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: 19
Grade code: 3.1.2.LI.5
Strand code: 1
Sub-strand code: 2
Content standard code: 3.1.2.CS.4
Indicator code: 3.1.2.LI.5
Theme: PHYSICAL CHEMISTRY
Subtheme: EQUILIBRIA
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This lesson explores the concept of corrosion, a natural but destructive process that affects metals all around us. We will focus specifically on the rusting of iron, which is a common sight on our roofing sheets, metal gates, bridges (like the Adomi or Adjin Kotoku bridges), and vehicles in Ghana. We will investigate this phenomenon not just as simple decay, but as a complex electrochemical process related to our study of redox reactions and equilibria. Understanding corrosion is crucial for developing ways to protect our valuable metal structures, saving money and ensuring safety in our communities.
This section breaks down the core scientific principles behind corrosion. 2.1 Defining Corrosion and Rusting Corrosion: This is the general term for the gradual destruction of a metal or alloy by chemical or electrochemical reaction with its environment. It is the process by which a refined metal is converted to a more chemically stable form, such as its oxide, hydroxide, or sulphide. Most metals corrode, for example, the green patina on old copper roofs or the tarnishing of silver. *Analogy:* Think of corrosion as the metal's way of returning to its natural state, similar to how it is found in the earth as an ore. Rusting: This is a specific type of corrosion that applies only to iron and its alloys, like steel. Rust is the common name for the reddish-brown substance that forms on iron when it is exposed to oxygen and moisture. The chemical name for rust is hydrated iron(III) oxide (Fe₂O₃·nH₂O). *Key Distinction:* All rusting is corrosion, but not all corrosion is rusting. A copper pipe turning green is corrosion, but not rusting. 2.2 The Electrochemical Mechanism of Rusting
Rusting is not a simple reaction; it is an electrochemical process that happens on the surface of the iron. The surface of an iron object acts like a collection of many tiny electrochemical cells (galvanic cells) in the presence of water, which acts as the electrolyte. Step 1: Anodic Region (Oxidation) At certain points on the iron surface, especially where there are impurities or stress, the iron acts as the anode (the site of oxidation). Here, iron atoms lose electrons and become iron(II) ions. Anode Half-Equation: `Fe(s) → Fe²⁺(aq) + 2e⁻` Step 2: Electron Flow The electrons released from the iron travel through the metal to another region on the surface. Step 3: Cathodic Region (Reduction) This other region, which has a higher electrode potential, acts as the cathode (the site of reduction). Here, oxygen from the air dissolves in the water and gains the electrons, forming hydroxide ions. Cathode Half-Equation: `O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)` Step 4: Formation of Rust The Fe²⁺ ions from the anode and the OH⁻ ions from the cathode diffuse towards each other and combine to form a precipitate of iron(II) hydroxide. `Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s)` (Greenish precipitate)
This iron(II) hydroxide is unstable and is quickly oxidized further by dissolved oxygen to form iron(III) hydroxide. `4Fe(OH)₂(s) + O₂(g) + 2H₂O(l) → 4Fe(OH)₃(s)` (Reddish-brown precipitate)
Finally, iron(III) hydroxide loses water to form the familiar reddish-brown, flaky solid we call rust. `2Fe(OH)₃(s) → Fe₂O₃·nH₂O(s)` (Hydrated Iron(III) Oxide - RUST)