In simple terms
A friendly intro before the formal notes — no formulas yet.
The Electron Exchange
Redox reactions are all about the transfer of electrons from one chemical species to another. We use a system called oxidation numbers to keep track of this electron movement.
Imagine a seesaw on a playground. For one person to go up (oxidation, losing the 'weight' of electrons), another person must go down (reduction, gaining the 'weight' of electrons). The two movements are linked and happen at the same time. Oxidation numbers are like measuring the height of each person on the seesaw to see who went up and who went down.
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Oxidation is electron loss; reduction is electron gain (OIL RIG).
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Oxidation number rules: element = 0; ion = charge; O usually −2, H usually +1.
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Redox reactions involve simultaneous oxidation and reduction.
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Half-equations balance atoms and charge; combine for full ionic equation.
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Full topic notes
Formal explanation with the rigour you need for the exam.
Defining Redox: Electron Transfer
At its heart, a redox reaction is all about electrons moving from one substance to another. We use the mnemonic OIL RIG to remember the definitions:
- Oxidation Is Loss of electrons.
- Reduction Is Gain of electrons.
Because electrons cannot be created or destroyed in a chemical reaction, if one species loses electrons, another must gain them. Therefore, oxidation and reduction always occur simultaneously.
Oxidation Numbers: A System for Electron Accounting
While we can think about direct electron transfer, it's often more practical to use a bookkeeping tool called oxidation numbers (or oxidation states). An oxidation number is the hypothetical charge an atom would have if all its bonds to different elements were 100% ionic. We assign these numbers using a clear set of rules.
Elements: The oxidation number of an atom in an uncombined element is 0 (e.g., , , ).
Simple Ions: The oxidation number of an atom in a monatomic ion is equal to the charge on the ion (e.g., is +1, is -2).
Compounds & Ions: The sum of oxidation numbers in a neutral compound is 0. The sum in a polyatomic ion equals the ion's charge.
Specific Elements (Hierarchy):
- Group 1 metals are +1, Group 2 metals are +2 in their compounds.
- Fluorine is always -1 in its compounds.
- Hydrogen is usually +1 (except in metal hydrides like , where it is -1).
- Oxygen is usually -2 (except in peroxides like where it's -1, and with fluorine like where it's +2).
Identifying Oxidising and Reducing Agents
Once we can assign oxidation numbers, we can easily identify what is being oxidised and reduced. An increase in oxidation number signifies oxidation. A decrease in oxidation number signifies reduction. This leads to the definitions of oxidising and reducing agents. A reducing agent is the species that donates electrons and gets oxidised itself. An oxidising agent is the species that accepts electrons and gets reduced itself.
A very common mistake is to confuse the species being oxidised/reduced with the agent. Remember: the oxidising agent gets reduced, and the reducing agent gets oxidised. Write it down in the exam to be sure!
Constructing Redox Equations
For complex redox reactions, it is easiest to split the overall reaction into two half-equations: one for the oxidation process and one for the reduction process. We balance each half-equation for atoms and charge separately, then combine them to give the overall balanced ionic equation.
Step 1: Write the unbalanced equation for the species being oxidised/reduced.
Step 2: Balance all atoms except for O and H.
Step 3: Balance oxygen atoms by adding molecules.
Step 4: Balance hydrogen atoms by adding ions (for acidic/neutral solutions).
Step 5: Balance the charges by adding electrons () to the more positive side.
Step 6: Multiply the half-equations by integers so the number of electrons lost in oxidation equals the number of electrons gained in reduction.
Step 7: Add the balanced half-equations and cancel out any species that appear on both sides.
Disproportionation: One Element, Two Fates
Disproportionation is a specific type of redox reaction where an element in a single reactant is simultaneously oxidised and reduced. A classic example is the reaction of chlorine gas with cold, dilute sodium hydroxide solution:
Here, the chlorine atom (oxidation number 0 in ) is reduced to -1 in and oxidised to +1 in .
Worked examples
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Determine the oxidation number of chromium in the dichromate(VI) ion, .
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The overall charge of the ion is -2, so the sum of oxidation numbers must be -2.
Consider the reaction: . Identify the oxidising agent and the reducing agent, justifying your answer with oxidation numbers.
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Assign oxidation numbers to Fe:
How it all connects
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Glossary
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Quick check
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Revision flashcards
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What is a redox reaction?
A reaction involving both reduction and oxidation, characterised by the transfer of electrons and changes in oxidation numbers.
Key takeaways
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Elements: The oxidation number of an atom in an uncombined element is 0 (e.g., , , ).
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Simple Ions: The oxidation number of an atom in a monatomic ion is equal to the charge on the ion (e.g., is +1, is -2).
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Compounds & Ions: The sum of oxidation numbers in a neutral compound is 0. The sum in a polyatomic ion equals the ion's charge.
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Specific Elements (Hierarchy):
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- Group 1 metals are +1, Group 2 metals are +2 in their compounds.
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- Fluorine is always -1 in its compounds.
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- Hydrogen is usually +1 (except in metal hydrides like , where it is -1).
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- Oxygen is usually -2 (except in peroxides like where it's -1, and with fluorine like where it's +2).
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