Electrochemistry
study of how electricity produces
chemical reactions and chemical
reactions produces electricity
• Involves redox reactions
• Electrochemical cell: any device which converts
chemical energy into electrical energy or vs.
Electrochemistry
study of how electricity produces chemical
reactions and chemical reactions produces
electricity
• Involves redox reactions
• Electrochemical cell:
any device which
converts chemical energy into electrical energy
or vs.
Introduction
Electrochemistry is a branch of chemistry
that studies chemical reactions which take
place in a solution at the interface of an
electron conductor (a metal or a
semiconductors) and an ionic conductor (the
electolyte), and which involve electron
transfer between the electrode and the
electrolyte or species in solution.
Alessandro Volta's discovery, in 1793, that
electricity could be produced by placing two
dissimilar metals on opposite sides of a
moistened paper.
In 1800, Nicholson and Carlisle, using
Volta’s primitive battery as a source,
showed that an electric current could
decompose water into oxygen and hydrogen.
By 1812, the Swedish chemist Berzelius
could propose that all atoms are electrified,
hydrogen and the metals being positive, the
nonmetals negative.
Humphry Davy prepared the first elemental
sodium by electrolysis of a sodium hydroxide melt.
Michael Faraday, to show that there is a direct
relation between the amount of electric charge
passed through the solution and the quantity of
electrolysis products
Chemical reactions where electrons are transferred
between molecules are called oxidation/reduction
(redox) reactions. In general, electrochemistry
deals with situations where oxidation and reduction
reactions are separated in space or time, connected
by an external electric circuit to understand each
process.
Electron Transfer Reactions
Electron transfer reactions are oxidation-reduction or redox
reactions.
Results in the generation of an electric current (electricity)
or be caused by imposing an electric current.
Therefore, this field of chemistry is often called
ELECTROCHEMISTRY.
Terminology for Redox Reactions
OXIDATION :loss of electron(s) by a species; increase in
oxidation number; increase in oxygen.
REDUCTION: Gain of electron(s); decrease in oxidation
number; decrease in oxygen; increase in hydrogen.
OXIDIZING AGENT: Electron acceptor; species is
reduced.
REDUCING AGENT: Eelectron donor; species is oxidized.
OXIDATION-REDUCTION REACTIONS
Direct Redox Reaction
Oxidizing and reducing agents in direct contact.
Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)
Why Study Electrochemistry?
• Batteries
• Corrosion
• Industrial production of
chemicals such as Cl2,
NaOH, F2 and Al
• Biological redox
reactions
The heme group
Classification of Conductors
These may be divided into three main categories;
they are:
(I) gaseous (II) metallic or
electronic (III) electrolytic.
Gases conduct electricity with difficulty and only
under the influence of high potentials or if exposed
to the action of certain radiations.
Metallic or electronic conductors : Conductors
which transfer electric current by transfer of
electrons, without transfer of any matter, are known
as metallic or electronic conductors. Metals such
as copper, silver, aluminum, etc., non-metals like
carbon (graphite - an allotropic form of carbon) and
various alloys belong to this class.
Electrolytic conductors : (a) Conductors
like aqueous solutions of acids, bases and
salts in which the flow of electric current is
accompanied by chemical decomposition are
known as electrolytic conductors.
b)The substances whose aqueous solutions
do not conduct electric current are called
non-electrolytes. Solutions of cane sugar,
glycerine, alcohol, etc., are examples of nonelectrolytes
Fig. 1. Illustration of electrochemical terms 18 aug
Mechanisam of electrolytic conduction and electrolysis
H+ + 2e2Cl- - 2e-
2NaCl(aq) + 2H2O(l)
H2 (hydrogen gas at the (-)cathode).
Cl2 (chlorine gas at the (+)anode).
The overall reaction is
2Na+(aq) + 2OH-(aq) + Cl2(g)+ H2(g)
Electrolysis of sodium chloride solution
NaCl ↔ Na+ + ClH2O ↔ H+ + OHAt cathode
H+ + e- → H
2H → H2
At Anode
Cl- → Cl + e2Cl → Cl2
Electrolysis of copper sulphate solution using platinum electrodes
CuSO4 ↔ Cu2+ + SO42H2O ↔ H+ + OHAt cathode
At Anode
Cu2+ + 2e- → Cu
2OH- → H2O + O + 2eO + O→O2
The laws, which govern the deposition of substances
(In the form of ions) on electrodes during the process
of electrolysis, is called Faraday's laws of electrolysis.
These laws given by Michael Faraday in 1833.
Faraday's first law: It states that, the mass of any
substance deposited or liberated at any electrode is
directly proportional to the quantity of electricity
passed.
WαQ
W = Mass of ions liberated in gm,
Q = Quantity of electricity passed in Coulombs
= Current in Amperes ( i ) × Time in second (t)
Wαit
W=Zit
Where, Z = constant, known as electrochemical equivalent
(ECE) of the ion deposited
• Where Z is the consant known as the
Electrochemical
equivalent
of
the
substance (electrolyte).
• If I= 1 ampere and t = 1 scond, then m= Z
• Thus, the electrochemical equivalent is the
amount of a substance deposited by 1
ampere current passing for 1 second (I.e.,
one coulomb)
• The Electrical unit Faraday
It has been found experimentally that the
quantity of electricity required to liberate one
gram-equivalent of a substance is 96,500
coulombs. This quantity of electricity is known as
Faraday and is denoted by the symbol F.
• it is obvious that the quantity of electricity
needed to deposit 1 mole of the substance
is given by the expression.
Quantity of electricity = n x F
Where n is the valency of its ion. Thus the
quantity of electricity required to discharge.
one mole of Ag+ = 1 x F = 1F
one mole of Cu2+ = 2 x F = 2F
one mole of A13+ = 3 x F = 3F
we can represent the reaction on the
cathode as:
Ag+ + e = Ag
Cu2+ + 2e =Cu
A13+ + 3e = A1
• Moles of electrons required to discharge one
mole of ions Ag+, Cu2+ and a13+ is one, two
and three respectively. Therefore it means
that the quantity of electricity in one
Faraday is one mole of electrons. Now we
can say that.
1 Faraday = 96,500 coulombs = 1 Mole electrons
(1) Importance of first law
(2) With the help of first law of electrolysis we are
able to calculate:
(3) (1) the value of electrochemical equivalents of
different substances.
(4) (2) the masses of different substances
produced by passing a known quantity of
electricity through their solutions.
• Example No. (1) 0.1978 g of copper is
deposited by a current of 0.2 ampere
in
50
minutes.
What
is
the
electrochemical equivalent of copper?
• Example No. (2) what current strength
in amperes will be required to liberate
10 g of iodine from potassium iodide
solution in one hour ?
Faraday's second law: It states that, when the same
quantity of electricity is passed through different
electrolytes, the masses of different ions liberated at the
electrodes are directly proportional to their chemical
equivalents (Equivalent weights).

WαE
E α Z or E = FZ or E = 96500 × Z
Faraday's law for gaseous electrolytic product for the gases,
we use V = It Ve/96500
Where, V = Volume of gas evolved at S.T.P. at an electrode
Ve = Equivalent volume = Volume of gas evolved at an
electrode at S.T.P. by 1 Faraday charge
Importance of the second law
• The second law of electrolysis helps
to calculate:
• (1) the equivalent weights of metals
• (2) the unit of electric charge
• (3) the Avogadro’s number
•
conductance and its measurement
Ohm’s law
Metallic as well as electrolytic conductors obey Ohm’s law
which states the strength of current (I) flowing through a
conductor is directly proportional difference (V) applied
across the conductor and is inversely proportional to the
resistance (R ) of the conductor
I = V/R
R - Resistance in V/A = Ω (Ohm)
V - Voltage or potential difference in Volts, V
I - Current in Amperes, A
If a material has a resistance of 1 Ω, it means that when
applying a potential difference of 1 V, the current in the
material is 1 A.
For metals:
Ohm’s Law
R = V/I
R: resistance
Dimension: Ohm, 
Conductance is the ability of a material to pass electrons
C=1/R
Specific conductance or conductivity
The resistance of any conductor varies directly as its length
(l) and inversely as its cross sectional area (a), i.e.,
R α l/a or R = ρ l/a , Here ρ = specific resistance
If l = 1 cm and a = 1 cm2, then
R=ρ
Κ= 1/ρ, Κ = kappa - the specific conductance
ρ = a/l. R or 1/ρ = 1/a.1/R
K = 1/a×C (1/z = cell constant)
Specific conductance = cell constant x Conductance
The unit of specific conductance is ohm-1 cm-1.
Specific conductance or conductivity
-
+
-
anode
Solution
Cathode
+
1c
m
1c
m
Representation of specific conductance
Specific conductance depend on the number
of ions present in unit volume (1 ml ) of
solution
Equivalent conductance (/\)
To understand the meaning of equivalent conductance,
imagine a rectangular trough with two opposite sides
made of metallic conductor (acting as electrodes)
exactly 1 cm apart, If 1 cm3 (1 mL) solution containing
1 gram equivalent of an electrolyte is places in this
container is measured.
/\ = KV
In case, if the concentration of the solution is c g
equivalent per liter, then the volume containing 1 g
equivalent of the electrolyte will be 1000/C.
So equivalent conductance
/\ k 1000/c
/\ = k × 1000/N
Where N = normality
The unit of equivalent conductance is ohm-1 cm-2 equi-1.
One of the factors on which the conductance of an
electrolytic solution depends is the concentration of
the solution. In order to obtain comparable results for
different electrolytes, it is necessary to take
equivalent conductances.
Equivalent conductance is defined as the
conductance of all the ions produced by one gram
equivalent of an electrolyte in a given solution.
1 cc
m
1c
Representation of Equivalent conductance
1 cm
Molar conductance
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
Λ m = 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
Λ m = k × 1000/c
It units are ohm-1 cm2 mol-1.
Equivalent conductance = (Molar conductance)/n
Where
n = (Molecular mass) / (Equivalent mass)
Effect of dilution on equivalent conductance
Conductance’s of electrolytes of different type
Kohlrausch’s law of independent ionic mobilities
At time infinite dilution (m) , the molar
conductivity of an electrolyte can be expressed as the
sum of the contributions from its individual ions
Λ∞m = v+ λ∞ + + v- λ∞v+ and v- are the number of cations and anions
per formula unit of electrolyte respectively and,
λ∞+ and λ∞- are the molar conductivities of the
cation and anion at infinite dilution respectively
24 Aug
Applications of Kohlrausch's law
Determination of Λ∞m for weak electrolytes
Determination of the degree of ionization of a weak
electrolyte
Determination of the ionization constant of a weak
electrolyte
Determination of the solubility of a sparingly soluble salt
Charging a Battery
When you charge a battery, you
are
forcing
the
electrons
backwards (from the + to the -).
To do this, you will need a higher
voltage
backwards
than
forwards.
This is why the
ammeter in your car often goes
slightly higher while your battery
is charging, and then returns to
normal.
• In your car, the
battery charger is
called an alternator.
If you have a dead
battery, it could be
the battery needs to
be replaced OR the
alternator
is
not
charging the battery
properly. 24 (aug)
evn
H2 as a Fuel
Cars can use electricity generated by H2/O2 fuel cells.
H2 carried in tanks or generated from hydrocarbons