Editing Faraday constant (section)

== Derivation ==
The Faraday constant can be thought of as the conversion factor between the mole (used in chemistry) and the [[coulomb]] (used in physics and in practical electrical measurements), and is therefore of particular use in [[electrochemistry]]. Because there are exactly ''N''<sub>A</sub> = {{val|6.02214076|e=23}} entities per mole,<ref name="SI2019" /> and there are exactly {{math|1={{sfrac|1|''e''}} = {{sfrac|10<sup>19</sup>|{{val|1.602176634}}}}}} elementary charges per coulomb,<ref name="SI2019" /> the Faraday constant is given by the quotient of these two quantities:
: {{math|1=''F'' = {{sfrac|''N''<sub>A</sub>|1/''e''}} = {{val|9.64853321233100184|e=4|u=C.mol-1}}.}}

One common use of the Faraday constant is in [[electrolysis]] calculations. One can divide the amount of charge (the current integrated over time) by the Faraday constant in order to find the [[amount of substance|chemical amount]] of a substance (in moles) that has been electrolyzed.

The value of {{math|''F''}} was first determined in the 1800s by weighing the amount of [[silver]] deposited in an electrochemical reaction, in which a measured [[Current (electricity)|current]] was passed for a measured time, and using [[Faraday's law of electrolysis]].<ref>[http://physics.nist.gov/cuu/Constants/historical1.html NIST Introduction to physical constants]</ref> Until about 1970, the most reliable value of the Faraday constant was determined by a related method of electro-dissolving silver metal in [[perchloric acid]].<ref name="IUPAC">{{Cite journal |last=IUPAC |author-link=International Union of Pure and Applied Chemistry |year=1976 |title=Status of the Faraday constant as an analytical standard |journal=Pure and Applied Chemistry |volume=45 |issue=2 |pages=125–130 |doi=10.1351/pac197645020125 |doi-access=free}}</ref>