Co 1
Co 1
Co 1
For
The Course: Engineering Chemistry
A.Y. 2018-19
CO1
KLEF
Department of Chemistry
Basic concepts of electrochemistry-Electrode Potential, Single electrode
potential Galvanic cell, Measurement of Electrode Potential.
It is defined as the potential developed at the inter phase between the metal and the
solution, when a metal is dipped in a solution containing its own ions. It is
represented as E.
When a metal is dipped in a solution containing its own ions, the metal may undergo
oxidation by loosing electrons or the metal ions in solution may undergo reduction and
get deposited on the metal surface.
M Mn+ + ne-
Simultaneously the metal ions from the solution tend to deposit on the metal as metal
atoms (reduction)
Mn+ + ne- M
Mn+ + ne- → M
When a metal undergoes oxidation it loses positive ions into solution leaving behind a
layer of negative charges on its surface. This layer attracts positive changes and forms
an electric double layer (EDL) because of the formation of EDL electrode potential
arises.
When metal ions undergo reduction depositing metal atoms on the metallic surface the
metal surface becomes positively charged. The accumulated positive charge on the
metal surface attracts a layer of –ve charges and forms an electrical double layer or
Helmotz EDL which causes the origin of electrode potential.
The electrode potentials of any metal electrodes can be determined by using reference
electrodes like standard hydrogen electrode (SHE).
The SHE is coupled with the electrode whose electrode potential is to be determined and
the electrode potential of the electrode is determined by fixing the electrode potential and
SHE as zero [at all temperatures]
Example: Consider the determination of Single electrode potential of Zinc electrode using
Standard Hydrogen electrode.
To determine the Single electrode potential of Zinc electrode it is coupled with Standard
Hydrogen electrode as follows:
E cell = E SHE – E Zn
0.76 = 0 - E Zn
E Zn = -0.76 V
1) Galvanic or Voltaic cells (Reversible): These are the electrochemical cells, which
convert chemical energy into electrical energy.
2) Electrolytic cell (Irreversible): These are the electrochemical cells, which are used
to convert electrical energy into chemical energy.
It is defined as the potential difference between the two electrodes of a galvanic cell
which causes the flow of current from an electrode with higher reduction potential to the
electrode with lower reduction potential. It is denoted as E cell.
E cell = E right –E left.
Problems:
The Eqs (5) and (7) are known as Nernst Equation for single electrode potential.
Applications of Nernst Equation:
Electrode potential of unknown metal can be determined using the Nernst equation.
It is used to predict the corrosion tendency of metals.
Problems:
(1) Calculate the reduction potential of Cu/Cu 2+ electrode which is dipped in 0.5M of
its own salt solution at 250 C and the standard electrode potential or its is 0.337.
Find out the emf.
E = E0 + 0.0591/n log (Cu 2+) Volts n =2 ;
(Ans): Reduction Potential= 0.3230 V
(2) Find out the oxidation potential of Zn/Zn 2+ which is in contact with 0.2 M of its
own salt solution at 250C and the std. electrode potential is 0.763 V.
E = E0 + 0.0591/n log (Zn 2+) V n=2;
(Ans) : Oxidation potential = 0.7836 V
Concentrations cells
These are the galvanic cells consisting of same metal electrodes dipped in same
metal ionic solution in both the half cells but are different in the concentration of
the metal ions.
Ex: Consider the following concentration cell constructed by dipping two copper
electrodes in CuSO4 solutions of M2 molar and M1 molar where M2 > M1
The two half-cells are internally connected by a salt bridge and externally connected
by a metallic wire through voltmeter or ammeter.
The electrode, which is dipped in less ionic concentrations solutions (M1) act as
anode and undergoes oxidation. The electrode, which is dipped in more ionic
concentration(M2) act as cathode and undergoes reduction.
The standard hydrogen electrode consists of platinised platinum foil fused to the
glass tube. Mercury is placed at the bottom of the tube and a copper wire is used for
electrical connections. The platinum foil is immersed in a solution containing unit
molar hydrogen ions. Pure hydrogen gas is bubbled about the electrode through the H2
gas inlet at 1atm pressure.
Limitations of SHE
The calomel electrode acts as both anode and cathode depending upon the other
electrode used. The platinum wire is used for electrical connections. Salt bridge is used
to couple with other half cell.
Ag/AgCl electrode is a metal metal salt electrode. It consists of narrow glass tube at
the bottom of which agar is placed above which saturated solution of KCl is placed.
The silver wire is used for electrical connections and it is coated electrolytically with
AgCl. The cell can be
Electrode acts as both anode and cathode depending on the other electrode used.
When it acts as anode the electrode reaction is
Ag + Cl- Ag Cl + e-
When it acts as cathode the electrode reaction is
AgCl + e- Ag + Cl-.
Ion Selective Electrodes
= Eo Glass - 0.0591 pH
Session-6
2) Secondary Battery is the one in which cell reactions can be reversed by passing
external emf in opposite direction. i.e. it can be used for many cycles of charging and
discharging.
3) Flow battery is the one in which all the constituents of the battery flow throughout
the battery.
E.g. Hydrogen-Oxygen
Oxygen Fuel cell.
The anode is made of zinc (Zn) and separated from the cathode with a layer of paper or
other porous material soaked with electrolyte; this is known as a salt bridge. Two half-
reactions occur at the anode. The first consists of an electrochemical reaction step:
Zn + HgO → ZnO + Hg
In other words, during discharge, zinc is oxidized (loses electrons) to become zinc oxide
(ZnO) while the mercuric oxide gets reduced (gains electrons) to form elemental
mercury. A little extra mercuric oxide is put into the cell to prevent evolution of hydrogen
gas at the end of life.
Storage cell: it is the one which can act both as voltaic cell and electrolytic cell. When it
functions as voltaic cell, it supplies electric current and the process is known as
discharging. When it functions as electrolytic cell, it receives electric current and this
process is known as charging. So it can be used for a large no. of cycles of charging and
discharging. The best example for storage cell is lead acid battery or lead acid
accumulator.
Discharging:
Reactions taking place during discharging: During discharging, it acts as voltaic cell and
supplies electrical energy.
Charging:
2 PbSO4 ↓ + 2 H2O + Energy (> 2 V) → Pb + PbO2 + 4 H+ + 2 SO42-
The principles in which NiMH cells operate are based on their ability to absorb, release,
and transport (move) hydrogen between the electrodes within the cell. The following
sections will discuss the chemical reactions occurring within the cell when charged and
discharged and the adverse effe cts of overcharge and overdischarge conditions.
The success of the NiMH battery technology comes from the rare earth, hydrogen-
absorbing alloys (commonly known as Misch metals) used in the negative electrode.
These metal alloys contribute to the high energy density of the NiMH negative electrode
that results in an increase in the volume available for the positive electrode. This is the
primary reason for the higher capacity and longer service life of NiMH batteries over
competing secondary batteries.
When a NiMH cell is charged, the positive electrode releases hydrogen into the
electrolyte. The hydrogen in turn is absorbed and stored in the negative electrode. The
reaction begins when the nickel hydroxide (Ni(OH)2 in the positive electrode and
hydroxide (OH-) from the electrolyte combine.
This produces nickel oxyhydroxide (NiOOH) within the positive electrode, water (H2O)
in the electrolyte, and one free electron (e ̄). At the negative electrode the metal alloy (M)
in the negative electrode, water (H2O) from the electrolyte, and an electron (e-) react to
produce metal hydride (MH) in the negative electrode and hydroxide (OH̄ ) in the
electrolyte.
Li-ion battery, H2-O2 Fuel cell
The overall reaction has its limits. Over discharge supersaturates lithium cobalt oxide,
oxide
leading to the production of lithium oxide possibly by the following irreversible reaction:
In a lithium-ion
ion battery the lithium ions are transported to and from the cathode or anode,
with the transition metal, cobalt (Co),
( in LixCoO2 being oxidized from Co3+ to Co4+
during charging, and reduced from Co4+ to Co3+ during discharge.
It consists of two inert porous electrodes made of graphite impregnated with finely
divided Pt and 25% KOH solution as electrolyte. Hydrogen gas is bubbled through one
inert electrode, acts as anode. Oxygen gas is bubbled through
through another electrode, acts as
cathode. The hydrogen-oxygen
oxygen fuel cell produces water as a product and hence is an ideal
power source for zero-emission
ission vehicles. Hence it is called an eco friendly battery.
Redox Reaction in a Hydrogen- oxygen Fuel Cell