Chapter 11 — Carbon and its Compounds: Long Answer Questions
1. Tetravalency
Carbon has 4 valence electrons (electronic configuration: 2, 4) and therefore forms 4 covalent bonds. This means each carbon atom can bond to:
- 4 other carbon atoms
- Hydrogen atoms (H has 1 valence electron)
- Oxygen (2 bonds), nitrogen (3 bonds), sulphur, halogens, etc.
The 4 bonds can be:
- 4 single bonds (sp³ hybridisation — tetrahedral geometry)
- 1 double bond + 2 single bonds (sp² hybridisation — trigonal planar)
- 1 triple bond + 1 single bond (sp hybridisation — linear)
2. Catenation
Catenation is the ability of carbon atoms to bond with other carbon atoms in a chain.
Carbon can form:
- Long straight chains: CH₃−CH₂−CH₂−CH₂−CH₃ (pentane)
- Branched chains: 2-methylbutane
- Rings: Benzene (C₆H₆), cyclohexane (C₆H₁₂)
- Double bonds: Ethene (CH₂=CH₂)
- Triple bonds: Ethyne (CH≡CH)
Why only carbon? Because C−C bonds are both: 1. Strong (bond energy = 347 kJ/mol) 2. Non-polar (same electronegativity on both sides → no dipole → not easily attacked by chemicals)
Other elements like silicon also show catenation, but Si−Si bonds are weaker and more easily hydrolysed.
Versatility of Carbon Bonding
| Bond type | Example | General formula |
|---|---|---|
| C−C only (single bonds) | Alkanes | CₙH₂ₙ₊₂ |
| C=C (double bond) | Alkenes | CₙH₂ₙ |
| C≡C (triple bond) | Alkynes | CₙH₂ₙ₋₂ |
| Benzene ring | Aromatic compounds | C₆H₆ and derivatives |
Alkanes (Saturated hydrocarbons — only single bonds)
General formula: CₙH₂ₙ₊₂
| n | Name | Formula | Boiling point |
|---|---|---|---|
| 1 | Methane | CH₄ | −161°C |
| 2 | Ethane | C₂H₆ | −89°C |
| 3 | Propane | C₃H₈ | −42°C |
| 4 | Butane | C₄H₁₀ | −1°C |
| 5 | Pentane | C₅H₁₂ | 36°C |
Chemical properties: Undergo substitution reactions (one H replaced by another atom, e.g., Cl in the presence of sunlight)
Alkenes (Unsaturated — one C=C double bond)
General formula: CₙH₂ₙ
| n | Name | Formula |
|---|---|---|
| 2 | Ethene | C₂H₄ |
| 3 | Propene | C₃H₆ |
| 4 | Butene | C₄H₈ |
Chemical properties: Undergo addition reactions (molecules add across the double bond)
- Ethene + Br₂ → CH₂Br−CH₂Br (bromine water turns colourless → test for unsaturation)
- Ethene + H₂ → Ethane (hydrogenation — used in margarine production)
Alkynes (Unsaturated — one C≡C triple bond)
General formula: CₙH₂ₙ₋₂
| n | Name | Formula |
|---|---|---|
| 2 | Ethyne (acetylene) | C₂H₂ |
| 3 | Propyne | C₃H₄ |
Ethyne burns with a very hot flame → used in oxy-acetylene welding torches.
Structure
Benzene is a planar, cyclic compound with 6 carbon atoms forming a regular hexagon, each with one hydrogen atom.
Historical: Kekulé proposed alternating single and double bonds (cyclohexatriene). But this was wrong — benzene's actual structure involves delocalised electrons.
Actual structure (resonance/delocalisation model):
- 6 carbon atoms in a ring, each forming 2 σ-bonds to adjacent carbons and 1 σ-bond to hydrogen
- Each carbon has 1 remaining p-orbital perpendicular to the ring plane
- All 6 p-orbitals overlap to form a delocalised π-electron cloud above and below the ring
- The 6 π-electrons are shared equally by all 6 carbons — all C−C bonds are identical (bond order = 1.5)
- Represented by a circle inside a hexagon
Properties
- Physical: Colourless liquid, bp = 80°C, immiscible with water, volatile, highly flammable
- Aromatic smell — hence "aromatic compounds"
- Carcinogenic — known to cause leukaemia
Unusual Stability (Aromaticity)
Benzene is much more stable than expected for a compound with three double bonds:
- Does not decolourise bromine water readily (unlike alkenes)
- Does not readily undergo addition reactions
- Prefers substitution reactions (one H replaced by another group) — the ring is preserved
Reason: The delocalised π-electrons lower the energy of benzene (resonance stabilisation energy = 150 kJ/mol). Breaking the delocalised system (addition reaction) would destroy this stability — so benzene avoids it.
Hückel's rule of aromaticity: A cyclic compound is aromatic if it has 4n+2 π-electrons (n = 0, 1, 2...). Benzene has 6 π-electrons (n=1) → aromatic.
Saponification
Saponification is the process of making soap by the alkaline hydrolysis of fats or oils with a strong base (NaOH for solid soap; KOH for liquid soap).
Reaction:
Example (simplified):
Where R = long hydrocarbon chain (e.g., C₁₇H₃₅ for sodium stearate)
Structure of Soap
Each soap molecule has:
- Hydrophilic (water-loving) head: The carboxylate group −COO⁻Na⁺ — ionic, attracted to water
- Hydrophobic (water-fearing) tail: The long hydrocarbon chain −CH₂...CH₃ — non-polar, repels water, attracted to oil/grease
How Soap Cleans (Micelle Formation)
1. When soap is added to dirty water:
2. This forms a micelle — a spherical cluster with grease trapped inside and ionic heads on the outside 3. The micelle is water-soluble (ionic exterior) → grease is lifted off the surface and washed away
- The hydrophobic tails of soap molecules surround grease/oil (like dissolves like — non-polar interacts with non-polar)
- The hydrophilic heads point outward towards water
Soap vs Detergent
| Feature | Soap | Detergent |
|---|---|---|
| Preparation | Saponification of natural fats/oils | Synthetic; from petroleum-derived raw materials |
| Performance in hard water | Ineffective — reacts with Ca²⁺/Mg²⁺ ions to form insoluble scum | Effective in hard water — does not form scum |
| Biodegradability | Biodegradable | Some are non-biodegradable (cause pollution) |
| Lathering | Good lather in soft water | Good lather in all water |
| pH | Slightly alkaline | Can be neutral, acidic, or alkaline |
| Examples | Sodium stearate, sodium palmitate | Sodium lauryl sulphate (SLS), sodium dodecylbenzene sulphonate |
| Application | Personal hygiene, washing | Laundry, dishwashing, industrial cleaning |
Hard water contains dissolved Ca²⁺ and Mg²⁺ ions. These react with soap to form insoluble calcium/magnesium stearate (scum):
This wastes soap and leaves a sticky residue. Detergents do not have this problem.
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