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A nineteenth-century electrolytic cell, designed for the forceful production of oxyhydrogen. A quaint reminder that we've been bullying atoms for a long time.
An electrolytic cell is a type of electrochemical cell that fundamentally operates on coercion. It harnesses an external source of electrical energy to force a chemical reaction to proceed in a direction it naturally would not—a process bluntly named electrolysis.[1]:64, 89 [2]:GL7 Imagine a boulder that refuses to roll uphill. The electrolytic cell is the machine you build to shove it.
Inside this contraption, a voltage is imposed across two electrodes—an anode, which is assigned a positive charge, and a cathode, assigned a negative charge. These are submerged in an electrolyte solution, a conductive soup of ions waiting to be told where to go.[1]:89 [3][page needed] This setup is the philosophical opposite of a galvanic cell, the far more agreeable device that forms the basis of batteries. A galvanic cell taps into a spontaneous chemical reaction, converting the inherent chemical energy into a useful electrical flow.[1]:64
The defining difference lies in thermodynamics, a concept people find complicated but is really just about whether the universe approves. In an electrolytic cell, the net reaction is non-spontaneous, meaning its Gibbs free energy is positive. It's an energetically unfavorable process that will not happen without a forceful input of energy. Conversely, in a galvanic cell, the reaction is spontaneous; its Gibbs free energy is negative, and it proceeds with the universe's blessing, releasing energy as it goes.[3][page needed] One cell is a power plant; the other is an energy sink.
Principles
Look, the universe prefers order to descend into chaos, not the other way around. To reverse that trend, you need to apply force. In an electrolytic cell, an external voltage pushes a current through the cell, initiating a chemical reaction that would otherwise remain stubbornly inert. In a galvanic cell, it’s the spontaneous reaction itself that eagerly pushes a current out into a circuit.
Between these two states of doing and being done to lies the delicate, almost poetic, state of electrochemical equilibrium. Here, the cell's natural tendency to generate a current is perfectly counteracted by an externally applied voltage, a so-called back-EMF. The result is a standoff. No current flows. No net reaction occurs. It's a moment of perfect tension. If you increase that external voltage just slightly, you overwhelm the cell's natural inclination, and it becomes an electrolytic cell, forcing the reaction backward. If you decrease it, the cell's spontaneous nature takes over, and it behaves as a galvanic cell.[4]:354
Every electrolytic cell requires three basic components:
- An electrolyte: This isn't just any liquid. It's typically a solution, often using water or other solvents, in which salts or acids have dissociated into mobile ions. Alternatively, molten salts, like liquefied sodium chloride, can serve the same purpose, providing a chaotic sea of charged particles.
- Two electrodes: A cathode and an anode, the conductive terminals that connect the external power source to the electrolyte.
When the external voltage is applied, the system lurches into motion. The ions within the electrolyte, previously wandering aimlessly, are now drawn inexorably toward the electrode bearing the opposite charge. At these electrode surfaces, the real business happens: charge-transfer reactions, more formally known as redox reactions. It is only with an external voltage of the correct polarity and sufficient strength that an electrolytic cell can crack open a normally stable or chemically inert compound. The supplied electrical energy is converted directly into chemical energy, stored in the bonds of the newly formed products.
Michael Faraday, who had a talent for naming things, defined the cathode as the electrode where cations (ions with a positive charge, like Ag+) migrate to be reduced. Reduction is the gain of electrons, so these cations take electrons from the cathode. He defined the anode as the electrode where anions (ions with a negative charge, like Cl−) migrate to be oxidized, which means they surrender their electrons to the anode.
This is where people get confused. In an electrolytic cell, the anode is positive (to attract anions) and the cathode is negative (to attract cations). In a galvanic cell, this polarity is reversed because the cell is producing the current, not consuming it. In a galvanic cell, the anode is the source of electrons (negative terminal) and the cathode is where they end up (positive terminal). Just remember: in electrolysis, you are defining the charges with your power supply.
Applications
A video describing the process of electrolytic reduction as used on Captain Kidd's Cannon at The Children's Museum of Indianapolis. Even pirate artifacts can be subjected to forced chemistry.
The primary use of electrolytic cells is to decompose chemical compounds, a process aptly named electrolysis—from electro, for electricity,[5] and the Greek word lysis, meaning "to break up." It is chemistry via brute force.
Key examples include:
- Decomposition of Water: By passing a current through water (usually with a bit of salt or acid to make it conductive), it can be split into its constituent elements: hydrogen gas and oxygen gas. A fundamentally important reaction, especially if you're trying to produce hydrogen fuel.
- Aluminum Production: The Hall-Héroult process uses electrolysis on a massive industrial scale to reduce alumina, which is extracted from bauxite ore, into pure aluminum. This process is so energy-intensive that aluminum was once more valuable than gold.
- Electroplating: A technique for applying a thin metallic coating to an object. A cheap base metal can be plated with a layer of copper, silver, nickel, or chromium to improve its appearance or protect it from corrosion. It is the chemical equivalent of putting on a fancy coat.
- Industrial Metal Production: Electrolytic cells are the workhorses for the electrorefining and electrowinning of many non-ferrous metals. The vast majority of the world's high-purity aluminum, copper, zinc, and lead are churned out by these energy-hungry industrial cells.
Let's dissect a few of these processes more closely.
When you electrolyze water, especially saline or acidic water, the external voltage drives a specific set of reactions. Positively charged hydrogen ions (H+) are herded to the negatively charged cathode. There, they accept electrons and are reduced to form hydrogen gas (H₂). Meanwhile, at the positively charged anode, negatively charged hydroxide ions (OH−) are stripped of their electrons in an oxidation reaction, producing oxygen gas (O₂).
The electrolysis of molten sodium chloride (NaCl) is a starkly simple example. At a temperature high enough to melt table salt, a current is passed through the liquid. The anode, being positive, attracts the chloride ions (Cl−). It oxidizes them, forcing them to give up their electrons and form chlorine gas. The cathode, being negative, attracts the sodium ions (Na+). It reduces them, forcing them to accept electrons and deposit as pure, highly reactive sodium metal.
Things get more complicated when sodium chloride is dissolved in water. Now, water molecules are also present and can compete in the reactions. At the anode, chloride ions are still oxidized to produce chlorine gas. Depending on the pH, this can also lead to the formation of Hypochlorous acid. But at the cathode, a crucial difference emerges. Water is easier to reduce than sodium ions. So, instead of producing sodium metal, the cathode reduces water molecules, producing hydrogen gas and hydroxide ions (OH−). The final tally is chlorine gas and an acidic anolyte at the anode, and hydrogen gas and a basic catholyte of aqueous sodium hydroxide (NaOH) at the cathode. This entire procedure is known as the chloralkali process, a cornerstone of the modern chemical industry, producing three critical industrial chemicals from simple saltwater.