Safety


Working with chemicals inherently involves some risk. In addition to the standard chemistry laboratory safety precautions, safety of electrochemistry has some of its own unique aspects.

For an organic electrochemistry reaction, the electrical current (mostly below 1 A) and voltage (rarely outside the window of –3 to +3 V) are typically within the capacity of a household battery. However, the potentiostats that deliver the power to a reaction are fed by the normal power grid, and are capable of distributing power at a level that can be harmful.

A compromised potentiostat can be a risk of its own. Do not use a potentiostat if it appears physically damaged or unable to perform a designated task (e.g., delivers 11 mA when it is supposed to deliver 10 mA). In general, the equipment should be fully disconnected before any attempts to replace or adjust electrodes or any other component. Never attempt maintenance or repair while instrument is connected to the power source. Do not touch the electrodes while conducting an experiment. Ensure electrodes do not come into contact with metal or with each other.

As with all chemical reactions, the Electrasyn 2.0 should be kept in a well-ventilated chemical fume hood in order to minimize exposure of chemicals to the operator and the risk associated with a buildup of organic solvent vapor. Adequate ventilation is essential for electrochemical oxidations coupled with cathodic proton reduction at large scale to avoid hydrogen gas buildup. In addition, solvents prone to peroxide generation should be avoided in electrochemical reactions.

It is recommended that all service and repair be performed by the equipment supplier. Unauthorized instrument modification is not recommended.

Please refer to the user manual of your potentiostat for additional safety precautions.

Electrodes


Electrodes must be conductive to carry the electric charge as part of the circuit; must elicit the surface electron exchange; and must be durable, such that they do not disintegrate easily. They are typically made of metals or carbon-based materials. Common electrode materials include platinum, silver, nickel, iron, copper, lead, mercury, zinc, aluminum, magnesium, and various conducting forms of carbon.

There are three basic forms of carbon commonly used as electrodes:
  1. Graphite, as in a pencil lead, is the most widely-found used. Many electrochemistry reactions are performed with graphite because the material is readily available and inexpensive. Graphite can be oxidized to graphene oxide in a strong oxidative environment, and it is known to adsorb organic compounds due to its porous nature.
  2. Vitreous carbon, also called glassy carbon, is a form of carbon with a fullerene-related structure, as opposed to layers of flat graphene sheets found in graphite. It is more stable to oxidative stress than graphite. Glassy carbon does not adsorb organic compounds, as it is not porous. Vitreous carbon produced as a foam is reticulated vitreous carbon (RVC). RVC electrodes come with a much larger surface area compared with the normal glassy carbon or graphite electrodes of the same geometric dimension. The cost of vitreous carbon is much more than graphite, but still insignificant compared with some of the expensive metal electrodes (e.g., Pt). RVC electrodes are among the most commonly used in organic electrochemistry.
  3. Boron-doped diamond (BDD) is a material derived from doping diamond with boron to obtain electrical conductivity. It is highly durable under redox stress and offers an extraordinarily large potential range and electrochemical power for both oxidation and reduction reactions. BDD electrodes are relatively new, expensive, and not yet commonly used in synthetic organic electrochemistry.

The electrodes commercially available on the market are primarily designated for analytical work. They are usually made with small and well-defined surface area for accuracy and sensitivity. Electrodes for synthetic organic chemistry, however, require large surface area. It is a common practice in the synthesis community to customize electrodes due to a lack of commercial availability. For example, one may use a coil made out of a piece of stainless steel wire from a hardware store as an electrode. The coil made by one chemist is rarely identical to the one made by another. Furthermore, there are many different grades of stainless steel wires available from suppliers around the world. Indeed, issues of reproducibility of an electrochemical reaction stems from important details like these. Part of IKA’s electrochemistry platform is its standardized electrodes, which are manufactured with precision from materials guaranteed to meet IKA’s rigid procurement and quality assurance protocols.

The standard hydrogen electrode (SHE) is a reference electrode with a potential declared zero to set up the basis of a thermodynamic scale of oxidation-reduction potentials, much like setting the freezing point of water at 0 °C to define the temperature scale. In practice, several other kinds of reference electrodes are used for convenience and precision, silver / silver(I) chloride electrode being one of the most popular. The silver / silver(I) chloride electrode is made of a silver wire that is coated with a thin layer of silver chloride and submerged in an aqueous potassium chloride solution of a defined concentration, all confined in a glass tube with a porous plug on one end. The layer of silver chloride coat can be applied by either electroplating or physically dipping the wire into molten silver chloride.

The potential of some reference electrodes are listed below:
  • Standard hydrogen electrode (SHE): 0.00 V
  • Silver / Silver(I) chloride electrode: 0.197 V (saturated KCl)
  • Saturated calomel electrode (SCE): 0.241 V
  • Copper / copper (II) sulfate electrode (CSE): 0.314 V

For a constant voltage reaction, the electrode potential is set with the help of a reference electrode. For example, a Shono oxidation at a constant +2.5 V anode potential in reference to a silver / sliver(I) chloride electrode means the anode potential is actually +2.697 V in reference to the SHE. In the literature, however, many constant voltage reactions are actually performed under the condition of a constant cell voltage rather than a constant electrode voltage. For example, a constant cell voltage of 5 V means the voltage difference between the anode and cathode is 5 V, which may or may not correlate to a +2.5 V anode potential in reference to the silver / sliver(I) chloride electrode or the SHE.

In a constant current reaction, a common oversight is letting the reaction run over an extended period of time without making sure the potential stays within reasonable boundaries. As the redox active species are depleted from the system, the voltage of a constant current reaction increases. If left unattended, this voltage will eventually reach a point to destroy the electrodes and likely obliterate products.

Solvents


There are two important considerations in the selection of a solvent used in an organic electrochemistry reaction:
  1. It must have sufficient dielectric constant and dissolution power to effectively dissolve the ionic electrolytes (as well as the organic substrates and various reagents involved). This is essential to facilitate the productive movement of radical-anions and -cations, and to maintain a balanced overall charge distribution in the reaction mixture such that electrons can freely flow from the anode to cathode.
  2. The solvent itself must be sufficiently stable to sustain the electrical redox environment within the electrochemistry cell, without significant participation in the redox process and without generating reactive species that may interfere with the intended reaction. The usable potential range of a solvent should cover the parameters of the intended electrochemical reaction.

For an anodic oxidation which is coupled with a proton reduction, the pH of the reaction media becomes a relevant factor. If a protic solvent is used and it is the only source of proton in such a system, then the pKa of the solvent becomes a consideration as well. Insufficient acidity can lead to sluggish cathodic reduction which may hamper the intended anodic transformation.

Common solvents used in electrochemistry organic reactions include acetonitrile, dimethyl formamide (DMF), dimethyl acetamide (DMA), N-methyl pyrrolidinone (NMP), dimethyl sulfoxide (DMSO), methanol, ethanol, 2,2,2-trifluoroethanol, 1,1,1,3,3,3-hexafluoroisopropanol (HFIP), acetic acid, and dichloromethane (DCM). Chart below provides additional details on solvent selection (Nicewicz, et al., Synlett. 2016, 27, 214):

Electrolytes


Charged species formed at the surface of the electrodes need to move into the bulk solution in order to participate in the desired transformation. For example, the radical cations formed on the surface of an anode in an electrochemical oxidation have a tendency to move toward the cathode (due to attraction of opposite charges), where the radical cations can be reduced and effectively nullify the chemical transformation. The presence of cations from electrolyte in the system slows down this migration, thwart the unintended cathodic reduction, and promote a more productive outcome for the electrochemical oxidation.

Electrolytes, ionic salts with adequate solubility and redox stability, are essential for electrochemical transformations. The ions of electrolytes are typically bulky, weakly coordinating, and inert to electrophiles or nucleophiles. Common anions include perchlorate, tetrafluoroborate, hexafluorophosphate, nitrate, aromatic sulfonate, carboxylate, while common cations include lithium ions and tetraalkylammonium ions. In general, lithium perchlorate (LiClO4), tetrabutylammonium tetrafluoroborate (Bu4NBF4), tetraethylammon ium tosylate (Et4NOTs), and tetrabutylammonium acetate (Bu4NOAc) are the most widely used electrolytes. Ionic liquids, with attributes of electrolyte as well as the ability to function as solvent, have also found use in electrochemical reactions.

One question people might ask is how big of an inconvenience it is to remove the electrolytes during product isolation. If an aqueous wash is part of the workup procedure, the electrolyte is typically completely removed. If a reaction is performed in a polar solvent (e.g., DMSO) and reversed-phase column chromatography is employed for product isolation (a common practice for small-scale organic reactions), the reaction mixture can often be loaded directly on to a (semi)preparative column without any workup procedures. Electrolytes are typically removed from the column in the solvent void in both analytical and preparative reversed phase HPLC.

Overpotentials


Overpotential is the potential difference between a half-reaction’s thermodynamically-determined potential and the potential at which the redox event is experimentally observed.

For example, hydrogen reduction potentials on a platinized platinum electrode (part of the SHE) and other materials (e.g., zinc) are not the same, despite the similarities of the thermodynamic reaction. This potential comes at –0.07 V for platinized platinum and –0.77 V for zinc. This makes sense intuitively if we remember the drastic difference of platinum vs. zinc as catalysts for hydrogenation. This deviation reflects the different surface phenomenon kinetics of the various species of hydrogen (hydrogen gas, radicals, and protons) on a specific material, platinized platinum and zinc in this case.

If an organic electro-oxidation is coupled with a cathodic proton reduction (e.g., Shono oxidation in MeOH), hydrogen overpotential must be considered. Cathode materials with a low hydrogen overpotential (smaller negative number) can help to ensure an adequate counter reaction without an overly harsh reductive environment around the cathode which could complicate reaction outcomes. Platinum (–0.07 V), palladium (–0.07 V), gold (–0.09 V), iron (–0.15 V), silver (–0.22 V) and nickel (–0.28 V) are preferred choices of cathode in this case. On the other hand, if the intended organic reaction is a cathodic reduction, then a cathode with high hydrogen overpotential will help to suppress the competition of proton reduction. Preferred cathode materials in this case include zinc (-0.77 V), lead (-0.71 V), and graphite (-0.62 V).

For anodic oxidation it is also common to use an anode made of materials with large oxygen overpotentials to suppress competitive water oxidation. Water may be introduced via electrolyte or the open atmospheric environment many electrochemical reactions are conducted under. These materials with large oxygen overpotentials include gold (+1.02 V), graphite (+0.95 V), shiny platinum (+0.95 V), palladium (+0.93 V), silver (+0.91 V), and lead (+0.81 V). If the intended organic reaction is a cathodic reduction coupled with water oxidation, materials with low oxygen overpotentials are justified (e.g., nickel +0.56 V). In practice, because water oxidation comes at a relatively high standard potential already (+1.23 V), this reaction may not be the first choice as a counter reaction in an organic electrochemistry reduction.

How-to Guides


How do I select electrodes?
(Courtesy of Dr. Yu Kawamata, Scripps Research)




How do I clean electrodes?
Download Electrodes Manual

How do I set up an electrochemical reaction on ElectraSyn 2.0?


How do I run a CV on ElectraSyn 2.0?


How do I optimize an electrochemical reaction?
(Courtesy of Dr. Yu Kawamata, Scripps Research)