Chemistry
The Principle of Electrogravimetry for Quantitative Metal Analysis
Quick fact
In electrogravimetry, the mass of the deposited metal is used to calculate the original concentration, and the technique can achieve accuracies on the order of 0.1%–0.5% when performed correctly. Unlike many modern instrumental methods, it requires no calibration curve because the measurement is directly linked to the fundamental chemistry.
Why this is interesting
You have seen electroplating—like a shiny chrome bumper or a gold-plated connector—but did you know that the very same principle can be used to measure the exact amount of a metal in a solution?
Read the full explanation
Understanding The Principle of Electrogravimetry for Quantitative Metal Analysis
Imagine you have a solution containing dissolved copper ions and you want to know exactly how much copper is present. In electrogravimetry, you immerse two electrodes into the solution and apply a voltage. The copper ions are reduced at the cathode, forming a solid copper deposit that sticks to the electrode. If you weigh the cathode before and after the experiment, the increase in mass tells you how much copper was in the solution. To ensure that only copper deposits and not other metals, you carefully control the voltage, choosing a potential that is sufficient to reduce copper but not other ions. The process continues until there is no more copper left in the solution, which you can check by testing a drop of the solution with a chemical that reacts with copper ions. Once the deposition is complete, you wash, dry, and reweigh the electrode. The difference in mass is directly the mass of copper in your original sample. This method is time-honored, precise, and relies on the stoichiometry of the electrochemical reaction: each copper ion (Cu2+) requires two electrons to become a copper atom.
A deeper explanation
The principle of electrogravimetry is rooted in Faraday's law of electrolysis. During electrodeposition, the amount of substance deposited at an electrode is directly proportional to the quantity of electricity (charge) passed. For a metal ion Mn+ that is reduced to metal M, the half-reaction is Mn+ + ne− → M(s). The charge Q (in coulombs) is the product of current (I) and time (t), and it relates to the number of moles of electrons (n) by Q = nF, where F is the Faraday constant (96,485 C/mol). The moles of metal deposited are Q/(nF), and the mass is moles multiplied by the molar mass. In practice, the experiment is set up with a three-electrode system (working, reference, and auxiliary electrodes) to maintain a constant working electrode potential. The potential is chosen from a voltammogram to lie within the limiting-current plateau of the target metal, ensuring that the current is diffusion-limited and the deposition efficiency approaches 100%. The solution is stirred to maintain a steady supply of metal ions to the electrode surface. As the metal ions are depleted, the current decays to a low background level, signaling completion. After deposition, the electrode is washed (often with distilled water and sometimes a volatile solvent) and dried to constant mass. The mass difference gives the analyte amount, and from the original sample volume or mass, the concentration is calculated. Key factors that affect accuracy include co-deposition of interfering ions, incomplete deposition, and loss of deposited metal during rinsing. These are mitigated by careful potential selection, use of a suitable electrolyte and complexing agents, and proper electrode pretreatment. Electrogravimetry is particularly well-suited for metals that form adherent, smooth deposits, such as copper, nickel, silver, and lead, and it serves as a reliable benchmark for validating other analytical methods.