Dissolve a carbamate-protected pyrrolidine in methanol with a little supporting salt. Dip two carbon electrodes into the solution and connect them to a power supply. A few hours later, the carbon next to nitrogen carries a methoxy group. No chromium, peroxide, or hypervalent iodine was added. The oxidant was the anode itself.

This is the Shono oxidation, and it raises questions worth answering with care. How does a surface held at a fixed voltage select one C–H bond out of many nearly identical ones? Where do the removed electrons go? And what reaction on the other electrode keeps the circuit, and the stoichiometry, balanced?

Electrochemical synthesis treats the electron as a reagent. It has no molecular weight, leaves no spent byproduct, and has a reactivity you can adjust with a dial. The technique is old, going back to Faraday and Kolbe in the nineteenth century, but it is becoming a serious tool in modern process chemistry. Following the pathway from applied potential to bond formation shows why.

Every chemical oxidant or reductant is, in effect, a stored electrode potential sold in a bottle. Cerium(IV) sits around +1.6 V versus the standard hydrogen electrode in acidic media. Lithium metal reaches roughly −3.0 V. An electrode replaces the bottle with a variable. The relation ΔG = −nFE converts voltage into driving force. Each volt applied to a one-electron transfer corresponds to about 96.5 kJ/mol, so small shifts in potential produce large changes in reactivity.

Matching conditions to a transformation starts with cyclic voltammetry. Sweeping the potential reveals where a substrate begins to give up or accept electrons. The goal is to work just beyond that onset, while staying short of the potentials at which the product, the solvent, or other functional groups react. Acetonitrile offers a wide window. Methanol oxidizes more readily, but it can serve as a nucleophile. A supporting electrolyte such as Et₄NBF₄ carries the current without taking part in the chemistry.

There are two ways to control a cell. Potentiostatic operation fixes the electrode potential against a reference electrode, which gives the best selectivity. Galvanostatic operation fixes the current. It is simpler and scales more easily, but the potential drifts to whatever species is easiest to oxidize or reduce at that moment. The electrode material matters as much as the voltage. Platinum, graphite, and lead impose very different kinetic barriers, or overpotentials, on the same reaction. Choosing an electrode is therefore much like choosing a catalyst.

Two accounting rules govern every electrolysis. First, stoichiometry is measured in charge. A two-electron oxidation of one mole of substrate needs 2 faradays, about 193,000 coulombs, or roughly 54 hours at one ampere. Second, every oxidation at the anode is paired with a reduction at the cathode. That counter-reaction is often hydrogen evolution or metal dissolution, and it must be designed with the same care as the target step. It may also require a divided cell to keep intermediates apart.

Takeaway

A chemical reagent offers one fixed potential, while an electrode offers a continuous scale. Selectivity becomes a matter of choosing the right voltage rather than finding the right molecule.

At the anode, a substrate near the electrode surface transfers an electron from its highest occupied orbital into the electrode. The product is a radical cation, a species that is strongly acidic at adjacent C–H bonds and strongly electrophilic. Its fate follows a recurring pattern often written as ECEC: electron transfer, chemical step, electron transfer, chemical step. Deprotonation, a second oxidation, and nucleophilic capture are the usual moves. The sequence in which they occur determines the product.

The Shono oxidation shows the pattern clearly. Removing an electron from the carbamate nitrogen lone pair weakens the neighboring α-C–H bonds through hyperconjugation. Loss of a proton gives an α-amido radical. That radical is easier to oxidize than the starting material, so it immediately loses a second electron to form an N-acyliminium ion, which methanol traps. At the cathode, methanol is reduced to hydrogen and methoxide, and the methoxide supplies the base the mechanism needs. The resulting α-methoxy carbamates are versatile precursors in alkaloid and pharmaceutical synthesis.

The Kolbe electrolysis shows how the electrode can redirect a pathway. Oxidizing a carboxylate gives a carboxyl radical, which rapidly loses CO₂ to leave an alkyl radical. On platinum at high current density, these radicals build up at the surface and dimerize. On graphite, the radical can instead be oxidized a second time to a carbocation, which leads to Hofer–Moest products such as ethers and alcohols. The substrate is the same, but a different surface produces a different reaction manifold.

Many oxidations work better indirectly, through a mediator. TEMPO is oxidized at the anode to its oxoammonium form, which converts primary alcohols to aldehydes and is then regenerated. Halide mediators work in a similar way, producing electrophilic bromine or iodine species in situ. These approaches replace Jones reagent, which leaves chromium(III) waste, as well as Dess–Martin periodinane and Swern conditions. Mediators also lower the effective overpotential and prevent electrode fouling. The main byproduct is hydrogen gas released at the opposite electrode.

Takeaway

An anode does not simply remove electrons. It creates reactive radical cations whose fate the chemist steers through choice of solvent, base, nucleophile, and electrode surface.

At the cathode, the flow of electrons reverses. An electron enters the lowest unoccupied orbital of an electron-poor substrate, such as a carbonyl, imine, activated alkene, or aryl halide, to form a radical anion. Proton availability then decides the outcome. Protonation followed by a second electron transfer gives a carbanion that is protonated again, which amounts to net hydrogenation. The main competitor is reduction of protons to H₂. For this reason, cathodes with high hydrogen overpotential, such as lead or glassy carbon, are often preferred.

The largest organic electrosynthesis ever commercialized relies on this logic. In the 1960s, Manuel Baizer at Monsanto developed the hydrodimerization of acrylonitrile to adiponitrile, a precursor to nylon-6,6. Reduction at the cathode generates a nucleophilic intermediate that adds to a second acrylonitrile molecule. Protonation then completes the C–C bond-forming dimerization. Quaternary ammonium salts form a hydrophobic layer at the electrode surface. This layer suppresses hydrogen evolution and limits premature protonation. The process runs on a scale of hundreds of thousands of tonnes per year.

Dissolving-metal reductions such as the Birch reduction depend on lithium or sodium in liquid ammonia at −33 °C to generate solvated electrons. These conditions are hazardous, cryogenic, and hard to scale. Electrochemical versions generate equivalent reducing power at the cathode instead. Recent work has shown that lithium ions, in the presence of suitable additives, can reduce arenes at room temperature without ammonia. In these systems, the cathode potential and the proton source together set how far the reduction proceeds.

For carbonyls, the cathode offers an alternative to LiAlH₄ and its pyrophoric handling and aluminium salt workup. When protons are scarce, the ketyl radical anion dimerizes to a pinacol. When protons are plentiful, it goes on to the alcohol. The counter-reaction usually involves a sacrificial magnesium or zinc anode, or the oxidation of an amine. A more ambitious design is paired electrolysis, in which both electrodes make useful products and no electron is wasted.

Takeaway

In cathodic chemistry, the proton is the hidden second reagent. Controlling where and when protons arrive determines whether an electron makes a new bond or simply adds hydrogen.

Seen mechanistically, electrosynthesis is the same organic chemistry with a different source of driving force. Radical cations, radical anions, iminium ions, and ketyls are familiar intermediates. What changes is that the energy creating them comes from a tunable potential instead of a stoichiometric reagent.

Several design principles follow from this. Measure substrate potentials before choosing conditions. Treat the electrode material as a catalyst. Plan the counter-electrode reaction as deliberately as the target step. Use mediators when direct electron transfer is too slow or too unselective.

Flow cells with narrow electrode gaps are now addressing the old problems of scale and resistance. As renewable electricity becomes cheaper, the electron is shifting from a curiosity to a practical reagent with almost no waste.