01 - electrochemistry - ramesh- gec...

Electrochemistry

Dr. A. R. Ramesh

Assistant Professor of Chemistry

Govt. Engineering College, Kozhikode

1Dr. A. R. Ramesh - GEC CLT -

ELECTROCHEMISTRY - 2015

Electro Chemistry : Chemistry of flow of electrons

Electrochemical Cell

Chemical to Electrical

Electrolytic Cell

Electrical to Chemical

(Reverse of Electrochemical Cell)

Anode = -ve Anode = +ve

The electrons flow through the external circuit

Redox Reaction



A voltaic cell or Galvanic Cell

Two-half cells separated by a porous boundary with solid electrodes

connected by an external circuit

Electrons always travel in the external circuit from anode to cathode

Internally, cationsmove

toward the cathode, anions

move toward the anode,

keeping the solution neutral

(ionic movement through

electrolyte)



Single Electrode Potential(Origin of electrode potential)

When a metal rod is dipped in a solution of its own ions

Either oxidation or Reduction takes place

Oxidation

� Layer of negative charge (e-) at the electrode surface

� -vely charged electrode surface attract a layer of +vely charged ions

at the interface

� Develop an electrical double layer (EDL) at the metal-solution interface

� The potential difference between the metal and solution at the

interface (EDL) is the single electrode potential



Oxidation Reduction

(Reverse case)

Standard Electrode Potential (E0) is the electrode potential when the

electrode is in contact with a solution of unit concentration at 298 K.

It measures the tendency of the metallic electrode to lose (oxidation

potential) or gain (reduction potential) electrons, when it is in contact

with its own salt solution of 1M concentration at 250C



The electrode potential depends upon

� The nature of the metal and its ions

� Concentration of the ions in the solution and

� Temperature

Helmholtz Double Layer

A Helmholtz double layer (HDL) is an electrical double layer (EDL) of

positive and negative charges one molecule thick. This occurs at the

surface of a metal immersed in a solution.

Potential difference

ε = dielectric constant of the medium

σ = charge density

a = distance between the layersLayer of aligned ions, which is one

particle thick and then immediately

next to that, free solution.6

Dr. A. R. Ramesh - GEC CLT -


Gouy-Chapman Model

There is not a simple layer of ions but, an ionic

distribution that extends some distance from the surface

- called diffused layer

Stern Model

Rigid Helmholtz layer

Dispersed outside the Helmholtz

plane



Terms Used for Voltaic CellsTerms Used for Voltaic Cells

Half Cell & Cell



Electromotive Force (emf)

• Water only

spontaneously flows

one way in a waterfall.

• Likewise, electrons only

spontaneously flow one

way in a redox

reaction—from higher

to lower potential

energy.



Electromotive Force (emf)

• The potential difference between the cathode and

anode in a cell is called the electromotive force

(emf).

• It is also called the cell potential, and is designated

Ecell.



Daniell Cell

CuSO4 (aq)ZnSO4 (aq)

Cu metalZn metal

salt bridge

Cathode (reduction)

+ive

Anode (oxidation)

–ive

Zn(s) → Zn2+(aq) + 2e– Cu2+(aq) + 2e– → Cu(s)



Galvanic Cells (cont.)

• In turns out that we still will not get electron flow

in the example cell. This is because charge build-

up results in truncation of the electron flow.

• We need to “complete the circuit” by allowing positive

ions to flow as well.

• We do this using a “salt bridge” which will allow charge

neutrality in each cell to be maintained.



Salt Bridge

Functions – Complete the inner electrical circuit,

maintain electrical neutrality

Carefully merge Carefully merge

two solutions. two solutions.

Make CuSOMake CuSO44

more dense than more dense than

ZnSOZnSO44. Sheath . Sheath

Cu electrode in Cu electrode in

glass.glass.

Liq

uid

Liq

uid

--liq

uid

in

terf

ace

liq

uid

in

terf

ace

DaniellDaniell Cell without salt bridgeCell without salt bridge

Salt bridge makes Salt bridge makes cell construction cell construction

and operation and operation easier.easier.

Pack tube with a viscous, aqueousPack tube with a viscous, aqueous

solution of KCl or KNOsolution of KCl or KNO33. The. The

viscosity prevents mixing withviscosity prevents mixing with

the electrolytes. The ions permitthe electrolytes. The ions permit

exchange of charge. The chosenexchange of charge. The chosen

ions have similar mobility toions have similar mobility to

minimize junction potentials.minimize junction potentials.

Salt bridge is an inverted U tube filled with a concentrated solution

of KCl or KNO3 or NH4NO3 in agar-agar or gelatin



Electrochemical Series

Is an arrangement of elements in the increasing order of their

reduction potential



Electrochemical Series - Applications

�To know relative ease of oxidation and reduction

E0 = +ve (Reduction)

= -ve (Oxidation)

�To predict whether metal react with acid to give hydrogen

E0 = -ve only react with H2

�To calculate standard EMF of cell

E0Cell = E0

R –E0L = E0

Cathode – E0Anode



Nernst Equation

Note the difference between using natural logarithms and base10 logarithms.Note the difference between using natural logarithms and base10 logarithms.

Be aware of the significance of “n” Be aware of the significance of “n” –– the number of moles of electrons the number of moles of electrons

transferred in the process according to the stoichiometry chosen.transferred in the process according to the stoichiometry chosen.

E =Eo

−0.0257

nlnQ

E =Eo

−0.0592

nlogQ



EMF measurements(Poggendorf’scompensation method)

Cannot measured using voltmeter- cause current flow- change in

concentration

Principle: The EMF of the cell is opposed by an external source of EMF.

When there is no net flow of current in the circuit the imposed

potential will be equal to the EMF of the cell.



E = Battery (whose EMF is greater than Cell)

Ex = Cell (unknown EMF)

Es = Standard Cell (known EMF, Weston Cell)

AB = Potentiometer wire length

D = null deflection for Ex (AD)

D’ = null deflection for Es (AD’)



Standard Hydrogen Electrode

The convention is to select a particular electrode and assign its standard

reduction potential the value of 0.0000V. This electrode is the Standard

Hydrogen Electrode.

2H+(aq) + 2e– → H2(g)

H2

H+

Pt

The “standard” aspect to this cell is that the The “standard” aspect to this cell is that the

activity of Hactivity of H22(g) and that of H(g) and that of H++(aq) are both 1. This (aq) are both 1. This

means that the pressure of Hmeans that the pressure of H22 is 1 atm and the is 1 atm and the

concentration of Hconcentration of H++ is 1M, given that these are our is 1M, given that these are our

standard reference states.standard reference states.



Standard Hydrogen Electrode

• E° = 0 V (by

definition; arbitrarily

selected)

• 2H+ + 2e- → H2



Calculating Cell Potential

Because we tabulate reduction potentials, the cell potential is calculated (from

those tabulated numbers) as

Ecell = Ecathode - Eanode

The minus sign is present only because we are using reduction potential tables

and, by definition, an anode is where oxidation occurs.



Concentration Cells

• . . .a cell in which both compartments

have the same components but at

different concentrations



Concentration Cells

• Notice that the Nernst equation implies that a cell

could be created that has the same substance at both

electrodes.

• For such a cell, would be 0, but Q would not.Ecell°

• Therefore, as long as the concentrations are

different, E will not be 0.27

Dr. A. R. Ramesh - GEC CLT -


e–

e–e–

e –

Ag

1 M Ag+

1 M NO3–

Anode Cathode

Porous

diskAg

0.1 M Ag+

0.1 M NO3–



Batteries

• A battery is a galvanic cell or,

more commonly, a group of

galvanic cells connected in series.



How Does a Battery Work

cathode (+)

anode (-)

Electrolyte

Paste

Seal/cap

Assume a generalized battery



Battery

cathode (+):

Reduction occurs

here

anode (-):

oxidation

occurs here

e- flow

Electrolyte paste:

ion migration occurs

here

Placing the battery into a flashlight,

etc., and turning the power on

completes the circuit and allows

electron flow to occur



How Does a Battery Work

• Battery reaction when producing electricity

(spontaneous):

Cathode: O1 + e- → R1

Anode: R2 → O2 + e-

Overall: O1 + R2 → R1 + O2

• Recharging a secondary cell

– Redox reaction must be reversed, i.e., current is reversed (nonspontaneous)

Recharge: O2 + R1 → R2 + O1

– Performed using electrical energy from an external power source



Alkaline Dry Cell



Alkaline Dry Cell

Brass rod

Plated steel (-)

Anode:

Mixture of Zn

and KOH(aq)

Plated steel (+)

Cathode:

Mixture of

MnO2 and C

(graphite)

Paper or fabric

Separator

Insulators



Alkaline Dry Cell

Half-reactions

anode: Zn(s) + 2OH-(aq) --> ZnO(s) + H2O(l) + 2e-

cathode: 2MnO2(s) + H2O(l) + 2e- -->

Mn2O3(s) + 2OH-(aq)

overall: Zn(s) + 2MnO2(s) --> Mn2O3(s) + ZnO(s)

Ecell = 1.54 V



Batteries are Galvanic Cells• Car batteries are lead storage batteries.

• Pb +PbO2 +H2SO4 →PbSO4(s) +H2O



Lead Storage Battery

(cathode)

(anode)

6 x 2V = 12 V 39Dr. A. R. Ramesh - GEC CLT -



Half-reactions

anode: Pb(s) + HSO42-(aq) --> PbSO4(s) + H+ + 2e-

cathode: PbO2(s) + 3H+(aq) + HSO42-(aq) + 2e- -->

PbSO4(s) + 2H2O(l)

overall: Pb(s) + PbO2(s) + 2H+ + 2HSO4-(aq) -->

2PbSO4(s) + 2H2O(l)

Cell reaction reversed during recharging.




Half-reactions during recharging (nonspontaneous)

cathode: PbSO4(s) + H+ + 2e- --> Pb(s) + HSO42-(aq)

anode: PbSO4(s) + 2H2O(l) -->

PbO2(s) + 3H+(aq) + HSO42-(aq) + 2e-

overall: 2PbSO4(s) + 2H2O(l) -->

PbO2(s) + Pb(s) + 2H+ + 2HSO4-(aq)

Cell converted into electrolytic cell via application of

external electrical energy.



NiNi--Cad BatteryCad Battery

Anode (Anode (--))

Cd + 2 OHCd + 2 OH-- ------> Cd(OH)> Cd(OH)22 + 2e+ 2e--

Cathode (+) Cathode (+)

NiO(OH) + HNiO(OH) + H22O + eO + e-- ------> Ni(OH)> Ni(OH)22 + +

OHOH--



Fuel Cells

• Voltaic-like cell that operates with continuous

supply of energetic reactants (fuel) to the

electrodes

– utilize combustion reactions

– do not store chemical energy

• Not self-contained since reactants must be supplied to

the electrodes

– Example: Hydrogen-Oxygen fuel cell



Hydrogen-Oxygen Fuel Cell



Hydrogen-Oxygen Fuel Cell

Half-reactions

anode: 2H2(g) + 4OH-(aq) --> 4H2O(l) + 4e-

cathode: O2(g) + 2H2O(l) + 4e- --> 4OH-(aq)

overall: 2H2(g) + O2(g) --> 2H2O(l)



Fuel Cells

• Galvanic cells

• Reactants are continuously supplied.

• 2H2(g) + O2(g) → 2H2O(l)

• anode: 2H2 + 4OH− → 4H2O + 4e−

• cathode: 4e− + O2 + 2H2O → 4OH−



01 - electrochemistry - ramesh- gec...

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