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Electrolysis and Electrolytic Cell
Deniell Cell
Electrochemical Series
Electrode Potential
Arrhenius Theory of Electrolytic Dissociation
Faradays Laws of Electrolysis
Nernst Equation
Electrolytic Conductance
Kohlarusch’s Law
Batteries
Concentration Cells
Applications of Electrolysis
Commercial Production of Chemicals
Solved Examples on Electrochemistry
Electrochemistry is related to our biological system also. The
transmission of sensory signals in our body from brain to other body
parts takes place through neurons. This transmission of signals in our
body has electrochemical origin. Therefore, electrochemistry is not only
limited up to chemistry but its braches extend to physics and biology
also. In this chapter we will explore the concepts of electrochemistry in
more details under following subtopics:
Substances around us can be divided
into two classes based on their ability of conduct electricity:
Conductors and Non Conductors
Non-Conductors: Those substances which do not allow electric current to
pass through them are called non-conductors or insulators. Example: - wood,
plastic glass, rubber etc.
Conductors: Those substances which allow electric current to flow through
tem are called conductors. Examples: Copper, Iron, Gold, Silver, Graphite,
salt solution etc.
Conductors can further be divided into two groups
Metallic Conductors: These conductors conduct electricity or electric current
by movement of electrons without undergoing any chemical change during
the process. These conduct electricity in both solid as well as molten state.
Example: All the metals and Graphite
Electrolytes: Those substances which conduct electricity only when they are
present in aqueous solution and not in solid form are called electrolytes.
For a substance to conduct electricity; it must either have free
electrons or ions which carry electricity with them. Electrolytes
neither have free electrons nor free ion in solid state although
they are ionic compound. This is because the oppositely charged
ions are held together by strong electrostatic attraction and are
not free to move. But when they are dissolved in water, the two
ions split up and become free to move in solution and now they
are free to conduct electricity. Examples of electrolytes are: NaCl,
KCl, Na2SO4 etc
Non-ionic compound or covalent compounds do not conduct electricity in aqueous
solution and hence they are called non-electrolytes. Examples of non- electrolytes
are: Urea, Glucose, Sugar etc.
Electrolytes can further be divided into strong and weak electrolytes:
Strong Electrolytes are those electrolytes which dissociate completely in
aqueous solution to give constituent ions. For example: Inorganic salts like
NaCl, KCl, Strong Acid like HCl, H2SO4, Strong bases like NaOH, KOH etc.
Weak Electrolytes are those electrolytes which partially dissociate in
aqueous solution to give constituent ions. For example: weak acid like
CH3COOH and Weak bases like NH3.
Comparison of Electrolytic and Metallic Conduction
S.No
Metallic Conduction
Electrolytic Conduction
1
Electric current flows by
movement of electrons.
Electric current flows by
movement of ions.
2
No chemical change
occurs.
Ions are oxidized or reduced at
the electrodes.
3
It does not involve the
transfer of any matter.
It involves transfer of matter in
the form of ions.
4
Ohm's law is followed.
Ohm's low is followed.
5
Resistance increases with
increase of temperature.
Resistance decreases with
increase of temperature.
6
Faraday law is not
followed.
6Faraday law is followed.
Pure water does not conduct electricity. Conductivity of tap water
is due to dissolved salts and minerals.
Conductivity of metallic conductors increases while that of
electrolytes decreases with increase in temperature.
Common salt which we use daily in our food is a strong
electrolyte.
Electrolytes are very important in our biological system. Our body
requires electrolytes for functioning of nervous system and other
life processes.
Semiconductors: Semiconductors are those substances whose conductivity is
intermediate to those of conductors and insulators i.e. Conductivity is more than
insulator and less then
conductors.
Question 1: Substances which do not conduct electricity in solid state but in
question solution is known as..
a. insulators
b. conductors
c. electrolytes
d. strong acids
Question 2: Which of the following substances is not an electrolyte?
a. Common salt
b. Sulphuric Acid
c. Acetic Acid
d. Glucose
Question 3: Which of the following substances is not a non-electrolyte?
a. Gold
b. Mercury
c. Graphite
d.Ammonium Chloride
Question 4: Which of the following compounds will not conduct electricity in its
aqueous solution?
a. Carbontetrachloride
b. Silver chloride
c. Sodium acetate
d. Sulphuric Acid
Question 5: Which of the following electrolytes is not strong one?
a. HCl
b. NaCl
c. H2SO4
d. CH3COOH
Q.1
Q.2
Q.3
Q.4
Q.5
c
d
d
a
d
Concentration Cells The cells in which EMF arises...
Molar Conductivity
Electrolytic Conductance The conductance is the...
Commercial Production of Chemicals
Commercial Production of Chemicals As we discussed...
Arrhenius Theory of Electrolytic Dissociation
Arrhenius Theory of Electrolytic dissociation In...
Nernst Equation
Electrode and Cell Potentials EMF of A Galvanic...
Electrode Potential
Electrode Potential What is electrode potential?...
Electrochemical Series
Electrochemical Series By measuring the potentials...
Electrolysis and Electrochemical Cells...
KOHLRAUSCH S LAW
Kohlrausch’s law Kohlrausch’s law...
Batteries
Batteries Any cell or Battery (more than one cells...
Electrolytic cells: In this type of cells electrical energy is used to
carry out a non-spontaneous reaction.
In simple words,one can say that in galvanic cells, chemical energy is
converted into electrical energy, while in electrolytic cell electrical energy
is converted into chemical energy.
Difference in Electrolytic Cell and Galvanic Cell
Electrolytic cell
Galvanic cell
Electrical energy is converted into
chemical energy.
Chemical energy is converted into
electrical energy.
Anode positive electrode. Cathode
negative electrode.
Anode negative electrode. Cathode
positive electrode.
Ions are discharged on both the
electrodes.
Ions are discharged only on the
cathode.
If the electrodes are inert,
concentration of the electrolyte
decreases when the electric current
is circulated.
Concentration of the anodic halfcell increases while that of cathodic
half-cell decreases when the two
electrodes are joined by a wire.
Both the electrodes can be fitted in
the same compartment.
The electrodes are fitted in different
compartments.
Electrolytic cell
An electrolytic cell is an arrangement in which electricity is conducted
through a solution or a molten salt by the movement of ions. It can be
said that electrical energy is converted to chemical energy.
The principles of electrolytic conduction are best illustrated by reference
to an electrolytic cell such as that shown in figure for the electrolysis of
molten NaCl between inert electrodes.
In order to pass the current through an electrolytic conductor (aqueous
solution or fused electrolyte), two rods or plates (metallic conductors) are
always needed which are connected with the terminals of a battery.
These rods or plates are known as electrodes. The electrode through
which the current enters the electrolytic solution is called the anode
(positive electrode) with the electrode through which the current leaves
the electrolytic solution is known as cathode (negative electrode).
The entire assembly except that of the external battery is known as the
cell.The electrons are received from the negative end of the external
battery by the negative electrode of the cell.
These are used up in the reduction reaction at this electrode.
The numbers of electrons received at the negative electrode are given
back to the positive end of the external battery from the positive electrode
of the cell where electrons are released as a result of oxidation reaction.
Within the cell, the current is carried by the movements of ions; cations
towards the negative electrode (cathode) and anions towards the positive
electrode (anode).
This movement of ions gives rise to what is known as the electrolytic
conduction.
Let us now take a situation where more than one type of cation is
present. The ability of cation to move towards the negative electrode and
get reduced depends upon the size, mass, positive charge, negative
charge etc.
It is therefore not possible to predict, qualitatively, the order of reduction
of cations, as one factor might enhance it while the another factor might
hamper it.
The only way we can predict this is by giving a quantitative value based
on the cumulative effect of all the factors responsible for a cation ability
to get reduced.
This quantitative value is called the standard reduction potential (SRP). A
cation with a higher value of SRP would get reduced in preference to a
cation with a lower value of SRP.
The standard reduction potential values at 25°C are given below for some
reactions
Standard Reduction Potentials at 25° C
Reducation half reaction
E°, V
F2 + 2e– → 2F–
2.87
S2O82-+2e- → 2SO42-
2.0
Co3+ + e– → Co+2
H2O2 + 2H+ + 2e– → 2H2O
MnO4- + 4H+ + 3e– → MnO2 +
2H2O
PbO2 + 4H+ + SO42+ 2e– → PbSO4
+ 2H2O
Ce4+ + e– → Ce3+
MnO4- + 8H+ + 5e– → Mn+2 +
4H2O
1.82
1.77
Au3+
+
3e–
→ Au
Cl2 + 2e– → 2Cl–
Cr2O72- + 14H+ + 6e– → 2Cr+3 +
7H2O
Ti3+ + 2e– → Ti+
MnO2 + 4H+ + 2e– → Mn2+ +
1.70
Reducation half
reaction
AgCl + e– → Ag + Cl–
PdI42- + 2e– → Pd +
4I–
Cu2+ + e– → Cu+
Sn4+ + 2e– → Sn2+
Ag(S2O3)23-+e– → Ag+
2S2O32-
E°, V
0.222
0.18
0.15
0.13
0.017
1.70
2H+ + 2e– → H2
0.000
1.70
Pb2+ + 2e– → Pb
– 0.126
1.51
Sn2+ + 2e– → Sn
– 0.14
1.36
2CuO + H2O + 2e– →
– 0.15
Cu2O + 2OH–
AgI + e– → Ag + I–
– 0.151
1.33
CuI + e– → Cu + I–
– 0.17
1.26
1.23
Ni2+ + 2e– → Ni
Co2+ + 2e– → Co
– 0.25
– 0.28
1.50
2H2O
O2 +
4H+
+
4e–
→ 2H2O
1.229
2IO3- + 12H+ + 10e– → I2 + 6H2O 1.20
Br2 + 2e– → 2Br–
1.09
AuCl4- + 3e– → Au + 4Cl–
OCl– + H2O + 2e– → Cl– + 2OH–
Pd2+ + 2e– → Pd
2Hg2+ + 2e– → Hg22+
1.00
0.94
0.92
0.92
Cu2+ + I– + e– → CuI
0.85
Ag+ + e– → Ag
Fe3+ + e– → Fe2+
O2 + 2H+ + 2e– → H2O2
Cu2+ + Cl– + e– → CuCl
I2 + 2e–- → 2I–
Cu+ + e– → Cu
Cu2+ + 2e– → Cu
Hg2Cl2 + 2e– → 2Hg + 2Cl–
Hg2Cl2 + 2e– → 2Hg + 2Cl– (satd
KCl)
0.799
0.771
0.69
0.566
0.535
0.52
0.34
0.270
0.244
PbSO4 + 2e– → Pb +
SO42Ti+ + e– → Ti
Cu2O + H2O + 2e–
→ 2Cu + 2OH–
Cd2+ + 2e– → Cd
Fe2+ + 2e– → Fe
Cr3+ + 3e– → Cr
Zn2+ + 2e– → Zn
2H2O + 2e– → H2 +
2OH–
Mn2+ + 2e– → Mn
H2 + 2e– → 2H–
Mg2+ + 2e– → Mg
Ce3+ + 3e– → Ce
Na+ + e– → Na
Ca2+ + 2e– → Ca
K+ + e– → K
Li+ + e– → Li
– 0.31
– 0.336
– 0.34
–
–
–
–
0.403
0.44
0.74
0.7628
– 0.828
– 1.18
– 2.25
– 2.37
– 2.48
– 2.713
– 2.87
– 2.93
–3.03
Ge2+ + 2e– → Ge
0.23
For anions the ability to get oxidized is given by the standard oxidation
potential which is the reverse of the standard reduction potential of a
molecule to form the anion.
According to the table, if we take an aqueous solution of NaCl and do its
electrolysis, H+ would be reduced to H2 gas (the H+ ions are present since
the solution is aqueous) at the cathode, while Cl– ions would be oxidised
to Cl2 gas at the anode.
Though what we have stated just now is used in solving problem, it is not
always valid. This is because the ability of a cation to be reduced or an
anion to be oxidized not only depends on their SRP’s, but also depends
on their concentrations. This means that it is possible to reduce a cation
in preference to another cation even though the SRP of the former may
be less than that of the latter, just by adjusting concentrations.
A most remarkable feature of oxidation - reduction reactions is that they
can be carried out with the reactants separated in space and linked only
by an electrical connection. That is to say, chemical energy is converted
to electrical energy. Consider figure ,a representation of a galvanic cell
which involves the reaction between metallic
zinc and cupric ion:
The cell consists of two beakers, one of which contains a solution of Cu2+
and a copper rod, the other a Zn2+ solution and a zinc rod. A connection
is made between the two solutions by means of a “salt bridge”, a tube
containing a solution of an electrolyte, generally NH4NO3 or KCl.
Flow of the solution from the salt bridge is prevented either by plugging
the ends of the bridge with glass wool, or by using a salt dissolved in a
gelatinous material as the bridge electrolyte.
When the two metallic rods are connected through an ammeter, a
deflection is observed in ammeter which is an
evidence that a chemical reaction is occurring.
The zinc rod starts to dissolve, and copper is
deposited on the copper rod. The solution of Zn2+
becomes more concentrated, and the solution of Cu2+ becomes more
dilute.
The ammeter indicates that electrons are flowing from the Zinc rod to the
copper rod. This activity is continuous as long as the electrical
connection and the salt bridge are maintained, and visible amounts of
reactants remain.
Now let us analyze what happens in each beaker more carefully. We note
that electrons flow from the Zinc rod through the external circuit, and
that Zinc ions are produced as the Zinc rod dissolves. We can summarize
these observations by writing,
Zn → Zn2+ + 2e– (at the zinc rod).
Also, we observe that electrons flow to the copper rod as cupric ions leave
the solution and metallic copper is deposited. We can represent these
occurrences by
2e– + Cu2+ (aq) → Cu (at the copper rod).
In addition, we must examine the purpose of the salt bridge. Since Zinc
ions are produced as electrons leave the zinc electrode, we have a
process which tends to produce a net positive charge in the left beaker.
The purpose of the salt bridge is to prevent any net charge accumulation
in either beaker, diffuse through the bridge, and enter the left beaker.
At the same time, there can be a diffusion of positive ions from left to
right.
If this diffusional exchange of ions did not occur, the net charge
accumulating in the beakers would immediately stop the electron flow
through the external circuit, and the oxidation reduction reaction would
stop.
Thus, while the salt bridge does not participate chemically in the cell
reaction, it is necessary if the cell is to operate.
Significance of salt bridge:
The following are the functions of the salt bridge:


It connects the solutions of two half-cells and completes the cell
circuit.


It prevents transference or diffusion of the solutions from one halfcell to the other.


It keeps the solutions in two half-cells electrically neutral. In anodic
half cell, positive ions pass into the solution and there shall be
accumulation of extra positive charge in the solution around the
anode which will prevent the flow of electrons from anode. This does
not happen because negative ions are provided by salt bridge.
Similarly, in cathodic half-cell negative ions will accumulate around
cathode due to deposition of positive ions by reduction. To neutralize
these negative ions, sufficient number of positive ions is provided by
salt bridge. Thus, salt bridge maintains electrical neutrality.


It prevents liquid-liquid junction-potential, i.e., the potential
difference which arises between two solutions when in contact with
each other.



A broken vertical line or two parallel vertical lines in a cell reaction
indicates the salt bridge.
Zn|Zn2+||Cu2+|Cu


Salt bridge can be replaced by a porous partition which allows the
migration of ions without allowing the solutions to intermix.
IUPAC Cell Representation
The galvanic cell mentioned above is represented in a short IUPAC cell
notation as follows:
It is important to note that:


First of all the anode (electrode of the anode half cell) is written. In
the above case, it is Zn.


After the anode, the electrolyte of the anode should be written with
concentration. In this case it is ZnSO4 with concentration as C1
moles/litre.


A slash (|) is put in between the Zn rod and the electrolyte. This
slash denotes a surface barrier between the two as they exist in
different phases.


Then we indicate the presence of a salt bridge by a double slash (||).


Now, we write the electrolyte of the cathode half-cell which is CuSO4
with its concentration which is C2 moles/ l.


Finally we write the cathode electrode of the cathode half – cell .


A slash (/) between the electrolyte and the electrode in the cathode
half – cell.


In case of a gas, the gas to be indicated after the electrode in case of
anode and before the electrode in case of cathode. Example: Pt,
H2/H+ or H+|H2, Pt.
Electrolysis
Electrolysis can be defined as the process of process of separating any
compound into its constituent elements by passing an electric current
through its aqueous solution.
Preferential Discharge Theory
If an electrolytic solution consists of more than two ions and the
electrolysis is done, it is observed that all the ions are not discharged at
the electrodes simultaneously but certain ions are liberated at the
electrodes in preference to others. This is explained by preferential
discharge theory.
It states that if more than one type of ions are attracted towards a
particular electrode, then the one discharged is the ion which requires
least energy.
The potential at which the ion is discharge or deposition potential.
The values of discharge potential are different for different ions. For
example, the discharge potential of H+ ions is lower than Na+ ions when
platinum or most of the other metals are used as cathodes. Similarly,
discharge potential of Cl- ions is lower than that of OH- ions. This can be
explained by some examples given below:
Electrolysis of sodium chloride solution:
The solution of sodium chloride besides Na+ and Cl- ions possesses
H+ and OH-ions due to ionization of water. However, the number is small
as water is a weak electrolyte. When potential difference is established
across the two electrodes, Na+ and H+ ions move towards cathode and Cland OH- ions move towards anode. At cathode H+ ions are discharged in
preference to Na+ ions as the discharge potential of H+ ions is lower than
Na+ ions. Similarly at anode, Cl- ions are discharged in preference to OHions.
NaCl
H2 O
Na+ + Cl-
H+ + OH-
At cathode
H+ + e - → H
At Anode
Cl- → Cl + e-
2H→ H2
2Cl → Cl2
Thus, Na+ and OH- ions remain in solution and the solution when
evaporated yields crystals of sodium hydroxide.
Refer to the following video for electrolysis of sodium chloride
Electrolysis of copper sulphate solution using platinum electrodes:
CuSO4
H2 O
Cu2+ + SO42-
H+ + OH-
At cathode
At Anode
Cu2+ + 2e- → Cu
2OH- → H2O + O + 2eO + O → O2
Copper is discharged at cathode as Cu2+ ions have lower discharge
potential than H+ ions. OH- ions are discharged at anode as these have
lower discharge potential than ions. Thus, copper is deposited at
cathode and oxygen gas is evolved at anode.
Electrolysis of sodium sulphate solution using inert electrodes:
Na2SO4
H2 O
2Na+ + SO42-
H+ + OH-
At cathode
H+ + e - → H
At Anode
OH- → H2O + O + 2e-
2H→ H2
O + O → O2
Hydrogen is discharged at cathode as H+ ions have lower discharge
potential than Na+ ions. OH- ions are discharged at anode as these have
lower discharge potential than ions. Thus, hydrogen is evolved at
cathode and oxygen is evolved at anode, i.e., the net reaction describes
the electrolysis of water. The ions of Na2SO4 conduct the current through
the solution and take no part in the overall chemical reaction.
The decreasing order of discharge potential or the increasing order of
deposition of some of the ions is given below:
For cations: K+> Na+> Ca2+> Mg2+> Al3+>Zn2+> H+> Cu2+> Hg2+> Ag+
For anions: SO42-> NO3-> OH->Cl-> Br-> I-
Electrolysis of copper sulphate solution using copper electrodes:
CuSO4
Cu2+ + SO42-
At cathode, copper is deposited.
Cu2+ + 2e- → Cu
At anode, the copper of the electrode is oxidised to Cu2+ ions or ions
solution dissolve equivalent amount of copper of the anode.
Cu → Cu2+ + 2e-
Thus, during electrolysis, copper is transferred from anode to cathode.
Electrolysis of silver nitrate solution using silver electrodes:
AgNO2
Ag+ + NO42-
At cathode, silver is deposited.
Ag+ + e- → Ag
At anode, the silver of the electrode is oxidised to Ag+ ions which go into
the solution or ions dissolve equivalent amount of silver of the electrode.
Ag → Ag+ + e-
Ag + NO3- → AgNO3 + e-
Some more examples of electrolysis
Electrolyte
Electrode
Cathodic
reaction
Aqueous acidified
CuCl2 solution
Pt
Cu2+ + 2e-→ Cu
Pt
Pb2+ + 2e- → Pb
Molten PbBr2
Hg
Sodium chloride
solution
Silver nitrate
solution
Pt
2Na+ + 2e- →
2Na
Ag+ + e- --> Ag
Anodic reaction
2Cl- → Cl2 + 2e-
2br- --> Br2 +
2e-
2Cl- → Cl2 + 2e-
2OH- →
1/2 O2 + H2O +
2e-
Sodium nitrate
solution
Pt
2H+ + 2e- → H2
Question 1: Electrolytic cell converts..
a. electrical energy into chemical energy.
b. chemical energy into electrons.
2OH- → 1/2
O2 + H2O + 2e-
c. chemical energy into electrical energy.
d. electrical energy into ions.
Question 2: Which of the following objects is not part of electrochemical
cell
a. Anode
b. Battery
c. Cathode
d. Electrolyte
Question 3: Which of the following statements is incorrect regarding salt
bridge?
a. It connects the solutions of two half-cells and completes the cell
circuit.
b. It prevents transference or diffusion of the solutions from one half-cell
to the other.
c. It keeps the solutions in two half-cells electrically neutral.
d. It causes liquid-liquid junction-potential
Question 4: Which of the following gases is produced at anode during
electrolysis of sodium chloride?
a. Oxygen
b. Hydrogen
c. Chlorine
d. Nitrogen
Question 5: Which of the following ions has maximum discharge
potential?
a. K+
b. Na+
c. Ca2+
d. Mg2
Q.1
a
Q.2
b
Related Resources:-
Q.3
d
Q.4
c
Q.5
a