Welcome to the World of Redox!

Welcome to one of the most important chapters in your A Level Chemistry journey: Redox I. If you’ve ever wondered how the battery in your phone works, why a silver spoon tarnishes, or how your body gets energy from food, the answer lies in Redox reactions.

At its heart, Redox is simply about the "shipping and receiving" of electrons. Don't worry if it seems like a lot of rules at first—we will break it down step-by-step until you're a Redox pro!

1. What is an Oxidation Number?

Think of an oxidation number (or oxidation state) as a "bookkeeping" tool. It is a number assigned to an element in a chemical combination that represents the number of electrons lost or gained by an atom of that element in the compound.

Analogy: Imagine atoms are like people in a business deal. The oxidation number tells us if someone has "invested" (lost) electrons or "acquired" (gained) them.

The Rules of the Game

To calculate oxidation numbers, you need to follow a specific set of rules. It’s like learning the rules of chess—once you know them, the game becomes much easier!

  • Rule 1: The oxidation number of any uncombined element is always 0.
    Example: \(Cl_2\), \(Na\), and \(O_2\) all have an oxidation number of 0.
  • Rule 2: For a simple ion, the oxidation number is the same as its charge.
    Example: \(Mg^{2+}\) is +2; \(Cl^-\) is -1.
  • Rule 3: In a neutral compound, the sum of all oxidation numbers must be 0.
  • Rule 4: In a polyatomic ion, the sum of all oxidation numbers must equal the overall charge of the ion.
  • Rule 5: Some elements are very predictable:
    • Group 1 metals: Always +1.
    • Group 2 metals: Always +2.
    • Fluorine: Always -1 (it's the most "greedy" element!).
    • Hydrogen: Usually +1, except in metal hydrides (like \(NaH\)) where it is -1.
    • Oxygen: Usually -2, except in peroxides (like \(H_2O_2\)) where it is -1, or when bonded to Fluorine.

Using Roman Numerals

When we name compounds where an element could have different oxidation states (like transition metals), we use Roman numerals in brackets.
Example: In Iron(II) chloride, the iron has an oxidation number of +2. In Iron(III) chloride, it is +3.

Quick Review Box:
Sum of Ox. Nos. in a molecule = 0
Sum of Ox. Nos. in an ion = Charge of ion

2. Defining Oxidation and Reduction

The word Redox comes from two processes that always happen at the same time: Reduction and Oxidation.

The "OIL RIG" Mnemonic

This is the most famous memory aid in Chemistry. Use it every time you get confused!

  • Oxidation Is Loss (of electrons)
  • Reduction Is Gain (of electrons)

Changes in Oxidation Number

Sometimes it’s easier to look at the numbers rather than the electrons:

  • Oxidation is an increase in oxidation number (e.g., from 0 to +2).
  • Reduction is a decrease in oxidation number (e.g., from +1 to 0).

Did you know? When a metal reacts, it almost always undergoes oxidation to form positive ions. Non-metals usually undergo reduction to form negative ions.

3. Oxidising and Reducing Agents

This is a common area where students trip up, but here is a simple trick: think of an "agent" as someone who makes something happen to someone else.

  • An oxidising agent oxidises something else. To do this, it must gain electrons (it gets reduced itself).
  • A reducing agent reduces something else. To do this, it must lose electrons (it gets oxidised itself).

Encouragement: Think of a "Travel Agent." They don't go on holiday themselves; they make the holiday happen for you! Similarly, an oxidising agent "gives" oxidation to another atom by taking its electrons.

Key Takeaway:
Oxidising Agent = Electron Taker = Gets Reduced
Reducing Agent = Electron Giver = Gets Oxidised

4. Disproportionation Reactions

Usually, one element is oxidised and a different one is reduced. However, in a disproportionation reaction, the same element in a single species is simultaneously oxidised and reduced.

Example: The decomposition of hydrogen peroxide
\(2H_2O_2 \rightarrow 2H_2O + O_2\)
In \(H_2O_2\), Oxygen is -1.
In \(H_2O\), Oxygen is -2 (Reduced).
In \(O_2\), Oxygen is 0 (Oxidised).
Because the Oxygen went from -1 to both -2 and 0, this is disproportionation.

5. Writing and Balancing Redox Equations

To write a full ionic equation, we often start with two half-equations. One shows the loss of electrons (oxidation) and the other shows the gain of electrons (reduction).

Step-by-Step: Combining Half-Equations

Let's look at the reaction between Zinc and Copper(II) ions:

  1. Write the Oxidation half-equation: \(Zn \rightarrow Zn^{2+} + 2e^-\)
  2. Write the Reduction half-equation: \(Cu^{2+} + 2e^- \rightarrow Cu\)
  3. Check the electrons: Both equations have \(2e^-\). If they were different (e.g., 1 and 2), you would have to multiply the equations so the electrons cancel out.
  4. Add them together: \(Zn + Cu^{2+} + 2e^- \rightarrow Zn^{2+} + 2e^- + Cu\)
  5. Cancel the electrons: \(Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu\)

Common Mistake to Avoid: Never leave electrons in your final full ionic equation! They must always cancel out completely.

6. Summary of Key Concepts

  • Oxidation Number: The charge an atom would have if the compound were purely ionic.
  • Oxidation: Loss of electrons; increase in oxidation number.
  • Reduction: Gain of electrons; decrease in oxidation number.
  • Disproportionation: Simultaneous oxidation and reduction of the same element.
  • Reducing Agent: Lays down (loses) electrons to reduce another species.
  • Oxidising Agent: Picks up (gains) electrons to oxidise another species.

Don't worry if this seems tricky at first! The key to mastering Redox is practice. Start by assigning oxidation numbers to every element you see in an equation, and soon it will become second nature.