chem 40a: ch 2

Organic Chemistry with a Biological Emphasis

Introduction

  • Author: Tim Soderberg

  • Affiliation: University of Minnesota, Morris

  • License: Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License

  • URL for License: https://creativecommons.org/licenses/by-nc-sa/4.0/

Chapter 2: Structure and Bonding II

Overview

  • The chapter introduces significant aspects of organic structure and bonding.

  • The chapter will detail concepts such as covalent bonding, hybridization, and various types of intermolecular forces impacting organic compounds.

Learning Outcomes

  • Sketch 1s, 2s, and 2p atomic orbitals.

  • Draw s (sigma) and p (pi) molecular orbitals.

  • Identify resonance structures and their stability.

  • Determine hybridization and geometry in molecules.

  • Explain trends in melting point, boiling point, density, and solubility of organic molecules.

  • Identify and explain the impact of hydrogen bonding.

Section 2.1: Covalent Bonding in Organic Molecules

Valence Bond Theory

  • Definition: The theory is used to describe how atoms in organic molecules form bonds via overlap of atomic orbitals, each containing one electron.

Formation of Covalent Bonds

  • Covalent Bonding: Involves sharing of electrons between atoms.

  • Overlap of atomic orbitals leads to lower potential energy and the stability of the bonded electrons.

  • If atomic nuclei approach too far, repulsive force increases leading to instability.

Bond Length and Strength

  • Optimal distance for hydrogen (H2) molecule: 74 pm (picometers).

  • Bond Dissociation Energy (BDE) for H-H bond: 435 kJ/mol.

  • Examples of bond lengths and strengths in organic molecules:

    • C–C single bond: 150 pm

    • C=C double bond: 130 pm, BDE: generally stronger than single bonds

    • C≡C triple bond: 120 pm, BDE: highest among carbon bonds

    • C–H bond: 100-110 pm

Covalent Bond Characteristics

  • Covalent bonds can be pictured as springs; they're flexible and can change length under stress.

  • Sigma (σ) Bond: Interface of overlapping orbitals maintains cylindrical symmetry.

Section 2.1A: The s Bond in the H2 Molecule

  • Two hydrogen atoms form a covalent bond by overlapping their 1s orbitals, creating a stable bond.

Section 2.1B: sp3 Hybridization in Methane (CH4)

  • Carbon's electron configuration: 1s² 2s² 2p² indicates it must hybridize for four bonds with sp³ hybridization, leading to a tetrahedral geometry.

Section 2.1C: sp² and sp Hybridization

  • sp² Hybridization in Ethene (C2H4): Discusses planar structure, bond angles of 120°, and the presence of a double bond comprised of one σ bond and one π bond.

  • sp Hybridization in Ethyne (C2H2): Accounts for the linear geometry of acetylene with 180° angles and its triple bond consisting of one σ bond and two π bonds.

Section 2.2: Molecular Orbital Theory, Conjugation, and Aromaticity

Introduction to Molecular Orbital Theory

  • Molecular orbital (MO) theory explains the bonding characteristics, particularly for systems exhibiting resonance or conjugation.

Interaction of Atomic Orbitals

  • Atomic orbitals combine to form bonding and antibonding MOs based on constructive and destructive interference of wave functions.

Conjugation

Example: 1,3-Butadiene

  • Conjugated systems of alternating double/single bonds exhibit additional stability due to electron delocalization, leading to lower reactivity and unique physical properties.

Aromatic Compounds

  • Characterized by cyclic structure, planar geometry, sp² hybridization, and adherence to Hückel’s rule for stabilization.

Examples of Aromaticity in Molecules

  • Benzene: All bonds are equivalent (138 pm) and exhibit reduced reactivity due to delocalization of π electrons across the ring.

Biological Implications of Aromatic Compounds

  • Heterocycles like pyridine exhibit similar aromatic characteristics, enhancing their stability and solubility in biological systems.

Section 2.3: Resonance

What is Resonance?

  • Resonance describes the phenomenon of multiple Lewis structures that contribute to the actual structure of a molecule.

  • Resonance Contributors: Each structure is a snapshot that reflects different ways the π-electrons may be distributed.

Resonance Hybrid

  • The true molecular structure is a blend of all resonance contributors, leading to stabilization.

Drawing Rules

  • Only p-bonded or lone-pair electrons can be moved; s-bonds are fixed.

  • Resonance structures must have the same net charge and adhere to octet rules.

Section 2.4: Non-Covalent Interactions

Types of Non-Covalent Interactions

  • Dipole-Dipole Interactions: Occurs between polar molecules with a net dipole.

  • Ion-Dipole Interactions: Engaged when ions interact with polar molecules—crucial in biological systems.

  • Hydrogen Bonds: Strong attraction between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom.

Examples

  • Water as a solvent can stabilize charges and polar groups through H-bonding, significantly influencing biochemical reactions.

Applications in Organic and Biological Chemistry

  • Understanding non-covalent interactions allows predictions regarding the solubility, melting, and boiling points of organic compounds.

Section 2.5: Physical Properties of Organic Compounds

Solubility

  • Solubility of organic compounds in water depends on the balance of hydrophilic groups (increase solubility) and hydrophobic groups (decrease solubility).

Boiling and Melting Points

  • Melting and boiling points are influenced by the strength of non-covalent interactions. Higher interaction strength raises boiling/melting points.

Example: Lipids and Proteins

  • The rigidity and stability of various lipids at room temperature correlate with their degree of saturation (double bonds) and length of hydrocarbon chains.

Conclusion

  • Mastery of the structure and behavior of organic compounds is essential for understanding the chemical nature of living organisms and biochemistry.