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Electrolysis uses an external source of electrical work to drive redox chemical reactions at the electrodes of an electrolytic cell. Its characteristic application is forcing a chemical reaction in a direction that is nonspontaneous under the specified conditions: a discharging galvanic cell supplies electrical work, whereas electrolysis consumes electrical work. This makes electrolysis useful for producing otherwise difficult-to-obtain chemical substances, charging a rechargeable battery, and depositing a metal coating on a conducting object.^[ref]

Always identify the anode by oxidation and the cathode by reduction. In ordinary driven two-terminal operation, the source withdraws electrons from the positive anode and delivers electrons to the negative cathode. In a discharging galvanic cell, the anode is negative and the cathode positive. Thus terminal signs change with operating mode, while the reaction definitions remain valid. Through the electrolyte, ions transport electric charge; they need not all be consumed at an electrode. The electrolyte and electrolytic cell articles explain this transport and its conventional current direction.

For a net chemical equation transferring $z$ electrons per reaction event, the molar reaction Gibbs free energy obeyswhere $F$ is the Faraday constant and $E_{\mathrm{chem}}$ is the reversible cell electromotive force for the reaction in the written direction. A nonspontaneous direction has ${\Delta }_{r}G>0$ and $E_{\mathrm{chem}}<0$ at fixed temperature and pressure. The ideal opposing voltage magnitude is ${\Delta }_{r}G/(zF)$; sustaining a finite electric current generally needs additional drive because of electrical resistance and finite electrode reaction rates. This thermodynamic voltage concerns the actual composition, rather than necessarily standard conditions.^[ref] Not every use of electrolysis increases bulk chemical Gibbs free energy: transferring copper between equivalent surfaces can have zero bulk ${\Delta }_{r}G$ while the source still supplies dissipated electrical work. See the existing electrolytic cell example.

A particularly revealing experiment measures copper transfer instead of merely looking for a coating. Two separated copper electrodes stand in aqueous $\mathrm{CuS}{\mathrm{O}}_4$ electrolyte in a vessel, with insulated leads to a regulated direct-current power supply, a series ammeter and a switch. A balance reads each electrode's mass before and after a timed run. The Institute of Physics describes copper cells with ammeters, adjustable current and removable cathodes; its comparison uses three similar cells at $1\mathrm{A}$ for ten minutes, $0.5\mathrm{A}$ for ten minutes, and $0.5\mathrm{A}$ for twenty minutes. The specified observation to test is that deposited copper mass is proportional to electric current times time. These are documented experimental settings, not measurements generated here.^[ref]

The ideal half-reactions areEach deposited copper chemical atom requires two electrons, so $z=2$. The anode loses copper and the cathode gains it; ideally their mass changes have equal magnitudes, and the bulk copper cation inventory stays constant. Local concentrations near the electrodes can nevertheless differ. The sulfate anions participate in ionic transport without appearing in these net half-reactions; electronic electric current in the wires becomes ionic electric current in the electrolyte through the two interfaces.^[ref]

Clean, rinse, dry and identify the electrodes before their initial weighing. After interrupting the electric current, rinse away adhering electrolyte, dry again and weigh without rubbing off the deposit. KIT's university practical explicitly calls for distilled-water rinsing, careful drying and weighing both electrodes; it warns that loose copper grains can be lost during handling.^[ref] Useful additional controls are repeated drying and weighing until the reading stabilizes, and a matched unpowered immersion followed by the same rinse/dry cycle. These controls test retained liquid, residual dissolved salts, handling loss and chemical changes unrelated to passed electric charge; they are proposed checks, not additional observations reported by KIT. Keep immersed electrode area, spacing, electrolyte composition and temperature comparable across runs, and keep the electrodes from touching.

Record the electric current as a function of time, together with the exact switch-on and switch-off times. Taking $I(t)\ge 0$ in the chosen deposition direction, the transported electric charge isFor constant electric current, $Q=I(t_1-t_0)$; with sampled readings, use numerical integration rather than assuming that one ammeter reading remained constant. The two equal-charge Institute of Physics settings each give $600\mathrm{C}$, while its shorter $0.5\mathrm{A}$ setting gives $300\mathrm{C}$. With $M_{\mathrm{Cu}}=63.546\mathrm{g}\mathrm{mo}{\mathrm{l}}^{-1}$, the mass per amount of substance tabulated by KIT, and $F\approx 96485.33212\mathrm{C}\mathrm{mo}{\mathrm{l}}^{-1}$, ideal Faraday's laws of electrolysis predictThese are calculated predictions for complete conversion of the measured electric charge into retained copper, not experimental results. For a stable process, doubling the electric charge doubles the deposited mass; equal electric charge gives equal deposited mass even when the electric current and duration differ.

The balance reading is a test of the whole model. Competing reduction of water or other chemical species can divert electric current from copper deposition; transient electrode capacitance can also consume part of a short pulse without producing copper. A fragile deposit can detach even when its formation used the expected electric charge. Conversely, retained electrolyte or salts can increase an apparent mass gain. Comparing anode loss with cathode gain helps detect such problems but does not identify every cause. A slope below $M_{\mathrm{Cu}}/(2F)$ alone cannot distinguish competing chemical reactions from lost deposit. Current efficiency measures the fraction of transported electric charge assigned to a specified electrode reaction, rather than the fraction of electrical energy stored chemically.^[ref]

The beautiful connection is that an electrical measurement counts chemical events: integrating electric current and reading a balance provide two ways of following the same electron stoichiometry. This makes electrolysis useful for controlling coating quantity and for testing a charge-to-amount of substance relation. It does not determine coating uniformity, adhesion or the required voltage, which depend on transport and electrode chemical kinetics as well.

Faraday's laws of electrolysis state that the mass of material formed or consumed by a specified electrode reaction is proportional to the transported electric charge, and that equal electric charge produces masses proportional to $M/z$ for different reactions. Here $M$ is the product's mass per mole, and $z$ is the positive number of electrons transferred per product entity. For complete use of the electric charge in that reaction,The first relation uses the Faraday constant; the second uses half-reaction stoichiometry; the third converts amount of substance to mass. For equal transported electric charge in two ideal cells,The number $z$ belongs to the balanced half-reaction, rather than to the metal's name alone: copper deposited from ${\mathrm{Cu}}^{2+}$ needs $z=2$. These modern forms of both laws are given in KIT's practical.^[ref]

If other same-direction electrode reactions share the measured electric current, define an integrated current efficiency $\eta =Q_{\mathrm{product}}/Q$. Provided the product is neither subsequently consumed nor lost, the retained mass isHere $0\le \eta \le 1$ follows from the specified same-direction partition of electric current. It is not a universal bound on every signed partial-current convention: IUPAC notes that simultaneous oxidation and reduction at one electrode can produce a current-efficiency ratio above one.^[ref] When efficiency varies during a run, use the charge-weighted ratio above, rather than an unweighted average over time. If deposition removes material from a multicomponent chemical solution, specify which chemical species is being counted and balance its half-reaction before choosing $z$.

Michael Faraday's accessible 1834 paper, Experimental Researches in Electricity, Seventh Series, §783 states proportionality between chemical action and the quantity of electricity passed. In §§789–791 he describes connecting a molten tin-chloride cell in series with a water-decomposition charge indicator, weighing the receiving electrode before and after, and comparing the deposited metal with the indicator's gas readout. The common series electric current makes these two readouts comparable. This was a quantitative test of chemical equivalents, rather than merely an observation that a deposit appeared. It is historical evidence, not a suggested classroom procedure; the modern electron-counting derivation above uses concepts unavailable to Faraday.^[ref]

Michael Faraday connected chemistry with quantitative electricity through electrolysis. His experimental treatment related the amount of material transformed at an electrode to the electric charge passed through the apparatus, the basis of Faraday's laws of electrolysis.

His candle lectures made a different connection between apparatus and explanation: the wick supplies melted fuel through capillary action, while the flame's products reveal the chemical transformations involved in combustion.^[ref]

A chemist whose route into experimental research began with a bookbinding apprenticeship (1805–1812). He became a laboratory assistant at the Royal Institution in 1813, its laboratory director in 1825, and Fullerian Professor of chemistry from 1833 to 1867.^[ref]

Frictional electricity and voltaic electricity name different methods of producing electrical effects. Their identity is supported by showing common kinds of response and quantitative continuity between them, rather than merely observing that both can give a sensation.

Charges from either source show the same attraction and repulsion, can charge a conductor or capacitor, and can discharge through a circuit. A current from either source can generate a magnetic field, produce resistive heating, and cause appropriate electrolysis. Reversing the sign or current direction reverses the corresponding directional effects. Different observed strengths need not imply different electrical substances.

Michael Faraday carried out a comparison of common and voltaic electricity. In a specific magnetic test, he discharged a machine-charged jar battery through a galvanometer, using a wet thread in the discharge path; the needle's deflection connected ordinary static charge with a current's magnetic action. He also compared common and voltaic electricity quantitatively.^[ref]

The modern distinction is mainly in the source characteristics. A friction machine commonly supplies a small charge at high electric potential difference, while a chemical cell can sustain a much larger continuing electric current at a lower voltage. An insulating gap or a poorly conducting path can therefore affect the demonstrations differently. Once charge passes through a specified circuit, its electric and magnetic action is not determined by whether mechanical rubbing or a chemical reaction originally supplied the energy.

For a specified electrode reaction with controlled current efficiency, Faraday's laws of electrolysis express the quantitative relationHere $M$ is the product's molar mass, $z$ the electrons per reaction unit and $F$ the Faraday constant. Equal transferred charges give equal reaction equivalents under those conditions. Charge, voltage and current must be distinguished: equal voltages alone do not imply equal delivered quantities of electricity.

Compare the currents using the same calibrated detector and unchanged geometry. With a tangent galvanometer, a circular coil gives $B_c={\mu }_{0}NI/(2a)$ at its centre, perpendicular to the horizontal terrestrial field $B_H$. A small compass needle aligns with their resultant, soUse the same coil and location, set its plane in the magnetic meridian, and reverse each current to help identify zero offsets or unwanted local fields. The needle must be sufficiently small relative to the coil for the nearly uniform central field approximation. A finite deflection is not generally proportional to current; replacing the tangent ratio by the angle ratio requires a small-angle approximation or a different calibrated instrument.

Electrolysis offers another comparison. With identical suitable voltameters, the same reaction and constant currents acting for equal times, the deposited mass or collected reaction product is proportional to charge $Q=It$, according to Faraday's laws of electrolysis. Thus $I_1/I_2=m_1/m_2$ when the same product is collected with equal current efficiency. For unequal observation times, compare $m_1/t_1$ with $m_2/t_2$. Gas collection additionally requires matching pressure, temperature and vapour corrections. Current-dependent side reactions can invalidate a naive product ratio.

An essential circuit qualification is that adding the detector can change the current. Either compare the currents in the circuits as actually measured or account for detector resistance and source loading. Comparing two voltaic sources' currents through one specified circuit is not the same as establishing a ratio of their electromotive forces.