CHM102 – Introduction to Organic Chemistry

Historical Background of Organic Chemistry

• Definition: Organic chemistry is the study of carbon and its various compounds. Carbon is unique due to its ability to form a seemingly unlimited number of compounds.
• Impact: Organic compounds are vital to daily life, influencing medicine, agriculture, and general biological existence.
• Theoretical Origins (Oparin, 1923): Theory suggests organic chemistry began with the Big Bang, where components like ammonia, nitrogen, carbon dioxide, and methane combined to form amino acids. This was later verified in laboratory settings by Miller (1950).
• Ancient Usage: Romans and Egyptians utilized organic substances from natural sources as dyes, medicines, and poisons, though they did not understand their chemical composition.
• Chronological Development:
  • 1769 (Scheele): First isolation of organic compounds from nature in a pure state.
  • 1784 (Lavoisier): Development of analytical methods to determine elemental composition.
  • 1807 (Berzelius): Proposed the "vital force" theory, suggesting organic chemicals possessed a special force that directed their synthesis, making laboratory synthesis impossible.
  • 1828 (Frederich Wöhler): Disproved "vital force" by synthesizing urea (a natural component of urine) in a lab by heating ammonium cyanate. This discovery paved the way for modern synthetic organic chemistry.
  • 1858 (Kekulé and Couper): Introduced Valence Theory, the first theory regarding the bonding structures of organic chemistry.
  • Late 19th Century: Knowledge expanded into biological systems, specifically studying proteins and DNA.

The Chemical Bond and Atomic Theory

• Quantum Mechanics (1926): Developed by Heisenberg and Schroedinger, providing mathematical solutions for electronic energy levels in atoms.
• Electronic Structure: Electrons exist in energy levels (shells) surrounding the nucleus. Energy increases with distance from the nucleus.
• Shells and Orbitals: Within shells are subshells or orbitals containing up to two electrons.

Electron Energy and Orbitals Table

• Orbital Shapes:
  • 1s and 2s: Spherical shapes.
  • 2p: Barbell-type shapes aligned along the x, y, and z axes (designated $p_x$, $p_y$, and $p_z$).
• Aufbau Principle: Electrons fill the lowest energy levels first until all electrons (equal to the atomic number) are used.

Electron Configurations (1st and 2nd Row)

• H (1): $1s^1$
• He (2): $1s^2$
• Li (3): $1s^2, 2s^1$
• Be (4): $1s^2, 2s^2$
• B (5): $1s^2, 2s^2, 2p^1$
• C (6): $1s^2, 2s^2, 2p^2$
• N (7): $1s^2, 2s^2, 2p^3$
• O (8): $1s^2, 2s^2, 2p^4$
• F (9): $1s^2, 2s^2, 2p^5$
• Ne (10): $1s^2, 2s^2, 2p^6$ (Inert; completely filled).

Electronegativity and Types of Bonding

• Electronegativity: The ability of an atom to attract electrons.
  • Horizontal Trend: Increases from left to right (e.g., $Li < Be < B < C < N < O < F$).
  • Vertical Trend: Increases from bottom to top (e.g., $I < Br < Cl < F$).
• Electropositive Elements: Elements that easily lose electrons to gain a positive charge (e.g., Alkali metals).

Ionic Bonds

• Occurs between atoms with vastly different electronegativities.
• Includes the transfer of an electron from one atom to another so both achieve a stable noble gas configuration.
• Example: Lithium Fluoride (LiF): The $2s^1$ electron of Lithium is transferred to the $2p^5$ orbital of