The Faraday constant can be thought of as theproportionality factor between the charge incoulombs (used in physics and in practical electrical measurements) and theamount of substance inmoles (used in chemistry), and is therefore of particular use inelectrochemistry, particularly inelectrolysis calculations. Because the elementary charge is exactly1.602176634×10−19 C,[1] and there are exactlyNA = 6.02214076×1023 entities per mole,[1] the Faraday constant is given by the product of these two quantities:
F =e ×NA
=1.602176634×10−19 C ×6.02214076×1023 mol−1
=9.64853321233100184×104 C/mol.
The value ofF was first determined in the 1800s by weighing the amount ofsilver deposited in an electrochemical reaction, in which a measuredcurrent was passed for a measured time, and usingFaraday's law of electrolysis.[2] Until about 1970, the most reliable value of the Faraday constant was determined by a related method of electro-dissolving silver metal inperchloric acid.[3]
Related to the Faraday constant is the "faraday", a unit ofelectrical charge. Its use is much less common than of thecoulomb, but is sometimes used in electrochemistry.[4] One faraday of charge is the charge of onemole ofelementary charges (or of negative one mole of electrons), that is,
1 faraday = F × 1 mol = 9.64853321233100184×104 C = N0 × e = 6.02214076×1023e.
WhereN0 isAvogadro's number, the unitless counterpart toNA. Conversely, the Faraday constantF equals 1 faraday per mole. Thefarad is an unrelated unit ofcapacitance,1 farad = 1 coulomb / 1 volt.
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