Electrochemistry 4 Flashcards
what is a cell
Any battery (actually it may have one or more than one cell connected
in series) or cell that we use as a source of electrical energy is basically
a galvanic cell where the chemical energy of the redox reaction is
converted into electrical energy. However, for a battery to be of practical
use it should be reasonably light, compact and its voltage should not
vary appreciably during its use.
what are primary cells
In the primary batteries, the reaction occurs only once and after use
over a period of time battery becomes dead and cannot be reused
again.
describe leclanche cell
i) electrodes
ii) electrolyte
iii) electrode reactions
The most familiar example of this type is the dry
cell (known as Leclanche cell after its discoverer) which is used commonly in our transistors and clocks.
The cell consists of a zinc container that also acts as anode and the cathode is a carbon (graphite) rod surrounded by powdered manganese dioxide and carbon (Fig.2.8). The
space between the electrodes is filled by a moist paste of ammonium chloride (NH4Cl) and zinc chloride (ZnCl2).
The electrode reactions are complex, but they can be written
approximately as follows
Anode: Zn(s) → Zn2+ + 2e–
Cathode: MnO2+ NH4++ e– → MnO(OH) + NH3
In the reaction at cathode, manganese is reduced
from the + 4 oxidation state to the +3 state.
Ammonia produced in the reaction forms a complex with Zn2+ to give[Zn (NH3)4]2+. The cell has a potential of nearly 1.5 V.
describe mercury cell
i) uses
ii) electrodes/ electrolyte
iii) electrode reactions
iv) potential
Mercury cell, (Fig. 2.9) suitable for low current devices like hearing aids, watches, etc. consists of zinc – mercury amalgam as anode and a paste of HgO and carbon as the cathode. The electrolyte is a paste of KOH and ZnO. The electrode reactions for the cell are given below:
Anode: Zn(Hg) + 2OH– → ZnO(s) + H2O + 2e–
Cathode: HgO + H2O + 2e– → Hg(l) + 2OH–
The overall reaction is represented by
Zn(Hg) + HgO(s) → ZnO(s) + Hg(l)
The cell potential is approximately
1.35 V and remains constant during its
life as the overall reaction does not
involve any ion in solution whose
concentration can change during its life
time.
what is a secondary cell
A secondary cell after use can be recharged by passing current
through it in the opposite direction so that it can be used again. A
good secondary cell can undergo a large number of discharging
and charging cycles.
describe the lead storage battery
i) use
ii) electrodes and electrolyte
iii) reactions
iv) reversed
The most important secondary cell is the lead
storage battery (Fig. 2.10) commonly used in automobiles and
invertors.
It consists of a lead anode (lead grid filled with spongy lead) and a grid of lead packed with
lead dioxide (PbO2 ) as cathode. A 38% solution of sulphuric acid is used as an electrolyte.
The cell reactions when the battery is in use are given below:
Anode: Pb(s) + SO42–(aq) → PbSO4 (s) + 2e–
Cathode: PbO2(s) + SO42–(aq) + 4H+(aq) + 2e– → PbSO4(s) + 2H2O (l)
i.e., overall cell reaction consisting of cathode and anode reactions is:
Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)
On charging the battery the reaction is reversed and PbSO4 (s) on anode and cathode is converted into Pb and PbO2
, respectively.
describe nickel cadmium battery
Another important secondary
cell is the nickel-cadmium cell
(Fig. 2.11) which has longer life
than the lead storage cell but
more expensive to manufacture.
We shall not go into details of
working of the cell and the
electrode reactions during
charging and discharging.
The overall reaction during
discharge is:
Cd (s) + 2Ni(OH)3 (s) → CdO (s) + 2Ni(OH)2 (s) + H2O (l )
A Ni-Cd cell is in the form of a jelly roll arrangement, where the positive and negative sheets are seperated by a seperator ( a layer soaked in sodium/potassium hydroxide)
traditional method of energy production
Production of electricity by thermal plants is not a very efficient method
and is a major source of pollution. In such plants, the chemical energy
(heat of combustion) of fossil fuels (coal, gas or oil) is first used for
converting water into high pressure steam. This is then used to run
a turbine to produce electricity.
what are fuel cells
Galvanic cells that are designed to convert
the energy of combustion of fuels like hydrogen, methane, methanol,
etc. directly into electrical energy are called fuel cells.
hydrogen oxygen fuel cells
One of the most successful fuel cells
uses the reaction of hydrogen with oxygen
to form water (Fig. 2.12). The cell was
used for providing electrical power in the
Apollo space programme. The water
vapours produced during the reaction
were condensed and added to the
drinking water supply for the astronauts.
In the cell, hydrogen and oxygen are
bubbled through porous carbon
electrodes into concentrated aqueous
sodium hydroxide solution. Catalysts like
finely divided platinum or palladium
metal are incorporated into the electrodes
for increasing the rate of electrode
reactions.
cell reactions
Cathode: O2(g) + 2H2O(l) + 4e– → 4OH–(aq)
Anode: 2H2(g) + 4OH–(aq) → 4H2O(l) + 4e–
Overall reaction being:2H2(g) + O2(g) → 2H2O(l )
efficiency and outlook of fuel cells
The cell runs continuously as long as the reactants are supplied.
Fuel cells produce electricity with an efficiency of about 70 % compared to thermal plants whose efficiency is about 40%.
There has been
tremendous progress in the development of new electrode materials, better catalysts and electrolytes for increasing the efficiency of fuel cells.
These have been used in automobiles on an experimental basis. Fuel cells are pollution free and in view of their future importance, a variety
of fuel cells have been fabricated and tried.
what is meant by corrosion
Corrosion is the gradual destruction or deterioration of a material, typically a metal, due to a chemical reaction with its environment.
Corrosion slowly coats the surfaces of metallic objects with oxides or
other salts of the metal. The rusting of iron, tarnishing of silver,
development of green coating on copper and bronze are some of the
examples of corrosion. It causes enormous damage to
buildings, bridges, ships and to all objects made of
metals especially that of iron.
anode and cathode on iron metal
At a particular spot
(Fig. 2.13) of an object made of iron,
oxidation takes place and that spot
behaves as anode and we can write
the reaction
Oxidation: Fe (s)→ Fe2+ (aq) +2e–
Electrons released at anodic spot move through the metal and go
to another spot on the metal and reduce oxygen in the presence of H+
(which is believed to be available from H2CO3 formed due to dissolution
of carbon dioxide from air into water. Hydrogen ion in water may also
be available due to dissolution of other acidic oxides from the
atmosphere). This spot behaves as cathode with the reaction
Cathode: O2(g) + 4 H+(aq) + 4 e– → 2 H2O (l)
The overall reaction being:
2Fe(s) + O2(g) + 4H+(aq) → 2Fe2 +(aq) + 2 H2O (l)
atmospheric oxidation
The ferrous ions are further oxidised by atmospheric oxygen to
ferric ions which come out as rust in the form of hydrated ferric oxide
(Fe2O3
. x H2O) and with further production of hydrogen ions.
2Fe2+(aq) + 2H2O(l) + ½O2
(g) →→ Fe2O3 (s) + 4H+(aq)
