Organic Chemistry Fundamentals and Alkanes

Introduction to Organic Chemistry

• Definition: Organic chemistry is the study of the chemistry of carbon and its compounds.
• Exclusions from the Definition:
  • Sulphides of carbon
  • Oxides of carbon
  • Metallic carbonates
  • Cyanates
  • Hydrogen carbonates
• Elemental Composition: Organic compounds typically contain carbon and one or more of the following: oxygen, nitrogen, sulphur, phosphorus, and any of the halogens.
• Ubiquity and Importance: Millions of organic compounds exist, with many new ones discovered annually. Examples include:
  • Biological/Health: Proteins, carbohydrates, fats, oils, vitamins, hormones.
  • Industrial/Commercial: Perfumes, soaps, dyes, paints, textiles, paper.
  • Agricultural: Insecticides, herbicides, fungicides.

Historical Development of Organic Chemistry

• 18th and Early 19th Century Belief: Compounds were categorized by source.
  • Inorganic compounds: Obtained from minerals.
  • Organic compounds: Obtained from materials produced by living organisms (vegetable or animal sources).
• The Vital Force Theory:
  • Proposed By: Jacob Berzelius, a Swedish Chemist.
  • Core Concept: Scientists believed a "vital force" existed within organic materials that was absent in inorganic materials. They argued that organic substances could only be formed through a mysterious, God-given power within plants and animals.
• The Disproval of Vital Force Theory (1828):
  • Pioneer: Friedrich Wohler, a German Chemist.
  • The Experiment: Wohler synthesized urea (an organic compound found in animal urine) from inorganic starting materials.
  • Chemical Reaction: He mixed ammonium chloride and silver cyanate. While the expected product was ammonium cyanate (inorganic), the end result was urea.
  • Formula: $NH_4Cl + AgOCN \rightarrow NH_2OCN \rightarrow CO(NH_2)_2$
  • Significance: Urea was the first organic compound to be synthesized in a lab, proving that organic substances can be created without a biological "vital force."

The Uniqueness of Carbon

• Catenation: This is the unique property of carbon atoms to attach to one another to form chains (thousands of atoms long) or rings of varying sizes.
  • The chains/rings can have branches and crosslinks.
  • This ability is attributed to carbon's tetrahedral structure and its tetracovalency (covalent character).
• Bonding Versatility: Carbon can form multiple bonds with itself and other elements:
  • Self-bonding: $C=C$ (double bonds) and $C \equiv C$ (triple bonds).
  • With other elements: $C=O$ (oxygen), $C=S$ (sulphur), and $C \equiv N$ (nitrogen).
• Comparison with Silicon: Although Silicon ($Si$) can undergo catenation, silicon bonds are much shorter and less stable due to the large atomic size. Carbon bonds are strong and stable because the carbon atom is small.

Functional Groups and Homologous Series

• Functional Group Definition: The part of the organic molecule responsible for its chemical behavior (reactivity). The rest of the molecule is generally considered relatively inert.
• Homologous Series Definition: A family of related organic compounds that possess the same functional group and similar chemical structures.
• Characteristics of a Homologous Series:
  1. General Formula: Members share a common molecular formula (e.g., Alkenes = $C_nH_{2n}$).
  2. Chemical Behavior: Members exhibit similar chemical reactions.
  3. Preparation: Members can be prepared using similar synthetic methods.
  4. Molecular Mass Increment: Each member differs from the preceding compound by a $CH_2$ unit, which has a relative molecular mass of $14$.
  5. Gradation of Physical Properties: There is a regular change in physical state.
     • Example (Alkanes): $C_1$ to $C_4$ are gases; $C_5$ to $C_{17}$ are liquids (ranging from volatile to heavy like kerosene); $C_{18}$ and above are solids.

Classes of Organic Compounds

General Characteristics of Organic Compounds

1. Composition: Mainly $C, H, O$. Often contains halogens, $N, S$, and occasionally $P$.
2. Physical State: Usually gases, liquids, or low-melting volatile solids.
3. Solubility: Generally insoluble in water. Soluble in non-polar solvents (benzene, ether).
   • Exception: Compounds containing polar groups ($-OH, -COOH, -SO_3H$) are water-soluble.
   • Principle: "Like dissolves like" – sugar and alcohol dissolve in water due to the hydroxyl group ($-OH$) and hydrogen bonding.
4. Conductivity: They do not conduct electricity in solution or molten form because they exist as covalent molecules (not ions).
5. Reactivity: Reactions are generally slow, require energy (heat), and rarely go to completion. This necessitates careful purification of products.
6. Combustibility: Many are highly combustible, burning to yield carbon dioxide ($CO_2$) and water ($H_2O$).
7. Isomerism: This is common in organic chemistry. It is the phenomenon where compounds have the same molecular formula but different structural arrangements.

Electronic Configuration and Allotropy of Carbon

• Electronic Configuration: Carbon (Atomic Number 6) has the configuration $1s^2 2s^2 2p^2$.
• Allotropy: The ability of an element to exist in two or more different forms in the same physical state. Other elements exhibiting this include sulphur, phosphorus, and oxygen.
• Allotropes of Carbon:
  • Crystalline Forms: Atoms are in a highly ordered geometric crystal lattice (e.g., Diamond, Graphite, Fullerenes, Graphene).
  • Amorphous Forms: Atoms lack a defined crystalline structure and appear shapeless at the atomic level (e.g., Coal, Charcoal, Coke, Carbon Black/Lampblack, Gas carbon).

Alkanes (Paraffins)

• Introduction: Saturated hydrocarbons with the general formula $C_nH_{2n+2}$. They are called "paraffins" (meaning "little affinity") because they are relatively unreactive. Carbon atoms in alkanes are $sp^3$ hybridized (tetrahedral geometry).
• Synthetic Preparations:
  1. Hydrogenation of Alkenes: Passing a mixture of alkene and hydrogen over a catalyst.
     • Catalysts: Platinum ($Pt$) or Palladium ($Pd$) work at room temperature. Nickel ($Ni$) requires high temperature ($150^\circ C$) and pressure.
     • $C_nH_{2n} + H_2 \xrightarrow{Pt/Pd} C_nH_{2n+2}$
  2. Reduction of Haloalkanes: Using a zinc-copper ($Zn-Cu$) couple (copper-coated zinc) in aqueous alcohol at room temperature.
     • $2RX + Zn + 2H^+ \rightarrow 2RH + Zn + X^-$
  3. From Grignard Reagent ($RMgX$): Heating alkyl magnesium halide with aqueous acid ($HCl$ or $H_2SO_4$).
     • $RMgX + H_2O \rightarrow R-H + Mg(OH)X$
  4. Kolb Synthesis: Electrolysis of the sodium salt of a monocarboxylic acid in aqueous methanolic solution. It produces symmetrical alkanes ($R-R$) at the anode.
     • $2CH_3COO^- \rightarrow CH_3CH_3 + 2CO_2 + 2e^-$
  5. Wurtz Synthesis: Coupling two alkyl halides using metallic sodium in dry ether.
     • $2RX + 2Na \xrightarrow{\text{dry ether}} R-R + 2NaX$
  6. Method Specific to Methane: Decarboxylation of sodium ethanoate by fusing it with sodalime ($NaOH$ and $CaO$).
     • $CH_3COONa + NaOH \xrightarrow{sodalime} CH_4 + Na_2CO_3$

Reactions and Physical Properties of Alkanes

• Reactions:
  1. Combustion: Rapid exothermic oxidation producing $CO_2$ and steam. Burns with a non-smoky flame.
     • $C_nH_{2n+2} + \frac{3n+1}{2}O_2 \xrightarrow{heat} nCO_2 + (n+1)H_2O$
  2. Cracking: Breaking long-chain alkanes into smaller alkanes and alkenes.
     • $\text{Heptane } (C_7H_{16}) \rightarrow \text{Pentane } (C_5H_{12}) + \text{Ethene } (C_2H_4)$
  3. Halogenation: Substitution reaction occurring in UV light or at $250-400^\circ C$.
     • Methane + Chlorine yields a mixture: Chloromethane ($CH_3Cl$), Dichloromethane ($CH_2Cl_2$), Trichloromethane ($CHCl_3$), and Tetrachloromethane ($CCl_4$).
     • Reactivity: Fluorination is violent; Iodination does not occur. Reaction does not occur in the dark.
• Physical Properties:
  1. State: $C_1-C_4$ (gases), $C_5-C_{17}$ (liquids), $C_{18}+$ (solids) at $20^\circ C$.
  2. Solubility: Non-polar; immiscible in water, soluble in organic solvents (e.g., trichloromethane).
  3. Density: Less dense than water ($0.6-0.8\,g/cm^3$). In a mixture with water, the alkane floats on top. Density increases with molecular weight.
  4. Boiling Points: Increases with molecular weight due to increased surface area and van der Waals forces.
     • Branching effect: Branching decreases the boiling point (increases volatility) compared to straight-chain isomers because branching reduces surface area and intermolecular attractions.