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ARTICLES IN THE BOOK

  1. AAAA battery
  2. AAA battery
  3. AA battery
  4. A battery
  5. Absorbent glass mat
  6. Alessandro Volta
  7. Alkaline battery
  8. Alkaline fuel cell
  9. Aluminium battery
  10. Ampere
  11. Atomic battery
  12. Backup battery
  13. Baghdad Battery
  14. Batteries
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  16. B battery
  17. Bernard S. Baker
  18. Beta-alumina solid electrolyte
  19. Betavoltaics
  20. Bio-nano generator
  21. Blue energy
  22. Bunsen cell
  23. Car battery
  24. C battery
  25. Clark cell
  26. Concentration cell
  27. Coulomb
  28. 2CR5
  29. Daniell cell
  30. Direct borohydride fuel cell
  31. Direct-ethanol fuel cell
  32. Direct methanol fuel cell
  33. Dry cell
  34. Dry pile
  35. Duracell
  36. Duracell Bunny
  37. Earth battery
  38. Electric charge
  39. Electric current
  40. Electricity
  41. Electrochemical cell
  42. Electrochemical potential
  43. Electro-galvanic fuel cell
  44. Electrolysis
  45. Electrolyte
  46. Electrolytic cell
  47. Electromagnetism
  48. Electromotive force
  49. Energizer Bunny
  50. Energy
  51. Energy density
  52. Energy storage
  53. Flashlight
  54. Float charging
  55. Flow Battery
  56. Formic acid fuel cell
  57. Fuel cell
  58. Fuel cell bus trial
  59. Galvanic cell
  60. Gel battery
  61. Grove cell
  62. Half cell
  63. History of the battery
  64. Hybrid vehicle
  65. Lead-acid battery
  66. Leclanché cell
  67. Lemon battery
  68. List of battery sizes
  69. List of battery types
  70. List of fuel cell vehicles
  71. Lithium battery
  72. Lithium ion batteries
  73. Lithium iron phosphate battery
  74. Lithium polymer cell
  75. LR44 battery
  76. Luigi Galvani
  77. Manganese dioxide
  78. Memory effect
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  80. Metal hydride fuel cell
  81. Methane reformer
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  83. Michael Faraday
  84. Microbial fuel cell
  85. Molten carbonate fuel cell
  86. Molten salt battery
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  90. Nickel-zinc battery
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  95. Panasonic EV Energy Co
  96. Peukert's law
  97. Phosphoric acid fuel cell
  98. Photoelectrochemical cell
  99. Polymer-based battery
  100. Power density
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  103. PP3 battery
  104. Primary cell
  105. Prius
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  107. Proton exchange membrane fuel cell
  108. Protonic ceramic fuel cell
  109. Radioisotope piezoelectric generator
  110. Ragone chart
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  114. Reversible fuel cell
  115. Searchlight
  116. Secondary cell
  117. Short circuit
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  121. Sodium-sulfur battery
  122. Solid oxide fuel cell
  123. Super iron battery
  124. Thermionic converter
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  126. Vanadium redox battery
  127. Volt
  128. Voltage
  129. Voltaic pile
  130. Watch battery
  131. Water-activated battery
  132. Weston cell
  133. Wet cell
  134. Zinc-air battery
  135. Zinc-bromine flow battery
  136. Zinc-carbon battery

 

 
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This article is from:
http://en.wikipedia.org/wiki/Galvanic_cell

All text is available under the terms of the GNU Free Documentation License: http://en.wikipedia.org/wiki/Wikipedia:Text_of_the_GNU_Free_Documentation_License 

Galvanic cell

From Wikipedia, the free encyclopedia

 

The Galvanic cell, named after Luigi Galvani, consists of two different metals connected by a salt bridge or a porous disk between the individual half-cells. It is also known as a voltaic cell and an electrochemical cell.

History

In 1780, Luigi Galvani discovered that when two different metals (copper and zinc for example) were connected together and then both touched to different parts of a nerve of a frog leg at the same time, they made the leg contract. He called this "animal electricity". The Voltaic pile, invented by Alessandro Volta in the 1800s, is similar to the galvanic cell. These discoveries paved the way for all electrical batteries.

Description

Scheme of a galvanic cell
Scheme of a galvanic cell

The Galvanic cell's metals dissolve in the electrolyte at two different rates. Metals become positive ions upon dissolving, and leave electrons behind. As a result, the metal acquires a negative net charge while the electrolyte becomes equally positive. Each metal undergoes a different half-reaction, giving different dissolving rates, which builds up different electrode potentials between the electrolyte and each metal. If an electrical connection, such as a wire or direct contact, is formed between the two electrodes, an electric current appears in the metal. At the same time, an equal electric current composed of positive ions appears in the electrolyte. Ions of the more active metal which forms the anode are transferred to the electrolyte. Dissolved ions are also transferred to the less active metal, the cathode, and deposited there as a plating. In this way the anode is consumed or corroded. When the anode material corrodes entirely away, the cell's potential drops and the current halts. The metal may be regarded as the fuel which powers the device. A similar process is used in electroplating. The electric current in the electrolyte is equal to the current in the external circuit, so a complete circuit is formed with a path through the electrolyte.

There is a flow of electrons from the oxidized ion at the anode to the reduced atom (formerly an ion) at the cathode. It is this flow, due to this redox reaction which constitutes the current.

Electric potential of a Galvanic cell

The electrode potential of a cell can be easily determined by use of a standard potential table. An oxidation potential table could also be used, but the reduction table is more common. The first step is to identify the two metals reacting in the cell. Then one looks up the Eo (standard electrode potential, in volts) for each of the two half reactions. The electric potential for the cell is equal to the more positive Eo value minus the more negative Eo value.

For example, in the picture above the solutions are CuSO4 and ZnSO4. Each solution has a corresponding metal strip in it, and a salt bridge or porous disk connecting the two solutions and allowing SO42− ions to flow freely between the copper and zinc solutions. In order to calculate the electric potential one looks up copper and zinc's half reactions and finds that:

Cu2+ + 2e → Cu (E = +0.34 V)
Zn2+ + 2e → Zn (E = −0.76 V)

Thus the reaction that is going on is really

Cu2+ + Zn → Cu + Zn2+

The electric potential is then +0.34 V −(−0.76 V) = 1.10 V

If the cell is operated under non-standard conditions, the potentials must be adapted using the Nernst equation.

Galvanic corrosion

Main article: Galvanic corrosion

Galvanic corrosion is a process that degrades metals electrochemically. This corrosion occurs when two dissimilar metals are placed in contact with each other in the presence of an electrolyte, such as salt water, forming a galvanic cell. A cell can also be formed if the same metal is exposed to two different concentrations of electrolyte. The resulting electrochemical potential then develops an electric current that electrolytically dissolves the less noble material.

Cell types

  • Concentration cell
  • Electrolytic cell
  • Electrochemical cell

See also

  • Electrode potential
  • Galvanic series
  • Alessandro Volta
  • Voltaic pile
  • Volt
  • Battery (electricity)
  • Electrosynthesis

External links

  • "Galvanic (Voltaic) Cells and Electrode Potential". Chemistry 115B, Sonoma.edu.
  • "Making and testing a simple galvanic cell". Woodrow Wilson Leadership Program in Chemistry, The Woodrow Wilson National Fellowship Foundation.
  • "Galvanic Cell" Good Animation

Retrieved from "http://en.wikipedia.org/wiki/Galvanic_cell"