Chapter 1 Organic Chemistry Notes

Atomic Structure

1. Components of the Atom

  • The nucleus of the atom consists of:

    • Protons: Positively charged particles.

    • Neutrons: Uncharged particles.

  • The electron cloud contains:

    • Electrons: Negatively charged particles.

2. Atomic Number and Mass Number

  • Atomic Number: Defined as the number of protons in the nucleus, which equals the number of electrons in a neutral atom.

  • Mass Number: Total number of protons and neutrons in the nucleus.

    • Example: Carbon has 6 protons and 6 neutrons, hence its atomic number is 6, and mass number is 12.

  • Atomic Weight: The weighted average mass of an element's isotopes, reported in atomic mass unit (amu).

Ions

1. Types of Ions

  • Cations: Positively charged ions (fewer electrons than protons).

  • Anions: Negatively charged ions (more electrons than protons).

Isotopes

  • Definition: Atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

    • Most isotopes of carbon have 6 protons and 6 neutrons.

The Periodic Table

1. Arrangement of Elements

  • Elements in the same row (period) are similar in size.

  • Elements in the same column (group) exhibit similar electronic and chemical properties.

2. First Row Elements

  • Contains only one orbital in the first shell.

  • Maximum electrons: 2 (H and He).

3. Second Row Elements

  • Contains four orbitals: one 2s and three 2p orbitals.

  • Maximum capacity for valence electrons in this row: 8.

Atomic Orbitals

  • s Orbital: Spherical shape, lower in energy compared to other orbitals in the same shell.

  • p Orbital: Dumbbell shape; contains a node at the nucleus, higher in energy than s orbitals.

Bonding Concepts

1. Definition of Bonding

  • Bonding: The joining of two atoms to achieve a stable arrangement through the gain, loss, or sharing of electrons to attain the nearest noble gas configuration.

    • Types of Bonds:

    • Ionic Bonds: Formed by the transfer of electrons from one atom to another (e.g., NaCl).

    • Covalent Bonds: Formed by sharing electrons between atoms (e.g., H₂).

2. Ionic Bonding

  • Typically involves elements on the far left and right of the periodic table.

  • A positive cation (left side) attracts a negative anion (right side).

3. Covalent Bonding

  • Occurs between elements with similar electronegativities, cementing the presence of bonds between identical atoms.

  • Molecules are covalent compounds formed by shared electron pairs.

    • Example: Molecular Hydrogen (H₂) where each hydrogen satisfies its valence by forming one bond.

Valence Electrons and Bond Formation

  • Second Row Elements: Max eight electrons allowed.

  • Molecules:

    • Elements with 1-4 valence electrons form an equivalent number of bonds in neutral conditions (e.g., BF₃, CH₄).

    • Elements with 5 or more valence electrons adjust bonding to achieve an octet (NH₃).

Nonbonded Electrons

  • Unshared electrons are referred to as lone pairs.

  • Lewis Structures: Representations that depict valence electrons, where each line signifies a bond.

Drawing Lewis Structures

General Steps

  1. Arrange Atoms: Place atoms next to each other with hydrogen and halogens on the perimeter.

  2. Count Electrons: Total all valence electrons, considering charges.

  3. Arrange Electrons: Fill valence shells with bonds and lone pairs appropriately.

  4. Assign Formal Charges: Determine charges for stability and correct structure representation.

Multiple Bonds

  • Criteria: Form multiple bonds when necessary to satisfy octet requirements.

Formal Charge Calculation

  • Formal Charge: Charge assigned to individual atoms within a Lewis structure, derived from:

    • Count of electrons “owned” by an atom (bonds and lone pairs).

Isomers

  • Isomers: Compounds with the same molecular formula but different arrangements of atoms.

    • Example: Ethanol vs. Dimethyl Ether.

Resonance Structures

  • Definition: Some molecules can’t be described by a single Lewis structure; instead, they use resonance structures to show different electron arrangements without altering atomic positions.

  • Resonance Hybrid: The actual structure is a composite of all resonance forms, displaying properties of both.

Basic Principles of Resonance Theory

  • Resonance structures are not real entities but representations of electron distribution.

  • Movement of electrons occurs only in the hybrids, not between individual resonance forms.

Drawing Resonance Structures

Rules

  1. Different structures must exhibit alternate placements of multiple bonds/nonbonded electrons.

  2. Must maintain the same number of unpaired electrons.

  3. Only valid Lewis structures allowed.

Molecular Shape Determination

  1. Bond Length and Bond Angle: Key factors that define a molecule's geometry.

  2. Bond Length Trends:

    • Decreases across the periodic table.

    • Increases down a group.

  3. Molecular Geometry: Based on the number of groups around central atoms (VSEPR theory).

Summary Geometry Table (Bonding Groups)

  • 2 Groups: Linear (180°)

  • 3 Groups: Trigonal Planar (120°)

  • 4 Groups: Tetrahedral (109.5°)

Drawing and Interpreting 3D Structures

  • Bonds represented with solid lines (in-plane), wedges (front of the plane), and dashed lines (behind the plane).

Hybridization

Concept Overview

  • Hybridization: The process of mixing atomic orbitals to form new hybrid orbitals optimized for bonding.

    • sp³, sp², and sp hybridization are defined by the number of hybrid and unhybridized orbitals.

Determining Hybridization


  • Based on the number of surrounding groups (atoms or lone pairs) which corresponds to the number of orbitals to hybridize.


Hybridization Types Table:

Number of groups

Type of hybrid orbital


2

sp


3

sp²


4

sp³

Bonding and Hybridization Examples

  • Carbon Skeletons:

    • Methane (CH₄): sp³ hybridized with tetrahedral geometry.

    • Ethylene (C₂H₄): sp² hybridized, with restricted rotation due to π bonds.

    • Acetylene (C₂H₂): sp hybridized, featuring triple bonds formed by σ and π bonds.

Electronegativity

Definition and Measurement

  • Electronegativity: The ability of an atom to attract electrons in a bond.

Bond Polarity

  • The difference in electronegativity values results in varying bond polarities:

    • Nonpolar Bonds: Equal sharing of electrons (C–C, C–H).

    • Polar Bonds: Unequal sharing leads to partial charges (C–O).

Assessing Molecular Polarity

  • Procedures include analyzing bond dipoles based on electronegativity differences and determining if they cancel based on molecular geometry.

    • Example: Water (H₂O) is polar, while carbon dioxide (CO₂) is nonpolar due to dipole cancellation.