Organic Chemistry Topics and Foundations 1

1. Foundations

• Atomic structure & bonding (s, p orbitals, hybridization: sp, sp², sp³)

• Resonance & formal charges

• Acids and bases (pKa, conjugate acids/bases, stability)

• Functional groups overview

ELECTRONIC CONFIGURATION

1. What is electronic configuration?

• It shows how electrons are distributed among the different orbitals (s, p, d, f).

• This arrangement follows rules that minimize energy and explain chemical behavior.

2. The 3 Main Rules

1. Aufbau principle

   • Electrons fill orbitals from lowest to highest energy.

   • Order:
     1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

2. Pauli exclusion principle

   • Each orbital can hold max 2 electrons, with opposite spins.

3. Hund’s rule

   • In a set of orbitals with the same energy (like the 3 p orbitals), electrons fill singly first (with parallel spins) before pairing up.

3. Example: Oxygen (O, atomic number 8)

• Total electrons = 8.

• Filling order:

  • 1s² (2 electrons)

  • 2s² (2 electrons)

  • 2p⁴ (4 electrons → each p orbital gets one first, then pairing starts)

👉 Configuration = 1s² 2s² 2p⁴

4. Example: Sodium (Na, atomic number 11)

• Total electrons = 11.

• Filling order:

  • 1s² 2s² 2p⁶ 3s¹

• Outer electron is in the 3s orbital → explains why sodium is very reactive (easily loses that 1 electron).

5. Shorthand (Noble Gas) Notation

Instead of writing all orbitals, use the previous noble gas in brackets.

• Oxygen: [He] 2s² 2p⁴

• Sodium: [Ne] 3s¹

How to write the electronic configuration

Step 1: Determine the total number of electrons

• This equals the atomic number of the element.

• Example: Carbon (C) → atomic number = 6 → 6 electrons.

Step 2: Fill orbitals using the Aufbau principle

• Follow the energy order:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

• Start from lowest energy orbital (1s) and keep adding electrons.

Step 3: Apply the Pauli Exclusion Principle

• Maximum 2 electrons per orbital.

• Each electron in an orbital must have opposite spins.

Step 4: Apply Hund’s Rule

• When filling degenerate orbitals (same energy, e.g., 2p has 3 orbitals), put one electron in each orbital first, then start pairing.

Step 5: Optional – Use Noble Gas Shorthand

• Find the previous noble gas (closed-shell) and write it in brackets, then add remaining electrons.

Step 6: Write the configuration

• Use superscripts for the number of electrons in each orbital.

• Example:

  • Oxygen (8 e⁻) → 1s² 2s² 2p⁴ → [He] 2s² 2p⁴

  • Sodium (11 e⁻) → 1s² 2s² 2p⁶ 3s¹ → [Ne] 3s¹

Quick Tips

1. Always check your total electrons match the atomic number.

2. Remember the special order for transition metals and f-block elements (e.g., 4s fills before 3d, but 4s electrons may be lost first in ions).

3. Practice with a few examples to memorize the orbital filling order.

ORBITALS

1. s Orbitals

• Shape: Spherical (like a ball around the nucleus).

• How many per energy level: 1 (e.g., 1s, 2s, 3s…).

• Electrons it can hold: 2.

• Gets bigger as the energy level increases (2s is bigger than 1s).

2. p Orbitals

• Shape: Dumbbell-shaped (two lobes with a node at the nucleus).

• How many per energy level: 3 (px, py, pz) – aligned along x, y, and z axes.

• Electrons they can hold: 6 total (2 in each).

• Start appearing in the 2nd energy level (2p, 3p, etc.).

3. d Orbitals

• Shape: More complex (cloverleaf shapes, with four lobes, except one shaped like a dumbbell with a donut).

• How many per energy level: 5 orbitals.

• Electrons they can hold: 10 total.

• Start appearing in the 3rd energy level (3d, 4d, etc.).

4. f Orbitals

• Shape: Even more complex (multi-lobed).

• How many per energy level: 7 orbitals.

• Electrons they can hold: 14 total.

• Start appearing in the 4th energy level (4f, 5f, etc.).

5. Summary Table

HYBRIDISATION

• Hybridization is the concept of mixing atomic orbitals (s, p, sometimes d) on the same atom to form new equivalent orbitals called hybrid orbitals.

• These hybrid orbitals form covalent bonds and help explain molecular shapes and bond angles.

Key idea: Orbitals “mix” to minimize energy and allow atoms to form stable bonds.

Why hybridization happens

• Carbon in methane (CH₄) has 1s² 2s² 2p².

• Carbon forms 4 bonds but has only 2 unpaired electrons in 2p.

• So, one 2s orbital mixes with three 2p orbitals → 4 equivalent sp³ orbitals.

• Each orbital forms a σ bond with hydrogen → tetrahedral geometry.

Types of Hybridization

How to Determine Hybridization

1. Count the number of regions of electron density around the central atom (bonds + lone pairs).

2. Match to hybridization type:

   • 2 regions → sp (linear, 180°)

   • 3 regions → sp² (trigonal planar, 120°)

   • 4 regions → sp³ (tetrahedral, 109.5°)

   • 5 regions → sp³d (trigonal bipyramidal)

   • 6 regions → sp³d² (octahedral)

Examples

• CH₄: 4 σ bonds → 4 regions → sp³

• C₂H₄: Each carbon has 3 σ bonds + 1 π bond → 3 regions → sp²

• C₂H₂: Each carbon has 2 σ bonds + 2 π bonds → 2 regions → sp

• PCl₅: 5 σ bonds → 5 regions → sp³d

• SF₆: 6 σ bonds → 6 regions → sp³d²

Key Points

• Hybrid orbitals only affect σ bonds; π bonds come from unhybridized p orbitals.

• Lone pairs are counted as regions of electron density.

• Hybridization explains bond angles and shapes much better than just looking at unhybridized orbitals.

Resonance

• Occurs when a molecule can be represented by two or more valid Lewis structures.

• True structure = resonance hybrid (blend of all forms).

Rules for resonance structures:

1. Only electrons move, not atoms.

2. Minimize formal charges.

3. Maintain octets for second-row elements (C, N, O, F).

4. Negative charge prefers more electronegative atoms.

Example:

• Carboxylate anion (R–COO⁻)

  • Two resonance forms: negative charge can be on either oxygen

  • Hybrid → charge delocalized over both O

b) Formal Charges

• Helps determine most stable resonance structures.

Formula:

Formal charge (FC)=Valence electrons−(nonbonding electrons+1/2bonding electrons).

Example: NH₄⁺

• N valence = 5

• Nonbonding electrons = 0

• Bonding electrons = 8 → 8/2 = 4

• FC = 5 – (0 + 4) = +1 → matches +1 charge

Tips:

• Resonance structures with least formal charge and negative on electronegative atoms are most stable.

• Resonance stabilizes molecules and affects reactivity, acidity, and basicity.

TOPICS

2. Structure & Reactivity

• Stereochemistry (chirality, R/S, optical activity)

• Conformations (Newman projections, cyclohexane chair conformations)

• Intermolecular forces (hydrogen bonding, dipole, London dispersion)

3. Functional Group Chemistry

• Alkanes (properties, free radical halogenation)

• Alkenes (electrophilic addition, Markovnikov vs anti-Markovnikov, polymerization)

• Alkynes (addition reactions, acidity, metalation)

• Aromatic compounds (benzene, electrophilic aromatic substitution, directing groups)

• Alcohols, ethers, epoxides

• Aldehydes & ketones (nucleophilic addition, oxidation, reduction)

• Carboxylic acids & derivatives (esters, amides, acid chlorides, anhydrides)

• Amines (basicity, reactions, synthesis)

4. Reaction Mechanisms

• Substitution (SN1, SN2)

• Elimination (E1, E2)

• Radical reactions

• Rearrangements (hydride, methyl shifts)

5. Spectroscopy & Analysis

• IR spectroscopy (functional group identification)

• NMR spectroscopy (¹H, ¹³C: chemical shifts, splitting, integration)

• Mass spectrometry (fragmentation patterns)

6. Advanced Topics

• Organometallics (Grignard, organolithium, Gilman reagents)

• Pericyclic reactions (Diels–Alder, electrocyclic, sigmatropic)

• Biomolecules (carbohydrates, amino acids, peptides, nucleic acids)

• Green chemistry & modern applications