Electrochemistry Lecture 1
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Transcript of 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.