= Electrode
{wiki}
An <electrode> is an electronic <electrical conductor> in contact with an <electrolyte>. In a cell carrying sustained <redox> <electric current>, <chemical reactions> at this interface exchange <Electrons> with <chemical species>. An <electrode> can participate in the <chemical reaction>, as the dissolving <zinc> <anode> of a <Daniell cell> does, or provide a conducting surface without appearing in the net <chemical equation>. This distinction is illustrated by active and inert <electrodes> in https://openstax.org/books/chemistry-2e/pages/17-2-galvanic-cells[Flowers, Theopold, Langley and Robinson, Chemistry 2e, §17.2].
The interface connects two kinds of <electric charge> transport: <Electrons> in the external <electrical circuit> and <ions> in the <electrolyte>. The same <Electron> does not travel through the wire, swim through the <salt bridge>, and complete a circuit! The <half-reactions> account for how electronic and ionic transport meet. This is the useful idea behind both obtaining electrical <work> in a <galvanic cell> and supplying electrical <work> in an <electrolytic cell>.
Name an <electrode> from its net <chemical reaction>: <oxidation> means <anode>, and <reduction> means <cathode>. State the operating mode before assigning terminal signs. For a <galvanic cell> supplying a passive load, the <anode> is negative and the <cathode> positive; for an ordinary two-terminal <electrolytic cell> driven by a source, the <anode> is positive and the <cathode> negative. Here positive means higher terminal <electric potential> than the other <electrode>, not a claim that the bulk material has a large positive net <electric charge>. See https://doi.org/10.1515/pac-2018-0109[Pingarrón et al., Terminology of electrochemical methods of analysis (IUPAC Recommendations 2019), §§3.1 and 3.4].
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