Organic Chemistry Basics

Overview

This lecture covers foundational concepts in organic chemistry, focusing on bonding, molecular structures, key functional groups, and how to draw and interpret Lewis structures of organic molecules.

Basic Principles of Organic Chemistry

• Organic chemistry studies compounds containing carbon atoms.
• Carbon typically forms four bonds; other elements have characteristic bonding preferences (e.g., H:1, O:2, N:3, F/Cl/Br/I:1).
• Understanding bonding preferences aids in drawing correct Lewis structures.

Lewis Structures and Bonding Types

• Lewis structures represent bonds (shared electron pairs) and lone pairs (non-bonding electrons).
• A bond between H and O (as in H₂O) is a covalent bond, with added lone pairs for O to satisfy the octet rule.
• Hydrogen bonds occur when H is attached to N, O, or F, explaining phenomena like water’s high boiling point.

Polarity and Bond Types

• A polar bond has unequal electron sharing due to differences in electronegativity (ΔEN ≥ 0.5).
• Nonpolar bonds (ΔEN < 0.5) share electrons equally (e.g., C-H).
• Covalent bonds involve shared electrons; ionic bonds involve electron transfer.
• Cations are positively charged; anions are negatively charged.

Alkanes, Alkenes, Alkynes & Saturation

• Alkanes (saturated hydrocarbons) follow the formula CₙH₂ₙ₊₂ (e.g., methane, ethane, propane).
• Alkenes contain at least one C=C double bond (e.g., ethene/ethylene).
• Alkynes have a C≡C triple bond (e.g., ethyne/acetylene).
• Alkenes/alkynes are unsaturated; alkanes are saturated.

Bond Length, Strength, and Types

• Single bonds (154 pm) are longer than double (133 pm) or triple bonds (120 pm).
• Triple bonds are strongest, single bonds are weakest.
• Single bonds are sigma (σ); double bonds have one sigma and one pi (π); triple have one sigma and two pi bonds.
• Sigma bonds are stronger than pi bonds.

Hybridization

• Count attached atoms + lone pairs to determine hybridization:
  • 4 groups: sp³
  • 3 groups: sp²
  • 2 groups: sp
• Bond hybridization is determined by the atoms involved (e.g., sp³–s).

Sigma and Pi Bonds Counting

• Each single bond is one sigma bond.
• Double bonds: one sigma, one pi.
• Triple bonds: one sigma, two pi.

Formal Charge Calculation

• Formal charge = valence electrons − (number of bonds + number of dots/lone electrons).
• Example: Carbon with three bonds and no dots has a +1 formal charge (carbocation). With three bonds and two dots: −1 (carbanion).
• Radicals have odd numbers of electrons.

Functional Groups & Common Structures

• Alcohols (—OH group): ethanol.
• Aldehydes (—CHO group): ethanal (acetaldehyde).
• Ketones (internal carbonyl): propanone (acetone).
• Ethers (C—O—C): dimethyl ether.
• Esters (—COO—): methyl ethanoate.
• Carboxylic acids (—COOH): pentanoic acid.

Drawing and Expanding Structures

• Methyl groups (CH₃) usually at ends; CH₂ in the middle; CH branches can have substituents like Br or OH.

Key Terms & Definitions

• Alkane — Saturated hydrocarbon with single bonds only.
• Alkene — Unsaturated hydrocarbon with at least one double bond.
• Alkyne — Unsaturated hydrocarbon with a triple bond.
• Functional Group — Specific group of atoms conferring characteristic properties to molecules.
• Sigma bond (σ) — Strongest type of covalent bond, formed by head-on orbital overlap.
• Pi bond (π) — Weaker covalent bond, formed by side-on overlap after sigma bond.
• Hybridization — Mixing of atomic orbitals to form new hybrid orbitals.
• Formal Charge — Apparent charge on an atom calculated by a specific formula.

Action Items / Next Steps

• Practice drawing Lewis structures and identifying functional groups.
• Memorize alkane names up to decane.
• Calculate formal charges for sample atoms and molecules.
• Review and expand condensed structural formulas.