Electrochemical cell

An electrochemical cell combines electrode interfaces, where charge-transfer reactions occur, with an ion-conducting electrolyte. Separating oxidation and reduction allows electrons to pass through an external circuit instead of being transferred directly between reactants at one location. Ionic movement inside the cell complements electronic conduction outside it.

In a two-beaker demonstration, a salt bridge connects the solutions. It is not an electron-carrying wire: moving ions compensate for charge changes caused by the electrode reactions. Practical cells may use separators and other arrangements. The familiar pair of beakers and U-shaped tube illustrates a principle rather than defining a required shape.

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A galvanic, or voltaic, cell uses a spontaneous redox reaction under the relevant conditions to perform electrical work. In a Daniell cell, zinc becomes zinc ions while copper ions become copper metal. The overall reaction is Zn + Cu²⁺ → Zn²⁺ + Cu. Electrons travel through the external connection from the zinc side toward the copper side.

The electron balance is informative: one zinc atom releases two electrons and one copper ion takes up two. When the half-reactions are added, those electrons cancel. They link the separated reactions rather than appearing as a newly created final product. A cell is therefore not simply a container holding a finite pile of electrons to be poured out. For two zinc atoms reacting in the same way, four electrons must be transferred and two copper ions reduced. Doubling the reaction changes the amount of material and charge involved, but does not change which species is oxidized or which species is reduced.

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The anode is the electrode at which oxidation occurs; the cathode is where reduction occurs. These names describe chemical roles. Positive and negative describe relative electrode potential. During galvanic operation, the anode is negative and the cathode positive.

An electrolytic cell is driven by an external source. Its anode is connected to the source's positive terminal and its cathode to the negative terminal, reversing the polarity association found in a discharging galvanic cell. The reaction definitions remain unchanged: oxidation at the anode, reduction at the cathode. Identifying the operating process first avoids confusing these two naming systems.

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Electrolysis uses electrical work to drive chemical change, as in electroplating, production of some chemicals and battery charging. Its principal energy-conversion direction is opposite to galvanic discharge. An external source must supply sufficient driving force; merely placing electrodes in an electrolyte does not guarantee a desired reaction.

Rechargeable cells use reactions that can be reversed under suitable conditions, restoring their chemical composition to some extent. Primary cells are not designed for repeated charging. The fact that electrolysis can reverse some reactions does not establish that every battery is safe to recharge: practical reversibility also depends on materials and construction.

References: [2] [3]