Structure and Bonding 1

Overview

  • Course: PHAR201 Towards unbounded thinking.

  • Topic: Structure & Bonding

  • Instructor: Dr. Mohamed Salah

Learning Objectives

  • Importance of Structure and Bonding:

    • Structure and bonding are crucial for drug molecules.

    • Understanding molecular shape and its impact on drug efficacy.

  • Physicochemical Properties:

    • Key properties include polarity, formal charge, and resonance.

  • Structural Representations:

    • Recognize Lewis, condensed, and bond-line structures of organic molecules.

Module Learning Outcomes (LOs)

  • a1:

    • Describe the role of structure and bonding in drug molecules.

    • Understand their impact on molecular shape, physical/chemical properties, reactions, interactions, and pharmaceutical applications.

Atomic Structure

  • Atomic Number (Z):

    • The number of protons in an atom.

  • Mass Number (A):

    • Total number of protons and neutrons. (A = Z + N)

Bonding – Atoms, Electrons, and Orbitals

  • Atomic Characteristics for Carbon:

    • Z = 6 (number of protons),

    • A = 12 (total of protons and neutrons).

  • Isotopes:

    • Atoms with the same number of protons but different neutrons.

  • Electron Configuration:

    • Fixed number of protons and electrons; variable number of neutrons.

  • Relative Atomic Mass of Carbon: 12.011 g mol-1

Periodic Table Understanding

  • History of the Periodic Table:

    • Created by Dmitri Mendeleev in 1869.

  • Organization:

    • Rows (periods) - indicate the number of electron shells.

    • Columns (groups) - indicate the number of valence electrons.

  • Categories:

    • Metals, Metalloids, Non-metals, and Noble Gases.

Electron Distribution in Atomic Orbitals

  • First Shell:

    • Can hold 2 electrons (1s orbital).

  • Second Shell:

    • Contains 1 s orbital (2s) and 3 p orbitals (2p), holding 8 electrons total.

  • Third Shell:

    • Contains 1 s orbital (3s), 3 p orbitals (3p), and 5 d orbitals (3d), holding 18 electrons total.

Electron Filling Rules

  1. Aufbau Principle: Electrons fill the lowest-energy orbitals first.

  2. Pauli Exclusion Principle: Each orbital can hold two electrons with opposite spins.

  3. Hund’s Rule: When multiple orbitals are available, one electron enters each orbital before pairing starts.

Electron Configuration Example for Sulfur (S)

  • Configuration:

    • S: 1s² 2s² 2p⁶ 3s² 3p⁴

    • Notation: [Ne] 3s² 3p⁴

Structural Representations of Molecules

  • Molecular Formula:

    • Example: C₃H₈O

  • Condensed Formula:

    • Example: Ethoxyethane (CH₃CH₂OCH₂CH₃).

  • Lewis Structures:

    • Illustrates connectivity, useful for small molecules but takes time.

Bond-Line and Skeletal Formulas

  • Bond-Line Formulas:

    • Bonds are shown by lines; hydrogen atoms bonded to carbons are omitted.

  • Skeletal Formulas:

    • Carbon atoms are at the ends of lines or where lines intersect.

Covalent Bonding Overview

  • Electron Sharing: Covalent bonds form when atoms share electrons for a complete shell.

  • Lewis Structures in Covalent Bonds:

    • A single covalent bond is represented by a line between nuclei.

    • Illustrate two electrons being shared.

Carbon's Bonding Characteristics

  • Carbon's Valence Electrons:

    • Carbon has four valence electrons available for bonding.

  • Bonding in Methane (CH₄):

    • All four bonds in methane are identical.

References

  • Recommended Textbook:

    • McMurry, J. Organic Chemistry, 9th edition, 2019.

    • Available in the library and as a softcopy.