Learn & Review: Lewis Structures | Organic Chemistry Complete Course

Organic Chemistry Playlist Introduction and Lewis Structures

This document summarizes the first lesson of an organic chemistry playlist, focusing on Lewis structures as a review of general chemistry concepts. The playlist will cover Lewis structures, formal charges, hybridization, valence bond theory, molecular orbital theory, polarity, and intermolecular forces. The instructor will use a whiteboard format for instruction.

Introduction to Organic Chemistry

• Definition: Organic chemistry is primarily the study of carbon-containing compounds. While historically linked to living systems, the modern definition emphasizes carbon as the central element.
• Contrast with Inorganic Chemistry: Inorganic chemistry deals with all other elements not primarily focused on in organic chemistry.
• Playlist Approach: The course aims for a "middle-of-the-road" approach, acknowledging variations in teaching methods across different institutions. The instructor will highlight areas where multiple presentations exist and advise students to pay attention to their professor's approach.

Lewis Structures and Valence Electrons

• Core Concept: Lewis structures represent molecules by showing valence electrons.
• Determining Valence Electrons: The number of valence electrons for an element can be determined by its position on the periodic table:
  • Group 1: 1 valence electron
  • Group 2: 2 valence electrons
  • Group 13 (Boron's column): 3 valence electrons
  • Group 14 (Carbon's column): 4 valence electrons
  • Group 15 (Nitrogen's column): 5 valence electrons
  • Group 16 (Oxygen's column): 6 valence electrons
  • Group 17 (Halogens): 7 valence electrons
  • Group 18 (Noble Gases): 8 valence electrons (generally not involved in bonding in this context).

The Octet Rule and Its Exceptions

• The Octet Rule: Most atoms strive to achieve eight valence electrons to attain a stable electron configuration, similar to noble gases. This arises from the capacity of electron shells (s and p subshells) to hold a total of eight electrons (2 in the s orbital + 6 in the p orbitals).
• Exceptions to the Octet Rule:
  • Expanded Octet: Atoms in the third period and below can accommodate more than eight valence electrons due to the availability of d subshells. This is common for elements like sulfur but less relevant for core organic chemistry elements (C, N, O, H).

    Example: Sulfuric acid (H₂SO₄) shows sulfur with 12 electrons around it.

  • Under the Octet Rule:
    • Hydrogen: Only needs two electrons to achieve the stable configuration of helium.
    • Beryllium (Be): Typically forms two bonds, resulting in only four valence electrons.
    • Boron (B) and Aluminum (Al): Often form three bonds, resulting in six valence electrons.

      Example: BH₃ has boron with only 6 electrons.

  • Odd Number of Valence Electrons: Species with an odd number of total valence electrons will have at least one atom with an incomplete octet. This is rare in organic chemistry.

    Example: NO (Nitric Oxide).

Achieving a Filled Octet: Ionic vs. Covalent Bonds

• Ionic Bonds: Occur between a metal and a nonmetal. Electrons are transferred from the metal to the nonmetal, allowing both to achieve a stable electron configuration.

  Example: Sodium (Na) with 1 valence electron and Chlorine (Cl) with 7 valence electrons. Chlorine gains an electron from sodium, forming Na⁺ and Cl⁻. Chlorine achieves an octet, and sodium's full second shell is considered stable.

• Covalent Bonds: Occur between two nonmetals. Electrons are shared between atoms to achieve a stable electron configuration. This is the predominant type of bonding in organic chemistry.

  Example: Two chlorine atoms sharing a pair of electrons to form a Cl-Cl bond. Each chlorine atom counts the shared pair, giving them a total of 8 electrons.

Drawing Lewis Structures: Step-by-Step Examples

The general process for drawing Lewis structures involves:

1. Count total valence electrons.
2. Identify the central atom (usually the least electronegative, excluding hydrogen).
3. Connect atoms with single bonds to form a skeleton.
4. Fill the octets of outside atoms first.
5. Place remaining electrons on the central atom.
6. Check if the central atom has a filled octet. If not, form multiple bonds (double or triple) by moving lone pairs from outer atoms to create bonds with the central atom.

Examples:

• CH₄ (Methane):

  • Total valence electrons: 4 (C) + 4(H) = 8
  • Carbon is central.
  • Single bonds connect C to each H.
  • Hydrogens are satisfied with 2 electrons each.
  • No electrons remaining for carbon.
  • Carbon has 8 electrons (2 from each bond). Structure is complete.

• NH₃ (Ammonia):

  • Total valence electrons: 5 (N) + 3(H) = 8
  • Nitrogen is central.
  • Single bonds connect N to each H.
  • Hydrogens are satisfied.
  • 2 electrons remaining, placed as a lone pair on nitrogen.
  • Nitrogen has 8 electrons (6 from lone pair + 2 from each bond). Structure is complete.

• H₂CO (Formaldehyde):

  • Total valence electrons: 2(H) + 4(C) + 6(O) = 12
  • Carbon is central.
  • Single bonds connect C to 2 H's and 1 O.
  • Hydrogens are satisfied. Oxygen needs electrons.
  • No electrons remaining for carbon.
  • Carbon only has 6 electrons. Oxygen has lone pairs.
  • A lone pair from oxygen forms a double bond with carbon.
  • Carbon now has 8 electrons (2 from each single bond + 4 from double bond). Oxygen has 8 electrons (4 from lone pairs + 4 from double bond). Structure is complete.

• CH₃CHO (Acetaldehyde):

  • Total valence electrons: 2(C) + 4(H) + 6(O) = 18
  • Condensed formula indicates CH₃ group bonded to the CHO group.
  • Skeleton: CH₃ bonded to C, which is bonded to H and O.
  • Fill outside atoms (hydrogens satisfied, oxygen gets lone pairs).
  • Central carbon in CH₃ has 8 electrons. The other carbon has 6 electrons.
  • Oxygen shares a lone pair to form a double bond with the second carbon.
  • Both carbons and oxygen now have octets.

• CH₃COOH (Acetic Acid):

  • Total valence electrons: 2(C) + 4(H) + 2(O) = 24
  • Skeleton based on condensed formula: CH₃ bonded to C, which is bonded to two O's, one of which is bonded to H.
  • Initial attempt might lead to atoms not having their typical number of bonds or octets.
  • The correct structure involves the carbon bonded to both oxygens, with one oxygen double-bonded to the carbon and the other single-bonded to the carbon and the hydrogen.
  • This structure ensures all atoms have their typical number of bonds (Carbon: 4, Oxygen: 2, Hydrogen: 1) and octets. The instructor notes that formal charge (discussed in the next lesson) helps determine the most stable structure.