Electrolysis Concepts and Applications

Electrolysis

• Electrolysis is the process of using electricity to break down or decompose a compound (usually an ionic compound in the molten state or in aqueous solution).
• Electrolysis takes place in an electrolytic cell.

Learning Objectives

• Define electrolysis
• Describe electrolysis as evidence of the structure of ionic compounds in the solid, molten, and aqueous states.
• Predict the products of the electrolysis of a molten ionic compound.
• Predict the products of the electrolysis of an aqueous electrolyte using the idea of selective discharge of ions.
• Construct ionic equations for the reactions occurring at the electrodes during the electrolysis, given relevant information.

Parts of an Electrolytic Cell

• Battery: Acts as an electron pump, drawing electrons away from the anode and supplying them to the cathode. Anode becomes positively charged and cathode becomes negatively charged.
• Electrodes:
  • Conduct electricity
  • Usually carbon rods or metal plates
  • Electrode connected to the positive terminal of the battery is the anode
  • Electrode connected to the negative terminal of the battery is the cathode
• Electrolyte:
  • Conducts electricity
  • Contains free-moving ions to allow electricity to flow
  • Is a molten ionic compound or an aqueous solution
  • Is decomposed to form positive ions (cations) and negative ions (anions)

How Does Electrolysis Work?

• At the anode:

  • Anions move to the anode, giving up electrons.
  • Oxidation occurs at the anode

• At the cathode:

  • Cations move to the cathode, receiving electrons.
  • Reduction occurs at the cathode

• Within the electrolyte:

  • The flow of ions towards the electrodes is the flow of electric current through the electrolyte

• In the external circuit:

  • During electrolysis, electrons flow from the negative terminal to the positive terminal of the battery.

• Examples of anion: Cl⁻, OH⁻

• Examples of cation: Na⁺, H⁺

What happens to ions at the electrodes?

• Cations receive electrons at the negatively-charged cathode
• Anions give up electrons at the positively-charged anode.
• When cations/anions gain/lose electrons at the electrodes, they form atoms or molecules. This is called discharging.

Electrolysis of molten Ionic Compounds

• When a molten binary compound (containing only 2 elements) undergoes electrolysis, a metal and a non-metal are formed as products.
• The molten compound does not contain water molecules

Electrolysis of molten sodium chloride

• At the anode:
  • Negatively-charged Cl⁻ ions are attracted to the anode.
  • Cl⁻ ions lose electrons to form chlorine gas.
  • Oxidation occurs
• 2Cl⁻(l) → Cl₂(g) + 2e⁻
• The Cl ions are discharged.
• At the cathode:
  • Positively-charged Na⁺ ions are attracted to the cathode.
  • Each Na⁺ ion gains one electron to form a sodium atom, which is reduced.
  • Na⁺(l) + e⁻ → Na(l)
  • The Na⁺ ions are discharged.
• The overall reaction: 2NaCl(l) → 2Na(l) + Cl₂(g)

Inert Electrodes

• Inert carbon electrodes are used to prevent chlorine from reacting with the electrode.
• Inert electrodes do not react with the products of electrolysis or with the electrolyte.
• Carbon (graphite) and platinum electrodes are considered inert because they are rarely involved in electrolytic reactions.

Electrolysis of dilute NaCl Solution

• At cathode:
  • H⁺ and Na⁺ ions are attracted to cathode.
  • Sodium is higher in the reactivity series. Therefore, H⁺ ions are preferentially discharged as hydrogen gas.
• 2H⁺ + 2e⁻ → H₂(g)
• Na+ ions remain in solution.
• At anode:
  • OH⁻ and Cl⁻ ions are attracted to the anode.
  • OH ions give up electrons more readily than Cl⁻ions.
  • 4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻
  • Cl⁻ ions remain in solution
• Overall reaction: 2H₂O(l) → 2H₂(g) + O₂(g)

Electrolysis of concentrated NaCl Solution

• At cathode:
  • H+ and Na+ ions are attracted to the cathode.
  • Sodium is higher in the reactivity series. Therefore, H+ ions are preferentially discharged as hydrogen gas.
  • 2H⁺ + 2e⁻ → H₂(g)
  • Na⁺ ions remain in solution.
• At anode:
  • OH⁻ and Cl⁻ ions are attracted to the anode.
  • Since Cl⁻ ions are more numerous than OH⁻ ions, Cl⁻ ions are discharged to form chlorine gas.
• 2Cl⁻(aq) → Cl₂(g) + 2e⁻
• OH⁻ ions remain in solution

Electrolysis of Aqueous Solutions

• An aqueous solution of a compound is a mixture of the compound with water.
• In an aqueous solution, more than one type of cation or anion is present in the electrolyte.
• E.g., aqueous sodium chloride contains Na⁺, H⁺, Cl⁻, and OH⁻ ions.
• Only one cation and one anion are preferentially discharged during electrolysis. This is known as selective discharge of ions.

Selective Discharge of Ions

• The ions discharged during electrolysis depend on three factors:
  • Selective discharge of cations
  • Selective discharge of anions
  • Effect of concentration on the selective discharge of anions
• The more reactive the metal, the greater its tendency to form ions.
• Therefore, in electrolysis, the ease of discharge of cations depends on the position of the metal in the reactivity series.
  • H⁺ ions are discharged in preference to ions of metals above hydrogen in the reactivity series.
• The OH⁻ ions will give up electrons preferentially to form water and oxygen gas during electrolysis.
• Sulfate (SO₄²⁻) and nitrate (NO₃⁻) ions normally remain in solution during electrolysis.

Predicting the Products of Electrolysis

• Identify the cations and anions in the electrolyte

• Determine the anion that is discharged at the anode (OH⁻ is often easier if not a high concentration of Cl⁻, Br⁻ or I⁻ anions)

• Determine the cation that is discharged at the cathode (using activity series)

• The cations and anions remaining in the solution after electrolysis form the product remaining in the solution.

Electrolysis of CuSO₄ Solution (inert electrodes)

• At anode:
  • OH⁻ and SO₄²⁻ ions are attracted to the anode.
  • SO₄²⁻ ions are not discharged.
  • OH⁻ ions are preferentially discharged to form water (H₂O) and oxygen (O₂) gas.
  • 4OH⁻(aq) → 2H₂O(l) + O₂(g) + 4e⁻
  • SO₄²⁻ ions remain in solution
• At cathode:
  • H⁺ and Cu²⁺ ions are attracted to the cathode.
  • Copper is lower than hydrogen in the reactivity series. Therefore, Cu²⁺ ions are preferentially discharged as copper metal.
  • Cu²⁺(aq) + 2e⁻ → Cu(s)
  • H⁺ ions remain in solution.
• Summary:
  • Copper metal is deposited at the cathode, and oxygen gas is given off at the anode.

Purification of Impure Copper

• To refine impure copper, the impure copper is used as the anode, while a thin sheet of pure copper is used as the cathode.
• The electrolyte must be an aqueous solution of copper (II) sulfate (CuSO₄).

Electroplating

• The object to be electroplated is made the cathode, and the anode is the source of the plating metal.
• The electrolyte is an aqueous solution of a salt of the plating metal (e.g., CuSO₄ for copper plating.).

Industrial Applications of Electrolysis

• Purify metals (e.g. refining copper)
• Deposit a layer of metal on another substance (electroplating) for decorative purposes or preventing corrosion (e.g., silver or gold plating).

This quiz covers the fundamental concepts of electrolysis, including the definition, processes involved, and the structure of ionic compounds. It also addresses the predictions regarding the products of electrolysis in various states and the ionic equations relevant to these reactions. Test your understanding of electrolytic cells and their components.