Lesson Notes By Weeks and Term v4 - SHS 3
EQUILIBRIA
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EQUILIBRIA
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Subject: Chemistry
Class: SHS 3
Term: 2nd Term
Week: 18
Grade code: 3.1.2.LI.4
Strand code: 1
Sub-strand code: 2
Content standard code: 3.1.2.CS.4
Indicator code: 3.1.2.LI.4
Theme: PHYSICAL CHEMISTRY
Subtheme: EQUILIBRIA
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This lesson introduces the concept of electrolysis, a process that uses electrical energy to drive non-spontaneous chemical reactions. While it might sound abstract, electrolysis is fundamental to many industrial processes that shape our daily lives in Ghana. From the aluminium roofing sheets on our houses (produced via electrolysis) and the shiny chrome parts on tro-tros and cars, to the purification of metals, this topic connects directly to our economy and environment. We will also discuss how electrolysis offers a cleaner alternative for metal refining compared to harmful practices like 'galamsey'.
2.1 What is Electrolysis? Electrolysis is the process of using a direct electric current (DC) to cause a chemical decomposition of a compound. The compound, called an electrolyte, must be in a molten state or dissolved in a suitable solvent (like water) so its ions are free to move. Electrolytic Cell vs. Voltaic (Galvanic) Cell: Voltaic Cell: A spontaneous chemical reaction *produces* electrical energy (e.g., a simple battery). Electrolytic Cell: Electrical energy from an external source is *used to force* a non-spontaneous chemical reaction to occur. 2.2 Components of an Electrolytic Cell An electrolytic cell has four main components: Power Supply: A source of direct current (DC), like a battery or a power pack. It acts as an "electron pump". Electrodes: These are electrical conductors (usually metals or graphite) immersed in the electrolyte. Anode: The positive electrode. It is connected to the positive terminal of the power supply. Oxidation occurs here (loss of electrons). Cathode: The negative electrode. It is connected to the negative terminal of the power supply. Reduction occurs here (gain of electrons). *Mnemonic:* PANIC -> Positive Anode, Negative Is Cathode. Electrolyte: A compound that conducts electricity when molten or in aqueous solution due to the presence of mobile ions. Examples: Molten NaCl, aqueous CuSO₄. Electrolytic Tank/Vessel: An inert container that holds the electrolyte and electrodes.
Diagram of a simple Electrolytic Cell: *(Teacher should draw this on the board and explain the flow)*
``` +-------[ DC Power Supply ]-------+ | | | | (+) Anode (-) Cathode | | |---------------------------------| | Electrolyte (e.g., | | molten NaCl with Na⁺ | | and Cl⁻ ions) | | | +---------------------------------+ ``` How it works (Flow of Charge): The power supply pushes electrons to the cathode, making it negative. Positive ions (cations) in the electrolyte are attracted to the negative cathode, where they gain electrons and are reduced. Negative ions (anions) in the electrolyte are attracted to the positive anode, where they lose electrons and are oxidized. The electrons flow from the anode, through the external wire, back to the positive terminal of the power supply, completing the circuit. 2.3 Predicting the Products of Electrolysis (Qualitative Analysis) What gets formed at the electrodes depends on several factors.
Factor 1: State of the Electrolyte (Molten vs. Aqueous) Molten Electrolytes: This is the simplest case. Only two types of ions are present: the cation and the anion from the salt. The cation moves to the cathode and is reduced. The anion moves to the anode and is oxidized. Example: Electrolysis of Molten Lead(II) Bromide (PbBr₂) Ions present: Pb²⁺ and Br⁻ At Cathode (-): Pb²⁺(l) + 2e⁻ → Pb(l) (Lead metal is formed) At Anode (+): 2Br⁻(l) → Br₂(g) + 2e⁻ (Bromine gas is formed) Aqueous Electrolytes: This is more complex because water can also be electrolyzed. Water can be reduced or oxidized. Reduction of water: 2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq) Oxidation of water: 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻ Therefore, at each electrode, there is a competition to see which species will be discharged.