Lecture Notes #1

Course Introduction

Welcome and Course Structure

This course, Organic Chemistry I (2311), introduces students to the foundational principles of organic chemistry, focusing on the structure and bonding of organic molecules. The course is organized into several units, beginning with the basics of molecular structure and progressing to reactivity and three-dimensional visualization.

• Required Textbook: Organic Chemistry, 9th edition, Leroy G. Wade and William Simek

• Recommended Materials: Molecular model kit, solutions manual

Unit 1: Learning the Organic Language

Course Learning Goals

• Reading, understanding, and speaking the language of organic chemistry – e.g., naming molecules such as ethane ($\text{Me-Me} = \text{H}_3\text{C}-\text{CH}_3$)

• Visualizing molecules in 3D – understanding spatial arrangement and molecular geometry

• Predicting and evaluating reactivity – determining how, where, and in what order reactions occur, and what products form

Structure and Bonding

What is Organic Chemistry?

Organic chemistry is the study of carbon-containing molecules. Carbon's versatility allows for a vast array of molecular structures and reactivities, making it central to biological and chemical processes.

• Example: Retinal (in vision) – sets off a reaction cascade that allows you to see

• Example: Caffeine (in coffee, tea) – affects biological processes such as stomach acidity

Bonding Basics

• Ionic Bonds: Formed by electron transfer, held together by electrostatic interactions. Example: $\text{NaCl} \rightarrow \text{Na}^+ :\text{Cl}^-:$ Key Point: The larger the electronegativity difference, the more ionic the bond.

• Covalent Bonds: Electrons are shared between atoms. Example: $\text{CH}_4\text{O}$ (methanol) – simplest visualization is the Lewis structure.

Lewis Structures: The Guide

Steps to Drawing Lewis Structures

1. Determine the total number of valence electrons ($ve^-$) in the molecule. Formula: Add all valence electrons per atom, adjust for charge (+1 per cation, -1 per anion).

2. Draw a skeletal structure – place the least electronegative element in the center (except H), arrange atoms around it.

3. Draw single bonds from the center atom to terminal atoms.

4. Fill in lone pairs on terminal atoms until octet is complete or all electrons are used up.

5. Check every atom for octet (or duet for H). If atoms need an octet, use lone pairs from terminal atoms to share.

6. Calculate formal charge: Formula: $\text{FC} = v - b - d$ where $v$ = valence electrons, $b$ = bonds, $d$ = lone pairs

Example: Methanol ($\text{CH}_4\text{O}$)

1. Count valence electrons: $4 + 2(1) + 6 = 12 ve^-$

2. Draw initial structure with single bonds

3. Assign lone pairs to oxygen

4. Check octet for carbon and oxygen

5. Calculate formal charges for each atom

Example: Formate Ion ($\text{HCO}_2^-$)

1. Count valence electrons: $4 + 1 + 2(6) + 1 = 18 ve^-$

2. Draw initial structure

3. Assign lone pairs to oxygen

4. Check octet for carbon and oxygen

5. Calculate formal charges

Bond Preferences of Common Elements

How Many Bonds Do Atoms Like?

Element	# Bonds	# Lone Pairs	Example	Additional Info
H	1	0	-H	Never more than 1 bond
C	4	0	-C-	Never make 5 bonds
N	3	1	-N-	3 bonds, 1 lone pair
O	2	2	-O-	2 bonds, 2 lone pairs
F, Cl, Br, I	1	3	-F-	1 bond, 3 lone pairs

Expanded Octets

Element	# Bonds	# Lone Pairs	Example
P	3, 5	1, 0	Phosphorus structures
S	2, 4, 6	2, 1, 0	Sulfur structures

Practice and Application

Lewis Structure Practice

Given the formula CH4O, which is the correct Lewis structure?

• Correct Answer: Structure B (as shown in the notes)

Additional Info

• Atoms with the same number of valence electrons tend to act similarly in bonding.

• Organic conventions: Charges of +1 are often shown, -1 is implied.

• Expanded octets are possible for elements in period 3 and beyond (e.g., P, S).