Electrolytic Processes

Overview of Electrolytic Processes

Electrolytic processes are electrochemical reactions that are driven by an external source of electrical energy. In contrast to galvanic (voltaic) cells, where a spontaneous redox reaction produces electrical energy, in electrolytic cells a non‑spontaneous redox reaction is forced to occur by applying a suitable voltage.

In this chapter the focus is on:

• How an electrolytic cell is set up and operated.
• How to predict what is formed at each electrode.
• How much substance is produced or consumed during electrolysis.
• Typical technical and everyday examples of electrolysis.

Basic ideas such as charge carriers in electrolytes, electrodes, and electrode potentials are assumed from the preceding chapters on electrochemical processes.

Construction and Operation of an Electrolytic Cell

An electrolytic cell consists essentially of:

• A power source (DC voltage supply or battery).
• Two electrodes (anode and cathode), which may be inert (e.g. Pt, graphite) or reactive.
• An electrolyte, containing mobile ions in either molten or aqueous solution form.
• External connections (wires) and, if needed, a diaphragm or membrane to separate products.

In electrolysis, the electrode definitions are determined by the direction of electron flow imposed by the power source:

• Cathode: Electrode connected to the negative pole of the power source (in most simple cells). Electrons are supplied here; reduction occurs at the cathode.
• Anode: Electrode connected to the positive pole. Electrons are withdrawn here; oxidation occurs at the anode.

Remember: cathode = reduction, anode = oxidation. In electrolytic cells, the cathode is typically negative, whereas in galvanic cells the cathode is positive. The sign changes, the redox roles do not.

Types of Electrolytes in Electrolysis

Two main forms are important:

• Molten salts (e.g. molten NaCl, Al$_2$O$_3$/cryolite mixtures): Contain only the ions of the salt; no solvent. Used at elevated temperatures.
• Aqueous electrolytes (salt solutions, acids, bases): Contain solute ions and water. Both solute ions and water can be oxidized or reduced, depending on the applied potential.

The presence of water in aqueous systems introduces competing half-reactions, making product prediction more involved than for molten salts.

Electrolysis of Molten Salts

For molten salts, the only redox-active species are the cation(s) and the anion(s). Predicting products is comparatively straightforward: cations are reduced at the cathode, anions are oxidized at the anode.

Example: Electrolysis of molten sodium chloride, NaCl(l)

• Composition: Na$^+$ and Cl$^-$ only.
• At the cathode (reduction of cations):
  $$\text{Na}^+ + e^- \rightarrow \text{Na}$$
• At the anode (oxidation of anions):
  $\ \text{Cl}^- \rightarrow \text{Cl}_2 + 2\,e^-$$

Overall reaction:
$$2\,\text{NaCl(l)} \rightarrow 2\,\text{Na(l)} + \text{Cl}_2(\text{g})$$

Key features:

• Metals that are too reactive to be obtained by direct chemical reduction (e.g. Na, Mg, Al) are often produced by electrolysis of their molten salts.
• The process requires high temperatures to melt the salt.
• No side reactions from solvents occur, since there is no water.

Electrolysis of Aqueous Solutions

In aqueous solutions, in addition to solute ions, water can also be oxidized or reduced. Thus, several possible half-reactions may compete at each electrode. The actual products depend on:

• The nature and concentration of the ions present.
• The electrode material (inert vs reactive).
• The applied voltage (overpotential, kinetics).

Competing Reactions at the Cathode

At the cathode, reduction may involve:

1. Reduction of metal cations (e.g. Cu$^{2+}$, Ag$^+$).
2. Reduction of water to hydrogen:
   $\,\text{H}_2\text{O} + 2\,e^- \rightarrow \text{H}_2 + 2\,\text{OH}^-$$

Empirically, for many aqueous solutions:

• Cations of less active metals (e.g. Cu$^{2+}$, Ag$^+$, Ni$^{2+}$) are more easily reduced than water. Result: metal deposition at the cathode.
• Cations of very active metals (e.g. Na$^+$, K$^+$, Ca$^{2+}$, Mg$^{2+}$) are much harder to reduce than water. Result: hydrogen gas is evolved, and hydroxide ions are formed; the solution becomes basic near the cathode.

Example: Electrolysis of aqueous CuSO$_4$ with inert electrodes

• Main cation: Cu$^{2+}$
• At cathode:
  $$\text{Cu}^{2+} + 2\,e^- \rightarrow \text{Cu}(\text{s})$$
  Copper metal is plated onto the cathode.

Example: Electrolysis of aqueous NaCl with inert electrodes

Possible reductions:

• Na$^+$ + e$^-$ → Na (very unfavorable in water)
• Water reduction:
  $\,\text{H}_2\text{O} + 2\,e^- \rightarrow \text{H}_2 + 2\,\text{OH}^-$$

Cathodic product: H$_2$(g) and OH$^-$ in solution, not sodium metal.

Competing Reactions at the Anode

At the anode, oxidation may involve:

1. Oxidation of anions (e.g. Cl$^-$, Br$^-$, I$^-$).
2. Oxidation of water to oxygen:
   $\,\text{H}_2\text{O} \rightarrow \text{O}_2 + 4\,\text{H}^+ + 4\,e^-$$

General trends with inert anodes in aqueous solutions:

• Halide ions (Cl$^-$, Br$^-$, I$^-$) are often oxidized to halogens if present in sufficiently high concentration.
• Example:
  $\,\text{Cl}^- \rightarrow \text{Cl}_2 + 2\,e^-$$
• If no easily oxidizable anion is present (e.g. with NO$_3^-$, SO$_4^{2-}$), water is oxidized to O$_2$.

Example: Electrolysis of dilute aqueous NaCl (brine) with inert electrodes

Cathode:
$$2\,\text{H}_2\text{O} + 2\,e^- \rightarrow \text{H}_2 + 2\,\text{OH}^-$$

Anode (at sufficiently high [Cl$^-$]):
$$2\,\text{Cl}^- \rightarrow \text{Cl}_2 + 2\,e^-$$

Overall reaction (simplified):
$$2\,\text{NaCl} + 2\,\text{H}_2\text{O} \rightarrow \text{H}_2 + \text{Cl}_2 + 2\,\text{Na}^+ + 2\,\text{OH}^-$$

The resulting solution around the cathode contains Na$^+$ and OH$^-$: effectively sodium hydroxide is produced.

Effect of Electrode Material

Electrodes can be:

• Inert (e.g. Pt, graphite, some oxides): Do not significantly participate chemically; only conduct electrons. Products arise from electrolyte components (ions, water).
• Reactive (e.g. Cu, Ni, Fe, Ag under some conditions): Can dissolve (oxidize) at the anode or be plated at the cathode.

Example: Electrorefining of copper

• Anode: impure Cu (reactive), oxidized:
  $$\text{Cu}(\text{s}) \rightarrow \text{Cu}^{2+} + 2\,e^-$$
• Cathode: pure Cu sheet; Cu$^{2+}$ from solution is reduced and deposited.

Here, the electrolyte (CuSO$_4$ solution) mostly carries Cu$^{2+}$ back and forth; the anode material itself supplies cations.

Quantitative Aspects: Faraday’s Laws of Electrolysis

Electrolysis allows controlled conversion between electrical charge and chemical amount. The quantitative relationships are summarized by Faraday’s laws of electrolysis.

First Law of Electrolysis

The mass $m$ of a substance produced (or consumed) at an electrode is proportional to the total charge $Q$ that has passed through the electrolyte:
$$m \propto Q$$

The total charge is:
$$Q = I \cdot t$$
with

• $I$ = current (in A = C·s$^{-1}$)
• $t$ = time (in s)

To relate charge to amount of substance, the concept of the Faraday constant $F$ is used:
$$F \approx 96485\ \text{C mol}^{-1}$$

One mole of electrons carries a charge of $F$ coulombs.

Second Law of Electrolysis

For different substances, if the same charge $Q$ passes through the cell, the amount of substance deposited or dissolved is proportional to the ratio of moles of electrons involved per mole of substance.

For a general half-reaction:
$$\text{M}^{n+} + n\,e^- \rightarrow \text{M}$$

• $n$ = number of moles of electrons required per mole of M.
• If $z$ is the number of electrons transferred per ion (often $z = n$), then the amount of substance $n_\text{M}$ (in mol) is:
  $$n_\text{M} = \frac{Q}{z\,F} = \frac{I \cdot t}{z\,F}$$

The corresponding mass $m$ is:
$$m = n_\text{M} \cdot M = \frac{I \cdot t \cdot M}{z\,F}$$

where $M$ is the molar mass.

Thus:

• Doubling the current or time doubles the amount produced.
• A species that requires more electrons per ion ($z$ larger) yields less substance for the same charge.

Example Structure (No Full Numerical Calculation)

Electroplating of copper from CuSO$_4$:

Cathode reaction:
$$\text{Cu}^{2+} + 2\,e^- \rightarrow \text{Cu}(\text{s})$$

Here $z = 2$. For a given $I$ and $t$,
$$n_{\text{Cu}} = \frac{I \cdot t}{2F}$$
and
$$m_{\text{Cu}} = \frac{I \cdot t \cdot M_{\text{Cu}}}{2F}$$

This approach is typical for all quantitative electrolysis problems.

Overvoltage (Overpotential) and Practical Cell Voltage

In ideal thermodynamic terms, electrode potentials and equilibrium cell voltages can be calculated from standard data. In real electrolytic cells, however, the actual applied voltage must be higher than the theoretical minimum to overcome:

• Activation barriers for electron transfer.
• Gas bubble formation effects at gas-evolving electrodes (H$_2$, O$_2$, Cl$_2$).
• Ohmic losses (resistance of electrolyte, contacts).
• Mass transport limitations (diffusion of reactants to the electrode).

This additional required voltage beyond the equilibrium (or reversible) potential is called overvoltage or overpotential.

Consequences:

• Hydrogen and oxygen evolution often require considerably more voltage than predicted.
• Choice of electrode material affects overvoltage substantially (e.g. Pt vs Hg vs Ni for H$_2$ evolution).
• In practice, cell voltages in industrial electrolysis are optimized as a compromise between energy efficiency and production rate.

Typical Applications of Electrolytic Processes

Electrolysis is central to many industrial, technical, and everyday processes. Only the electrolysis-specific aspects are emphasized here.

Industrial Metal Production

Some metals cannot be obtained economically by chemical reduction and are instead produced by electrolysis.

• Aluminum:
• Produced by electrolyzing molten Al$_2$O$_3$ dissolved in molten cryolite (Na$_3$AlF$_6$).
• Cathode: Al$^{3+}$ + 3 e$^-$ → Al(l)
• Anode: Oxidation reactions involving O$^{2-}$; CO$_2$ formation from carbon anodes is common.
• Sodium, magnesium and other light metals:
• Electrolysis of molten chlorides (NaCl, MgCl$_2$, etc.). Molten salt electrolysis avoids the competing reduction of water.

Key features: high temperature, large currents, and significant energy demand.

Electrolytic Refining and Recovery of Metals

Electrolytic methods are used to:

• Purify metals (electrorefining), especially Cu, Ni, Pb, Ag.
• Recover valuable metals from solutions (hydrometallurgy, waste treatment).

In electrorefining:

• Impure metal anode dissolves; pure metal is deposited at the cathode.
• Less noble impurities dissolve and later precipitate, or remain in solution; noble impurities (e.g. Ag, Au, Pt group) fall off the anode and form “anode mud,” which can be further processed.

Electroplating and Surface Treatment

Electrolytic processes allow deposition of a metal coating onto a conductive object:

• Electroplating (e.g. Cu, Ni, Cr, Zn plating):
• The object to be coated is the cathode.
• Metal ions in solution are reduced and deposited.
• Used for:
• Corrosion protection (zinc on steel).
• Decorative purposes (gold or silver plating).
• Functional surfaces (hard chromium, conductive layers on electronic connectors).

Control of current density, bath composition, temperature and pH determines coating quality (thickness, adhesion, grain size).

Production of Chemicals

Electrolysis is important for some bulk chemicals (details of specific industrial processes are addressed elsewhere):

• Water electrolysis:
• Splits water into hydrogen and oxygen:
• Cathode: $2\text{H}_2\text{O} + 2e^- \rightarrow \text{H}_2 + 2\text{OH}^-$
• Anode: $2\text{H}_2\text{O} \rightarrow \text{O}_2 + 4\text{H}^+ + 4e^-$
• Net: $2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2$
• Provides hydrogen without fossil fuels when powered by renewable electricity.
• Chloro-alkali electrolysis (electrolysis of brine):
• Produces Cl$_2$, H$_2$ and aqueous NaOH, as outlined earlier.
• Various cell designs (diaphragm, membrane, mercury) manage separation of products and environmental aspects.

Electrolytic Cleaning and Polishing

Electrolytic techniques can:

• Remove oxide layers, rust, or scale from metal surfaces (electrolytic cleaning or pickling).
• Electropolish metals:
• The workpiece acts as an anode.
• Material is removed preferentially from microscopic peaks, resulting in a smoother, brighter surface.

Electrolysis in Everyday Devices

Elementary forms of electrolysis appear in:

• Electrolytic capacitors (formation of oxide layers).
• Electrosynthesis in small-scale organic and inorganic lab processes.
• Sensors and analytical devices that rely on controlled electrolysis of analytes.

Summary

• Electrolytic processes use an external power source to drive non‑spontaneous redox reactions.
• In electrolytic cells, reduction occurs at the cathode (usually negative) and oxidation at the anode (usually positive).
• Molten salt electrolysis generally produces simple, predictable products: cations → metal at cathode; anions → nonmetal (often gas) at anode.
• In aqueous electrolysis, water can be oxidized or reduced, leading to competing half-reactions; actual products depend on ion type, concentration, electrode material and overvoltage.
• Faraday’s laws relate passed charge ($Q = I t$) to the amount of substance produced or consumed: $n = Q/(zF)$.
• Overvoltage and other practical factors mean real cells require higher voltages than thermodynamic predictions.
• Electrolytic processes are essential for metal production and refining, electroplating, production of key chemicals (e.g. H$_2$, Cl$_2$, NaOH), surface treatment, and numerous technical applications.