Electrochemistry
Plan
1. Electrode processes. Electrode potential.
2. Different types of electrodes.
3. Cell potential.
4. Galvanic cells. Cell potential or EMF.
5. The kinetics of electrochemistry
processes.
Assistant Kozachok S.S prepared
The changes in which electrical energy is produced as a result
of chemical change. The devices used to produce electrical
energy from chemical reactions are called electrical cells,
galvanic or voltic cells.
In these cells, oxidation and reduction reaction reactions occur
in separate containers called half cells and the red-ox reaction
is spontaneous.
The arrangement consists of two beakers, one of with contains
1,0 M solution of zinc sulphate and the other 1,0 M solution of
copper sulphate. A zinc rod is dipped into ZnSO4 solution
while a copper rod is dipped into CuSO4 solution. These
metallic rods are known as electrodes.
The metallic rods in the beaker are connected to the
ammeter by means of an insulated wire through a
key. Ammeter is used to know the passage of current
which moves in opposite direction to the flow of
electrons. The solution in the two beakers are
connected by an inverted U-tube containing
saturated solution of some electrolyte such as KСl,
KNO3, NH4OH which does not undergo a chemical
change during the process.
The two openings of the U-tube are plugged with
some porous material such as glass wood or cotton.
The U-tube which connects the two glass beakers is
called a salt-bridge.
When the circuit is completed by inserting the key in
the circuit, it is observed that electric current flows
through external circuit as indicated by the ammeter.
The following observations are made:
Therefore, the current flows from copper to zinc
N.B. The flow of electric current is taken opposite to
the flows of electrons
There observation can be explained as:
During the reaction, zinc is oxidized to Zn2+ ions
which go into the solution. Therefore,
the zinc rod gradually loses its weight. The electrons
released at the zinc electrode move towards the other
electrode through outer circuit. Here, these are
accepted by Cu2+ ions of CuSO4 solution which are
reduced to copper.
The zinc electrode where electrons are released or
oxidation occurs s called anode while the copper
electrode where electrons are accepted or reduction
occurs is called cathode.
The two containers involving oxidation and reduction
half reactions are called half cells. The zinc rod
dipping into a ZnSO4 solution is oxidation half cell
and the copper electrode dipping into a CuSO4
solution is reduction half cell
N.B. The galvanic cells which consists of the zinc rod dipping
into a ZnSO4 solution and the copper electrode dipping
into a CuSO4 solution is Daniell cell.
Its formula is:
Salt bridge and its function. It’s usually an inverted U-tube
filled with concentrated solution of inert electrolyte. The
essential requirements of electrolyte are:
a) The mobility of the anion and cation of the electrolyte
should be almost same.
b) The ions of the electrolyte are not involved in
electrochemical change.
c) The ions do not react chemically with the species of the
cell.
Generally, salts like KCl, KNO3, etc. are used. The seturated
solutions of these electrolytes are prepared in agar agar
jelly or gelatin. The jelly keeps the electrolyte in semisolid phase and thus prevents mixing.
The important functions of the salt bridge are:
a) Salt bridge completes the electrical circuit.
b) Salt bridge maintains electrical neutrality of two half cell
solution.
The accumulation of charges in the two half cells
(accumulation of extra positive charge in the solution
around the anode according to the realizing of Zn2+ in
excess and accumulation of extra negative charge in the
solution around the catode due to excess of SO42- ) is
prevented by using salt bridge, which provides a passage
for the flow of the charge in the internal circuit.
REPRESENTATION OF AN ELECTROCHEMICAL CELL
An electrochemical cells or galvanic cell consists of two
electrodes: anode and cathode. The electrolyte solution
containing these electrodes are called half cells.
The following conventions are used in representing an
electrochemical cell:
1. A galvanic cell is represented by writing the anode (where
oxidation occurs) on the left hand side and cathode (where
reduction occurs) on the right hand side.
2. The anode of the cell is represented by writing metal or
solid phase first and then the electrolyte (or the cation of
the electrolyte) while the cathode is represented by writing
the electrolyte first and then metal or solid phase.
3. The salt bridge which separates the two half cells is
indicated by two vertical lines.
Electrode Potential and E.M.F. of a galvanic cell
Electrode Potential
The flow of electric current in an electrochemical cell
indicates that a potential difference exists between two
electrodes.
If the metal has relatively high tendency to get oxidised, its
atom will lose electrons readily and form Cu2+ ions, which
go into the solution. The electrons lost on the electrode
would be accumulated on the metal electrode and the
electrode acquires a slight negative charge with respect to
the solution. Some of the Cu2+ ions from the solution will
take up electrons and become Cu atoms. After some time,
an equilibrium will be established as:
When such an equilibrium is attained, it results in separation
of charges (negative on the electrode with respect to the
solution).
Forming the double layer
Cu
1- metal
2 - solution
Similarly, if the metal ions have relatively greater
tendency to get reduced, they will take electrons
from the electrode. As a result, a net positive
charge will be developed on the electrode with
respect to the solution. This will also result into
separation of charges (positive on the electrode
with respect to the solution).
Due to separation of charges between the electrode
and the solution, an electrical potential is set up
between metal electrode and its solution.
The electrical potential difference set up between the
metal and its solution is known as electrode
potential.
The mechanism of the double layer forming
For example of the copper electrode is dipped into CuSO4
solution.
The chemical potential of copper’s ions in the metal and
in the solution is not equal. The chemical potential of
copper’s ions in the metal at the given temperature is
stable value, the chemical potential of copper’s ions in
the solution depends on the solution’s concentration.
If the at the given concentration of solution the chemical
potential of copper’s ions in the solution is greater than
the chemical potential of these ions in the metal . Then
at the dipping of the metal in the solution some quantity
of Cu2+ ions are hydrated and transferred on the metal
according to that positive charge forming on the metal’s
surface. Sulfate’s anions are attracted to metal’s surface,
they courses the negative charge. These processes cause
the double electrical layer and related with it the
potential difference.
Forming the double layer
Cu
+
-
+
+
+
+
+
+
-
( -)
- - - --CuSO4
The electrode potential may be of two types:
1. Oxidation potential:
The tendency of an electrolyte to lose electrons
or to get oxidised
2. Reduction potential. The tendency of an
electrode to gain electrons or to get reduced.
E.M.F. or Cell Potential of a Cell
The difference between the electrode potentials of the two
electrodes constituting an electrochemical cell is known as
electromotive force (e.m.f.) or cell potential of a cell. This
acts as a driving force for the cell reaction. The potential
difference is expressed in volts.
Therefore, the cell potential or e.m.f. arises from the difference
in the tendencies of the two ions to get reduced.
It is equal to the reduction potential for the substance that
actually undergoes reduction minus the reduction
potential of the substance that undergoes oxidation.
Thus, e.m.f. of a cell may be defined as the potential
difference between two electrodes of the cell when either no
or negligible current is allowed to flow in the circuit.
Standard electrode potential
Since a half cell in an electrochemical cell can work only in
combination with the other half cell and does not work
independently, it is not possible to determine the absolute
electrode potential of an electrode. We can, therefore, find
only the relative electrode potential.
This difficulty can be solved by selecting one of the electrodes
as a reference electrode and arbitrarily fixing the potential
of this electrode as zero. For this purpose, reversible
hydrogen electrode has been universally accepted as a
reference electrode. It is called standard hydrogen
electrode (S.H.E) or normal hydrogen electrode (N.H.E.)
Standard hydrogen electrode. It consists of platinum
wire sealed in a glass tube and has a platinum foil
attached to it. The foil is coated with finely divided
platinum and acts as platinum electrode. It is dipped
into an acid solution containing H+ ions in 1 M
concentration (1M HCl). Pure hydrogen gas at 1
atmospheric pressure is constantly bubbled into
solution at constant temperature of 298K. The
surface of the foil acts as a site for the reaction.
The electrode potential of an electrode can be
determined by connecting this half cell with a
standard hydrogen electrode. The electrode
potential of the standard hydrogen electrode is taken
as zero.
The electrode potential of a metal electrode as
determined with respect to a standard or normal
hydrogen electrode is called standard electrode
potential (E0). Standard electrode potentials are
always associated with the reduction occurring at the
electrodes.
In this case, the electrons flow from zinc electrode to
hydrogen electrode and therefore, the zinc electrode acts as
anode and S.H.E. acts as a cathode. The cell may be
represented as:
2. Measurement of Electrode Potential of Cu2+|Cu
Electrode
In this case, the hydrogen has greater tendency to lose
electrons. Therefore, oxidation occurs at hydrogen electrode
and reduction occurs at copper electrode.
It may be noted that it is not always convenient to use
standard hydrogen electrode as reference electrode
because of experimental difficulties in its
preparation and use.
N.B. The standard electrode potentials given in the
following table are measured in their standard states
when the concentration of the electrolyte solutions are
fixed as 1M and temperature is 298 K.
Application of the Electrochemical
(electromotive) series:
1. Relative strengths of oxidising and reducing
agents.
The substances which have lower reduction potentials
are stronger reducing agents while have higher
reduction potentials are stronger oxidations agent.
2. Calculation of the E.M.F. of the cell.
E0cell = E0(cathode) - E0(anode)
3. Predicting feasibility of the reaction.
In general, a red-ox reaction is feasible only if the
species to release electrons must have lower
reduction potential as compared to the species
which is to accept electrolytes.
4. To predict whether a metal can liberate
hydrogen from acid or not.
In general, only those metals can liberate hydrogen
from the acid which have negative values of
reduction potentials , - E0 values.
Dependence of electrode and cell potentials on
concentration: Nernst equation
The electrode potentials depend on the concentration of the
electrolyte solutions.
or
E  E0 
0.059
log [ Mn  (aq) ]
n
Nerst equation
ЕOx / Re d  Е
0
Ox / Re d
RT

ln a
nF
RT
2.3
 0.059, that
F
ЕOx / Re d  Е
0
Ox / Re d
0.059

lg a
n
In general, for an electrochemical cell reaction:
ne
aA  bB  cC  dD
The Nerst equation may be written as:
The value of a, b, c, d and n are obtained from the balanced
cell reactions.
N.B. It must be remembered that while writing the Nerst
equation for the overall cell reaction, the log term is the
same as the expression for the equilibrium constant for the
reaction. However, some books use the expression in the
reverse form as the expression for the equilibrium constant
but, sign after E0 is changed.
Equilibrium constant from Nernst equation
The e.m.f. of the cell may be used to calculate the
equilibrium constant for the cell reaction. At
equilibrium, the electrode potentials of the two
electrodes become equal so that e.m.f. of the cell is
zero. Consider the following redox reaction:
Significance of Kc. The value of Kc gives the extent of
the cell reaction. If the value of Kc is large, the
reaction proceeds to large extent.
Electrochemical cell and free energy
In electrochemical cells, the chemical energy is
converted into electrical energy. The cell potential is
related to free energy change. In an electrochemical
cell, the system does work by transferring electrical
energy through an electric circuit.
Where
is the standard free energy for the reaction.
Significance. The above equation helps us to predict the
feasibility of the cell reaction. For a cell reaction to be
spontaneous,
must be negative. This means that E must
be positive for a spontaneous cell reaction.
SOME COMMERCIAL CELLS
One of the main uses of galvanic cells is the
generation of portable electrical energy. These
cells are also popularly known as batteries. The
term battery is generally used for two or more
galvanic cells connected in series. Thus, a battery
is an arrangement of electrochemical cells used as
an energy source. The basis of an electrochemical
cell is an oxidation-reduction reaction.
Types of commercial cells.
1) Primary cells
2) Secondary cells
Primary cells. In these cells, the electrode reactions
cannot be reversed by an external electric energy
source. In these cells, reactions occur only once and
after use thaey become dead. Therefore, they are not
chargeable. Examples are dry cell, mercury cell.
Secondary cells (storage cells or accumulators). In
the secondary cells, the reaction can be reversed by
an external electric source. Therefore, these cells can
be recharged by passing electric current and used
again and again. Examples are lead storage battery
and nickel-cadmium storage cell.
The most popular example is of lead storage cell
which is used in automobiles.
Each battery consists of a number of voltaic cells connected in
series. Three to six such cells are generally combined to get
6 to 12 volt battery. In each cell, the anode is a grind of lead
packed with divided spongy lead and the cathode is a grind
of lead packed with PbO2.
The electrolyte is aqueous solution of sulfuric acid (38% by
mass) having a density 1,30 g ml-1 sulfuric acid. When the
lead plates are kept for sometimes, a deposit of lead
sulphate is formed on them.
At the anode, lead is oxidised to Pb2+ ions and insoluble
PbSO4 is formed. At the cathode PbO2 is reduced to to Pb2+
ions and PbSO4 is formed.
The following reactions take place in the lead storage cell:
During the working of the cell, PbSO4 is formed at each
electrode and sulphuric acid is used up. As a result, the
concentration of H2SO4 decreases and the density of the
solution also decreases. When the density of H2SO4 falls
below 1.2 g ml-1, the battery needs recharge.
Recharge the Battery
The cell can be charge by passing electric current of a suitable
voltage in the opposite direction. The electrode reaction gets
reversed. As a result, the flow of electrons gets reversed and
lead is deposited on anode and PbO2 on the cathode. The
density of sulphuric acid also increases. The reaction can be
written as:
The most important types of electrodes are:
1. The first reference electrode: Metal-metal ion
electrodes and gas-ion electrodes
2. The second reference electrode: Metal-insoluble
salt-anion electrodes
3. The third reference electrode: inert "oxidationreduction" electrodes
4. Membrane electrodes
The metal - metal ion electrode consists of а metal in contact with
its ions in solution.
An example is а piece of silver metal immersed in а solution of
silver nitrate. The diagram for such an electrode serving as а
cathode (it would appear at the right in а cell diagram) is:
Ag+(aq)  Ag(s)
and the cathode half-reaction is: Ag+(aq) + e-Ag(s)
in which the electrons соmе from the external circuit. When this
electrode serves as an anode, it is diagramed as: Ag(s)  Ag+(aq)
(as it would appear at the left in а cell diagram), and its halfreaction equation is:
Ag(s)  Ag+(aq) + еIn general the first reference electrons can be represented as:
Mz+/M. The half reduction reaction is:
Mz+ + ze  M
Following convention the half reaction that occurs on the electrode
is written as a reduction reaction
Nernst equation for these type electrodes is:
0.059
z
E E 
log [ Mn (aq) ]
n
0
0.059
E E 
log a Mn z
n
0
0.059
[Mn(s) ]
E E 
log
n
[Mn z  (aq) ]
0
a Mn(s)
0.059
E  E 
log
n
a Mn z 
0
The gas-ion electrode (Standard hydrogen electrode)
Hydrogen electrode that works at the following conditions:
a H = 1, p = 101,3 kPa, T = 250C = 298K is called standard.
H+

Electrochemical potential of this electrode depends on
the hydrogen ions concentration. However the
standard potential of this electrode equals o and the
valency of hydrogen equals 1 (n=1) Nernst equation
is:
Ecell = 0,059 log [H+] = - 0,059 pH
Measuring of pH to use potentiometric method of the
determination of hydrogen ions concentration.
This method is based on the measuring of e.m.f of
the cell which consists of the reduction electrode
(calomel) and the electrode which has dependence
on the hydrogen ions concentration (gas-ion
electrode, glass electrode, quinonhydrone electrode
2. In the metal-insoluble salt-anion electrode, а metal is in
contact with one of its insoluble salts and also with а
solution containing the anion of the salt.
0.059
E E 
log a A z
n
0
An example is the so-called silver - silver chloride electrode,
written as а cathode as:
Cl-(aq)  AgCl(s)  Ag(s)
for which the cathode half-reaction is:
AgCl (s) + е-  Ag(s) + Cl- (aq)
EAg,AgCl | Cl- = E0Ag+/AgCl - 0.059 lg aClE0Ag+/AgCl = 0,2224
Silver - silver chloride electrode
Ag, is covered by the layer of nonsoluble AgCl
КCl solution
KCl, AgCl | Ag
Calomel electrode consists of the mercury and
calomel past that is dipped in potassium chloride
solution. It is often used as a reference electrode to
determine the standard electrode potential ( more
often than hydrogen electrode). Its scheme is:
Cl-|Hg2Cl2, Hg
The half reaction is:
Hg2Cl2 + 2e  2Hg + 2ClEcell = E0 - 0.059 lg aClAs a rule to use the calomel electrodes that contain
0,1 M, 1 M and saturated solution of potassium
chloride. Their standard potential at 298K equal;
0,337; 0,2801; 0,2412 V.
3. An inert oxidation-reduction electrode
It consists of а strip, wire, or rod of an inert materiel (Pl, Au,
Ir…) in contact with а solution, which contains ions of а
substance is two different oxidation states (oxidation and
reduction form). The difference between general metal
electrode and ox-red electrode is that ox/red electrode does
not take place in ox-red reaction which exist in solution but is
the electrons conductor. For example: Pt| Sn2+, Sn4+ or Pt|
Fe2+, Fe3+
a
0.059
E  E0 
log Red
n
a Ox
Ox +ze → Red
There are two types of ox-red electrodes:
1.Simple: Fe2+, Fe3+ | Pt Fe3+ + e → Fe2+
a Fe 2
0.059
0
E E 
log
n
a Fe3
Scheme of ox-red electrode
(the third reference electrode)
Pt
Fe3+ , Fe2+ | Pt
Fe3+ + e = Fe2+
FeCl3 + FeCl2
EFe3 / Fe2  E
0
Fe3 / Fe2
 0.059 lg
aFe3
aFe2
2. Complex ox-red electrode there is changing the charge and
the composition of the ions
Mn2+, MnO4-, H+ | Pt MnO4- + 8H++5e→Mn2++4H2O
Example is quinonhydrone electrode.
It is prepared by the platinum strip or wire which is contained
in the glass tube. The electrode is dipped in the solution
with unknown pH that is needed to determine and to add
some quinonhydrone’s crystalls in this solution.
Quinonhydrone is a crystalline product which consists of
quinone (benzoquinone) С6Н4О2 and hydroquinone
C6H4(OH)2. It is less solubility in water and decomposes
into quinone and hydroquinone in the solution. In the
saturated solution equal molar mixture of quinone and
hydroquinone is formed.
Quinonhydrone electrode
(the third reference electrode)
С6Н4О2, С6Н4(ОН)2, H+ | Pt
С6Н4О2 + 2Н+ + 2е = С6Н4(ОН)2
Ecell
aC6 H 4 (OH )2
0.059
E 
lg
2
аC6 H 4O2  а 2 Н 
0
Including that the activity of
quinone and hydroquinone is equal
in the seturated solution, we have
Ecell  E  0.059 lg aH 
0
Quinonhydrone
The scheme of quinonhydrone cell
with one electrolyte
Pt, Н2 | quinhydr, H+ | KCl | KCl,Hg2Cl2| Hg
Ecell = E quinhydr - Ecalomel
The glass electrode for pH measurements. The cell consists of
a glass indicator electrode and а saturated calomel reference
electrode, both immersed in the solution whose pH is to be
determined. The indicator electrode consists of а thin, рНsensitive glass membrane sealed onto one end of а heavywalled glass or plastic tube. А small volume of dilute
hydrochloric acid saturated with silver chloride is contained in
the tube (in some electrodes this solution is а buffer containing
chloride ion). A silver wire in this solution forms а silver/silver
chloride reference electrode, which is connected to one of the
terminals of а potential-measuring device. The calomel
electrode is connected to the other terminal.
Show that the system contains two reference electrodes: (1) the
external – calomel electrode and (2) the internal silver/silver
chloride electrode. Although the internal reference electrode is
part of the glass electrode, it is not the pH-sensing element.
Instead, it is the thin glass membrane at the tip of the electrode
that responds to pH.
The make the cell with glass and calomel electrodes
and measuring its e.m.f can be determined pH of
solution.
E glass  const  0.059 lg a Na 
glass electrode’s constant which is depended on the
electrode nature. The constant is fined according to the
graph which is plotted between the Ecell and pH
ordinates.
The mechanism of the diffusion
potential
HCl 1 н
H+
ClEcell
uKt   u An 0.059 a1

lg
uKt   u An n
a2
where a1 > a2
HCl 0.1 н
-
+
+
+
+
+
+
+
+
+
The scheme of concentrated cell
_
+
(+)Ag
| AgNO3 || AgNO3 | Ag(-)
C
1
>
C
2
Ecell = E2 Ag+/Ag – E1 Ag+/Ag =
= 0.059 lg (a2 / a1)
NO3AgNO3, C2
AgNO3, C1
The electric circuit with transfer and without it
The electric circuit without transfer is when the
electrodes are dipped in one solution.
Examples: Pt|H2|HCl|AgCl, Ag – hydrogen-silversilver chloride electrode;
Pb | PbSO4 || H2SO4(aq) | PbO2 | Pb – the lead storage
battery
The electric circuit with transfer is when the
electrodes are dipped in different solution which
contact with each other.
Examples: Zn|ZnSO4||CuSO4|Cu – Daniell cell
Pt|H2|HCl||KCl|Hg2Cl2,Hg – hydrogen-calomel cell