1.4 - Electrolytic Cell

Electrolysis, Electrolytes and Electrodes

• Electrolysis is the decomposition of an electrolyte by electric current.

• An electrolyte conducts electricity in the molten or aqueous state and undergoes chemical change because it contains freely moving ions. A non-electrolyte contains no freely moving ions and does not conduct.

• In an electrolytic cell, electrical energy is converted to chemical energy:

  • anode: positive electrode; anions move here and are oxidised;
  • cathode: negative electrode; cations move here and are reduced.

• Carbon and platinum are inert electrodes unless otherwise stated.

• Ionic compounds such as lead(II) bromide and sodium chloride conduct when molten or dissolved; their solid ions are fixed in a lattice. Acids and alkalis conduct in water because ions are present; without water, there are no freely moving ions.

• Acetamide and glucose are non-electrolytes: they consist of molecules and have no freely moving ions.

Conductors Compared with Electrolytes

Conductor	Electrolyte
Conducts in the solid or molten state without chemical decomposition	Conducts in molten or aqueous form and undergoes chemical decomposition
Charge carried by free electrons	Charge carried by freely moving ions
Conductivity decreases as temperature rises	Conductivity increases as temperature rises
Metals and graphite	Ionic compounds, acids and alkalis

Electrolysis of a Molten Compound

Investigation of Molten Lead(II) Bromide

• Solid $\mathrm{PbB}{\mathrm{r}}_2$ does not conduct because its ions are fixed in a lattice. When molten, $\mathrm{P}{\mathrm{b}}^{2+}$ and $\mathrm{B}{\mathrm{r}}^{-}$ move freely and the bulb lights.
• At the cathode, grey lead forms: $\mathrm{P}{\mathrm{b}}^{2+}(\mathrm{l})+2{\mathrm{e}}^{-}\to \mathrm{Pb}(\mathrm{l})$.
• At the anode, brown bromine gas forms: $2B{\mathrm{r}}^{-}(\mathrm{l})\to \mathrm{B}{\mathrm{r}}_2(\mathrm{g})+2{\mathrm{e}}^{-}$.
• Overall: $\mathrm{P}{\mathrm{b}}^{2+}(\mathrm{l})+2B{\mathrm{r}}^{-}(\mathrm{l})\to \mathrm{Pb}(\mathrm{l})+\mathrm{B}{\mathrm{r}}_2(\mathrm{g})$.
• Aim: Study electrolysis of molten lead(II) bromide.
• Material: Lead(II) bromide powder.
• Apparatus: Beaker, crucible, bulb, battery, carbon electrodes, connecting wires with crocodile clips, switch, tripod, pipe-clay triangle and Bunsen burner.
• Procedure:
  1. Half-fill a crucible with lead(II) bromide and support it on a pipe-clay triangle over a tripod. Connect the electrodes, battery, bulb and switch.
  2. Heat until molten, immerse the electrodes and switch on the circuit.
  3. After five minutes, switch off and carefully pour the melt into a beaker to inspect the products.
• Precaution: The brown bromine gas is poisonous.

Electrolysis of Aqueous Solutions

• An aqueous electrolyte contains ions from the solute plus ${\mathrm{H}}^{+}$ and $\mathrm{O}{\mathrm{H}}^{-}$ from partial dissociation of water.
• Three factors determine which ions are discharged:

Factor	Effect
$E^0$ value	Cathode: more positive/less negative $E^0$ favours cation reduction. Inert anode: more negative/less positive $E^0$ favours anion oxidation.
ion concentration	Applies at the anode when halide ions are present: high halide concentration can favour their discharge despite a more positive $E^0$.
electrode type	an active metal anode can dissolve and form its ions instead of an anion being discharged

Aqueous Copper(II) Sulphate and Sulphuric Acid

• Aim: Study electrolysis of $0.1\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ copper(II) sulphate and $0.1\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ sulphuric acid with carbon electrodes.
• Apparatus: Electrolytic cell, battery, carbon electrodes, connecting wires with crocodile clips, switch, ammeter, test tubes and wooden splinter.
• Procedure:
  1. Half-fill the cell with copper(II) sulphate. Fill and invert two test tubes over the electrodes in the solution.
  2. Connect the electrodes, switch, ammeter and battery; switch on for a few minutes.
  3. Examine both electrodes and test collected gases with a glowing splinter at the anode and a burning splinter at the cathode.
  4. Repeat with sulphuric acid.
• Copper(II) sulphate: $\mathrm{C}{\mathrm{u}}^{2+}$, ${\mathrm{H}}^{+}$, $\mathrm{S}{\mathrm{O}}_4^{2-}$ and $\mathrm{O}{\mathrm{H}}^{-}$ are present. Cations move to the cathode; anions move to the anode.
  • Cathode: brown copper deposits, $\mathrm{C}{\mathrm{u}}^{2+}+2{\mathrm{e}}^{-}\to \mathrm{Cu}$, because $E^0=+0.34\mathrm{V}$ is more positive than $E^0({\mathrm{H}}^{+}/{\mathrm{H}}_2)=0.00\mathrm{V}$.
  • Anode: colourless oxygen bubbles form, $4O{\mathrm{H}}^{-}\to {\mathrm{O}}_2+2{\mathrm{H}}_2\mathrm{O}+4{\mathrm{e}}^{-}$; $E^0=+0.40\mathrm{V}$ is less positive than the sulphate-system value $+2.01\mathrm{V}$.
• Dilute sulphuric acid: Hydrogen forms at the cathode; oxygen forms at the anode. Cathode half-equation: $2{\mathrm{H}}^{+}+2{\mathrm{e}}^{-}\to {\mathrm{H}}_2$.

Experiment 1B — Concentration of Hydrochloric Acid

• Aim: Study the effect of ion concentration on selective discharge.
• Problem statement: Does HCl concentration affect the ions discharged at the anode?
• Hypothesis: $1.0\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ HCl discharges chloride ions; $0.0001\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ HCl discharges hydroxide ions.
• Variables: Manipulated = HCl concentration; responding = ion discharged at the anode; fixed = hydrochloric acid and carbon electrodes.
• Materials: The two HCl solutions above.
• Apparatus: Electrolytic cell, battery, carbon electrodes, connecting wires with crocodile clips, switch, ammeter, test tubes, wooden splinter and blue litmus paper.
• Procedure:
  1. Half-fill the cell with $1.0\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ HCl. Invert two tubes filled with the same acid over the electrodes and complete the circuit.
  2. Switch on for a few minutes; collect the anode gas and test with damp blue litmus paper.
  3. Repeat with $0.0001\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ HCl and test the anode gas with a glowing splinter.
• Products: Concentrated chloride gives chlorine, $2C{\mathrm{l}}^{-}\to \mathrm{C}{\mathrm{l}}_2+2{\mathrm{e}}^{-}$; very dilute acid gives oxygen, $4O{\mathrm{H}}^{-}\to {\mathrm{O}}_2+2{\mathrm{H}}_2\mathrm{O}+4{\mathrm{e}}^{-}$.
• Conclusion: High halide concentration can override the $E^0$ preference at the anode; the cathode selection still follows $E^0$.

Experiment 1C — Type of Electrode

• Aim: Study how electrode type affects selective discharge.
• Problem statement: Does electrode type affect the ions discharged?
• Material: $0.5\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ copper(II) sulphate.
• Apparatus: Beaker, battery, carbon and copper electrodes, connecting wires with crocodile clips, switch, ammeter, test tubes, wooden splinter, electronic balance and sandpaper.
• With carbon electrodes, copper deposits at the cathode and oxygen forms at the anode; the blue solution becomes paler as copper(II) ions are removed.
• With copper electrodes, the anode dissolves, $\mathrm{Cu}\to \mathrm{C}{\mathrm{u}}^{2+}+2{\mathrm{e}}^{-}$, while copper deposits at the cathode, $\mathrm{C}{\mathrm{u}}^{2+}+2{\mathrm{e}}^{-}\to \mathrm{Cu}$. The blue colour remains unchanged because copper(II) ions are replenished.

Voltaic Cell Compared with Electrolytic Cell

Voltaic cell	Electrolytic cell
chemical $\to$ electrical energy	electrical $\to$ chemical energy
spontaneous redox reaction	non-spontaneous reaction driven by a power supply
anode negative; cathode positive	anode positive; cathode negative
commonly uses different electrodes	may use electrodes of the same material

• In both cells, oxidation occurs at the anode, reduction at the cathode and electrons move through the external circuit from anode to cathode.

Electroplating and Purification of Metals

• Electroplating: object to be plated = cathode; plating metal = anode; electrolyte contains ions of the plating metal.
  • At the anode, plating metal atoms form ions.
  • At the cathode, those ions gain electrons and coat the object.
  • The object is cleaned with sandpaper; a rheostat controls the current during electroplating.
• Purification of copper: impure copper = anode; pure copper sheet = cathode; copper(II) salt solution = electrolyte.
  • Anode: $\mathrm{Cu}\to \mathrm{C}{\mathrm{u}}^{2+}+2{\mathrm{e}}^{-}$.
  • Cathode: $\mathrm{C}{\mathrm{u}}^{2+}+2{\mathrm{e}}^{-}\to \mathrm{Cu}$.
  • Pure copper builds up at the cathode; insoluble impurities collect as anode sludge.

Electroplating Investigation

• Aim: Study coating an iron spoon with copper.
• Materials: $1.0\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ copper(II) sulphate, iron spoon, copper plate and sandpaper.
• Apparatus: Beaker, battery, connecting wires with crocodile clips, ammeter, switch and rheostat.
• Procedure:
  1. Clean the spoon and copper plate. Half-fill a beaker with copper(II) sulphate.
  2. Connect the spoon to the negative terminal and copper to the positive terminal; immerse both in the solution.
  3. Adjust the current to $0.2\mathrm{A}$ with the rheostat; switch off after 30 minutes, remove and dry the spoon.
• Results: Copper deposits on the spoon while the anode becomes thinner. The blue solution remains unchanged because the anode replaces the copper(II) ions discharged at the cathode.
• Purpose: Electroplating improves appearance and resistance to corrosion.

Copper-purification Investigation

• Aim: Study purification of copper by electrolysis.
• Materials: $1.0\mathrm{mol}\mathrm{d}{\mathrm{m}}^{-3}$ copper(II) nitrate, impure copper plate and pure copper plate.
• Apparatus: Beaker, battery, connecting wires with crocodile clips, ammeter and switch.
• Procedure:
  1. Half-fill the beaker with copper(II) nitrate.
  2. Connect pure copper to the negative terminal and impure copper to the positive terminal; immerse both in the electrolyte.
  3. Pass current for 30 minutes and examine both electrodes.
• Results: The impure anode becomes thinner, the pure cathode becomes thicker and impurities collect below the anode.
• Copper from ore smelting is about $99.5\%$ pure; electrical wiring requires $99.99\%$ purity because impurities reduce conductivity.