Electrolysis principles and products with inert electrodes

Two conducting electrodes dip into the electrolyte without touching. Each is connected to a terminal of the DC supply. Inert electrodes, such as graphite or platinum, conduct electricity without reacting with the electrolyte or its products under the conditions used.

The electrode connected to the negative terminal is the cathode. Positively charged ions, called cations, are attracted to it. The electrode connected to the positive terminal is the anode. Negatively charged ions, called anions, are attracted to it.

Charge travels in different ways in different parts of the apparatus: electrons carry charge through the wires and electrodes, while moving ions carry charge through the electrolyte. At the electrodes, ions gain or lose electrons and form products. This is called being discharged.

Molten lead bromide: following each ion

Molten lead bromide contains lead ions, $\mathrm{Pb^{2+}}$, and bromide ions, $\mathrm{Br^-}$. With inert electrodes, it produces lead at the cathode and bromine at the anode.

Lead ions move to the negative cathode, where they gain electrons and become lead atoms. Bromide ions move to the positive anode, where they lose electrons. The resulting bromine atoms pair up to form bromine molecules, $\mathrm{Br_2}$, released as a gas.

Lead bromide must be heated to keep it molten. This is a teacher demonstration: the high temperature, toxic lead compounds and harmful bromine require appropriate safety controls.

Predicting products from other molten compounds

A binary ionic compound contains just two elements. When it is molten, its ions come only from that compound, so each electrode product follows directly from the ions present:

• The positive metal ions form the metal at the cathode.
• The negative non-metal ions form the non-metal at the anode.

For example, molten sodium chloride contains $\mathrm{Na^+}$ and $\mathrm{Cl^-}$ ions. Sodium forms at the cathode and chlorine, $\mathrm{Cl_2}$, forms at the anode. There is no water present to provide an alternative product.

Why aqueous solutions can give different products

An aqueous electrolyte contains ions from the dissolved compound, but water also provides hydrogen ions, $\mathrm{H^+}$, and hydroxide ions, $\mathrm{OH^-}$. These provide competing possibilities at the electrodes.

At the cathode, compare the metal with hydrogen in the reactivity series. If the metal is less reactive than hydrogen, as copper is, its ions form the metal. If the metal is more reactive than hydrogen, as sodium is, hydrogen forms instead of the metal.

At the anode, chloride, bromide or iodide ions can form their halogen. In the chloride solutions considered here, chlorine is the expected product. Concentration can affect this choice: very dilute chloride solutions can produce oxygen instead. When no halide ions are present, oxygen normally forms from hydroxide ions or water. Sulfate ions are not discharged in these examples.

The named electrolytes and their products

The following products are formed using inert electrodes:

Copper chloride solution: copper ions move towards the cathode and form a copper deposit because copper is less reactive than hydrogen. Chloride ions move towards the anode and form chlorine gas.

Sodium chloride solution: sodium is more reactive than hydrogen, so hydrogen forms at the cathode rather than sodium. Chloride ions form chlorine at the anode. Sodium ions remain in solution, along with hydroxide ions, leaving sodium hydroxide solution.

Sodium sulfate solution: sodium ions are not discharged at the cathode; hydrogen forms instead. Sulfate ions are not discharged at the anode; oxygen forms instead. The sodium and sulfate ions remain in solution, while water is broken down.

Water acidified with sulfuric acid: pure water conducts poorly because it contains very few ions. Adding dilute sulfuric acid provides more ions so current can flow readily. Hydrogen forms at the cathode and oxygen at the anode; sulfate ions are not discharged. The overall change is:

$$\mathrm{2H_2O(l) \rightarrow 2H_2(g) + O_2(g)}$$

Bubbles form at both electrodes. At the same temperature and pressure, the expected hydrogen volume is twice the oxygen volume. Hydrogen gives a squeaky pop with a lit splint; oxygen relights a glowing splint.

Higher tier: oxidation, reduction and half equations

In electron-transfer terms, reduction is gain of electrons and oxidation is loss of electrons. During electrolysis, reduction occurs at the cathode and oxidation occurs at the anode. The ions do not simply arrive and stick to an electrode: electron transfer changes them into products.

A half equation shows the reaction at one electrode. Start by writing the reacting ion and the product, balance the atoms, then add electrons, $\mathrm{e^-}$, so the total charge also balances.

For lead at the cathode, one $\mathrm{Pb^{2+}}$ ion needs two electrons to become a neutral atom:

$$\mathrm{Pb^{2+} + 2e^- \rightarrow Pb}$$

The total charge on the left is $+2-2=0$, matching the neutral lead on the right. This is reduction.

For bromine at the anode, two bromide ions are needed to make one diatomic bromine molecule:

$$\mathrm{2Br^- \rightarrow Br_2 + 2e^-}$$

Both sides have a total charge of $-2$. Electrons are released, so this is oxidation.

The other cathode reactions needed for the named solutions are:

$$\mathrm{Cu^{2+} + 2e^- \rightarrow Cu}$$ $$\mathrm{2H^+ + 2e^- \rightarrow H_2}$$

Chlorine formation at the anode is:

$$\mathrm{2Cl^- \rightarrow Cl_2 + 2e^-}$$

Oxygen formation can be represented using hydroxide ions:

$$\mathrm{4OH^- \rightarrow O_2 + 2H_2O + 4e^-}$$

Here, four oxygen atoms and four hydrogen atoms appear on each side. The charge is $-4$ on each side too. This illustrates why balancing a half equation requires checking both atoms and charge.

• Electrolysis uses electrical energy from a DC supply to decompose an electrolyte.
• Electrolytes conduct when molten or aqueous: their ions can move. Solid ionic compounds have fixed ions.
• Cations → negative cathode. Anions → positive anode.
• Inert electrodes conduct without reacting. Ions carry charge through the electrolyte; electrons carry charge through the external circuit.

Products with inert electrodes

Molten binary compound: metal at the cathode, non-metal at the anode.

Aqueous solution: water provides competing products. Metals less reactive than hydrogen form at the cathode; otherwise hydrogen forms. Halides can form halogens at the anode; without halides, oxygen forms. Sulfate ions are not discharged. Very dilute chloride solutions can give oxygen rather than chlorine.

Acidified water: $\mathrm{2H_2O \rightarrow 2H_2 + O_2}$; hydrogen : oxygen volume ratio 2 : 1.

Higher tier: electron transfer

Cathode: reduction, electron gain.

$\mathrm{Pb^{2+} + 2e^- \rightarrow Pb}$

$\mathrm{Cu^{2+} + 2e^- \rightarrow Cu}$

$\mathrm{2H^+ + 2e^- \rightarrow H_2}$

Anode: oxidation, electron loss.

$\mathrm{2Br^- \rightarrow Br_2 + 2e^-}$

$\mathrm{2Cl^- \rightarrow Cl_2 + 2e^-}$

$\mathrm{4OH^- \rightarrow O_2 + 2H_2O + 4e^-}$

Every half equation must balance atoms and total charge.