Electrolytic Cells & Electrolysis

What is an Electrolytic Cell

An electrolytic cell is a device in which electrical energy from an external direct-current source is used to drive a chemical reaction that would not occur on its own. The process itself is called electrolysis: the decomposition of an electrolyte, in the molten state or in solution, by the passage of electric current. Because the reaction being driven is non-spontaneous, its Gibbs energy change is positive, and the external source must continuously supply the energy that the reaction cannot release by itself.

The simplest illustration from the NCERT text uses two copper strips dipping in aqueous copper sulphate. When a DC voltage is applied, $\ce{Cu^2+}$ ions are reduced and deposited at the cathode while copper from the anode is oxidised back into solution:

Cathode (reduction): $\ce{Cu^2+ (aq) + 2e^- -> Cu(s)}$
Anode (oxidation): $\ce{Cu(s) -> Cu^2+ (aq) + 2e^-}$

Copper is dissolved at the anode and deposited at the cathode. This single observation already contains the seed of an entire industry: it is the basis of the process by which impure copper is refined into copper of high purity. Metals such as sodium, magnesium and aluminium, for which no convenient chemical reducing agent exists, are likewise produced on a large scale by electrolytic reduction of their cations.

Galvanic vs Electrolytic Cells

The relationship between the two cell types is best seen as one continuous experiment. A Daniell cell drives the spontaneous reaction $\ce{Zn(s) + Cu^2+ (aq) -> Zn^2+ (aq) + Cu(s)}$ and develops about 1.1 V. If an opposing external potential is applied and slowly raised, the reaction continues until the opposing voltage reaches 1.1 V, at which point no current flows. Increase the external potential beyond 1.1 V and the same cell now runs in reverse: it has become an electrolytic cell. The chemistry has not changed, only the direction in which energy flows.

The constant that survives the switch is the chemistry of the electrodes: in both cells oxidation occurs at the anode and reduction at the cathode. Only the sign of each electrode is reversed, because the electron flow is now imposed from outside rather than generated within. For a full treatment of how the spontaneous version operates, study the sibling note on electrochemical cells.

Construction of an Electrolytic Cell

A laboratory electrolytic cell needs only three essential parts: an electrolyte that supplies mobile ions, two electrodes dipping into it, and an external DC source connected across them. The electrode wired to the negative terminal becomes the cathode, where cations migrate and are reduced; the electrode wired to the positive terminal becomes the anode, where anions migrate and are oxidised. Within the electrolyte, current is carried by the movement of ions; in the external wires it is carried by electrons.

For electrolysis to begin, the applied voltage must exceed a minimum value called the decomposition (discharge) potential of the electrolyte. Below this voltage essentially no sustained current flows; above it, products begin to appear at the electrodes. The decomposition potential is governed by the electrode potentials of the species present, modified, as we shall see, by overpotential effects for gaseous products.

Molten vs Aqueous Electrolytes

The single most important question in this subtopic is: what species are actually available to be discharged? The answer differs sharply between a molten salt and its aqueous solution, and that difference dictates the products.

In a molten electrolyte there is no solvent. Molten sodium chloride furnishes only its own ions, $\ce{Na+}$ and $\ce{Cl-}$, so each electrode has exactly one candidate reaction:

Cathode: $\ce{Na+ (l) + e^- -> Na(l)}$
Anode: $\ce{2Cl- (l) -> Cl2(g) + 2e^-}$

Sodium metal collects at the cathode and chlorine gas is evolved at the anode. There is no competition, which is precisely why molten salts are the route to reactive metals — sodium and magnesium are won from their fused chlorides, and aluminium from molten alumina dissolved in cryolite.

In an aqueous solution, water itself becomes a competitor at both electrodes. Besides $\ce{Na+}$ and $\ce{Cl-}$, the solution contains $\ce{H+}$ and $\ce{OH-}$ from the self-ionisation of water, along with the $\ce{H2O}$ molecules. Now each electrode has two or more possible reactions, and which one wins is decided by electrode potentials and overpotential.

Preferential Discharge & Overpotential

When more than one ion can be discharged at an electrode, the products of electrolysis depend on the standard electrode potentials of the competing species. The governing principle, stated in both the NCERT and NIOS texts, is that the easiest process occurs:

• At the cathode, the species with the higher (more positive) reduction potential is reduced preferentially.
• At the anode, the species with the lower reduction potential (i.e. greater tendency to be oxidised) is oxidised preferentially.

This electrode-potential rule is the first filter. The second filter is overpotential: in many cases the voltage actually required to discharge a species, especially a gas, exceeds the value predicted from its electrode potential. The excess is the overpotential, and it can be large enough to overturn the prediction of the simple potential comparison. The classic figure is that the overpotential for $\ce{H2}$ evolution is essentially zero on platinum but about 1.5 V on a mercury cathode — large enough that on mercury, sodium is discharged in preference to hydrogen.

Electrolysis Worked Out: NaCl, Water, H₂SO₄

Aqueous sodium chloride (brine)

At the cathode two reductions compete: $\ce{Na+ + e^- -> Na}$ with $E^\circ = -2.71\,\text{V}$ and $\ce{H+ + e^- -> 1/2 H2}$ with $E^\circ = 0.00\,\text{V}$. The higher potential is preferred, so hydrogen is liberated; since $\ce{H+}$ comes from water dissociation, the net cathode reaction is written:

Cathode: $\ce{H2O(l) + e^- -> 1/2 H2(g) + OH^-(aq)}$

At the anode two oxidations compete: $\ce{Cl- -> 1/2 Cl2 + e^-}$ with $E^\circ = 1.36\,\text{V}$ and $\ce{2H2O -> O2 + 4H+ + 4e^-}$ with $E^\circ = 1.23\,\text{V}$. On potential alone water should be oxidised first, but the high overpotential of oxygen makes chloride oxidation the easier route, so chlorine is evolved. The net cell reaction is:

$\ce{NaCl(aq) + H2O(l) -> Na+(aq) + OH^-(aq) + 1/2 H2(g) + 1/2 Cl2(g)}$

The solution is left enriched in $\ce{Na+}$ and $\ce{OH-}$, i.e. sodium hydroxide — the industrial chlor-alkali process, which yields $\ce{NaOH}$, $\ce{Cl2}$ and $\ce{H2}$ simultaneously.

Water and dilute sulphuric acid

During electrolysis of sulphuric acid with inert platinum electrodes, the possible anode oxidations are:

$\ce{2H2O(l) -> O2(g) + 4H+(aq) + 4e^-}$ ($E^\circ = +1.23\,\text{V}$)
$\ce{2SO4^2-(aq) -> S2O8^2-(aq) + 2e^-}$ ($E^\circ = +1.96\,\text{V}$)

For dilute $\ce{H2SO4}$ the first reaction is preferred and oxygen is evolved at the anode, while hydrogen is liberated at the cathode. The net result is the electrolysis of water itself, producing $\ce{H2}$ and $\ce{O2}$ in a 2 : 1 volume ratio. Only at high concentrations of $\ce{H2SO4}$ does sulphate oxidation to peroxodisulphate ($\ce{S2O8^2-}$) become preferred.

Inert vs Reactive Electrodes

Products of electrolysis depend not only on the electrolyte but on the nature of the electrode. An inert electrode such as platinum or gold does not take part in the reaction; it serves only as a source or sink for electrons, so the products come entirely from the ions and water present. A reactive electrode participates directly, and this changes the outcome.

The contrast is sharpest for aqueous $\ce{CuSO4}$. With inert Pt electrodes, $\ce{Cu^2+}$ is deposited at the cathode and water is oxidised to $\ce{O2}$ at the anode. With copper electrodes, the anode metal itself is oxidised more easily than water or sulphate, so the copper anode dissolves:

Anode (Cu): $\ce{Cu(s) -> Cu^2+(aq) + 2e^-}$
Cathode (Cu): $\ce{Cu^2+(aq) + 2e^- -> Cu(s)}$

No oxygen is evolved; copper simply transfers from anode to cathode. This anode-dissolution behaviour is exactly what makes electrorefining and electroplating possible.

Applications: Electroplating & Electrorefining

The reactive-electrode behaviour of the copper–copper sulphate cell scales directly into two major industrial uses.

Electrorefining purifies a metal. The impure metal is made the anode and a thin sheet of the pure metal the cathode, both immersed in a solution of a salt of that metal. On passing current the impure anode dissolves, and pure metal deposits on the cathode. More electropositive (soluble) impurities remain in solution, while noble impurities such as silver and gold, which are not oxidised, fall to the bottom as anode mud — itself a valuable by-product. This is how the impure copper of the introductory example is converted into copper of high purity.

Electroplating deposits a thin protective or decorative layer of one metal onto another. The article to be coated is made the cathode and is immersed in a solution containing ions of the plating metal; current deposits the metal uniformly on its surface. Chromium plating, for instance, is both decorative and protective, guarding the underlying metal against corrosion. The same Faraday relationships that govern refining govern the thickness of a plate, linking this subtopic back to Faraday's laws of electrolysis.

Beyond these, electrolysis underpins the extraction of metals that resist chemical reduction — sodium and magnesium from molten chlorides, aluminium from molten alumina in cryolite — and the chlor-alkali industry built on aqueous brine. In each case the choice between molten and aqueous feed, and between inert and reactive electrodes, is dictated by the discharge principles developed above and by the underlying ordering of electrode potentials in the standard hydrogen electrode and electrochemical series.