageConstructionThe electrolytic solution is taken in the H shaped vessel. There are two electrodes one of which is platinum dipped in electrolyte. The cell is then connected to a reference electrode generally calomel electrode through a salt bridge containing the saturated solution of a salt. The whole apparatus is then connected to a variable resistance D through a galvanometer. The potential of the electrodes is determined by connecting it with a potentiometer.

Fig 1.20

WorkingThe electrolyte used is a dilute Sulphuric acid. The current density is maintained at a desired level. The potential of the cathode at a given current density is measured potentiometrically by combining it with a calomel electrode. The difference between this potential and the theoretical potential gives the hydrogen overvoltage at the particular metal at the given current density.

Factors affecting the phenomenon of over voltage:1.Current density : The over voltage increases with increasing current density (I) according to the following equation = a + b log Iwhere a and b are constants.

The value of b may be represented byb = 2 x 2.303 RT/F

Its value is equal to 0.118 at ordinary temperatures.

2.Hydrogen ion concentration In the absence of strongly absorbed ions the over voltage is independent of hydrogen ion concentration. In the case of strongly acidic or alkaline solutions some deviations occur.3.Temperature Over voltage decreases with increase in temperature.4.Impurities Impurities present in cathode material greatly influence the over voltage.5.Pressure At lower pressure, the over voltage increases rapidly on the cathodes like copper, mercury or nickel.At high pressures, the over voltage is affected very slightly.

Applications of overvoltage1.For solutions of equivalent ionic concentration, the higher the electrode potential involved, the more easily the cations will be discharged.2.Used in the electrolysis of a solution which is molar w.r.t zinc sulphate as well as sulphuric acid. Where electroplating of zinc from acidic solution is achieved due to the overvoltage only.3.It is used in electrolysis of neutral solution of cadmium salts. Since the overvoltage of hydrogen over cadmium is 0.99V more, cadmium deposition could be achieved before the evolution of hydrogen. 4. Metal plating is achieved because of the high value of hydrogen overvoltage.5.In lead accumulators due to over voltage the lead is deposited on the negative plate accumulators while changing the cell.6.In the electrolytic reduction of organic compounds. In the reduction of nitrobenzene, the electrodes with high over voltage i.e. lead are used instead of platinum.7.It is utilized in the industrial production of chlorine and NaOH by the electrolysis of NaCl solution.

1.12 Polarisation

The phenomenon of back EMF brought about by the product of electrolysis is termed as polarization.

Explanation of polarizationIn a voltaic cell, zinc and copper electrodes are connected externally by a wire. Hydrogen ions (carrying a positive charge) will move towards the copper electrode. They give their charges at the electrode and escape through the solution as bubbles. In this process, a film of neutral hydrogen is gradually deposited at the copper electrode. This with time increases in thickness. So it provides a greater resistance to the flow of the current through the cell. At this stage the current is diminished.

As the process proceeds, the film of neutral hydrogen increases in thickness. As hydrogen is positive towards zinc, an electric field from the layer of hydrogen to that of zinc is set up. This is known as back emf. When the field is sufficiently high, it stops further traveling of positively charged hydrogen ions towards the copper electrode, and hence the current stops altogether. This is the stage at which the cell is said to be completely polarized.

Causes of polarization1. The liberated gases offer resistance to the normal flow of current through the cell which requires high voltage to keep the normal flow.2. The product of electrolysis may convert the inert electrode into active electrodes which can exercise a back emf.3. The deposited metals may also form a cell functioning in the opposite direction similar to the gas electrodes.MODERN VIEW1. Polarisation occurs due to the slowness of one or more processes taking place at the electrodes.2. Due to slowness of the diffusion of ions in the solution. Concentration polarization.

Factors which affect electrode polarization :

1.Concentration of electrolyte High concentration of electrolyte increases the polarization effects.2.Nature of ions deposited on electrodes : Extent of polarization increases when there is formation of the adherent and non-porous film on the electrodes.3.Size of the electrode Small surface area of the electrode increases the polarization effects.4.Nature of electrode surface Rough surface of the electrode decreases polarization effects than smoother of surface.For example, platinum black electrode exhibits less polarization effect than the smooth platinum electrode.5.Stirring of the electrolytic solution Polarization effects can be minimized by stirring.6Temperature The rate of diffusion of ions increases with increase in temperature, which reduces the polarization effects.

Effects of electrode polarization1.Polarized galvanic cells develop smaller potential than theoretically predicted.2.In an polarized electrolytic cell, more than theoretically required potential has to be used to maintain a given current strength.

Elimination of polarization 1.Mechanical method By brushing off the hydrogen from time to time from the electrode.By roughening the surface of the electrode, so the bubbles may not stick to the surface.

2.Electro-chemical methodIn these electro-chemical methods, the two solutions are taken such that the hydrogen meets with a second solution, from which ions of the same metal as that of the positive plate or gas is liberated.

3.Chemical methodUsage of depolarizers. The chemicals which are used to eliminate polarization are known as depolarizers.

The depolarizers are usually strong oxidizing agents (like nitric acid or chromic acid). They convert evolved hydrogen into water thus not allowing the hydrogen to get absorbed on the electrode.

1.13 Decomposition potential

The minimum voltage necessary to produce continuous electrolysis of an electrolyte is called the decomposition potential of the electrode.

ExampleWhen two smooth platinum electrodes are placed in dilute Sulphuric acid and a low voltage is applied, practically no current flows through the circuit.

ABEDCApplied Voltage (V)Current(C)Fig 1.21 Decomposition Potential

But as the applied voltage is gradually increased by means of the external battery, to electrolyse the solution, the current increases in the manner as shown in the current density vs potential plot.

It could be seen that an appreciable voltage is required for free flow of electricity through the cell. The emf applied at the point D at which the electrolysis restarts and proceeds continuously, is called the decomposition voltage

Measurement of Decomposition potentialPrincipleThe decomposition potential can be measured by plotting current density against the applied voltage. The point where there is sudden increase in the current gives the measure of decomposition potential.

ConstructionThe apparatus consists of two platinum electrodes A and A dipped in an electrolyte as shown in fig. S is a stirrer for stirring the solution constantly. The two electrodes are connected to a voltmeter V through ammeter C to measure the strength of current. B is the external source of E.M.F and R is the variable resistance. The voltage may be measured by voltmeter V.

First of all a resistance is applied from R and the ammeter and voltmeter readings are noted. Then at another reduced resistance from R, again the values of V and C are taken. The process is continued until the electrolysis in the cell is not visible.A series of current density is taken versus the applied voltage above and below the decomposition potential. i.e., The current (C ) is plotted against the applied voltage (V).

Importance and significance of decomposition potential1.It is of greater importance in the controlled deposition of metals or other electrolytic products from an solution containing different electrolytes. The one having the lowest decomposition potential is electrolysed first.2.The knowledge of decomposition potential are of use in electro refining, electroplating and electrometallurgy of metals.3.Used in separation of copper and zinc electrically.

1.14 Kohlrauschs law

Kohlrauschs law states that at infinite dilution when dissociation is complete and all interionic effects vanish, each ion moves independently of its co-ion and contributes to the total molar conductance of an electrolyte a definite share.

Thus the equivalent conductivity of an electrolyte at infinite dilution (0) is the sum of two parts one due to the cation + and other due to the anion

0 = 0+ + 0In terms of molar conductivities the Kohlrauschs law may be mathematically stated as

0m = 0+ + 0where 0+ = contribution of cation towards the molar conduction at dilution. 0 = contribution of anion towards the molar conduction at dilution.

Applications of Kohlrauschs law

i) Determination of molar conductivity of weak electrolytesThe equivalent conductivity for a weak electrolyte does not become constant even at very high dilutions, so it is not possible to determine the value of 0 for these directly. Kohlrauchs law can be used to determine 0 for weak electrolytes.Example : The equivalent conductance of a weak electrolyte like acetic acid can be determined by determining the equivalent conductances of strong electrolytes like HCl , CH3COONa and NaCl at infinite dilutions as below :

i.e.,

Similarly

The molar conductance at infinite dilution of a sparingly soluble substances like silver chloride, can also be obtained from similar considerations.

ii) To find net Ionic product of water

Water dissociates feebly. Dissociation of water may be written as:H2O H+ + OH-

Its equilibrium is given by the relation

Since the dissociation of water is very low, concentration of undissociated molecules [H2O] may be considered to remain constant . Hence

[H+] [OH] is also a constant.This constant is called Ionic product of water and is represented as Kw . Thus ,

Kw = [H+] [OH]

In pure water the product of the concentration of hydrogen and hydroxide ion is constant. The conductivity of purest distilled water is first determined experimentally. Let it be k.The equivalent conductivity of water is given by

Applying the formula = k V (Volume in cm3 containing 1 eq of electrolyte)where k = conductivity of water = 18 k ( V = 18 in case of water)

The degree of dissociation of water ,

From the dissociation of water H2O H+ + OH [H+] = c = [OH]where c is moles of H2O per litre of water = 1000/18 . Knowing the values of c and , [H+] and [OH] is computed. And from this

Kw = [H+] [OH] = c2 2 is worked out.

iii) Determination of solubility of sparingly soluble salt :

There are number of salts , such as silver chloride, barium sulphate, lead sulphate etc., which are so sparingly soluble in water.

The conductance of the solution containing sparingly soluble salt in water is determined using conductivity meter

The conductance of water used in the preparation of the solution is also determined

Specific conductance cell constant

Let the specific conductance be z Sm-1.

Suppose the solubility of silver chloride = x mole /m3Then concentration of AgCl in aquous solution = x mol/m3

Molar conductance is given by

= At infinite solution, from Kohlrausch law,

or

Solubility of AgCl in water is =

iv) Degree of Dissociation of Weak electrolyte : The degree of dissociation of weak electrolyte such as NH4OH , acetic acid can be determined by measuring the equivalent conductivity , of the solution of the electrolyte at any given dilution. Since 0 is the equivalent conductance when electrolyte is fully dissociated , the ratio /0 would give the degree of dissociation of the weak electrolyte /0 = 0 could be obtained from Kohlrauschs law.

Similarly, if 0 and the degree of dissociation are known, can be calculated.

v) Limiting (or absolute )mobilities of ions . The mobility of an ion u is its speed under potential gradient of 1 volt/cm .At infinite dilution,

Example 1 : The molar conductance of sodium acetate, hydrochloric acid and sodium chloride at infinite dilution are 91104, 426.16104 and 126.45 104 S m2 mol1 respectively at 250C Calculate the molar conductance at infinite dilution for acetic acid.Solution :

= 91104 S m2 mol1

= 426.16104 S m2 mol1

= 126.45 104 S m2 mol1

(or)

= (91 + 426.16 126.45 ) 104= 390.71 104 S m2 mol1

Example 2: The molar conductance of a 0.01 M solution of acetic acid was found to be 16.30 104 S m2 mol1 at 250C. The molar ionic conductances of hydrogen and acetate ions at infinite dilution are 349.8104 and 40.9104 S m2 mol1 respectively, at the same temperature. What percentage of acetic acid is dissociated at this concentrationSolution :

= ( 40.9 +349.8) 104 = 390.7 104 S m2 mol1

0.01 M acetic acid is 4.172 percent dissociated.

Example 3 : The conductivity of a saturated solution of barium sulphate at infinite dilution is 3.06 10-6 ohm-1 cm-1 The equivalent conductance of BaSO4 at infinite dilute is 143 ohm-1 cm2 eq-1. What is the solubility of BaSO4 at 250C in gm equivalent per litre. ( AU Jan 2005)Solution :Conductivity of saturated solution of Barium sulphateK = 3.06 106 ohm1 cm1Equivalent conductivity of BaSO4 =

.

Concentration of BaSO4 = C =

= = 2.14105 gm/equ/litwe know equivalent weight of BaSO4 = 116.5Solubility of BaSO4 = C equivalent weight of BaSO4 = 2.14105 116.5 = 2.49 103 gm lit1

Example 4 : A conductivity cell , when filled with an aqueous solution of 0.02 M KCl at 250C, had a resistance of 250 ohm. Its resistance, when filled with 6105 M NH4OH solution was 105 ohm . The specific conductance of 0.02 M KCl was 0.277 S m-1 . The molar conductances at infinite dilution of NH4+ and OH ions are 73.4104 and 198104 S m2 mol-1, respectively. Calculate the degree of dissociation of 6105 M NH4OH solution.Solution :Since specific conductance = k = cell constant/R , henceCell constant = kR = 0.277 Sm1 ( 250 ) = 69.2 m1For NH4OH solution,c = 6 105 M = 6 105 mol dm3 = 6 102 mol m3

= 115 104 S m2 mol1

According to Kohlrauschs law,0m = 0+ + 0 = (73.4 +198) 104 S m2 mol1 = 271.4 104 S m2 mol1

Example 5: At 250C the specific conductance of a saturated solution of AgCl is 2.68 104 S m-1 and that of water with which the solution was made is 0.86 104 Sm-1. If molar conductances at infinite dilution of AgNO3 , HNO3 and HCl are respectively, 133 104, 421 104 and 426 104 S m2 mol1, calculate the solubility of AgCl in grams per dm3 in water at the given temperature.Solution : ksolution = kAgCl + kwater kAgCl = ksolution kwater = (2.68 0.86) 104 S m1 = 1.82 104 S m1Since AgCl is formed according to the reaction

AgNO3 + HCl AgCl + HNO3

Hence using kohlrauschs law,

= (133.0 + 426.0 421.0 ) 104 S m2 mol1 = 138.0 104 S m2 mol1

and for the saturated solution of the salt,

= 1.32 105 S dm3Mm(AgCl) = 143.5 g mol1Solubility of AgCl = (1.32 105 mol dm3)(143.5g mol1) = 1.89 103 g dm3

Example 6 : At 250C the specific conductance of a 0.01M aqueous solution of acetic acid is 1.63 102 S m1 and the molar conductance at infinite dilution is 390.7 102 S m2 mol1. Calculate the degree of dissociation and hence the dissociation constant of the acid.Solution : k = 1.63 102 S m1, c = 0.01 mol dm1 = 0.01 103 mol m3

The degree of dissociation is given by

The dissociation constant of the acid =

Review Questions1)i) What is single electrode potential ii) Derive Nernst equation . Explain the various terms involved ( Anna University May 2003 )2)How is emf of electrochemical cell determined by potentiometry . (Anna Univ Campus Nov2002 )3)Define the term single electrode potential .Derive the Nernst equation and give its applications . (Anna Univ. Model Q.P)4)How is calomel electrode constructed ? . Discuss how this electrode may be used for the determination of pH of a solution ( Anna University Model Q.P )

5) What is emf series . Explain its significance .6)What are galvanic cells and concentration cell .7)The emf of the cell Zn/ZnCl2 // AgCl /Ag is 1.005 V at 250C . Calculate the heat of reaction H at the given temperature.

8)Explain different kinds of single electrodes with examples .9)What are reversible and irreversible calls . Give examples 10)What is electrochemical series , what are its applications .

11)Define e.m.f , Derive an expression for emf of a cell

12)Explain sacrificial anode cell

13)Determine the emf or Zn-Fe cell at 2980K . The oxidation potentials of zinc and iron are +0.76V and +0.44V respectively .14)Determine the potential developed when a zinc rod is kept immersed in zn(NO3)2 solution of strength 0.5M at 250C E0Zn2+/Zn = -0.76 V

15) Determine the emf of the cell Zn/Zn2+// Ag+/Ag , when Zn2+ = 0.1 M and Ag+ = 10M , Ecell at 250C = 1.56 V.

16. What do you understand by a half cell ? What are the half cells present in a Daniel cell ?

17. What for you understand by a standard hydrogen electrode and what its significance ?

18.What is a salt bridge and what is its role in a Galvanic cell ?19. What is electrochemical series and what are its important features ? 20. Write short notes on the Following :i) Potential decompositionii) Over voltage

Example Problems

Example 1: ( Anna SE , Jan2001)Determine the emf of a Daniel cell at 250C , when the concentration of ZnSO4 and CuSO4 are 0.001M and 0.1 M respectively . The standard emf of the cell is 1.1 volt Solution:Given : E0 = 1.1 Volts C2 = 0.1 , C1 = 0.001The emf of the cell is given by

= 1.1591 VoltsExample 2Determine the emf of a concentration cell at 250 C consists of two Zn electrodes immersed in a solution of zinc ions of 0.1M and 0.01M concentration.

Solution :Given : n = 2 C1 = 0.01 C2 = 0.1

The reaction isZn2+ + 2e- ZnThe emf of the cell is given by

= 0.0296 Volts

Example 3 ( Anna University Dec2001 )Determine the reduction potential of Cu/Cu2+ is 0.5M at 250 C E0Cu/Cu2+ = 0.337 V .Solution :E0 = 0.337 Volts , [Cu2+] = 0.5The reduction potential is given by

log [Cu2+] E = 0.0591 n

log 0.5 = 0.0591 2 = 0.328 Volts

Example 4Determine the emf of the concentration cell given below: Cu Cu2+ Cu2+Cu (0.2M) (2M)

Solution :Given : n = 2 C1 = 0.2 M C2 = 2 M

The reaction isCu2+ + 2e- CuThe emf of the cell is given by

= 0.029 Volts

Example 5Determine the emf of the following cellZn Zn2+ Ag+ Ag , [Zn2+ ] = 0.1 M , [Ag+] = 0.1M , E0Ag = +0.8Volts , E0Zn = 0.76 Volts

Solution :Given :[Zn2+ ] = 0.1 M , [Ag+] = 0.1M , E0Ag = +0.8Volts , E0Zn = 0.76 Volts

The emf of the cell is given by

Ecell = Ecathode Eanode ( EAg+/Ag) (EZn2+/Zn )

= (0.8 (0.76) ) 0.02955 = 1.53045 VoltsExample 6 Determine the emf of the cell at 250C concentration of ZnSO4 and CuSO4 are 0.01 M and 0.1 M respectively . The standard e.m.f of the cell is 1.1 volts .Solution:

Given :[ZnSO4] = 0.01 M , [CuSO4] = 0.1 M , Ecell = Ecathode Eanode = ECu2+/Cu EZn2+/Zn

( We have E0Cu2+/Cu - E0Zn2+/Zn = 1.1 volts )

= 1.129 Volts

Example 7:Determine the emf of a concentration cell consisting of silver electrodes immersed in 0.01M and 0.1 M solutions of its ions at 250 C .Solution : For the concentration cell with silver electrode ,The electrical reaction involved is Ag Ag+ + en = 1 , C2 = 0.1 M and C1 = 0.01MThe emf of the cell is given by

= 0.0591 Volts

Example 8 :The standard electrode potentials of lead and silver are 0.18V and + 0.8V respectively . Determine the emf .Solution : E0Pb2+/pb = 0.18 V E0Ag+/Ag = + 0.8 V Ecell = E0cathode E0anode = +0.8 ( 0.18 ) = 0.98 Volts

Example 9 ( Anna University Nov2001)Determine the concentration of H+ in the following cell . Pt, H2 ( P= 1 atm) H+ (C = 10-6 M)