This document provides information about electrochemistry. It discusses how electrochemistry deals with oxidation-reduction reactions that produce or utilize electrical energy. There are two main types of electrochemical cells: galvanic cells which convert chemical energy to electrical energy, and electrolytic cells which use electrical energy to drive non-spontaneous chemical reactions. The document explains the basic setup and workings of these two cell types.

1. ELECTROCHEMISTRY
Presented To:- Presented By:-
SIR EZAZ GILLANI SHEHMAN ASAD
Department of Chemistry Roll No: 03
MINHAJ UNIVERSITY MPHIL II Semester

2. ELECTROCHEMISTRY
 Electrochemistry is the branch of chemistry concerned with the
interrelation of electrical and chemical changes that are caused by the
passage of current.
 Study of chemical processes that cause electrons to move.
 Study of electricity and how it relates to chemical reactions.
 Chemical transformation occurring owing to the external applied electrical
current or leading to generation of electrical current is studied in
electrochemistry.
 Chemical reactions that involve the input or generation of electric currents
are called electrochemical reactions.
 Such reactions are broadly classified into two categories:
1. Production of chemical change by electrical energy i.e. the
phenomenon of electrolysis.
2. Conversion of chemical energy into electrical energy.

3. ELECTROCHEMISTRY
 Electrochemistry deals with oxidation-reduction
reactions that either produce or
utilize electrical energy
and electrochemical reactions take place in cells.
 A chemical reaction brought about by an electric
current. (Electrolytic Cells)
Electrical → Chemical
 Electric current produced by chemical reactions.
(Electrochemical Cells)
Chemical → Electrical

4. ELECTROCHEMICAL CELLS
 An electrochemical cell is a device that can
generate electrical energy from the chemical
reactions occurring in it, or use the electrical
energy supplied to it to facilitate chemical
reactions in it. These devices are capable of
converting chemical energy into electrical
energy, or vice versa.
 A common example of an electrochemical cell
is a standard 1.5-volt cell which is used to
power many electrical appliances such as TV
remotes and clocks.

5. Types of Electrochemical Cells
 The two primary types of electrochemical cells are:
1. Galvanic cells (also known as Voltaic cells)
2. Electrolytic cells
 Galvanic cells: Cells capable of generating an electric
current from the chemical reactions occurring in them
care called Galvanic cells or Voltaic cells.
 Electrolytic cells: The cells which cause chemical
reactions to occur in them when an electric current is
passed through them are called electrolytic cells.

6. DIFFERENCE B/W Galvanic
&Electrolytic Cell
Galvanic Cell / Voltaic Cell Electrolytic Cell
Chemical energy is transformed into
electrical energy in these electrochemical
cells.
Electrical energy is transformed into
chemical energy in these cells.
The redox reactions that take place in
these cells are spontaneous in nature.
An input of energy is required for the
redox reactions to proceed in these cells,
i.e. the reactions are non-spontaneous.
In these electrochemical cells, the anode
is negatively charged and the cathode is
positively charged.
These cells feature a positively charged
anode and a negatively charged
The electrons originate from the species
that undergoes oxidation.
Electrons originate from an external
source (such as a battery).

7. GALVANIC CELL
 The working of a galvanic cell is quite simple. It involves a chemical
reaction that makes the electric energy available as the end result. During
a redox reaction, a galvanic cell utilizes the energy transfer between
electrons to convert chemical energy into electric energy.
 Galvanic cell utilizes the ability to separate the flow of electrons in the
process of oxidization and reduction, causing a half reaction and
connecting each with a wire so that a path can be formed for the flow of
electrons through such wire. This flow of electrons is essentially called a
current. Such current can be made to flow through a wire to complete a
circuit and obtain its output in any device .
 A galvanic cell can be made out of any two metals. These two metals can
form the anode and the cathode if left in contact with each other.

8. SETUP OF GALVANIC CELL
 In order to create a galvanic cell, one would have to go through the
following setup.
 The cell would ideally include two electrodes. One of these electrodes,
the cathode, shall be a positively charged electrode while the other, shall
be the anode, the negatively charged electrode.
 These two electrodes shall form the two essential components of the
galvanic cell. The chemical reaction related to reduction shall take place
at the cathode while the oxidation half-reaction shall take place at the
anode.
 Any two metals can be used to create the chemical reaction.

9. SETUP OF GALVANIC CELL
 The two metals involved in the chemical reaction are zinc and copper. As
the chemical reaction takes place, Zinc would end up losing two
electrons. This will be taken up by copper to become elemental copper.
 The two metals will be placed in two separate containers and would be
connected by a conducting wire, an electric current would be formed,
which would transfer all electrons from one metal to another.
 At the same time, the two metals shall be immersed in a salt solution, say,
Zinc sulphate and Copper sulphate in this case. In this case, the two
solutions are not mixed together directly but can be joined using a bridge
or a medium. This medium shall be responsible for the transfer of ions
but also make sure that the two solutions do not come to mix with each
other.

10. SETUP OF GALVANIC CELL
 Such bridge helps in completing the
circuit for carrying the electric charge
and also makes sure that the solutions in
the containers with the metals remain
neutral and do not mix with each other.
As long as the salt bridge does not
interfere with the redox reaction, under
which oxidization and reduction are
taking place, it does not matter which
salt bridge is being used in the chemical
reaction.

11. ELECTRODE
 An electrode is strip of metal on which the reaction takes place.
 In a voltaic cell, the oxidation and reduction of metals occurs at the
electrodes.
 There are two electrodes in a voltaic cell, one in each half-cell.
The cathode is where reduction takes place and oxidation takes place at
the anode.
 An oxidation-reduction equilibrium is established between the metal and
the substances in solution.
 When electrodes are immersed in a solution containing ions of the same
metal, it is called a half-cell.
 Electrolytes are ions in solution, usually fluid, that conducts electricity
through ionic conduction.

12. WORKING OFGALVANIC CELL
 Two possible interactions can occur between the metal atoms on the
electrode and the ion solutions.
 Metal ion Mn+ from the solution may collide with the electrode, gaining
"n" electrons from it, and convert to metal atoms. This means that the
ions are reduced.
 Metal atom on the surface may lose "n" electrons to the electrode and
enter the solution as the ion Mn+ meaning that the metal atoms are
oxidized.
OIL = Oxidization is Loss (of e-)
RIG = Reduction is Gain (of e-)
 The purpose of the salt bridge is to keep the solutions electrically neutral
and allow the free flow of ions from one cell to another.

13. CELL NOTATION
 Electrons always flow from the anode to the cathode
or from the oxidation half cell to the reduction half
cell.
 In terms of Eo cell of the half reactions, the electrons
will flow from the more negative half reaction to the
more positive half reaction.
 The anode is always placed on the left side, and the
cathode is placed on the right side.
 The salt bridge is represented by double vertical lines
(||).
 The difference in the phase of an element is
represented by a single vertical line (|), while changes
in oxidation states are represented by commas (,).

14. ELECTROLYTIC CELL
 An electrolytic cell can be defined as an electrochemical device that uses
electrical energy to facilitate a non-spontaneous redox reaction.
 Electrolytic cells are electrochemical cells that can be used for the
electrolysis of certain compounds.
 For example, water can be subjected to electrolysis (with the help of an
electrolytic cell) to form gaseous oxygen and gaseous hydrogen.
 This is done by using the flow of electrons (into the reaction environment)
to overcome the activation energy barrier of the non-spontaneous redox
reaction.
 The electrolyte provides the medium for the exchange of electrons
between the cathode and the anode.
 Commonly used electrolytes in electrolytic cells include water (containing
dissolved ions) and molten sodium chloride.

15. WORKING OF ELECTROLYTIC CELL
 Molten sodium chloride (NaCl) can be subjected to electrolysis with the
help of an electrolytic cell.
 Two inert electrodes are dipped into molten sodium chloride (which
contains dissociated Na+ cations and Cl– anions).
 When an electric current is passed into the circuit, the cathode becomes
rich in electrons and develops a negative charge.
 The positively charged sodium cations are now attracted towards the
negatively charged cathode. This results in the formation of metallic
sodium at the cathode.
 Simultaneously, the chlorine atoms are attracted to the positively charged
cathode.
 This results in the formation of chlorine gas (Cl2) at the anode (which is
accompanied by the liberation of 2 electrons, finishing the circuit).

16. WORKING OF ELECTROLYTIC CELL
 Reaction at Cathode:
Na+ + e– → Na
 Reaction at Anode:
2Cl– → Cl2 + 2e–
 Cell Reaction:
2NaCl → 2Na + Cl–
 Molten sodium chloride can be subjected to
electrolysis in an electrolytic cell to generate
metallic sodium and chlorine gas as the products.

17. ELECTROLYSIS
 The use of electric current to stimulate a non-spontaneous reaction.
 Electrolysis can be used to separate a substance into its original
components/elements and it was through this process that a number of
elements have been discovered and are still produced in today's industry.
 In Electrolysis, an electric current it sent through an electrolyte and into
solution in order to stimulate the flow of ions necessary to run an
otherwise non-spontaneous reaction.
 Processes involving electrolysis include:
 electro-refining, electro-synthesis, and the chloro-alkali process.
 Example: When we electrolyze water by passing an electric current
through it, we can separate it into hydrogen and oxygen.
2H2O(l)→2H2 (g)+O2(g)

18. ELECTROLYSIS OF WATER
 The Hoffman electrolysis apparatus is filled with Na2SO4 solution
containing universal indicator and is started.
 The solution turns blue at the cathode (basic) and red at the
anode (acidic). Twice as much gas is evolved at the cathode as at
the anode. When the solutions in the two electrodes
compartments are mixed, the indicator turns green (neutral).
 Two reactions can take place at the cathode.
2H2O (l) + 2 e - → H2 (g) + 2 OH-
(aq) E° = -
0.8277 V
Na+
(aq) + e - → Na(s) E° = -2.7109 V
 Two reactions can also occur at the anode.
2 SO4
2-
(aq) → S2O8(aq) + 2e- E° = -2.05 V
2 H2O(l) → O2(g) + 4H+
(aq) + 4e- E° = -1.229 V

19. FACTORS WHETHER OR NOT ELECTROLYSIS WILL
TAKE PLACE
 There are four primary factors that determine whether or not electrolysis will take
place even if the external voltage exceeds the calculated amount:
 An over-potential or voltage excess is sometimes needed to overcome
interactions at the electrode surface. This case happens more frequently with
gases. E.g. H2 (g) requires a 1.5 V over-potential, while Pt (s) requires 0 V over-
potential.
 There might be more than one electrode reaction that occurs meaning that there
may be more than one half-reaction leaving two or more possibilities for the cell
reaction.
 The reactants may be in nonstandard conditions which means that the voltage for
the half cells may be less or more than the standard condition amount. For
Example:
Concentration of chloride ion = 5.5M not the unit activity of 1M. This means
that the reduction of chloride = 1.31V not 1.36V
 An inert electrode’s ability to electrolysis depend on the reactants in the
electrolyte solution while an active electrode can run on its own to perform the
oxidation or reduction half reaction.

20. CONDUCTANCE
 Electrical conductance measures how easily electricity flows
through electrical components for a given voltage difference. The SI unit
of conductance is siemens.
 Electrical conductance is closely related to electrical conductivity.
 Electrical conductance is a property of a particular electrical component
(like a particular wire), while conductivity is a property of the material itself
(like silver).
 Conductors are the materials or substances which allow electricity to flow
through them.
 They conduct electricity because they allow electrons to flow easily inside
them from atom to atom.
 Conductors allow the transmission of heat or light from one source to
another.

21. CONDUCTORS
 Some examples of conductors of electricity are:
 Copper
 Aluminium
 Silver
 Gold
 Graphite
 Platinum
 Water
 People
 Semiconductors
 Although semiconductors are not as good at conducting electricity as
conductors, they still have their uses. Examples of semiconductors are
Germanium (Ge) and Silicon (Si).

22. TYPES OF CONDUCTION
 Metallic conduction involves
drifting of electrons through
vacancies of conduction band
in random fashion.
 Electrolytic conduction is the
movement of free ions in
electrolyte which is either the
substance in molten state or
their aqueous solution.
S.No. Metallic conduction Electrolytic conduction
1.
The flow of electricity takes place
without decomposition of the
substance.
The flow of electricity takes place
with decomposition of the
substance.
2. There is no actual transport of matter.
Transfer of matter takes place in
the form of ions.
3.
The conduction of electricity is due to
the ow of electrons.
The conduction of electricity is
due to the flow of ions in the
solution.
4.
Metallic conductivity decreases with
increase in temperature.
Electrolytic conduction increases
with increases in temperature.
5.
Metallic conduction is a physical
process that means no new substance
is formed.
Electrolytic conduction is both
physical and chemical change.

23. CONDUCTIVITY
 Conductivity is the measure of the ease at which an electric charge or
heat can pass through a material.
 The greater the electrical conductivity within the material the higher the
current density for a given applied potential difference.
 For example, the wire conductors need to let current flow as easily as
possible. W
 Some other minerals required to restrict the flow of the current, as in the
case of the resistor.
 On the other hand, some other materials are required not to conduct
electricity as in the case of the insulators.
 The conductivity formula is the inverse of the resistivity that is:
σ=1/ρ

24. RESISTANCE
 The resistance of the conductor is proportional to its length and is
inversely proportional to its area.
R ∝ L
R ∝ 1/A
R is the resistance in ohms (Ω), L is the length in metres (m), A is the area in
square metres (m2), and where the proportional constant ρ (the Greek letter
“rho”) is known as Resistivity.

25. SPECIFIC CONDUCTANCE.
 The reciprocal of specific resistance is termed the specific conductance.
 Specific Conductance or conductivity is the conductance of a given
solution enclosed in a cell having two electrodes of unit area and are
separated by 1cm.
 It is denoted by the symbol k(kappa).
k = 1/p (R = p l/a , p = Ra/l)
k = 1 = l/Ra
Ra/l

26. Equivalent conductance & molar
conductance
 Equivalent conductance is defined as the conductance of all the ions
produced by one gram equivalent of an electrolyte in a given solution.
 It is denoted by A.
 The molar conductance is defined as the conductance of all the ions
produced by ionization of 1 g mole of an electrolyte when present in V
mL of solution. It is denoted by:
 Molar conductance μ = k ×V
 where V is the volume in mL containing 1 g mole of the electrolyte. If c is
the concentration of the solution in g mole per litre, then
 μ = k × 1000/c

27. ELECTROMOTIVE FORCE
 Electromotive force is defined as the electric potential produced by either
electrochemical cell or by changing the magnetic field.
 A generator or a battery is used for the conversion of energy from one
form to another.
 In these devices, one terminal becomes positively charged while the
other becomes negatively charged.
 An electromotive force is a work done on a unit electric charge.
 Maximum potential difference between two electrodes of a cell.
 Also called cell potential and denoted by Ecell
Ecell = Ered,cathode – Ered, anode
 Positive cell potential = Spontaneous processes
 Negative cell potential = Non-spontaneous processes

28. APPLICATIONS OF EMF
MEASUREMENTS
 Measurement of the standard Emf of the cell, E°Cell, enables one to
evaluate the equilibrium constant for the electrode reaction.
 To evaluate the solubility product of sparingly soluble salt such as AgCI.
 Determination of the pH of solution using hydrogen electrode, saturated
calomel electrode (SCE) or any other reference electrode.
 Determination of Activity Coefficients.
 To detect the end point of a titration by measuring the emf of a cell
consisting of an indicator electrode (electrode, whose potential depends
on the concentration of the reactant ions) and a reference electrode.
 Determination of transport number of ions.

29. ELECTRODE
 An electrode is a solid electric conductor that carries electric
current into non-metallic solids, or liquids, or gases, or
plasmas, or vacuums. Electrodes are typically good electric
conductors, but they need not be metals.
 An early version of an electrode was the electrophore which
was used to study static electricity. It was invented by Johan
Wilcke.
 The word was coined by William Whewell at the request of
the scientist Michael Faraday from two Greek words: elektron,
meaning amber, and hodos, a way.
 Electrodes are commonly used in electrochemical
cells, semiconductors like diodes, and in medical devices.

30. TYPES OF ELECTRODES
 There are mainly two types of electrodes namely reactive and inert
electrodes.
 An inert type does not participate in any reaction while reactive types
participate actively in reactions.
 Some commonly used inert electrodes include platinum, gold,
graphite(carbon), and rhodium.
 Some reactive electrodes include zinc, copper, lead, and silver.
 Electrodes are vital components in electrochemical cells as they transport
produced electrons from one half-cell to another, which results in the
production of an electrical charge.

31. USES OF ELECTRODES
 The main use of electrodes is to generate electrical current and pass it
through non-metal objects to basically alter them in several ways.
Electrodes are also used to measure conductivity.
 Electrodes are used in different battery types, electroplating and
electrolysis, welding, cathodic protection, membrane electrode assembly,
for chemical analysis, and Taser electroshock weapon.
 In the medical field, electrodes are also used in ECG, ECT, EEG, and
defibrillator.
 Electrodes are further used for electrophysiology techniques in
biomedical research.

32. TYPES OF ELECTRODES
 There are four types of electrodes:
 Gas electrodes
 Metal–sparingly soluble metal salt
electrodes
 Metal – metal ion electrodes
 Redox electrodes

33. METAL – METAL ION ELECTRODES
 This is the simplest type of electrode where metal is
dipped into a solution of as ions. These electrodes
consist of a pure metal (M) in contact with a solution
of its cation (Mn+). For example, a silver rod
immersed in a solution of Ag+ ions or copper rod in
a copper sulphate solution. Examples are silver in
silver nitrate solution, copper in a copper sulphate
solution. The silver-silver ion electrode is represented
as, Ag│Ag+ and the copper-copper ion electrode as
Cu│Cu2+.
 The electrode reactions are Ag+ + e– ↔ Ag and
Cu2+ + 2e– ↔ Cu

34. HYDROGEN ELECTRODE
 The Standard Hydrogen Electrode is often
abbreviated to SHE, and its standard electrode
potential is declared to be 0 at a temperature of
298K.
 It acts as a reference for comparison with any other
electrode.
 The redox half cell of SHE is where the following
reaction takes place:
2H+ (aq) + 2e– → H2 (g)
 The reaction given above generally takes place on
a platinum electrode.
 As an indicator electrode for calculating pH values in
early studies.

35. REDOX ELECTRODES
 In these electrodes, inert metal like Pt rod is dipped in
a solution containing ions of an active metal in two
different oxidation states.
Pt l Fe+2 , Fe3+
Fe +3 + e - Fe +2
 Platinum wire immersed in a mixture of ferrous and
ferric ion is a typical example .

36. CALOMEL ELECTRODE
 Calomel electrode is a type of half cell in which the
electrode is mercury coated with calomel (Hg2Cl2) and the
electrolyte is a solution of potassium chloride and saturated
calomel. In the calomel half cell the overall reaction is
 Hg2Cl2(s) + 2e- →← 2Hg(l) + 2Cl-
 SCE is used in pH measurement, cyclic voltammetry and
general aqueous electrochemistry.
 This electrode and the silver/silver chloride reference
electrode work in the same way. In both electrodes, the
activity of the metal ion is fixed by the solubility of the
metal salt.
 The calomel electrode contains mercury, which poses much
greater health hazards than the silver metal used in the
Ag/AgCl electrode.

37. ELECTRODE POTENTIAL
 When a metal is placed in a solution of its ions, the metal acquires either
a positive or negative charge with respect to the solution. On account of
this, a definite potential difference is developed between the metal and
the solution. This potential difference is called electrode potential.
 For example, when a plate of zinc is placed in a solution having Zn2+ ions,
it becomes negatively charged with respect to solution and thus a
potential difference is set up between zinc plate and the solution. This
potential difference is termed the electrode potential of zinc.
 The potential difference developed between metal electrode and the
solution of its ions of unit molarity (1M) at 25°C (298 K) is called standard
electrode potential.

38. ELECTROLYTE
 An electrolyte is a substance that produces an electrically conducting
solution when dissolved in a polar solvent, such as water. The dissolved
electrolyte separates into cations and anions, which disperse uniformly
through the solvent. Electrically, such a solution is neutral.
 The most familiar electrolytes are acids, bases, and salts, which ionize
when dissolved in such solvents as water or alcohol.
 Many salts, such as sodium chloride, behave as electrolytes when melted
in the absence of any solvent.
 Silver iodide electrolyte even in the solid state.
 Potassium, sodium, and calcium are electrolytes, or salts that help
conduct electric current in the body.

39. FACTORS AFFECTING ELECTROLYTIC
CONDUCTANCE
1. Concentration of ions
The sole reason for the conductivity of electrolytes is the ions present in
them. The conductivity of electrolytes increases with increase in the
concentration of ions as there will be more charge carriers if the concentration
of ions is more and hence the conductivity of electrolytes will be high.
2. Nature of electrolyte
Electrolytic conduction is significantly affected by the nature of electrolytes. The
degree of dissociation of electrolytes determines the concentration of ions in the
solution and hence the conductivity of electrolytes. Substances such
as CH3COOH, with a small degree of separation, will have less number of ions in
the solution and hence their conductivity will also below, and these are called
weak electrolytes. Strong electrolytes such as KNO3 have a high degree of
dissociation and hence their solutions have a high concentration of ions, and so
they are good electrolytic conductance.

40. FACTORS AFFECTING ELECTROLYTIC
CONDUCTANCE
3. Temperature
Temperature affects the degree to which an electrolyte gets dissolved in
solution. It has been seen that higher temperature enhances the solubility of
electrolytes and hence the concentration of ions which results in increased
electrolytic conduction.
4. Ionic size
There is an inverse relationship observed, which means the larger the size of
ion the lesser the conductance.
5. Nature of solvent and viscosity
In the case where the nature of the solvent has greater polarity then there is
the presence of higher conductance. An inversely proportional relationship
has been observed for viscosity and electrolytic conduction. When the
viscosity of the solvent is high then the conductance is a ected as it gets
reduced.

41. DEBYE-HUCKEL-ONSAGER
EQUATION
 Equivalent conductance increases with dilution in the
case of weak electrolytes.
 In case of weak electrolytes increase in conduction is
due to increase of dissociation of weak electrolyte.
 But strong electrolytes are completely dissociated even
at moderate concentration.
 In this theory we have to explore those factors which
increase the conductance of strong electrolytes.
Λ = Λ0 – A√C

42. KOHLRAUSCH’S LAW
 Kohlrausch’s law states that the equivalent conductivity of an electrolyte at
infinite dilution is equal to the sum of the conductances of the anions and
cations.
 Molar conductivity of a solution at a given concentration is the
conductance of the volume of solution containing one mole of electrolyte
kept between two electrodes with the unit area of cross-section and
distance of unit length.
 The molar conductivity of a solution increases with the decrease in
concentration. This increase in molar conductivity is because of the
increase in the total volume containing one mole of the electrolyte. When
the concentration of the electrolyte approaches zero, the molar
conductivity is known as limiting molar conductivity, Ëm°.

43. KOHLRAUSCH’S LAW
 Kohlrausch observed certain regularities while comparing the values of
limiting molar conductivities of some strong electrolytes.
 On the basis of his observations, Kohlrausch proposed “limiting molar
conductivity of an electrolyte can be represented as the sum of the
individual contributions of the anions and cations of the electrolyte”.
 This law is popularly known as Kohlrausch law of independent migration
of ions.
 For example, limiting molar conductivity, Ëm° of sodium chloride can be
determined with the knowledge of limiting molar conductivities of sodium
ion and chloride ion.

44. KOHLRAUSCH’S LAW
Uses of Kohlrausch’s law
 Calculation of Degree of dissociation
 Calculation of solubility of sparingly soluble salt
 Calculation of Dissociation Constant for week electrolytes
 Calculation of Molar Conductivity for week electrolytes at infinite dilution

45. NERNST EQUATION
 The Nernst equation provides a relation between the cell potential of an
electrochemical cell, the standard cell potential, temperature, and the reaction
quotient.
 Even under non-standard conditions, the cell potentials of electrochemical cells can be
determined with the help of the Nernst equation.
 The equation was introduced by a German chemist named Walther Hermann Nernst.
Ecell = E0 – [RT/nF] ln Q
 Ecell = cell potential of the cell
 E0 = cell potential under standard conditions
 R = universal gas constant
 T = temperature
 n = number of electrons transferred in the redox reaction
 F = Faraday constant
 Q = reaction quotient

46. NERNST EQUATION APPLICATIONS
The Nernst equation can be used to calculate:
 Single electrode reduction or oxidation potential at any conditions
 Standard electrode potentials
 Comparing the relative ability as a reductive or oxidative agent.
 Finding the feasibility of the combination of such single electrodes to
produce electric potential.
 Emf of an electrochemical cell
 Unknown ionic concentrations
 The pH of solutions and solubility of sparingly soluble salts can be
measured with the help of the Nernst equation.

47. Applications of electrochemistry
 Electrochemistry has a number of different uses, particularly in industry.
 The principles of cells are used to make electrical batteries.
 In science and technology, a battery is a device that stores chemical energy
and makes it available in an electrical form.
 Batteries are made of electrochemical devices such as one or more galvanic
cells or fuel cells. Batteries have many uses including in:
 torches
 electrical appliances such as cellphones (long-life alkaline batteries)
 digital cameras (lithium batteries)
 hearing aids (silver-oxide batteries)
 digital watches (mercury/silver-oxide batteries)
 military applications (thermal batteries)

48. Applications of electrochemistry
 The electrolytic cell can be used for electroplating.
 The process of coating an electrically conductive
object with a thin layer of metal using an electrical
current.
 Electroplating occurs when an electrically
conductive object is coated with a layer of metal
using electrical current.
 Sometimes, electroplating is used to give a metal
particular properties or for aesthetic reasons:
 corrosion protection
 abrasion and wear resistance
 the production of jewellery
a) An electroplated piece of aluminium
artwork and b) a wax stool electroplated
in copper.

49. Applications of electrochemistry
 Electrochemical processes are used in many ways and their use is likely to
increase because they can replace polluting chemical situations with
nonpolluting electrochemical ones.
 All technologically important metals, except iron and steel, are either
obtained or refined by electrochemical processes.
 For example, aluminum, titanium, alkaline earth, and alkali metals are
obtained by electrodeposition from molten salts, and copper is refined
by electrolysis in aqueous copper sulfate solutions.
 The energy of chemical reactions is converted into electrical energy in
fuel cells. In these, the fuel (e.g., hydrogen, hydrazine) is fed continuously
to one electrode, while oxygen from the air is reacting at the other one.
The efficiency of energy conversion in fuel cells is more than twice that
attainable by conventional means—for example, by means of
internal combustion.

50. Applications of electrochemistry
 In analytical chemistry, most modern automated instrumental analysis is
based on electrode processes—for example, potentiometry, used to
measure ionization constant.
 In biology the idea that many biological processes, from blood clotting to
the transfer of nerve impulses, are electrochemical in nature continues to
spread. The biological conversion of the chemical energy of food
to mechanical energy takes place at an efficiency so high that it is difficult
to explain without electrochemical mechanisms. Intensive research is
developing in various directions in bioelectrochemistry.
 Certain diabetes blood sugar meters measure the amount of glucose in
the blood through its redox potential.
 Electrochemistry has also important applications in the food industry, like
the assessment of food/package interactions, the analysis of milk
composition, the characterization and the determination of the freezing
end-point of ice-cream mixes, the determination of free acidity in olive
oil.

51. LATEST RESEARCH TOPICS
 Artificial dual solid-electrolyte interfaces based on in situ organothiol
transformation in lithium sulfur battery
 Laser-driven growth of structurally defined transition metal oxide
nanocrystals on carbon nitride photoelectrodes in milliseconds
 Electrosynthesis of 1,4-bis(diphenylphosphanyl) tetrasulfide via sulfur
radical addition as cathode material for rechargeable lithium battery
 Chemical analysis and computed tomography of metallic inclusions in
Roman glass to unveil ancient coloring methods
 Boosting oxygen reduction activity and enhancing stability through
structural transformation of layered lithium manganese oxide
 Comparison of electrochemical impedance spectra for electrolyte-
supported solid oxide fuel cells (SOFCs) and protonic ceramic fuel cells
(PCFCs)