BATTERY

An electrical battery is one or more electrochemical cells that convert

stored chemical energy into electrical energy. It can be used as a source of
direct electric current at a constant voltage. A battery has advantage of being
completely self-contained and requiring no auxiliary components such as salt
bridges.According to a 2005 estimate, the worldwide battery industry
generates US$48 billion in sales each year, with 6% annual growth.

There are two types of batteries: primary batteries (disposable batteries),

which are designed to be used once and discarded when they are exhausted,
and secondary batteries (rechargeable batteries), which are designed to be
recharged and used multiple times. Miniature cells are used to power devices
such as hearing aids and wristwatches; larger batteries provide standby
power for telephone exchanges or computer data centers.

HISTORY

The first electrochemical cell was developed by Italian physicist

Alessandro Volta in 1792, and he invented the first battery in
1800.

Benjamin franklin in 1748 described multiple Leyden jars (early

electrical capacitors) which predated Volta’s use of multiple
galvanic cells. Volta's work was stimulated by the Italian
anatomist and physiologist Luigi Galvani, who in 1780 noticed
that dissected frog's legs would twitch when struck by a spark
from a Leyden jar which was an external source of electricity. In
1791 he published a report on "animal electricity." He created
an electric circuit consisting of the frog's leg (FL) and two
different metals A and B, each metal touching the frog's leg and
each other, thus producing the circuit A-FL-B-A-FL-B...etc. The
frog's leg served both as the electrolyte and the sensor, and
the metals served as electrodes. He noticed that even though
the frog was dead, its legs would twitch when he touched them
with the metals.
Within a year, Volta realized the frog's moist tissues could be
replaced by cardboard soaked in salt water, and the frog's
muscular response could be replaced by another form of
electrical detection. He already had studied the electrostatic
phenomenon of capacitance, which required measurements of
electric charge and of electrical potential ("tension"). Building
on this experience, Volta was able to detect electric current
through Galvanic cell. The terminal voltage of a cell that is not
discharging is called its electromotive force (emf), and has the
same unit as electrical potential, named (voltage) and
measured in volts, in honor of Volta. In 1800, Volta invented the
battery by placing many voltaic cells in series, piling them one
above the other. This voltaic pile gave enhanced net emf for
the combination with a voltage of about 50 volts for a 32-cell
pile.
Volta did not appreciate that the voltage was due to chemical
reactions. He thought that his cells were an inexhaustible
source of energy,[15] and that the associated chemical effects
(e.g. corrosion) were a mere nuisance, rather than an
unavoidable consequence of their operation, as Michael
Faraday showed in 1834.[16] According to
Faraday, cations (positively charged ions) are attracted to
thecathode,[17] and anions (negatively charged ions) are
attracted to the anode.[18]
Although early batteries were of great value for experimental
purposes, in practice their voltages fluctuated and they could
not provide a large current for a sustained period. Later,
starting with the Daniel cell in 1836, batteries provided more
reliable currents and were adopted by industry for use in
stationary devices, particularly in telegraph networks where
they were the only practical source of electricity, since
electrical distribution networks did not exist at the time.
[19]
These wet cells used liquid electrolytes, which were prone to
leakage and spillage if not handled correctly. Many used glass
jars to hold their components, which made them fragile. These
characteristics made wet cells unsuitable for portable
appliances. Near the end of the nineteenth century, the
invention of dry cell batteries, which replaced the liquid
electrolyte with a paste, made portable electrical devices.
PRINCIPAL OF OPERATION
It consists of a number of voltaic cells; each voltaic cell consists of two half cells connected in
series by a conductive electrolyte containing anions and cations. Electrolyte is connected
internally through a salt bridge. One half-cell includes electrolyte and the electrode to which
anions (negatively charged ions) migrate, i.e., the anode or negative electrode; the other half-cell
includes electrolyte and the electrode to which cations (positively charged ions) migrate, i.e., the
cathode or positive electrode. In the redox reaction that powers the battery, reduction (addition of
electrons) occurs to cations at the cathode, while oxidation (removal of electrons) occurs to
anions at the anode.[23] The electrodes do not touch each other but are electrically connected by
the electrolyte. The two half cells are connected by a metallic wire through a voltmeter and a
switch externally. Many cells use two half-cells with different electrolytes. In that case each half-
cell is enclosed in a container, and a separator that is porous to ions, but not the bulk of the
electrolytes, prevents mixing. At each electrode-electrolyte interface there is a tendency of metal
ions from the solution to deposit on the metal electrode trying to make it positively charged. At
equilibrium, there is a separation of charges and depending on the tendency of the two opposing
reactions, the electrode may be positively or negatively charged with respect to the solution. A
potential difference develops between the electrode and the electrolyte which is called electrode
potential. When the concentrations of all the species involved in a half cell is unity then the
electrode potential is known as standard electrode potential. The difference between the
electrode potential of the two half cells is called electromotive force (EMF). It is measured with
a voltmeter and expressed in volts. Net emf is €2 - €1. The electrical driving force across the
terminals of cell is known as terminal voltage. The terminal voltage of a cell that is neither
charging nor discharging is known as open circuit voltage and is equal to emf of cell. The
terminal voltage of a cell that is discharging is smaller in magnitude than the open-circuit voltage
and the terminal voltage of a cell that is charging exceeds the open-circuit voltage because of
internal resistance. An ideal cell has negligible internal resistance, so it would maintain a
constant terminal voltage of until exhausted, then dropping to zero. If such a cell maintained
1.5 volts and stored a charge of one coulomb then on complete discharge it would perform
1.5 joule of work.[25] In actual cells, the internal resistance increases under discharge,[26] and the
open circuit voltage also decreases under discharge. If the voltage and resistance are plotted
against time, the resulting graphs typically are a curve; the shape of the curve varies according to
the chemistry and internal arrangement employed.[28]