Chapter 1: Structure and Bonding – Foundations of Organic Chemistry

Structure and Bonding in Organic Chemistry

Introduction to Organic Chemistry

Organic chemistry is the study of the structure, properties, composition, reactions, and synthesis of carbon-containing compounds. This field is foundational for understanding the chemistry of living things and many synthetic materials.

• Organic compounds are primarily composed of carbon and hydrogen, often with oxygen, nitrogen, sulfur, phosphorus, and halogens.

• Examples include proteins (hair, enzymes), DNA, foods, and medicines.

• Over 50 million known chemical compounds contain carbon.

Properties of Carbon

Carbon is unique among the elements due to its ability to form a vast array of compounds.

• Carbon is a group 4A element in the periodic table.

• It has four valence electrons and can form four covalent bonds.

• Carbon atoms can bond to each other, forming long chains and rings.

• This versatility leads to immense diversity in organic compounds.

The Position of Carbon in the Periodic Table

Carbon is located in group 4A (14) of the periodic table, alongside elements like silicon and germanium. Its position explains its tetravalency and bonding behavior.

Atomic Structure – The Nucleus

Atoms consist of a dense, positively charged nucleus surrounded by a cloud of negatively charged electrons.

• The nucleus contains protons (positively charged) and neutrons (electrically neutral).

• The volume around the nucleus is occupied by orbiting electrons.

• The typical diameter of an atom is about $2 \times 10^{-10}$ m (200 pm).

• 1 ångström (Å) = $10^{-10}$ m = 100 pm.

Atomic Number, Mass Number, and Isotopes

• Atomic number (Z): Number of protons in the nucleus.

• Mass number (A): Number of protons plus neutrons.

• Atoms of the same element have the same atomic number but may have different mass numbers (isotopes).

• Atomic weight: Weighted average mass of an element’s naturally occurring isotopes (in amu).

Atomic Orbitals and Electron Configuration

Electrons occupy regions of space called orbitals, which are grouped into shells around the nucleus.

• Types of orbitals: s (spherical), p (dumbbell-shaped), d, and f.

• Organic chemistry primarily involves s and p orbitals.

• Each orbital can hold two electrons (with opposite spins).

• Electron shells are filled in order of increasing energy (Aufbau principle): $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d$

• First shell: 1s orbital (2 electrons); Second shell: 2s and 2p orbitals (8 electrons); Third shell: 3s, 3p, 3d (18 electrons).

Hybridization and Molecular Geometry

Carbon forms different types of hybrid orbitals to achieve stable bonding arrangements:

• sp3 hybridization: Combination of one s and three p orbitals forms four equivalent, tetrahedrally oriented orbitals (bond angle ≈ 109.5°). Example: methane (CH4).

• sp2 hybridization: Combination of one s and two p orbitals forms three planar orbitals (bond angle ≈ 120°), with one unhybridized p orbital perpendicular to the plane. Example: ethylene (C2H4).

• sp hybridization: Combination of one s and one p orbital forms two linear orbitals (bond angle = 180°), with two unhybridized p orbitals. Example: acetylene (C2H2).

Example: Methane (CH4)

• All C–H bonds are identical and tetrahedral.

• Bond length: 109 pm; Bond strength: 439 kJ/mol.

Example: Ethane (C2H6)

• Two carbons bond via sp3–sp3 overlap; each C–H bond is sp3–1s.

• C–C bond length: 154 pm; Bond strength: 377 kJ/mol.

Example: Ethylene (C2H4)

• Each carbon is sp2 hybridized; forms a double bond (one σ and one π bond).

• Bond angles: ≈ 120°.

Example: Acetylene (C2H2)

• Each carbon is sp hybridized; forms a triple bond (one σ and two π bonds).

• Bond angle: 180°.

Development of Chemical Bonding Theory

Atoms form bonds to achieve greater stability, often by attaining a noble gas electron configuration.

• Ionic bonds: Formed by transfer of electrons and electrostatic attraction between ions.

• Covalent bonds: Formed by sharing electrons between atoms (common in organic compounds).

• Molecule: Neutral collection of atoms held together by covalent bonds.

Representing Chemical Structures

• Electron-dot (Lewis) structures: Show valence electrons as dots.

• Line-bond (Kekulé) structures: Show covalent bonds as lines between atoms.

• Condensed structures: Omit some or all bonds, grouping atoms together (e.g., CH3CH2OH).

• Skeletal structures: Carbon atoms are implied at line ends and intersections; hydrogens on carbon are omitted.

Bonding Electrons and Lone Pairs

• Bonding electrons: Shared between atoms in covalent bonds.

• Lone pairs: Valence electrons not involved in bonding; influence molecular shape and reactivity.

• Example: In ammonia (NH3), nitrogen has one lone pair and three bonding pairs.

Valency and Octet Rule

• Number of covalent bonds an atom forms depends on the number of electrons needed to complete its octet.

• Carbon: 4 bonds; Nitrogen: 3 bonds; Oxygen: 2 bonds; Hydrogen: 1 bond.

Summary Table: Hybridization and Geometry

Hybridization	Geometry	Bond Angle	Example
sp3	Tetrahedral	109.5°	Methane (CH4)
sp2	Trigonal planar	120°	Ethylene (C2H4)
sp	Linear	180°	Acetylene (C2H2)

Practice Example

• Draw the line-bond structure for propane (C3H8), predict bond angles (≈109.5°), and indicate the overall shape (tetrahedral at each carbon).

• Identify the number of hydrogens bonded to each carbon in a given skeletal structure and determine the molecular formula.

Additional info: These notes summarize foundational concepts in atomic structure, bonding, and molecular geometry, which are essential for further study in organic chemistry.