Electrochemical cell
A demonstration electrochemical cell setup resembling the Daniell cell. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit.

An electrochemical cell is a device capable of either deriving electrical energy from chemical reactions, or facilitating chemical reactions through the introduction of electrical energy. A common example of an electrochemical cell is a standard 1.5-volt "battery". (Actually a single "Galvanic cell"; a battery properly consists of multiple cells, connected in either parallel or series pattern.)

Contents

Half-cells

The Bunsen cell, invented by Robert Bunsen.

An electrochemical cell consists of two half-cells. Each half-cell consists of an electrode, and an electrolyte. The two half-cells may use the same electrolyte, or they may use different electrolytes. The chemical reactions in the cell may involve the electrolyte, the electrodes or an external substance (as in fuel cells which may use hydrogen gas as a reactant). In a full electrochemical cell, species from one half-cell lose electrons (oxidation) to their electrode while species from the other half-cell gain electrons (reduction) from their electrode. A salt bridge (e.g. filter paper soaked in KNO3) is often employed to provide ionic contact between two half-cells with different electrolytes—to prevent the solutions from mixing and causing unwanted side reactions. As electrons flow from one half-cell to the other, a difference in charge is established. If no salt bridge were used, this charge difference would prevent further flow of electrons. A salt bridge allows the flow of ions to maintain a balance in charge between the oxidation and reduction vessels while keeping the contents of each separate. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the Bunsen cell (right).

Equilibrium reaction

Each half-cell has a characteristic voltage. Different choices of substances for each half-cell give different potential differences. Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions—when equilibrium is reached the cell cannot provide further voltage. In the half-cell which is undergoing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide. Similarly, in the reduction reaction, the further the equilibrium lies to the ion/atom with the more negative oxidation state the higher the potential.

Cell potential

The cell potential can be predicted through the use of electrode potentials (the voltages of each half-cell). (See table of standard electrode potentials). The difference in voltage between electrode potentials gives a prediction for the potential measured.

Cell potentials have a possible range of about zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts, because the very powerful oxidizing and reducing agents which would be required to produce a higher cell potential tend to react with the water.

Electrochemical cell types

Main types

Cells are classified into two broad categories,

  • Primary cells irreversibly (within limits of practicality) transform chemical energy to electrical energy. When the initial supply of reactants is exhausted, energy cannot be readily restored to the electrochemical cell by electrical means.[1]
  • Secondary cells can be recharged; that is, they can have their chemical reactions reversed by supplying electrical energy to the cell, restoring their original composition.[2]

Primary electrochemical cells

Primary electrochemical cells can produce current immediately on assembly. Disposable cells are intended to be used once and discarded. Disposable primary cells cannot be reliably recharged, since the chemical reactions are not easily reversible and active materials may not return to their original forms.

Common types of disposable cells include zinc-carbon cells and alkaline cells. Generally, these have higher energy densities than rechargeable cells,[3] but disposable cells do not fare well under high-drain applications with loads under 75 ohms (75 Ω).[4]

Secondary electrochemical cells

Secondary electrochemical cells must be charged before use; they are usually assembled with active materials in the discharged state. Rechargeable electrochemical cells or secondary electrochemical cells can be recharged by applying electric current, which reverses the chemical reactions that occur during its use. Devices to supply the appropriate current are called chargers or rechargers.

The oldest form of rechargeable cell is the lead-acid cell.[5] This electrochemical cell is notable in that it contains a liquid in an unsealed container, requiring that the cell be kept upright and the area be well ventilated to ensure safe dispersal of the hydrogen gas produced by these cells during overcharging. The lead-acid cell is also very heavy for the amount of electrical energy it can supply. Despite this, its low manufacturing cost and its high surge current levels make its use common where a large capacity (over approximately 10Ah) is required or where the weight and ease of handling are not concerns.

An improved type of liquid electrolyte cell is the sealed valve regulated lead acid (VRLA) cell, popular in the automotive industry as a replacement for the lead-acid wet cell. The VRLA cell uses an immobilized sulphuric acid electrolyte, reducing the chance of leakage and extending shelf life.[6] VRLA cells have the electrolyte immobilized, usually by one of two means:

  • Gel cells contain a semi-solid electrolyte to prevent spillage.
  • Absorbed Glass Mat (AGM) cells absorb the electrolyte in a special fibreglass matting

Other portable rechargeable cells are (in order of increasing power density and cost): nickel-cadmium cells (NiCd), nickel metal hydride cells (NiMH), and lithium-ion cells(Li-ion).[7] By far, Li-ion has the highest share of the dry cell rechargeable market.[8] Meanwhile, NiMH has replaced NiCd in most applications due to its higher capacity, but NiCd remains in use in power tools, two-way radios, and medical equipment.[8]

Special types

See also

References

  1. ^ Dingrando 675.
  2. ^ Fink, Ch. 11, Sec. "Batteries and Fuel Cells."
  3. ^ Alkaline Manganese Dioxide Handbook and Application Manual (PDF). Energizer. Retrieved 25 August 2008.
  4. ^ Buchmann, Isidor. Will secondary batteries replace primaries?. Battery University. Retrieved 6 January 2008.
  5. ^ Buchmann, Isidor. Can the lead-acid battery compete in modern times?. Battery University. Retrieved 2 September 2007.
  6. ^ Dynasty VRLA Batteries and Their Application. C&D Technologies, Inc. Retrieved 26 August 2008.
  7. ^ What's the best battery?. Battery University. Retrieved 26 August 2008.
  8. ^ a b Buchmann, Isidor. Battery statistics. Battery University. Retrieved 11 August 2008.

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