Organic Chemistry Basics - First Semester

Overview

• Focus: Organic compounds with carbon atoms.
• Key Concept: Carbon likes to form 4 bonds.

Bonding Preferences of Elements

• Group 1 Elements: Form 1 bond (e.g., Hydrogen)
• Beryllium: Forms 2 bonds
• Boron: Forms 3 bonds
• Carbon: Forms 4 bonds (4 valence electrons)
• Nitrogen: Forms 3 bonds
• Oxygen: Forms 2 bonds
• Halogens (Fluorine, Chlorine, Bromine, Iodine): Generally form 1 bond
  • Note: Can form up to 7 bonds (covered in Lewis Structures video)

Drawing Lewis Structures

• Water (H2O): Oxygen forms 2 bonds
  • Structure: Oxygen in the center with 2 hydrogens and 2 lone pairs.
• Methyl Fluoride (CH3F): Carbon forms 4 bonds, Hydrogen 1 bond, Fluorine 1 bond
  • Structure: Carbon in the center, 3 Hydrogens, 1 Fluorine with 3 lone pairs

Bond Types and Polarity

• Hydrogen Bonds: H attached to N, O, or F (e.g., H2O)
• Polar Bonds: Electronegativity difference ≥ 0.5
  • Example: Carbon-Fluorine bond (Electronegativity: C=2.5, F=4.0)
  • Concept: Charge separation in a polarized bond
    • Fluorine (partial negative), Carbon (partial positive)
• Non-Polar Bonds: Electronegativity difference < 0.5
  • Example: Carbon-Hydrogen bond (Electronegativity: H=2.1, C=2.5)

Covalent vs. Ionic Bonds

• Covalent Bonds: Electrons are shared
  • Non-Polar: Electrons shared equally (e.g., H2)
  • Polar: Electrons shared unequally (e.g., HF)
• Ionic Bonds: Electrons are transferred
  • Example: Sodium (Na) gives an electron to Chlorine (Cl), forming Na+ and Cl-
  • Attraction between positive (cation) and negative (anion) ions

Saturated and Unsaturated Compounds

• Alkanes: Saturated compounds (only single bonds)
  • Formula: CₙH₂ₙ₊₂ (e.g., Methane CH4, Ethane C2H6)
  • Example Names and Formulas:
    • Methane: CH4
    • Ethane: C2H6
    • Propane: C3H8
    • Butane: C4H10
    • Etc.
• Alkenes: Unsaturated compounds (at least one double bond)
  • Example: Ethene (C2H4)
• Alkynes: Unsaturated compounds (at least one triple bond)
  • Example: Ethyne (C2H2)

Bond Length and Strength

• Single Bonds: Longest and weakest
  • Length: ~154 pm
• Double Bonds: Intermediate
  • Length: ~133 pm
• Triple Bonds: Shortest and strongest
  • Length: ~120 pm

Sigma and Pi Bonds

• Single Bond: 1 sigma bond
• Double Bond: 1 sigma and 1 pi bond
• Triple Bond: 1 sigma and 2 pi bonds
• Strength: Sigma bonds are stronger than pi bonds

Bond Order

• Single Bond: Bond order = 1
• Double Bond: Bond order = 2
• Triple Bond: Bond order = 3

Hybridization

• Determining Hybridization:
  • Count attached atoms and lone pairs
  • Types:
    • sp3 (4 groups)
    • sp2 (3 groups)
    • sp (2 groups)
• Example:
  • CH bond in sp3 hybridized carbon and s hybridized hydrogen is sp3-s
  • For sp hybridized carbon, CH bond is sp-s

Formal Charge Calculation

• Formula: Formal charge = (Valence electrons) - (Bonds + Dots)
• Examples:
  • For Carbon with 3 bonds and no lone pairs: FC = +1
  • For Sulfur with 1 bond and 3 lone pairs: FC = -1
  • For Nitrogen in NH4+: FC = +1

Bonding and Non-Bonding Electrons

• Bonding Electrons: In bonds (2 per bond)
• Non-Bonding Electrons: Lone pairs (2 per pair)

Functional Groups and Organic Structures

• Alcohol: -OH group
  • Example: Ethanol (C2H5OH)
• Aldehyde: -CHO group
  • Example: Ethanal (CH3CHO)
• Ether: -O- between carbons
  • Example: Dimethyl Ether (CH3OCH3)
• Ketone: Carbonyl group (C=O) in the middle
  • Example: Propanone (CH3COCH3)
• Carboxylic Acid: -COOH group
  • Example: Pentanoic Acid (C5H10O2)
• Ester: Combination of carbonyl and ether groups
  • Example: Methyl Ethanoate (CH3COOCH3)

Expanding Organic Structures

• Steps:
  • Identify number and type of each group (CH3, CH2, etc.)
  • Arrange to ensure carbon has 4 bonds, hydrogen 1 bond
• Example:
  • Given: CH3(CH2)3COOH
  • Expand: CH3-CH2-CH2-CH2-COOH