BatterY...

HiStOrY Of BaTtEry....

The first battery was invented in 1800 by Alessandro Volta. Although it was of great value for experimental purposes, its limitations made it impractical for large current drain. Later batteries, starting with John

Frederic Daniell's wet cell in 1836, provided more reliable currents and were adopted by industry for use in stationary

devices, particularly in telegraph networks where, in the days before electrical distribution networks, they were the only practical source of electricity. These so-called wet cells used liquid electrolytes, and were thus prone to leaks and spillage if not handled correctly. Some, like the gravity cell, could only function in a certain orientation. Many used glass jars to hold their components, which made them fragile. These practical flaws made them unsuitable for portable appliances. Near the end of the 19th century, the invention of dry cell batteries,which replaced liquid electrolyte with a paste, made portable electrical

devices practical.

Daniell Cell

A British chemist named John Frederic Daniell searched for a way to eliminate the hydrogen bubble problem found in the Voltaic Pile, and his solution was to use a second electrolyte to cons

ume the hydrogen produced by the first. In 1836, he invented the Daniell cell, which consisted of a copper pot filled with a copper sulphate solution, in which was immersed an unglazed earthenware container filled with sulphuric acid and a zinc electrode. The earthenware barrier was porous, which allowed ions to pass through but kept the solutions from mixing. Without this barrier, when no current was drawn the copper ions would drift to the zinc anode and undergo reduction without producing a current, which would destroy the battery's life.After a while copper buildup would block the pores in the earthenware barrier and cut short the battery's life. Nevertheless, the Daniel cell provided a longer and more reliable current than the Voltaic cell because the electrolyte deposited copper (a conductor) rather than hydrogen (an insulator) on the cathode. It was also safer and less corrosive. It had an operating voltage of roughly 1.1 volts. It saw widespread use in telegraph networks until it was supplanted by the Leclanché cell in the late 1860s.

CarBon

Zink Batteries

A zinc-carbon dry cell or battery is packaged in a zinc can that serves as both a container and negative terminal. It was developed from the wet Leclanché cell(pronounced /lɛklɑːnˈʃeɪ/). The positive terminal is usually a carbon rod or graphite rod surronded by a mixture of manganese dioxide and carbon powder . Theelectrolyte used is a paste of zinc chloride and ammonium chloride dissolved in water. Zinc chloride cells are an improved version from the original ammonium chloride variety. Zinc-carbon batteries are the least expensive primary batteries and thus a popular choice by manufacturers when devices are sold with batteries included. They are commonly labeled as "General Purpose" batteries. They can be used in remote controls, flashlights, clocks, or transistor radios, since the power drain is not too heavy. A zinc-carbon dry cell is described as a primary cell because as the cell is discharged, it is not intended to be recharged and must be discarded. "Battery Rejuvenators" were once marketed to restore partially discharged zinc-carbon cells by applying a reverse current to them. However the effects of such devices were only temporary and prone to cause the cell to leak or burst. Zinc-carbon cells

are more likely to leak as the anode is the container. The first ever zinc-carbon battery was invented in Selsdon, Croydon in Greater London in 1834.

CheMiCaL ReaCTiOn...

In a dry cell, the outer zinc container is the negative terminal. The zinc oxidise according to the following half-equation.

Zn(s) → Zn²⁺(aq) + 2 e^-

A graphite rod surrounded by a powder containing manganese(I

V) oxide is the positive terminal. The manganese dioxide is mixed with carbon powder to increase the conductivity. The reaction is as follows:

2MnO₂(s) + H₂(g)→ M

n₂O₃(s) + H₂O(l)

The H_{2 comes from the NH4⁺(aq):}

2NH₄⁺(aq) + 2 e^- → H₂(g) + 2NH₃(aq)

and the NH₃ combines with the Zn²⁺.

In this half-reaction, the manganese is reduced from an oxidation state of (+4) to (+3).

There are other possible side-reactions, but the overall reaction in a zinc-carbon cell can be represented as:

Zn(s) + 2MnO₂(s) + 2NH₄⁺(aq) → Mn₂O₃(s) + Zn(NH₃)₂²⁺(aq) + H₂O(l)

The battery has an e.m.f. of about 1.5 V. The approximate nature of the e.m.f is related to the complexity of the cathode reaction. The anode (zinc) reaction is comparatively simple with a known potential. Side reactions and depletion of the active chemicals increases the internal resistance of the battery, and this causes the e.m.f. to drop.