Topic 4.9 Notes – Oxidation-Reduction (Redox) Reactions

1. What Redox Reactions Are

A redox reaction involves the transfer of electrons.

• Oxidation = loss of electrons → oxidation number increases
• Reduction = gain of electrons → oxidation number decreases
• Electrons flow from the species oxidized → to the species reduced
• Both always happen together

Two quick memory hooks:

• OIL RIG (Oxidation Is Loss, Reduction Is Gain)
• LEO says GER (Loss of Electrons = Oxidation; Gain of Electrons = Reduction)

The key tool here is the oxidation number. Even in covalent compounds where electrons are shared, we assign oxidation numbers as if electrons were completely transferred. That lets us see where electrons are “moving” on paper.

If no oxidation numbers change, it is not a redox reaction.

2. Assigning Oxidation Numbers

Before you can balance anything, you must confidently assign oxidation numbers.

Core Rules You Need

• Free elements (like $\ce{O2}$, $\ce{Na}$, $\ce{Cl2}$) = 0
• Monatomic ions = their charge
  • $\ce{Fe^{3+}}$ = +3
  • $\ce{Br^{-}}$ = -1
• Sum rule
  • Neutral compound → total = 0
  • Polyatomic ion → total = overall charge
• Oxygen
  • Usually -2
  • Peroxides ($\ce{H2O2}$, $\ce{O2^{2-}}$) = -1
  • Superoxide ($\ce{O2^{-}}$) = -½
• Hydrogen
  • Usually +1
  • Metal hydrides ($\ce{NaH}$) = -1
• Fluorine
  • Always -1
  • Other halogens usually -1 unless bonded to O or F

Quick Example

In $\ce{Cr2O7^{2-}}$:

• O is -2 (7 × -2 = -14)
• Total charge is -2
• Let Cr = x

$$2x + (-14) = -2$$

$$2x = 12 \Rightarrow x = +6$$

Chromium is +6.

When comparing reactants and products:

• If oxidation number increases → oxidized
• If oxidation number decreases → reduced

This comparison is how you decide what your half-reactions are.

3. The Half-Reaction Method

This is the skill the AP exam expects. You must balance mass and charge.

Let’s outline the structure clearly.

Balancing in Acidic Solution

1. Assign oxidation numbers.
2. Separate into oxidation and reduction half-reactions.
3. Balance all elements except O and H.
4. Balance O using $\ce{H2O}$.
5. Balance H using $\ce{H^{+}}$.
6. Balance charge using electrons $\ce{e^{-}}$.
7. Multiply half-reactions so electrons cancel.
8. Add them together.
9. Cancel anything that appears on both sides.

Two things students forget:

• You are balancing charge with electrons, not guessing coefficients.
• When you multiply a half-reaction, multiply every species in it.

When finished:

• Atoms balanced ✔️
• Total charge balanced ✔️
• No electrons left ✔️

If electrons remain, you missed cancellation.

4. Balancing in Basic Solution

Basic solution uses the same method, with one extra adjustment.

First, balance the reaction as if it were acidic.

After combining:

• If $\ce{H^{+}}$ appears, add the same number of $\ce{OH^{-}}$ to both sides.
• $\ce{H^{+} + OH^{-} -> H2O}$
• Cancel excess water if possible.

Final equation in basic solution should:

• Contain no $\ce{H^{+}}$
• Usually contain $\ce{OH^{-}}$ and/or $\ce{H2O}$**

Students often forget to add $\ce{OH^{-}}$ to both sides. That breaks charge balance.

5. What AP Questions Expect You to Do

You may be asked to:

• Identify which species is oxidized/reduced.
• Determine oxidation numbers.
• Write properly balanced half-reactions.
• Balance a full redox reaction in acidic or basic solution.

Where students lose points:

• Balancing atoms but ignoring charge.
• Forgetting to multiply half-reactions before adding.
• Leaving electrons in the final answer.
• Not removing $\ce{H^{+}}$ in basic solution.

On FRQs, graders look for clean cancellation of electrons and explicit balancing of charge. If charge doesn’t match on both sides, it’s automatically wrong even if atoms are balanced.