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OXIDATION AND REDUCTION
▢ Basic Terms:
Table 1: Oxidation-Reduction Terminology
Term | Change in oxidation number | Change in electrons |
|---|---|---|
Oxidation | Increase | Loss of electrons |
Reduction | Decrease | Gain of electrons |
Oxidising agent | Decrease | Accepts electrons |
Reducing agent | Increase | Donates electrons |
Oxidised substance | Increase | Loses electrons |
Reduced substance | Decrease | Gains electrons |
❖ Memory: LEO-GER = Loss of Electron → Oxidation; Gain of Electron → Reduction
▢ Redox Reaction:
❖ Definition: Reaction involving oxidation + reduction simultaneously
❖ Classification:
- •Redox reaction
- •Disproportionation reaction / auto-oxidation reaction
❖ Concept:
- •One reagent = oxidising agent → gets reduced
- •Another reagent = reducing agent → gets oxidised
- •Oxidant oxidises others and itself gets reduced
- •Reductant reduces others and itself gets oxidised
❖ Uses:
- •Metallurgy
- •Non-metals
- •Electrochemistry
- •Organic reactions
- •Energy production in body
❖ Biological Example: Energy in body mainly from oxidation of glucose
▢ Oxidation and Reduction Definitions:
Table 1: Electron Concept
Process | Meaning |
|---|---|
De-electronation | Oxidation |
Electronation | Reduction |
Reducing agent | Species de-electronated during chemical change |
Oxidising agent | Species electronated during chemical change |
❖ Electrolysis:
- •Electrolysis = redox reaction
- •Oxidation occurs at anode
- •Reduction occurs at cathode
❖ Lewis Concept:
- •Lewis acids = oxidants
- •Lewis bases = reductants
▢ Disproportionation Reaction:
❖ Definition: Same species acts as oxidising agent and reducing agent in same reaction
❖ Also Called: Auto-oxidation reaction
❖ Condition: Element must be in intermediate oxidation state
❖ Examples:
- •S, P, Cl, Br, I react with alkali → disproportionation
- •Fluorine does not show disproportionation due to strongest oxidising nature
- •H₂O₂ can undergo disproportionation
- •Bromine in alkaline medium can be both oxidised and reduced
Table 1: Lowest / Intermediate / Highest Oxidation State
Oxidation state | Behaviour |
|---|---|
Lowest O.S. | Reductant only |
Highest O.S. | Oxidant only |
Intermediate O.S. | Both oxidant and reductant |
Table 2: Nitrogen Example
Compound | O.S. of N | Behaviour |
|---|---|---|
NH₃ | −3 | Reductant only |
HNO₂ | +3 | Both oxidant and reductant |
HNO₃ | +5 | Oxidant only |
❖ Reaction Examples:
- •H₂O₂ + H₂O₂ → H₂O + O₂
- •O in H₂O₂: −1 → −2 = reduction
- •O in H₂O₂: −1 → 0 = oxidation
- •P₄ + NaOH + H₂O → NaH₂PO₂ + PH₃
- •P: 0 → +1 and 0 → −3
▢ Oxidation Number and Oxidation State:
Table 1: Meaning
Term | Meaning |
|---|---|
Oxidation number | Charge appearing by loss/gain/sharing of electrons in combined state during bond formation |
Oxidation state | Charge appearing per atom by loss/gain/sharing of electrons in combined state |
Valency | Combining capacity |
❖ Important Points:
- •Oxidation state and oxidation number are not necessarily equal
- •In electrovalent/ionic compounds: oxidation number, oxidation state and valency are same
- •Maximum O.N. of element ≈ number of valence electrons
- •Lowest O.N. of atom ≈ number of valence electrons − 8
❖ Examples:
- •MnO₄⁻: O.S. of Mn = +7
- •Cr₂O₇²⁻: O.S. of each Cr = +6; total O.N. of Cr atoms = +12
- •Pb₃O₄ = 2PbO·PbO₂ → Pb shows +2 and +4
- •Fe₃O₄ = FeO·Fe₂O₃ → Fe shows +2 and +3; average O.S. = +8/3
▢ Rules for Calculating Oxidation Number:
Table 1: General Rules
Rule | Oxidation number |
|---|---|
Free element | 0 |
Allotropes of element | 0 |
Ozone O₃ | O = 0 |
Fluorine | Always −1 |
Oxygen normally | −2 |
Oxygen in peroxides | −1 |
Oxygen in superoxides | −1/2 |
Oxygen in OF₂ | +2 |
Hydrogen normally | +1 |
Hydrogen in metallic hydrides | −1 |
Group IA | +1 |
Group IIA | +2 |
Ion | O.N. = electrical charge on ion |
Neutral molecule | Algebraic sum of O.N. of all atoms = 0 |
Polyatomic ion | Algebraic sum of O.N. = ionic charge |
❖ Special Notes:
- •In KO₃, O.S. of oxygen = −1/3
- •Superoxide ion is orange and paramagnetic
- •Ozonide ion is also coloured and paramagnetic
- •Hydride ion H⁻ has O.S. −1 and is isoelectronic with helium
- •H₂O₂ acts as oxidant, reductant and acid
▢ Oxidation State Trends:
Table 1: Periodic Trends
Property | Trend |
|---|---|
Oxidising power / tendency to gain electron / electronegative character | Decreases down group; increases left → right |
Reducing power / tendency to lose electron / electropositive character | Increases down group; decreases left → right |
Halogen oxidising property | F₂ > Cl₂ > Br₂ > I₂ |
Lithium | Strongest reducing agent in solution |
Caesium | Strongest reducing agent in air |
Manganese | Highest O.S. in 3d series = +7 |
Ruthenium and Osmium | Highest O.S. = +8 |
❖ Extra Points:
- •Starch-iodide paper → test for oxidising agent
- •HNO₂ shows oxidising, reducing and complex-forming properties
- •Oxygen shows positive O.N. only with fluorine, e.g. OF₂
- •Bleaching by chlorine is permanent due to oxidation
- •Bleaching by SO₂ is temporary due to reduction
▢ Acidic Nature and Oxidation State:
Table 1: Metal Oxides
Series | Nature |
|---|---|
FeO < Fe₂O₃ | Acidic nature increases with O.S. of metal |
CrO < Cr₂O₃ < CrO₃ | Basic < amphoteric < acidic |
Table 2: Non-metal Oxides / Oxyacids
Series | Nature |
|---|---|
SO₂ < SO₃ | Acidic nature increases with O.S. of non-metal |
CO < CO₂ | Neutral < acidic |
H₂N₂O₂ < HNO₂ < HNO₃ | Hyponitrous < nitrous < nitric acid |
HClO < HClO₂ < HClO₃ < HClO₄ | Hypochlorous < chlorous < chloric < perchloric acid |
❖ Thermal Stability: Thermal stability of oxyacids increases with increase in oxidation number of central element
▢ Equivalent Weight:
❖ Formula: Equivalent weight = Molecular weight / n-factor
Table 1: KMnO₄ Equivalent Weight in Different Media
Medium | Product | Change in O.N. of Mn | Equivalent weight |
|---|---|---|---|
Acidic | Mn²⁺ | +7 → +2 = 5 | M/5 = 158/5 = 31.6 |
Basic / alkaline | MnO₄²⁻ | +7 → +6 = 1 | M/1 = 158 |
Neutral | MnO₂ | +7 → +4 = 3 | M/3 = 158/3 = 52.6 |
Table 2: Common Equivalent Weights
Species | Equivalent weight |
|---|---|
K₂Cr₂O₇ | M/6 = 294/6 = 49 |
H₂O₂ | 34/2 = 17 |
Oxalic acid H₂C₂O₄ | Mol. wt./2 |
Na₂S₂O₃ in iodine reaction | Mol. wt./1 |
▢ Equivalent Weight of Few Oxidants and Reductants:
Table 1: n-factor Table
Species | Initial O.N. | Product | Final O.N. | Total change | Eq. wt. |
|---|---|---|---|---|---|
Cr₂O₇²⁻ | +6 | Cr³⁺ | +3 | 3 × 2 = 6 | Mol. wt./6 |
C₂O₄²⁻ | +3 | CO₂ | +4 | 1 × 2 = 2 | Mol. wt./2 |
S₂O₃²⁻ | +2 | S₄O₆²⁻ | +2.5 | 0.5 × 2 = 1 | Mol. wt./1 |
H₂O₂ as oxidant | −1 | H₂O | −2 | 1 × 2 = 2 | Mol. wt./2 |
H₂O₂ as reductant | −1 | O₂ | 0 | 1 × 2 = 2 | Mol. wt./2 |
I₂ | 0 | I⁻ | −1 | 1 × 2 = 2 | Mol. wt./2 |
▢ Important Oxidation States in Compounds:
Table 1: Common Oxidation States
Element | Compound | Formula | O.S. |
|---|---|---|---|
Cr | Potassium dichromate | K₂Cr₂O₇ | +6 |
Fe | Potassium ferrocyanide | K₄[Fe(CN)₆] | +2 |
Fe | Prussian blue / ferriferrocyanide | Fe₄[Fe(CN)₆]₃ | +2 |
Fe | Turnbull's blue / ferro ferricyanide | Fe₃[Fe(CN)₆]₂ | +3 |
C | Cane sugar | C₁₂H₂₂O₁₁ | 0 |
C | Glucose | C₆H₁₂O₆ | 0 |
C | Dichloromethane | CH₂Cl₂ | 0 |
C | Carbon tetrachloride | CCl₄ | +4 |
Fe | Brown ring | FeSO₄·NO | +2 |
Fe | Dark brown ring | [Fe(H₂O)₅NO]SO₄ | +1 |
S | Sodium tetrathionate | Na₂S₄O₆ | 0 for S-S and +5 |
S | Dimethyl sulphoxide | (CH₃)₂SO | 0 |
S | Perdisulphuric acid / Marshall acid | H₂S₂O₈ | +6, not +7 |
S | Hypo | Na₂S₂O₃ | +5 and −1 |
Cl | Bleaching powder | CaOCl₂ | +1 and −1 |
S | Peroxo monosulphuric acid / Caro's acid | H₂SO₅ | +6 |
S | Oleum | H₂S₂O₇ | +6 |
Co | Fischer's salt | K₃[Co(NO₂)₆] | +3 |
Cr | Chromium pentoxide | CrO₅ | +6 |
I | Potassium tri-iodide | KI₃ | 0 and −1 |
C and N | Cyanide / prussic acid | H-CN | +2 and −3 |
C and N | Isocyanide | H-N≡C | +2 and −3 |
N | Hydrazine | NH₂-NH₂ | −2 |
N | Nitrosyl chloride | NOCl | +3 |
Fe | Sodium nitroprusside | Na₂[Fe(CN)₅NO] | +2 |
Fe | Sodium sulphoniitroprusside | Na₄[Fe(CN)₅NOS] | +2 |
N | Hydrazoic acid | N₃H | −1/3 |
Ni | Nickel tetracarbonyl | Ni(CO)₄ | 0 |
Fe | Ferrous oxide | FeO | +2 |
Fe | Ferric oxide | Fe₂O₃ | +3 |
Fe | Ferrosoferric oxide | Fe₃O₄ | +8/3 average |
Cu | Schweitzer reagent | [Cu(NH₃)₄]SO₄ | +2 |
Fe | Haemoglobin | Hb | +2 |
Fe | Methaemoglobin | Hb | +3 |
▢ Special Compound Rules:
Table 1: Special Cases
Case | Point |
|---|---|
Metal carbonyls | O.N. of metal atom = 0 |
Examples of metal carbonyls | Ni(CO)₄, Fe(CO)₅, Cr(CO)₆, Mn₂(CO)₁₀ |
Alloys | O.N. of constituent metals = 0 |
Amalgam | O.N. of metal = 0 |
NaHg | O.N. of Na and Hg = 0 |
Carbohydrates with CH₂O empirical formula | Carbon O.S. = 0 |
Neutral ligands | O.S. = 0 |
Neutral ligand examples | C₂H₄, CO, NH₃, N₂ |
Fe in haemoglobin | +2 |
Fe in methaemoglobin | +3 |
▢ Only Oxidising and Reducing Agents:
Table 1: Classification
Type | Examples |
|---|---|
Only oxidising agents | H₂SO₄, SO₃, HNO₃, N₂O₅, P₂O₅, H₃PO₄, Cl₂O₇, HClO₄, CO₂, halogens, K₂Cr₂O₇, KMnO₄, FeCl₃, SnCl₄, HgCl₂, CuO |
Only reducing agents | HI, H₂S, N₂O, carbon, IA group elements, IIA group elements, H₃PO₂ |
Both oxidising and reducing agents | NO, N₂O₂, N₂O₄, N₂O₃, SO₂, H₂SO₃, HNO₂, CO, P₂O₃, H₃PO₃, H₂O₂, O₂, Na₂SO₃, NaNO₂, Na₂S₂O₃, HClO, HClO₂, Cl₂O₃, Cl₂O₅ |
❖ Hint: Check through maximum, intermediate and lowest oxidation state concept
▢ Read and Digest:
Table 1: Important Points
Fact | Answer |
|---|---|
SnCl₂ + HgCl₂ | HgCl₂ is reduced; precipitate forms |
NaCl dissolved in water | Na⁺ becomes hydrated, not hydrolysed |
3Br₂ + 6CO₃²⁻ + 3H₂O → 5Br⁻ + BrO₃⁻ + 6HCO₃⁻ | Bromine is both reduced and oxidised |
C₁₂H₂₂O₁₁ → CO₂ | Oxidation process |
Aluminothermic process | Aluminium acts as reducing agent |
Among HNO₂, H₂S, H₂SO₃, SnCl₂ | Strongest reducing agent = H₂S |
Reason for H₂S strongest reducing agent | Sulphur is in lowest oxidation state |
Carbon in CH₂O | O.N. = 0 |
Two Cl atoms in bleaching powder CaOCl₂ | +1 and −1 |
M³⁺ loses 3 electrons | Final O.N. = +6 |
P in Ba(H₂PO₂)₂ | +1 |
SO₂ + H₂S | SO₂ acts as oxidising agent; sulphur precipitates |
F₂ | Strongest oxidising agent; superhalogen |
Oxidation state of F | Always −1 |
F⁻ among F⁻, Cl⁻, Br⁻, I⁻ | Can bring highest oxidation state of transition metal |
K₂Cr₂O₇ + KOH | Orange dichromate changes to lemon-yellow chromate |
KMnO₄ | Acts as oxidising agent |
KMnO₄ + HCl | Reduction of Mn and production of Cl₂ |
K₂Cr₂O₇ reduced by 1 mole Sn²⁺ | Moles of K₂Cr₂O₇ = 1/3 |
Fe + SO₂ reaction | Fe is oxidised: 3Fe + SO₂ → 2FeO + FeS |
▢ High-Yield Recall:
Table 1: One-Liners
Clue | Answer |
|---|---|
Oxidation | O.N. ↑, electron loss |
Reduction | O.N. ↓, electron gain |
Oxidising agent | Accepts e⁻, gets reduced |
Reducing agent | Donates e⁻, gets oxidised |
Anode in electrolysis | Oxidation |
Cathode in electrolysis | Reduction |
LEO | Loss of Electron = Oxidation |
GER | Gain of Electron = Reduction |
Disproportionation | Same species oxidised + reduced |
F₂ | Strongest oxidising agent |
Li in solution | Strongest reducing agent |
Cs in air | Strongest reducing agent |
KMnO₄ acidic medium | n-factor = 5 |
KMnO₄ basic medium | n-factor = 1 |
KMnO₄ neutral medium | n-factor = 3 |
K₂Cr₂O₇ | n-factor = 6 |
H₂O₂ | n-factor = 2 |
O in peroxide | −1 |
O in superoxide | −1/2 |
O in ozonide | −1/3 |
O in OF₂ | +2 |
H in metallic hydride | −1 |
Mn highest O.S. in 3d series | +7 |
Ru/Os highest O.S. | +8 |
Chlorine bleaching | Permanent oxidation |
SO₂ bleaching | Temporary reduction |
Metal carbonyl metal O.N. | 0 |
Fe₃O₄ average O.S. of Fe | +8/3 |
Pb₃O₄ | Mixed oxide of PbO and PbO₂ |
HNO₂ | Oxidising + reducing + complex forming |
Starch iodide paper | Test for oxidising agent |
Q1.
How many electrons are needed to balance the equation: XO3− + H2O + e− → XH4 + OH−?
📅MOE 2001
Q2.
In the equation 2S2O3^2− → S4O6^2− + X, the value of X is
📅IOM 2005
Q3.
When SO2 is passed through acidified KMnO4 solution, the solution turns green because
📅MOE 2004
Q4.
In FeSO4 + KMnO4 + H2SO4 → Fe2(SO4)3 + K2SO4 + MnSO4 + H2O, which element is reduced?
📅MOE 2064
Q5.
A solution of SO2 in water reacts with H2S, precipitating sulphur. Here SO2 acts as
📅MOE 2058
Q6.
In the reaction Cr2O7^2− + H+ + I− → Cr3+ + H2O + I2, which one is reduced?
📅MOE 2008
Q7.
Which element is oxidized in the reaction 2CuS + 3O2 → 2CuO + 2SO2?
📅MOE 1997
Q8.
Oxidation state of oxygen atom in potassium superoxide KO2 is
📅BPKIHS 1994
Q9.
In the aluminothermic process, aluminium acts as
📅BPKIHS 2004
Q10.
Oxygen has positive oxidation state in
📅BPKIHS 2005
Q11.
The reaction 4P + 3NaOH + 3H2O → 3NaH2PO2 + PH3 is
📅BPKIHS 2000
Q12.
Oxidation can be defined as
Q13.
Reduction can be defined as
Q14.
A reducing agent is a substance which can
Q15.
An oxidizing agent is a substance which can
Q16.
In C + H2O → CO + H2, H2O acts as
Q17.
When NaCl is dissolved in water, the solution is
Q18.
In the bromine reaction producing bromide and bromate ions, bromine is
Q19.
One gas bleaches colour of flowers by reduction while the other by oxidation. These two gases are
Q20.
The oxidation state of carbon in phosgene gas COCl2 is
Q21.
In which of the following reactions is there no change in valency?
Q22.
Which of the following is the strongest oxidizing agent?
Q23.
The strongest reducing agent among the following is
📅MOE
Q24.
An element which never has positive oxidation number in any of its compounds is
Q25.
Oxidation state of chlorine in perchloric acid HClO4 is
Q26.
Oxidation number of carbon in CH3OH is
Q27.
Oxidation state of oxygen in hydrogen peroxide is
Q28.
The oxidation state of Cr in K2Cr2O7 is
Q29.
The oxidation number of Mn in KMnO4 is
Q30.
Phosphorus has oxidation state of +3 in
Q31.
Oxidation numbers of two chlorine atoms in bleaching powder CaOCl2 are
Q32.
Oxidation number of oxygen in F2O is
Q33.
The oxidation state of nitrogen in NH3 is
Q34.
Oxidation number of iron in potassium ferrocyanide K4[Fe(CN)6] is
Q35.
If three electrons are lost by a metal ion M3+, its final oxidation number would be
Q36.
Oxidation state of oxygen atom in potassium peroxide is
Q37.
The oxidation number of Ni in Ni(CO)4 is
Q38.
In MnO4− + SO3^2− + H+ → SO4^2− + Mn2+ + H2O
Q39.
Oxidation number of carbon in C6H12O6 is
Q40.
In coordination compound K4[Ni(CN)6], oxidation state of nickel is
Q41.
Oxidation number of S in S8, S2F2 and H2S respectively are
Q42.
When KMnO4 is reduced with oxalic acid in acid medium, oxidation number of Mn changes from
Q43.
The highest oxidation state of Mn is shown by
Q44.
The oxidation number of iodine in HIO3 is
Q45.
Oxidation number of phosphorus in Ba(H2PO2)2 is
Q46.
In haemoglobin, iron is present in