Introductory Organic Chemistry Concepts

Overview

• Introductory organic chemistry lecture covering bonding, Lewis structures, polarity, hydrocarbons, bond properties, hybridization, formal charge, and common functional groups.
• Emphasis on rules for drawing structures and naming small organic molecules.

Bonding Basics

• Carbon typically forms four bonds; other common elements: H (1), Be (2), B (3), N (3), O (2), halogens (1).hubbiug8h
• Second-row elements (C, N, O, F) prefer an octet (8 electrons) when possible.

Lewis Structures & Examples

• Water (H2O): O forms two bonds + two lone pairs to satisfy octet.
• Methyl fluoride (CH3F): C in center with 3 H and 1 F; F has three lone pairs to complete octet.
• Always ensure each atom has its preferred number of bonds and appropriate lone pairs.

Polarity And Hydrogen Bonding

• Electronegativity: C = 2.5, H = 2.1, F = 4.0.
• Polar covalent bond: electronegativity difference ≥ 0.5.
• C–F is polar (F partial negative, C partial positive); C–H is essentially nonpolar (difference 0.4).
• A polarized bond has charge separation (partial + and -).
• Hydrogen bonds occur when H is directly attached to N, O, or F; explains high boiling point of water.

Covalent vs Ionic Bonds

• Covalent: electrons are shared (equally → nonpolar, unequally → polar).
  • H2: nonpolar covalent (equal sharing).
  • HF: polar covalent (H partial +, F partial -); special case often called hydrogen bond when interacting between molecules.
• Ionic: electrons are transferred (example: Na + Cl → Na+ + Cl-), forming electrostatic attraction between cation and anion.

Hydrocarbons: Alkanes, Alkenes, Alkynes

• Alkanes: saturated (only single C–C bonds), general formula CnH2n+2.
  • 1 C: methane (CH4); 2: ethane (C2H6); 3: propane (C3H8); 4: butane (C4H10); 5: pentane; 6: hexane; 7: heptane; 8: octane; 10: decane.
• Alkenes: contain at least one double bond (C=C). Example: C2H4 = ethene.
• Alkynes: contain at least one triple bond (C≡C). Example: C2H2 = ethyne (acetylene).
• Unsaturated compounds: alkenes and alkynes (do not have maximum H per C). Saturated: alkanes.

Bond Length, Strength, and Order

• Bond length order (longest → shortest): single > double > triple.
  • C–C single: ~154 pm (1.54 Å)
  • C=C double: ~133 pm
  • C≡C triple: ~120 pm
• Bond strength order (weakest → strongest): single < double < triple (triple has more bonds to break).
• Bond order: single = 1, double = 2, triple = 3.

Sigma (σ) and Pi (π) Bonds

• Single bond = 1 σ bond.
• Double bond = 1 σ + 1 π.
• Triple bond = 1 σ + 2 π.
• σ bonds are stronger than π bonds; breaking a π is easier than breaking the σ.

Hybridization Rules

• Determine hybridization by counting groups (atoms + lone pairs) around the atom:
  • 4 groups → sp3 (s1 p3)
  • 3 groups → sp2 (s1 p2)
  • 2 groups → sp (s1 p1)
• For bonds like C–H, specify both atom hybridizations (e.g., sp3–s or sp–s).
• Example: carbon with four attachments = sp3; carbon with three attachments = sp2; with two = sp.

Counting σ and π Bonds (Method)

• Count every single bond as one σ bond.
• Add one σ for each double and one σ for each triple (already included in single/bond counting).
• Count π bonds: +1 for each double, +2 for each triple.
• Example molecule counted in lecture: 6 σ bonds and 2 π bonds.

Formal Charge Calculation

• Formula: Formal charge = (valence electrons) − (number of bonds + nonbonding electrons/dots).
• Carbon example cases:
  • C with 3 bonds, 0 dots: FC = 4 − 3 = +1 (carbocation).
  • C with 3 bonds, 1 lone pair electron pair (2 dots): FC = 4 − 4 = 0.
  • C with 3 bonds, 2 dots (one lone pair): FC = 4 − 5 = −1 (carbanion).
• Sulfur example: S (group 6A, 6 valence e-) with 1 bond + 3 lone pairs (6 dots): FC = 6 − 7 = −1.
• Nitrogen in ammonium: N (5 valence e-) with 4 bonds, 0 dots: FC = 5 − 4 = +1.
• Bonding electrons: each bond = 2 bonding electrons. Nonbonding electrons: each lone pair = 2 nonbonding electrons.

Functional Groups: Structures And Names

• Alcohol (–OH): example CH3CH2OH = ethanol; suffix “-ol”.
• Aldehyde (–CHO): carbonyl (C=O) at chain end bonded to H; two-carbon example = ethanal (common: acetaldehyde); suffix “-al”.
• Ether (R–O–R): oxygen between two carbons; CH3OCH3 = dimethyl ether.
• Ketone (C=O in the middle): carbonyl within chain, no H on carbonyl carbon; three-carbon example = propanone; suffix “-one”.
• Ester (R–C(=O)–O–R): contains carbonyl and an O–C bond; example named as methyl ethanoate.
• Carboxylic acid (R–C(=O)–OH): carbonyl plus hydroxyl on same carbon; five-carbon example = pentanoic acid; suffix “-oic acid”.

Structure Expansion Rules / Condensed Form Interpretation

• CH3 groups (methyl) usually at ends; CH2 groups typically in the middle.
• CH groups often branch off and can carry substituents (e.g., halogens, OH).
• To expand condensed formulas, ensure carbon has four bonds and correct placement of double bonds or oxygen to satisfy valency.

Key Terms And Definitions

Action Items / Next Steps (If Studying)

• Practice drawing Lewis structures and assigning lone pairs for common molecules.
• Memorize alkane names up to at least 10 carbons.
• Do example problems counting σ and π bonds, and calculating formal charges.
• Practice hybridization determination by counting groups around atoms.
• Review functional group recognition and correct naming (alcohols, aldehydes, ketones, ethers, esters, carboxylic acids).