Unit 6
General and Physical Chemistry
Class 11 Chemistry
Oxidation and Reduction
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NEB/CDC syllabus scope: Unit 6 – Oxidation and Reduction is a 5-teaching-hour General and Physical Chemistry unit. It covers oxidation and reduction through the electronic concept; oxidation number and its rules; identification of oxidation, reduction, oxidizing agents and reducing agents from oxidation-number changes; balancing redox reactions by oxidation-number and ion-electron/half-reaction methods; and qualitative plus quantitative treatment of Faraday’s laws of electrolysis.
1. Oxidation and Reduction: Redox Overview
Redox reaction
A reaction in which oxidation and reduction occur simultaneously.
A reaction in which oxidation and reduction occur simultaneously.
Electrons lost by one species must be gained by another. Therefore oxidation cannot occur independently of reduction in a complete redox process.
Oxidation = loss of electrons = increase in oxidation number
Reduction = gain of electrons = decrease in oxidation number
Diagram 1: Oxidation and reduction are coupled processes
2. Classical Concepts of Oxidation and Reduction
Before the electronic definition became standard, oxidation and reduction were described in terms of oxygen and hydrogen transfer.
| Concept | Oxidation | Reduction |
|---|---|---|
| Oxygen | Addition of oxygen | Removal of oxygen |
| Hydrogen | Removal of hydrogen | Addition of hydrogen |
Example: Oxidation by Addition of Oxygen
2Mg + O₂ → 2MgOExample: Reduction by Removal of Oxygen
CuO + H₂ → Cu + H₂OWhy the electronic concept is better
Many redox reactions involve no oxygen or hydrogen at all. Electron transfer and oxidation-number changes provide a more general definition.
Many redox reactions involve no oxygen or hydrogen at all. Electron transfer and oxidation-number changes provide a more general definition.
3. Electronic Concept of Oxidation and Reduction
Oxidation
Loss of one or more electrons by an atom, ion or species.
Loss of one or more electrons by an atom, ion or species.
Reduction
Gain of one or more electrons by an atom, ion or species.
Gain of one or more electrons by an atom, ion or species.
Example
Zn → Zn²⁺ + 2e⁻Zinc loses electrons, so zinc is oxidized.
Cu²⁺ + 2e⁻ → CuCopper(II) ion gains electrons, so Cu²⁺ is reduced.
Combining the half-reactions:
Zn + Cu²⁺ → Zn²⁺ + CuMemory device
OIL RIG: Oxidation Is Loss; Reduction Is Gain — of electrons.
OIL RIG: Oxidation Is Loss; Reduction Is Gain — of electrons.
Diagram 2: Electron transfer between zinc and copper(II)
4. Oxidizing and Reducing Agents
Oxidizing agent / oxidant
A species that causes another species to be oxidized. The oxidizing agent itself is reduced.
A species that causes another species to be oxidized. The oxidizing agent itself is reduced.
Reducing agent / reductant
A species that causes another species to be reduced. The reducing agent itself is oxidized.
A species that causes another species to be reduced. The reducing agent itself is oxidized.
For:
Zn + Cu²⁺ → Zn²⁺ + Cu- Zn is oxidized → Zn is the reducing agent.
- Cu²⁺ is reduced → Cu²⁺ is the oxidizing agent.
High-yield rule
The agent is named by what it does to the other species. Oxidizing agent gets reduced; reducing agent gets oxidized.
The agent is named by what it does to the other species. Oxidizing agent gets reduced; reducing agent gets oxidized.
5. Oxidation Number / Oxidation State
Oxidation number
A formal number assigned to an atom in a compound or ion according to agreed electron-bookkeeping rules, representing the charge the atom would have under an idealized ionic assignment of bonding electrons.
A formal number assigned to an atom in a compound or ion according to agreed electron-bookkeeping rules, representing the charge the atom would have under an idealized ionic assignment of bonding electrons.
Oxidation number helps identify electron redistribution even in covalent reactions where actual complete electron transfer may not occur.
Oxidation → oxidation number increases
Reduction → oxidation number decreases
6. Rules for Assigning Oxidation Number
| Rule | Oxidation number | Example / exception |
|---|---|---|
| Free element | 0 | Na, O₂, Cl₂, S₈ |
| Monatomic ion | Equal to ion charge | Fe³⁺ = +3, S²⁻ = −2 |
| Group 1 metals | +1 | Na in NaCl = +1 |
| Group 2 metals | +2 | Mg in MgO = +2 |
| Fluorine | −1 in compounds | Highest electronegativity |
| Oxygen | Usually −2 | Peroxide: −1; superoxide: −1/2; OF₂: O = +2 |
| Hydrogen | Usually +1 | Metal hydrides such as NaH: H = −1 |
| Cl, Br, I | Usually −1 | Can be positive with O or F |
| Neutral compound | Sum = 0 | H₂SO₄ total = 0 |
| Polyatomic ion | Sum = ion charge | SO₄²⁻ total = −2 |
Do not memorize oxygen = −2 without exceptions
Peroxides, superoxides and oxygen fluorides are standard exceptions.
Peroxides, superoxides and oxygen fluorides are standard exceptions.
Diagram 3: Systematic oxidation-number assignment
7. Calculating Oxidation Numbers
Example 1: Sulfur in H₂SO₄
Let oxidation number of S = x.
2(+1) + x + 4(−2) = 0 2 + x − 8 = 0 x = +6
Example 2: Mn in KMnO₄
(+1) + x + 4(−2) = 0
x = +7
Example 3: Cr in Cr₂O₇²⁻
2x + 7(−2) = −2
2x − 14 = −2
x = +6
Example 4: N in NH₄⁺
x + 4(+1) = +1
x = −3
Example 5: Oxygen in H₂O₂
Hydrogen peroxide is a peroxide, so oxygen is −1, not −2.
Example 6: Oxygen in OF₂
Fluorine is always −1 in ordinary compounds:
x + 2(−1) = 0 O = +28. Oxidation Number vs Valency
| Feature | Oxidation Number | Valency |
|---|---|---|
| Meaning | Formal electron-bookkeeping number | Combining capacity in a bonding context |
| Sign | May be positive, negative or zero | Usually expressed without +/− sign in traditional school use |
| Can be zero? | Yes | Generally used as a positive combining capacity |
| Can vary in same element? | Yes | Yes for variable-valency elements |
| Fractional average possible? | Yes in some mixed-valence/average assignments | Usually whole number in introductory treatment |
9. Identifying Oxidation, Reduction and Redox Agents
Consider:
2FeCl₂ + Cl₂ → 2FeCl₃- Fe changes from +2 to +3 → oxidation.
- Cl in Cl₂ changes from 0 to −1 → reduction.
- FeCl₂ contains the species oxidized → acts as reducing agent.
- Cl₂ is reduced → acts as oxidizing agent.
Increase in ON = oxidation
Decrease in ON = reduction
Diagram 4: Oxidation-number change identifies the redox process
10. Common Types of Redox Reactions
| Type | General idea | Example |
|---|---|---|
| Combination | Substances combine and oxidation states change | 2Mg + O₂ → 2MgO |
| Decomposition | One reactant gives products with redox changes | 2KClO₃ → 2KCl + 3O₂ |
| Displacement | One element displaces another | Zn + CuSO₄ → ZnSO₄ + Cu |
| Disproportionation | Same element is simultaneously oxidized and reduced | 2H₂O₂ → 2H₂O + O₂ |
Disproportionation
In H₂O₂, O has oxidation number −1. It becomes −2 in H₂O and 0 in O₂, so oxygen is both reduced and oxidized.
In H₂O₂, O has oxidation number −1. It becomes −2 in H₂O and 0 in O₂, so oxygen is both reduced and oxidized.
11. Balancing Redox Reactions by Oxidation-Number Method
General Steps
- Write the skeletal equation.
- Assign oxidation numbers to elements that change.
- Find the increase and decrease in oxidation number.
- Choose coefficients so total increase = total decrease.
- Balance remaining atoms.
- Balance O with H₂O and H with H⁺ when working in acidic medium, if needed.
- Check both atoms and net charge.
Example: Fe²⁺ + Cr₂O₇²⁻ → Fe³⁺ + Cr³⁺ in Acidic Medium
Oxidation changes:
- Fe: +2 → +3, increase = 1 per Fe.
- Cr: +6 → +3, decrease = 3 per Cr; two Cr atoms give total decrease 6.
Therefore 6 Fe²⁺ are required per Cr₂O₇²⁻:
6Fe²⁺ + Cr₂O₇²⁻ → 6Fe³⁺ + 2Cr³⁺Balance oxygen with water:
6Fe²⁺ + Cr₂O₇²⁻ → 6Fe³⁺ + 2Cr³⁺ + 7H₂OBalance hydrogen with H⁺:
14H⁺ + 6Fe²⁺ + Cr₂O₇²⁻ → 6Fe³⁺ + 2Cr³⁺ + 7H₂OFinal check
Atoms balance and total charge is +24 on each side.
Atoms balance and total charge is +24 on each side.
Diagram 5: Oxidation-number balancing method
12. Ion-Electron / Half-Reaction Method in Acidic Medium
General Steps
- Split the equation into oxidation and reduction half-reactions.
- Balance all atoms except H and O.
- Balance O atoms using H₂O.
- Balance H atoms using H⁺.
- Balance charge using electrons.
- Multiply half-reactions to equalize electrons.
- Add and cancel common species.
Example: MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺
Reduction half:
MnO₄⁻ → Mn²⁺ MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂OOxidation half:
Fe²⁺ → Fe³⁺ + e⁻Multiply the Fe half-reaction by 5 and add:
MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺
13. Half-Reaction Method in Basic Medium
A reliable method is first to balance the equation as if it were acidic, then neutralize every H⁺ by adding an equal number of OH⁻ ions to both sides.
Procedure after Acidic Balancing
- Add OH⁻ to both sides equal to the number of H⁺ present.
- Combine H⁺ + OH⁻ into H₂O.
- Cancel water molecules appearing on both sides.
- Check atoms and charge.
Example Half-Reaction: MnO₄⁻ → MnO₂ in Basic Medium
Acidic-form balance:
MnO₄⁻ + 4H⁺ + 3e⁻ → MnO₂ + 2H₂OAdd 4OH⁻ to both sides:
MnO₄⁻ + 4H₂O + 3e⁻ → MnO₂ + 2H₂O + 4OH⁻Cancel 2H₂O:
MnO₄⁻ + 2H₂O + 3e⁻ → MnO₂ + 4OH⁻
Do not insert OH⁻ randomly
Balance systematically. The acidic-first conversion is less error-prone than guessing coefficients directly in basic medium.
Balance systematically. The acidic-first conversion is less error-prone than guessing coefficients directly in basic medium.
Diagram 6: Half-reaction method in acidic and basic media
14. Electrolysis: Qualitative Concept
Electrolysis
The use of electrical energy to drive a non-spontaneous chemical change in an electrolyte.
The use of electrical energy to drive a non-spontaneous chemical change in an electrolyte.
Electrolyte
A molten ionic substance or solution containing mobile ions that can conduct electric current through ionic motion.
A molten ionic substance or solution containing mobile ions that can conduct electric current through ionic motion.
Electrolytic Cell
- Anode: electrode where oxidation occurs.
- Cathode: electrode where reduction occurs.
- In an electrolytic cell, the anode is connected to the positive terminal and the cathode to the negative terminal of the external power supply.
- Cations migrate toward the cathode.
- Anions migrate toward the anode.
Universal electrode rule
AnOx, RedCat: oxidation always occurs at the anode; reduction always occurs at the cathode.
AnOx, RedCat: oxidation always occurs at the anode; reduction always occurs at the cathode.
Diagram 7: Ion migration and electrode reactions during electrolysis
15. Faraday’s First Law of Electrolysis
Faraday’s first law
The mass of a substance deposited, liberated or consumed at an electrode is directly proportional to the quantity of electric charge passed through the electrolyte.
The mass of a substance deposited, liberated or consumed at an electrode is directly proportional to the quantity of electric charge passed through the electrolyte.
m ∝ Q
Q = It
m = ZIt
where:
- m = mass deposited/liberated,
- Q = electric charge,
- I = current,
- t = time,
- Z = electrochemical equivalent of the substance.
Electrochemical equivalent, Z
The mass of substance deposited or liberated by one coulomb of charge under the specified electrode reaction.
The mass of substance deposited or liberated by one coulomb of charge under the specified electrode reaction.
16. Faraday’s Second Law of Electrolysis
Faraday’s second law
When the same quantity of electricity passes through different electrolytes, the masses of substances deposited or liberated are proportional to their chemical equivalent masses.
When the same quantity of electricity passes through different electrolytes, the masses of substances deposited or liberated are proportional to their chemical equivalent masses.
m₁/m₂ = E₁/E₂
where equivalent mass:
E = M/n
M is molar mass and n is the number of electrons involved per formula unit/ion in the relevant electrode reaction.
One Faraday
Approximately 96485 C of charge corresponds to one mole of electrons.
Approximately 96485 C of charge corresponds to one mole of electrons.
F ≈ 9.6485 × 10⁴ C mol⁻¹ e⁻
17. Quantitative Electrolysis Formulae
Suppose an ion is discharged by:
Mⁿ⁺ + ne⁻ → MOne mole of M requires nF coulombs. Therefore:
m = (MIt)/(nF)
Since E = M/n:
m = EIt/F
Z = E/F = M/(nF)
Moles of Electrons
n(e⁻) = Q/F = It/F
Time Required
t = mnF/(MI)
Diagram 8: Charge-to-mass relationship in electrolysis
18. Worked Examples and Numericals
Example 1: Identify Oxidation and Reduction
For:
Zn + CuSO₄ → ZnSO₄ + CuZn: 0 → +2 = oxidation. Cu: +2 → 0 = reduction.
Reducing agent: Zn. Oxidizing agent: Cu²⁺.
Example 2: Oxidation Number of Mn in MnO₄⁻
x + 4(−2) = −1
x = +7
Example 3: Oxidation Number of S in SO₃²⁻
x + 3(−2) = −2
x = +4
Example 4: Oxidation-Number Change
In:
2Fe²⁺ + Cl₂ → 2Fe³⁺ + 2Cl⁻Each Fe loses 1e⁻; two Fe atoms lose 2e⁻ total. Cl₂ gains 2e⁻ total.
Example 5: Charge Passed
A current of 2.0 A flows for 30 min.
t = 30 × 60 = 1800 s.
Q = It = 2.0 × 1800 = 3600 C
Example 6: Mass of Copper Deposited
Find Cu deposited by 2.0 A for 30 min from Cu²⁺. Take M(Cu)=63.55 g mol⁻¹, n=2, F=96485 C mol⁻¹.
m = MIt/(nF) m = 63.55×2.0×1800 /(2×96485) m ≈ 1.19 g
Example 7: Silver Deposition
Ag⁺ + e⁻ → Ag. A charge of 965 C passes through AgNO₃ solution. M(Ag)=107.87 g mol⁻¹.
m = MQ/(nF) = 107.87×965 /(1×96485) m ≈ 1.08 g
Example 8: Moles of Electrons
Charge passed = 48242.5 C.
n(e⁻) = Q/F = 48242.5/96485 = 0.500 mol
Example 9: Faraday’s Second Law
The same charge passes through Ag⁺ and Cu²⁺ electrolytes.
Equivalent masses:
E(Ag)=107.87/1=107.87 E(Cu)=63.55/2=31.775 m(Ag)/m(Cu)=107.87/31.775≈3.39
Example 10: Time Required
How long must 1.5 A pass to deposit 1.00 g Cu from Cu²⁺?
t = mnF/(MI) t = 1.00×2×96485 /(63.55×1.5) ≈ 2024 s≈ 33.7 min.
19. Formula and Concept Sheet
| Topic | Key relation |
|---|---|
| Oxidation | Loss of e⁻; oxidation number increases |
| Reduction | Gain of e⁻; oxidation number decreases |
| Oxidizing agent | Gets reduced |
| Reducing agent | Gets oxidized |
| Charge | Q = It |
| Faraday first law | m = ZIt |
| Equivalent mass | E = M/n |
| Faraday second law | m₁/m₂ = E₁/E₂ for same Q |
| Faraday constant | F ≈ 96485 C mol⁻¹ e⁻ |
| Moles of electrons | n(e⁻) = It/F |
| Mass deposited | m = MIt/(nF) |
| Electrochemical equivalent | Z = M/(nF) = E/F |
20. Common Exam Mistakes
- Confusing oxidation with oxygen addition only; the electronic definition is more general.
- Writing oxidation as gain of electrons. Oxidation is electron loss.
- Calling the oxidizing agent the species that is oxidized. The oxidizing agent is reduced.
- Calling the reducing agent the species that is reduced. The reducing agent is oxidized.
- Assigning non-zero oxidation number to a free element such as O₂ or Cl₂.
- Always assigning oxygen −2 and ignoring peroxide/superoxide/OF₂ exceptions.
- Always assigning hydrogen +1 and ignoring metal hydrides.
- Forgetting the oxidation-number sum equals the overall ion charge for a polyatomic ion.
- Confusing oxidation number with actual ionic charge in covalent compounds.
- Confusing oxidation number and valency.
- Balancing atoms but not charge in an ionic redox equation.
- Failing to equalize electrons lost and gained.
- Using H⁺ directly in the final basic-medium equation without converting with OH⁻.
- Adding electrons to the wrong side of a half-reaction.
- Forgetting that oxidation occurs at the anode and reduction at the cathode.
- Confusing electrolytic-cell electrode signs with the universal oxidation/reduction rule.
- Using minutes directly in Q = It instead of converting time to seconds when I is in amperes.
- Using molar mass instead of equivalent mass without accounting for n-electron transfer.
- For Cu²⁺ deposition, taking n=1 instead of n=2.
- Using F as 96500 C total rather than approximately 96485 C per mole of electrons.
- Applying Faraday’s second law to different charges; its simple mass ratio applies when the same quantity of electricity passes.
21. Important Exam Questions
Very Short / Short Questions
- Define oxidation and reduction by the electronic concept.
- Explain oxidation and reduction in terms of oxidation-number change.
- Define redox reaction.
- Define oxidizing agent and reducing agent.
- Why do oxidation and reduction always occur together?
- Define oxidation number.
- State the rules for assigning oxidation number.
- Find oxidation numbers of S in H₂SO₄, Mn in KMnO₄, Cr in Cr₂O₇²⁻ and N in NH₄⁺.
- State important exceptions for oxygen and hydrogen oxidation numbers.
- Differentiate oxidation number and valency.
- Identify the oxidized/reduced species and agents in given reactions.
- What is disproportionation? Give an example.
- State the steps of the oxidation-number balancing method.
- State the steps of the ion-electron method in acidic medium.
- How is a redox equation balanced in basic medium?
- Define electrolysis.
- Define electrolyte, anode and cathode.
- State the electrode processes at anode and cathode.
- State Faraday’s first law of electrolysis.
- Define electrochemical equivalent.
- State Faraday’s second law of electrolysis.
- Define equivalent mass in an electrochemical reaction.
- What is one Faraday of charge?
- Write the equation m = MIt/(nF) and define every symbol.
Long / Descriptive Questions
- Explain the classical and electronic concepts of oxidation and reduction with examples.
- Explain the rules for assigning oxidation number with suitable exceptions.
- Balance a redox reaction by the oxidation-number method.
- Balance a redox reaction by the half-reaction method in acidic medium.
- Balance a redox reaction by the half-reaction method in basic medium.
- Explain the qualitative mechanism of electrolysis using an electrolytic cell.
- State and explain Faraday’s first and second laws of electrolysis.
- Derive m = MIt/(nF) from Faraday’s law.
Numerical Practice
- Calculate oxidation numbers in neutral compounds and polyatomic ions.
- Calculate current, charge or time using Q = It.
- Calculate mass deposited from current and time.
- Calculate moles of electrons from charge.
- Use Faraday’s second law to compare deposited masses.
- Find the current or time needed to deposit a specified mass.
Exam Strategy
Master Unit 6 in four blocks: electronic redox concept → oxidation-number rules → balancing methods → electrolysis/Faraday numericals. Always check both atoms and charge after balancing a redox equation.
Master Unit 6 in four blocks: electronic redox concept → oxidation-number rules → balancing methods → electrolysis/Faraday numericals. Always check both atoms and charge after balancing a redox equation.
22. One-Minute Revision
- Unit 6: Oxidation and Reduction — 5 teaching hours.
- Oxidation = loss of electrons.
- Reduction = gain of electrons.
- Oxidation number increases during oxidation.
- Oxidation number decreases during reduction.
- Oxidizing agent is reduced.
- Reducing agent is oxidized.
- Free elements have oxidation number 0.
- Monatomic-ion oxidation number equals ion charge.
- F = −1 in ordinary compounds.
- O is usually −2; peroxide O = −1.
- H is usually +1; metal-hydride H = −1.
- Neutral-compound oxidation-number sum = 0.
- Polyatomic-ion oxidation-number sum = ion charge.
- Equalize total oxidation-number increase and decrease when balancing.
- Half-reaction method balances mass and charge separately.
- In acidic medium use H₂O, H⁺ and e⁻ systematically.
- For basic medium, neutralize H⁺ with OH⁻ and simplify H₂O.
- Electrolysis uses electrical energy to drive chemical change.
- Oxidation occurs at anode.
- Reduction occurs at cathode.
- Cations move to cathode; anions move to anode.
- Charge Q = It.
- Faraday first law: m ∝ Q.
- m = ZIt.
- Equivalent mass E = M/n.
- Faraday second law: same Q gives masses proportional to equivalent masses.
- F ≈ 96485 C mol⁻¹ electrons.
- n(e⁻) = It/F.
- m = MIt/(nF).
23. Diagram Practice
- Oxidation–reduction electron-transfer concept.
- Zn/Cu²⁺ electron transfer.
- Oxidation-number rule map.
- Oxidation-number change diagram.
- Oxidation-number balancing workflow.
- Half-reaction balancing workflow.
- Electrolytic-cell diagram.
- Faraday current–time–mass relationship.
Source handling: The original Nepal eNotes Oxidation and Reduction PDF remains embedded above using the Google Drive file directly linked as embedded content from the source page. The typed section follows the current Grade 11 Chemistry Unit 6 syllabus and includes electronic redox concepts, oxidation numbers, balancing by oxidation-number and half-reaction methods, and qualitative/quantitative Faraday laws of electrolysis. It is designed as a searchable, responsive study companion and is not claimed to be a word-for-word transcription of the handwritten PDF.
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