Electrolysis can occur in nature
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Faraday already noticed that when a voltage is applied to salt or acid solutions in an electrochemical cell, chemical changes occur at the electrodes and that chemical reactions can be carried out with electricity. This process is known as electrolysis.
Separate reaction equations can be set up for the individual reactions at the electrodes in an electrochemical cell. A quantitative relationship is found between the chemical conversion and the amount of charge transported through the electrolyte in the form of electrons that have been taken up or released. If a current flows through the external circuit during the time, a quantity of charge is transported through it.
It must be equal to the charge exchanged at the cathode or anode at the same time, which is transported through the electrolyte solution of the internal circuit. If ions are discharged at an electrode over time, this corresponds to exchanged electrons. The following relationship results for the charge transported in the electrolyte.
- Faraday constant
- numerically corresponds to the charge of 1 electron. In recognition of Faraday's work, this quantity is called the Faraday constant. The Faraday constant is just the amount of charge of 1 electron.
In electrolysis, 1 ion equivalents, i.e. 1 ion equivalents (,), ½ ions (,) and ⅓ ions (,), are deposited by the amount of charge of a Faraday.
Summing up the equations for and one obtains:
This quantitative connection was already found empirically by Faraday. In 1832 and 1833 Faraday formulated two laws based on his experiments:
- First Faraday's Law
- The mass of the substances deposited on the two electrodes is proportional to the charge (current times time) conducted through the cell.
The determination of the mass of the deposited electrolysis product is a convenient means of measuring the amount of charge (electrogravimetry).
- Second Faraday's law
- The masses of different substances separated by the same amount of charge behave like the molar masses divided by the number of charges of the cell reaction (equivalent weights).
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