Electrochemistry
What are half reactions?
Let us consider the reaction 2Na + Cl2 → 2Na+ + 2Cl–. It occurs by the transfer of electrons from
Na to Cl. Na loses an electron and is said to be oxidized to Na+ ion. At the same time,
Cl gains an electron and is reduced to Cl– ion. Such a reaction which is brought about by loss of
electrons
(oxidation)
and
gain
of
electrons
(reduction)
simultaneously
is
called
an
Oxidation-Reduction reaction or Redox reaction in brief. It may be noted that in the overall redox
reaction no free electrons are generated.
The redox reaction can be considered as made up of two reactions. For example, the redox
reaction 2Na + Cl2 → Na+ + 2Cl– is composed of two half-reactions:
2Na → 2Na+ + 2e– (oxidation)
Cl2 + 2e– → 2Cl– (reduction)
Each of the two reactions shows just its oxidation or just the reduction portion of the overall
redox reaction. Being half components of the redox reaction, these reactions are called Half-reactions.
The first half-reaction that proceeds by oxidation is often referred to as the Oxidation half-reaction.
The second half-reaction that occurs by reduction, is referred to as the Reduction half-reaction.
When the two half-reactions are added together, the sum is the net redox reaction.
Electrochemical cells
A device for producing an electrical current from a chemical reaction (redox reaction) is called
an electrochemical cell. How a Redox reaction can produce an electrical current? When a bar of zinc is
dipped in a solution of copper sulphate, copper metal is deposited on the bar (See Figure 1).
The net reaction is Zn + Cu2+ ⎯⎯→ Zn2+ + Cu
Fig. 1: Redox reaction taking place on zinc bar itself
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This is a redox reaction and the two half-reactions are:
Zn ⎯⎯→ Zn2+ + 2e–
Cu2+ + 2e– ⎯⎯→ Cu
In this change, Zn is oxidized to give Zn2+ ions and Cu2+ ions are reduced to Cu atoms. The
electrons released in the first half-reaction are used up by the second half-reaction. Both the half
reactions occur on the zinc bar itself and there is no net charge.
Now, let the two half-reactions occur in separate compartments which are connected by a wire
(See Fig. 2) The electrons produced in the left compartment flow through the wire to the other
compartment. However the current will flow for an instant and then stop. The current stops flowing
because of the charge build up in the two compartments. The electrons leave the left compartment and
it would become positively charged. The right compartment receives electrons and becomes negatively
charged. Both these factors oppose the flow of electrons (electrical current) which eventually stops.
Fig. 2: Separate half-reactions cause flow of
electrons (Current) in the wire connecting them
This problem can be solved very simply. The solutions in the two compartments may be
connected, say, by a salt bridge. The salt bridge is a U-tube filled with an electrolyte such as NaCl,
KCl, or K2SO4. It provides a passage to ions from one compartment to the other compartment without
extensive mixing of the two solutions. With this ion flow, the circuit is complete and electrons pass
freely through the wire to keep the net charge zero in the two compartments.
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Voltaic Cells
A Voltaic cell, also known as a galvanic cell is one in which electrical current is generated by a
spontaneous redox reaction. A simple voltaic cell is shown in Fig. 3. Here the spontaneous reaction of
zinc metal with an aqueous solution of copper sulphate is used.
Zn + Cu2+ ⎯⎯→ Zn2+ + Cu
A bar of zinc metal (anode) is placed in zinc sulphate solution in the left container. A bar of
copper (cathode) is immersed in copper sulphate solution in the right container. The zinc and copper
electrodes are joined by a copper wire. A salt bridge containing potassium sulphate solution
interconnects the solutions in the anode compartment and the cathode compartment.
The oxidation half-reaction occurs in the anode compartment. Zn ⎯⎯→ Zn2+ + 2e–
The reduction half-reaction takes place in the cathode compartment. Cu2+ + 2e– ⎯⎯→ Cu
Fig. 3: A simple voltaic (galvanic) cell
When the cell is set up, electrons flow from zinc electrode through the wire to the copper
cathode. As a result, zinc dissolves in the anode solution to form Zn2+ ions. The Cu2+ ions in the
cathode half-cell pick up electrons and are converted to Cu atoms on the cathode. At the same time,
SO42– ions from the cathode half-cell migrate to the anode half-cell through the salt bridge.
Likewise, Zn2+ ions from the anode half-cell move into the cathode half-cell. This flow of ions from
one half-cell to the other completes the electrical circuit which ensures continuous supply of current.
The cell will operate till either the zinc metal or copper ion is completely used up.
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Current is the flow of electrons through a wire or any conductor.
Electrode is the material: a metallic rod/bar/strip which conducts electrons into and out of a
solution.
Anode is the electrode at which oxidation occurs. It sends electrons into the outer circuit. It has
negative charge and is shown as (–) in cell diagrams.
Cathode is the electrode at which electrons are received from the outer circuit. It has a positive
charge and is shown as (+) in cell diagrams.
Electrolyte is the salt solutions in a cell.
Anode compartment is the compartment of the cell in which oxidation half-reaction occurs.
It contains the anode.
Cathode compartment is the compartment of the cell in which reduction half-reaction occurs.
It contains the cathode.
Half-cell. Each half of an electrochemical cell, where oxidation occurs and the half where
reduction occurs, is called the half cell.
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Daniel Cell
It is a typical voltaic cell. It was named after the British chemist John Daniel. It is a simple zinc
copper cell like the one described above. In this cell the salt-bridge has been replaced by a porous pot.
Daniel cell resembles the above voltaic cell in all details except that Zn2+ ions and SO42– ions flow to
the cathode and the anode respectively through the porous pot instead of through the salt-bridge.
Inspite of this difference, the cell diagram remains the same.
Cell reaction
The flow of electrons from one electrode to the other in an electrochemical cell is caused by the
half-reactions taking place in the anode and cathode compartments. The net chemical change obtained
by adding the two half-reactions is called the cell reaction. Thus, for a simple voltaic cell described
above, we have
(a) Half-reactions:
Zn(s) ⎯⎯→ Zn2+ (aq) + 2e–
Cu2+ (aq) + 2e– ⎯⎯→ Cu(s)
(b) Cell reaction by adding up the half-reactions:
Zn(s) + Cu2+ (aq) ⎯⎯→ Zn2+ (aq) + Cu(s)
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Cell potential or emf
In a Zn-Cu voltaic cell, electrons are released at the anode and it becomes negatively charged.
The negative electrode pushes electrons through the external circuit by electrical repulsions. The copper
electrode gets positive charge due to the discharge of Cu2+ ions on it. Thus electrons from the outer
circuit are attracted into this electrode. The flow of current through the circuit is determined by the
‘push’, of electrons at the anode and ‘attraction’ of electrons at the cathode. These two forces constitute
the ‘driving force’ or ‘electrical pressure’ that sends electrons through the circuit. This driving force is
called the electromotive force (abbreviated emf) or cell potential. The emf of cell potential is measured
in units of volts (V) and is also referred to as cell voltage.
Cell diagram or Representation of a Cell
A cell diagram is an abbreviated symbolic depiction of an electrochemical cell. For this purpose,
we will consider that a cell consists of two half-cells. Each half-cell is again made of a metal electrode
contact with metal ions in solution.
(1) A single vertical line
(|) represents a phase boundary between metal electrode and ion solution
(electrolyte). Thus the two half-cells in a voltaic cell are indicated as
It may be noted that the metal electrode in anode half-cell is on the left, while in cathode halfcell it is on the right of the metal ion.
(2) A double vertical line
(||) represents the salt bridge, porous partition or any other means of
permitting ion flow while preventing the electrolyte from mixing.
(3) Anode half-cell is written on the left and cathode half-cell on the right.
(4) In the complete cell diagram, the two half-cells are separated by a double vertical line
(salt bridge) in between. The zinc-copper cell can now be written as
(5) The symbol for an inert electrode, like the platinum electrode is often enclosed in a bracket.
For example,
6
(6) The value of emf of a cell is written on the right of the cell diagram. Thus a zinc-copper cell has
emf 1.1V and is represented as
Convention regarding sign of emf value
The magnitude of the emf of a cell reflects the tendency of electrons to flow externally from one
electrode to another. The electrons are transported through the cell solution by ions present and pass
from the positive electrode (Cu in case of Daniel cell) to the negative electrode. This corresponds to a
clockwise flow of electrons through the external circuit. Thus the emf of the cell is given the +ve sign.
The negative sign indicates that the cell is not feasible in the given direction. The reaction will take
place in the reverse direction.
Calculating the emf of a cell
The emf of a cell can be calculated from the half-cell potentials of the two cells (anode and
cathode) by using the following formula
where ER and EL are the reduction potentials of the right-hand and left-hand electrodes
respectively. It may be noted that absolute values of these reduction potentials cannot be determined.
These are found by connecting the half-cell with a standard hydrogen electrode whose reduction
potential has been arbitrarily fixed as zero.
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Measurement of emf of a cell [Poggendorf’s compensation method]
The emf of an unknown cell can be measured with the help of a potentiometer (Fig. 4).
It consists of a wire AB which is about one metre long. The two ends of this wire are connected to a
working battery W. A standard cell C1 (i.e., a cell of known emf) is connected to the end A.
At the other end, the cell C1 is connected to a galvanometer through a key K1. The galvanometer is then
joined to a sliding contact that moves on the wire AB. The cell C2 whose emf is to be measured is
similarly connected to the key K2, the galvanometer and then the sliding contact. By using the key K1,
the cell C1 is put into the circuit and the contact is moved to and fro along AB. When no current flows
through the galvanometer, the point of contact X1 is recorded. Then by using the key K2, the cell C2 is
put into the circuit and the procedure is repeated to find the corresponding point X2. The emf of the cell
C2 is calculated by using the following equation:
Fig. 4: Measuring the emf of a cell with a potentiometer
Potential differences at interfaces
The transition region between two phases consists of a region of charge unbalance known as the
electric double layer [Helmholtz double layer]. As its name implies, this consists of an inner
monomolecular layer of adsorbed water molecules and ions, and an outer diffuse region that
compensates for any local charge unbalance that gradually merges into the completely random
arrangement of the bulk solution. In the case of a metal immersed in pure water, the electron fluid
within the metal causes the polar water molecules to adsorb to the surface and orient themselves so as
to create two thin planes of positive and negative charge. If the water contains dissolved ions, some of
the larger (and more polarizable) anions will loosely bond (chemisorb) to the metal, creating a negative
inner layer which is compensated by an excess of cations in the outer layer.
8
Electrochemistry is the study of reactions in which charged particles (ions or electrons) cross the
interface between two phases of matter, typically a metallic phase (the electrode) and a conductive
solution, or electrolyte. A process of this kind can always be represented as a chemical reaction and is
known generally as an electrode process. Electrode processes take place within the double layer and
produce a slight unbalance in the electric charges of the electrode and the solution. Much of the
importance of electrochemistry lies in the ways that these potential differences can be related to the
thermodynamics and kinetics of electrode reactions. In particular, manipulation of the interfacial
potential difference affords an important way of exerting external control on an electrode reaction.
The interfacial potential differences which develop in electrode-solution systems are limited to
only a few volts at most. This may not seem like very much until you consider that this potential
difference spans a very small distance. In the case of an electrode immersed in a solution, this distance
corresponds to the thin layer of water molecules and ions that attach themselves to the electrode
surface— normally only a few atomic diameters. Thus a very small voltage can produce a very large
potential gradient.
Interfacial potentials are not confined to metallic electrodes immersed in solutions; they can in
fact exist between any two phases in contact, even in the absence of chemical reactions. In many forms
of matter, they are the result of adsorption or ordered alignment of molecules caused by non-uniform
forces in the interfacial region. Thus colloidal particles in aqueous suspensions selectively adsorb a
given kind of ion, positive for some colloids, and negative for others. The resulting net electric charge
prevents the particles from coming together and coalescing, which they would otherwise tend to do
under the influence of ordinary van der Waals attractions.
Fig. 5: Electric double layer at an electrode surface
9
Interfacial potential differences are not directly observable. The usual way of measuring a
potential difference between two points is to bring the two leads of a voltmeter into contact with them.
It’s simple enough to touch one lead of the meter to a metallic electrode, but there is no way you can
connect the other lead to the solution side of the interfacial region without introducing a second
electrode with its own interfacial potential, so you would be measuring the sum of two potential
differences. Thus single electrode potentials, as they are commonly known, are not directly observable.
Single electrode potential
An electrochemical cell consists of two half-cells. With an open-circuit, the metal electrode
in each half-cell transfers its ions into solution. Thus an individual electrode develops a potential
with respect to the solution. The potential of a single electrode in a half-cell is called the
Single electrode potential. Thus in a Daniel cell in which the electrodes are not connected externally,
the anode Zn/Zn2+ develops a negative charge and the cathode Cu/Cu2+, a positive charge.
The amount of the charge produced on individual electrode determines its single electrode potential.
The single electrode potential of a half-cell depends on: (a) concentration of ions in solution; (b)
tendency to form ions; and (c) temperature.
Standard emf of a cell [E°]
The emf generated by an electrochemical cell is given by the symbol E. It can be measured with
the help of a potentiometer. The value of emf varies with the concentration of the reactants
and products in the cell solutions and the temperature of the cell. When the emf of a cell is determined
under standard conditions, it is called the standard emf. The standard conditions are (a) 1 M solutions
of reactants and products; and (b) temperature of 25°C. Thus standard emf may be defined as:
the emf of a cell with 1 M solutions of reactants and products in solution measured at 25°C.
Standard emf of a cell is represented by the symbol E°. With gases 1 atm pressure is a standard
condition instead of concentration. For a simple Zn-Cu voltaic cell, the standard emf, E°, is 1.10 V.
This means that the emf of the cell operated with [Cu2+] and [Zn2+] both at 1 M and 25°C is 1.10 V.
Determination of emf of a half-cell
By a single electrode potential, we also mean the emf of an isolated half-cell or its half-reaction.
The emf of a cell that is made of two half-cells can be determined by connecting them to a voltmeter.
However, there is no way of measuring the emf of a single half-cell directly. A convenient procedure to
do so is to combine the given half-cell with another standard half-cell. The emf of the newly
constructed cell, E, is determined with a voltmeter. The emf of the unknown half-cell, E° can then be
calculated from the expression
10
If the standard half-cell acts as anode, the equation becomes
On the other hand, if standard half-cell is cathode, the equation takes the form
The standard hydrogen half-cell or Standard Hydrogen Electrode (SHE) is selected for
coupling with the unknown half-cell. It consists of a platinum electrode immersed in a 1 M solution of
H+ ions maintained at 25°C. Hydrogen gas at one atmosphere enters the glass hood and bubbles over
the platinum electrode. The hydrogen gas at the platinum electrode passes into solution,
forming H+ ions and electrons. H2 ⎯⎯→ 2H+ + 2e–
Fig. 6: Standard Hydrogen Electrode (SHE)
The emf of the standard hydrogen electrode is arbitrarily assigned the value of zero volts.
So, SHE can be used as a standard for other electrodes. The half-cell, whose potential is desired, is
combined with the hydrogen electrode and the emf of the complete cell determined with a voltmeter.
The emf of the cell is the emf of the half-cell.
For example, it is desired to determine the emf of the zinc electrode, Zn | Zn2+. It is connected
with the SHE as shown in Fig. 29.10. The complete electrochemical cell may be represented as:
11
Fig. 7: The zinc electrode (Zn, Zn2+) coupled with hydrogen electrode
The emf of the cell has been found to be – 0.76V which is the emf of the zinc half-cell.
Similarly, the emf of the copper electrode, Cu2+ | Cu can be determined by pairing it with the
SHE when the electrochemical cell can be represented as:
The emf of this cell has been determined to be 0.34 V which is the emf of the copper half-cell.
The two situations are shown in Fig. 8
Fig. 8: SHE can act both as cathode and anode when joined with another half-cell
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When it is placed on the right-hand side of the Zinc electrode, the hydrogen electrode reaction is
2H+ + 2e– ⎯⎯→ H2. The electrons flow to the SHE and it acts as the cathode. When the SHE is placed
on the left hand side, the electrode reaction is H2 ⎯⎯→ 2H+ + 2e–. The electrons flow to the copper
electrode and the hydrogen electrode as the anode. Evidently, the SHE can act both as anode and
cathode and, therefore can be used to determine the emf of any other half-cell electrode (or single
electrode).
IUPAC convention places the SHE on the left-hand side
In the procedure for determining the emf of a given half-cell, the standard hydrogen electrode
can be placed on the left-hand or the right-hand. In the convention adopted by the IUPAC (International
Union of Pure and Applied Chemistry), the SHE is always placed on the left-hand side of the half-cell
under study. The electrons flow from left-to-right and the given half-cell electrode gains electrons
(reduction). The observed emf of the combined electrochemical cell is then the emf of the half-cell on
the
right-hand. Such
emf values of half-cells,
or half reactions,
are known
as the
Standard reduction potentials or Standard potentials. However, if the SHE be placed on the right
hand side of the given half-cell, the potential so obtained is called as the Standard oxidation potential.
The latter potentials are the standard potentials with the sign reversed, the only difference being that
cells have been turned around.
According to IUPAC convention, the standard reduction potentials alone are the standard
potentials. The values of the standard potentials at 25°C (298 K) for some common Reduction Half
reactions are listed in Table.
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Electrochemical Series
The standard electrode potentials (reduction) of a number of electrodes are given in Table. These
values are said to be on the hydrogen scale since in these determinations, the potential of the standard
hydrogen electrode used as the reference electrode has been taken as zero. When the various electrodes
are arranged in the order of their increasing values of standard reduction potentials on the hydrogen
scale, then this arrangement is called electrochemical series.
Applications of Electrochemical Series or Electrode Potentials: The following are the
applications of electrochemical series:
[1]The standard emf of a cell can be calculated if the standard electrode potential values are known.
[2] The relative tendencies of metals to go into solution can be noted by the help of electrochemical
series. Metals on the top are more easily ionized into solution.
[3] The metal with negative reduction potential will displace hydrogen from an acid solution:
Zn + H2SO4 → ZnSO4 + H2
EZn0 = -0.76 volt
The metal with positive reduction potential will not displace hydrogen from an acid solution.
Ag + H2SO4 → no reaction EAg0 = +0.80 volt
[4] Metals which lie higher in the series can displace from the solution those elements which lie below
them in the series. Thus, zinc can displace copper from the solution.
[5] The anodic or more active metals in the series are more prone to corrosion. The cathodic or more
noble metals are less prone to corrosion. The galvanic series is used for this purpose.
Galvanic Series: the galvanic series is used to provide sufficient information in predicting the corrosion
behaviour in a particular set of environmental conditions. Oxidation potential measurements of various metals and alloys
have been made using the standard calomel electrode as the reference electrode and immersing the metals and alloys in
sea water. These are arranged in decreasing order of activity and this series is known as the galvanic series. The galvanic
series gives more practical information on the relative corrosion tendencies of the metals and alloys. The speed and
severity of corrosion depends upon the difference in potential between the anodic and cathodic metals in contact.
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Using Electrochemical Series or Electrode Potentials
In Table the standard reduction potentials (E°) are arranged in the order of increasing potentials.
The relative position of electrodes (M/M+) in the table can be used to predict the reducing or oxidising
ability of an electrode.
The electrodes that are relatively positive indicate that reduction reaction involving addition of
electrons, M+ + e– ⎯⎯→ M is possible. In case of relatively negative potential involving loss of
electrons, M ⎯⎯→ M+ + e– is indicated. It also follows that the system with higher electrode potential
will be reduced by the system with lower electrode potential.
Predicting the Oxidising or Reducing Ability
Let us consider a series of elements Cu, H2, Ni, Zn and their ions. These four elements could act
as reducing agents. On the other hand, their ions Cu2+, H+, Ni2+ and Zn2+ can act as electron acceptors
or oxidising agents. If we list the respective half-reactions (or electrodes) in order of descending E°
values, we will have placed the oxidising agents in descending order of their ability to attract electrons.
It is noteworthy that the value of E° becomes more negative down the series. This means that
Cu2+ is the best oxidising agent (most electron-attracting ion) of those in the list. That is, Cu2+ shows
the greatest tendency to be reduced. Conversely, Zn2+ is the worst oxidising agent, being the least
electron-attracting ion. Of the elements Cu, H2, Ni and Zn, Zn is the best reducing agent (best electron
donor), since E° for the half-reaction, Zn ⎯⎯→ Zn2+ + 2e–
E° = + 0.76 V has the most positive value.
By the reasoning, Cu is the worst reducing agent. The table of half reaction potentials above tells us
that at standard conditions the following reactions occur spontaneously.
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Some important points concerning the Table of Standard Reduction Potentials are:
(1) The more positive the value of E°, the better the oxidising ability (the greater the tendency to be
reduced) of the ion or compound, on moving upward in the Table.
(2) The more negative the value of E° the better the reducing ability of the ions, elements or
compounds on moving downward in the Table.
(3) Under standard conditions, any substance in this Table will spontaneously oxidise any other
substance lower than it in the Table.
Predicting cell emf
The standard emf, E°, of a cell is the standard reduction potential of right-hand electrode (cathode)
minus the standard reduction potential of the left-hand electrode (anode). That is,
Let us predict the emf of the cell
by using the E° values from the Table
The answer is so clear from Figure 9
Fig. 8: Diagrammatic representation of Cell emf.
Predicting Feasibility of Reaction
The feasibility of a redox reaction can be predicted with the help of the electrochemical series.
The net emf of the reaction, Ecell can be calculated from the expression
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Predicting whether a metal will displace another metal from its salt solution or not
As already shown, the metals near the bottom of the electrochemical series are strong reducing
agents and are themselves oxidised to metal ions. On the contrary, the metals lying higher up in the
series are strong oxidising agents and their ions are readily reduced to the metal itself. For example,
zinc lying down below the series is oxidised to Zn2+ ion, while copper which is higher up in the series
is produced by reduction of Cu2+ ion.
Zn ⎯⎯→ Zn2+ + 2e–
Cu2+ + 2e– ⎯⎯→ Cu↓
Thus when zinc is placed in CuSO4 solution, Cu metal gets precipitated. In general we can say that
a metal lower down the electrochemical series can precipitate the one higher up in the series.
Silver cannot precipitate Cu from CuSO4, solution, since both metals have positions higher up in the
series and are strong oxidising agents.
Predicting whether a metal will displace hydrogen from a dilute acid solution
Any metal above hydrogen in the electrochemical series is a weaker reducing agent than
hydrogen itself and cannot reduce H+ to H2. Any metal lying below hydrogen is a stronger reducing
agent than hydrogen and will convert H+ to H2. This explains why Zn lying below hydrogen reacts
with dil H2SO4 to liberate H2, while Cu lying above hydrogen does not react.
The Nernst Equation
We know experimentally that the potential of a single electrode or half-cell varies with the
concentration of ions in the cell. In 1889 Walter Nernst derived a mathematical relationship which
enable us to calculate the half-cell potential, E, from the standard electrode potential, E°, and the
temperature of the cell. This relation known as the Nernst equation can be stated as
where
E° = standard electrode potential
R = gas constant
T = Kelvin temperature
n = number of electrons transferred in the half-reaction
F = Faraday of electricity
K = equilibrium constant for the half-cell reaction as in equilibrium law
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Derivation
Consider a reversible cell, e.g. Daniel cell. The overall cell reaction occurring in the cell is
represented by the equation
In general, for a reversible cell the equation is
The electrical energy of a reversible cell can be measured by the free energy decrease (-∆G) of
the reaction taking place in the cell. In the cell, if the reaction involves the transfer of n number of
electrons, then n Faradays of electricity will flow. If E is the emf of the cell, then the total electrical
energy produced in the cell is
where -∆G = decrease in free energy.
At standard conditions,
where -∆Go = standard free energy change.
Standard free energy change is the change in free energy when the concentration of reactants
and products are unity. EJ is the standard emf of the cell in which the reactants and products are kept at
1 molar concentration at 25°C, ∆G and ∆G0 are related as
where K is the equilibrium constant of the reaction defined as the ratio of the concentration of the
products to the concentration of the reactants.
Divide the above equation by –nF
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At 25°C, T = 298 K, R = 8.314 11K/mol, F = 96500 coulombs. Therefore,
This is known as the Nernst equation. This equation may be used to calculate the emf of cells
when concentration of reactants and products of the cell reactions are known.
Applications of the Nernst Equation
[1] Calculation of Half-cell potential: For an oxidation half-cell reaction when the metal
electrode M gives Mn+ ion,
The Nernst equation takes the form
The concentration of solid metal [M] is equal to zero. Therefore, the Nernst equation can be written as
Substituting the values of R, F and T at 25°C, the quantity 2.303 RT/F comes to be 0.0591. Thus
the Nernst equation can be written in its simplified form as
This is the equation for a half-cell in which oxidation occurs. In case it is a reduction reaction,
the sign of E will have to be reversed.
[2] Calculation of Cell potential: The Nernst equation is applicable to cell potentials as well. Thus,
K is the equilibrium constant of the redox cell reaction.
[3] Calculation of Equilibrium constant for the cell reaction: The Nernst equation for a cell is
19
At equilibrium, the cell reaction is balanced and the potential is zero. The Nernst equation may,
now, be written as
Concentration Cells
A concentration cell is an electrochemical cell which produces electrical energy by the transfer
of material from a system of higher concentration to a system of lower concentration. The difference in
concentration may be due to the difference in concentration of the electrodes or electrolytes. Based on
this, concentration cells are divided into two categories; (i) Electrode Concentration Cells and
(ii) Electrolyte Concentration Cells
Electrode Concentration Cells
In electrode concentration cells, two electrodes of the same metal with different concentrations
are dipped in the same solution of the electrolyte, e.g. amalgam concentration cells. Amalgam
electrodes are produced by mixing various proportions of any metal and mercury. Depending on the
amount of metal taken during the preparation of amalgam, the concentration of electrodes will vary.
Electrolyte Concentration Cells
Cell potentials depend on concentration of the electrolyte. Thus a cell can be constructed by
pairing two half-cells in which identical electrodes are dipping in solution of different concentrations of
the same electrolyte. Such a cell called Electrolyte concentration cell. It may be described as:
a cell in which emf arises as a result of different concentrations of the same electrolyte in the
component half-cells.
The electrode dipped in the solution of lower concentration is an anode and the electrode
dipped in solution of higher concentration is a cathode. The overall reaction involves the transfer
of material from higher concentration to lower concentration. The two solutions are connected together
by a salt bridge.
A typical concentration cell is shown in Figure. It consists of two silver electrodes, one
immersed in 0.1 M silver nitrate solution and the other in 1 M solution of the same electrolyte. The two
solutions are in contact through a membrane (or a salt bridge). When the electrodes are connected by a
wire, it is found experimentally that electrons flow from the electrode in more dilute (0.1M) solution to
that in the more concentrated (1 M) solution.
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Explanation: The concentration of Ag+ ions in the left compartment is lower (0.1M) and in the
right compartment it is higher (1M). There is a natural tendency to equalize the concentration of Ag+
ions in the two compartments. This can be done if the electrons are transferred from the left
compartment to the right compartment. This electron transfer will produce Ag+ ions in the right
compartment by the half-cell reactions:
Fig. 9: A typical Electrolyte Concentration Cells
Thus in a bid to equalize concentration of Ag+ ions in the two compartments the cell will
develop emf (Le Chatelier’s Principle) to cause the transfer of electrons. Eventually, the solutions in
two compartments will have equal Ag+ ion concentration and there will be no emf recorded.
Emf of Concentration Cell
Suppose the concentrations in the two half-cells are C1 and C2 at 25°C, C2 being greater than C1.
Then emf, E, of the concentration cell will be given by the difference between the two electrode
potentials. In terms of Nernst equation
21
where EM is the standard electrode potential of the metal M and n is the valence of the ions in contact
with it. For example, the emf, E, of the concentration cell
can be found by substituting the values in the above equation.
Applications of Concentration Cells
[1] Determination of Solubility of Sparingly soluble salts: The solubility of a sparingly soluble
salt like AgCl can be calculated by measuring the emf of the cell.
The cell can be constructed very easily by placing one of the silver electrodes in contact with O.OIN
solution of silver nitrate and the other in contact with 0.0l N solution of potassium chloride. The two
solutions are connected by a salt bridge. A drop of silver nitrate solution is added to the potassium
chloride solution. The small amount of silver chloride formed u sufficient to give its saturated solution.
The cell so constructed is a concentration cell with respect to silver ions. The emf of the above cell is
given by the relation
Here n =1 and C2 =0.0l N. Hence,
From the above expression, we can calculate the concentration of AgCI in gram equivalent/ litre.
Multiplying this by 143.5, the equivalent weight of AgCI, we get the solubility of AgCl in gram/litre.
[2] Determination of the valency of ions in doubtful cases: By substituting the measured value
of emf of a concentration cell and the values of C1 and C2 in the equation
The value of n, the number of electrons involved in the cell reaction or the valency of ions can
be calculated. For example, from the concentration cell the value of n involving mercurous ions has
been found to be 2. Hence, the formula of mercurous ion is Hg22+.
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Types of Electrodes (M/M+, M/MX/X-, M/A , A , Pt/H2,H+, glass electrode)
[1] Metal-metal ion electrode {M (s) /Mn+ (aq)}: It consists of a pure metal [M (s)] in contact
with a solution of its cation [Mn+ (aq)]. It is represented as M (s) /Mn+ (aq). The electrode reaction can
be written as Mn+ (aq) + ne- → M (s). Electrode potential is
E/ E/
RT
C
nF C
where C is the concentration of M and Mn+ respectively.
Example: Zinc electrode dipping in ZnSO4 solution or Copper electrode dipping in Cu SO4 solution.
[2] Metal-metal insoluble salt electrode {M (s) /MX (s)/X- (aq)}: It consists of a pure metal
[M (s)] coated by a layer of its sparingly soluble salt [MX (s)] and kept immersed in a solution
containing a common anion [X- (aq)]. It is represented as M (s) /MX (s)/X- (aq). The electrode reaction
can be written as MX (s) + ne- → M (s) + X- (aq). Electrode potential is
E// E//
RT C . C
nF
C
Example is Calomel electrode: The standard calomel electrode, SCE, consists of a wide glasstube with a narrow side-tube. It is set up as illustrated in Fig. 10. A platinum wire is dipping into liquid
mercury covered with solid mercurous chloride (Hg2Cl2, calomel). The tube is filled with a 1 M
solution of KCl (or saturated KCl solution). The side-tube containing KCl solution provides the salt
bridge which connects the electrode to any other electrode. The calomel electrode is represented as
The half-cell reaction is
Calomel electrode
Fig. 10: The calomel electrode coupled determine its emf with
zinc electrode to
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[3] Redox electrode/ , !: It contains an inert electrode (i.e., M = Pt) dipped in a
solution
containing
ions
of
the
same
substance
in
two
oxidation
states "A , A # .
The electrode is represented as M$s&/A $C' &, A $C( & . The electrode reaction can be written as
A $C' & ) ne+ , A $C( & where C1 and C2 are the concentrations of two ions A and A
respectively. The electrode potential is
E
// ,/
E
// ,/
RT C(
nF C'
Example is electrolyte concentration cell. It consists of two silver electrodes, one immersed in
0.1 M silver nitrate solution and the other in 1 M solution of the same electrolyte.
[4] Gas electrode {Pt/H2, H+}: It contains an inert electrode like platinum dipped in a solution
containing
ions
(H+
ions)
to
which
the
gas
(H2
gas)
is
bubbled
continuously.
The electrode is represented as Pt (s)/H2 (P = x atm), H+ (aq). The electrode reaction can be written as
2H+ (aq) + 2e- → H2 (g). The electrode potential is
E01/2/2 E01/2
/2
C2
RT
nF $C2 &(
Example is Standard Hydrogen Electrode (SHE). It consists of a platinum electrode immersed in
a 1 M solution of H+ ions maintained at 25°C. Hydrogen gas at one atmosphere enters the glass hood
and bubbles over the platinum electrode.
[4] Glass electrode: It consists of a glass tube having a thin-walled bulb at the lower end.
The bulb contains a 1M HCl solution. Sealed into the glass-tube is a silver wire coated with silver
chloride at its lower end. The lower end of this silver wire dips into the hydrochloric acid, forming
silver-silver chloride electrode. The glass electrode may be represented as
When placed in a solution, the potential of the glass electrode depends on the H+ ion
concentration of the solution. The potential develops across the glass membrane as a result of a
concentration difference of H+ ions on the two sides of the membrane. This happens much in the same
way as the emf of a concentration cell develops.
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A commonly used secondary standard electrode is the so-called glass electrode.
Its emf is determined by coupling with a standard calomel electrode (SCE). The glass electrode
provides one of the easiest methods for measuring the pH of a given solution. A simple type of glass
electrode is shown Fig. 11.
Fig. 11: Glass electrode
Additional Note:
Relation between emf and free energy: When a cell produces a current, the current can be
used to do work–to run a motor, for instance. Thermodynamic principles can be employed to derive a
relation between electrical energy and the maximum amount of work, W456 obtainable from the cell.
The maximum amount of work obtainable from the cell is the product of charge flowing per mole and
maximum potential difference, E, through which the charge is transferred.
W456 nFE …………….. (1)
where n is the number of moles of electrons transferred and is equal to the valence of the ion
participating in the cell reaction. F stands for Faraday and is equal to 96,500 coulombs and
E is the emf to the cell.
According to thermodynamics, the maximum work that can be derived from a chemical reaction
is equal to the free energy (∆G) for the reaction.
W456 ∆G ……………. (2)
Therefore from (1) and (2), we can write
∆G nFE
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Batteries
Introduction
Batteries are considered as storehouses for electrical energy on demand. An electrochemical
power source or battery is a device which converts the chemical energy derived from a chemical
reaction into electrical energy.
The chemical reaction involved in a battery is a redox reaction and some of these reactions are
reversible. The reverse reaction can be brought out by supplying energy, i.e. by applying the current to
the system from an external source. This process is called charging of the battery. Such batteries in
which the cell reactions can be reversed by passing direct electric current in opposite direction are
called secondary batteries. Lead-acid accumulators, nicad battery, etc. are secondary batteries.
Batteries that cannot be charged in this manner, because the cell reactions are irreversible, are called
primary batteries. Primary batteries are the most common batteries available today because they are
cheap and simple to use. Carbon-zinc dry cells, alkaline batteries, mercury batteries, lithium batteries,
etc. are examples for primary batteries. Table shows the differences between primary and secondary
batteries. In the following sections we will discuss in detail primary and secondary batteries.
Primary Batteries
Lechlanche Cell or Zinc-Carbon Dry Cell: The zinc-carbon cell uses a zinc anode, a
manganese dioxide cathode and an electrolyte of ammonium chloride dissolved in water. Powdered
carbon is mixed with manganese dioxide to improve conductivity and retain moisture.
In the dry cell which is a modified Lechlanche cell consists of a zinc anode shaped as a container
for the electrolyte and a graphite rod as a cathode surrounded by MnO2 and a paste of NH4CI and ZnCl2
as an electrolyte. The dry cell is so called since there is no fluid phase presents (Figure 1).
The voltage of the cell is about 1.5 volt. It is commonly used in flash lights and radios.
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Chemistry: As the cell is discharged zinc is oxidized and manganese dioxide is reduced:
Figure 1: Interior section of a dry cell
Drawbacks
of
dry
cells:
The
following
are
the
drawbacks
of
dry
cells:
[1] Dry cells do not have an indefinite life period. Since the electrolyte medium is acidic,
zinc metal dissolves slowly there by reducing the life time even if it is not in use.
[2] When current is rapidly drawn from the cell, due to the building up of products on the electrodes
there is a drop in voltage.
Alkaline
Battery:
An
alkaline
battery
is
an
improved
form
of the
dry
cell.
In this cell, the electrolyte is a concentrated solution of potassium hydroxide (35-52%).
As the electrolyte in this cell is an alkali, it is called alkaline battery. The cathode is made from a
mixture of manganese dioxide and carbon. The anode mix consists of alkaline electrolyte, zinc powder
and small quantity of gelling agent (starch) to immobilize the electrolyte and suspend the zinc powder.
The alkaline cell derives power from the oxidation of zinc anode and the reduction of MnO2 cathode.
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The emf of the cell is 1.5 volt. The advantages of this battery over a dry battery are as follows:
[1] As potassium hydroxide is used as an electrolyte, zinc does not readily dissolve.
[2] Since there is no corrosion of zinc, the life of alkaline battery will be longer.
[3] Its output capacity is high.
Secondary Batteries or Storage Batteries or Accumulators
Lead-acid Accumulator or Acid Battery: A storage cell is one which can operate both as a
voltaic cell and as an electrolytic cell. When it acts as a voltaic cell it supplies electrical energy. On
recharging it acts as an electrolytic cell.
An acid battery consists of a negative electrode of porous lead (lead sponge) as the anode and a
positive electrode of lead dioxide as the cathode. A number of such electrode pairs are immersed in an
aqueous solution of 20% sulphuric acid (specific gravity: 1.15 at 25°C) which is the electrolyte
(Figure 2).
Figure 2: Lead-acid storage cell
Discharging: The electrode reactions that occur during the discharge of the cell, i.e. when
current is drawn from the cell are as follows:
Anode
Cathode
The overall cell reaction is as follows:
The PbSO4 formed gets precipitated on the cathode and in the solution.
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Charging: When both anode and cathode become covered with PbSO4, the cell stops its
functioning. Recharging is done by applying a voltage across the electrodes that is slightly higher than
the voltage that the battery can deliver. The net reaction during discharging is as follows:
During the discharge process the consumption of sulphuric acid is replaced by an equivalent
quantity of water and the sulphuric acid concentration decreases. On charging the reverse reaction takes
place. During the reverse reaction water is consumed and sulphuric acid is regenerated. Hence the
original strength of acid is restored. Since both of these changes are associated with variations in the
specific gravity of the acid, the extent of charge or discharge of the cell at any time can be determined
by testing the specific gravity of the acid.
Uses: The following are the uses of lead-acid storage cells: [1] Lead-acid storage cells are used
in automobiles, hospitals, telephone exchanges, etc. [2] As it is rechargeable it is used in UPS
(uninterrupted power supply) a power system which maintains current flow without even a momentary
break in the event of current failure.
Maintenance: If the lead-acid batteries are properly maintained in the following ways,
they can function for long periods: [1] Avoid over discharging of the battery. [2] Maintain the
electrolyte at the proper level by adding water (whenever required). [3] Keep the battery clean.
[4] Avoid overheating of the battery.
Nickel-Cadmium Battery or Nicad Battery: A nickel-cadmium battery is a type of alkaline
storage battery. This battery consists of a cadmium anode, nickel oxyhydroxide cathode and an alkaline
electrolyte (potassium hydroxide).
During discharge cadmium metal oxidizes to cadmium hydroxide at the anode:
By accepting the electrons, nickel oxyhydroxide [NiO(OH)] is reduced to nickel hydroxide [Ni(OH)2]
at the cathode:
The overall cell reaction is:
The emf of the cell is 1.3 volt. This is a rechargeable battery. When the discharged battery is connected
to an external voltage source, the cell reaction is reversed.
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Advantages: Nickel--cadmium cells are characterized by long life, relatively high rates of
discharge and charge and ability to operate at low temperatures.
Uses: Nicad cells are used in calculators, electrical shavers, etc. The good low
temperature performance has lead to wide use of nickel-cadmium batteries in aircraft and space
satellite power systems.
Fuel cells
A fuel cell is an electrochemical cell in which the chemical energy of the fuel-oxidant system is
directly converted into electrical energy. It is an energy conversion device or electricity generator.
A fuel cell operates like a galvanic cell with the exception that the reactants are supplied from
outside. It is an example for a primary cell. This is capable of supplying current as long as it is provided
with the supply of reactants.
Hydrogen-Oxygen Fuel Cell: A fundamental and important example of a fuel cell is the
hydrogen/oxygen cell. Like an electrochemical cell, the fuel cell is also having two electrodes and an
electrolyte. The two electrodes are made up of porous graphite admixed with nickel powder.
The electrolyte used is potassium hydroxide solution maintained at 200°C and 20-40 atmospheres.
Hydrogen and oxygen gases are bubbled through the anode and cathode compartments respectively.
Hydrogen is oxidized at the anode whereas the oxygen gets reduced at the cathode (Figure 3).
Figure 3: Hydrogen-oxygen fuel cell
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The cell reaction is the same as combustion of hydrogen in air or oxygen. Generally a large
number of these cells are stacked together in series to make a battery called fuel cell battery
or fuel battery.
In the fuel cells, gaseous fuels used are hydrogen, alkanes and co. Among the liquid fuels
methanol, ethanol, etc. are very important. Oxygen, air, hydrogen peroxide, etc. are some of the
oxidants used.
Advantages
of
fuel
cells:
The
following
are
the
advantages
of
fuel
cells:
[1] The energy conversion efficiency is very high (75-83%). [2] They are used as power sources in
spacecrafts. [3] The product of a hydrogen-oxygen fuel cell is pure water which can be used for
drinking purpose. [4] Noise and thermal pollution are very low. [5] The maintenance cost is very low.
[6] It saves fossil fuels.
Limitations: The following are the limitations of fuel cells: [1] The cost of power from a fuel
cell is high as result of the cost of electrodes and pure hydrogen gas. [2] As the fuels used are gases,
they have to be stored in big tanks under high pressures
Solar Cells or Photovoltaic Cells
Photovoltaic
cells
convert
solar
energy
directly
into
electrical
energy.
Description of solar cells: The conventional solar cells made up of p-type doped semiconductor
(i.e., silicon doped with boron) and n-type doped semiconductor (i.e., silicon doped with phosphorus) is
shown in Figure 4
Figure 4: Solar cell
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The surface layer is made up of n-type semiconductor and it is extremely thin (~ 0.2 µm) so that
sun light can penetrate through it. The bottom layer is made up of p-type semiconductor. The electrodes
made of Ti-Ag solder are attached to both the sides to provide electrical contact. The electrode on the
top surface is in the form of a metal grid with fingers which permit sun light to pass through. On the
back side, the electrode completely covers the surface. An anti-reflection coating of silicon oxide
having a thickness of 0.1 µm is also put on the top surface.
Working: When p-type and n-type semiconductors are placed together, electrons from
n-type side diffuse across p-n junction to combine with holes present in p-type semiconductor side.
As a result positive ions (on n-type side) and negative ions (on p-type side) are created near
the junction to certain thickness. The separation of charges produces an electric field across
p-n junction which is about 0.6-0.7 volt. This potential at p-n junction prevents the charges moving
across it further.
When light strikes on the p-n junction, electron-hole pair is produced. As the potential barrier
resists the flow of charge carriers across it, they flow through the conductor/load connected externally
to produce electric current. Since the emf of a single solar cell is about 0.6 V, they are arranged into
larger groupings called arrays in solar panels.
Applications: The following are the application of solar cells: [1] Solar cells are used to provide
power supply for space satellites. [2] They are used for the distillation of water to get pure drinking
water. [3] Solar cells provide thermal energy for solar cookers, solar furnaces, etc.
[4] They provide electricity for street lighting in remote areas, and to run water pumps and radios in
desert areas. [5] Solar cells provide electric power to light houses.
Chemical sensors
A chemical sensor is a transducer which provides direct information about the chemical
composition of its environment. It can warn occupants of potentially toxic agents in air. It consists of a
physical transducer and a chemically selective layer.
Actually, it contains an array of tiny micro-hot-plates in conjunction with thin metalized films
such as tin, titanium, or zinc oxides. Both the hot plates and sensing films are incorporated into an
integrated circuit device. Sensitive films create a sensitive surface for detecting ambient chemicals.
If a specific chemical of interest is present, the resistance of the device changes. It produces a type of
“signature” for a specific chemical that can be matched up against a library of chemical signatures to
identify both the type and concentration of the gas in the ambient air. Metal-oxide semiconductors
(MOS) structures are of particular interest for chemical sensors.
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Syllabus
Electrochemistry – single electrode potential – Helmholtz double layer – Nernst equation –
derivation – types of electrodes (M/M+, M/MA/A-, M/A+,A2+, Pt/H2,H+, glass electrode) Electrochemical cells, concentration cells - salt bridge – emf measurement – Poggendorf’s
compensation method – Electrochemical series – applications
Storage cells – Lead acid accumulator – alkaline cells – Nickel cadmium – fuel cells –
H2/O2 fuel cell – solar cells – Chemical sensors.
Previous University Questions
[1] Write the redox reactions taking place in the lead acid accumulators {2 marks}
[2] How is EMF of an electrochemical cell determined through Poggendorf’s compensation
method? {5 marks}
[3] What is electrochemical series? What are its applications? {5 marks}
[4] Write a descriptive account on Fuel cells and solar cells {10 marks}
[5] What is reference electrode? Give examples {5 marks}
[6] What is the effect of electrolyte concentration on electrode potential {2 marks}
[7] Derive the Nernst equation for electrode potential {5 marks}
[8] Write a descriptive account on Fuel cells and Nickel-Cadmium cell {10 marks}
ietchemistry@gmail.com
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