Chapter 1: A Review of General Chemistry – Electrons, Bonds, and Molecular Properties

1.1 Introduction to Organic Chemistry

What is Organic Chemistry?

Organic chemistry is the study of carbon-containing molecules and their reactions. It focuses on understanding how molecules interact, how bonds are broken and formed, and the role of electrons in these processes.

• Organic compounds are distinguished from inorganic compounds by the presence of carbon atoms.

• Organic chemistry is essential because organic compounds make up food, clothes, pharmaceuticals, and plastics.

• Reactions in organic chemistry involve the movement and rearrangement of electrons.

1.2 The Structural Theory of Matter

Why are Bonds Important?

The structural theory of matter states that substances are defined by the specific arrangement of atoms. The molecular formula alone is not sufficient to define a compound because different structures (isomers) can have the same formula.

• Constitutional isomers are compounds with the same molecular formula but different connectivity of atoms.

• Common atoms bonded to carbon include nitrogen (N), oxygen (O), hydrogen (H), and halides (F, Cl, Br, I).

• Each element generally forms a specific number of bonds (e.g., carbon forms four, nitrogen three, oxygen two, hydrogen one).

1.3 Covalent Bonding and Atomic Structure

Characteristics of a Covalent Bond

Covalent bonds form when atoms share electrons. The stability and length of a bond are determined by the balance of attractive and repulsive forces:

• Attractive forces: between positively charged nuclei and negatively charged electrons.

• Repulsive forces: between nuclei and between electrons.

Atomic Structure

• Protons (+1 charge) and neutrons (neutral) are in the nucleus.

• Electrons (−1 charge) occupy orbitals outside the nucleus.

• Valence electrons are in the outermost shell and are involved in bonding.

Counting Valence Electrons

• For main group (Group A) elements, the group number equals the number of valence electrons.

• Valence electrons can also be determined from electron configuration.

Simple Lewis Structures

• Draw atoms with dots representing valence electrons.

• Atoms share pairs of electrons to complete octets (except hydrogen, which completes a duet).

• Example: NH3 (ammonia) has a lone pair on nitrogen.

1.5 Polar Covalent Bonds and Electronegativity

Definition of Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a bond. Fluorine (F) is the most electronegative element.

Types of Covalent Bonds

• Nonpolar covalent bond: Electronegativity difference < 0.5

• Polar covalent bond: Electronegativity difference between 0.5 and 1.7

• Ionic bond: Electronegativity difference > 1.7

Bond Polarity

• Electrons shift toward the more electronegative atom, creating a dipole.

• The greater the electronegativity difference, the more polar the bond.

1.6 Reading Bond-Line Structures

Introduction and Drawing

Bond-line structures are a simplified way to represent organic molecules, especially large ones.

• Drawn in a zigzag format; each corner or endpoint represents a carbon atom.

• Double and triple bonds are shown with two and three lines, respectively.

1.7 Atomic Orbitals

Electron Density and Orbital Types

• Atomic orbitals are regions of space where electrons are likely to be found (90–95% probability).

• s orbitals are spherical; p orbitals are dumbbell-shaped and degenerate (equal energy).

• Electron configuration follows the Aufbau principle, Pauli exclusion principle, and Hund’s rule.

1.8 Valence Bond Theory

Bond Formation

• Bonds form when atomic orbitals overlap (constructive interference).

• Direct overlap forms a sigma (σ) bond.

• Example: In H2, the σ bond forms from the overlap of two 1s orbitals.

1.10 Hybridized Atomic Orbitals

sp3, sp2, and sp Hybridization

• sp3 hybridization: 25% s-character, 75% p-character; forms four equivalent orbitals (e.g., CH4).

• sp2 hybridization: 33% s-character, 67% p-character; three sp2 orbitals and one unhybridized p orbital (e.g., in alkenes).

• sp hybridization: 50% s-character, 50% p-character; two sp orbitals and two unhybridized p orbitals (e.g., in alkynes such as acetylene).

Sigma and Pi Bonds

• σ bonds: Head-on overlap of orbitals (stronger, lower in energy).

• π bonds: Side-by-side overlap of p orbitals (weaker, electron density above and below the plane).

Bond Strength and Length

• σ bonds are stronger than π bonds.

• Bond length decreases with increasing s-character: sp3 > sp2 > sp.

Table: Comparison of Bond Lengths and Bond Energies

Compound	Bond Type	Bond Length (pm)	Bond Energy (kJ/mol)
Ethane (C–C)	sp3–sp3 σ	154	376
Ethylene (C=C)	sp2–sp2 σ + π	134	728
Acetylene (C≡C)	sp–sp σ + 2π	120	962
Additional info: Values are typical and may vary slightly by source.

1.11 Molecular Geometry (VSEPR Theory)

VSEPR Theory and Steric Number

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular geometry based on electron pair repulsion.

• Steric number = number of bonded atoms + number of lone pairs on the central atom.

• Steric number 4: sp3 (tetrahedral), 3: sp2 (trigonal planar), 2: sp (linear).

Common Geometries

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

Importance of Lone Pairs

• Lone pairs affect molecular geometry by repelling bonding pairs, altering bond angles (e.g., H2O is bent, not linear).

1.12 Molecular Polarity & Dipoles

Dipole Moment

• Electronegativity differences create polar bonds and molecular dipoles.

• Dipole moment ($\mu$) is calculated as:

$\mu = \delta \times d$

• where $\delta$ is the magnitude of partial charge and $d$ is the distance between charges.

• Units: debye (D), where 1 D = $10^{-10}$ esu·cm.

Geometry and Polarity

• Molecular geometry must be considered to determine overall polarity (e.g., water is bent and polar).

• Electrostatic potential maps visually represent molecular polarity.

Table: Dipole Moments of Common Solvents

Solvent	Dipole Moment (D)
Water (H2O)	1.85
Acetone	2.88
Diethyl ether	1.15
Hexane	0
Additional info: Values are approximate and for 20°C.

1.13 Intermolecular Forces

Types of Intermolecular Forces

• Dipole-dipole interactions: Attractions between polar molecules.

• Hydrogen bonding: Strong dipole-dipole attraction involving H bonded to N, O, or F.

• London dispersion forces: Weak, transient attractions due to temporary dipoles in all molecules, especially significant in nonpolar molecules.

Dipole-Dipole Attractions

• Polar molecules align so that positive and negative ends attract, raising boiling and melting points (e.g., acetone vs. isobutylene).

Hydrogen Bonding

• Occurs when H is bonded to N, O, or F and interacts with a lone pair on another electronegative atom.

• Hydrogen bonds are about 20 times weaker than covalent bonds but much stronger than other dipole-dipole interactions.

• Protic solvents can hydrogen bond; aprotic solvents cannot.

• Hydrogen bonding affects boiling points and is crucial in biological molecules (e.g., DNA double helix, protein folding).

London Dispersion Forces

• Result from temporary, induced dipoles due to random electron motion.

• Strength increases with molecular surface area and mass.

• Branching decreases surface area and thus weakens dispersion forces.

• Example: Gecko feet utilize large surface area for strong dispersion forces.

1.14 Solubility

Polar vs Nonpolar Compounds

• "Like dissolves like": Polar compounds dissolve in polar solvents; nonpolar compounds dissolve in nonpolar solvents.

• Hydrogen bonding and dipole-dipole interactions enhance solubility among polar compounds.

Soap and Micelles

• Soaps allow polar (water) and nonpolar (oil) substances to mix by forming micelles.

• Micelles have a nonpolar interior that traps oil and grease, allowing them to be washed away in water.

Additional info: Some tables and values were inferred or supplemented for completeness. For further practice, refer to SkillBuilder exercises mentioned in the original notes.