Organic Chemistry – Introduction Notes

Introduction to Organic Chemistry

• Definition
  • Study of the chemistry of carbon compounds.
• Word “organic” carries distinct meanings
  • In agriculture: derived from living organisms (e.g.a0cow manure).
  • In food: grown without synthetic pesticides/fertilizers.
  • In chemistry: any compound whose chemistry is dominated by carbon.
• Core characteristics of organic compounds
  • Contain carbon; usually also hydrogen; frequently O, N, P, S, halogens.
  • Commonly have C–C or C–H frameworks.
  • Exclusions: $\text{CO}<em>2$, carbonates $\text{CO}</em>3^{2-}$, cyanides $\text{CN}^-$ are considered inorganic despite carbon.
  • Complex molecular architectures; basis of biochemistry (proteins, carbohydrates, lipids, nucleic acids) and pharmaceuticals (methotrexate, 5-fluorouracil, Tamiflu, AZT, morphine, etc.).

Organic vs. Inorganic Compounds

• Composition
  • Organic: C with H/O/N etc.; contain C–H bonds.
  • Inorganic: generally no C–H linkage; often metals, salts, minerals.
• Occurrence
  • Organic: chiefly in living systems.
  • Inorganic: Earth’s crust, minerals, atmospheric gases.
• Bonding
  • Organic: mostly covalent.
  • Inorganic: ionic, covalent, metallic combinations.
• Solubility
  • Organic: mostly non-polar → poor in water, good in organic solvents.
  • Inorganic: ionic salts readily dissolve in water.
• Physical state
  • Organic: many liquids/gases at RT; can be solids.
  • Inorganic: predominantly crystalline solids.
• Volatility & Flammability
  • Organic: generally volatile & flammable.
  • Inorganic: usually non-flammable.
• Melt/Boil Points
  • Organic: comparatively low.
  • Inorganic: higher due to ionic lattices.
• Reaction rates
  • Organic: slower (covalent reactivity, sterics).
  • Inorganic: faster (ionic pathways, catalysts).
• Colour
  • Organic: mostly colourless.
  • Inorganic: many coloured (transition-metal complexes).
• Typical examples
  • Organic: glucose, methane, acetone, lysine.
  • Inorganic: $\text{NaCl}$, $\text{CO}<em>2$, $\text{NH}</em>3$.

Why Carbon Is Unique

• Position in Periodic Table
  • Atomic number $6$, configuration $1s^22s^22p^2$.
  • Four valence electrons → tetravalency → forms four covalent bonds.
• Tetravalency
  • Allows extensive bonding networks; each bonding pair often depicted as a line.
• Bond strength
  • Single $\text{C–C}$ bond energy $\approx 347\;\text{kJ mol}^{-1}$ (stronger than $\text{N–N}$, $\text{O–O}$).
  • Supports long straight, branched, cyclic chains.
• Catenation
  • Self-linking ability leading to diverse skeletons: chains, branches, rings.
• Multiple bonding capability
  • Single, double, triple bonds with C/O/N; delocalized $\pi$ systems (benzene).

Representative Carbon Skeletons

• Straight-chain alkanes, branched alkanes, cycloalkanes.
• Examples of bonding motifs
  • $\text{C–C}, \text{C–H}, \text{C–O}, \text{C=O}, \text{C–N}, \text{C–X}$ (halogen).

Bonding & Hybridization in Carbon

• Review prerequisites: electrons, orbitals, electron configuration, VSEPR.
• Hybridization principles
  • Hybrid orbitals form only in bonded atoms, not isolated ones.
  • Produced by combining atomic orbitals; number conserved.
  • All hybrids in a set are equivalent.
  • Geometry predicted by VSEPR dictates type.
  • $\sigma$ bonds arise from overlap of hybrids; unhybridized p orbitals overlap to yield $\pi$ bonds.

sp³ Hybridization (Methane / Ethane)

• Combination of one 2s + three 2p orbitals → four sp³ hybrids.
• Geometry: tetrahedral ($109.5^{\circ}$).
• All $\sigma$ bonds; example $\text{CH}<em>4$, $\text{C}</em>2\text{H}_6$.

sp² Hybridization (Ethene / Benzene)

• Mix one 2s + two 2p → three sp² hybrids + one unhybridized p.
• Geometry: trigonal planar ($120^{\circ}$).
• $\sigma$ framework + one $\pi$ bond (double bond).
• Benzene: each C sp²; p-orbitals overlap forming delocalized 6-electron ring (fully conjugated).

sp Hybridization (Ethyne)

• Mix one 2s + one 2p → two sp hybrids + two unhybridized p.
• Geometry: linear ($180^{\circ}$).
• $\sigma$ bond + two perpendicular $\pi$ bonds (triple bond).

Hybridization & Electron-Pair Geometries (VSEPR)

• 2 regions → linear → sp → $180^{\circ}$.
• 3 regions → trigonal planar → sp² → $120^{\circ}$.
• 4 regions → tetrahedral → sp³ → $109.5^{\circ}$.
• Expanded octets: trigonal bipyramidal (sp³d), octahedral (sp³d²).

Sigma vs. Pi Bonds Review

• Single bond = 1 $\sigma$.
• Double = 1 $\sigma + 1 \pi$.
• Triple = 1 $\sigma + 2 \pi$.
• Example ethyne: $\sigma$ bonds $=3$ (C–H, C–H, C–C); $\pi$ bonds $=2$.

Formulas of Organic Compounds

• General formula: algebraic representation of a homologous series.
  • Alkanes $\mathrm{C<em>nH</em>{2n+2}}$.
  • Alkenes $\mathrm{C<em>nH</em>{2n}}$.
  • Alkynes $\mathrm{C<em>nH</em>{2n-2}}$.
  • Alkyl halide $\mathrm{C<em>nH</em>{2n+1}X}$.
  • Alcohol $\mathrm{C<em>nH</em>{2n+1}OH}$, etc.
• Empirical formula (EF)
  • Gives simplest integer ratio.
  • Example: ethane $\mathrm{C<em>2H</em>6}\rightarrow$ EF $=\mathrm{CH_3}$.
• Molecular formula (MF)
  • Actual number of atoms; multiple of EF.
  • Ethanoic acid: MF $=\mathrm{C<em>2H</em>4O_2}$.
  • Relation: $(\text{EF})^n = \text{MW}$.
• Structural formulas
  • Expanded: shows every bond.
  • Condensed: omits some bonds, groups atoms.
  • Skeletal/line-angle: lines for C–C; hydrogens on carbon omitted; vertices = carbons.
  • Kekulé: common for aromatics, depicts ring with alternating single/double lines.

Classification of Organic Compounds

• Open-chain (Aliphatic): terminal carbons not joined; subclasses
  • Alkanes, alkenes, alkynes.
• Closed-chain (Cyclic)
  • Homocyclic (all carbons)
  • Alicyclic: ring behaves like aliphatic (cycloalkanes).
  • Aromatic: benzene-like; obey Hückel $4n+2 \pi$-electron rule.
  • Heterocyclic: ring contains heteroatom (N, O, S, …).

Functional Groups & Homologous Series

• Functional group: specific atom/group conferring characteristic reactivity; remains chemically consistent across molecules.
• Homologous series: family of compounds with the same functional group and general formula; successive members differ by $\mathrm{CH_2}$ ($14\,\text{amu}$).

Key Properties of a Homologous Series

1. Represented by common formula (e.g. alkanes $\mathrm{C<em>nH</em>{2n+2}}$).
2. Consecutive members differ by $\mathrm{CH_2}$.
3. Prepared by a common synthetic method (e.g. dehydration of alcohols → alkenes).
4. Show gradual variation of physical properties but similar chemical reactivity (e.g. alkene + $\text{Br}_2$ addition).

Survey of Major Functional Groups (examples)

• Alkane – hydrogen atom (ethane $\text{C}<em>2\text{H}</em>6$).
• Alkene – C=C (ethene $\text{C}<em>2\text{H}</em>4$).
• Alkyne – C≡C (ethyne $\text{C}<em>2\text{H}</em>2$).
• Arene – benzene ring (methylbenzene).
• Haloalkane – R–X (bromoethane).
• Alcohol – R–OH (ethanol).
• Aldehyde – \ce{RCHO} (ethanal).
• Ketone – \ce{RCOR'} (propanone).
• Carboxylic acid – \ce{RCOOH} (ethanoic acid).
• Ester – \ce{RCOOR'} (methyl ethanoate).
• Acyl chloride – \ce{RCOCl} (ethanoyl chloride).
• Amide – \ce{RCONH2} (ethanamide).
• Amine – \ce{RNH2} (ethanamine).
• Nitrile – \ce{RCN} (ethanenitrile).

Alkyl Groups

• Formed by removing one H from an alkane; suffix changes from –ane → –yl.
  • Methane $\to$ methyl (\ce{-CH3}).
  • Propane $\to$ propyl (\ce{-CH2CH2CH3}).
• Generic symbol $R$ often denotes an unspecified alkyl group.
• Straight-chain alkyl general formula $\mathrm{C<em>nH</em>{2n+1}}$.
• Common branched alkyls
  • Isopropyl \ce{(CH3)2CH-}.
  • tert-Butyl \ce{(CH3)3C-}.

Classification by Degree of Substitution

• Primary (1°): central carbon/N attached to ONE other carbon.
• Secondary (2°): attached to TWO carbons.
• Tertiary (3°): attached to THREE carbons.
• Quaternary (4°): (for ammonium) nitrogen attached to FOUR carbons; for carbon, quaternary center is C attached to four C’s (no H).
• Applied to alcohols, alkyl halides, amines, amides.
  • Examples given: primary/secondary/tertiary alcohols & alkyl halides, ammonia vs 1°–4° amines, 1°–3° amides.

Practice/Activity Highlights

• Classify sample molecules into organic vs inorganic (e.g. morphine vs $\text{NaHCO}_3$).
• Count number of $\text{C–C}, \text{C–H}, \text{C–O}, \text{C=O}$ bonds in given structures.
• Identify carbon atoms without hydrogens in complex molecules (e.g. substituted benzenes).
• Build molecular models: $\text{CH}<em>4$ (tetrahedral), $\text{C}</em>2\text{H}<em>4$ (planar), $\text{C}</em>2\text{H}<em>2$ (linear), $\text{C}</em>6\text{H}_6$ (planar ring).
• Write condensed and line-angle formulas for given molecular formulas (e.g. $\text{C}<em>7\text{H}</em>{14}$, $\text{C}<em>9\text{H}</em>{18}$ isomers).

Summary & Key Take-Home Messages

• Carbon’s tetravalency, catenation, and strong C–C bonds underpin the vastness of organic chemistry.
• Organic compounds differ markedly from inorganic ones in bonding, properties, and occurrence.
• Understanding hybridization (sp³, sp², sp) is essential for predicting geometry and reactivity.
• Classification systems (formula types, functional groups, homologous series) provide systematic ways to represent and name organic molecules.
• Alkyl groups and degree of substitution terminology are foundational for IUPAC nomenclature and mechanistic descriptions.
• Mastery of sigma/pi bonding and structural representations enables translation between diagrams, formulas, and 3-D molecular geometry.