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Chemistry studies the vast diversity of matter and the transformations that change one type of matter into another. These transformations occur through various kinds of chemical reactions, and among these, redox reactions form one of the most important categories. Redox reactions are processes in which oxidation and reduction occur simultaneously. Below, we have provided the NCERT Notes Class 11 Chemistry (Part-2) Chapter 7: Redox Reactions notes for your thorough exam prep.
Contents
- 1 Introduction
- 2 Classical Idea of Redox Reactions – Oxidation and Reduction
- 3 Redox Reactions in Terms of Electron Transfer
- 4 Oxidation Number
- 5 Types of Redox Reactions
- 6 Balancing of Redox Reactions
- 7 Redox Reactions as the Basis for Titrations
- 8 Redox Reactions and Electrode Processes
- 9 Important Formulas in NCERT Notes Class 11 Chemistry (Part-II) Chapter 7: Redox Reactions
- 10 FAQs
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Introduction
Redox processes are widespread in both physical and biological systems.
- Physical Processes involving Redox Reactions:
- Combustion of fuels (e.g., coal, petrol, natural gas) for energy generation.
- Corrosion of metals.
- Electrochemical processes such as those in batteries.
- Biological Processes:
- Cellular respiration.
- Photosynthesis.
- Metabolism of nutrients.
- Industrial and Metallurgical Applications:
- Extraction of highly reactive metals (e.g., aluminium, sodium) and non-metals through electrochemical methods.
- Manufacture of important chemical compounds (e.g., caustic soda, chlorine).
- Agricultural Relevance:
- Redox principles are involved in the production of fertilisers and pesticides.
Everyday and Technological Importance
- Energy Production: Burning of fuels for domestic heating, transport, and industrial purposes is a redox process.
- Electrochemical Applications: Dry cells, lead–acid batteries, and fuel cells function via redox principles.
- Metal Protection and Corrosion: Understanding corrosion as an electrochemical redox process allows the design of preventive measures.
- Environmental Issues:
- Hydrogen Economy: Proposes using liquid hydrogen as a clean fuel, based on redox concepts.
- Ozone Layer Depletion: Involves redox changes in atmospheric chemistry leading to the ‘ozone hole’.
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Classical Idea of Redox Reactions – Oxidation and Reduction
The classical idea of redox reactions- oxidation and reduction is discussed below.
Original Definition of Oxidation
The original definition of oxidation is discussed below.
- Oxidation: Addition of oxygen to an element or compound.
- Reason: Presence of ~20% oxygen in the atmosphere; many elements occur naturally as oxides.
- Examples:
- 2Mg (s) + O₂ (g) → 2MgO (s)
- S (s) + O₂ (g) → SO₂ (g)
Extension of the Definition
The extension of the definition is given below.
- Observation: In CH₄ oxidation, hydrogen is replaced by oxygen.
- Broader view: Oxidation includes the removal of hydrogen from a substance.
- Examples:
- CH₄ (g) + 2 O₂ (g) → CO₂ (g) + 2H₂O (l)
- 2H₂S (g) + O₂ (g) → 2 S (s) + 2H₂O (l)
Further Extension
The concept of oxidation can be further extended as follows
- Oxidation includes the addition of electronegative elements other than oxygen.
- Examples:
- Mg (s) + F₂ (g) → MgF₂ (s)
- Mg (s) + Cl₂ (g) → MgCl₂ (s)
- Mg (s) + S (s) → MgS (s)
Removal of Electropositive Elements as Oxidation
Example: 2K₄[Fe(CN)₆] (aq) + H₂O₂ (aq) → 2K₃[Fe(CN)₆] (aq) + 2KOH (aq)
Original Definition of Reduction
Reduction: Removal of oxygen from a compound.
Extended Definition of Reduction
Includes:
- Removal of oxygen or any electronegative element.
- Addition of hydrogen or any electropositive element.
- Removal of oxygen:
- 2HgO (s) → 2Hg (l) + O₂ (g)
- Removal of electronegative element:
- 2FeCl₃ (aq) + H₂ (g) → 2FeCl₂ (aq) + 2HCl (aq)
- Addition of hydrogen:
- CH₂ = CH₂ (g) + H₂ (g) → C₂H₆ (g)
- Addition of an electron-positive element:
- 2HgCl₂ (aq) + SnCl₂ (aq) → Hg₂Cl₂ (s) + SnCl₄ (aq)
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Redox Reactions in Terms of Electron Transfer
In this section, we have discussed the redox reactions in terms of electron transfer.
Examples of Redox as Electron Transfer
The following are examples of redox reactions as electron transfer.
- Reactions:
- 2Na (s) + Cl₂ (g) → 2NaCl (s)
- 4Na (s) + O₂ (g) → 2Na₂O (s)
- 2Na (s) + S (s) → Na₂S (s)
- In each reaction:
- Sodium is oxidised (addition of oxygen or a more electronegative element).
- Cl₂, O₂, and S are reduced (addition of electropositive sodium).
- Compounds formed are ionic: NaCl⁻, (Na⁺)₂O²⁻, (Na⁺)₂S²⁻.
Half-Reaction Concept
The half-reaction concept is discussed below.
- Each redox process can be split into:
- Oxidation half: loss of electrons.
- Reduction half: gain of electrons.
- Example (formation of NaCl):
- Oxidation: 2 Na (s) → 2 Na⁺ (g) + 2 e⁻
- Reduction: Cl₂ (g) + 2 e⁻ → 2 Cl⁻ (g)
- Overall: 2 Na (s) + Cl₂ (g) → 2 NaCl (s)
Definitions (Electron Transfer Concept)
The following are the definitions of oxidation, reduction, oxidising, and reducing agent.
- Oxidation: Loss of electron(s) by a species.
- Reduction: Gain of electron(s) by a species.
- Oxidising agent: Acceptor of electron(s).
- Reducing agent: Donor of electron(s).
Oxidation Number
Example: Formation of water- 2H₂ (g) + O₂ (g) → 2H₂O (l)
- H changes from 0 (in H₂) to +1 (in H₂O).
- O changes from 0 (in O₂) to –2 (in H₂O).
- Indicates electron shift from H to O (partial in covalent bonds).
- Oxidation number method:
- Assumes complete transfer of electrons from the less electronegative to the more electronegative atom.
- Used for bookkeeping in redox reactions.
Oxidation Number Representation
Oxidation number: Oxidation state of an element in a compound, calculated assuming full electron transfer to the more electronegative element.
Examples:
- 2H₂ (0) + O₂ (0) → 2H₂(+1)O(–2)
- H₂ (0) + Cl₂ (0) → 2H(+1)Cl(–1)
- CH₄ (C–4H+1) + 4Cl₂ (0) → CCl₄ (C+4Cl–1) + 4H(+1)Cl(–1)
Rules for Determining Oxidation Number
Below, we have discussed the rules for determining oxidation number.
- Elemental state → Oxidation number = 0 (e.g., H₂, O₂, Cl₂, Na).
- Monatomic ions → Equal to charge (Na⁺ = +1, Cl⁻ = –1).
- Alkali metals: always +1.
- Alkaline earth metals: always +2.
- Al in compounds: +3.
- Oxygen: Usually –2.
- In peroxides: –1.
- In superoxides: –½.
- In compounds with fluorine: positive (+2 in OF₂, +1 in O₂F₂).
- Hydrogen: +1 (except in metal hydrides: –1).
- Fluorine: –1 in all compounds.
- Other halogens: –1 unless combined with oxygen.
- Sum Rule:
- Neutral compound → Sum = 0.
- Polyatomic ion → Sum = ion charge.
Oxidation-Reduction Definitions (Oxidation Number Concept)
The oxidation and reduction reactions are defined on the basis of oxidation number concepts.
- Oxidation: Increase in oxidation number.
- Reduction: Decrease in oxidation number.
- Oxidising agent: Increases the oxidation number of other species (gets reduced).
- Reducing agent: Decreases the oxidation number of other species (gets oxidised).
- Redox reaction: Involves a change in the oxidation number of species.
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Types of Redox Reactions
The types of redox reactions are discussed below.
Combination Reactions
The combination reactions are discussed as follows.
- General: A + B → C
- Either/both reactants in elemental form.
- Examples:
- C (0) + O₂ (0) → CO₂ (C+4O–2)
- 3Mg (0) + N₂ (0) → Mg₃N₂ (Mg+2N–3)
- CH₄ (C–4) + 2O₂ (0) → CO₂ (C+4) + 2H₂O (H+1O–2)
Decomposition Reactions
The decomposition reactions are discussed as follows.
- Opposite of combination: one compound → 2+ components, at least one elemental.
- Examples:
- 2H₂O (H+1O–2) → 2H₂ (0) + O₂ (0)
- 2NaH (Na+1H–1) → 2Na (0) + H₂ (0)
- 2KClO₃ (K+1Cl+5O–2) → 2KCl (K+1Cl–1) + 3O₂ (0)
Displacement Reactions
The displacement reactions are discussed as follows.
a) Metal Displacement
- More reactive metal displaces less reactive metal from compound.
- Examples:
- CuSO₄ + Zn → ZnSO₄ + Cu
- V₂O₅ + 5Ca → 2V + 5CaO
- TiCl₄ + 2Mg → Ti + 2MgCl₂
- Cr₂O₃ + 2Al → Al₂O₃ + 2Cr
b) Non-Metal Displacement
(i) Hydrogen Displacement
- Active metals (alkali, alkaline earth) from water:
- 2Na + 2H₂O → 2NaOH + H₂
- Ca + 2H₂O → Ca(OH)₂ + H₂
- Less active metals with steam:
- Mg + 2H₂O → Mg(OH)₂ + H₂
- 2Fe + 3H₂O → Fe₂O₃ + 3H₂
- Metals with acids:
- Zn + 2HCl → ZnCl₂ + H₂
- Mg + 2HCl → MgCl₂ + H₂
- Fe + 2HCl → FeCl₂ + H₂
(ii) Halogen Displacement
- Reactivity order: F₂ > Cl₂ > Br₂ > I₂.
- Examples:
- 2H₂O + 2F₂ → 4HF + O₂
- Cl₂ + 2KBr → 2KCl + Br₂
- Cl₂ + 2KI → 2KCl + I₂
- Br₂ + 2I⁻ → 2Br⁻ + I₂
Disproportionation Reactions
The disproportionation reactions are discussed as follows.
- An element in one oxidation state is simultaneously oxidised and reduced.
- Condition: Element can exist in ≥3 oxidation states.
- Examples:
- 2H₂O₂ → 2H₂O + O₂ (O: –1 → –2 and 0)
- P₄ + 3OH⁻ + 3H₂O → PH₃ + 3H₂PO₂⁻ (P: 0 → –3 and +1)
- S₈ + 12OH⁻ → 4S²⁻ + 2S₂O₃²⁻ + 6H₂O (S: 0 → –2 and +2)
- Cl₂ + 2OH⁻ → ClO⁻ + Cl⁻ + H₂O (Cl: 0 → +1 and –1)
- Fluorine exception: Cannot disproportionate (most electronegative, no positive oxidation state).
- 2F₂ + 2OH⁻ → 2F⁻ + OF₂ + H₂O
Balancing of Redox Reactions
Two main methods are used to balance redox equations:
Oxidation Number Method
The steps of the oxidation number method are:
Steps:
- Write correct formulas for all reactants and products.
- Assign oxidation numbers to all elements and identify the atoms whose oxidation numbers change.
- Calculate the change in oxidation number per atom and for the entire molecule/ion. Multiply by suitable numbers to make total increases equal total decreases.
- Adjust charges:
- If in an acidic medium, add H⁺ ions.
- If in basic medium, add OH⁻ ions.
- Balance hydrogen atoms by adding H₂O molecules, then check and balance oxygen atoms.
- Ensure the equation is balanced for both atoms and charges.
Half Reaction Method
Steps (Example: Oxidation of Fe²⁺ to Fe³⁺ by Cr₂O₇²⁻ in acidic medium):
- Write the unbalanced ionic equation:
Fe²⁺ + Cr₂O₇²⁻ → Fe³⁺ + Cr³⁺ - Split into half reactions:
- Oxidation: Fe²⁺ → Fe³⁺
- Reduction: Cr₂O₇²⁻ → Cr³⁺
- Balance atoms other than O and H:
Cr₂O₇²⁻ → 2Cr³⁺ - Balance oxygen by adding H₂O and balance hydrogen by adding H⁺:
Cr₂O₇²⁻ + 14H⁺ → 2Cr³⁺ + 7H₂O - Balance charges by adding electrons:
- Oxidation: Fe²⁺ → Fe³⁺ + e⁻
- Reduction: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O
- Multiply oxidation half reaction by 6 to equalise electrons:
6Fe²⁺ → 6Fe³⁺ + 6e⁻ - Add half reactions and cancel electrons:
6Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ → 6Fe³⁺ + 2Cr³⁺ + 7H₂O - Verify the balance of atoms and charges.
- For basic medium, first balance as acidic, then add OH⁻ ions equal to H⁺ ions on both sides, combining OH⁻ and H⁺ to form H₂O.
Redox Reactions as the Basis for Titrations
Redox titrations are used to determine the strength of a reductant or oxidant using a redox-sensitive indicator.
Types of Indicators:
- Self-indicator reagents:
- Example: KMnO₄ (MnO₄⁻) – acts as its own indicator.
- Endpoint: appearance of a faint pink colour when all reductant is consumed.
- External indicators:
- Example: K₂Cr₂O₇ with diphenylamine indicator.
- Diphenylamine turns blue just after the equivalence point.
- Iodometric method:
- Oxidant liberates iodine from iodide ions:
2Cu²⁺ + 4I⁻ → Cu₂I₂ + I₂ - I₂ reacts with starch to give an intense blue colour.
- I₂ is titrated with sodium thiosulfate:
I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻ - Blue colour disappears at the endpoint.
- Oxidant liberates iodine from iodide ions:
Limitations of the Concept of Oxidation Number
The oxidation number approach assumes complete electron transfer, which is not always the case in covalent compounds (electron shift is often partial).
- The modern view:
- Oxidation = decrease in electron density around the atom.
- Reduction = increase in electron density around the atom.
- This reflects the gradual shift in understanding redox processes beyond simple integer oxidation states.
Redox Reactions and Electrode Processes
The redox reactions and electrode processes are discussed below.
Direct vs. Indirect Electron Transfer
- Example: Zn rod in CuSO₄ solution.
- Direct reaction:
- Zn is oxidised to Zn²⁺.
- Cu²⁺ is reduced to Cu metal.
- Electrons transfer directly from Zn to Cu²⁺.
- Heat is evolved.
- Modification for indirect electron transfer:
- Separate the two half reactions into different containers.
- Use zinc sulphate solution with Zn rod and copper sulphate solution with Cu rod in separate beakers.
Redox Couples
A Pair consisting of oxidised and reduced forms of a species taking part in a half reaction. Oxidised form / Reduced form (vertical line represents phase/interface). Examples in this experiment
- Zn²⁺/Zn
- Cu²⁺/Cu
- The oxidised form is always written first.
Daniell Cell Setup
The Daniell cell setup is discussed as follows.
- Beaker 1: Zn rod in ZnSO₄ solution.
- Beaker 2: Cu rod in CuSO₄ solution.
- Connected by:
- Salt bridge (U-tube with KCl or NH₄NO₃ gelled with agar-agar).
- Maintains electrical neutrality.
- Allows ion migration without mixing solutions.
- Metallic wire between rods (with ammeter and switch).
Observations When the Switch is ON
The following are the observations when the switch is on.
- Electrons flow through the metallic wire from Zn to Cu.
- Ionic current between solutions flows via the salt bridge.
- The flow of current is possible due to a potential difference between the electrodes.
Electrode Potential
Potential associated with an electrode is called electrode potential. Standard conditions for Standard Electrode Potential (E°):
- All species have unit concentration (1 mol dm⁻³).
- Gases at 1 atm pressure (if present).
- Temperature: 298 K.
- Reference electrode: Standard Hydrogen Electrode (SHE) has E° = 0.00 V (by convention)
Interpretation of E° values
The interpretation of E° values is given below.
- Negative E°: The Redox couple is a stronger reducing agent than H⁺/H₂.
- Positive E°: The Redox couple is a weaker reducing agent than H⁺/H₂.
- E° values indicate the relative tendency of species to exist in oxidised or reduced form.
Important Formulas in NCERT Notes Class 11 Chemistry (Part-II) Chapter 7: Redox Reactions
Here are the important formulas and equations from NCERT Notes Class 11 Chemistry (Part-2) Chapter 7: Redox Reactions
- Oxidation Number = (Charge on atom in free state) or as per rules of electronegativity
- Oxidation: Increase in oxidation number of an element.
- Reduction: Decrease in oxidation number of an element.
- Redox reaction: A Change in oxidation number of interacting species indicates a redox reaction.
- Balancing via Oxidation Number Method: Increase in ON = Decrease in ON (per atom basis × number of atoms)
- Oxidation Half Reaction: Fe²⁺ → Fe³⁺ + e⁻
- Reduction Half Reaction: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O
- Overall Redox Reaction: 6Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ → 6Fe³⁺ + 2Cr³⁺ + 7H₂O
- Electrode Potential: E°cell = E°cathode − E°anode
- Nernst Equation: Ecell = E°cell − (0.0591 / n) log Q
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FAQs
Ans: Oxidation number denotes the oxidation state of an element in a compound, determined according to rules based on the assumption that the electron pair in a covalent bond belongs entirely to the more electronegative element.
Ans: A redox couple consists of the oxidised and reduced forms of a substance taking part in an oxidation or reduction half-reaction, represented as oxidised form/reduced form (e.g., Zn²⁺/Zn, Cu²⁺/Cu).
Ans: Oxidation is an increase in the oxidation number of an element in a given substance, whereas reduction is a decrease in the oxidation number of an element in a given substance.
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