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UNIT – 5 ELECTRO ANALYSIS AND SURFACE MICROSCOPY
UNIT V ELECTRO ANALYSIS AND SURFACE MICROSCOPY 9
Electrochemical cells- Electrode potential cell potentials –
potentiometry- reference electrode – ion selective and molecular
selective electrodes – Instrument for potentiometric studies –
Voltametry – Cyclic and pulse voltametry- Applications of voltametry
. Study of surfaces – Scanning probe microscopes – AFM and STM.
1
Syllabus

• A dc electrochemical cell consists of two electrical conductors called
electrodes, each immersed in a suitable electrolyte solution.
• For a current to develop in a cell, it is necessary
1) that the electrodes be connected externally with a metal
conductor
2) that the two electrolyte solutions be in contact to permit
movement of ions from one to the other
3) that an electron- transfer reaction can occur at each of the two
electrodes.
Figure below shows an example of a simple electrochemical cell.
• It consists of a silver electrode immersed in a solution of silver nitrate
and a copper electrode in a solution of copper sulfate.
• The two solutions are joined by a salt bridge, which consists of a tube
filled with a solution that is saturated with potassium chloride or,
sometimes, a different electrolyte.
2
5 . 1 Electrochemical cells

3

• The two ends of the tube are fitted with porous plugs or disks that
permit the movement of ions across them but prevent siphoning of
liquid from one electrolyte solution to the other.
• The purpose of the bridge is to isolate the contents of the two halves
of the cell while maintaining electrical contact between them.
• Isolation is necessary to prevent direct reaction between silver ions
and the copper electrode.
• The cells in Figure contain two so-called liquid junctions, one being
the interface between the silver nitrate solution and the salt bridge;
the second is at the other end of the salt bridge where the electrolyte
solution of the bridge contacts the copper sulfate solution.
• A small junction potential at each of these interfaces may influence
significantly the accuracy of the analysis.
4

Conduction in a Cell :
• The high impedance meter in Figure -a allows the voltage developed
in the cell to be measured without drawing significant current from
the cell.
• If we replace the high-impedance meter in Figure - a with a
low-resistance wire or load as shown in Figure ‑ b, the circuit is
completed, and charge flows.
• Charge conduction results from three distinct processes in various
parts of the cell shown in Figure –b.
5
1. In the silver and copper electrodes, as well as in the external
conductor, electrons are the charge carriers, moving from the
copper electrode through the external conductor to the silver
electrode.

6

2. Within the solutions the flow of charge is the result of migration of
both cations and anions. In the halfcell on the left, copper ions migrate
away from the electrode into the solution, and sulfate and hydrogen
sulfate ions move toward it; in the half-cell on the right, silver ions move
toward the electrode and anions move away from it. Inside the salt
bridge, charge is carried by migration of potassium ions to the right and
chloride ions to the left. Thus, all of the ions in the three solutions
participate in the flow of charge.
3. A third process occurs at the two electrode surfaces. At these
interfaces, an oxidation or a reduction reaction couples the ionic
conduction of the solution with the electron conduction of the electrode
to provide a complete circuit for the flow of charge. The two electrode
processes are described by the reactions
7

Galvanic and Electrolytic Cells :
• The net cell reaction that occurs in the cell shown in Figure -b is the
sum of the two half-cell reactions shown as Equations (1 ) and (2)
That is,
8
• The voltage of this cell is a measure of the tendency for this reaction
to proceed toward equilibrium.
• Thus, as shown in Figure (a), when the copper and silver ion
concentrations are 0.0200 M, the cell voltage is 0.412 V, which
shows
that the reaction is far from equilibrium.
• If we connect a resistor or other load in the external circuit as shown
in Figure b, a measurable current in the circuit results and the cell
reaction occurs.

• As the reaction proceeds, the voltage becomes smaller and smaller,
and it ultimately reaches 0.000 V when the system achieves
equilibrium.
• Cells, that produce electrical energy are called galvanic cells.
• In contrast, electrolytic cells consume electrical energy.
• For example, the Cu-Ag cell can operate in an electrolytic mode by
connecting the negative terminal of a battery or dc power supply to
the copper electrode and the positive terminal to the silver
electrode, as illustrated in Figure -c.
• If the output of this supply is adjusted to be somewhat greater than
0.412 V, as shown, the two electrode reactions are reversed and the
net cell reaction is
9

10
(c) an electrolytic cell.
A cell in which
reversing the
direction of the
current simply
reverses the reactions
at the two electrodes
is termed a chemically
reversible cell.

Anodes and Cathodes :
• By definition, the cathode of an electrochemical cell is the electrode
at which reduction occurs.
• The anode is the electrode where oxidation takes place.
• These definitions apply to both galvanic cells under discharge and to
electrolytic cells.
• For the galvanic cell shown in Figure - b, the silver electrode is the
cathode and the copper electrode is the anode.
• On the other hand, in the electrolytic cell of Figure - c, the silver
electrode is the anode and the copper electrode the cathode because
the half-cell reactions are reversed.
11

Cells without Liquid Junctions :
• The cell shown in Figure -1 has two liquid junctions, one between the
silver nitrate solution and one end of the salt bridge, the other
between the copper sulfate solution and the salt bridge.
• Sometimes it is possible and advantageous to prepare cells in which
the electrodes share a common electrolyte and thus eliminate the
effect of junction potentials.
• An example of a cell of this type is shown in Figure below.
• If the voltmeter were removed and replaced by a wire, silver would
behave as the cathode.
• The reaction at the cathode would be
12

Under discharge, hydrogen is consumed at the platinum anode:
13
The overall cell reaction is then obtained by multiplying each term in the
first equation by 2 and adding the two equations.
That is,
• The direct reaction between hydrogen and solid silver chloride is so
slow that the same electrolyte can be used for both electrodes
without significant loss of cell efficiency because of direct reaction
between cell components.

14
FIGURE A galvanic cell without a liquid junction.

• The cathode reaction in this cell is interesting because it can be
considered the result of a two-step process described by the
equations
15
• The slightly soluble silver chloride dissolves in the first step to provide
an essentially constant concentration of silver ions that are then
reduced in the second step.
• The anodic reaction in this cell is also a two-step process that can be
formulated as

• Hydrogen gas is bubbled across the surface of a platinum electrode
so that the concentration of the gas at the surface is constant at
constant temperature and constant partial pressure of hydrogen.
• It is to be Noted that in this case the inert platinum electrode plays
no direct role in the reaction but serves only as a surface where
electron transfer can occur.
• The cell in Figure above is a galvanic cell with a potential of about
0.46 V.
• This cell is also chemically reversible and can be operated as an
electrolytic cell by applying an external potential of somewhat
greater than 0.46 V.
• It is to be Noted that you cannot tell whether a given electrode will
be a cathode or an anode unless you know whether the cell is
galvanic under discharge or electrolytic.
16

Solution Structure: The Electrical Double Layer :
• An electrode can donate or accept electrons only from a species that
is present in a layer of solution that is immediately adjacent to the
electrode, usually within a few angstroms.
• Thus, as a result of the chemical and physical changes that occur at
the electrode-solution interface, this layer may have a composition
that differs significantly from that of the bulk of the solution.
• For example, let us consider the structure of the solution
immediately adjacent to an electrode in an electrolytic cell when a
positive voltage is first applied to the electrode (e.g., the Ag
electrode)
• Immediately after applying the voltage, there is a momentary surge
of current, which rapidly decays to zero if no reactive species is
present at the surface of the electrode.
17

• This current is a charging current that creates an excess (or a
deficiency) of negative charge at the surface of the two electrodes.
• Because of ionic mobility, however, the layers of solution
immediately adjacent to the electrodes acquire a charge of the
opposite sign.
• This effect is illustrated in Figure below:
18
FIGURE : Electrical double
layer formed at electrode
surface as a result of an
applied potential.

• The surface of the metal electrode has an excess of positive charge
because of an applied positive voltage.
• The charged solution layer consists of two parts:
(1) a compact inner layer (d0
to d1
), in which the potential decreases
linearly with distance from the electrode surface and
(2) a diffuse layer (d1
to d2
), within which the decrease is approximately
exponential .
• This entire array of charged species and oriented dipoles (such as
water molecules) at the electrode-solution interface is called the
electrical double layer.
Faradaic and Nonfaradaic Currents :
• Two types of processes can conduct charge across an electrode
solution interface.
19

• One type involves a direct transfer of electrons via an oxidation
reaction at one electrode and a reduction reaction at the other.
• Processes of this type are called faradaic processes because they are
governed by Faraday’s law, which states that the amount of chemical
reaction that occurs at an electrode is proportional to the current,
called a faradaic current.
• Under certain conditions a range of voltages may be applied to a cell
that do not produce faradaic processes at one or both of the
electrodes.
• Faradaic processes may be prevented either because electrons do
not have sufficient energy to pass over the potential energy barrier
at the electrode-solution interface (thermodynamic reasons) or
because the electron-transfer reaction is not fast enough on the
time scale of the experiment (kinetic reasons). 20

• To understand the basic difference between a faradaic and a
nonfaradaic current, imagine an electron traveling down the external
circuit to an electrode surface.
• When the electron reaches the solution interface, it can do one of
only two things.
• It can remain at the electrode surface and increase the charge on the
double layer, which constitutes a nonfaradaic current.
• Alternatively, it can leave the electrode surface and transfer to a
species in the solution, thus becoming a part of the faradaic current.
Mass Transfer in Cells with the Passage of Current :
• Because an electrode can probe only a very thin layer of solution at
the electrode surface (d0
to d1
) a faradaic current requires continuous
mass transfer of reactive species from the bulk of the solution to the
electrode surface. 21

• Three mechanisms bring about this mass transfer:
✔ convection,
✔ migration,
✔ and diffusion.
• Convection results from mechanical motion of the solution as a result
of stirring or the flow of the solution past the surface of the
electrode.
• Migration is the movement of ions through the solution brought
about by electrostatic attraction between the ions and the electrode
of the opposite charge.
• Diffusion is the motion of species caused by a concentration gradient.
22

• The cell reaction of an electrochemical cell as being made up of two
half-cell reactions, each of which has a characteristic electrode
potential associated with it.
• These electrode potentials measure the driving force for the two
half-reactions when, by convention, they are both written as
reductions. Thus, the two half-cell or electrode reactions for the cell is
shown as
23
5 . 2 Electrode potential cell potentials
To obtain the spontaneous cell reaction, the second half-reaction is
subtracted from the first to give

• If the electrode potentials EAgCl/Ag
and EH+/H2
are known for the two
half-reactions, we may then find the cell potential Ecell by
subtracting the electrode potential for the second reaction from the
first.
• That is,
24
A more general statement of the last relationship is
where Eright
is the potential of the half-cell written on the right in the
diagram of the cell or in the shorthand representation, and Eleft
is the
electrode potential for the half-reaction written on the left.

Nature of Electrode Potentials:
• The potential of an electrochemical cell is the difference between the
potential of one of the electrodes and the potential of the other.
• This potential is a measure of an electrode’s electron energy.
• For a metallic conductor immersed in a solution of an electrolyte, all
excess charge resides on its surface, and it is possible to adjust the
charge density on this surface by adjusting the output of an external
power supply attached to the conductor.
• As this external source forces electrons onto the surface of an
electrode, the electrons become more crowded, and their energy
increases because of coulombic repulsion.
• The potential of the electrode thus becomes more negative.
• If the external circuitry withdraws enough electrons from the
electrode, the surface will have a positive charge, and the electrode
becomes more positive. 25

• It is also possible to vary the electron energy, and thus the potential,
of a metallic electrode by varying the composition of the solution
that surrounds it.
• For example, there is a potential at a platinum electrode immersed in
a solution containing hexacyanoferrate(II) and hexacyanoferrate(III)
as result of the equilibrium.
26
• If the concentration of Fe(CN)6
4-
is made much larger than that of
Fe(CN)6
3-
, hexacyanoferrate(II) has a tendency to donate electrons to
the metal, thus creating a negative charge at the surface of the
metal.
• Under this condition, the potential of the electrode is negative.

• On the other hand, if Fe(CN)6
3-
is present in large excess, there is a
tendency for electrons to be removed from the electrode, causing
surface layer ions to form in the solution and leaving a positive
charge on the surface of the electrode. The platinum electrode then
exhibits a positive potential.
• We must emphasize that no method can determine the absolute
value of the potential of a single electrode, because all voltage-
measuring devices determine only differences in potential.
• One conductor from such a device is connected to the electrode
under study.
• To measure a potential difference, however, the second conductor
must make contact with the electrolyte solution of the half-cell under
study.
• This second contact inevitably creates a solid-solution interface and
hence acts as a second half-cell in which a chemical reaction must
also take place if charge is to flow. 27

• A potential is associated with this second reaction. Thus, we
cannot measure the absolute value for the desired half-cell
potential.
• Instead, we can measure only the difference between the
potential of interest and the half-cell potential for the
contact between the voltage-measuring device and the
solution.
• Our inability to measure absolute potentials for half-cell
processes is not a serious problem because relative half-cell
potentials, measured versus a common reference electrode,
are just as useful.
• These relative potentials can be combined to give real cell
potentials. In addition, they can be used to calculate
equilibrium constants of oxidation-reduction processes. 28

The Standard Hydrogen Electrode :
• Hydrogen gas electrodes were widely used in early electrochemical
studies not only as reference electrodes but also as indicator
electrodes for determining pH.
• The half-cell shown on the left in Figure below shows the
components of a typical hydrogen electrode.
• The conductor is made of platinum foil that has been platinized.
• Platinum electrodes are platinized by coating their surfaces with a
finely divided layer of platinum by rapid chemical or electrochemical
reduction of H2
PtCl6
.
• The finely divided platinum on the surface of the electrode does not
reflect light as does polished platinum, so the electrode appears
black.
29

30

• Because of its appearance, the deposited platinum is called platinum
black. Platinum black has a very large surface area to ensure that the
reaction
31
is rapid at the electrode surface. The stream of hydrogen simply keeps
the solution adjacent to the electrode saturated with the gas.
• The hydrogen electrode may act as the positive or the negative
electrode, depending on the half-cell with which it is coupled with
the salt bridge shown in Figure above.
• When the cell is short circuited, hydrogen is oxidized to hydrogen ions
when the electrode is an anode; the reverse reaction takes place
when the electrode is a cathode.

• Under proper conditions, then, the hydrogen electrode is
electrochemically reversible.
• When the cell is connected as shown in Figure above, there is
essentially no current in the cell because of the very high impedance
of the meter.
• Under short circuit, the meter is replaced with a wire or
low-resistance load, and the reaction proceeds.
• The potential of a hydrogen electrode depends on the temperature,
the hydrogen ion activity in the solution, and the pressure of the
hydrogen at the surface of the electrode.
Practical Reference Electrodes :
• One of the most common of the Practical Reference Electrode is the
silver–silver chloride electrode.
32

• This electrode can be prepared by applying an oxidizing voltage to a
silver wire immersed in a dilute solution of hydrochloric acid.
• A thin coating of silver chloride forms that adheres tightly to the
wire.
• The wire is then immersed in a saturated solution of potassium
chloride.
• A salt bridge connects the potassium chloride solution to the
electrode system being studied.
• The potential of this electrode is about +0.22 V with respect to the
SHE(Standard Hydrogen Electrode ).
• The electrode half-reaction is
33
• A second widely used reference electrode is the saturated calomel
electrode (SCE), which consists of a pool of mercury in contact with a
solution that is saturated with mercury(I) chloride (calomel) as well as
potassium chloride.

• Platinum wire dipping in the mercury provides electrical contact to
the other conductor, and a salt bridge to the second electrolyte
completes the circuit.
• The potential of this reference is about 0.24 V positive with respect
to the SHE. The electrode reaction is
34
Definition of Electrode Potential
• An electrode potential is defined as the potential of a cell in which the
electrode under investigation is the right-hand electrode and the SHE
is the left-hand electrode
• Electrode potentials could more properly be called relative electrode
potentials
• This cell potential may be positive or negative depending on the
electron energy of the electrode under study.

• Thus, when this energy is greater than that of the SHE, the electrode
potential is negative; when the electron energy of the electrode in
question is less than that of the SHE, the electrode potential is
positive.
• The cell in Figure above illustrates the definition of the electrode
potential for the half-reaction.
35
• In this figure, the half-cell on the right consists of a strip of the metal
M in contact with a solution of M2+
.
• The half-cell on the left is a SHE.
• By definition, the potential E observed on the voltmeter is the
electrode potential for the M2+
/M couple.

• In this general example, we assume that the junction potentials
across the salt bridge are zero.
• If we further assume that the activity of M2+
in the solution is exactly
1.00, the potential is called the standard electrode potential for the
system and is given the symbol E0
.
• That is, the standard electrode potential for a half-reaction is the
electrode potential when the reactants and products are all at unit
activity.
• If M in the figure is copper, and if the copper ion activity in the
solution is 1.00, the compartment on the right is positive, and the
observed potential is +0.337 V.
• The spontaneous cell reaction, which would occur if the voltmeter
were replaced by a wire, is
36

• Because the hydrogen electrode is on the left, the measured
potential is, by definition, the electrode potential for the Cu2+
/Cu
half-cell.
• It is to be noted that the copper electrode is positive with respect to
the hydrogen electrode; thus, we write
37
• If M is Cadmium instead of copper, and the solution has a cadmium
ion activity of 1.00, the potential is observed to be -0.403 V .
• In this case, the cadmium electrode is negative, and the cell potential
has a negative sign. The spontaneous cell reaction would be
and we may write

• A zinc electrode in a solution of zinc ion at unity activity exhibits a
potential of -0.763 V when coupled with the SHE.
• The zinc electrode is the negative electrode in the galvanic cell, and
its electrode potential is also negative.
• The standard electrode potentials for the four half-cells just
described can be arranged in the order
38

• The magnitudes of these standard electrode potentials show the
relative strengths of the four ionic species as electron acceptors
(oxidizing agents).
• In other words, we may arrange the ions in order of their decreasing
strengths as oxidizing agents,
39
• Alternatively, the elements in order of their increasing strengths as
reducing agents,
Sign Conventions for Electrode Potentials:
• When we consider a normal chemical reaction, we speak of the
reaction occurring from reactants on the left side of the arrow to
products on the right side.

• By the International Union of Pure and Applied Chemistry (IUPAC)
sign convention, when we consider an electrochemical cell and its
resulting potential, we consider the cell reaction to occur in a certain
direction as well.
• This rule implies that we always measure the cell potential by
connecting the positive lead of the voltmeter to the right-hand
electrode in the schematic or cell drawing (for example, the Ag
electrode in the following example) and the common, or ground, lead
of the voltmeter to the left-hand electrode (the Cu electrode).
40
• By the International Union of Pure and Applied Chemistry (IUPAC)
sign convention, when we consider an electrochemical cell and its
resulting potential, we consider the cell reaction to occur in a certain
direction as well.

• That is, the direction of the overall process has Cu metal being
oxidized to Cu2+
in the left hand compartment and Ag+
being reduced
to Ag metal in the right-hand compartment.
• In other words, the reaction for the process occurring in the cell is
considered to be
41
Effect of Activity on Electrode Potential
Let us consider the half-reaction
where the capital letters represent formulas of reacting species
(whether charged or uncharged),
e-
represents the electron, and the lowercase italic letters indicate the
number of moles of each species (including electrons) participating in
the half-cell reaction.

• By invoking the same arguments that we used in the case of the
silver–silver chloride–SHE cell, we obtain
42
At room temperature (298 K), the collection of constants in front of the
logarithm has units of joules per coulomb (volts). Therefore,
When we convert from natural (ln) to base ten logarithms (log) by
multiplying by 2.303, the previous equation can be written as

• Equation (1) is a general statement of the Nernst equation, which
can be applied to both half-cell reactions and cell reactions.
43
The Standard Electrode Potential, E0
• The standard electrode potential is an important physical constant that
gives a quantitative description of the relative driving force for a half-
cell reaction.
• Keep in mind the following four facts regarding this constant.
1) The electrode potential is temperature dependent; if it is to have
significance, the temperature at which it is determined must be
specified.

(2) The standard electrode potential is a relative quantity in the sense
that it is really the potential of an electrochemical cell in which the
left electrode is a carefully specified reference electrode—the
SHE—whose potential is assigned a value of zero.
(3) The sign of a standard potential is identical with that of the
conductor in the right-hand electrode in a galvanic cell, with the left
electrode being the SHE.
(4) The standard potential is a measure of the driving force for a
half-reaction.
44
Some Limitations to the Use of Standard Electrode Potentials :
• Standard electrode potentials are of great importance in
understanding electroanalytical processes.
• There are certain inherent limitations to the use of these data

Substitution of Concentrations for Activities :
• As a matter of convenience, molar concentrations—rather than
activities—of reactive species are usually used in making calculations
with the Nernst equation.
• Unfortunately, these two quantities are identical only in dilute
solutions.
• With increasing electrolyte concentrations, potentials calculated
using molar concentrations are often quite different from those
obtained by experiment.
Effect of Other Equilibria:
• The application of standard electrode potentials is further
complicated by solvation, dissociation, association, and
complex-formation reactions involving the species of interest.
45

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