Organic Chemistry Chapter 1: Structure and Bonding – Comprehensive Study Notes

Chapter 1. Structure and Bonding

1. Introduction to Organic Chemistry

Organic chemistry is the study of carbon-containing compounds, originally defined as substances derived from living organisms. The definition has evolved over time, especially after Friedrich Wöhler's synthesis of urea from inorganic materials, disproving the need for a 'vital force'.

• Historical Definition: Study of compounds extracted from living organisms; required a 'vital force'.

• Modern Definition: Chemistry of carbon compounds; no need for vital force.

• Example: Wöhler’s synthesis of urea from ammonium cyanate.

Additional info: Organic chemistry now includes synthetic compounds and covers a vast array of molecules beyond those found in nature.

2. Examples of Organic Compounds in Daily Life

Organic compounds are ubiquitous in everyday life, found in food, medicines, beverages, and cosmetics.

• Caffeine (coffee, tea)

• Acetaminophen (Tylenol)

• Nicotine (tobacco)

• Fragrances (cosmetic scents)

Application: Understanding organic chemistry helps explain the structure, function, and reactivity of these compounds.

3. Carbon Element

Properties and Position in the Periodic Table

Carbon is central to organic chemistry due to its ability to form four covalent bonds and its position in the second row of the periodic table.

• Valence Electrons: Carbon has 4 valence electrons.

• Octet Rule: Atoms transfer or share electrons to achieve a filled shell (8 electrons).

• Tetravalency: Carbon always forms four bonds.

Element	Valence Electrons	Bonding Capacity
Carbon (C)	4	4
Nitrogen (N)	5	3
Oxygen (O)	6	2
Hydrogen (H)	1	1

4. Chemical Bonds

Types of Bonds

Bonds are formed by the interaction of electron densities between atoms.

• Covalent Bond: Sharing electrons in space between atoms.

• Ionic Bond: Attraction of opposite charges due to electron transfer.

Example: Formation of NaCl (ionic) vs. H2 (covalent).

5. Covalent Bond

Electron Sharing and Atomic Orbitals

Covalent bonds involve the sharing of electrons between two atoms, described by atomic orbitals.

• Atomic Orbital: Region in space with high probability of finding an electron.

• Bohr Model: Electrons occupy discrete energy levels (n=1,2,3...)

6. Covalent Bond – Atomic Orbitals and Wave Functions

Schrödinger Equation

The behavior of electrons is described by wave functions, solutions to the Schrödinger equation:

• $\left(\frac{h^2}{2m} \nabla^2 + V\right) \psi = E \psi$

• $H \psi = E \psi$

For hydrogen atom:

• 1s: $\psi_{100} = \psi_{1s}$

• 2p: $\psi_{210}, \psi_{211}, \psi_{21-1}$

• 3d: $\psi_{322}, \psi_{321}, \psi_{320}, \psi_{32-1}, \psi_{32-2}$

7. Covalent Bond – Energy Orbitals

Atomic Orbital Energy Levels

Atomic orbitals are mathematical functions describing electron probability around the nucleus. Energy increases with principal quantum number and orbital type.

• Carbon Configuration: $1s^2 2s^2 2p^2$

• Energy Order: $1s < 2s < 2p < 3s < 3p < 3d$

8. Atomic Orbitals – Geometry

Shapes and Electron Density

• s orbital: Spherical, 2 electrons

• p orbital: Dumbbell-shaped, 6 electrons (px, py, pz)

• d orbital: Complex shapes, 10 electrons (transition metals)

9. Atomic Orbitals – Molecular Orbitals

s-s Overlap and σ Bonds

Formation of molecular hydrogen (H2) involves s-s overlap, creating a σ (sigma) bond.

• Constructive Overlap: Lower energy, bonding MO

• Destructive Overlap: Higher energy, antibonding MO

For every σ-bonding, there is a corresponding σ*-antibonding orbital.

10. Atomic Orbitals – p-p and s-p Overlap

σ Bonds from p-p and s-p Overlap

• p-p Overlap: Linear overlap forms σ bond (e.g., F2)

• s-p Overlap: s and p orbitals overlap to form σ bond (e.g., HF)

11. Atomic Orbitals – π Bonds

π Bonds from Sideways p-p Overlap

• π Bond: Formed by sideways overlap of two parallel p orbitals (e.g., C=C in ethylene)

• Strength: π bonds are weaker than σ bonds

12. Covalent Bonds – σ and π Bonds

Bond Types and Overlap

• σ Bonds: s-s, s-px, px-px overlap (linear)

• π Bonds: py-py, pz-pz overlap (sideways)

• σ Bond Strength: σ bonds are stronger and less reactive than π bonds

13. Molecular Geometry of Organic Molecules

Bond Angles and Orbital Combination

Simple s and p orbitals cannot explain observed bond angles in organic molecules. Hybridization is required.

• Methane (CH4): 109.5° (tetrahedral)

• Ethylene (C2H4): 120° (trigonal planar)

• Acetylene (C2H2): 180° (linear)

14. Molecular Geometry – VSEPR and Hybridization

Explaining Bond Angles

• VSEPR Theory: Electron pairs around central atom determine geometry

• 4 pairs: tetrahedral

• 3 pairs: trigonal planar

• 2 pairs: linear

• Hybridized Orbitals: sp3 (tetrahedral), sp2 (trigonal), sp (linear)

15–20. Hybrid Atomic Orbitals

sp3, sp2, and sp Hybridization

• sp3 Hybridization: Combination of one s and three p orbitals; tetrahedral geometry (e.g., methane)

• sp2 Hybridization: Combination of one s and two p orbitals; trigonal planar geometry (e.g., ethylene)

• sp Hybridization: Combination of one s and one p orbital; linear geometry (e.g., acetylene)

Hybrid Orbitals	Hybridization	Geometry	Bond Angle
2	sp	Linear	180°
3	sp2	Trigonal planar	120°
4	sp3	Tetrahedral	109.5°

22. Covalent Bond – Polarity

Polar and Nonpolar Covalent Bonds

• Nonpolar Covalent Bond: Electrons shared equally (ΔE < 0.4)

• Polar Covalent Bond: Electrons shared unequally (0.4 < ΔE < 1.7)

• Dipole Moment: $\mu$ indicates bond polarity

Element	Electronegativity (E)
H	2.1
C	2.5
N	3.0
O	3.5
F	4.0

Application: C–H bonds are considered nonpolar due to similar electronegativities.

23. Ionic Bonding

Electron Transfer and Ionic Bonds

• Ionic Bond: Formed by transfer of electrons; ΔE > 1.7

• Example: LiF (ΔE = 3)

24. Lewis Structures

Drawing Lewis Structures

• Rule 1: Draw molecular skeleton

• Rule 2: Count total valence electrons

• Rule 3: Provide octets (duets for H) around all atoms

• Rule 4: Assign formal charges

Formal Charge Formula:

• $\text{Formal charge} = \text{Nb of valence electrons} - \text{Nb of lone pair electrons} - \frac{1}{2} \text{Nb of bonding electrons}$

25. Resonance

Resonance Structures

• Valid Lewis structures only

• Only electrons move, not nuclei

• Number of unpaired electrons must be constant

• Delocalization stabilizes ions (e.g., acetate ion)

26. Non-equivalent Resonance Forms

Major and Minor Resonance Contributors

• Major forms have full octets and minimal formal charges

• Minor forms may have incomplete octets or higher formal charges

27. Acid and Base

Definitions and Theories

• Arrhenius: Acids produce H3O+, bases produce OH- in water

• Brønsted-Lowry: Acids donate protons, bases accept protons

• Lewis: Acids accept electron pairs, bases donate electron pairs

Terminology: Lewis base = nucleophile; Lewis acid = electrophile

28. Acid Strength

Acid Dissociation Constant and pKa

• $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• $pK_a = -\log K_a$

• Stronger acid = smaller pKa

Compound	pKa
CH4	~50
NH3	~36
H2O	~15.7
HF	~3.2

29. Acid Strength – Comparative Analysis

Factors Affecting Acid Strength

• Electronegativity of substituents

• Resonance stabilization

• Inductive effects

Example: Carboxylic acids with different X substituents (F, Cl, Br, I, H) – more electronegative X increases acid strength.

30. Summary

• Covalent bond: Nonpolar, polar

• Carbon hybridization: sp3 (tetrahedral), sp2 (trigonal), sp (linear)

• Lewis structure, resonance, acid/base strength

Key skill: Determining the shape of organic molecules is essential.

31. Problems

Practice Questions

• For BH3, BH4, NH3, NH4:

  • Draw Lewis structure

  • Determine hybridization of central atom

  • Describe geometry

  • Calculate formal charge

• Compare acidity of CH4 vs. CICH3, CH3CH2OH vs. CH3COOH

• Draw resonance structures for H2CNN and identify the major contributor