University of Benin CHM212 intermediate

Oxidation and reduction reaction

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Course Syllabus

  1. Fundamentals of Oxidation and Reduction
  2. Oxidation Numbers and Redox Systems
  3. Electrochemistry: Galvanic Cells
  4. Electrochemistry: Nernst Equation and Applications
  5. Electrochemistry: Electrolytic Cells and Corrosion

Study Notes

Fundamentals of Oxidation and Reduction

Fundamentals of Oxidation and Reduction

TL;DR

Oxidation is losing electrons, and reduction is gaining electrons; together these are called redox reactions. We can remember which is which using simple mnemonics like OIL RIG. These reactions are fundamental in chemistry, powering everything from batteries to digestion.

1. The Mental Model

Think of chemical reactions as a game of electron 'hot potato'. Sometimes an atom passes an electron to another (oxidation), and sometimes it catches one (reduction). These electron transfers are always linked – one can't happen without the other.

2. The Core Material

Oxidation and reduction are everywhere in chemistry, from how your body processes food to how batteries produce electricity. You'll often hear them referred to jointly as redox reactions.

What are Oxidation and Reduction?

The most fundamental way to think about oxidation and reduction is in terms of electron transfer.

  • Oxidation: This is the loss of electrons. When an atom or ion loses electrons, its charge becomes more positive.
  • Reduction: This is the gain of electrons. When an atom or ion gains electrons, its charge becomes more negative.

A highly useful mnemonic to remember this is:
OIL RIG
Oxidation Is Losing (electrons)
Reduction Is Gaining (electrons)

Remember, for every electron lost by one species, another species must gain it. Oxidation and reduction always happen simultaneously. You can't have one without the other. The substance that gets oxidized is called the reducing agent because it causes the other substance to be reduced. The substance that gets reduced is called the oxidizing agent because it causes the other substance to be oxidized.

Oxidation Numbers (or Oxidation States)

Sometimes, electron transfer isn't completely obvious, especially when we're dealing with covalent compounds (where electrons are shared, not fully transferred). To help track electron shifts in these situations, we use oxidation numbers (or oxidation states). Think of these as a way to assign imaginary charges based on electronegativity, even if no full electron transfer occurs.

Here are the basic rules for assigning oxidation numbers:

  1. Elemets in their elemental form have an oxidation number of 0.
    • Examples: H₂ (0), O₂ (0), Na (0), Fe (0).
  2. The oxidation number of a monatomic ion equals its charge.
    • Examples: Cl⁻ (-1), Na⁺ (+1), Ca²⁺ (+2)
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