Organic Chemistry: Structure and Bonding

Organic Chemistry Chapter 1: Structure and Bonding

Learning Objectives

  • (1.1) Atomic Structure: The Nucleus

  • (1.2) Atomic Structure: Orbitals

  • (1.3) Atomic Structure: Electron Configurations

  • (1.4) Development of Chemical Bonding Theory

  • (1.5) Describing Chemical Bonds: Valence Bond Theory

  • (1.6) sp³ Hybrid Orbitals and the Structure of Methane

  • (1.7) sp³ Hybrid Orbitals and the Structure of Ethane

  • (1.8) sp² Hybrid Orbitals and the Structure of Ethylene

  • (1.9) sp Hybrid Orbitals and the Structure of Acetylene

  • (1.10) Hybridization of Nitrogen, Oxygen, Phosphorus, and Sulfur

  • (1.11) Describing Chemical Bonds: Molecular Orbital Theory

  • (1.12) Drawing Chemical Structures

What is Organic Chemistry?

  • Organic chemistry focuses on compounds containing carbon, which are essential to life.

  • Examples include:

    • Proteins (e.g., those making up hair)

    • DNA

    • Foods

    • Medicines

Origins of Organic Chemistry

  • The foundation of organic chemistry dates back to the mid-1700s.

  • Early organic compounds were primarily obtained from living sources (plants and animals).

  • These compounds were often low-melting solids and challenging to isolate or purify.

  • Initially, organic compounds were thought to possess a "vital force" due the presence of life, leading to the belief that they could not be synthesized artificially.

  • Significant milestones:

    • 1816: Chevreul discovered that soap could be separated into various organic compounds, termed fatty acids.

    • 1828: Wöhler demonstrated the synthesis of urea from ammonium cyanate, an inorganic salt, disproving the vital force theory.

Characteristics of Organic Compounds

  • Organic chemistry: study of compounds primarily containing carbon.

  • Over 50 million known carbon-containing compounds exist.

  • Carbon is classified as a group 4A element.

  • Carbon can share four valence electrons, forming four covalent bonds, capable of bonding with itself to create long chains and rings, facilitating immense diversity in compound structure.

Atomic Structure - The Nucleus

  • An atom’s nucleus is positively charged and surrounded by negatively charged electrons.

  • The nucleus consists of:

    • Protons: Positively charged particles.

    • Neutrons: Electrically neutral particles.

Atomic Number and Atomic Mass

  • Atomic number (Z): Number of protons in an atom's nucleus.

  • Atoms of a given element all share the same atomic number.

  • Mass number (A): Total number of protons and neutrons.

  • Isotopes: Atoms with the same atomic number but differing mass numbers.

  • Atomic mass (atomic weight): The weighted average mass of an element’s naturally occurring isotopes expressed in atomic mass units (amu).

Atomic Structure - Orbitals

  • The wave equation describes the behavior of electrons in atoms.

  • Wave function or Orbital (Ψ): Solution of the wave equation that predicts an electron's location.

  • The plot of ext{Ψ}^2 indicates the probability density of an electron's location; there is no defined boundary for the electron cloud.

Types of Orbitals

  • Different orbitals include:

    • s orbitals: Spherical and centered around the nucleus.

    • p orbitals: Dumbbell-shaped, with the nucleus at the middle.

    • d orbitals: Elongated dumbbell shapes centered on the nucleus.

Energy Levels of Electrons in an Atom

  • Electron orbitals are organized into shells around the nucleus:

    • 1st shell: capacity of 2 electrons

    • 2nd shell: capacity of 8 electrons

    • 3rd shell: capacity of 18 electrons

Atomic Structure: Electron Configurations

  • Ground-state electron configuration: Arrangement of an atom’s electrons in its lowest energy state.

  • Following the Aufbau principle, electrons fill the lowest-energy orbitals first (1s → 2s → 2p → 3s → 3p → 4s → 3d).

Spin and Orbital Filling

  • Electrons spin around an axis, oriented up (↑) or down (↓).

  • Pauli exclusion principle: No more than two electrons can occupy an orbital, both must have opposite spins.

  • Hund's rule: When multiple orbitals are of equal energy, one electron occupies each orbital singly before pairing begins.

Worked Examples

  • For Magnesium (Mg): Ground-state configuration 1s^2 2s^2 2p^6 3s^2, indicating 2 outermost electrons.

  • For Sulfur (S): Atomic number 16, (1s^2 2s^2 2p^6 3s^2 3p^4) configuration shows 6 valence electrons.

Development of Chemical Bonding Theory

  • Notable contributors: Kekulé and Couper recognized carbon's tetravalency.

  • Van't Hoff and Le Bel described specific spatial directions for carbon's four bonds.

  • Valence shell: The outermost shell of an atom contributing to bonding.

  • Atoms form bonds that yield more stability than isolated atoms.

Types of Bonds

  • Ionic bonds: Formed through electron transfer resulting in electrostatic attractions between ions.

  • Covalent bonds: Result from electron sharing, typical in organic compounds.

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

Structural Representations

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

  • Line-bond structures (Kekulé structures): Use lines to represent covalent bonds.

Counting Covalent Bonds

  • The number of covalent bonds formed is based on how many additional valence electrons are needed to achieve a stable octet:

    • Carbon: 4 valence electrons, forms 4 bonds.

    • Nitrogen: 5 valence electrons, forms 3 bonds.

Non-Bonding Electrons

  • Lone pairs: Valence electrons not involved in bonding (e.g., nitrogen in ammonia ext{(NH}_3 ext{)}) has two lone pairs).

Worked Examples for Structures

  • Draw both electron-dot and line-bond structures for various compounds (e.g., chloromethane, water, ethane, formaldehyde, ethylene).

Valence Bond Theory

  • Description of covalent bond formation: occurs when orbitals from two atoms overlap.

  • H-H bond formed from overlapping single 1s orbitals yielding a $ ext{sigma (σ) bond}$.

  • Bond strength for H-H is 436 kJ/mol. Bond length is ideal for maximum stability.

Hybridization and Structures

  • sp³ hybridization (e.g., in methane ext{(CH}_4 ext{)}): Combination of one s orbital with three p orbitals to form four equivalent tetrahedral orbitals.

  • sp² hybridization (e.g., in ethylene ext{(C}2 ext{H}4 ext{)}): Combination of one s and two p orbitals to form three planar orbitals arranged at 120°.

  • sp hybridization: Forms linear configurations (e.g., in acetylene ext{(C}2 ext{H}2 ext{)}) with triple bonds.

Bonding Comparisons

  • Table of C-C and C-H bonds in various hydrocarbons showing bond strengths and lengths. Highest bond strength occurs in acetylene.

Hybridization of Other Elements

  • Elements like nitrogen, oxygen, phosphorus, and sulfur exhibit hybridization influencing molecular shapes and bond angles (e.g., 107.1° for methylamine).

Formal Charges

  • Formal charge calculated by comparing bonded atoms to the free atom's valence electron structure.

  • Summary of common formal charges for various atoms.

Resonance

  • Molecules with multiple structures represented as resonance forms contribute to a hybrid structure.

  • Example: Acetate ion exhibits two resonance forms.

Drawing Chemical Structures

  • Shorthand methods for writing structures include condensed and skeletal forms; non-carbon atoms are explicitly shown while carbon is typically implied at line intersections.

Worked Examples

  • Draw structures for selected hydrocarbons, indicating the number of hydrogen atoms bonded to each carbon, and determining molecular formulas.