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Cyclic voltammetry sweeps an electrode potential forward and then reverses the sweep while recording electric current. The resulting current-potential curve probes electron transfer and chemical reactions coupled to it.

An oxidation peak followed by a reduction peak can indicate that an electrode-generated species survives long enough to react on the return sweep. Changing the scan rate changes the experimental timescale; loss of a return peak can therefore reveal a subsequent reaction, although transport and electrode kinetics must also be considered.

Electrocapillarity is the dependence of interfacial surface tension on an electrical potential difference. Charging an electrode–electrolyte interface changes its interfacial free energy.^[ref]

At fixed temperature, pressure and chemical potentials, the electrocapillary relation reduces towhere $\gamma$ is interfacial surface tension, $E$ the appropriately defined electrode potential and $\sigma$ the electrode's free surface charge density. Thus an experimental curve of $\gamma$ versus $E$ measures charge through its slope. If $C=d\sigma /dE>0$, then $d^2\gamma /d{E}^2=-C<0$; zero charge gives a local maximum under these conditions. Adsorption or changing solution composition adds chemical-potential terms, so a measured maximum is not an unconditional universal charge marker.

Oxidize zinc at the anode and reduce nitrate to nitric oxide at the inert platinum cathode. Multiplying the oxidation half-reaction by three and the reduction half-reaction by two cancels six electrons:The standard cell potential is the cathode reduction potential minus the anode reduction potential:The written direction is spontaneous under standard conditions, since ${\Delta }_r{G}^{\circ }=-6F{E}^{\circ }<0$. Multiplying a half-reaction changes its extensive Gibbs energy, but does not multiply its intensive electrode potential.

nickel is less noble than copper and can dissolve anodically, but its reduction requires a much more negative cathodic electrode potential. It therefore normally accumulates in solution rather than codepositing with copper. lead can also oxidize, but sulfate can precipitate the resulting lead ions as poorly soluble lead sulfate, contributing to anode slime or passivation. gold normally remains in the solid residue. silver is normally recovered predominantly in anode slime, but its path can include dissolution and reprecipitation or cementation: a permanent refusal of every silver atom to dissolve is not the mechanism. Experiments with controlled Cu–Ag–Pb anodes, a fixed sulfate electrolyte and SEM/EDS plus diffraction analysis explicitly find metallic silver and sulfate-containing slime phases.^[ref]

Selective electrode potential, solution management and physical exclusion of slime permit purer copper deposition. However, the unusually large printed anodic slope prevents an unconditional silver-retention claim here. The preceding calculation places the anode near $0.34+0.6788=1.0188$ V versus the standard hydrogen electrode, above the tabulated ${\mathrm{Ag}}^{+}$/Ag value. silver oxidation is then thermodynamically possible; any silver ions reaching the cathode can be reduced preferentially to copper. Actual activities, precipitation, cementation and transport determine contamination. The requested 99.999 wt% purity cannot be established from the listed standard potentials alone. Monitor anode potential against a reference, assay the electrolyte and deposit, and prevent slime entrainment; do not silently replace the printed 0.295 V/decade with a smaller number to remove this issue.

Gas formation is decided by each interfacial electrode potential and reaction kinetics, rather than by comparing the terminal cell voltage directly with 1.23 V. On the conditional model above, the cathode potential is $E_c\simeq 0.34-0.12854=+0.21146$ V versus the standard hydrogen electrode. At unit proton activity and unit hydrogen pressure this is above the hydrogen equilibrium potential, so hydrogen evolution is not favored there. The anode is at $E_a\simeq 1.0188$ V, below the unit-activity oxygen equilibrium potential of 1.23 V, so oxygen evolution is likewise not favored at the intended point. Neither gas is expected in that model, despite a terminal voltage near 2 V; 1.2 V is lost in the electrolyte.

Without the missing cathodic kinetic information, the exact cathode potential is not fixed, so the hydrogen conclusion is conditional. Locally depleted copper ions, increased current, electrode blocking or strong passivation can move an interface into gas evolution. Actual proton/oxygen/hydrogen activities shift equilibrium through the Nernst equation, and gas nucleation and activation barriers further affect observable bubbles. A reference electrode at each interface and direct gas observation distinguish these effects from an increased solution-resistance drop.

The species that oxidizes nickel is gold(III) ion, rather than metallic gold. The half-reactions areMultiplying the first by three and the second by two givesStandard electrode potentials are not multiplied by the balancing factors. Thus $E_{\mathrm{cell}}^{\circ }=1.50-(-0.25)=1.75\mathrm{V}$ and six moles of electrons are transferred per mole of this balanced reaction:The negative sign predicts a thermodynamically favorable galvanic replacement under standard conditions.