Redox Reactions is built on a single idea: oxidation is loss of electrons and reduction is gain of electrons, and the two always happen together. The chapter teaches you to assign oxidation numbers, identify oxidising and reducing agents, classify reactions into combination, decomposition, displacement and disproportionation, and balance equations by the ion-electron method. It is a high-yield, formula-light NEET topic that also sets up Electrochemistry, so the oxidation-number and electrode-potential skills you build here pay off across the whole inorganic and physical chemistry syllabus.
Table of Contents
- Key Concepts - Oxidation & reduction, oxidation number, redox couples, types of redox reactions, balancing methods, electrode processes
- Weightage in Board & Entrance Exams
- Important Definitions
- Solved Examples
- Important Questions for Board Exams
- Quick Revision Points
Key Concepts
1. Oxidation and Reduction - the Classical View
The oldest way to spot a redox reaction is to watch oxygen and hydrogen move around. When a magnesium ribbon burns to form MgO, magnesium gains oxygen - it is oxidised. When CuO is heated with hydrogen to give copper, the oxide loses oxygen - it is reduced.
- Oxidation: addition of oxygen OR removal of hydrogen.
- Reduction: removal of oxygen OR addition of hydrogen.
This view works for simple cases but fails for reactions with no oxygen or hydrogen at all - so chemists upgraded to the electronic concept.
2. The Electron-Transfer (Modern) Concept
Every redox reaction is really electrons leaving one species and arriving at another. When zinc reacts with copper sulphate, zinc atoms hand over two electrons each - that is oxidation. Copper ions accept those electrons - that is reduction.
Oxidation = loss of electrons. Reduction = gain of electrons. The memory hook is OIL RIG - Oxidation Is Loss, Reduction Is Gain.
- Oxidation half: Zn → Zn²⁺ + 2e⁻ (electrons lost)
- Reduction half: Cu²⁺ + 2e⁻ → Cu (electrons gained)
Oxidation and reduction always happen together - electrons cannot simply vanish - which is why these are called redox (reduction–oxidation) reactions.
3. Oxidising and Reducing Agents
The species that accepts electrons is itself reduced, and by stripping electrons off the other species it acts as the oxidising agent. The species that donates electrons is itself oxidised, so it acts as the reducing agent.
- Oxidising agent (oxidant): gains electrons, gets reduced (e.g., KMnO₄, K₂Cr₂O₇, O₂, F₂).
- Reducing agent (reductant): loses electrons, gets oxidised (e.g., H₂, C, Zn, SO₂).
Memory tip: the oxidising agent is the one that gets reduced - it does the opposite to itself of what its name suggests.
4. Oxidation Number (Oxidation State)
Oxidation number is the imaginary charge an atom would carry if every bond in the molecule were assumed to be fully ionic. It lets us track electron transfer even in covalent compounds where no real ions exist.
Rules for Assigning Oxidation Number
- Free elements (O₂, H₂, Na, Cl₂) have oxidation number 0.
- For a monatomic ion, it equals the ionic charge (Na⁺ = +1, S²⁻ = −2).
- Oxygen is usually −2 (but −1 in peroxides like H₂O₂, and +2 in OF₂).
- Hydrogen is usually +1 (but −1 in metal hydrides like NaH).
- Fluorine is always −1; alkali metals always +1; alkaline earth metals always +2.
- Sum of oxidation numbers = 0 for a neutral molecule, and = charge for a polyatomic ion.
Worked rule example: In KMnO₄, K = +1 and O = −2 (×4 = −8). So +1 + Mn + (−8) = 0 → Mn = +7.
5. Redox in Terms of Oxidation Number
The electron-transfer idea is restated cleanly using oxidation numbers, which is the form examiners love.
- Oxidation: an increase in oxidation number.
- Reduction: a decrease in oxidation number.
In Zn + Cu²⁺ → Zn²⁺ + Cu, zinc goes 0 → +2 (oxidised) and copper goes +2 → 0 (reduced). Spotting the change in oxidation number is the fastest way to identify redox in an exam.
6. Types of Redox Reactions
CBSE classifies redox reactions into four neat families. Recognising the family often gives you the products instantly.
(a) Combination Reactions
Two or more substances combine into one, with at least one element changing oxidation number.
C + O₂ → CO₂ (carbon: 0 → +4, oxidised)
(b) Decomposition Reactions
A single compound breaks into two or more products; it is the reverse of combination.
2H₂O → 2H₂ + O₂ (hydrogen +1 → 0, oxygen −2 → 0)
(c) Displacement Reactions
One element displaces another from its compound. These can be metal or non-metal displacements.
Zn + CuSO₄ → ZnSO₄ + Cu (more reactive Zn displaces Cu)
(d) Disproportionation Reactions
The same element in an intermediate oxidation state is simultaneously oxidised and reduced.
2H₂O₂ → 2H₂O + O₂ - oxygen in H₂O₂ (−1) goes to −2 in water (reduced) and to 0 in O₂ (oxidised).
[TABLE: Type → key feature → example. Combination: many→one; Decomposition: one→many; Displacement: element swaps in; Disproportionation: same element up and down.]
7. Balancing Redox Reactions - Oxidation Number Method
This method balances the total increase and decrease in oxidation number, since electrons lost must equal electrons gained.
Steps:
- Write the skeletal equation and assign oxidation numbers.
- Identify atoms whose oxidation number changes; find the increase and decrease.
- Multiply by suitable integers so total increase = total decrease.
- Balance the remaining atoms (O by adding H₂O, H by adding H⁺), then balance charge.
Example skeleton: in MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺, Mn falls +7 → +2 (gain 5e⁻) and Fe rises +2 → +3 (lose 1e⁻), so 5 Fe²⁺ are needed per MnO₄⁻.
8. Balancing Redox Reactions - Half-Reaction (Ion-Electron) Method
Here the reaction is split into an oxidation half and a reduction half, each balanced separately, then recombined. This is the cleaner method for solutions.
Steps (acidic medium):
- Split into two half-reactions.
- Balance atoms other than O and H first.
- Balance O by adding H₂O, balance H by adding H⁺.
- Balance charge by adding electrons (e⁻).
- Equalise electrons in both halves, then add and cancel.
Basic medium: do the same, then add OH⁻ to both sides to neutralise H⁺ into H₂O.
Reduction half example: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. Oxidation half: Fe²⁺ → Fe³⁺ + e⁻. Multiply the iron half by 5 and add.
9. Redox Reactions and Electrode Processes
If we physically separate the two half-reactions, the electrons are forced to travel through an external wire instead of jumping directly - and a flowing electron is electricity. This is the leap from chemistry to a working cell.
- Anode: oxidation occurs here (electrons released); it is the negative terminal in a galvanic cell.
- Cathode: reduction occurs here (electrons consumed); it is the positive terminal in a galvanic cell.
Memory hook: An Ox (Anode = Oxidation) and Red Cat (Reduction = Cathode).
10. Electrochemical (Galvanic) Cells - an Introduction
A galvanic (voltaic) cell converts the chemical energy of a spontaneous redox reaction directly into electrical energy. The classic example is the Daniell cell, built from the Zn–Cu reaction.
[DIAGRAM: Daniell cell - Zn rod in ZnSO₄ (anode, −) connected through a voltmeter to a Cu rod in CuSO₄ (cathode, +), with a salt bridge joining the two beakers.]
- At the anode: Zn → Zn²⁺ + 2e⁻ (oxidation).
- At the cathode: Cu²⁺ + 2e⁻ → Cu (reduction).
- Salt bridge: completes the circuit and keeps both solutions electrically neutral.
The electrode potential of each electrode measures its tendency to lose or gain electrons; the difference between the two gives the cell’s EMF. Standard hydrogen electrode (SHE) is taken as the zero reference (0 V).
11. Applications of Redox Reactions
Redox is not just an exam topic - it runs the modern world.
- Batteries and cells: dry cells, lead storage batteries, fuel cells.
- Corrosion: rusting of iron is the oxidation of iron - a costly redox process.
- Metallurgy: extraction of metals by reduction of their ores (e.g., iron from haematite).
- Bleaching and disinfection: by oxidising agents like Cl₂ and KMnO₄.
- Photosynthesis and respiration: the redox engines of life.
Read the rest of the chapter →Hide the rest ↑
Weightage in Board & Entrance Exams
| Exam | Typical Weightage | Most-Tested Areas |
|---|---|---|
| CBSE Board (Class 11) | 6–8 marks | Oxidation number, identifying oxidant/reductant, balancing redox equations |
| JEE Main / Advanced | 1–2 questions | Oxidation states, disproportionation, balancing in acidic/basic medium |
| NEET | 1–2 questions | Oxidation number, types of redox reactions, oxidising/reducing agents |
[TABLE: Question-type split - VSA (1 mark): definitions, oxidation-number calculation; SA (2–3 marks): identify oxidant/reductant, classify reaction type; LA (5 marks): balancing redox equations by both methods, Daniell cell.]
Important Definitions
| Term | Definition |
|---|---|
| Oxidation | Loss of electrons / increase in oxidation number / addition of oxygen |
| Reduction | Gain of electrons / decrease in oxidation number / removal of oxygen |
| Redox reaction | A reaction in which oxidation and reduction occur simultaneously |
| Oxidising agent | Species that gains electrons and is itself reduced |
| Reducing agent | Species that loses electrons and is itself oxidised |
| Oxidation number | Imaginary charge on an atom if all bonds were assumed ionic |
| Disproportionation | Same element in one oxidation state is both oxidised and reduced |
| Anode | Electrode where oxidation occurs (negative terminal in a galvanic cell) |
| Cathode | Electrode where reduction occurs (positive terminal in a galvanic cell) |
| Galvanic cell | Device that converts chemical energy of a spontaneous redox reaction into electrical energy |
Solved Examples
Example 1
Find the oxidation number of sulphur in H₂SO₄.
Answer: H = +1 (×2 = +2), O = −2 (×4 = −8). So 2 + S − 8 = 0 → S = +6.
Example 2
In the reaction Zn + CuSO₄ → ZnSO₄ + Cu, identify the oxidising and reducing agents.
Answer: Zn goes 0 → +2 (oxidised) so it is the reducing agent; Cu²⁺ goes +2 → 0 (reduced), so CuSO₄ is the oxidising agent.
Example 3
Find the oxidation number of chromium in K₂Cr₂O₇.
Answer: K = +1 (×2 = +2), O = −2 (×7 = −14). So 2 + 2Cr − 14 = 0 → 2Cr = 12 → Cr = +6.
Example 4
Show that 2H₂O₂ → 2H₂O + O₂ is a disproportionation reaction.
Answer: Oxygen in H₂O₂ is −1. In H₂O it becomes −2 (reduced) and in O₂ it becomes 0 (oxidised). The same element is both oxidised and reduced, so it is disproportionation.
Example 5
Balance the reduction half-reaction for MnO₄⁻ to Mn²⁺ in acidic medium.
Answer: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. (O balanced by 4 H₂O, H by 8 H⁺, charge by 5 electrons.)
Example 6
Find the oxidation number of nitrogen in NH₄⁺ and in NO₃⁻.
Answer: In NH₄⁺: N + 4(+1) = +1 → N = −3. In NO₃⁻: N + 3(−2) = −1 → N = +5.
Important Questions for Board Exams
1-Mark Questions (VSA)
- Define oxidation in terms of electrons.
- What is the oxidation number of oxygen in OF₂?
- Name the species that gets reduced in a redox reaction.
- What is a disproportionation reaction? Give one example.
- At which electrode does oxidation occur in a galvanic cell?
2–3-Mark Questions (SA)
- Identify the oxidising and reducing agents in: MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O.
- Explain why the same element can be both oxidised and reduced in a disproportionation reaction, using an example.
- Calculate the oxidation number of phosphorus in H₃PO₄ and of manganese in KMnO₄.
- Distinguish between a combination and a displacement redox reaction with one example each.
5-Mark Questions (LA)
- Balance the equation MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ in acidic medium by the ion-electron (half-reaction) method.
- Describe the construction and working of a Daniell cell. Write the electrode reactions and state the function of the salt bridge.
- Balance Cr₂O₇²⁻ + SO₃²⁻ → Cr³⁺ + SO₄²⁻ in acidic medium and identify the oxidising agent.
Quick Revision Points
- OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons)
- Oxidation = increase in oxidation number; reduction = decrease
- Oxidising agent is itself reduced; reducing agent is itself oxidised
- Oxidation number of free elements = 0; O is usually −2; H usually +1; F always −1
- Four types: combination, decomposition, displacement, disproportionation
- Disproportionation: same element both oxidised and reduced (e.g., H₂O₂)
- Balancing: oxidation-number method (balance net change) or half-reaction method (split, balance, recombine)
- In acidic medium balance O with H₂O and H with H⁺; in basic medium add OH⁻
- An Ox, Red Cat: Anode = Oxidation, Cathode = Reduction
- Galvanic cell converts chemical energy of a spontaneous redox reaction into electricity; salt bridge keeps it neutral
- SHE is the reference electrode with potential 0 V
Next Chapter: Equilibrium Class 11 Notes
Class 11 Chemistry · Chapter 8 – swipe through all 9 cards to understand the whole chapter.
Oxidation & Reduction (OIL RIG)
Redox is all about electron transfer, and the two halves always happen together.
No oxidation without simultaneous reduction.
- Oxidation = loss of e⁻; reduction = gain of e⁻
- Zn → Zn2⁺ + 2e⁻ (oxidised); Cu2⁺ + 2e⁻ → Cu (reduced)
- Old views: gain O / lose H = oxidation, but e⁻ view is the boss
Oxidation Number Rules
Oxidation number is the charge an atom would have if every bond were ionic.
Apply rules in order; O.N. can be fractional (an average).
- Free element = 0; monatomic ion = its charge; F always −1
- O = −2 (peroxide −1, superoxide −½, OF2 = +2); H = +1 (−1 in metal hydrides)
- KMnO4: Mn = +7; H2SO4: S = +6; S4O62⁻: S = +2.5 (avg)
Oxidising vs Reducing Agents
Each agent does the opposite to itself of what its name suggests.
Highest O.N. = oxidant only; lowest = reductant only.
- In MnO4⁻ + Fe2⁺: MnO4⁻ is oxidant (Mn +7→+2), Fe2⁺ is reductant (+2→+3)
- S in H2SO4 / Mn in KMnO4 (max O.N.) can only oxidise
- Intermediate states (H2O2, SO2) act as both — a NEET trap
Types of Redox Reactions
Most redox reactions sort into four named patterns.
All involve a change in oxidation number.
- Combination C + O2 → CO2; Decomposition 2KClO3 → 2KCl + 3O2
- Displacement: Zn + CuSO4 → ZnSO4 + Cu; Cl2 + 2KBr → 2KCl + Br2
- Disproportionation: same element oxidised AND reduced
Disproportionation
One element in an intermediate oxidation state splits to both higher and lower states.
Reverse process (two states → one) = comproportionation.
- 2H2O2 → 2H2O + O2 : O goes −1→−2 (reduced) and −1→0 (oxidised)
- Needs an INTERMEDIATE O.N. to occur
- Element at its highest or lowest O.N. cannot disproportionate
Balancing: Ion-Electron Method
Split into half-reactions, balance each, then combine equal electrons.
Basic medium: add OH⁻ to both sides to neutralise H⁺.
- Balance O with H2O, H with H⁺, charge with e⁻
- Equalise e⁻ in both halves before adding
- Final atom + charge check — H⁺/H2O is the common slip
Electrode Potential & E°cell
Each half-reaction has a standard electrode potential measured against the SHE.
ΔG° = −nFE°, so E°cell > 0 ⇒ spontaneous.
- More positive E° = stronger oxidant (F2 = +2.87 V)
- More negative E° = stronger reductant (Li = −3.05 V)
- Reaction spontaneous when E°cell is positive
Activity Series & Displacement
A metal displaces any metal ion lying below it in the activity series.
Lower / more-negative E° metal is more easily oxidised.
- Zn is oxidised more easily than Cu, so it displaces Cu2⁺
- Order of E° gives the reactivity (activity) series
- Predicts which metal-metal-ion reactions actually happen
Redox Titrations
Strong oxidants like KMnO4 and K2Cr2O7 are used to titrate reductants.
Titrations rely directly on electrode-potential strengths.
- KMnO4 and K2Cr2O7 are common volumetric oxidants
- Acidic KMnO4 needs no external indicator
- Links E° strength to a high-yield NEET theme
📝 Practice Redox Reactions — 10 NEET PYQs
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Frequently Asked Questions
A redox reaction is one in which electrons are transferred from one species to another, so oxidation (loss of electrons) and reduction (gain of electrons) occur at the same time. A simple memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons.
Apply the rules in order: a free element is 0, a monatomic ion equals its charge, fluorine is always minus 1, oxygen is minus 2 (minus 1 in peroxides), and hydrogen is plus 1 (minus 1 in metal hydrides). The sum of oxidation numbers is 0 in a neutral molecule and equals the charge in an ion, and an oxidation number can be a fractional average, as for sulphur in S4O6 two minus.
The oxidising agent (oxidant) gains electrons and is itself reduced, so its oxidation number decreases, while the reducing agent (reductant) loses electrons and is itself oxidised, so its oxidation number increases. For example, in MnO4 minus reacting with Fe two plus, MnO4 minus is the oxidant (Mn goes plus 7 to plus 2) and Fe two plus is the reductant (plus 2 to plus 3).
In disproportionation the same element in one intermediate oxidation state is simultaneously oxidised and reduced, as when chlorine forms both Cl minus and ClO minus in alkali. In a displacement reaction one element pushes out another from its compound, such as zinc displacing copper from copper sulphate, and the two elements are different.
Yes, Redox Reactions is a high-yield Class 11 chemistry chapter that is firmly in the NEET syllabus and also forms the base for Electrochemistry. Focus on quick oxidation-number assignment, spotting agents, the four reaction types, and balancing by the ion-electron method, since most questions are application based rather than memory based.