electrochemistry lecture 1

7/23/2019 Electrochemistry Lecture 1

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Electrochemistry

Metals in HCl(aq)

2H+ + 2e- H2(g) E0 = 0.000 V

Cu2+ + 2e- Cu Pb2+ + 2e- Pb

E0 = 0.340 V E0 = - 0.125 V

Ni2+ + 2e- Ni Zn2+ + 2e- Zn

E0 = - 0.257 V E0 = - 0.762 V


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Standard Reduction (Half-Cell)

PotentialsStrong

oxidants:easily

reduced(forwardreaction)

Strongreductants:

easilyoxidised(reversereaction)

Voltaic Cells

In spontaneous oxidation-reduction (redox)

reactions, electrons are transferred and

energy is released.


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Voltaic Cells We can use that energy to do

work if we make the electrons

flow through an external

device.

We call such a setup a voltaic

cell.

The oxidation occurs at the anode.

The reduction occurs at the cathode.

Voltaic Cells

Once even one

electron flows

from the anode

to the cathode,

the charges ineach beaker

would not be

balanced and

the flow of

electrons would

stop.


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Voltaic Cells Therefore, we use

a salt bridge,

usually a U-

shaped tube that

contains a salt

solution, to keep

the charges

balanced:

Cations movetoward the

cathode.

Anions move

toward the anode.

Voltaic Cells

In the cell, electrons leave the anode and flowthrough the wire to the cathode.

As the electrons leave the anode, the cationsformed dissolve into the solution in the anode

compartment. As the electrons reach the cathode, cations in the

cathode are attracted to the now negative cathode.

The electrons are taken by the cation, and theneutral metal is deposited on the cathode.


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Standard Cell Potentials

The cell potential at standard conditions

can be found through this equation:

Because cell potential is based on the

potential energy per unit of charge, it is anintensive property.

Ecell = Ered (cathode) Ered (anode)

Cell Potentials

For the oxidation in this cell:

For the reduction: Ered = +0.34 V

Ered = 0.76 V


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How do we measure these values?

Half-cell potentials

In practice, measuring the potential of such a half-cell isnear-impossible

It requires measuring the potential difference betweenthe electrode and the solution without putting another

electrode in the solutionThe best we can do is put two half-cells back-to-backand measure thedifference between their potentials

If we take some particular half-cell as standard we canrefer all others to it


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Standard Hydrogen Electrode

Their values are referenced to a standard

hydrogen electrode (SHE).

By definition, the reduction potential for

hydrogen is 0 V.

2 H+ (aq, 1M) + 2 e H2 (g, 1 atm)

Potentiometry

The electromotive force, emf, of anelectrochemical cell is the potential differencethat must be applied to the cell in order to stopthe spontaneous cell reaction from occurring.

Cd|Cd2+

(1 M)||KCl(satd)|Hg2Cl2|Hg E0

= 0.64 V If a potential difference of Eappl = 0.64 V is

applied, the spontaneous cell reaction

Hg2Cl2 + Cd 2 Hg + 2 Cl- + Cd2+

is stopped, there is no current flow, the systemis at equilibrium, and its properties aredetermined by thermodynamics.


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Cell emf and GG for a redox reaction can be found byusing the equation:

G = nFE

A positive value of E and a negative value of

G both indicate that a reaction is

spontaneous.

Consequently, under standard conditions:

G =

nFEn is the number of moles of electrons transferred.

F is called Faradays constant: 1 F = 96,485 C/mol

= 96,485 J/V mol

The Nernst Equation

Remember that:

G = G + RT ln Q

This means:nFE = nFE + RT ln Q

E = E RTnF ln Q

E = E 2.303 RT

nF log Q


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At room temperature (298 K):

Thus the equation simplifies to:

E = E 0.0592

n log Q

2.303 RTF = 0.0592 V

Electrolytic CellsCd|Cd2+(1 M)||KCl(satd)|Hg2Cl2|Hg E

0 = 0.64 V

When Eappl < 0.64 V, electrons flow from

left to right in the external circuit, the

spontaneous cell reaction occurs - the cell

behaves as a voltaic cell (or galvanic cell).

When Eappl > 0.64 V, electrons flow in the

reverse direction, the reverse of the

spontaneous cell reaction occurs:

2 Hg + 2 Cl- + Cd2+ Hg2Cl2 + Cd

The cell behaves as an electrolytic cell.


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Voltaic and Electrolytic Cells


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Rate of Reaction

Since the current, i, which represents the number of coulombsof charge flowing per second, is stoichiometrically related tothe number of mole of (say) Cd2+ reacting per second, it is ameasure of the rate of the electrochemical reaction

v (mol s-1 cm-2) = i/nFA = j/nF

where n is the number of electrons transferred, F is Faradaysconstant,A is the area of the electrode, andj = i/A is thecurrent density (A m-2).

Kinetics, rather than thermodynamics rule here!

Polarisation

Theoretically, an applied potential, Eappl, slightly in

excess of the cell emf would cause the reverse of the

spontaneous cell reaction to occur.

In practice, the applied potential may need to exceed

the cell emf by anything up to a couple of tenths of a

volt before this is achieved!

The departure of the electrode potential from the

equilibrium value on passage of a current is called

polarisation.


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Ideal Polarised Electrodes An ideal polarised electrode shows a very large change

in potential upon the passage of a small current, and ischaracterised by a horizontal i-E profile.

An inert electrode (eg., Hg, Pt, Au) in a solutioncontaining only electro-inactive species approaches thisideal.

Ideal polarized (polarizable) electrode:An electrode iscalled "ideal polarizable" if no electrode reactions canoccurwithin a fairly wide electrode potential range.Consequently, the electrode behaves like a capacitorand only capacitive current ( no faradaic current) isflowing upon a change of potential. Many electrodes canbehave as an ideal polarized electrode but only within anelectrode potential range called the "double-layer range."Also called "completely-polarizable electrode" and"totally-polarized electrode." Contrast with ideal non-polarizable electrode.

Ideal polarized (polarizable) electrode


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Ideal Nonpolarisable Electrode The potential of an ideal nonpolarisable electrode does

not change on passage of current. It is an electrode offixed potential.

Nonpolarisable electrodes have a vertical i-E profile.

Reference electrodes (silver-silver chloride and SCE)approach non-polarisability at low current densities.

Reference electrodesReference electrode is an electrode which has a stable and well-

known electrode potential

The Standard Hydrogen Electrode (SHE) forms the basis of the

thermodynamic scale of oxidation-reduction potentials

Based on

2H+(aq) + 2e-H2(g)

Often impractical to use

Large area required : platinised platinum

Cumbersome, can be hazardous


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Reference electrodesReference electrode is an electrode which has a stable and

well-known electrode potential

Therefore to form a basis for comparison with all other

electrode reactions, Hydrogen's standard electrode potential

(E0) is declared to be zero at all temperatures.

Potentials of any other electrode is compared with that of the

SHE at the same temperature.

Common ref. electrodes : Ag/AgCl,

saturated calomel electrode (SCE)

Reference electrodes

Ag/AgCl (3M NaCl) is one of the most commonly used

Based on

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

-(aq)

Ideal non-polarizableelectrode

E = 0.220 V vs SHE

Unit activity at standard conditions

For Ag/AgCl (3M KCl)

E = 0.196 V


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Overpotential

The extent of polarisation is measured by

the overpotential,

= |E(i) - E(0)|,

the (absolute) difference between the cell

potential when there is no current flow,

E(0), and when there is current flow, i -

E(i).

The overpotential increases as the current

flowing through the system increases.

Overpotential

Overpotential is always deleterious to

performance - it decreases the potential

available during discharge:

Edischarge = E(0) -

and increases the potential required for

charging:

Echarge = (Eappl) = E(0) +


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Overpotential An electrode reaction O + ne R can be thought of

as composing a series of steps: mass transfer of O to and R away from the electrode

electron transfer at the electrode

chemical reactions before and after the electron transfer.

surface reactions - adsorption, desorption,electrodeposition.

Overpotentials can be associated with each of thesesteps. Overpotential serves as an activation energy

required to drive the processes at the rate reflectedby the current.

Overpotential means you must apply a greaterpotential before redox chemistry occurs

Overpotential


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iR Drop

With the passage of a current, i, through

a cell of resistance R, there is a potential

drop, iR (Ohms Law), that has the same

effect as an overpotential - it decreases

the potential available during discharge:

Edischarge = E(0) - - iR

and increases the potential required for

charging:

Echarge = (Eappl) = E(0) + + iR

Tafel Equation

The overpotential, , increases as the

current flowing through the system

increases.

Tafel (1905) found that the overpotentialis related to the logarithm of the current:

= a + b logi

where a and b are empirical constants.


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Metal-electrolyte interface


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A plane of charge due to complexation or dissociation exists

on a surface. For example, consider in this case that the

total charge is positive as shown in the previous figure.

Counterions attracted to the surface by electrostatics shield

the surface charge. The counterions are not all in the same

plane, however, because they are held in a dynamic

balance by electrostatic attraction and the tendency to

diffuse away. The concentration of ions in the diffuse layer

decays with distance from the surface. Thus the surfacecharge forms one layer and the diffuse shielding charge

forms the other layer, hence the term double layer.

The double layer has a certain structure. The Inner

Helmholtz Plane (IHP) is the plane cutting through the center

of the solvent molecules or specifically adsorbed ions.

The Outer Helmholtz plane (OHP) is the plane cutting

through the solvated negative ions at their position of closest

approach.

Ions in the diffuse layer are always being exchanged at thesurface, but the surface excess of counterions always

exactly balances the total charge associated with the solid.


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Helmholtz model

Gouy-Chapman Model


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Gouy-Chapman Model

Stern model


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The Capacitance Current

The charging or capacitance current, ic , is due to

the presence of the electrical double layer and it is

always present. This current, of course, is not

related to any movement of ions.

Ic = Cdl x V

Where:

Cdl = the capacitance of the electrical double layer

V = voltage scan rate

The capacitance current makes its presence felt

when measuring charge transfer (Faradaic)

processes at concentrations of 10-5 M.

electrochemistry lecture 1

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