unit 1 electrochemistry

Chapter-1 Engineering chemistry

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Electrochemistry

Electrochemistry is the branch which deals with the relationship between chemical, electrical

properties phenomena and the laws of interaction phenomena.

The laws of electrochemistry form the basis of electrolysis and electro synthesis.

Electrolysis: Electrolysis is the process where electrical energy causes chemical changes.

Example: Electrolysis of water yields H2 and O2

2H2O 2H2+O2

Electrical conduction: The substances can be divided into four different types depending upon the

capability of flow of electrons. They are electrical conductors, insulators, semiconductors and

superconductors.

Electrical conductors: Generally all the good thermal conductors are good electrical conductors except

Mica. They allow electrons to pass through then with minimum resistance and show no preference of

direction of flow.

Example: Metals, Metal sulphides, acids, bases, fused salts.

Insulators: Substances of this category do not allow the electricity to pass through them.

Example: Pure water, Synthetic organic compounds like Benzene, CCl4, ether.

In an insulator, the gap between the valence and conduction bond is consider large making flow

of electrons difficult.

Semiconductors: Elements of 4A group, especially ‘Si’ and ‘Ge’ have properties that are intermediate

between those of metal and non metals and therefore called semiconductors. Their conducted

properties are considerable enhanced by the edition of certain impurities like phosphorus that have

more valance electron than Si generally called donor impurities give n type semi conductors. Similarly

acceptor impurities like Boron (B), Aluminum(Al) that have less number of valency electrons produce p

type semi conductors.

The most valuable property derived out of doping and unidirectional flow of electrons is

replacement of cumbersome value versions with small semi conductor pieces of thickness of pencil

eraser called chips. This led to invention of solid state device like calculators, mini computers etc.

Superconductors: Led was fond to conduct electricity with zero résistance bellow 7.K. The mixture of

Copper, barium and rare earth metal oxides exhibit zero resistance to as high as 130K the property of

certain metals alloys and compound by virtue of which they conduct electricity with zero resistance is

known as super conductivity.


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Because of zero resistance of super conductors attain the capacity to store infinity amount of

current. This allows transfer of power from the place of generation to utility. Super conductivity is

expected to create vey high field strength electromagnets that can be used for many purposes like lifting

super heavy weight.

Electrical resistance and Ohm’s law: The metallic conduction increases with fall in temperature. With

approach of absolute zero resistance will be almost equal to zero and therefore all substances behave

like perfect conductors.

G.S.Ohm in 1827 found that resistance(R) is related to the current passed (I) and potential

differences across the conductor (E).

I=E/R

Ohm’s law may thus be defined “that the current strength flowing through a conductor at

uniform temperature is directly proportional to the potential differences applied across the conductor

and inversely proportional to the résistance of the material”.

When current is measured I amperes and potential differences in volts the resistance offered a

fixed value ‘Ohm’ represented by “Ω”

Ohm=Volt/Ampere

Ohm’s law is strictly obeyed by all metallic conductors.

In addition ohm also found that the resistance of a conductor of uniform cross section is

directly proportional to its length and inversely proportional to the area of cross section.

R α l/a or R=ρ l/a

Where ρ (rho) is a proportionality constant called “specific resistance” or “resistivity”

The units of specific resistance are ρ =Ohm x cm2/cm = ohm cm

Conductance – Electrolyte in solution, specific Conductance

Resistance and specific resistance are commonly used for metallic conductors, where the atoms

are static in case of electrolytic solutions the electricity is virtually conducted by constantly moving ions.

So it was though more meaningful to define another quantity called conductance. It gives the ease with

which electricity flows through an electrolytic conductor. The conductance of electrolytic solutions is

defined as the reciprocal of its resistance.

Conductance (L) = 1/ R Ohm -1 or Siemens(S).


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The reciprocal of specific resistance is defined as the conductivity or electrolytic Conductivity or

specific conductance. It is denoted by the letter ‘K’. In other words electrolytic Conductivity is the

conductivity of a solution taken in between electrodes of unit area of cross Section (a) and separated by

unit distance.

Units: Specific Conductance in CGS system S cm-1

Conductivity K= 1/ ρ

1/ Ohm cm = S cm-1.

In SI system meter being the fundamental unit of distance, the conductivity is expressed as Sm-1.

Equivalent Conductivity

In 1876 F.W.Kohlraush suggested the use of equivalent conductance. It is represented by ‘Λv’ (Greek-

Capital lambda)

It is defined as “the conductance of all ions produced by the dissociation of one gram

equivalent of an electrolyte in a certain volume (V), of the solvent at a constant temperature”.

Mathematically it is expressed as Λv= k × 1000/c

Where c is the concentration of the solution in gram equivalent /liter

Units: In CGS System ohm-1 cm2 equivqlent-1

In SI System ohm-1 Sm2 equivqlent-1

Molar Conductivity:

In recent times it has become a practice to express the concentration in moles per liter rather

the Equivalent per liter. Keeping in the pace with this new practice the term molar conductivity/molar

conductance was introduced.

Molar conductance is defined as “the conductance of all ions produced by the dissociation of

one gram molecular weight of an electrolyte dissolved in a certain volume (V) of the solvent at

constant temperature”.


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Molar conductivity represented by ‘Λ m’

Mathematically it is expressed as Λm= k × 1000/c

Units: In CGS System ohm-1 cm2 mol-1

In SI System ohm-1 Sm2 mol-1

Experimental determination of specific conductance:

Specific conductance is determined experimentally by using the principle of Wheatstone bridge

the following diagram will illustrate the method of determination of specific conductance

R is the resistance box,

AB is a uniform wire with a sliding contact,

X. Conductivity cell is represented by C. On the sliding wire,

T is the telephone head to detect flow of current.

From the above diagram it is clear that a resistant box with standard resistances is connected to meter

bridge.

An electrolytic solution of known concentration is taken in a conductive cell which contains

platinum electrodes having an area of cross section 1 cm2 and the distance between two electrodes in 1

cm.The terminals are connected to meter bridge then alternate current is passed from induction coil

and the detector is moved across the scratched wire and the point ‘x’ is noted where minimum sound

was heard then it follows that

Form this equation we can calculate resistance of the standard solution and 1/ R=conductance of the

solution Resistance of the solution

Resistance R

Length of XB

Length of AX=


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Specific conduction=observed conduction x l/a

Due to the repeated usage of conductivity cell for many numbers of times the specification of the

electrode may change. Then value of l and a may also change but l/a is constant value which is called cell

constant.

Kohlraush as determined the specific conductance value of KCl solutions at different concentrations and

different temperatures.

Concentration Specific conductance

00C 180C 250C

1.0 N (Normal) 6.543 9.820 11.1732

0.1N (decinormal ) 0.7154 1.1192 1.2886

0.01N(Centinormal) 0.07751 0.12227 0.14114

KOHLRAUSCH’S LAW:

On the basis of specific conductive value (k) electrolytes are classified into strong electrolytes,

and weak electrolytes. Those with high k values are considered as strong electrolytes and those with low

k values are considered as “weak electrolytes”. Strong electrolytes are generally completely ionized in

aqueous solutions. Ex: [Salt acids, alkalies] while weak electrolytes are partially ionized in aqueous

solutions [Ex: Carboxylic acids].

The linear variation of equivalent conductivity with √c of difference strong electrolytes, led to

the formulation of an empirical relation called KOHLRAUSCH’S equation.

Λ v= Λ 0 - b√c

Where’ b ‘is constant

KOHLRAUSCH’S equation suggests that for strong electrolytes the intercept of Λ v vs. √c plots is equal

to Λ0 weak electrolytes do not obey kohlraush relation and therefore this method of calculation of Λ0 is

not suitable for weak electrolytes this is achieved through direct calculation based on kohlrausch’slaw

of independent migration of ions.

Cell Constant =Specific conductance

Observed conductance


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KOHLRAUSCH’S LAW OF INDEPENDENT MIGRATION OF IONS:

DEFINITION:

At infinite dilution, each ion of an electrolyte makes a definite contribution to the total

equivalent conductance of the electrolyte and this contribution is independent of the presence of other

ionic species.

Thus, the equivalent conductance of the electrolyte at infinite dilution is the sum of the

equivalent conductance of the constituent ions at infinite dilution i.e.

λ∞ = λ0 (+) V+ + λ0 (−) V−

Where V+ and V− are the number of cations and anions per formula of the electrolyte.

λ0(+) and λ0(−) are the ionic conductance of cation and anion respectively.

Ionic Conductance and Ionic Mobility: As explained earlier ionic conductance of an ion (cation or anion)

is defined as the contribution made by it towards the equivalent conductance of the electrolyte at

infinite dilution. These are represented by λc and λa.

Since λ∞ = λc + λa

λ c = λ∞ - λa and λa = λ∞ - λc .

Ionic Mobility: Ionic mobility of an ion is defined as its absolute velocity (i.e. the Distance travel in

cm/sec) under a potential gradient of one volt per cm. Thus it must be noted that ionic conductance

ionic mobility are not the same thing. However these are found to be directly proportional to each

other.

λ c α μc and λ a α μa

λ c = k μc and λ a =k μa

μc and μa are the ionic mobilites of the cation and anion respectively. And K is

proportional constant the value of K is always found to be 96500.

Ionic Mobility = Ionic Conductance / 96500


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Applications of Kohlrausch’s law:

(1) Determination of molar or equivalent conductance at infinite dilution (Λ 0) of a weak

electrolyte:

Weak electrolytes do not ionize completely even at very great dilution. Hence the value of λ∞ in such

cases can be calculated with the help of Kohlrausch’s law in this following way.

Consider the example acetic acid

Λ∞ CH3

CooNa+ Λ∞ Hcl- Λ∞ Nacl

λ∞CH

3Coo-+ λ∞

Na++ λ∞ H+ + λ∞

cl-- λ∞Na λ- λ∞

cl

λ∞CH

3Coo-+ λ∞

H+

λ∞ CH3

COOH

(2)Calculation of Degree of dissociation and dissociation constant of weak electrolyte:

The ratio of equivalent conductivity at any dilution Λ v to that at infinite dilution Λ∞ is called

conductance ratio

α = Λ v/ Λ ∞

For weak electrolytes conductance ratio is quite small and is a characteristic quantity called degree of

dissociation. This in fact represents the fraction of molecules dissociated out of unit concentration of the

electrolyte.

Initial concentration 1 0 0

Equilibrium concentration 1- α α α

When c moles are taken initially (1- α) c αc αc

Dissociation constant

KD= [CH3Coo-][H+] [αc][αc]

------------------- = ----------------

[CH3COOH] (1- α) c

KD Of univalent electrolyte = α2c/1- α

Where α = Λ v/ Λ ∞Where c is the concentration of the weak

electrolyte and Λ v is the corresponding equivalent conductivity.

CH3COOH CH

3COO H

+-+


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3).Calculation of ionic product of water:

Water is weak electrolyte, which dissociates partially to give H+ and OH-

Equilibrium concentration 1- α α α

When c moles are taken initially (1- α) c αc αc

Ionic product of water Kw= [H+][OH-]

(αc)(αc)=C2 α2

` Kw=C2 α2

Determination of solubility of products: Salts like Agcl, Pbso4, Agi, Baso4, and Pbi2 are soluble to a

negligible extent and whatever that gets solution dissociates completely.

If S is the solubility, in gram equivalent per liter of a given salt and K is specific conductance of the

saturated solution then Λ v of the solution given by

These Solutions are usually so dilute that their equivalent conductivity Λ v is almost equal to Λ∞

Since the solutions are extremely dilute, the conductance contribution due to pure water is

appreciable part of total conductance and needs to be subtracted from the total value to get the

contribution of dissolved salt.

Where Kw is the specific conductance of pure water. In case of univalent electrolyte like Agcl

The solubility product Ksp= (αAg+) (αCl-)

The solution is dilute enough to regard the activity coefficient are unity

αAg=αCl=S

Therefore Ksp=S2

H+H20 OH

α

v

=1000κ

S

α

v=1000κ

S

α

v

=1000κ

S

α

v=1000 κ

Sκ w-( )


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Conductometric titrations:

This is another application which exploits the variation of the conductance of a solution during a

titration. This technique can be used advantageously with

Colored solutions where selection of proper indicator is difficult.

Acids and bases and their mixture

In and mixtures are readily precipitated by the addition of titrant.

A grave disadvantage of conductivity based titrations is their non-applicability with solutions

contaminated with high concentrations of electrolytic impurities. Presence of excess of electrolytic

(ions) and hence high electrolytic conductivity masks the minor changes in conductivity, taking place as a

result of addition of titration.

Example: Acid-base and precipitation titrations

Titration on of strong acid Hcl with strong base NaOH

Strong acids dissociate completely to give hydrogen’s (H+) ions. The ionic conductance data

suggest that protons (H+) are almost two times more conducting than hydroxyl (OH-) ,which themselves

are more conducting than all other ions. When NaOH is added to a solution of Hcl .The reaction is

(H+ + cl-) + (Na++ OH-) (Na+ + Cl-) H2O

During the course of neutralization initially high conducting H+ ions present in the acid are

replaced by sodium ions having a much lower conductance. Therefore addition of alkali is accompanied

by a fall in the conductance till all the H+ ions are totally neutralized. Further addition of alkali may be

interpreted as only an addition

of OH- ions which are no longer used up in the acid-base reaction.

Hence after neutralization conductance of the reaction solution

increase linearly (ON) with volume of alkali added. Interpolation of

both the linear segments MO and ON to meet at O gives the end point.

Conducta

nce

Volume of the alkali

Neutralization point

M

N

O


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Titration of weak acid (CH3COOH) by strong base (NaOH):

A weak acid, like is characterized through its insignificant dissociation.

Therefore at the beginning of titration the acid solution has a low conductance.

Addition of alkali to this moderately weak acid solution, initially

result in a small fall of conductance owing to the replacement of free

H+ ions by acetate ions. If acid is very weak acid,

Example: Boric acid and phenol, the initial decrease

Of the conductance upon neutralization is not observed due to

Insignificant dissociation. After the addition of the few drops

of the alkali, the conductance increase in regular fashions.

This is due to suppression of ionization equilibrium of acetic acid by the acetate ions formed and

increase in ionic population in the solution.

(CH3COOH) + (Na++OH-) (CH3COO-+Na +) +H2O

Beyond neutralization point, additions of alkali result in a steep raise in conductance, the strong

base being a better conductor than the salt.

Conductometric titration of a mixture of strong acid (HCl) and weak acid (CH3COOH) with strong base

(NaOH):

One of the valuable features of conductometric is the

titration of a mixture strong acid and a weak acid without

employing any indicator. In presence of a strong acid,

the dissociation of weak acid is completely suppressed

due to’ common ion effect’. Therefore addition of a base

to such a mixture will first results in the neutralization of

the strong acid .The weak acid starts reacting only after neutralization of strong acid. In titration graph

the line segment MX represents the neutralization of strong acid and the segment XY represents the

neutralization of the weak acid.

Conducta

nce



M

N

O

Conducta

nce



M

N

X


Y


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Electro motive force: The difference of potential, which causes the current flow from at higher potential

to one of lower potential is called Electro motive force (emf) of the cell.

Cell: a cell may be defined as a single arrangement consisting of two electrodes and capable to

producing electricity due of chemical reaction.

There are two types of cells

Galvanic (or) voltaic cells: It is uses the energy released from a spontaneous chemical reaction to

generate electricity.

An electrolytic cell: It is uses the electrical energy to causes a non-spontaneous chemical reaction occurs

in spontaneous manner and the process is commonly called electrolysis.

Modern society depends on the galvanic cells for running cars, radios, calculators, medical

instruments, computers, emergency flash lights and so on.

Galvanic cells: In order to tap the energy of the reaction in the form of electrical energy, the reactants

must be physically separated so as to prevent direct transfer of electrons. Electrons are then allowed to

go from Zn to Cl2 through an external circuit, made out of a wire or metallic conductor. In the current

example both Zn rod and Cl2 taken in two separate compartments which are separated either through a

porous partition or salt bridge. The porous partition or salt bridge prevents direct mixing of reactants

and at the same time are permeable to ions.

Since the Cl2 itself is a non-conductor unlike Zn rod, to conduct electricity in the Cl2

compartment an inert metal wire like ‘Pt ‘is placed. The Zn rod and Pt wire are called electrodes. By the

definition of electrode at which oxidation occurs is called anode Zn-electrode, the electrode at which

reduction occurs is known as cathode Cl2-gas electrode.

The individual oxidation and reduction reactions at the electrodes are called half cell reaction.

Zn(s) Zn2++2e- (half cell reaction at anode)

Cl2+2e_ 2Cl- (half cell reaction at cathode)

-----------------------------------------

Zn+Cl2 ……………Zn2++2Cl- (net reaction)

----------------------------------------


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In order to maintain electrical neutrality of both compartments the Cl and Zn2+ moves across the

porous partions.For the movements of each Zn2+ ion anodic compartment to cathodic compartment

there is a current migration of two Cl- ions from cathode compartment to anodic compartment.

Representation of a cell: A symbolic representation is used to describe a cell. This representation is

called cell diagram. Following rules are adopted in these representations.

Example: Daniel cell is represented as

The anode, where oxidation occurs is represented at the left and at cathode, where reduction

takes place, is represented on the right of the cell diagram.

A single vertical line represents the boundary between different phases that constitute an

electrode.

A double vertical line indicates that the liquid junction potential has been to minimum by

suitable means like using a salt bridge.

The concentrations of the dissolved substances are indicated in parenthesis.

In case of gases at equilibrium it is necessary to represent the pressure of the gas

Pt/H2 (1 atm)/H+ (0.1 M)//Cu2+ (0.1M)/Cu(s)

When two oxidation states of the same species are present in a homogenous Medium, they are

usually separated by a comma.

Pt/H2 (1 atm)/H+ (0.1 M)//Fe3+ (0.2 M), Fe2+ (0.1 M)/Pt

Single electrode potential: Each of these electrodes exhibits its characteristic electrode potential. This

potential; is called Single electrode potential.

Ecell=Ecathode-Eanode

Standard electrode potential: The potential exhibited by a single electrode at unit concentration of the

concerned metal ion or non-metal ion solution at 25o C is called standard electrode potential.

This is expressed by ‘E0’

Standard electrode potential ‘E0’value of a simple electrode is determined experimentally by combining

the standard electrode with standard hydrogen electrode (S.H.E) .The EMF of the resulting.

Standard hydrogen electrode: To determine the potential difference of a single electrode

experimentally, it is combined with a standard hydrogen electrode and the EMF of cell so constructed is


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measured with a potentiometer. Standard hydrogen electrode is constructed and used as standard

electrode or reference electrode.

Pure hydrogen gas is bubbled into a solution of 1 M Hcl along ‘Pt’ electrode coated with

platinum black. A platinum black placed in the solution at atmospheric pressure as shown in the

diagram.

Generally the electrode is fitted into a tube. The tube will have two circular small holes. This

tube is immersed in the acid solution such that one half of the circular hole is exposed to air and other

half is in the solution.

The following equilibrium exists at the electrode

½ H2 (1 atm) ………………………H+ (aq) (1 M) +e-

Types of electrodes:

Standard calomel electrode [S.C.E]: Calomel electrode is particularly very simple to construct, free from

surface sensitivity and accurate to use even in a very normal laboratory.

A calomel electrode consists of inner glass tube and outer jacket. In the inner glass tube a

platinum wire is dipped into mercury which rests on a paste is in contact with KCl present in the outer

jacket, through the glass frit plug fixed at the bottom inner glass tube. The calomel electrode potential

depends on the concentration of KCl taken in the outer jacket.

Electrode reaction is

½ Hg2Cl2+2e-………………………2Hg+2Cl-


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And the corresponding reaction is

E Hg2Cl2/Cl- = E 0Hg2Cl2/Cl

- -2.303/F log a2cl-

Quinhydrone electrode: Quinhydrone is a 1:1 molar mixture of quinine and hydroquinone. The

electrode consists of a shiny platinum (Pt) electrode dipped in an acid/base test solution which is

saturated with quinhydrone.

Electrode reaction is

The electrode potential at 250

E Pt/Q,H+,QH2= E0 Pt/Q,H+,QH2 -0.0592/2 log aQh2/aQa2

H+

Since quinone and hydroquinone are taken in equimolar amounts

E=E0-0.0592/2 log 1/ a2H

+

E=E0+2 x 0.0592/2 log aH+

E=E0+ 0.0592 log aH+

EQ, QH2=E0Q, QH2-0.0592 PH

Quinhydrone electrode can thus be used to measure PH of a solution. Due to instability of quinone in

strong alkaline medium; this hydrogen ion indicator electrode is suitable only up to a PH of 8.

Advantages:

PH values of solutions containing reducible substances like Cu2+,Cd 2+,unsaturated acids ,NO3- etc and

catalytic poisons can be measured using Quinehydrone electrode.

Limitations:

1) The electrode cannot be used at PH values greater than 8.

2) Even this electrode fails in presence of strong oxidizing and reducing agents.

Ion Selective Electrodes:

Ion Selective electrode is the one which selectively responds to a specific ion in a mixture and

the potential developed at electrode is a function of the concentration of that ion in the solution. The

electrode generally consists of a Membrane which is capable of exchanging the specific ions with the

solution with which it is in contact. Therefore these electrodes are also referred to as Membrane

electrodes.

The Ion Selective electrode generally consists of ion selective membrane in contact with an analyte

solution on one side and an internal reference solution on the other side. An internal reference

electrode constituted in contact with the reference solution.

The electrode can be represented as follows;


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Solution to be // Internal Internal

Analysed // standard solution reference electrode

[Mn+]=C1 // [Mn+]=C2

The mechanism of potential development in ion selective electrodes is different from that in

other types of electrodes. The observed potential of an ion selective electrode is a kind of junction

potential that develops across a membrane that separates the analyte solution from reference solution.

The potential developed across the membrane is a function of concentration of analyte and reference

solution.

Ej=2.303RT/ nF log C1/C2

Where C1 and C2 are the concentration of the analyte and reference solution respectively. If

the concentration of the inner standard solution is kept constant, then

Ej=k+ (2.303RT/ nF) log C1, where k is a constant k=-2.303RT/ nF log C2

Em=Ej+Eref

=k+ (2.303RT/ nF) log C1+ Eref

=E0M+ (2.303RT/ nF) log C1, Where E0

M =k+ Eref

At 298 K EM= E0M+ (0.0591/ n) log C1

The electrode is combined with an external reference electrode, and emf of the so formed cell is

determined by potentiometric method. Knowing the electrode potential of the reference electrode ,the

electrode potential of the ion selective electrode is calculated, which in turns gives the ion

concentrations. The potentiometers are generally to read the ion content directly.

Applications:

Used to determine the concentrations of a number of cations and anions such as

H+,Li+,Na+,K+,Pb2+,F- etc.


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Glass electrode:

Principle:when the glass electrode is immersed in another solution, whose pH value is to be determined,

there develops a potential between the two surfaces of the membrane. The potential difference

developed is proportional to the difference in pH value.

pH electrodes are called glass electrodes. Glass electrode is made of glass tube ended with small

glass sensitive to proton. In side of the electrode is usually filled with buffered solution of chlorides in

which silver wire covered with AgCl2 is immersed.

Active part of the electrode is the glass tube while tube has strong and thick walls bubble is

made to be as thin as possible .Surface of the glass is protonated by both external and internal solution

till the equilibrium achived.Both sides of the glass are charged by the adsorbed protons, this charge is

responsible for potential difference.

Determination of pH:

To determine the pH of a given solution .the glass electrode is dipped in a solution whose pH need to be

determined.It is combined with a saturated calomel electrode.

Hg/Hg2Cl2/Cl-//solution (pH=?)/glass/0.1 N HCl/AgCl/Ag

The emf of the so formed cell is determined potentiometrically.

Ecell=EG-Ecal

=E0G-0.0591 pH-Ecal

pH=(E0G-Ecal-Ecell)/0.0591

E0G Value is evaluated bu dipping the glass electrode in a solution of knows pH and measuring

the emf of the cell formed when combined with a calomel electrode.


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Advantages:

Can be used in oxidizing and reducing environments.

Electrodes are not get poisoned.

Concentration cells:

A Concentration cell is an electro chemical cell that as two equivalent half cells of the same material

different only in Concentration.

The Concentration difference could be affected in the electrode material or in the electrolyte.

Further they could be divided into cells with transference or without transference.

Electrode Concentration Cells:

In these Cells the potential difference is developed between at different concentration dipped in

the same solution of the electrolyte.

Example: Two Hydrogen Electrodes at different gas pressure in the same solution of hydrogen ions

constitute a cell of this type.

Pt/H2 (Pressure P1) / H+ (a) / H2 (Pressure P2) / Pt

If P1>P2 Oxidation occurs at L.H.S Electrode and reduction occurs at R.H.S electrode.

The Nernst equation can be used to derive an expression for the potential of this electrode

concentration cell is

H2(P1) = 2H++2e-

2H++2e- = H2 (P2)

----------------------------------------

H2 (P1) = H2 (P2)

----------------------------------------

E Cell = 0.0591 / 2 log (P1/P2) at 250 C

The Standard cell potential is ‘0’ because a cell cannot derive a current through a circuit with identical

electrodes.


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Electrolyte Concentration Cell:

In these Cell electrodes are identical but these are immersed in the solution of same electrolyte of

different concentration. The source of the electrical energy in the cell is the tendency of the electrolyte

to defuse from a solution of higher concentration to that of lower concentration. With the expiry of

time, the two concentrations tend to becomes equal. Thus at the starting the emf of the cell is

maximum and it gradually falls to zero.

Such cells are represented bellow;

M/Mn+ (C1) // Mn+ (C2) / M (C2>C1)

Zn/Zn2+ (C1) // Zn2+ (C2) / Zn (C2>C1)

The emf of the cell is given by the following expression;

E Cell = 0.0591 / n log C2/ C1 at 250 C

The above examples are typical example of electrolyte concentration cell with transference.

Transference indicates the presence of salt bridge.

Examples of Electrolyte Concentration cell without transference is difficult to comprehend at

this stage and are not considered now.

Applications:

1).The concentration cells are used to determine the solubility of sparingly soluble salts valency of the

cation of the electrolyte and transition of the two allotropic forms of a metal used as electrodes etc.

2).Concentration cell corrosion occurs when two or more areas of a metal surface are in contact with

different concentration of the same solution.

Titrations: the mixing of two substances or two solutions in specified manner is called as titrations.

Examples;

Potentiometric titrations

Conduct metric titrations

pH metric titrations

Potentiometric titrations:

Potentiometric titration is a very interesting application of electrode potential. They involve a

study of variation of emf with the volume of titrant added. Here we discuss two types of titrations.

1).Precipitation titrations’ 2) acid-base titrations.


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Precipitation titrations: Consider the titration of Ag+(Silver) ions with KCl .On addition of any chloride

salt the silver ions fall out of the aqueous solution forming an insoluble precipitate of silver(silver

chloride).Hence we study the variation of silver electrode potential with the change in concentration of

silver ions.

EAg/Ag+ = E0Ag/Ag+-0.0592/2 log a Ag+

As the EAg/Ag+ varies in a logarithmic scale of the concentration of silver ions(Ag+) intiallt even for a

hundred fold variation the emf variation is only 0.118 V .

With the approach of neutralization point, the EAg+ gets smaller and smaller, while the emf, a logarithmic

value of Ag+ increases rapidly.

Thus we notice a sigmoid curve with the steep portion indicating the neutralization point.

The Ag/Ag+ electrode which is reversible to silver ion concentration is termed as the indicator.

Acid-base titrations: In case of acid-alkali titration a hydrogen electrode may be used as an electrode

indicator.

The suitable cell may be represented as;

(Pt)Hg, Hg2Cl2(s)/KCl (saturated)//H+/H2 (1 atm), (Pt)

The emf of indicator electrode is given as

E H2/H+ =E0

H2/H++0.0595 log [H+]

Since the standard emf of hydrogen electrode is zero

We have E H2/H+ =- 0.0595 log [H+]

E H2/H+ =-0.0595 pH

From the above expression we draw the conclusion that the emf decreases with decrease in the

concentration of H+ ions (or) increase in the Ph of the solution. On plotting the electrode potential (or)

cell potential against the volume of titrant added. A sigmoid curve is obtained. The volume correspond

to zero emf gives the titre value of the acid solution.


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Batteries

Battery is a device which transforms chemical energy into electrical energy. The term battery is

usually applied to a group of two or more electrical cells connected together electrically in series.

In general there are two types of batteries:

Primary batteries and secondary storage or accumulator batteries.

Primary batteries are constructed so that only one continuous or intermittent discharge can be

obtained. These batteries are used as a source of ‘dc’ power where the following requirements are

important:

Convenience is of major importance

The cost of a discharge is not much

Electrical charging equipment or power is not readily available.

Stand-by power is desirable without cell deterioration during periods of nonuse for days

or years

Secondary storage or accumulator batteries they are constructed in such a manner that they can be

recharged after partial or complete discharge. They are used as a source of ‘dc’ power where the

following requirements are important:

The battery is the primary source of power and large numbers of discharge-recharge cycles are

required, as in industrial hand or rider trucks, submarines, mine or switching locomotives.

The very large capacity is beneficial to the circuit, as in telephone exchanges.

Primary batteries: example:Dry (Leclanche) cell:

It consists of a Zinc anode which is shaped as a container for the electrolyte a carbon cathode

surrounded by Mno2 and a paste of NH4Cl andZnCL2 as a cathodic depolarizer. Cathode depolarizer

facilitates the H+ discharge reaction by removing the adsorbed hydrogen atoms. Such a cell known as

“dry cell” because of the absence of any mobile or liquid phase .Electrolyte consists of NH4Cl ,ZnCL2and

Mno2 to which starch is added to make it thick paste like so that it is less likely to leak .The cell is

enclosed in polypropylene cylinder and given a plastic coating to reduce leakage.

The Zn- Mno2 (dry) cell is represented as:

Zn/ Zn2+, NH4+/Mno2/C

The oxidation reaction at anode is

Zn(S) Zn2+ (aq) +2e-,


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The reduction reaction at cathode is

Mno2(S) +2H2O+2e_ Mn2O3+2OH-(aq)

But an reaction between OH- thus formed at cathode and NH4 +evolves NH3 (g).This liberated

NH3 (g) disrupts the current flow

NH4 + (aq) +OH-(aq) NH3 (g) +H2O(l)

This is prevented by a reaction of NH3 (g) with Zn2+ to form the complex [Zn (NH3)2]Cl2(s).

Thus the reaction at cathode can be written as:

2Mno2(s) + NH4 + (aq) +2e- [Zn(NH3)2]Cl2(s).

So the net reaction is;

Zn(s) + 2NH4 + (aq) + 2Cl-(aq) + 2Mno2(s) Mn2O3+

-[Zn (NH3)2] Cl2(s) +2H2O (l)

The various reactions involved in the dry cell cannot be reversed by passing electricity back

through the cell. Hence, the dry cell is primary cell.

Advantages:

1) These cells have voltage ranging from 1.25v to 1.50v.

2) Primary cells are used in the torches, radios, transistors, hearing aids, pacemakers, watches etc.

3) Price is low.

Disadvantages:

This cell does not have a long life, because the acidic NH4Cl corrodes the container even when the cell

is not in use.

Uses: It finds uses in calculators, transistor radios, flash-lights e.t.c..

Secondary batteries:

Example: The Lead –acid storage cell: A storage cell can operate both as a voltaic cell and as an

electrical cell .It has ability to work both ways, to receive electrical energy and also to supply it...When it

operates as a voltaic cell, it supplies electrical energy and as a result it eventually becomes run-down. It

then needs to be recharged .When being re-charged, the cell operates as an electrolytic cell.

Lead acid storage cell is the common example of storage. It is so classified because the

electrolyte is an acid and plates are largely leads. It consists of a lead-antimony alloy coated with

lead di oxide as cathode and spongy as lead as anode. The electrolyte is a 20 % solution of H2SO4 .In fact,

a lead accumulator for car consists of six lead-acid storage cells in series i.e. in electrolyte (H2SO4), six

pairs, with inert porous partitions in between, are dipped.


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Discharging: when the lead accumulator is used for supplying electrical energy, it is said to be

discharging. The lead electrode loses electrons, which flow through the wire.

I.e. at anode, oxidation takes place

Pb Pb2++2e-

The so formed Pb2+ ions then combine with SO42- ions.

Pb2++ SO42- PbSO4

The released electrons flow to the cathode, where PbO2 forms Pb2+.in other words, lead

undergoes reduction at cathode.

PbO2+2H++2e- Pb2++2H2O

Pb2+ ions then combine with SO42- ions.

Pb2++ SO42- PbSO4

Thus, the net reaction during use (discharging) is:

Pb+PbO2+4H++2SO42- 2PbSO4+2H2O+ENERGY.

The voltage of each cell is about 2.0 Volta at a concentration of 21.4 % H2SO4 at 250 C. Thatswhy;

a lead accumulator for car (consisting of six lead-acid storage cells in series) is capable of delivering 12

volts.

Charging: During discharging, PbSO4 is precipitated at both the electrodes. When PbSO4 covers

completely both anode and cathode, the cell stops functionating as a voltaic cell. Further use, it needs to

be re-charged.

Re-charging is done by passing an external e.m.f greater than 2 volts so that the reactions taking

place during discharging are reversed.

Reaction at anode; PbSO4+2H2O PbO2+SO42- +2H++2e-

Reaction at cathode; PbSO4+2H2O Pb+SO42-

Net reaction during charging is:

2PbSO4+2H2O+ENERGY Pb+PbO2+4H++2SO42-


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Applications:

These cells are used for supplying current to

electricalvehicals,railways,mines,laboratories,hospitals,automobiles,power stations, broadcasting

stations, in telephone exchanges, gas engine ignition, stand-by supplies(UPS) etc.

Nickel-Cadmium cell:

Nickel cadmium cell consists of a nickel wire gauze electrode. The anode consists of a mixture of spongy

cadmium with (78 %) cadmium hydroxide,18% iron,1 % nickel and 1% graphite. The cathode contains

nickel hydroxide (80%) cobalt hydroxide 2%, graphite 18% and traces of barium compound. Graphite

increases the conductivity, the cobalt and barium compounds increases the efficiency of active material

and also the life cycle.6M potassium hydroxide (KOH) is the electrolyte.

Electrode reaction: During discharging

At the anode, Cd+2OH- Cd (OH) 2 +2e-

At the cathode 2Ni (OH) 3 +2e- 2 Ni (OH)2+2 OH-

The net cell reaction is

Cd+Ni (OH) 3 2 Ni (OH) 2+2 Cd(OH)2

During charging, the above reaction is reversed.

Applications:

Nickel cadmium cells are used in battery operated appliances such as pocket calculators,

photo flash units, cordless garden tools, electric shaver’s instruments, alarm systems,

transmitters, receivers, emergency lighting, hearing aids, telemeters, etc.

Lithium batteries:

Lithium metal offers an attractive option be used as a battery anode material because of

its light weight, low electrode potential, high electrochemical equivalence and good

conductivity .For these reasons ,the use of lithium has predominated in the development of

high performance, high energy density primary and secondary batteries.


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Advantages of lithium batteries:

High cell voltage, up to 4 V, depending on the cathode material. This is because of the

very negative electrode potential of Li/Li+.

High energy density due to the low atomic mass of lithium.1 F is released by the

dissolution of 7 g of the metal.

Operation over a wide temperature range, from about 70-400 C.

Fuel cells:

A fuel cell is a galvanic cell in which the chemical energy contained in a readily available

fuel oxidant system is converted directly into electrical energy by means of electrochemical

process in which the fuel is oxidized at the cathode.

A fuel cell essentially consists of the following arrangements:

Fuel/electrode/electrolyte/electrode/oxidant.

Hydrogen oxygen fuel cell:

Hydrogen oxygen fuel cell is a simplest type in which hydrogen gas is used as a fuel and

oxygen as oxidant..A schematic diagram of H2-O2 fuel cell is shown below;

The cell consists of a porous carbon electrode impregnated with catalysts such as finely

divided platinum or palladium as anode. The cathode is also a porous carbon electrode

impregnated with platinum or silver as catalyst. The electrolyte is an aqueous solution of

KOH.The hydrogen gas fuel is continuously supplied at the anode and oxygen gas is supplied at

the cathode. As hydrogen gas diffuses through the anode, it is adsorbed on the electrode

surface and reacts with hydroxyl ions to form water. At the cathode oxygen diffusing through

the electrode is adsorbed and reduced to hydroxyl ions. These electrode reactions are

summarized bellow;


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At the anode: 2H2 4H++4e-

4H++4OH- 4H2O

-------------------------------

2H2+4OH- 4H2O +4e-

-----------------------------

At the cathode: O2+2H2O4e- 4OH-

The net cell reaction is, 2H2+O2 H20

Advantages of fuel cells:

Theoretically the efficiency can be 100 %.But actually it is about 50-80 %,owing to over

potential and resistance of the cell.

High efficiency of the energy conversion process.

No moving parts and so elimination of wear and tear.

Silent operation.

Absence of harmful waste products.

No need of charging.

unit 1 electrochemistry

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