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| title | chunk | source | category | tags | date_saved | instance |
|---|---|---|---|---|---|---|
| Electrolysis | 2/5 | https://en.wikipedia.org/wiki/Electrolysis | reference | science, encyclopedia | 2026-05-05T10:47:31.576904+00:00 | kb-cron |
Solvation or reaction of an ionic compound with a solvent (such as water) to produce mobile ions An ionic compound melted by heating The electrodes are immersed separated by a distance such that a current flows between them through the electrolyte and are connected to the power source which completes the electrical circuit. A direct current supplied by the power source drives the reaction causing ions in the electrolyte to be attracted toward the respective oppositely charged electrode. Electrodes of metal, graphite and semiconductor material are widely used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and manufacturing cost. Historically, when non-reactive anodes were desired for electrolysis, graphite (called plumbago in Faraday's time) or platinum were chosen. They were found to be some of the least reactive materials for anodes. Platinum erodes very slowly compared to other materials, and graphite crumbles and can produce carbon dioxide in aqueous solutions but otherwise does not participate in the reaction. Cathodes may be made of the same material, or they may be made from a more reactive one since anode wear is greater due to oxidation at the anode.
=== Process of electrolysis === The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons due to the applied potential. The desired products of electrolysis are often in a different physical state from the electrolyte and can be removed by mechanical processes (e.g. by collecting gas above an electrode or precipitating a product out of the electrolyte). The quantity of the products is proportional to the current, and when two or more electrolytic cells are connected in series to the same power source, the products produced in the cells are proportional to their equivalent weight. These are known as Faraday's laws of electrolysis. Each electrode attracts ions that are of the opposite charge. Positively charged ions (cations) move towards the electron-providing (negative) cathode. Negatively charged ions (anions) move towards the electron-extracting (positive) anode. In this process electrons are effectively introduced at the cathode as a reactant and removed at the anode as a product. In chemistry, the loss of electrons is called oxidation, while electron gain is called reduction. When neutral atoms or molecules, such as those on the surface of an electrode, gain or lose electrons they become ions and may dissolve in the electrolyte and react with other ions. When ions gain or lose electrons and become neutral, they will form compounds that separate from the electrolyte. Positive metal ions like Cu2+ deposit onto the cathode in a layer. The terms for this are electroplating, electrowinning, and electrorefining. When an ion gains or loses electrons without becoming neutral, its electronic charge is altered in the process. For example, the electrolysis of brine produces hydrogen and chlorine gases which bubble from the electrolyte and are collected. The initial overall reaction is thus:
2 NaCl + 2 H2O → 2 NaOH + H2 + Cl2 The reaction at the anode results in chlorine gas from chlorine ions:
2 Cl− → Cl2 + 2 e− The reaction at the cathode results in hydrogen gas and hydroxide ions:
2 H2O + 2 e− → H2 + 2 OH− Without a partition between the electrodes, the OH− ions produced at the cathode are free to diffuse throughout the electrolyte to the anode. As the electrolyte becomes more basic due to the production of OH−, less Cl2 emerges from the solution as it begins to react with the hydroxide producing hypochlorite (ClO−) at the anode:
Cl2 + 2 NaOH → NaCl + NaClO + H2O The more opportunity the Cl2 has to interact with NaOH in the solution, the less Cl2 emerges at the surface of the solution and the faster the production of hypochlorite progresses. This depends on factors such as solution temperature, the amount of time the Cl2 molecule is in contact with the solution, and concentration of NaOH. Likewise, as hypochlorite increases in concentration, chlorates are produced from them:
3 NaClO → NaClO3 + 2 NaCl Other reactions occur, such as the self-ionization of water and the decomposition of hypochlorite at the cathode, the rate of the latter depends on factors such as diffusion and the surface area of the cathode in contact with the electrolyte.
=== Decomposition potential === Decomposition potential or decomposition voltage refers to the minimum voltage (difference in electrode potential) between anode and cathode of an electrolytic cell that is needed for electrolysis to occur. The voltage at which electrolysis is thermodynamically preferred is the difference of the electrode potentials as calculated using the Nernst equation. Applying additional voltage, referred to as overpotential, can increase the rate of reaction and is often needed above the thermodynamic value. It is especially necessary for electrolysis reactions involving gases, such as oxygen, hydrogen or chlorine.
=== Oxidation and reduction at the electrodes === Oxidation of ions or neutral molecules occurs at the anode. For example, it is possible to oxidize ferrous ions to ferric ions at the anode:
Fe2+(aq) → Fe3+(aq) + e− Reduction of ions or neutral molecules occurs at the cathode. It is possible to reduce ferricyanide ions to ferrocyanide ions at the cathode:
Fe(CN)3-6 + e− → Fe(CN)4-6 Neutral molecules can also react at either of the electrodes. For example: p-benzoquinone can be reduced to hydroquinone at the cathode:
- 2 e− + 2 H+ → In the last example, H+ ions (hydrogen ions) also take part in the reaction and are provided by the acid in the solution, or by the solvent itself (water, methanol, etc.). Electrolysis reactions involving H+ ions are fairly common in acidic solutions. In aqueous alkaline solutions, reactions involving OH− (hydroxide ions) are common. Sometimes the solvents themselves (usually water) are oxidized or reduced at the electrodes. It is even possible to have electrolysis involving gases, e.g. by using a gas diffusion electrode.