Course Content
Module 1: Foundations of Redox Reactions
Oxidation and Reduction: Understanding redox in terms of oxygen transfer, hydrogen transfer, and electron transfer. Oxidation Numbers: Rules for assigning oxidation states and identifying oxidizing/reducing agents. Balancing Redox Equations: Half-reaction method and oxidation number method.
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Part 1: Defining Oxidation and Reduction
Part 2: Oxidation Numbers
Part 3: Balancing Redox Equations
Practice Questions on Redox Reactions
00:20:00
TEst For Oxidizing And Reducing Agents
Module 2: Electrolysis (The Electrolytic Cell)
Components of Electrolysis: Electrolytes (strong vs. weak), electrodes (inert vs. active), and the electrolytic tank. Preferential Discharge of Ions: Factors affecting the discharge of ions (position in electrochemical series, concentration, and nature of electrodes). Practical Examples: Electrolysis of acidified water, brine, and copper(II) sulfate solution.
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1. Components of an Electrolytic Cell
2. Preferential Discharge of Ions
3. Practical Examples of Electrolysis
Module 3: Quantitative Electrochemistry (Faraday’s Laws)
Module 3: Quantitative Electrochemistry (Faraday’s Laws)Faraday’s First Law: Relationship between mass ($m$), current ($I$), and time ($t$). Formula: $m = zIt$.Faraday’s Second Law: Chemical equivalents and the Faraday constant ($96,500 text{ C mol}^{-1}$).Calculations: Determining mass deposited, volume of gases at STP, and the charge on ions.
Module 4: Electrochemical Cells (Galvanic/Voltaic)
Cell Construction: Half-cells, salt bridges, and external circuits. Standard Electrode Potential (E ∘ ): The Standard Hydrogen Electrode (SHE) and the Electrochemical Series. Cell Notation: Writing and interpreting cell diagrams (e.g., Zn (s)∣Zn(aq)2+ ∣∣Cu (aq)2+ ∣Cu (s)). Electromotive Force (e.m.f.): Calculating E cell ∘ ​ =E cathode ∘ ​ −E anode
Module 5: Industrial Applications & Corrosion
Electroplating and Purification: How electrolysis is used to coat metals and refine copper. Primary and Secondary Cells: Lead-acid accumulators, dry cells (Leclanché), and fuel cells. Corrosion of Metals: The chemistry of rusting and methods of prevention (cathodic protection, galvanizing).
Module 6: Exam Prep (WAEC/NECO/UTME Focus)
Electrochemistry

Part 2: Oxidation Numbers

Oxidation numbers are a bookkeeping tool. They allow us to track electron movement in reactions where electrons aren’t explicitly shown. A change in oxidation number is the definitive proof that a redox reaction has occurred.

Rules for Assigning Oxidation Numbers (in order of priority)

  1. Free Elements: The oxidation number of an atom in its elemental form is 0.

    • Examples: ( $$ O_2, Zn, S_8, H_2 $$ ) → all atoms = 0.

  2. Monatomic Ions: The oxidation number equals the ion’s charge.

    • Examples: ( $$ Al^{3+} = +3 $$); ( $$ Cl^- = -1 $$ ).

  3. Fluorine: Fluorine (F) is always -1 in all its compounds (it’s the most electronegative element).

  4. Group 1 & 2 Metals: Group 1 metals are always +1. Group 2 metals are always +2.

  5. Hydrogen: Usually +1. Exception: In metal hydrides (e.g., NaH, CaH₂), it is -1.

  6. Oxygen: Usually -2. Exceptions: In peroxides (e.g., H₂O₂) it is -1; in compounds with fluorine (e.g., OF₂) it is +2.

  7. Sum of Oxidation Numbers:

    • For a neutral molecule, the sum of all oxidation numbers = 0.

    • For a polyatomic ion, the sum = the charge of the ion.

Identifying Oxidizing and Reducing Agents

Once you have assigned oxidation numbers, you can identify the key players:

  • Oxidizing Agent (Oxidant): The substance that is reduced (gains electrons). It causes another substance to be oxidized.

    • Its own oxidation number decreases.

  • Reducing Agent (Reductant): The substance that is oxidized (loses electrons). It causes another substance to be reduced.

    • Its own oxidation number increases.

Example:

$$ \overset{0}{Zn} + \overset{+1}{H_2}\overset{-2}{SO_4} \rightarrow \overset{+2}{ZnSO_4} + \overset{0}{H_2} $$

  • Zn: (0 $$ \rightarrow +2 $$) (oxidation, loses electrons). Zn is the reducing agent.

  • H: ( $$+1 \rightarrow 0 $$) (reduction, gains electrons). ( $$ H_2SO_4 $$ ) (the source of H⁺) is the oxidizing agent.

 

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