Electrochemistry-19.2.14.pdf

8/10/2019 Electrochemistry-19.2.14.pdf

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Electrochemistry

Electrochemistry

Electrochemistry is a branch of chemistry that studies chemical reactions

which take place in a solutionat the interface of an electron conductor(the

electrode: a metal) and an ionic conductor (the electrolyte), and which

involve electron transfer between the electrode and the electrolyte.

It deals with the chemical reactions produced by passing electric current or

the production of electric current through chemical reactions.

Electrochemical Processes The chemical reactions involving electricity are called electrochemical

reaction

Electrochemical reactions are redox reactions where an oxidation and a

reductionreaction go side by side and are separated in space.

The reactions are two types

1. Induced electrolytic reactions- these non spontaneous reactions are

forced by the passage of electricity through the reactants

2. spontaneous redox reaction that can produce electricity

REDOX REACTIONS

Oxidation-reduction or redox reactionstake place in electrochemical cells.

The term redox comes from the two concepts of reduction and oxidation.

Oxidationdescribes the lossof electrons:

Zn (s) Zn2+

(aq) + 2 e -

Reductiondescribes the gain of electrons

Cu2+

(aq) + 2 e- Cu (s)

Each of the reaction is known as half- reaction or half-cell and both

must always go side by side

The net reaction: Zn (s) + Cu2+

(aq) Zn2+

(aq) + Cu (s)


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Electrochemical Cells

An electrochemical cell is a device capable of either deriving electricalenergy

from chemical reactions, or facilitating chemical reactions through the

introduction of electrical energy.

There are two types of electrochemical cells.

TYPES:

1. Electrochemical cell and Electrolytic Cells

2.Galvanic cell and concentration cell

3. Reversible and irreversible cell

1. Spontaneous reactions occur in galvanic (voltaic) cells;

2. Nonspontaneous reactions occur in electrolytic cells.

Both types of cells contain electrodes where the oxidation and reduction

reactions occur.

Oxidation occurs at the electrode is termed as the anode and reductionoccurs at the electrode is called as the cathode.

1. Galvanic or Voltaic cell: Produces energy by a spontaneous reaction which

produces electricity as a result of electron transfer

The cells used for the generation of electrical energy from chemical

reactions are called galvanic or voltaic cells.

The redox reaction in a galvanic cell is a spontaneous reaction.

For this reason, galvanic cells are commonly used as batteries. Galvanic

cell reactions supply energy which is used to perform work. The energy

is harnessed by situating the oxidation and reduction reactions in separate

containers, joined by an apparatus that allows electrons to flow.


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Ex. A common galvanic cell is the Daniell cell, batteries, corrosion, etc

2. Electrolytic Cells The redox reaction in an electrolytic cell is non-

spontaneous. Electrical energy is required to induce the electrolysis reaction.

The cells used for electrolysis are called electrolytic cells

ex. Charging of battery, electroplating ,etc

Electrodes & Charge

The anode of a galvanic cell is negatively charged, since the spontaneous

oxidation at the anode is the sourceof the cell's electrons or negative charge.

The cathode of a galvanic cell is its positive terminal.

The anodeof an electrolytic cell is positive (cathode is negative), since theanode attracts anions from the solution.

In both galvanic and electrolytic cells, oxidation takes place at the anode and

electrons flow from the anode to the cathode.

An electrochemical cell is obtained by coupling two half cells

1. Anode: oxidation half-cell reaction takes place

2.

Cathode: reduction half-cell reaction occurs

3. Salt bridge: A salt bridge is often employed to provide ionic contact

between two half-cells with different electrolytesto prevent the solutions

from mixing and causing unwanted side reactions.

Ex.: filter paper soaked in KNO3

Functions of salt bridge:

1. The salt bridge allows ions to move on either side and maintain the

electrical neutrality of the electrolyte on both sides.

2. It serves as a bridge to complete the electric circuit.

3.

It prevents the liquid junction potential between the two electrodes.


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Liquid junction potential occurs when two solutions of different

concentrations are in contact with each other. The more concentrated

solution will have a tendency to diffuse into the comparatively less

concentrated one. The rate of diffusion of each ion will be roughly

proportional to its speed in an electric field. If anion diffuses more

rapidly than the cation, it will diffuse ahead into the dilute solution

leaving the later negatively charged and the concentrated solution

positively charged. So an electrical double layer of positive and negative

charges will be produced at the junction of the two solutions. So at the

point of junction, a difference of potential will develop because of theionic transfer. This potential is called liquid junction potential or

diffusion potential. The magnitude of the potential depends on the

relative speeds of the ions.

4.

External circuit: These two half-cells joined together by wire through

which electrons flow.

5. Electrolyte: Internal pathway that allows ions to migrate between them so

as to preserve electro neutrality.


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Difference in electrolytic cell and galvanic cell

Electrolytic cell Galvanic cell

1. Conversion of electrical energy

into chemical energy

1. Chemical energy into electrical

energy

2. The anode carries positive

charge vice versa.

2. The anode carries negative

charge Vice versa.

3. Electrons are supplied to the cell

from the external power supply.

3. Electrons are drawn from the

cell.

4. Not a spontaneous reaction. Eo

cellis -ve, then the process is

nonspontaneous. E.g electroplating

4. Spontaneous reaction. Eo

cell is+ve, then the process is

spontaneous. eg. Corrosion

5. The extent of chemical reaction

occurring at the electrode depends

on the quantity of electricity passed

& is governed by Faradays law of

electrolysis.

5. The e.m.f of the cell depends

on the concentration of the

electrolyte and chemical

nature of the electrode (Nernst

Equation)

6. The amount of electricity passed

during electrolysis is measured

by Coulometer.

6. The e.m.f produced in the cell

is measured by potentiometer.


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ELECTRODE POTENTIAL (E)

When a metal (M) is placed in the solution of its own ion it may act as an

anode or cathode.

As an anode: Positive metal ions passing into solution (oxidation) leaving

behind e-s on the electode.

The anode attains negative charge due to accumulation of e-s which attracts

positively charged free ions (cations) from the solution. Due to the

attraction the positive ions remain close to the metal.

M Mn++ne-

As a cathode: Positive ions depositing on the metal electrode (reduction)

and it attracts negatively charged ions.

Mn+

+ne-----M

HELMHOLTZ ELECTRICAL DOUBLE LAYER

Thus a sort of electrical double layer (positive or negative ions) is formed

all around the metal. This layer is called Helmholtz electrical double layer.

This layer prevents further passing of or deposition of metal ions on the

metal. Consequently a difference in potential (Galvanic potential) set up

between the metal and its solution.

At equilibrium, the potential difference is constant, which is known as

Electrode potential of a metal.EMF: The diff. of potential between the two electrodes in a voltaic cell which

causes flow of current/electrons is called the electromotive force


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Electrode potential of the metal shows its tendency to under go loss or gain of

electrons


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Electrode potential (E):

The measure of tendency of an electrode to lose or gain electrons when it is

contact with a solution of its own ion is called electrode potential.

Standard electrode potential (E

o

):The measure of tendency of an electrode to lose or gain electrons, when it is

contact with a solution of its own salt of 1 molar concentration at 25oC is known

as standard electrode potential.

Oxidation electrode potential (Eoxid):

The measure of tendency of an electrode to lose electrons when it is contact with a

solution of its own ion is called oxidation electrode potential.

Reduction electrode potential (E red):The measure of tendency of an electrode to

gain electrons when it is contact with a solution of its own ion is called electrode

reduction potential.

Factors affecting electrode potential

1. Nature of the electrode metal

2. Temperature

3. Concentration of metal ions in solution

Measurement of Single electrode potential

Impossible to know the absolute value of single electrode potential.

To measure the ele. pot of one electrodes which is connected to reference

electrode to form a complete cell

The electrode whose electrode pot. is either exactly known or arbitrarily

fixed is called reference electrode

The pot. of one electrode is fixed arbitrarily then another electrode pot. can

be measured

EMF of the complete cell can be directly read from the potentiometer.

Ecell= Ecathode- E

anode


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Potential of the single electrode can be calculated by

Nernst Equation:

The final and the most fundamental form the Nernst equation is written as:E = E

o RT/nF ln ap/ar

where R is the universal gas constant, T is the absolute temperature in degrees

Kelvin, z is the charge number of the electrode reaction (which is the number of

moles of electrons involved in the reaction as written), and F is the Faraday

constant (96,500 C mole-1

). The notation aprepresents the chemical activitiesof

all of the species which appear on the product side of the electrode reaction and

the notation arrepresents the chemical activities of all of the species which appear

on the reactant side of the electrode reaction.

Derivation

Nernst equation was derived from Vant Hoff reaction isotherm for the

redox reaction, M2+

+ 2e---M

Vant Hoff Reaction Isotherm

G = - RT ln K +RT ln Q

Vant Hoff reaction isotherm is an equation for the change in free energy

during the chemical reaction.

it relates the G (free energy change) and K (equilibrium constant) for

the redox reaction as

G = - RT ln K +RT ln Q ---------------(1)

Q (Reaction Quotient) = aC. aD/ aA . aB = [P] /[R]

aC. aD = chemical activity of products

aA . aB = chemical activity of reactants


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The activities of gases are usually taken as their partial pressures and

the activities of solutes such as ions are usually taken as their molar

concentrations.

G = - RT ln K +RT ln [P] /[R] ---------------(1)

At standard conditions of T and P, at equilibrium [P] = [R]

Go = - RT ln K

So equation 1 becomes

G = Go +RT ln [P] /[R] ---------------------(2)

Electrical energy (nFE) arises from the expense of free energy of the system

(- G).

Let n faraday charge be taken out of a electrode or cell of emf E; then

work done by the cell will be calculated as:

Work = Charge x Potential = nFE

Work done by the cell is equal to decrease in free energy.

-G = nFE

-Go= nFE

o

G = Go+ RT ln [product] / [reactant]

-nFE = -nFEO + 2.303 RT log [product] / [reactant]

Ecell= Eocell

- 2.303 RT / nF log [P]/[R]

R = 8.314 J/K/mole

At 25oC, T =298 K

F= 96500 Coulombs

The oxidation potential of an electrode for the reaction

M----Mn+

+ ne-

Eoxid = Eooxid- 0.0592/n log [Mn+] / [M] ----- half cell

[M] = 1

Eoxid = Eo

oxid- 0.0592/n log [Mn+]


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ELECTROCTREMICAL SERIES (e.m.f series): The arrangement of metals in the

increasing order of their std. reduction electrode potential

Half-Reaction E0

Li++ e- Li

Na+ + e- Na

Mg2+ + e- Mg

Al3+

+ 3e- Al

Ti2+

+ 2e- Ti

Mn2+ +e- Mn

Zn2+

+ 2e- ZnFe

2++ 2e- Fe

Co2+

+ 2e- Co

Ni2+

+ 2e- Ni

Fe3+

+ 3e- Fe

2H++ 2e- H2 (g)

Sn4+

+ 2e- Sn2+

Cu2++ 2e- Cu

Fe3+

+ e- Fe2+

Ag++ e- Ag

Pt4+

+ e- Pt

Mn4+

+ 2e- Mn

2+

Cr6+

+ 3e- Cr

3+

Au+ + e- Au

Mn7++ 5e- Mn2+

Cr4+

+ 1e- Cr3+

Au3+

+ 3e- Au

F2+ 2e- 2F-

-3.05

-2.70

-2.40

-1.66

-1.63

-1.18

-0.76-0.44

-0.28

-0.25

-0.04

0.00

+.15

+0.34

+0.77

+0.80

+0.86

+1.23

+1.33

+1.50

+1.51

+1.60

+1.69

+2.87


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Application of electrochemical series

1.

Relative ease of oxidation or reduction can be predicted

The metals on the top with ve reduction pot. Can more easily undergooxidation & act as anode.

Ex. Zn = -0.76 V favours oxidation reaction

While the metals at the bottom with +ve reduction. Pot. has great

tendency to undergo reduction & act as cathode

Ex. Cu = +0.34 V favours reduction reaction.

2. Cell representation can be predicted

A galvanic cell formed by two half cells with 2 diff. metals.

From EMF series, the metal which is undergoing oxidation or reduction can

be predicted

The electrode undergoes oxidation (anode) is written at left

The electrode undergoes reduction (cathode) is written at right

Ex. For Danial cell with Zn and Cu electrodes

The cell representation

Zn / Zn2+

// Cu2+

/ Cu

The cell reaction can also be written

Anodic reaction:

Zn (s) Zn2+

(aq) + 2 e

Cathodic reaction:

Cu2+

(aq) + 2 e- Cu (s)

The net reaction:

Zn (s) + Cu2+

(aq) Zn2+

(aq) + Cu (s)


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3. Calculation of Std emf (Eocell) of the cell

Oxidn Half cell|| Redn half cell

Eocell= E

oright - E

oleft

Or E

ocell= E

ocathode (red.) E

oanode (red.)

Or Eocell= E

oref.) E

ounknown

Ex.

Std e.m.f (Eocell) of (Danial cell) Zn-Cu cell

Eocell= EoCu EoZn= +0.34 (-0.76) V = +1.1 V

4. Calculation of std. free energy (Go)

Go = - nFEo

and equilibrium constant (Keqm)

Go= - RT ln Keqm

- nFEo= - RT ln Keqm

Eo= __RT_ 2.303 log Keqm

nF

log Keqm= nEo/ 0.0592 at 25

oC

5. The spontaneity or feasibility of the cell reaction can be predicted

Spontaneity of the redox reaction can be predicted from the e.m.f value of

complete cell reaction.

If the value of E cell is + ve,

(std. free energy (Go) is negative, since G

o = - nFE

o) the reaction is

feasible.

If E cell is -ve, Go

is positive. Then the reaction is not feasible.


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6. Hydrogen liberating tendency of the metal can be predicted

The metal with low reduction potential will displace H2 from an acid

solution.

Zn+ H2SO4 ZnSO4+ H2The metal with +ve potential will not displace H2 from an acid solution

Ag + H2SO4 no reactn

7. Replacing tendency of a metal (M) by another M:

The metals with low reduction potential undergo oxidation and pass into the

solution and the M with high red. pot. Undergo reduction and get deposits on

electrode.

Zn, Ni undergo dissolution in CuSO4 solution and will displace Cu from solution.

8. Corrosion tendency of M:

The metals higher in the series are anodic or more active and they are more prone

to corrosion.

The metals lower in the series are noble metals and they are less prone to

corrosion.


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Reversible and irreversible cells

Daniell cell has the emf value 1.09 volt. If an opposing emf exactly equal to

1.09 volt is applied to the cell, the cell reaction stops

Zn + Cu2+

--> Cu + Zn2+

but if it is increased infinitesimally beyond 1.09 volt, the cell reaction is

reversed.

Cu + Zn2+

--> Zn + Cu2+

Such a cell is termed a reversible cell.

Thus, the following are the main characteristics of reversible cell:

(i) The chemical reaction of the cell stops when an exactly equal opposing

emf is applied.

(ii) The chemical reaction of the cell is reversed and the current flows in

opposite direction when the opposing emf is slightly greater than that of the

cell.

(iii) The cell produces a small emf if the opposing emf is infinity smaller

than that of the cell

Any other cell which does not obey the above two conditions is termed as

irreversible. A cell consisting of zinc and copper electrodes dipped into thesolution of sulphuric acid is irreversible. Similarly, the cell

Zn|H2S04(aq)|Ag

is also irreversible because when the external emf is greater than the emf of

the cell, the cell reaction,

Zn + 2H+--> Zn

2++ H2

is not reversed but the cell reaction becomes

2Ag + 2H+--> 2Ag

++ H2


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Concentration cell

A concentration cell is made up of two half cells having identical electrodes,

identical electrolyte, except that the concentration of the reactive ions at the

two electrodes are different.

Theelectrical energy in a concentration cell arises from the transfer of

electrons from the electrode in the lower concentration side to the electrode

in higher concentration side.

The two half - cells may be joined by a salt bridge.

When a metal (M) electrode is dipped in a solution containing its own ions

[Mn+], then the potential (E) is developed at the electrode at the electrode,

the value of which varies with conc. Of the ions in accordance with the

Nernsts equation:

E =Eo - 2.303 RT / nF log 1/ [Mn+

]

E =Eo + 2.303 RT / nF log C

Let us consider a general conc. cell represented as:M/M

n+(C1M) II M

n+(C2M) I M

+

Where C1 and C2 are the concentrations of active metals ions (Mn+

) in

contact with two electrodes respectively and C2> C1.

THE CELL REACTIONS ARE:

At left electrode (anode): M --Mn+

(C1.) + ne-

At right electrode (cathode): Mn+

(C2) + ne--M

THE NET CELL REACTION: Mn+

(C2) ---Mn+

(C1.)

E.M.Fcell. of cell = Eright - Eleft

= [Eo+ 2.303RT/nF log C2] [Eo+ 2.303RT/nF log C1]

= 2.303RT/nF log C2/C1


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Ecell = 0.0592/n log C2/C1

The e.m.f so developed is due to the mere transference of metal ions from

the solution of higher conc. (C2) to lower conc. (C1.).

Applications:

i. Determination of solubility of sparingly soluble salts:

ii.

Determination of the valency of an ion.


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REFERENCE ELECTRODS:

The electrode with exactly known standard potential or fixed as constants or zero,

with which we can compare the potentials of other electrodes

The Characteristics of reference electrodes1. The electrode potential should be fixed as constants or zero.

2. The temperature co-efficient should be very low.

3. It should be a reversible electrode. It can function both as an anode and a

cathode

4. It should be easy to handle and use in laboratory.

1. Primary reference electrode

Standard Hydrogen Electrode:

Construction

This is a gas electrode.

It consists of a thin rectangular platinum foil which is coated with fresh platinum

black to increase the adsorption capacity of H2gas.

The inner tube is enclosed in an outer jacket having an inlet tube for sending in H 2

gas and has a perforated wider base for the escape of excess of H2. This unit is

dipped in 1 M HCl taken in a beaker such that the metal foil remains in the

solution.

In the above system, when the H2gas at a pressure of 1atm is bubbled through 1M

HCl, the electrode (constructed) or formed is called STANDARD HYDROGEN

ELECTRODE (SHE) or Normal H2electrode (NHE).

Working

This is represented as Pt, H2(1 atm) / H

+

(IM) When pure and dry H2gas ispassed through the inlet tube, a part of the gas gets adsorbed and the excess

bubbles out through the perforations.

Depends on the other half cell connected to SHE, SHE can under go either

oxidation or reduction.


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The electrode is represented as:

Pt, H2(1 atm); H+(1M)

The electrod potential of SHE is fixed as zero.The either of the reaction takeplace

H+(aq) + e

--- H2(g)

H2(g) --H+(aq) + e-

Act both as anode and cathode based on another M connected, it is a

reversible electrode.

When M electrode with lower reduction potential than H2, is coupled with

SHE, M undergoes oxidation and act as anode. Ex. Zn and SHE acts as

cathode

If the M electrode possesses higher reduction potential than H2, is coupled

with SHE, undergoes reduction and act as cathode. Ex. Cu


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Determination of single electrode potential of a metal

1.

If the single electrode Zn in ZnSO4 is connected to SHE then the emf of the

complete cell determined by means of a potentiometer. Since emf of SHE is

zero, the observed e.m.f gives directly the e.m.f of half-cell containing thesolution under test.

Zn has the tendency to under go oxidation, act as anode.

Pot. of the Zn single electrode can be obtained by

E cell= Ec.-Ea

Ecell= Ecathode (red.) Eanode (red.)

Ecell= ESHE Eunknown(Zn2+

, Zn)

ESHE= 0

Eunknown(Zn2+

, Zn)= 0 Ecell

Eunknown(Zn2+

, Zn)= Ecell

2. If the single electrode Cu in CuSO4 is connected to SHE then the emf of the

complete cell determined by means of a potentiometer. Since emf of SHE is zero,

the observed e.m.f gives directly the e.m.f of half-cell containing the solution

under test.

Cu has the tendency to undergo reduction & act as cathode

Pot. of the Cu single electrode

Ecell= Eunknown(Cu2+

, Cu) - ESHE

ESHE= 0

Ecell= Eunknown(Cu2+

, Cu)


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Zn possess lower red. Pot. than H2. Zn undergoes oxidation and act as anode.

H2 undergoes reduction.

H+(aq) + e

--- H2(g) --- act as cathode

Cu possess higher red. Pot. than H2. Cu undergoes reduction and

act as cathode. H2undergoes oxidation,

H2(g) --H+(aq) + e

- --- act as anode


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The electrode potential for decinormal 0.1 N KCl is +0.3335 V for normal

KCl solution +0.2810 V and for saturated calomel electrode (SCE) is

+0.2422 V

The electrode is represented as: Hg, Hg2Cl2(s), KCl (sat.solution) / Pt

The electrode potential of SCE is arbitrarily fixed as + 0.2422 V.

The electrode is based on the redox reaction

Hg22+

(s)+ 2e

-


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Eunknown (Zn2+

, Zn) = +0.2422 Ecell

2. When Cu electrode is connected to SCE

Cu undergo reduction and act as cathode and SCE act as anodeAt SCE oxidation takes place

Hg Hg++ e

-

2 Hg+

+ 2Cl2 -Hg2Cl2

The single electrode pot. Can be calculated

Ecell = Eunknown(Cu2+

, Cu) - ESCE

ESCE = +0.2422 V

Eounknown(Cu

2+, Cu) = Ecell+ 0.2422

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Electrochemistry-19.2.14.pdf