Group 16 Elements - Oxygen Family

Real-Life Connection: From Breath to Batteries

Every breath you take is 21% oxygen - the element that keeps you alive! Ozone (O₃) high in the atmosphere protects us from UV rays, but at ground level it’s a pollutant. Sulfur makes matches ignite, vulcanizes rubber for tires, and powers lead-acid car batteries as sulfuric acid. From the bleaching powder that purifies water to the SO₂ that preserves dried fruits, Group 16 elements are everywhere!

Group 16 Elements Overview

Members: Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po)

Common Names: Chalcogens (ore-forming elements, from Greek “chalcos” = copper)

Electronic Configuration Pattern

• General configuration: ns² np⁴
• Oxygen: [He] 2s² 2p⁴
• Sulfur: [Ne] 3s² 3p⁴
• Selenium: [Ar] 3d¹⁰ 4s² 4p⁴

Memory Trick - “Oh! So Sexy Teacher, Please!”: O, S, Se, Te, Po

Key Trends Down the Group

Memory Trick for Oxidation States: “Negative Two, Plus Two-Four-Six” - Group 16 shows -2, +2, +4, +6

Interactive Demo: Visualize Group 16 in the Periodic Table

Explore the oxygen family (chalcogens) and their position in the periodic table.

Exception Alert: Electron gain enthalpy of S is more negative than O (due to small size of O, electron-electron repulsion in small 2p orbital)

Dioxygen (O₂) - The Life Supporter

Occurrence

• 20.95% of atmosphere by volume
• 46% of Earth’s crust by mass (in silicates, oxides)
• 89% of water by mass
• Most abundant element in Earth’s crust

Preparation

Laboratory (Thermal decomposition):

2KClO₃ --MnO₂--> 2KCl + 3O₂ (heat)
2KMnO₄ --heat--> K₂MnO₄ + MnO₂ + O₂
2H₂O₂ --MnO₂--> 2H₂O + O₂

Industrial:

Fractional distillation of liquid air
(b.p. of O₂ = -183°C, N₂ = -196°C)
O₂ distills first (higher b.p.)

Memory Trick: “KKKH = Kill KClO₃, KMnO₄ & H₂O₂” - All three decompose to give O₂

Properties of Oxygen

Physical:

• Colorless, odorless gas
• Paramagnetic (two unpaired electrons in π* orbitals)
• Slightly soluble in water (supports aquatic life)
• Liquid O₂ is pale blue and magnetic

Chemical (strong oxidizing agent):

1. With metals (forms oxides):

4Na + O₂ → 2Na₂O (oxide)
2Na + O₂ → Na₂O₂ (peroxide, excess O₂)
K + O₂ → KO₂ (superoxide, excess O₂)
2Mg + O₂ → 2MgO
3Fe + 2O₂ → Fe₃O₄ (at high temp)

Oxide Formation Trend:

• Small metals (Li, Be, Mg, Al) → Normal oxides (M₂O)
• Medium metals (Na) → Peroxides (M₂O₂)
• Large metals (K, Rb, Cs) → Superoxides (MO₂)

Memory Trick: “Small, Medium, Large → Oxide, Peroxide, Superoxide”

2. With non-metals:

C + O₂ → CO₂
S + O₂ → SO₂
4P + 5O₂ → P₄O₁₀
N₂ + O₂ → 2NO (at high temp/lightning)

3. With compounds:

2H₂S + 3O₂ → 2SO₂ + 2H₂O
2CO + O₂ → 2CO₂
CH₄ + 2O₂ → CO₂ + 2H₂O (combustion)

Types of Oxides

Oxidation states:

• Normal oxide: O = -2
• Peroxide: O = -1
• Superoxide: O = -1/2

Ozone (O₃) - The Protector

Preparation

Laboratory (Silent electric discharge):

3O₂ --silent electric discharge--> 2O₃ (ΔH = +142 kJ/mol, endothermic)

In stratosphere (UV radiation):

O₂ --UV--> 2O
O + O₂ → O₃

Memory Trick: “SED = Silent Electric Discharge” for O₃ preparation

Structure of Ozone

Shape: Bent (V-shaped) Bond angle: 117° Hybridization: sp² (central O) Bond length: 128 pm (intermediate between O=O and O-O) Resonance: Two equivalent structures

    O⁻
   / \\
  O⁺  O  ↔  O  O⁺
              \\ /
               O⁻

Formal charges: +1 on central O, -1 on one terminal O

Properties of Ozone

Physical:

• Pale blue gas, pungent smell
• Diamagnetic (all electrons paired)
• More soluble in water than O₂
• Thermally unstable: 2O₃ → 3O₂ (exothermic)

Chemical (powerful oxidizing agent):

1. Oxidation of metals:

2Ag + O₃ → Ag₂O + O₂
2Hg + O₃ → 2HgO + O₂ (used to detect O₃)

2. Oxidation of non-metals:

PbS + 4O₃ → PbSO₄ + 4O₂ (tarnish removal)
2I⁻ + O₃ + H₂O → I₂ + O₂ + 2OH⁻

3. Oxidation of organic compounds:

C₂H₄ + O₃ → C₂H₄O₃ (ozonide, unstable)
C₂H₄O₃ + Zn/H₂O → 2HCHO + ZnO (ozonolysis)

Ozonolysis: Used to determine position of double bonds in alkenes

4. Bleaching action:

Coloring matter + O₃ → Oxidized (colorless) + O₂
Permanent bleaching (oxidation, not reversible)

Comparison: Cl₂ + H₂O → HCl + HOCl (reversible bleaching) O₃ directly oxidizes (permanent)

Memory Trick: “O₃ Oxidizes Permanently, Cl₂ Chlorinates Reversibly”

Uses of Ozone

• Water purification (kills bacteria)
• Bleaching agent (oils, waxes, paper)
• Ozonolysis in organic chemistry
• Air purification
• Stratospheric O₃ layer protects from UV radiation

Ozone Layer Depletion

Cause: CFCs (Chlorofluorocarbons) - CCl₂F₂, CFCl₃

Mechanism:

CFCl₃ --UV--> CFCl₂ + Cl•
Cl• + O₃ → ClO• + O₂
ClO• + O → Cl• + O₂
(Cl• acts as catalyst, regenerated)

Result: Ozone hole over Antarctica Solution: Montreal Protocol (1987) - banned CFCs

Sulfur - The Brimstone

Allotropes of Sulfur

1. Rhombic Sulfur (α-S)

Structure: S₈ ring (puckered crown shape) Color: Yellow Stable below: 369 K Density: 2.06 g/cm³ Solubility: Soluble in CS₂

Preparation:

Evaporation of roll sulfur solution in CS₂

2. Monoclinic Sulfur (β-S)

Structure: S₈ ring (different crystal packing) Color: Pale yellow (needle-shaped crystals) Stable between: 369 K - 392 K Density: 1.98 g/cm³ Solubility: Soluble in CS₂

Preparation:

Melting rhombic S and cooling slowly (below 392 K)

Conversion:

Rhombic S ⇌ Monoclinic S (369 K = transition temperature)
     (α)           (β)

Above 392 K: Monoclinic melts to form yellow liquid (S₈ rings mobile) Above 433 K: Viscosity increases (rings break, form chains) Above 720 K: Viscosity decreases (chains break, S₂ molecules)

Memory Trick: “369 = 3×3×41 = Rhombic to Monoclinic”

Preparation of Sulfur

Occurrence: Native sulfur, H₂S in natural gas, sulfides (FeS₂, PbS), sulfates (CaSO₄)

Extraction:

1. Frasch process (underground sulfur):
   Superheated water (at 170°C, 10 atm) melts sulfur
   Compressed air brings molten S to surface
   99.5% pure sulfur obtained

2. From H₂S (Claus process):
   2H₂S + 3O₂ → 2SO₂ + 2H₂O
   2H₂S + SO₂ → 3S + 2H₂O

Properties and Reactions of Sulfur

With metals:

Fe + S → FeS
Hg + S → HgS (black, cinnabar when red)

With oxygen:

S + O₂ → SO₂

With halogens:

S + Cl₂ → S₂Cl₂ (chlorine deficient)
S₂Cl₂ + Cl₂ → 2SCl₂
S + 3F₂ → SF₆ (inert gas, excellent insulator)

With acids:

S + 2H₂SO₄(conc) → 3SO₂ + 2H₂O
S + 6HNO₃(conc) → H₂SO₄ + 6NO₂ + 2H₂O

Hydrogen Sulfide (H₂S)

Preparation

Laboratory (Kipp’s apparatus):

FeS + 2HCl → FeCl₂ + H₂S
FeS + H₂SO₄(dil) → FeSO₄ + H₂S

Industrial:

H₂ + S --673K--> H₂S

Memory Trick: “Fool’s gold (FeS₂) produces H₂S with acid”

Properties of H₂S

Physical:

• Colorless gas
• Rotten egg smell
• Poisonous (more toxic than HCN!)
• Soluble in water (weakly acidic)

Structure:

• Bent shape (like H₂O)
• Bond angle: 92° (less than H₂O’s 104.5°)
• Less hydrogen bonding than H₂O (S less electronegative)

Why is H₂S less acidic than H₂O? Actually, H₂S is MORE acidic than H₂O!

• H₂S → H⁺ + HS⁻ (Ka₁ = 10⁻⁷)
• H₂O → H⁺ + OH⁻ (Kw = 10⁻¹⁴)
• Larger S-H bond, easier to break
• Less electronegative S, can release H⁺ more easily

Acidic trends: H₂O < H₂S < H₂Se < H₂Te (acidity increases)

Memory Trick: “Down the group, Acidity DOwn-ward increases” (paradox name!)

Chemical Reactions of H₂S

Acidic nature:

H₂S + 2NaOH → Na₂S + 2H₂O (excess base)
H₂S + NaOH → NaHS + H₂O (limited base)

Reducing agent (S goes from -2 to higher):

H₂S + Cl₂ → 2HCl + S (turbidity)
H₂S + Br₂ → 2HBr + S
2H₂S + SO₂ → 3S + 2H₂O
H₂S + H₂SO₄(conc) → S + SO₂ + 2H₂O

Precipitation reactions (qualitative analysis):

H₂S + Pb(NO₃)₂ → PbS (black ppt) + 2HNO₃
H₂S + CuSO₄ → CuS (black ppt) + H₂SO₄
H₂S + CdCl₂ → CdS (yellow ppt) + 2HCl
H₂S + ZnSO₄ (+ NH₄OH) → ZnS (white ppt) + H₂SO₄

Combustion:

2H₂S + 3O₂ → 2SO₂ + 2H₂O (complete combustion)
2H₂S + O₂ → 2S + 2H₂O (limited O₂)

Oxides of Sulfur

Sulfur Dioxide (SO₂)

Preparation:

Laboratory:
Na₂SO₃ + H₂SO₄(dil) → Na₂SO₄ + H₂O + SO₂
Cu + 2H₂SO₄(conc) → CuSO₄ + 2H₂O + SO₂

Industrial:
S + O₂ → SO₂
4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂ (roasting of pyrites)

Structure:

• Bent shape
• Bond angle: 119°
• sp² hybridized S
• Resonance structures

Properties:

• Colorless gas, pungent smell
• Poisonous, acidic
• Soluble in water
• Reducing and oxidizing agent (amphoteric redox)

Reactions:

1. Acidic nature:

SO₂ + H₂O → H₂SO₃ (sulfurous acid)
SO₂ + 2NaOH → Na₂SO₃ + H₂O
SO₂ + NaOH → NaHSO₃ (excess SO₂)

2. Oxidizing agent (with strong reducing agents):

SO₂ + 2H₂S → 3S + 2H₂O

3. Reducing agent (more common):

SO₂ + 2H₂O + Cl₂ → H₂SO₄ + 2HCl
SO₂ + Br₂ + 2H₂O → H₂SO₄ + 2HBr
2SO₂ + O₂ + 2H₂O → 2H₂SO₄
SO₂ + [O] → SO₃

4. Bleaching action (temporary, reduction):

SO₂ + H₂O → H₂SO₃
H₂SO₃ + Coloring matter → Colorless (reduced form)
On exposure to air: Colorless + O₂ → Colored (re-oxidized)

Comparison: O₃ bleaches permanently (oxidation), SO₂ bleaches temporarily (reduction)

Uses:

• H₂SO₄ manufacture (Contact process)
• Bleaching agent (wool, silk, straw)
• Disinfectant, preservative
• Reducing agent

Sulfur Trioxide (SO₃)

Preparation:

2SO₂ + O₂ ⇌ 2SO₃ (Contact process: 450°C, 2 atm, V₂O₅ catalyst)

Structure:

• Planar triangular
• sp² hybridized S (but can expand octet using d-orbitals)
• Three S=O double bonds
• Bond angle: 120°

Forms:

• Gas: Individual SO₃ molecules
• Solid: Polymeric (asbestos-like or ice-like structures)

Properties:

• Colorless gas/liquid
• Highly reactive with water

Reactions:

SO₃ + H₂O → H₂SO₄ (very exothermic, violent)
SO₃ + NaOH → NaHSO₄
SO₃ + 2NaOH → Na₂SO₄ + H₂O

Sulfuric Acid (H₂SO₄) - King of Chemicals

Preparation: Contact Process

Steps:

1. Burning of sulfur:

S + O₂ → SO₂
or 4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂

2. Oxidation to SO₃ (key step):

2SO₂ + O₂ ⇌ 2SO₃
Conditions: 450°C, 2 atm pressure, V₂O₅ catalyst

3. Absorption (not direct with water!):

SO₃ + H₂SO₄(conc, 98%) → H₂S₂O₇ (oleum/fuming H₂SO₄)
H₂S₂O₇ + H₂O → 2H₂SO₄

Why not SO₃ + H₂O directly?

• Too exothermic, violent reaction
• Forms dense mist of H₂SO₄ droplets (difficult to condense)
• Oleum method is safer and efficient

Memory Trick: “Contact = Catalyst Oxidizes, Not Touches Aqua, Create Two sulfuric”

Properties of Sulfuric Acid

Physical:

• Colorless, odorless, oily liquid
• Highly viscous (due to extensive H-bonding)
• Boils at 611 K with decomposition
• Hygroscopic and dehydrating

Structure:

• Tetrahedral around S
• Two S=O and two S-OH bonds
• sp³ hybridized (but uses d-orbitals for π bonding)

Chemical Reactions

1. Acidic nature (dibasic):

H₂SO₄ + NaOH → NaHSO₄ + H₂O
H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O

With metals:
Mg + H₂SO₄(dil) → MgSO₄ + H₂

2. Dehydrating agent (removes H₂O):

C₁₂H₂₂O₁₁ --conc H₂SO₄--> 12C + 11H₂O (sugar to carbon)
HCOOH --conc H₂SO₄--> CO + H₂O
H₂C₂O₄ --conc H₂SO₄--> CO + CO₂ + H₂O
2HNO₃ + H₂SO₄ → 2NO₂⁺ + 2HSO₄⁻ + H₂O (nitration mixture)

3. Oxidizing agent (hot, concentrated):

With metals:

Cu + 2H₂SO₄(conc) --heat--> CuSO₄ + SO₂ + 2H₂O
Zn + 2H₂SO₄(conc) → ZnSO₄ + SO₂ + 2H₂O
3Ag + 2H₂SO₄(conc) → Ag₂SO₄ + SO₂ + 2H₂O

Exception: Cold conc. H₂SO₄ passivates Fe, Al (protective oxide layer)

With non-metals:

C + 2H₂SO₄(conc) → CO₂ + 2SO₂ + 2H₂O
S + 2H₂SO₄(conc) → 3SO₂ + 2H₂O

With compounds:

2HBr + H₂SO₄(conc) → Br₂ + SO₂ + 2H₂O
2HI + H₂SO₄(conc) → I₂ + SO₂ + 2H₂O
H₂S + H₂SO₄(conc) → S + SO₂ + 2H₂O

Why can’t we prepare HBr/HI using conc. H₂SO₄?

• H₂SO₄ oxidizes HBr to Br₂ and HI to I₂
• Only H₃PO₄ (non-oxidizing) can be used

Memory Trick: “CHORD = Can’t Handle Oxidizing Reactants, Dehydrates”

Uses of Sulfuric Acid

• Fertilizers (single largest use)
• Petroleum refining
• Paints and pigments
• Detergents
• Battery acid (car batteries)
• Chemical synthesis
• Metal processing

Comparison: Oxygen vs Sulfur

Memory Trick: “Oxygen Only does -2, Sulfur Shows Six” - O limited, S shows +6

Common Mistakes to Avoid

1. Mistake: O₂ is diamagnetic

   • Correct: O₂ is paramagnetic (two unpaired electrons in π*)

2. Mistake: Ozone bleaches by reduction

   • Correct: O₃ bleaches by oxidation (permanent), SO₂ by reduction (temporary)

3. Mistake: H₂S is less acidic than H₂O

   • Correct: H₂S is MORE acidic (Ka = 10⁻⁷ vs 10⁻¹⁴)

4. Mistake: SO₃ + H₂O directly in Contact process

   • Correct: SO₃ + H₂SO₄ → H₂S₂O₇, then H₂S₂O₇ + H₂O → 2H₂SO₄

5. Mistake: Conc. H₂SO₄ can prepare HBr/HI

   • Correct: Oxidizes to Br₂/I₂, use H₃PO₄ instead

6. Mistake: All metals react with conc. H₂SO₄

   • Correct: Fe, Al show passivation (cold conc. H₂SO₄)

7. Mistake: Rhombic and monoclinic S have different formulas

   • Correct: Both are S₈, differ only in crystal structure

Practice Problems

Level 1: JEE Main Basics

1. Why is O₂ paramagnetic while O₃ is diamagnetic?

2. Write balanced equations for: a) Laboratory preparation of O₂ from KClO₃ b) Preparation of H₂S from FeS c) Contact process for H₂SO₄

3. Arrange in increasing order of acidic strength: H₂O, H₂S, H₂Se, H₂Te

4. What is the oxidation state of S in: a) SO₂ b) SO₃ c) H₂SO₄ d) H₂S

5. Why is SO₂ bleaching action temporary while O₃ is permanent?

Level 2: JEE Main Advanced

6. Explain why conc. H₂SO₄ cannot be used to prepare HBr from NaBr. Which acid should be used?

7. Draw the structure of: a) O₃ b) SO₂ c) SO₃ d) H₂SO₄

8. Calculate the bond angle in: a) H₂O vs H₂S b) SO₂ vs SO₃ Explain the differences.

9. Why is SO₃ not directly dissolved in water in Contact process?

10. Conc. H₂SO₄ shows different behavior with cold and hot conditions. Explain with examples.

Level 3: JEE Advanced

11. Explain the following observations: a) H₂O is liquid while H₂S is gas at room temperature b) Electron gain enthalpy of S is more negative than O c) Rhombic sulfur converts to monoclinic at 369 K

12. In Contact process, the optimal temperature is 450°C despite the reaction being exothermic. Explain using Le Chatelier’s principle and reaction kinetics.

13. Complete and balance: a) Na₂O₂ + H₂O → b) KO₂ + H₂O → c) H₂S + SO₂ → d) Sugar + conc. H₂SO₄ →

14. When SO₂ is passed through acidified K₂Cr₂O₇ solution, the orange color changes to green. Write the balanced equation and explain the role of SO₂.

15. Calculate the oxidation number of S in: a) H₂S₂O₇ (oleum) b) H₂S₂O₈ (peroxodisulfuric acid) c) H₂S₂O₃ (thiosulfuric acid) d) Na₂S₄O₆ (sodium tetrathionate)

Cross-Links to Other Topics

Related to Periodic Classification

• Periodic Trends - Electron gain enthalpy anomaly
• Oxidation States - Variable oxidation of S

Related to Chemical Bonding

• Molecular Orbital Theory - O₂ paramagnetism
• Resonance - O₃, SO₂ structures
• Hybridization - SO₂ (sp²), SO₃ (sp²)

Related to Chemical Equilibrium

• Le Chatelier’s Principle - Contact process optimization

Related to Other Chapters

• Redox Reactions - SO₂ as reducing agent
• Qualitative Analysis - H₂S precipitation
• Environmental Chemistry - SO₂ pollution, ozone layer

Memory Palace for Group 16

Imagine an Industrial Complex:

Main Gate: Sign “Oh! So Sexy Teacher, Please!” (O, S, Se, Te, Po)

Oxygen Tower:

• Paramagnetic O₂ molecules spinning (two unpaired electrons)
• Three floors: Oxide (-2), Peroxide (-1), Superoxide (-1/2)
• Blue oxygen cylinder (liquid O₂ is blue)

Ozone Layer (on roof):

• Bent O₃ molecules at 117°
• UV shields protecting workers
• Pale blue gas warning signs
• Oxidizing power meter at maximum

Sulfur Factory:

• Yellow S₈ crown rings everywhere
• 369 K thermometer (rhombic to monoclinic transition)
• CS₂ storage tank (S solvent)

H₂S Warning Zone:

• Rotten egg smell detectors
• Black PbS, CuS, CdS samples (qualitative analysis display)
• Acidity increasing chart: H₂O < H₂S < H₂Se < H₂Te

Contact Process Plant (3 sections):

• Section 1: S burning (SO₂ yellow gas)
• Section 2: V₂O₅ catalyst chamber at 450°C, 2 atm (SO₃ production)
• Section 3: Oleum tanks (H₂S₂O₇), then H₂SO₄ storage

H₂SO₄ Applications Building:

• Floor 1: Dehydration chamber (sugar → carbon demo)
• Floor 2: Oxidation lab (Cu + H₂SO₄ → SO₂ brown fumes)
• Floor 3: Acid-base neutralization
• Basement: Battery acid storage

Quick Revision Checklist

• Group 16 configuration: ns² np⁴
• O₂ paramagnetic (two unpaired e⁻)
• Oxide types: M₂O (-2), M₂O₂ (-1), MO₂ (-1/2)
• O₃ preparation: Silent electric discharge
• O₃ structure: Bent, 117°, resonance
• O₃ bleaching: Permanent (oxidation)
• S allotropes: Rhombic ⇌ Monoclinic (369 K)
• H₂S: Rotten egg, acidic, reducing agent
• Acidity trend: H₂O < H₂S < H₂Se < H₂Te
• SO₂: Bent, 119°, amphoteric redox
• SO₂ bleaching: Temporary (reduction)
• Contact process: 2SO₂ + O₂ → 2SO₃ (450°C, 2 atm, V₂O₅)
• SO₃ + H₂SO₄ → H₂S₂O₇ (not direct H₂O)
• Conc. H₂SO₄: Dehydrating, oxidizing
• Can’t prepare HBr/HI with conc. H₂SO₄ (oxidizes)

Important Equations Summary

1. O₂ lab: 2KClO₃ --MnO₂--> 2KCl + 3O₂
2. O₃: 3O₂ --silent discharge--> 2O₃
3. Peroxide: Na₂O₂ + 2H₂O → 2NaOH + H₂O₂
4. Superoxide: 2KO₂ + 2H₂O → 2KOH + H₂O₂ + O₂
5. H₂S lab: FeS + 2HCl → FeCl₂ + H₂S
6. H₂S reducing: H₂S + Cl₂ → 2HCl + S
7. SO₂ lab: Na₂SO₃ + H₂SO₄ → Na₂SO₄ + H₂O + SO₂
8. SO₂ bleaching: SO₂ + H₂O → H₂SO₃ (reducing)
9. Contact 1: S + O₂ → SO₂
10. Contact 2: 2SO₂ + O₂ → 2SO₃ (450°C, 2 atm, V₂O₅)
11. Contact 3: SO₃ + H₂SO₄ → H₂S₂O₇
12. Contact 4: H₂S₂O₇ + H₂O → 2H₂SO₄
13. Dehydration: C₁₂H₂₂O₁₁ + H₂SO₄(conc) → 12C + 11H₂O
14. Oxidation: Cu + 2H₂SO₄(conc) → CuSO₄ + SO₂ + 2H₂O
15. H₂S + SO₂: 2H₂S + SO₂ → 3S + 2H₂O

Last updated: July 2025 Previous: Group 15 Elements | Next: Group 17 Elements