Electrochemistry Electrochemistry deals with the links between chemical reactions and electricity. This includes the study of chemical changes caused by the passage of an electric current across a medium, as well as the production of electric energy by chemical reactions. Electrochemistry also embraces the study of electrolyte solutions and the chemical equilibria that occur in them. The devices used for the inter-conversion between chemical and electrical forms of energy are called electrochemical cells. Electrochemical cells which generate an electric current are called voltaic cells or galvanic cells, and common batteries consist of one or more such cells. In other electrochemical cells an externally supplied electric current is used to drive a chemical reaction which would not occur spontaneously. Such cells are called electrolytic cells. Electrolytic cell An electrolytic cell decomposes chemical compounds by means of electrical energy, in a process called electrolysis. The result is that the chemical energy is increased. 1
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Electrochemistry
Electrochemistry deals with the links between chemical reactions and electricity.
This includes the study of chemical changes caused by the passage of an electric current
across a medium, as well as the production of electric energy by chemical reactions.
Electrochemistry also embraces the study of electrolyte solutions and the chemical
equilibria that occur in them.
The devices used for the inter-conversion between chemical and electrical forms
of energy are called electrochemical cells.
Electrochemical cells which generate an electric current are called voltaic cells or
galvanic cells, and common batteries consist of one or more such cells. In other
electrochemical cells an externally supplied electric current is used to drive a chemical
reaction which would not occur spontaneously. Such cells are called electrolytic cells.
Electrolytic cell
An electrolytic cell decomposes chemical compounds by means of electrical
energy, in a process called electrolysis. The result is that the chemical energy is
increased.
An electrolytic cell has three component parts: an electrolyte and two electrodes
(a cathode and an anode). The electrolyte is usually a solution of water or other solvents
in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes.
When driven by an external voltage applied to the electrodes, the electrolyte provides
ions that flow to and from the electrodes, where charge-transferring or redox, reactions
can take place.
The electrolytic cell is the industrial chloralkali cell in which brine (an aqueous sodium
chloride solution) is electrolytically converted to chlorine and caustic soda (sodium
hydroxide, NaOH). The external power source supplies electric energy to drive the
overall reaction.
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2Cl− + 2H2O → Cl2 + H2 + 2OH−
Chloride ion is oxidized to chlorine gas at the carbon electrode, and water is reduced to
hydrogen gas (H2) and hydroxide ion (OH−) at the iron electrode. The electrolytes are
maintained as electrically neutral by a flow of sodium ions through the separator (such as
an ion exchange membrane).
For the electrolytic cell, the external markings of anode and cathode are opposite the
chemical definition. That is, the electrode marked as anode for discharge acts as the
cathode while charging and the electrode marked as cathode acts as the anode while
charging.
Voltaic Cells
When zinc metal is placed in a solution of copper ions as described by the net
ionic equation shown below.
Cu+2 (aq) + Zn (s) -------> Cu(s) + Zn+2 (aq)
The zinc metal slowly "dissolves" as its oxidation produces zinc ions which enter into
solution. At the same time, the copper ions gain electrons and are converted into copper
atoms which coats the zinc metal or sediments to the bottom of the container. The energy
produced in this reaction is quickly dissipated as heat, but it can be made to do useful
work by a device called, a voltaic cell.
A voltaic cell is composed to two compartments or half-cells, each composed of
an electrode dipped in a solution of electrolyte. These half-cells are designed to contain
the oxidation half-reaction and reduction half-reaction separately as shown below.
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The half-cell, called the anode, is the site at which the oxidation of zinc occurs as shown
below.
Zn (s) ----------> Zn+2 (aq) + 2e-
During the oxidation of zinc, the zinc electrode will slowly dissolve to produce zinc ions
(Zn+2), which enter into the solution containing Zn+2 (aq) and SO4-2 (aq) ions.
The half-cell, called the cathode, is the site at which reduction of copper occurs as
shown below.
Cu+2 (aq) + 2e- -------> Cu (s)
When the reduction of copper ions (Cu+2) occurs, copper atoms accumulate on the surface
of the solid copper electrode.
The reaction in each half-cell does not occur unless the two half cells are connected to
each other.
In order for oxidation to occur, there must be a corresponding reduction reaction that is
linked or "coupled" with it. Moreover, in an isolated oxidation or reduction half-cell, an
imbalance of electrical charge would occur, the anode would become more positive as
zinc cations are produced, and the cathode would become more negative as copper
cations are removed from solution. This problem can be solved by using a "salt bridge"
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connecting the two cells as shown in the diagram below. A "salt bridge" is a porous
barrier which prevents the spontaneous mixing of the aqueous solutions in each
compartment, but allows the migration of ions in both directions to maintain electrical
neutrality. As the oxidation-reduction reaction occurs, cations ( Zn+2) from the anode
migrate via the salt bridge to the cathode, while the anion, (SO4-2), migrates in the
opposite direction to maintain electrical neutrality.
The two half-cells are also connected externally. In this arrangement, electrons
provided by the oxidation reaction are forced to travel via an external circuit to the site of
the reduction reaction.
The differences between galvanic and electrolytic cells can be summarised in a table.
Galvanic/Voltaic Cells Electrolytic Cells
chemical energy electrical energy electrical energy chemical energy
Two half-cells with separate electrolytes and a
salt bridge (or porous barrier).
Electrodes usually in the same electrolyte
chemical reaction is spontaneous chemical reaction is forced by applying a voltage -
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Eo total is positive it is not spontaneous
Eo total is negative
The electrode on which oxidation takes place
is called the anode and the electrode on which
reduction takes place is called cathode.
anode - negative terminal
cathode - positive terminal
The electrode which is connected to the negative
terminal of the battery is called cathode / the cations
migrate to it which gains electrons and hence
reduction takes place.
anode - positive electrode
Cathode - negative electrode
Reversible and Irreversible cells:
Reversible cells
Electrochemical cells can also be said as two types namely reversible and
irreversible cell. A cell works reversibly in the thermodynamic conditions i.e., during the
measurement of EMF no current flows through the cell and no chemical reaction takes
place. Such cells are called as reversible cells.
If the external EMF is infinitely greater than that of the cell emf, an extremely
small amount of current flows through the cell in the opposite direction and small amount
of the chemical reaction takes place in the reverse direction. E.g. Daniel cell is a
reversible cell. Its cell potential is 1.1 V. Thus in Daniel cell (a galvanic cell), zinc
undergoes dissolution and copper undergoes deposition to realize an emf of 1.1V, as per
the following reaction sequence:
Zn + Cu2+ === Zn2+ + Cu …….. 1.1 V
If an emf of –1.101 V is impressed on Daniel cell, copper undergoes dissolution and zinc
undergoes deposition.
Irreversible cells
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Cells which do not obey the conditions of thermodynamic reversibility are called
irreversible cells. Irreversible is a cell where the cell reaction cannot be reversed even on
applying infinitesimally small but excess applied emf i.e. the products produced during
the cell reaction are not available for recombination on reversal of voltage. Example of an
irreversible cell is the cell used for the electrolysis of brine or the dry cell used in pen-
torches. In the electrolysis of brine (aqueous NaCl solution) for example, on applying
voltage, Na+ ions move towards cathode, gain one electrode and become elemental
sodium atoms. But the sodium atoms immediately react with water to form sodium
hydroxide. Similarly, chloride ions move towards anode, loose one electron to form
chlorine atoms. These chlorine atoms recombine forming molecular chlorine, which is
evolved as a gas. The reaction sequence is given below:
Na+ + e → Na ; 2 Na + 2 H2O → 2 NaOH + H2
Cl- → Cl + e ; Cl + Cl → Cl2 ↑
Cell diagram or representation of a cell
In general, the electrode at which reduction takes place is written on the RHS of the salt
bridge and the electrode at which oxidation takes place is written on the LHS of the salt
bridge. The salt bridge linking the aqueous solutions is represented by two vertical
parallel lines having ions on both sides.
For Zn – Cu cell, Zn electrode is written on the LHS while the Cu electrode on the RHS
of the salt bridge.
Zn Zn2+ Cu2+ Cu
The symbol for an inert electrode, like the platinum electrode is often enclosed in a slash.
For example,
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Mg Mg2+ H+ H2 Pt
The value of emf of a cell is written on the right of the cell diagram. Thus a zinc-copper cell has emf 1.1 V and is represented as
Zn Zn2+ Cu2+ Cu E = +1.1 V
Electrode Potential
Consider a zinc rod being dipped in ZnSO4 solution. The zinc atoms on the surface of the
metal with in the solution have tendency to release Zn+2 into solution retaining the
electrons on the surface of the metal. This process is called dissolution or solution
pressure of the metal and it is oxidative in nature.
The zinc ions of the solution have a tendency to accept the electrons on the surface of
zinc rod to form neutral zinc atoms and get deposited on the zinc rod. This process is
called deposition or the osmotic pressure of solution is reductive in nature.
These two processes will be taking place simultaneously at different rates. In this case,
the rate of dissolution is found to be greater than the rate of deposition. Consequently, by
the time equilibrium is reached, more of dissolution would have occurred and the solution
becomes negatively charged. Due to the attractive electrostatic forces, the Zn ions
accumulate around the Zn rod and an electrical double layer of opposite charges is
formed and between the Zn rod and the solution the potential developed is called single
electrode potential or Electrode Potential.
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Tendency of an electrode to lose or gain electrons when it is in contact with its
own ions in solution is called electrode potential. Tendency to gain electrons means
reduction potential and tendency to lose electrons means the oxidation potential.
Standard Electrode Potential (Eo)
It is the potential developed when the pure metal is in contact with its own ions at
one molar concentration at a temperature of 25oC or 298 K.
Cell potential or Electro Motive Force (EMF)
The emf (electro motive force) of a cell is the algebraic sum of the potentials of
the two constituent single electrode systems. It is obvious that cell is made of two half-
cells / single electrode systems. A cell is generally represented with the negative
electrode / anode written first at the left and then the cathode / positive electrode at the
right. Thus the emf of a galvanic cell is calculated from the half-cell potentials using the
relation
Ecell = Eright - Eleft = Ecathode - Eanode
Here it is to be noted the values of std potentials (reduction potentials) of cathodes
are more positive and those of anodes are more negative, so that the cell potential is
positive.
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e.g. for Daniel cell, Ecell = Eright - Eleft = Ecathode - Eanode =
ECu2+
/ Cu – EZn2+
/ Zn = 0.34 – (-0.76) = 1.10 Volt
Measurement of EMF
The potential difference, which causes current flow from the electrode of
higher potential to the electrode of lower potential, is called electromotive force (emf) of
the cell. The emf of a cell cannot be directly determined by connecting across a
voltmeter, as some part of the cell current is drawn by the voltmeter during its
measurement. This results in the formation of reaction products at the electrodes and
hence a change in the electrolyte concentration around the electrodes. This difficulty can
be overcome by the measurement of excess applied opposite emf that just nullifies the
cell emf (Poggendorff’s external compensation method). Care is to be taken during the
measurement such that the current taken from the cell is negligibly small and the ionic
concentrations are not appreciably altered. The emf of the cell thus remains constant and
its value can be determined with high degree of accuracy / precision.
The emf of a cell can be determined by Poggendorff’s external compensation
method. In this potentiometric measurement of cell emf, a standard cell is used whose
emf is known and does not vary with time. Weston cadmium cell is the conventionally
used standard cell. Fig. below is the schematic of a simple potentiometer.
It consists of a uniform wire AB of high resistance. A storage battery C of constant but
large emf is connected to the ends A and B of the wire through a variable resistance
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G
X (Cell unknown)
S (Cell - std)
A
C
B
R
K
D D’
(rheostat) R. The cell X whose emf is to be determined is included in the circuit through a
galvanometer G and a sliding contact D. The circuit is closed using a plug key K. The
position of the sliding contact D is slowly changed along the wire AB till a point is
reached at which there is no net current flowing through the galvanometer (its deflector
points to zero). This null point position D is noted.
The standard cell S (whose emf is known) is then introduced into the circuit and
the circuit closed using the plug key K. The position of the sliding contact D’ is again
slowly changed along the wire AB till a point is reached at which there is no net current
flowing through the galvanometer (its deflector points to zero). This null point position
D’ is noted. The emf of the cell X (unknown) is determined using the relation mentioned
below:
Emf of cell X (Ex) α balancing length AD
Emf of standard cell S (Es) α balancing length AD’s
Thus Es / Ex = (length AD’s) / (length ADx), from which Ex can be determined, as the
value of Es is known. Nowadays the digital potentiometers are used which have the in-
built circuitry of potentiometer set up and standard cell with switching arrangement for
standard cell so that the unknown cell is connected externally to directly read the cell emf
as digital display.
Relation between Gibb's Free Energy and Cell Potential (EMF) – Nernst Equation.
When a cell reaction takes place electrical energy is produced which results in decrease
in the free energy of the system.
Electrical work = Decrease in free energy
In an electro chemical cell,
Electric work done = Quantity of current produced x E.M.F.
For one mole of electrons quantity of current is 1F (96500 coulomb)
Therefore for n moles it is nF.
Electric work done = nFEcell
and For a standard cell
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For a general cell reaction
Van’t-Hoff isotherm says
G = Go + RT
Hence G = -nFE
-nFE = -nFEo + RT
Divide by nF and reverse the sign
Substituting R = 8.314 J/K/mole; F = 96500 coulombs; T = 298 K and multiplying 2.303
for conversion of ln to log, the equation becomes
Applications:
1) Determination of potential of cell and electrode
2) Free energy change can be determined using the equation
3) Equilibrium constant can be calculated using the relation
-∆G0 = RTlnK
Reference Electrodes:
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The potential of an electrode system (electrode of interest or working electrode)
can be measured by coupling with other electrode with a voltmeter introduced between
them. The coupled electrode should not possess any charge transfer reaction (electrode
reaction) in the electrolyte used or it should not be polarized. Such ideally non-
polarizable electrodes used for the measurement of working electrodes are called
reference electrodes. Reference electrodes are two types namely primary and secondary
reference electrodes. Primary reference electrode is one that is universally used such as
standard hydrogen electrode (SHE) and its potential is arbitrarily taken as zero. But
SHE involves tedious and cumbersome construction. This difficulty is overcome by the
use of ‘secondary reference electrodes, which can be constructed easily and their
potentials can be determined with SHE as reference.
It consists of a platinum electrode immersed in a 1 M solution of H+ ions
maintained at 25C. 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. It is represented as;
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1 M HCl
Pt. foil
H2 (1 atm.)
Pt, H2 (1atm) H+ (1M) Eo = 0 V
Limitations
i) It requires hydrogen gas and is difficult to set up and transport
ii) It requires large volume of test solution
iii) The solution may become poison, if any impurities presents on the surface of
Pt electrode.
iv) The potential of the electrode is dependent on atmospheric pressure.
Secondary Reference Electrode
Saturated calomel electrode (SCE)
Examples of secondary reference electrodes are calomel electrode, silver-silver
electrode, glass electrode, quinhydrone electrode etc. Calomel electrode is set up with
mercurous chloride, mercury and potassium chloride electrolyte and represented as Hg |
Hg2Cl2 | KCl.
Depending on the concentration of KCl, calomel electrode is of three types namely
saturated normal and decinormal calomel electrodes. The electrode reaction for calomel
electrode is Hg2Cl2 (s) + 2e- → 2Hg(s) + 2Cl-
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Saturated KCl
Hg + Hg2Cl2
Hg
Pt wire
The potentials of the different calomel electrodes against SHE reference is given below:
[ KCl ] Saturated 1.0 N 0.1 N
Potential , V 0.2422 0.2810 0.3335
Merits and Demerits of Calomel electrode: 1. Gives relative pH values compared to
hydrogen electrode, which gives absolute values of pH.
Ag/AgCl electrode is prepared by depositing a thin layer of AgCl electrolytically on a Ag
or Pt wire and immersing in a solution containing the chloride ions. It is represented as
Ag | AgCl | Cl-( M)
The electrode reaction of Ag/AgCl electrode is : AgCl + e → Ag + Cl-
Ion – Selective Electrode (ISE)
An Ion-selective electrode (ISE) is a transducer (sensor) which converts the
activity of a specific ion dissolved in a solution into an electrical potential which can be
measured by a voltameter or pH meter. The voltage if theoretically dependent on the
logarithm of the ionic activity, according to the Nernst equation.
The sensitive part of the electrode is usually made as an ion-specific membrane
along with a reference electrode. Hence ISE is also known as a specific ion electrode
(SIE).
Ion-selective electrodes are used in biochemical and biophysical research, where
measurements of ionic concentration in an aqueous solution are required, usually on a
real time basis.
Principle
At equilibrium, the membrane potential is mainly dependent on the concentration
of the target ion outside the membrane and is described by the Nernst equation. Briefly,
the measured voltage is proportional to the Logarithm of the concentration, and the
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sensitivity of the electrode is expressed as the electrode Slope - in millivolts per decade
of concentration. Unknown samples can then be determined by measuring the voltage
and plotting the result on the calibration graph.
The use of metals directly as ISE has the following disadvantages: (i) slow electrode
response (ii) Nernst equation not followed (iii) change of electrode potential due to the
availability of electrons on the electrode surface (iv) no well defined electron change.
Hence various membranes are used in ISEs and such electrodes are called Ion selective
membrane electrodes (ISMEs). ISMEs show some degree of specificity and selectivity,
These electrodes utilize some membrane to confine an inner solution and the reference
electrode. Membranes in the ISE and reference electrode (RE) sides function by ion
exchange (IE) mechanism.
There are three main types of membranes used in the Ion selective electrodes; they are
(i) Glass membrane
(ii) Solid-state crystal membrane and
(iii) Liquid ion-exchange membrane
Based on the ion selective membranes used, the ISE are classified into three groups
a) Glass membrane electrode
b) Crystalline membrane electrode
c) Liquid membrane electrode
1. Glass Membrane Electrodes
Glass membranes are made from an ion-exchange type of glass (silicate of
chalcogenide). This type of ISE has good selectivity, but only for several single-charged
cations; mainly H+, Na+, and Ag+. The glass membrane has excellent chemical durability
and can work in very aggressive media. A very common example of this type of electrode
is the pH glass electrode.
The selectivity of glass membranes depends depends on the composition of glass.
Generally they are based on Na2O-Al2O3-SiO2 mixtures.
Two typical composition of glass membranes are:
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i) Na2O (22%), CaO (6%) and SiO2 (72%) - responds to H+ ions
ii) Na2O (11%), Al2O3 (18%) and SiO2 (71%) – responds to alkali metal ions
Glass membrane electrodes are further subdivided into three types based on the
selectivity characteristics
i) pH type glass electrode
ii) Cation sensitive type
iii) Sodium sensitive type
Glass Electrode:
The glass elctrode assembly consists of a thin glass bulb filled with 0.1 N HCl and a
silver wire coated with silver chloride immersed in it. The Ag/AgCl electrode here acts as
the internal reference electrode. The glass electrode is represented as
Ag | AgCl(s) | 0.1 M HCl | glass.
Determination of pH of an aqueous solution by glass electrode:
The determination pf pH of a solution is one of the important applications of the
EMF measurements. When glass electrode is immersed in the solution whose pH is to be
determined, a potential difference is set up between the two surfaces of the glass
membrane. The potential value developed is proportional to the pH of the test solution
(sample).
The magnitude of this difference of potential (single electrode potential) is given by
EG = EoG - 0.0591 log [H+]
EG = EoG + 0.0591 pH
Where EoG is constant for the given glass electrode and it depends upon the nature and
composition of glass membrane.
Actually, the glass membrane of the glass electrode undergoes an ion-exchange reaction
in which the sodium ions of the glass membrane are exchanged with protons of the
sample solution. The electrode reaction of the glass electrode immersed in the test
solution can be represented as
glass ---- Na + H+ = glass ---- H + Na+
To carry out the determination of pH of a solution, the glass electrode is connected with a