Electrolysis can change electrode mass because electrons drive reduction at one electrode and oxidation at the other. Faraday’s laws connect electrical charge to moles of electrons, letting you predict plating or dissolution quantitatively.

Mass changes at electrodes: the core idea

When a species is reduced onto an electrode, that electrode gains mass; when an electrode material is oxidised into solution, that electrode loses mass. In both cases, the link between electricity and chemical change is electron counting.

Electroplating vs dissolution

Electrode mass changes depend on which half-reaction occurs at that surface.

Electroplating typically involves a metal cation (like $\text{Ag}^+$, $\text{Cu}^{2+}$, $\text{Ni}^{2+}$) gaining electrons and becoming a metal solid, increasing the electrode’s mass.

Electroplating cell diagram showing a silver anode dissolving to form $\text{Ag}^+$ in solution while $\text{Ag}^+$ is reduced and deposited as Ag(s) onto the spoon cathode. The labeled electron flow through the external circuit highlights that reduction occurs at the cathode (mass gain) and oxidation occurs at the anode (mass loss). Source

Dissolution is the reverse: a metal solid becomes ions, decreasing the electrode’s mass.

Using Faraday’s laws to connect charge to mass

Faraday’s laws state that the amount of chemical change is proportional to the total charge passed. The essential workflow is: current and time → charge → moles of electrons → moles of substance → mass change.

A key point is that the balanced half-reaction tells you how many electrons per mole of metal (or other species) are transferred.

This calculation only works when you correctly identify $z$ from the relevant half-reaction at that electrode.

Writing the half-reaction to get the electron ratio

Mass changes are tied to the electrode reaction:

Diagram of an electrolytic cell with labeled half-reactions, showing oxidation at the anode and reduction at the cathode driven by an external power source. This kind of schematic is useful for mapping each electrode’s half-reaction to the electron count $z$, which is the key link between charge passed and predicted mass change. Source

Metal deposition at the cathode (mass increases)

• General form: $\text{M}^{z+}(aq) + z e^- \rightarrow \text{M}(s)$

• The electrode gains mass because solid metal forms on the surface.

• The moles of metal deposited equal $\dfrac{n(e^-)}{z}$.

Metal dissolution at the anode (mass decreases)

• General form: $\text{M}(s) \rightarrow \text{M}^{z+}(aq) + z e^-$

• The electrode loses mass because metal atoms leave the solid as ions.

• The moles of metal dissolved equal $\dfrac{n(e^-)}{z}$.

What the balance of electrons implies for both electrodes

In an electrolytic setup, the same current passes through both electrodes, so the same total moles of electrons are transferred at the anode and cathode. However, mass changes can differ if:

• different substances react at the two electrodes, or

• $z$ differs (e.g., $\text{Ag}^+$ needs 1 $e^-$, but $\text{Cu}^{2+}$ needs 2 $e^-$).

Common pitfalls that change the predicted mass

• Using the overall (molecular) equation instead of the electrode half-reaction to find $z$.

• Forgetting to convert time to seconds when using $q=It$.

• Confusing the sign of mass change:

  • cathode (reduction): typically mass increases for metal-ion plating

  • anode (oxidation): typically mass decreases for metal dissolution

• Ignoring that the deposited/dissolved species must match the half-reaction (use the correct molar mass $M$).