Hydrocarbons
Hydrocarbons are compounds of carbon and hydrogen only—the primary components of fossil fuels, plastics, and synthetic fibers.
Start with the simplest version: this lesson is about Hydrocarbons. If you can explain the core idea to a friend using everyday language, examples, and one clear reason why it matters, you have moved from memorising to understanding.
Hydrocarbons are compounds of carbon and hydrogen only—the primary components of fossil fuels, plastics, and synthetic fibers. Understanding hydrocarbon structure, nomenclature, and reactions reveals how petroleum becomes fuel and materials, and why certain hydrocarbons are more useful than others for specific applications.
General Characteristics of Hydrocarbons
Hydrocarbons have distinctive properties stemming from carbon-hydrogen bonding:
Nonpolar: C-H bonds have similar electronegativities. Carbon and hydrogen are close in electronegativity (carbon 2.5, hydrogen 2.1), producing nonpolar molecules. Hydrocarbons don't dissolve in polar solvents like water but dissolve in nonpolar solvents like benzene.
Insolubility in water: Oil and water don't mix. Oil (hydrocarbon) floats because it's less dense and nonpolar. This property explains environmental concerns with oil spills—hydrocarbons don't readily dissolve into water, persisting in ecosystems.
Combustibility: Hydrocarbons burn in oxygen, releasing large amounts of energy. CH₄ + 2O₂ → CO₂ + 2H₂O + 890 kJ. This exothermic reaction is why natural gas and gasoline are fuels. Incomplete combustion (limited oxygen) produces carbon monoxide (toxic) and soot.
Relatively unreactive at room temperature (especially alkanes): C-C and C-H bonds are strong and nonpolar. Alkanes don't react with acids, bases, or oxidizing agents at room temperature. This stability makes them ideal for storage and transport as fuels. Alkenes and alkynes are more reactive—their multiple bonds are "reactive sites."
Alkanes: Saturated Hydrocarbons
Alkanes (CₙH₂ₙ₊₂ for acyclic) contain only single bonds—they're saturated with hydrogen. Each carbon bonds to four atoms (carbon or hydrogen); each hydrogen to one carbon. The smallest alkane is methane (CH₄); longer chains are ethane (C₂H₆), propane (C₃H₈), butane (C₄H₁₀), and so on.
Physical properties: Boiling points increase with carbon number—methane is gas, pentane is liquid, paraffin wax is solid. Longer chains have more van der Waals interactions between molecules, requiring more energy to separate. This is why natural gas (methane) is a gas at room temperature while crude oil (long alkanes) is liquid and paraffin is solid.
Nomenclature: Systematic IUPAC names use roots (meth-, eth-, prop-, but-, pent-, hex-...) with suffix -ane. Substituents are named as prefixes with position numbers. 2-methylbutane is a four-carbon chain with a methyl branch on carbon 2.
Isomerism: Alkanes exhibit structural isomerism. Butane (C₄H₁₀) is either n-butane (straight chain) or 2-methylpropane (branched). Pentane (C₅H₁₂) has three isomers. As carbon number increases, isomer count explodes—C₂₀ has over 300,000 isomers!
Reactions of alkanes: Limited at room temperature. Combustion is the primary reaction—burning in oxygen. Halogenation (reacting with Cl₂ or Br₂ under UV light or heat) substitutes hydrogens with halogens. CH₄ + Cl₂ (UV light) → CH₃Cl + HCl. Cracking (splitting long chains using heat and catalysts) converts fuel oil into more useful gasoline and gases. Isomerization rearranges bonds without changing molecular formula—straight-chain alkanes become branched (higher octane for better burning).
Alkenes: Unsaturated Hydrocarbons with Double Bonds
Alkenes (CₙH₂ₙ) contain one or more C=C double bonds. The simplest is ethene (C₂H₄), also called ethylene. Double bonds make alkenes much more reactive than alkanes—the "unshared" electrons in the double bond are targets for reactions.
Nomenclature: IUPAC names use suffix -ene. The position of the double bond is specified: prop-1-ene (double bond between carbons 1 and 2) versus prop-2-ene (though prop-1-ene and prop-2-ene are the same). The lowest number is assigned to the carbon starting the double bond.
Geometric isomerism: Alkenes exhibit cis-trans isomerism because double bonds prevent rotation. In but-2-ene, methyl groups can be on the same side (cis-but-2-ene) or opposite sides (trans-but-2-ene). Same atoms, different arrangement, different properties.
Reactions of alkenes:
Addition reactions are characteristic. The double bond breaks, and atoms add across it. Hydrogen can add: C₂H₄ + H₂ → C₂H₆ (ethene becomes ethane). Halogens add: C₂H₄ + Br₂ → C₂H₄Br₂ (decolorizes brown Br₂ solution—a test for unsaturation). Water adds: C₂H₄ + H₂O → C₂H₅OH (making ethanol industrially). Hydrogen halides add: C₂H₄ + HCl → C₂H₅Cl.
Polymerization links many alkene molecules. Under pressure with catalysts, ethene becomes polyethylene (thousands of ethene units linked). Propene becomes polypropylene. These plastics are ubiquitous—bags, containers, fibers.
Oxidation breaks the double bond. Mild oxidation (with KMnO₄) produces diols. Strong oxidation cleaves the molecule.
Alkynes: Unsaturated with Triple Bonds
Alkynes (CₙH₂ₙ₋₂) contain one or more C≡C triple bonds. The simplest is ethyne (C₂H₂), also called acetylene. Triple bonds are even more reactive than double bonds—more electrons available for reaction.
Nomenclature: Suffix -yne. Position specified like alkenes.
Reactions: Similar to alkenes but even more extensive. Hydrogen adds: C₂H₂ + 2H₂ → C₂H₆. Halogens add twice: C₂H₂ + 2Br₂ → C₂H₂Br₄. Water adds: C₂H₂ + H₂O → CH₃CHO (acetaldehyde). Multiple addition reactions can occur at each carbon.
Uses: Acetylene burns with oxygen producing extremely high temperatures—used for welding. Industrial synthesis uses acetylene as a starting material for many organic compounds.
Aromatic Hydrocarbons: Benzene and Derivatives
Benzene (C₆H₆) is the simplest aromatic hydrocarbon. It's a six-membered ring with alternating single and double bonds, but the electron distribution is delocalized—all C-C bonds are equivalent, intermediate between single and double. This delocalization creates exceptional stability (resonance stabilization).
Aromatic compounds contain benzene rings or similar structures. Benzene is remarkably unreactive—double bonds usually react readily, but benzene doesn't add hydrogen or halogens easily at room temperature. This stability comes from aromaticity: electrons delocalized around the ring resist disruption.
Nomenclature: Substituents are named as benzene derivatives. Toluene is methylbenzene (CH₃ bonded to benzene). Phenol is hydroxybenzene. For multiple substituents, positions are numbered 1,2,3 (ortho, meta, para).
Reactions: Aromatic compounds typically undergo substitution rather than addition. Chlorine, with a catalyst, substitutes for hydrogen: C₆H₆ + Cl₂ (catalyst) → C₆H₅Cl + HCl. Nitration (reaction with HNO₃) produces nitrobenzene. These substitution products are starting materials for explosives, dyes, and pharmaceuticals.
Fossil Fuels and Refining
Crude petroleum is a complex mixture of hydrocarbons. Fractional distillation separates components by boiling point. Petroleum gas (smallest alkanes) vaporizes first, then naphtha, gasoline (useful fuel), kerosene, diesel oil, and finally heavy fuel oil. Residue is bitumen.
Cracking converts large molecules into smaller, more useful ones. Thermal cracking (high temperature) breaks bonds. Catalytic cracking (temperature + catalyst) is more selective. Gasoline (C₅-C₁₂) is more valuable than heavy fractions (C₁₆⁺), so cracking maximizes gasoline yield.
Reforming rearranges straight-chain alkanes into branched isomers with higher octane ratings. Octane rating measures how evenly fuel burns—higher octane resists knocking (premature combustion) in engines.
Environmental Impact
Hydrocarbon combustion produces CO₂ (greenhouse gas) and, incompletely, CO and soot (pollutants). Extraction and transport of petroleum risk spills. Plastic (polymerized hydrocarbons) persists in environments, harming wildlife. Society faces balancing hydrocarbon benefits (energy, materials) against environmental costs.
Socratic Questions
- Alkanes don't react with acids, bases, or oxidizing agents at room temperature. Yet they burn in oxygen. Why is combustion special—what makes oxygen different?
- Alkene double bonds are "electron-rich." Why does having extra electrons make a bond more reactive to addition, not less reactive?
- Benzene has three double bonds in its structure, yet it doesn't react like alkenes (which readily add hydrogen). What explains benzene's exceptional stability?
- If we can crack petroleum (break large alkanes into smaller ones), why is it energetically favorable? Doesn't breaking bonds require energy input?
- In isomerization, straight-chain alkanes become branched with no change in molecular formula. Why do different structural arrangements have such different properties if atoms and bonds are identical?