Introduction to Organic Chemistry

Organic chemistry is the study of carbon compounds, fundamental to life and living organisms.
- Organic compounds include proteins, DNA, food, and medicines.
- Understanding organic chemistry is crucial for advances in medicine and biological sciences.

An atom consists of:
- A dense, positively charged nucleus containing protons (positively charged) and neutrons (neutral).
- Negatively charged electrons surrounding the nucleus.
Properties of atomic structure:
- Atoms are neutral overall with equal numbers of protons and electrons.
- Diameter of a typical atom: $2 imes 10^{-10} ext{ m}$ or 200 pm.
- Size of the nucleus: about $10^{-14} ext{ to } 10^{-15} ext{ m}$ in diameter.
Atomic number $Z$ indicates the number of protons.
Mass number $A$ is the total number of protons and neutrons.
Isotopes have the same atomic number but different mass numbers (e.g., Carbon-12, Carbon-13, Carbon-14).
- Natural abundance of carbon isotopes: Carbon-12 (98.89%), Carbon-13 (1.11%), Carbon-14 (negligible).
Atomic weight is based on the average mass of an element’s isotopes, expressed in unified atomic mass units (u or Da).

Electrons are distributed in orbitals, described mathematically via wave equations.
- Orbitals represent regions around the nucleus where electrons are likely to be found.
- The wave function denoted by the Greek letter psi (ψ) - the square of the wave function $ext{(ψ}^2$ ) indicates the electron density.
Types of orbitals:
- s Orbitals: Spherical shape.
- p Orbitals: Dumbbell shape.
- d and f Orbitals: More complex shapes.
Orbitals are grouped into shells:
- 1st shell: 1s orbital, holds up to 2 electrons.
- 2nd shell: 2s+2p orbitals, holds up to 8 electrons.
- 3rd shell: 3s+3p+3d orbitals, holds up to 18 electrons.
p orbitals have lobes separated by a node, indicating regions with zero electron density.

The ground-state electron configuration is determined by:
- Rule 1: Lowest-energy orbitals fill first (Aufbau principle).
- Rule 2: Electrons have spin (up or down), only two can occupy one orbital (Pauli exclusion principle).
- Rule 3: Electrons occupy degenerate orbitals singly before pairing (Hund’s rule).
Example configurations:
- Hydrogen: 1s
- Carbon: 1s² 2s² 2p²

In the mid-1800s, it was determined that carbon is tetravalent and forms four bonds.
Kekulé proposed that carbon atoms can form chains and rings.
Van’t Hoff introduced the concept of three-dimensionality in bonding (tetrahedral orientation).

Valence Bond Theory describes the formation of covalent bonds through orbital overlap.
- Example: In H2, the H–H bond results from overlapping 1s orbitals, forming a sigma (σ) bond.
- Energy and stability: Bond formation releases energy; bond breaking absorbs energy. H–H bond strength is 436 kJ/mol.

Methane (CH4) exhibits tetrahedral bonding due to sp3 hybridization of carbon.
- Bond angles are 109.5°, with identical C–H bonds resulting from the orientation of sp3 hybrid orbitals.
- Each C–H bond strength is 439 kJ/mol.

Ethane (C2H6) bonds through sp3 hybrid overlap between carbon atoms and hydrogen.
- C–C bond length is 153 pm, with a bond strength of 377 kJ/mol; C–H bond strength is 421 kJ/mol.
- Bond angles approach 109.5°.

Ethylene (C2H4) contains a double bond formed by overlapping sp2 hybrid orbitals and unhybridized p orbitals.
- C=C bond forms a pi (π) bond in addition to the σ bond, resulting in shorter distances and stronger bonds compared to single bonds.
- C–C bond length is 134 pm, bond strength is 728 kJ/mol.

Acetylene (C2H2) comprises a triple bond (σ and two π bonds) formed from sp hybrid orbitals.
- Acetylene exhibits linear geometry with bond angles of 180°; C–H bond length is 106 pm, bond strength 558 kJ/mol.

Other elements also use hybridization:
- Nitrogen forms sp3 hybrid orbitals in methylamine (NH3).
- Oxygen's bonding in methanol leads to sp3 hybridization, bond angle of 108.5°.
- Phosphorus can form five bonds and sulfur can form four bonds, illustrating the complexity of organic structures.

Molecular Orbital Theory uses mathematical combinations of atomic orbitals to create molecular orbitals, which can spread over the entire molecule.
- Bonding (lower energy) and antibonding (higher energy) orbitals are established.

Chemists use various shorthand methods to draw molecular structures:
- Condensed structures omit certain bonds for clarity.
- Skeletal structures represent carbons at line intersections and do not typically show hydrogen atoms.

The impact of understanding chemistry transcends academic boundaries, affecting health, environment, and even economic factors. Decisions around chemical use must balance benefits against potential risks, highlighting the need for ethical considerations in chemical production and application.

Understanding atomic structures and how they bond is fundamental in organic chemistry to describe reactions and compound formations. Hybridization aids in visualizing molecular shapes and the behavior of electrons, critical for further study in chemical interactions and applications.
The significance extends beyond the classroom, influencing various sectors through chemical production, drug development, and agricultural practices, necessitating a comprehensive understanding of concepts such as molecular structure and bonding mechanisms.