Bonding in Organic CompoundsBB2425

Bonding in Organic Compounds

Introduction

• The study focuses on various aspects of bonding in organic compounds, outlined by Samantha Drake.

Contents

• Key topics include:

• Revision

• Atomic Structure

• Electronic Structure

• Orbitals

• Intermolecular Bonding

• Sigma Bonding

• Pi Bonding

• Aromatic Hydrocarbons

Atomic Structure

• Atoms are composed of subatomic particles, characterized as follows:

  • Proton (p): ~1 amu, +1 charge

  • Neutron (n): ~1 amu, 0 charge

  • Electron (e-): ~0 amu, -1 charge

• The atomic structure includes a nucleus formed by protons and neutrons with electrons orbiting around it, influencing chemical behavior.

Electronic Structure

Heisenberg Uncertainty Principle

• The principle states it is not possible to precisely determine both the position and momentum of an electron at the same time. This uncertainty fundamentally impacts the understanding of electron locations and behaviors.

Orbitals

• Orbitals represent regions within an atom where there is a likelihood of finding electrons. Different types include:

  • s orbitals

  • p orbitals

  • d orbitals

  • f orbitals

Carbon's Electronic Configuration

• Carbon is represented as 1s²2s²2p².

• Only valence electrons (outer shell electrons) participate in bonding. Carbon has 4 valence electrons.

• Energy levels denote:

  • Inner shell (n=1)

  • Outermost shell (n=2), aka valence shell

• Key principles:

  • Aufbau Principle: Electrons fill the lowest energy orbitals first.

  • Hund’s Rule: Electrons will singly occupy degenerate orbitals before pairing.

Intramolecular Bonding

• Stability in atoms is achieved through full valence shells.

  • Duplet Rule: Applies to H and He, requires 2 electrons.

  • Covalent Bonding: Atoms share electrons to achieve stability based on the octet rule.

  • Different types: Single, double, and triple bonds.

Sigma (σ) and Pi (π) Bonding

Sigma Bonds

• All covalent bonds contain at least one σ bond. Examples:

  • Formed through head-on overlaps of orbitals.

  • Can occur between:

    • s orbitals

    • p orbitals

    • hybrids of s and p orbitals (sp, sp², sp³).

Pi Bonds

• Always occur alongside σ bonds and form from the lateral overlap of p orbitals. Pi bonds are weaker than sigma bonds.

  • They cannot exist without a preceding sigma bond.

Hybridization in Organic Molecules

Single Bonds

• sf and hybridization lead to tetrahedral shapes for saturated carbon compounds (example: Ethane).

Double and triple bonds

• Ethene (CH₂=CH₂):

  • Contains sp² hybridized carbons, featuring 3 σ bonds and 1 π bond. The molecule is planar and cannot rotate.

• Ethyne (C≡C):

  • Features sp hybridized carbons, 2 σ bonds, and 2 π bonds, with a linear configuration.

Aromatic Hydrocarbons

Definition and Properties

• Originally linked to fragrance, now denotes a specific type of bonding.

• Benzene (C₆H₆) is a fundamental aromatic compound characterized by a hexagonal ring with delocalized electrons, providing unusual stability.

Structure

• Benzene structure illustrates equal bond lengths among carbon bonds, providing resonance stability due to delocalized electrons.

Substituted Aromatics

• Methylbenzene (Toluene) has a methyl group substituted onto the benzene ring.

• Naming conventions depend on the number of carbon atoms in the substituent:

  • Alkyl groups with 6 or fewer carbons are named as alkyl-substituted benzenes.

  • IUPAC names commonly used include methylbenzene for toluene and hydroxybenzene for phenol.

Conclusion

• Understanding bonding types and structuring in organic compounds elucidates their behavior and properties, relevant in various chemical applications.