Lesson Notes By Weeks and Term v5 - Grade 12

Lesson Video

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Performance objectives

• Define and identify oxidation, reduction, oxidizing and reducing agents.
• Describe the operation of galvanic (voltaic) cells.
• Describe the operation of electrolytic cells.
• Use the standard electrode potential table to predict spontaneity.
• Explain electroplating and electrorefining applications.

Lesson summary

Electrochemical reactions are fundamental to many aspects of modern life, from the batteries that power our cellphones to the electroplating used to protect metal surfaces from corrosion. In South Africa, understanding these reactions is particularly relevant in contexts such as mining (where electrorefining is used to purify metals like copper and gold), water purification (electrolytic methods), and the development of sustainable energy sources (fuel cells, solar cells). Moreover, the principles behind electrochemical reactions are crucial for industries involved in the production of fertilizers and other essential chemicals.

Lesson notes

2.1 Redox Reactions: The Foundation At the heart of electrochemistry lies the concept of redox reactions, short for reduction-oxidation reactions. These reactions involve the transfer of electrons from one chemical species to another.

Oxidation: The loss of electrons by a chemical species. The oxidation number of the species increases.

Reduction: The gain of electrons by a chemical species. The oxidation number of the species decreases.

Oxidizing Agent: The species that accepts electrons and causes the oxidation of another species. The oxidizing agent itself is reduced.

Reducing Agent: The species that donates electrons and causes the reduction of another species. The reducing agent itself is oxidized.

Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain)

Example: Consider the reaction between zinc metal (Zn) and copper(II) ions (Cu 2+ ): Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s) Zinc (Zn) loses two electrons to become Zn 2+ . Thus, zinc is oxidized. Zinc is the reducing agent. Copper(II) ions (Cu 2+ ) gain two electrons to become Cu. Thus, copper(II) ions are reduced. Copper(II) ions are the oxidizing agent. 2.2 Electrochemical Cells: Harnessing Redox Reactions Electrochemical cells are devices that convert chemical energy into electrical energy (galvanic cells) or vice versa (electrolytic cells). 2.2.1 Galvanic Cells (Voltaic Cells) Galvanic cells utilize spontaneous redox reactions to generate electricity.

Key components include: Electrodes: Conductors where oxidation and reduction occur.

Anode: The electrode where oxidation occurs. It is the negative electrode in a galvanic cell.

Cathode: The electrode where reduction occurs. It is the positive electrode in a galvanic cell.

Electrolyte: A solution containing ions that conduct electricity. Each electrode is immersed in an electrolyte solution containing its own ions.

Salt Bridge: A tube containing an electrolyte (e.g., KCl, NaNO 3 ) that connects the two half-cells. Its function is to maintain electrical neutrality by allowing ions to flow between the half-cells, preventing charge buildup. Operation of a Galvanic Cell (e.g., Daniell Cell: Zn/Cu cell): Zinc electrode is placed in a solution of zinc sulfate (ZnSO 4 ) and a copper electrode is placed in a solution of copper(II) sulfate (CuSO 4 ). The two electrodes are connected externally by a wire, allowing electrons to flow. At the anode (Zn electrode), zinc atoms are oxidized: Zn(s) → Zn 2+ (aq) + 2e - . Electrons flow through the wire to the cathode. At the cathode (Cu electrode), copper(II) ions are reduced: Cu 2+ (aq) + 2e - → Cu(s). Copper ions from the solution are deposited on the copper electrode. The salt bridge allows ions to flow to balance the charge. K + ions migrate towards the CuSO 4 solution, and Cl - ions migrate towards the ZnSO 4 solution. The electron flow from the anode (Zn) to the cathode (Cu) generates an electric current.

Cell Notation (Shorthand Representation): Anode | Anode Solution || Cathode Solution | Cathode For the Daniell cell: Zn(s) | Zn 2+ (aq) || Cu 2+ (aq) | Cu(s) 2.2.2 Electrolytic Cells Electrolytic cells use electrical energy to drive non-spontaneous redox reactions (electrolysis). Key components are similar to galvanic cells: anode, cathode, electrolyte, and an external power source (e.g., a battery).

Anode: The electrode where oxidation occurs. It is the positive electrode in an electrolytic cell.

Cathode: The electrode where reduction occurs. It is the negative electrode in an electrolytic cell.

Key Difference: In electrolytic cells, an external power source forces electrons to flow in the opposite direction compared to galvanic cells.

Example: Electrolysis of Molten Sodium Chloride (NaCl) Molten NaCl (liquid state) contains Na + and Cl - ions. An external power source is connected to two inert electrodes (e.g., platinum or graphite). At the anode (+ electrode), chloride ions are oxidized: 2Cl - (l) → Cl 2 (g) + 2e - (Chlorine gas is produced). At the cathode (- electrode), sodium ions are reduced: Na + (l) + e - → Na(l) (Liquid sodium is produced).

Example: Electrolysis of Water Electrolysis of water (H 2 O) requires the addition of an electrolyte (e.g., dilute sulfuric acid, H 2 SO 4 ) to increase conductivity. At the anode (+ electrode), water is oxidized: 2H 2 O(l) → O 2 (g) + 4H + (aq) + 4e - (Oxygen gas is produced). At the cathode (- electrode), hydrogen ions are reduced: 4H + (aq) + 4e - → 2H 2 (g) (Hydrogen gas is produced).

Overall: 2H 2 O(l) → 2H 2 (g) + O 2 (g) 2.3 Standard Electrode Potentials The standard electrode potential (E°) is the measure of the potential difference between a half-cell and the standard hydrogen electrode (SHE) under standard conditions (298 K, 1 atm pressure, 1 M concentration). The SHE is assigned a potential of 0.00

V. Table of Standard Reduction Potentials: This table lists the reduction half-reactions and their corresponding standard reduction potentials. It's crucial for predicting the spontaneity of redox reactions.