Harnessing Electron Transfer in the Laboratory
Electrochemistry sits at the intersection of chemistry and electrical engineering. It studies chemical reactions that involve the transfer of electrons and the relationship between chemical energy and electrical energy. From batteries and fuel cells to corrosion protection and metal plating, electrochemical processes are ubiquitous in modern technology. In the laboratory, electrochemical techniques provide powerful methods for analyzing solutions, measuring thermodynamic properties, and synthesizing materials that cannot be made by other means.
Electrochemical Cell Construction
An electrochemical cell consists of two electrodes immersed in one or more electrolyte solutions, connected externally by a wire and internally by an ionic conductor (a salt bridge or porous frit).
Galvanic (voltaic) cells convert chemical energy to electrical energy — they are batteries. The spontaneous redox reaction drives electron flow through the external circuit. The anode (where oxidation occurs) is the negative terminal; the cathode (where reduction occurs) is the positive terminal.
Electrolytic cells use an external power supply to drive a non-spontaneous reaction. Electrode polarity is reversed: the anode is positive and the cathode is negative. Electrolysis is used for electroplating, water splitting, metal refining, and the industrial production of chlorine, sodium hydroxide, and aluminum.
Building a laboratory cell:
- Select electrode materials appropriate for the measurement. Platinum is the most versatile (inert, high overpotential for hydrogen and oxygen evolution). Glassy carbon is common for organic electrochemistry. Gold, silver, and mercury are used for specific applications.
- Polish solid electrodes with alumina slurry (0.3 and 0.05 micrometer) on a polishing pad, then rinse thoroughly with deionized water and sonicate for 30 seconds. Surface contamination dramatically affects electrochemical behavior.
- Prepare the electrolyte solution. For aqueous work, a supporting electrolyte (0.1-1.0 M KCl, KNO3, or H2SO4) provides ionic conductivity and minimizes migration effects. For non-aqueous work, tetrabutylammonium hexafluorophosphate (TBAPF6) in acetonitrile or dichloromethane is standard.
- Assemble the cell with three electrodes: working electrode (where the reaction of interest occurs), reference electrode (provides a stable potential), and counter electrode (completes the circuit). Use a potentiostat to control and measure the working electrode potential.
Reference Electrodes
A reference electrode maintains a constant, known potential against which the working electrode potential is measured. Without a stable reference, electrochemical measurements are meaningless.
Common reference electrodes:
| Reference Electrode | Abbreviation | Potential vs. SHE | Medium |
|---|---|---|---|
| Standard Hydrogen Electrode | SHE | 0.000 V (by definition) | H2(g) / H+(aq) |
| Saturated Calomel Electrode | SCE | +0.241 V | Hg / Hg2Cl2 / KCl(sat) |
| Silver/Silver Chloride | Ag/AgCl | +0.197 V (sat. KCl) | Ag / AgCl / KCl |
| Mercury/Mercurous Sulfate | MSE | +0.640 V | Hg / Hg2SO4 / K2SO4(sat) |
The Ag/AgCl electrode is the most widely used in modern laboratories because it is simple to construct, maintains a stable potential, and avoids the mercury used in SCE. It consists of a silver wire coated with silver chloride, immersed in saturated KCl solution, separated from the test solution by a porous frit.
Maintaining reference electrodes: Store immersed in their filling solution. Check the potential against a second reference periodically. Replace the filling solution regularly. A drifting reference electrode is the most common source of irreproducible electrochemical results.
Cyclic Voltammetry (CV)
Cyclic voltammetry is the single most informative electrochemical technique for characterizing redox-active species. The working electrode potential is swept linearly from an initial value to a switching potential, then swept back to the starting value, while the current is recorded.
What a cyclic voltammogram reveals:
- Peak positions — The anodic peak potential (Epa) and cathodic peak potential (Epc) indicate the formal potential of the redox couple: E0' = (Epa + Epc) / 2. For a reversible system, the peak separation is approximately 59/n mV at 25 degrees C, where n is the number of electrons transferred.
- Peak currents — Proportional to the concentration of the electroactive species and the square root of the scan rate (for diffusion-controlled processes). The Randles-Sevcik equation quantifies this relationship.
- Reversibility — A chemically reversible process shows both oxidation and reduction peaks with similar heights. An irreversible process shows only one peak or highly asymmetric peaks.
- Coupled chemical reactions — Changes in peak shape and position as a function of scan rate reveal chemical reactions coupled to the electron transfer (EC, CE, and ECE mechanisms).
Typical experimental conditions: Scan rates of 10-200 mV/s for initial characterization. Faster scans (1-100 V/s) probe reaction kinetics. Start with a freshly polished electrode and deaerated solution (purge with N2 or Ar for 10-15 minutes to remove dissolved O2, which has its own reduction peaks near -0.3 V vs. Ag/AgCl).
Electrolysis Setup
Preparative electrolysis uses sustained electrolysis at a controlled potential or current to transform a significant quantity of material. Applications include organic electrosynthesis, metal deposition, and water splitting.
Controlled-potential electrolysis uses a potentiostat to maintain the working electrode at a fixed potential. This ensures selectivity — only species that are reducible (or oxidizable) at that potential will react. The current decreases exponentially as the substrate is consumed.
Controlled-current electrolysis passes a constant current through the cell. It is simpler to implement but less selective — as the substrate concentration drops, the potential shifts to whatever value is needed to maintain the current, potentially triggering unwanted side reactions.
Practical considerations:
- Use high-surface-area electrodes (platinum mesh, reticulated vitreous carbon, or carbon felt) to maximize conversion rate.
- Separate anodic and cathodic compartments with a frit or membrane if the products of one electrode can be consumed at the other.
- Stir the solution vigorously to replenish the diffusion layer.
- Monitor charge passed (coulombs = current x time) to determine conversion. Faraday's law relates charge to moles of product: n = Q / (n_e x F), where F = 96,485 C/mol.
Safety in Electrochemistry
Electrochemistry introduces hazards beyond those of ordinary bench chemistry:
- Hydrogen gas — Generated at cathodes during aqueous electrolysis. Hydrogen is flammable and explosive at 4-75% concentration in air. Ensure adequate ventilation.
- Chlorine gas — Generated at anodes from chloride-containing electrolytes. Toxic and corrosive.
- Mercury — Used in some reference electrodes and polarographic methods. Handle with extreme care; clean spills immediately with commercial mercury spill kits.
- Electrical shock — Potentiostats operate at voltages that are generally safe, but high-current electrolysis setups require proper grounding and insulation.