Introduction to Organic Chemistry and Chemical Bonding

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Topic 1 — Key Facts for Kenyatta University (Kenya) Core concept: Organic chemistry is the study of carbon compounds; carbon’s unique ability to form four covalent bonds and catenation makes organic chemistry vast and central to life High-yield point: Know the difference between sigma (σ) and pi (π) bonds, sp³/sp²/sp hybridisation, and the tetravalent nature of carbon ⚡ Exam tip: Kenyatta University organic chemistry exams frequently test orbital hybridisation diagrams and functional group recognition — be able to draw sp³, sp², and sp carbon centres accurately

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Introduction to Organic Chemistry

Organic chemistry is the branch of chemistry that studies carbon compounds, particularly hydrocarbons and their derivatives. The name “organic” originated from the 19th-century belief that organic compounds could only be produced by living organisms — a theory disproved when Friedrich Wöhler synthesised urea (NH₂CONH₂) from inorganic ammonium cyanate in 1828. Today, we know that organic chemistry encompasses millions of compounds, both natural and synthetic.

Carbon is uniquely suited to forming diverse compounds due to three key properties:

1. Tetravalence: Carbon has four valence electrons and forms four covalent bonds
2. Catenation: Carbon atoms bond to each other to form long chains, rings, and complex network structures
3. Multiple bonding: Carbon forms double (C=C) and triple (C≡C) bonds, greatly expanding possible structures

Electronic Configuration and Bonding

Carbon’s ground state electronic configuration is 1s² 2s² 2p², giving it four valence electrons. However, the observed valency of carbon is four. This is explained by hybridisation — the mixing of atomic orbitals to form new hybrid orbitals.

sp³ Hybridisation (Tetrahedral) Occurs when one s orbital mixes with three p orbitals to form four equivalent sp³ hybrid orbitals, each at 109.5° to the others.

Example: Methane (CH₄)

• Carbon undergoes sp³ hybridisation
• Four C–H bonds are formed by overlap of sp³ orbitals with hydrogen 1s orbitals
• Geometry: Tetrahedral, bond angle 109.5°

sp² Hybridisation (Trigonal Planar) Occurs when one s orbital mixes with two p orbitals to form three sp² hybrid orbitals (120° apart) with one unhybridised p orbital perpendicular to the plane.

Example: Ethene (C₂H₄)

• Each carbon is sp² hybridised
• The C=C double bond consists of one σ bond (sp²–sp² overlap) and one π bond (lateral p orbital overlap)
• Plane geometry, bond angle approximately 120°

sp Hybridisation (Linear) Occurs when one s orbital mixes with one p orbital to form two sp hybrid orbitals (180° apart) with two unhybridised p orbitals.

Example: Ethyne (C₂H₂)

• Each carbon is sp hybridised
• The C≡C triple bond consists of one σ bond (sp–sp overlap) and two π bonds (p–p overlaps)
• Linear geometry, bond angle 180°

Types of Carbon Bonds

Bond Type	Hybridisation	Bond Length (pm)	Bond Strength (kJ/mol)
C–C (single)	sp³	154	347
C=C (double)	sp²	134	614
C≡C (triple)	sp	120	839

Sigma (σ) and Pi (π) Bonds

Sigma (σ) bonds:

• Formed by head-on overlap of orbitals along the bond axis
• Found in all single bonds
• Cylindrically symmetrical about the bond axis
• Can rotate freely (single bonds allow rotation)

Pi (π) bonds:

• Formed by side-to-side (lateral) overlap of p orbitals
• Found in double bonds (one π bond) and triple bonds (two π bonds)
• Electron density concentrated above and below the bond plane
• Cannot rotate (creates rigidity in structures)

Introduction to Functional Groups

A functional group is a specific atom or group of atoms within a molecule that determines the chemical behaviour of that compound. In organic chemistry, the functional group is the reactive centre.

Key Functional Groups to Know:

Functional Group	Structure	Suffix	Example Compound
Alkane	C–H	-ane	Methane (CH₄)
Alkene	C=C	-ene	Ethene (C₂H₄)
Alkyne	C≡C	-yne	Ethyne (C₂H₂)
Alcohol	–OH	-ol	Ethanol (C₂H₅OH)
Aldehyde	–CHO	-al	Ethanal (CH₃CHO)
Ketone	–CO–	-one	Propanone (CH₃COCH₃)
Carboxylic acid	–COOH	-oic acid	Ethanoic acid (CH₃COOH)
Ester	–COO–	-oate	Ethyl ethanoate (CH₃COOC₂H₅)
Amine	–NH₂	-amine	Ethanamine (C₂H₅NH₂)

⚡ Exam Tip: In Kenyatta University examinations, students who fail to identify functional groups correctly lose marks on both structural questions and reaction prediction questions. Always name the functional group before predicting reactivity.

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Detailed Orbital Theory and Bonding in Organic Chemistry

Molecular Orbital Theory

Molecular orbital (MO) theory provides a more complete description of bonding than valence bond theory. According to MO theory, atomic orbitals combine to form molecular orbitals that extend over the entire molecule. Electrons occupy these orbitals according to the Aufbau principle, Hund’s rule, and the Pauli exclusion principle.

Key Concepts:

• Bonding orbitals (lower energy): Occupied by bonding electrons, stabilise the molecule
• Antibonding orbitals (higher energy): Empty or occupied by anti-bonding electrons, destabilise the molecule
• Non-bonding orbitals: Have the same energy as atomic orbitals, localised on specific atoms

For the C=C double bond:

• σ bond: Lower energy, formed by in-phase overlap
• σ* antibonding orbital: Higher energy, empty in the ground state
• π bond: Formed by lateral overlap of p orbitals above and below the plane
• π* antibonding orbital: Higher energy, empty in the ground state

Resonance and Resonance Structures

Some molecules cannot be adequately described by a single Lewis structure. Resonance structures are two or more valid Lewis structures that differ only in the placement of electrons (never atoms).

Example: Benzene (C₆H₆)

• Two major resonance structures exist, each with alternating double bonds
• The actual structure is a hybrid of all resonance forms
• All C–C bonds are equivalent, intermediate in length between single and double bonds (140 pm)
• The resonance stabilisation energy of benzene is approximately 150 kJ/mol

⚡ Exam Tip: When asked to draw resonance structures, remember:

• Only electrons move; never move atoms
• The resonance hybrid is more stable than any individual contributing structure
• Structures with complete octets and minimum formal charges are major contributors

Formal Charge Calculations

Formal charge helps determine the most likely electron distribution in a molecule:

Formal Charge = (Valence electrons) − (Non-bonding electrons) − ½(Bonding electrons)

Example: Nitrate ion (NO₃⁻)

• Three major resonance structures exist
• One oxygen bears a formal charge of −1 in one resonance form
• All three oxygens are equivalent in the actual hybrid structure

Isomerism in Organic Chemistry

Structural Isomerism: Same molecular formula, different connectivity

• Chain isomerism: Different carbon skeleton (e.g., butane vs isobutane)
• Positional isomerism: Different position of functional group (e.g., propan-1-ol vs propan-2-ol)
• Functional group isomerism: Different functional group (e.g., dimethyl ether vs ethanol: both C₂H₆O)

Stereoisomerism: Same connectivity, different spatial arrangement

• Geometric isomerism (cis/trans): Due to restricted rotation about a double bond or in a ring (e.g., maleic acid vs fumaric acid)
• Optical isomerism (enantiomerism): Non-superimposable mirror images; presence of a chiral centre (asymmetric carbon — four different groups attached)
  • Chiral centre: Carbon with four different substituents
  • Example: Lactic acid (CH₃CH(OH)COOH) has one chiral centre and exists as two enantiomers
  • Enantiomers have identical physical properties (melting point, boiling point) except for rotation of plane-polarised light
  • D/L notation (old) vs R/S notation (IUPAC) for absolute configuration

R/S Nomenclature for Chiral Centres:

1. Priority: Rank four substituents by atomic number (highest = 1)
2. Orientation: Place lowest priority group away from viewer
3. Trace: Trace from priority 1 → 2 → 3
   • Clockwise = R (rectus, right)
   • Counter-clockwise = S (sinister, left)

⚡ Exam Tip: When asked to identify R/S configuration, always assign priorities correctly first — a common error is misranking substituents.

IUPAC Nomenclature: The Foundation

IUPAC naming follows these steps:

1. Identify the longest carbon chain containing the principal functional group
2. Number the chain from the end that gives the functional group the lowest possible number
3. Identify and name all substituents
4. Assemble the name: substituent positions + substituent names + parent chain name + principal group suffix

Key Nomenclature Rules:

• C₁–C₁₀: meth-, eth-, prop-, but-, pent-, hex-, hept-, oct-, non-, dec-
• Multiplicity of bonds: -an- (single), -en- (double), -yn- (triple)
• Position numbers always precede the group/prefix they describe
• Alphabetical order for substituent names (ignore prefixes: di-, tri-, n-, sec-, tert-)

Hydrocarbon Classification: A Summary

Saturated hydrocarbons (Alkanes): Contain only single bonds (C–C and C–H), formula CₙH₂ₙ₊₂ Unsaturated hydrocarbons:

• Alkenes: Contain at least one C=C double bond (CₙH₂ₙ)
• Alkynes: Contain at least one C≡C triple bond (CₙH₂ₙ₋₂)
• Aromatic hydrocarbons: Contain benzene ring; follow Hückel’s rule (4n+2 π electrons)

Environmental and Health Significance

Many organic compounds have critical environmental and health implications:

• CFCs (chlorofluorocarbons): Destroy stratospheric ozone
• VOCs (volatile organic compounds): Contribute to ground-level ozone and smog
• PCBs (polychlorinated biphenyls): Persistent organic pollutants
• Benzene: Known carcinogen; causes leukaemia

At Kenyatta University, understanding the environmental impact of organic chemistry is increasingly important, especially for students pursuing environmental chemistry pathways.

Laboratory Safety in Organic Chemistry

• Organic solvents (ethanol, acetone, dichloromethane) are flammable — keep away from open flames
• Many organic compounds are toxic — work in a fume hood for volatile substances
• Concentrated acids and bases must be handled with care
• Glassware should be checked for cracks before heating
• Never distill to dryness — explosions can result from peroxide formation

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