UNIT – III

ELECTROCHEMISTRY

INTRODUCTION

Electrochemistry is a branch of chemistry that analyse the phenomena

resulting from combined chemical and electrical effects. It covers two

processes. The electrolytic processes where chemical changes occur

on the passage of electric current and galvanic or voltaic processes

where production of electrical energy results due to chemical

reactions. Energy storage as a natural process is as old as the universe

itself. The energy stored in stars such as sun is used directly as solar

energy. To balance the supply and demand of energy, man started

storing energy. The introduction of electricity and chemical fuels like

gasoline and LPG makes the energy storage an important factor in the

economic development of the nation.

Electrolytic and metallic conduction (Distinction)

S. No METALLIC CONDUCTION ELECTROLYTIC

CONDUCTION

1. It involves the flow of electrons

in a conductor.

It involves the movement of ions in

a solution.

2. It does not involve any transfer

of matter.

It involves transfer of electrolyte in

the form of ions.

3. Conduction decreases with the

increase in temperature.

Conduction increases with the

increase in temperature.

4. No change in chemical

properties of the conductor.

Chemical reactions occur at the two

electrodes.

Specific conductance:

Specific conductance can be defined as the conductance of a material of 1

cm in length and 1cm2 in area of cross section, otherwise, generally it is the conductance of 1cm3 of any material. Unit of specific conductance = ohm-1 cm-1 or mho cm-1.

Equivalent conductance:

Equivalent conductance is defined as the conducting power of all the ions

produced by dissolving 1 gm equivalent of an electrolyte in a given volume of

solution. The equivalent conductance is calculated from the relation,

𝜆 =1000 𝜅

𝐶

λ = equivalent conductance at a given concentration

k = specific conductance at that concentration

C = concentration of the solution in Normality.

Unit of equivalent conductance is mho cm2eq-1. Molar conductance:

Molar conductance is the conducting power of all the ions produced by one gram mole of an electrolyte in a given volume of the solution.

𝜇 = 1000𝜅

𝐶

Where k is the specific conductance and C is the concentration of the solution in molarity. Measurement of conductance (Wheatstone’s Bridge method)

Wheatstone Bridge is as shown in figure. The assembly consists of two resistance arms. In the first arm a resistance box is included where a standard resistance R can be introduced. The solution whose resistance and in turn the conductance is to be measured is taken in a conductivity cell and introduced in the second arm. The Wheatstone Bridge is fed with an alternating source of

current through a key. If a direct current is used, electrolysis will occur at the conductivity cell and this will produce a back emf due to the accumulation of products at the electrode. This also changes the concentration of the solution near the electrodes. The middle of the arm is connected to a head phone which in turn is connected to a sliding contact “J” which can move along wire AB which is a uniform wire of high resistance. The length of the wire can be read out from the scale fixed below it. When the connections are made, a sound in the head phone is heard. On moving the sliding contact along the wire AB at a particular point a minimum sound is heard which indicates the null point. At this point,

𝐑 𝛂 𝐁𝐉

𝐑𝐱 𝛂 𝐀𝐉 𝐑

𝐑𝐱=

𝐁𝐉

𝐀𝐉𝐑𝐱 = 𝐑 ×

𝐀𝐉

𝐁𝐉

The conductance of the electrolyte solution, 𝐶 =𝟏

𝐑𝐱

GALVANIC CELL

Galvanic cell or voltaic cell is called electrochemical cell. It is a

device which produces electrical energy from a redox chemical reaction. The

decrease in the potential energy of the chemical reaction appears in the form of

electrical energy. Redox reaction is a combination of both oxidation and

reduction half reactions.Example: Daniel cell. Daniel cell consists of two chambers. In the first compartment a zinc plate

is immersed in Zinc sulphate solution (zinc electrode) and copper plate is

immersed in copper sulphate solution(copper electrode) in the second

compartment. The two compartments are connected to each other by means of a

salt bridge and externally by connecting wires. Salt bridge is a bent glass tube

containing a gel of K2SO4 and both the ends are plugged with glass wool. This will

provide the electrical contact between the electrolytes. The zinc rod and copper

rod are connected to an ammeter using connecting wire to check the production

of electrical current. When both electrodes are connected a deflection is noted in

the ammeter which shows the production of the electric current in the circuit.

Mechanism: Reactions at anode:Oxidation half reactionZn(𝑠) → Zn2+ + 2e−

Reaction at cathode:Reduction half reactionCu2+ + 2e− → Cu(𝑠) The overall cell reaction is

Zn(𝑠) + Cu2+(𝑎𝑞) ⇌ Zn2+

(𝑎𝑞) + Cu(𝑠)

Redox reaction

The electrode at which oxidation occurs is called anode while electrode at

which reduction occurs is called cathode. In Daniel Cell zinc electrode is the negative terminal as it gives out electrons and copper electrode is positive terminal as it accepts electrons. Cell representation or cell notation: By convention we always represent anode at the left side and cathode at the right side. The above galvanic cell represented by

Zn/Zn2+(M)//Cu2+(𝑀)/Cu Reversible and irreversible cells A cell which follows the following conditions is a reversible cell. (i) If an external emf exactly equal to the emf of the cell is applied, the cell

reaction is stopped. (ii) If an external emf slightly greater than the emf of the cell is applied, the cell

reaction gets reversed. (iii) If an external emf slightly lesser than the emf of the cell is applied, the

current flows from the cell. If any cell does not follow these conditions then it is called as an irreversible cell.

Example: Reversible cell: Daniel cell Zn/Zn2+//Cu2+/Cu. Irreversible cell: Zn / H2SO4 / Ag

Zn + H2SO4 → ZnSO4 + H2 ↑

In this case, Zinc dissolves with the liberation of H2 gas.Hydrogen has already escaped; the cell reaction cannot be reversed Emf

The difference in potential which causes the flow of current from an electrode of higher potential to another electrode of lower potential in a galvanic cell is called the electromotive force (emf) of the cell. Unit of emf is volt.

By convention we compare the reduction potential of all electrodes with

respect to each other and the reduction potential of the right hand electrode should be always higher than the reduction potential of the left hand electrode The emf of the cell

𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑅 − 𝐸𝐿 Standard electrode potential

The tendency of the electrode to lose or gain electrons when it is in contact with 1M concentration of its own salt solution at 250C is called as the standard electrode potential. Redox Potential

The potential difference that arises due to the presence of ions of a substance in two oxidation states is called redox potential. For example, a platinum wire

immersed in a solution containing 𝐹𝑒2+ 𝑎𝑛𝑑 𝐹𝑒3+ions. The reduction equation is,

𝐹𝑒3+ + 𝑒− ⇌ 𝐹𝑒2+ The redox potential of the electrode:

E𝑟𝑒𝑑𝑜𝑥 = E𝑟𝑒𝑑𝑜𝑥0 −

RT

nF𝑙𝑛

𝐹𝑒2+

𝐹𝑒3+

Nernst Equation for Electrode Potential Consider the following redox reaction,

Mn+ + ne− ⇌ M

For such a redox reversible reaction, the free energy change (ΔG)and its equilibrium constant (K) are interrelated as,

∆G = −RT𝑙𝑛𝑘 + RT𝑙𝑛[𝑝𝑟𝑜𝑑𝑢𝑐𝑡]

[𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡]

∆G = ∆G0 + RT𝑙𝑛[𝑝𝑟𝑜𝑑𝑢𝑐𝑡]

[𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡]

(1) Since, ∆G0 = −RT𝑙𝑛𝑘

Where, ΔG0 = standard free energy change.

The above equation (1) is known as Van’t Hoff isotherm. The decrease in free energy (-ΔG) in the above reaction will produce

electrical energy. In the cell, if the reaction involves the transfer of ‘n’ number of

electrons, then ‘n’Faraday of electricity will flow. If E is the emf of the cell, then

the total electrical energy nFE is produced in the cell. Where ‘F’ is the Faraday’s

constant which is equal to 96500 coulombs.

∆G = −nFE (or) −∆G0 = nF𝐸0 (2) Where, E = Electrode potential

E0 = Standard electrode potential

−∆G = decrease in free energy change

−∆G0 = decrease in standard free energy change Comparing equations (1) and (2), It becomes

−𝑛𝐹𝐸 = −𝑛𝐹𝐸0 + 𝑅𝑇𝑙𝑛[𝑀]

[𝑀𝑛+] (3)

Dividing the above equation (3) by - nF,

𝐸 = 𝐸0 − 𝑅𝑇

𝑛𝐹𝑙𝑛

1

[𝑀𝑛+] (4)

Since the activity of solid metal [M] = 1.

In general, 𝐸 = 𝐸0 +𝑅𝑇

𝑛𝐹𝑙𝑛[𝑀𝑛+]

𝐸 = 𝐸0 +2.303𝑅𝑇

𝑛𝐹𝑙𝑜𝑔[𝑀𝑛+]

When R =8.314 J/K/Mol, F = 96500 coulombs, T = 298K (250C), the above

equation becomes,

𝐸 = 𝐸0 +0.0591

𝑛𝑙𝑜𝑔[𝑀𝑛+]

This is known as the Nernst Equation. Calculate the reduction potential ofCu/Cu2+(0.5M)at 250E0 Cu2+/Cu = 0.337V. The reduction potential of copper electrode can be calculated as follows Given data: E0 Cu2+/Cu =0.337 V

Cu2+ + 2e− → Cu

ECu2+

Cu= E0

Cu2+

Cu−

RT

nF𝑙𝑛

1

[Cu2+]

Substitute the values

ECu2+

Cu= 0.337 −

8.314 × 298 × 2.303

2x96500𝑙𝑜𝑔

1

0.5

= 0.337 −0.591

2𝑙𝑜𝑔2

= 0.337 − 0.02955 × 0.3010

𝐄𝐂𝐮𝟐+

𝐂𝐮= 𝟎. 𝟑𝟐𝟖𝟏 𝐕

Electrochemical series

The reduction potential of most of the electrodes are found with respect to the standard hydrogen electrode. A series of various electrodes arranged in the increasing order of the standard reduction potentials is called electrochemical series or emf series. Significance:

(i) Relative Tendency of oxidation or reduction of electrodes can be predicted.

(ii) Predicting the feasibility of an electrochemical cell. Determination of emf (Poggendroff’s compensation principle)

The emf of a cell can be measured using poggendroff’s compensation principle. Here the emf of the cell is just opposed or balanced by an external emf (emf of a standard cell), so that no current flows in the circuit. A potentiometer is used to measure the emf of a cell.

The potentiometer consists of a uniform wire AB. A storage battery is connected

to the ends A and B of the wire through a rheostat (R). The cell of unknown

emf(x) is connected in the circuit by connecting its positive pole to A and the

negative pole is connected to a sliding contact D through a galvanometer G. The

sliding contact is freely moved along the wire ABtill no current flows through the

galvanometer. Then the distance AD is measured. The emf of unknown cell is

directly proportional to the distance AD.

E𝑥 α AD

Then the unknown cell(x) is replaced by a standard cell (c) in the circuit.

The sliding contact is again moved till there is null deflection in the galvanometer. Then the distance AD|is measured. The emf of standard cell Ec is directly proportional to the distance AD|

E𝑥 α AD|

Then the emf of the unknown cell can be calculated from the following equation. 𝑒𝑚𝑓 𝑜𝑓 𝑢𝑛𝑘𝑛𝑜𝑤𝑛 𝑐𝑒𝑙𝑙 𝐸𝑥

𝑒𝑚𝑓 𝑜𝑓 𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 𝑐𝑒𝑙𝑙 𝐸𝑐

=𝑙𝑒𝑛𝑔𝑡ℎ 𝐴𝐷

length AD|

𝐸𝑥

𝐸𝑐

=𝐴𝐷

AD|

Therefore, emf of the unknown cell,

𝐸𝑥 =𝐴𝐷

AD|× 𝐸𝑐

Application of EMF

1. Solubility of sparingly soluble salts:

The solubility of a sparingly soluble salt like AgCl can be determined by measuring the emf of the cell. For this purpose a concentration cell containing two Ag electrodes is constructed.

Ag/AgCl 𝑐2, 0.01N KCl //0.01N 𝑐1 AgNO3/Ag

The cell can be constructed by placing one of the silver electrode in

contact with 0.01 N solution of silver nitrate and the other in contact with

0.01N solution of potassium chloride. The two solutions are connected by a

salt bridge. A drop of silver nitrate solution is added to the KCl solution.

The small amount of AgCl formed is sufficient to give its saturated solution.

The cell so constructed is a concentration cell with respect to silver ions.

The emf of the above cell is given by,

E = E0 +0.0591

𝑛𝑙𝑜𝑔

C1

C2

E = emf of the cell

E0 = emf of the cell

n = 1

C2 = 0.01 𝑁 (𝑎𝑠𝑠𝑢𝑚𝑒)

C1 = 𝑇𝑜 𝑏𝑒 𝑓𝑜𝑢𝑛𝑑 𝑜𝑢𝑡

From E value the value of c1 is calculated.

Ksp, the solubility product of AgCl = [ Ag+] [Cl−] = C1 x 0.01

Solubility of AgCl = [ K𝑠𝑝]1

2⁄ = (C1 x 0.01)1

2⁄

2. Determination of thermodynamic functions

It is used to calculate H and S of a redox reaction taking place in a cell. From the relation

∆𝐺 = −𝑛𝐹𝐸

According to Gibb‟s Helmholtz equation,

∆𝐺 = ∆𝐻 + 𝑇 (𝜕(∆𝐺)

𝜕𝑇)

𝑃

We get,

−𝑛𝐹𝐸 = ∆𝐻 + 𝑇 (𝜕(−𝑛𝐹𝐸)

𝜕𝑇)

𝑃

÷ 𝑡ℎ𝑒 𝑒𝑞𝑢𝑎𝑡𝑖𝑜𝑛 𝑏𝑦 − 𝑛𝐹

𝐸 =∆𝐻

−𝑛𝐹+

𝑇𝜕𝐸

𝜕𝑇

𝐸 −𝑇𝜕𝐸

𝜕𝑇=

∆𝐻

−𝑛𝐹

−𝑛𝐹 (𝐸 −𝑇𝜕𝐸

𝜕𝑇) = ∆𝐻

∆𝐺 = ∆𝐻 − 𝑇∆𝑆

Comparing equation (7) and (3),

−𝑇∆𝑆 = 𝑇 [𝜕(−𝑛𝐹𝐸)

𝜕𝑇]

𝑃

−𝑇∆𝑆 = −𝑛𝐹𝐸 (𝜕𝐸

𝜕𝑇)

𝑃

∆𝑆 = 𝑛𝐹 (𝜕𝐸

𝜕𝑇)

𝑃

Calculate the emf of the following cell at 25℃ Zn/Zn2+, 0.1M // 𝐶𝑢2+0.5M /Cu, given that the standard emf of the cell is 1.1 V.

Cell reaction is 𝑍𝑛 + 𝐶𝑢2+ → 𝐶𝑢 + 𝑍𝑛2+ The Nernst equation for emf of the cell is

𝐸 = 𝐸0 −𝑅𝑇

𝑛𝐹𝑙𝑛

[𝑍𝑛2+]

[𝐶𝑢2+]

Substitute the values

𝐸0 = 1.1𝑉 𝑅 = 8.314

𝑇 = 25 = 273 + 25 = 298 𝐾

𝑛 = 2

𝐹 = 96500 𝑐𝑜𝑢𝑙𝑜𝑚𝑏𝑠

𝐸 = 1.1 −8.314 × 298 × 2.303

2 × 96500𝑙𝑜𝑔

0.1

0.5

𝐸 = 1.1 −0.0591

2 × 96500𝑙𝑜𝑔

0.1

0.5

𝐸 = 1.1207 𝑉

TYPES OF ELECTRODES

i) Metal-Metal ion electrode.

eg: Zn-Zn 2+

Cu-Cu 2+

ii) Metal- Metal sparingly soluble salt

electrode

eg: Calomel electrode.

iii) Gas electrode.

eg: Hydrogen electrode.

iv) Redox electrode

Pt/Fe 2+, Fe 3+

STANDARD HYDROGEN ELECTRODE

Hydrogen electrode consists of a platinum foil that is connected to a platinum

wire and sealed in a glass tube. Hydrogen gas is passed through the side arm of

the glass tube. This electrode, when dipped in a 1N HCl and hydrogen gas at 1

atmospheric pressure is passed forms the standard hydrogen electrode. The

electrode potential of SHE is taken as zero It is represented as,

𝑃𝑡, 𝐻2 (1 𝑎𝑡𝑚) /𝐻 + (1𝑀); 𝐸0 = 0𝑉

In a cell, when this electrode acts as anode, the electrode reaction can be written as

H2(g)→ 2 H+ + 2e-

When this electrode acts as cathode, the electrode reaction can be written as

2 H+ + 2e-→ H2(g)

Limitations (i) It

requires hydrogen gas and is difficult to set up and transport.

(ii) It requires considerable volume of test solution.

(iii) The solution may poison the surface of the platinum electrode.

(iv) The potential of the electrode is altered by changes in barometric pressure.

CALOMEL ELECTRODE

A calomel electrode is commonly used as a secondary reference electrode

for electrode potential measurements. A calomel electrode consists of a glass tube

with side tubes on both sides of it. The tube consists of a layer of pure mercury at

the bottom, over which mercurous chloride is placed. The remaining portion of

the tube is filled with saturated solution of KCl. The bottom of the tube is sealed.

A platinum wire is inserted for electrical contact.

This electrode is represented as

𝐻𝑔, 𝐻𝑔2𝐶𝑙2(𝑠) //𝐾𝐶𝑙 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

𝐻𝑔2𝐶𝑙2 + 2𝑒− → 2𝐻𝑔(𝑙) + 2𝐶𝑙−( 𝑎𝑞) (𝑅𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛)

2𝐶𝑙− (𝑎𝑞) + 2 𝐻𝑔(𝑙) → 𝐻𝑔2𝐶𝑙2(𝑙) + 2𝑒− (𝑂𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛)

The standard electrode potential of saturated calomel electrode is 0.2422 vlots. The electrode is reversible with respect to chloride ions. The potential of the

electrode depends upon the concentration of KCl solution taken. Determination of pH using calomel electrode. The cell of following type is constructed.

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

𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑅 − 𝐸𝐿

= 0.2422 − 𝐸𝐿

𝐸𝐿 = 𝐸𝐿0 − 0.0591𝑙𝑜𝑔 1

[𝐻+]⁄

𝐸𝐿 = 0.0591log[𝐻+]

𝐸𝐿 = −0.0591𝑝𝐻

𝐸𝑐𝑒𝑙𝑙 = 0.2422 + 0.0591 𝑝𝐻

𝑝𝐻 =𝐸𝑐𝑒𝑙𝑙 − 0.2422

0.0591

QUINHYDRONE ELECTRODE

Quinhydrone is a 1:1 molecular complex of quinone (represented by Q) and

hydroquinone (represented by H2Q). The electrode may be represented as Pt/Q, H2Q, H+. It is reversible with respect to H+ ions. The reduction reaction at this electrode may be represented as:

𝑄 + 2𝐻+ + 2𝑒− → 𝐻2𝑄 The emf of the electrode is 0.6994V Determination of pH using quinhydrone electrode The Quinhydrone electrode is coupled with a calomel electrode. The cell may be represented as

𝐻𝑔, 𝐻𝑔2𝐶𝑙2/𝐾𝐶𝑙// 𝐻 + 𝑄, 𝑄𝐻2, 𝑃𝑡 The emf of the cell is given as

𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑅 − 𝐸𝐿 = 𝐸𝑄 − 𝐸𝑐𝑒𝑙𝑙

= 𝐸𝑄 − 0.2422

𝐸𝑄 = 𝐸0𝑄 −

0.0591

2𝑙𝑜𝑔

[𝑄𝐻2]

[𝐻+]2[𝑄]

𝐸𝑄 = 𝐸0𝑄 −

0.0591

2𝑙𝑜𝑔

1

[𝐻+]2

𝐸𝑄 = 𝐸0𝑄 +

0.0591

2𝑙𝑜𝑔[𝐻+]2

𝐸𝑄 = 𝐸0

𝑄 + 0.0591 𝑙𝑜𝑔[𝐻+]

𝐸𝑄 = 𝐸0𝑄 − 0.0591 𝑝𝐻

𝐸𝑄 = 0.6694 − 0.0591 𝑝𝐻

𝐸𝑐𝑒𝑙𝑙 = 0.6694 − 0.0591 𝑝𝐻 − 0.2422

𝑝𝐻 =0.6694 − 0.2422−𝐸𝑐𝑒𝑙𝑙

0.0591

𝑝𝐻 =0.4572−𝐸𝑐𝑒𝑙𝑙

0.0591

ENERGY STORAGE DEVICES

Battery storage technology provides the

most wide spread satisfactory method as storage device in the current

scenario. Electrochemical batteries are of several types. Depending on

the type of battery their usage also varies. There is a growing trend

and need for the rechargeable batteries.

Lead storage battery Lead storage battery is also known as lead – acid battery.

A battery or storage cell is a combination of two or more cells arranged in series

or parallel in which electrical energy is stored as chemical energy. When

required, this chemical energy can be reconverted into electrical energy. Ex:

Lead – storage cell.

Lead storage cell consists of a lead anode and lead dioxide cathode (lead dioxide

is packed on a metal plate) immersed in 20% sulphuric acid solution. Actually a

number of lead and lead oxide plates are arranged alternatively, with insulating

material in between them. The lead storage battery consists of six identical cells

joined together in series.

Discharging

When the storage cell acts as a voltaic cell it is said to be discharging. In this

process sulphuric acid is consumed and water is generated. The following

reaction takes place during discharging.

Anode: Pb(s)+SO4(aq)2-→ PbSO4(s)+2e-

Cathode: 𝑃𝑏𝑂2(𝑠) + 4𝐻(𝑎𝑞)+ + 𝑆𝑂4(𝑎𝑞)

2− + 2𝑒− → 𝑃𝑏𝑆𝑂4(𝑆)+ 2𝐻2𝑂(𝑙)

Cell reaction:

𝑃𝑏(𝑠) + 𝑃𝑏𝑂2(𝑆)+ 4𝐻(𝑎𝑞)

+ + 2𝑆𝑂4(𝑎𝑞)2− → 2𝑃𝑏𝑆𝑂4(𝑆)

+ 2𝐻2𝑂(𝑙) + 𝑒𝑛𝑒𝑟𝑔𝑦

Recharging:

The lead storage battery is rechargeable. This is done by applying a voltage

slightly higher than the voltage of the battery, across the electrodes. In this

process the sulphuric acid consumed during discharging is reformed. The

recharging involves exactly the reverse process of the normal cell reaction. The

recharging reactions are

Cathode: 𝑃𝑏𝑆𝑂4(𝑠) + 2𝑒− → 𝑃𝑏(𝑠) + 𝑆𝑂4(𝑎𝑞)2−

Anode:𝑃𝑏𝑆𝑂4(𝑠) + 2𝐻2𝑂(𝑙) → 𝑃𝑏𝑂2𝑆+ 4𝐻(𝑎𝑞)

+ + 𝑆𝑂4(𝑎𝑞)2− + 2𝑒−

Cell reaction:

2𝑃𝑏𝑆𝑂4(𝑆) + 2𝐻2𝑂(𝑙) + 𝑒𝑛𝑒𝑟𝑔𝑦 → 𝑃𝑏(𝑠) + 𝑃𝑏𝑂2(𝑠) + 4𝐻(𝑎𝑞)

+ + 2𝑆𝑂4(𝑎𝑞)2−

The emf of each cell is 2 volts. In Automobiles six such cells are connected in series to form a battery with an emf of 12 volts. Nickel Cadmium cell or Nicad Battery Like a lead storage cell this is a rechargeable battery. Anode : Cd (metal) Cathode : NiO2 Electrolyte : KOH Nicad battery consists of a cadmium anode and a metal grid containing a paste of NiO2 acting as cathode. KOH solution is the electrolyte. It gives a constant voltage of 1.4 V. The cell can be represented as Cd/Cd(OH)2 // KOH(aq)/NiO2/Ni Working: Discharging (or current production) At Anode : Cadmium is oxidized to Cd2+ and this combines with OH- ions to form Cd(OH)2. In this reaction two electrons are released at the anode. Cd(s) + 2 OH- → Cd(OH)2 + 2e-

At Cathode: NiO2 receives the two electrons from the circuit and undergoes reduction

(Ni4+ → Ni2+) . The Ni2+ ions combines with OH- ions to form Ni(OH)2 NiO2(s) + 2 H2O + 2e- → Ni(OH)2 + 2 OH- Overall cell reaction is Cd(s) + NiO2(s) + 2 H2O ⇄ Cd(OH)2 + Ni(OH)2 As no gaseous products are produced, the cell reaction is completely reversible. Recharging: Like the lead –acid battery, Nicad battery can be recharged by sending the current in the opposite direction. The electrode reaction gets reversed and as a result Cd metal gets deposited on the anode and NiO2 at the cathode. Charging Reaction: Cd(OH)2 + Ni(OH)2 + energy ⇄ Cd(s) + NiO2(s) + 2 H2O Advantages:

1. It is lighter and smaller.

2. It has longer life than lead storage battery.

3. Like a dry cell, it can be sealed inside a container.

4. Gives a constant voltage of 1.4 V Disadvantage:

1. It is more expensive than lead-acid battery.

Uses:

It is used in calculators, Electronic camera flashes, rechargeable flash lights s

and cordless electronic appliances.

Lithium Battery Lithium battery is a rechargeable battery. It is considered to be the cell of the future. It is a solid state battery. Lithium is the anode and TiS2 (Titanium

disulfide) is the cathode. The electrolyte is solid made of polymer. The polymer allows the passage of ions but not the electrons.

Anode Solid Electrolyte Cathode

Li(s) → Li+ + e- TiS2(s) + e- → TiS2

-

Cell Reactions: Anode : Li(s) → Li+ + e- Cathode: TiS2(s) + e- → TiS2

-

Overall reaction: Li(s) + TiS2(s) → Li+ + TiS2

-

The cell is rechargeable and produces a cell voltage of 3V. Advantages:

1. Lithium is a light weight metal (7g). One mole of material is enough to produce one mole of electrons. It is rechargeable.

2. Cell voltage is high (3V). 3. The constituents are solids and there is no risk of leakage. 4. Battery can be made into various shapes and sizes.

Hydrogen – Oxygen fuel cell

Fuel cells are galvanic cells in which chemical energy of fuels is directly converted into electrical energy. It is an energy conversion device or electricity generator. Unlike a storage cell it cannot be reversed. It is similar to an electric generator set which converts chemical energy of fuels into electricity.

Example: Hydrogen – Oxygen fuel cell. In this cell combustion of H2 in O2 takesplace to from water

2𝐻2 + 𝑂2 → 2𝐻2𝑂

Description:

The cell consists of two electrodes made of porous graphite,

impregnated with platinum catalyst. These electrodes are placed in aqueous

KOH or NaOH solution. Oxygen and hydrogen gases are continuously fed in it

at high pressure of 50 atmospheres. The reaction taking place is,

Anode 2𝐻2 + 4𝑂𝐻− → 4𝐻2𝑂 + 4𝑒− Cathode 𝑂2 + 4𝑒− + 2𝐻2𝑂 → 4𝑂𝐻− Cell reaction 2𝐻2 + 𝑂2 → 2𝐻2𝑂

The emf of the cell is found to be 1V . Advantages of fuel cells

1. Very efficient and converts 75% of chemical energy to electrical energy. 2. The cell is compact and easy to maintain. 3. The fuels hydrogen and oxygen are easily available and cheap 4. The product is only H2O vapours and hence does not cause pollution. 5. Because of its light weight it is used in space vehicles.

Disadvantages

1. The catalyst is easily losing their activity. 2. The catalyst is very expensive.