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Zinc–chloride battery

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Zinc–chloride battery


A zinc–carbon battery is a battery packaged in a zinc can that serves as both a container and negative terminal. It was developed from the wet Leclanché cell /lɛklɑːnˈʃ/. The positive terminal is a carbon rod surrounded by a mixture of manganese dioxide and carbon powder. The electrolyte 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, while Zinc chloride cells are often labeled "Heavy Duty". They were the first commercial dry battery and made flashlights and other portable devices possible, because the battery can function in any position. They can be used in low drain or intermittent devices such as remote controls, flashlights, clocks, or transistor radios. They are replaced, in many usages, by alkaline cells, and rechargeable NiMH batteries.

Zinc–carbon batteries account for only 6% of all primary battery sales in Japan and only 7% of all types of batteries sold in Switzerland. In the UK 20% of all portable batteries sold are zinc–carbon and in the EU 18% (by volume).[1][2][3][4]

History

By 1876, the wet Leclanché cell was made with a compressed block of manganese dioxide. In 1886 Dr. Carl Gassner patented a "dry" version by using a zinc cup as the anode and making the electrolyte with a paste of plaster of Paris (and later, wheat flour) to gel and immobilize the electrolyte. In 1898 Conrad Hubert used consumer batteries manufactured by W. H. Lawrence to power what was the first flashlight, and subsequently the two formed the Ever Ready battery company.[5] In 1900 Gassner demonstrated dry cells for portable lighting at the World's Fair in Paris. Continual improvements were made to the stability and capacity of zinc–carbon cells throughout the 20th Century; by the end of the century the capacity of a zinc–carbon cell had increased fourfold over the 1910 equivalent.[6] Improvements include the use of purer grades of manganese dioxide, better sealing, and purer zinc for the negative electrode.


Chemical reactions

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

Zn(s) → Zn2+(aq) + 2 e [E° = +0.7626 volts]

A graphite rod surrounded by a powder containing manganese(IV) oxide is the positive terminal. The manganese dioxide is mixed with carbon powder to increase the electrical conductivity. The reaction is as follows:

2MnO2(s) + 2 e + 2NH4Cl(aq) → Mn2O3(s) + 2NH3(aq) + H2O(l) + 2 Cl [E° ≈ +0.5 v]

and the Cl 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) + 2NH4Cl(aq) → Mn2O3(s) + Zn(NH3)2Cl2 (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.

Although carbon is an important element of the battery's construction, it takes no part in the electrochemical reaction, instead only serving to collect current and reduce the resistance of the manganese dioxide mix. The cell could more properly be called a "zinc–manganese" cell.

A zinc–carbon dry cell is considered as a primary cell because the cell is not intended to be recharged.

Construction

The container of the zinc–carbon dry cell is a zinc can. The can contains a layer of NH4Cl or ZnCl2 aqueous paste impregnating a paper layer that separates the zinc can from a mixture of powdered carbon (usually graphite powder) & manganese (IV) oxide (MnO2) which is packed around a carbon rod. Carbon is the only practical conductor material because every common metal will quickly corrode away in the positive electrode in salt based electrolyte.

Early types, and low-cost cells, use a separator consisting of a layer of starch or flour. A layer of starch-coated paper is used in modern cells, which is thinner and allows more manganese dioxide to be used. Originally cells were sealed with a layer of asphalt to prevent drying out of the electrolyte; more recently a thermoplastic washer sealant is used. The carbon rod is slightly porous, which allows accumulated gas to escape while retaining the water for the electrolyte. The ratio of manganese dioxide and carbon powder in the cathode paste affects the characteristics of the cell; more carbon powder lowers the internal resistance, but more manganese dioxide improves capacity.[6]

Flat cells are also made for assembly into batteries with higher voltages, up to about 450 volts. A number of flat cells are stacked up, and the whole assembly is coated in wax to prevent evaporation of water from the electrolyte.

Leakage

Generally, the materials of the cell are inexpensive and cheap to reproduce even though the cell is non-rechargeable. These cells have a low energy density; the voltage falls during use due to a drop in electrolyte concentration around the cathode and the time required for the Mn2O3 to diffuse away from the cathode.

These cells have a short shelf life as the zinc is attacked by ammonium chloride. The zinc container becomes thinner as the cell is used, because zinc metal is oxidized to zinc ions. When the zinc case thins enough, zinc chloride begins to leak out of the battery. The old dry cell is not leak proof and becomes very sticky as the paste leaks through the holes in the zinc case. The zinc casing in the dry cell gets thinner even when the cell is not being used, because the ammonium chloride inside the battery reacts with the zinc. An "inside-out" form with a carbon cup and zinc vanes on the interior, while more leak resistant, has not been made since the 1960s.[6]

This picture shows the zinc container of fresh batteries at (a), and discharged batteries at (b) and (c). The battery shown at (c) had a polyethylene protection film (mostly removed in the photo) to keep the zinc oxide inside the casing.

Environmental impact

Thousands of tons of zinc–carbon batteries are discarded every year around the world, and are often not recycled.

The inside surface of the zinc can was originally treated with mercury to form an amalgamated electrode. This decreased 'local action' where impurities in the zinc set up small areas of galvanic action which would cause the zinc to react with the electrolyte more quickly than it otherwise would. Mercury in discarded cells could escape into the environment. Legislation in several countries (such as the European Union Battery Directive and the United States Mercury-Containing and Rechargeable Battery Management Act) now restricts the use of mercury in batteries. Manufacturers must now use more highly purified zinc to prevent local action and self-discharge.

Conveniently manufactured, and disposed, a single zinc–carbon dry cell has low environmental impact on disposal, compared with some other battery types. The manganese (III) is readily oxidized and so becomes immobilized; minute amounts of zinc, carbon and ammonium salts are also harmless.[7] In the United States, federal law does not require zinc-carbon batteries to be treated as hazardous waste.

The zinc chloride cell

The zinc chloride cell is an improvement on the original zinc–carbon cell, using purer chemicals and giving a longer life and steadier voltage output as it is used. These cells are often marketed as heavy-duty cells, to differentiate them from general-purpose zinc–carbon cells. This has been a source of consumer confusion after the introduction of alkaline cells, which last longer than the zinc-chloride heavy-duty cell. Instead of an electrolyte mixture containing much NH4Cl, it is largely only ZnCl2 paste. The cathode reaction is thus a little different:

MnO2(s) + H2O(l) + e- → MnO(OH)(s) + OH-(aq)

as is the overall reaction:

Zn(s) + 2 MnO2(s) + ZnCl2(aq) + 2 H2O(l) → 2 MnO(OH)(s) + 2 Zn(OH)Cl(aq)

Electrodes

Anode and cathode of an electrochemical device are defined by the direction of current flow, not the polarity of voltage. In electrolytic cells, the anode is referred as the positive terminal since all the anions (negative ions) will migrate to the anode to be selectively discharged while the cathode is the negative terminal because the cations (positive ions) will move to the cathode to be selectively discharged. Meanwhile, for voltaic cells, the anode is the negative terminal, while the cathode is the positive terminal. This is due to the convention which states that all anodes are terminals that undergo oxidation or release of electrons, and all cathodes are terminals which undergo reduction. (Conventional) current departs at a cathode and enters a device at the anode, independent of the polarity of the voltage of the device.

Storage

Manufacturers recommend storage of zinc–carbon batteries at room temperature; storage at higher temperatures reduces the expected service life.[8] While batteries may be frozen without damage, manufacturers recommend that they be returned to normal room temperature before use, and that condensation on the battery jacket must be avoided. By the end of the 20th century, the storage life of zinc–carbon cells had improved fourfold over expected life in 1910.[6]

See also

References

External links

  • Eveready: Carbon Zinc Application Notes
  • Rayovac: Alkaline and Heavy Duty Application Notes
  • Power Stream Battery Chemistry FAQs
  • Cell Construction
  • Malaysia Chemistry Text Book for Secondary School Form 4.
  • http://www.snopes.com/oldwives/battery.asp Cool Tips
sv:Brunstensbatteri
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