Lesson Notes By Weeks and Term v5 - Grade 12

Lesson Video

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Performance objectives

• Define oxidation and reduction (Redox).
• Explain how voltaic (galvanic) cells work.
• Calculate cell potential (E°cell) to predict spontaneity.
• Describe the process and applications of electrolysis.

Lesson summary

Electrochemical reactions are at the heart of many processes that shape our daily lives, from the batteries powering our cellphones to the electroplating used to protect metal structures from corrosion. In South Africa, understanding these reactions is crucial for industries like mining (extracting metals), manufacturing (creating corrosion-resistant products), and energy (developing efficient batteries and fuel cells). The principles of electrochemistry allow us to harness the energy released or consumed during chemical reactions in a controlled manner, leading to technological advancements and solutions for various challenges.

Lesson notes

2.1 Redox Reactions: Oxidation and Reduction Oxidation: Oxidation is the loss of electrons by a substance. The oxidation number of the substance increases.

Reduction: Reduction is the gain of electrons by a substance. The oxidation number of the substance decreases.

A helpful mnemonic is OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

Oxidizing Agent: The substance that causes oxidation by accepting electrons. The oxidizing agent itself is reduced.

Reducing Agent: The substance that causes reduction by donating electrons. The reducing agent itself is oxidized.

Oxidation Numbers: Oxidation numbers are assigned to atoms in a molecule or ion to keep track of electron distribution.

Here are some rules: The oxidation number of an element in its elemental form is 0 (e.g., Na(s), O₂(g), Cu(s)). The oxidation number of a monatomic ion is equal to its charge (e.g., Na⁺ is +1, Cl⁻ is -1). Oxygen usually has an oxidation number of -2, except in peroxides (like H₂O₂) where it is -1, and when bonded to fluorine, where it can be positive. Hydrogen usually has an oxidation number of +1, except when bonded to metals in binary compounds (metal hydrides) where it is -1 (e.g., NaH). The sum of the oxidation numbers in a neutral compound is zero. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.

Example 1: Identifying Oxidation and Reduction Consider the reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Zinc (Zn): Zn(s) goes from oxidation number 0 to +2 (loses 2 electrons). Zinc is oxidized. It's the reducing agent. Copper (Cu²⁺): Cu²⁺(aq) goes from oxidation number +2 to 0 (gains 2 electrons). Copper is reduced. It's the oxidizing agent. 2.2 Voltaic (Galvanic) Cells Voltaic cells are electrochemical cells that use spontaneous redox reactions to generate electrical energy. They consist of two half-cells connected by a salt bridge.

Half-Cell: Consists of an electrode (usually a metal) immersed in an electrolyte solution (containing ions of the same metal).

Electrodes: Anode: The electrode where oxidation occurs. Electrons are released at the anode. It's considered the negative terminal.

Cathode: The electrode where reduction occurs. Electrons are consumed at the cathode. It's considered the positive terminal.

Electrolyte: A solution containing ions that conduct electricity.

Salt Bridge: A tube containing an electrolyte solution (e.g., KNO₃ or KCl) that connects the two half-cells. The salt bridge allows ions to flow between the half-cells, maintaining electrical neutrality and completing the circuit. Without it, the reaction would quickly stop.

Example 2: The Zinc-Copper Voltaic Cell In the Zn-Cu cell: Anode (Oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻ Cathode (Reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)

Salt Bridge: Allows NO₃⁻ ions to flow into the Zn²⁺ solution to neutralize the build-up of positive charge and K⁺ ions to flow into the Cu²⁺ solution to neutralize the decrease in positive charge.

Electron Flow: Electrons flow from the zinc anode (where they are produced) to the copper cathode (where they are consumed) through an external circuit (e.g., a wire connected to a light bulb). 2.3 Standard Electrode Potentials (E°) The standard electrode potential (E°) is the potential of a half-cell under standard conditions (298 K, 1 atm pressure for gases, and 1 M concentration for solutions) relative to the Standard Hydrogen Electrode (SHE), which is assigned a potential of 0

V. E°cell Calculation: The standard cell potential (E°cell) for a voltaic cell is calculated as: E°cell = E°(cathode) - E°(anode)

Where: E°(cathode) is the standard reduction potential of the cathode. E°(anode) is the standard reduction potential of the anode.

Important: You use the reduction potential value, even for the anode.* Spontaneity: If E°cell > 0, the reaction is spontaneous (galvanic/voltaic cell). If E°cell 0, the reaction is spontaneous. This voltaic cell will work. 2.4 Electrolysis Electrolysis is the process of using an electrical current to drive a non-spontaneous redox reaction. This requires an external power source.

Electrolytic Cell: Consists of two electrodes immersed in an electrolyte solution and connected to an external power source (e.g., a battery).

Anode (Electrolytic): Oxidation occurs. It's the positive terminal (connected to the positive terminal of the power supply).

Cathode (Electrolytic): Reduction occurs. It's the negative terminal (connected to the negative terminal of the power supply). Key Difference between Voltaic and Electrolytic Cells: In voltaic cells, the reaction is spontaneous and produces electricity. In electrolytic cells, electricity is used to force a non-spontaneous reaction to occur.

Applications of Electrolysis: Electroplating: Coating a metal object with a thin layer of another metal (e.g., chrome plating car bumpers). The object to be plated is the cathode.