General Fundamentals of Organic Chemistry

What Makes Organic Chemistry Special?

Organic chemistry is the chemistry of carbon compounds, especially those containing carbon–hydrogen (C–H) bonds. It focuses on:

• Structures of carbon-based molecules
• How these structures determine properties
• Typical ways organic molecules react and transform

In this chapter, you get the basic “toolkit” needed for all later organic chemistry topics: how we write formulas, draw structures, and think about bonds and electron distribution in organic molecules.

(Details on specific reaction types, functional groups, and particular compound classes appear in later chapters.)

Carbon as the Central Element

Carbon is unique because:

• It forms four covalent bonds (it is tetravalent).
• It can bond to itself repeatedly, giving:
• Chains (linear or branched)
• Rings (small or large)
• Networks (very extended structures)
• It forms strong, stable bonds with many elements: H, O, N, halogens (F, Cl, Br, I), S, P, etc.

The most important bonding patterns in basic organic compounds are:

• C–C single bonds
• C=C and C≡C multiple bonds
• C–H bonds
• C–X bonds, where X is often O, N, halogen, S, etc. (functional groups)

Names, Formulas, and Structures (Overview)

Later chapters treat naming and structures in detail. Here you just need the basic idea that one and the same molecule can be represented in several ways that carry different amounts of structural information.

Types of formulas

1. Molecular formula

Shows only the number of each type of atom:

• Methane: $ \mathrm{CH_4} $
• Ethanol: $ \mathrm{C_2H_6O} $

It does not show how atoms are connected.

2. Structural formula

Shows the connectivity (which atom is bonded to which):

• Methane: $ \mathrm{H - C - H} $ (often written with 4 H around C)
• Ethanol: $ \mathrm{CH_3 - CH_2 - OH} $

3. Condensed structural formula

Compresses groups of atoms but preserves order:

• Ethane: $ \mathrm{CH_3CH_3} $
• Propane: $ \mathrm{CH_3CH_2CH_3} $
• 2-Propanol: $ \mathrm{(CH_3)_2CHOH} $

4. Skeletal (line) formula
• Carbon atoms are implied at the ends and intersections of lines.
• Hydrogen atoms bonded to carbon are usually omitted.
• H atoms attached to heteroatoms (O, N, etc.) are drawn explicitly.

This is the standard representation in organic chemistry because it is quick and emphasizes the “carbon skeleton” and functional groups.

Structural information vs. formula information

The same molecular formula can correspond to many different possible structures. For example, $ \mathrm{C_4H_{10}} $ can be:

• A straight chain of four carbons
• A branched chain (three carbons in a main chain with one as a side branch)

This idea (different structures, same formula) underlies the concept of isomerism, treated in its own chapter.

How We Draw Organic Molecules

Bond-line conventions

In skeletal formulas:

• Lines represent bonds between carbon atoms.
• Each vertex (corner) or line end represents a carbon atom (unless another element symbol is written).
• Hydrogens on carbon are not shown; you assume enough H to give each C four bonds.
• Other atoms (O, N, halogens, S, etc.) are written explicitly with their H atoms.

Example: Ethanol

• Condensed: CH3CH2OH
• Skeletal: a two-carbon zig-zag line ending in an OH group

Three-dimensional representation (basic idea)

Organic reactions and properties often depend on the 3D arrangement of atoms.

Common 3D drawing conventions:

• Solid wedge: bond coming out of the plane toward you
• Dashed wedge: bond going behind the plane
• Simple line: bond in the plane of the paper

For instance, a carbon atom with four different substituents can be drawn with one wedge, one dashed wedge, and two plain lines to show a tetrahedral geometry.

(Full treatment of stereoisomerism appears under “The Concept of Isomerism”.)

Bonds and Electron Distribution in Organic Molecules

Covalent bonding and bond types

In organic molecules, covalent bonding is dominant. Between carbon atoms you typically see:

• Single bonds: sigma ($\sigma$) bonds
• Double bonds: one $\sigma$ + one pi ($\pi$) bond
• Triple bonds: one $\sigma$ + two $\pi$ bonds

Multiple bonds (double, triple) are:

• Shorter and stronger than single bonds
• More electron-rich and thus often more reactive

Hybridization (qualitative view)

To understand shapes and bond angles in organic molecules, the concept of hybridization is widely used. Without going into quantum-mechanical detail, you should know these typical patterns for carbon:

• sp³ hybridization
• Four $\sigma$ bonds
• Tetrahedral geometry
• Approximate bond angle: $109.5^\circ$
• Example: methane $ \mathrm{CH_4} $, saturated carbon atoms in alkanes
• sp² hybridization
• Three $\sigma$ bonds + one $\pi$ bond (part of a double bond)
• Trigonal planar geometry
• Approximate bond angle: $120^\circ$
• Example: carbon in a $ \mathrm{C=C} $ double bond, as in ethene
• sp hybridization
• Two $\sigma$ bonds + two $\pi$ bonds (part of a triple bond, or two double bonds in line)
• Linear geometry
• Approximate bond angle: $180^\circ$
• Example: carbon in acetylene (ethyne), $ \mathrm{HC \equiv CH} $

Hybridization connects directly to the geometry of organic molecules, which in turn influences reactivity.

Polar and nonpolar bonds

Because different atoms have different electronegativities, bonds in organic molecules can be:

• Nonpolar covalent: roughly equal sharing of electrons
• Example: C–C in saturated hydrocarbons
• Weakly polar: small difference in electronegativity
• Example: C–H (slightly polarized)
• Strongly polar: large difference in electronegativity
• Example: C–O, O–H, C–Cl, C–N

Polar bonds create partial charges:

• More electronegative atom: partial negative charge ($\delta-$)
• Less electronegative atom: partial positive charge ($\delta+$)

These partial charges are central to understanding organic reaction mechanisms, because they define where electron-rich and electron-poor sites are.

Functional Groups (First Glimpse)

A functional group is a specific arrangement of atoms that:

• Defines the class of an organic compound
• Determines typical chemical behavior (reactivity pattern)

Examples (detailed treatment appears in later chapters):

• Hydroxyl group $(-\mathrm{OH})$: characterizes alcohols
• Carbonyl group $(\mathrm{C=O})$: present in aldehydes and ketones
• Carboxyl group $(-\mathrm{COOH})$: characterizes carboxylic acids
• Amino group $(-\mathrm{NH_2})$: basic unit in amines and amino acids
• Halogen substituents: $-\mathrm{Cl}$, $-\mathrm{Br}$, etc., in haloalkanes

Important ideas:

• Molecules are often described as a carbon skeleton plus attached functional groups.
• Most organic reactions occur at or near functional groups, not on “plain” C–C or C–H bonds.

Organic Molecules as Electron Systems

Electron-rich and electron-poor centers

Because of bond polarity and functional groups, organic molecules have positions that can be:

• Nucleophilic (electron-rich, “nucleus-seeking”)
• Tend to donate an electron pair
• Often negatively charged or with a lone pair on a relatively electronegative atom
• Electrophilic (electron-poor, “electron-seeking”)
• Tend to accept an electron pair
• Often positively charged or with a partial positive charge

Examples (qualitative):

• Carbon of a carbonyl group (C=O) is electrophilic ($\delta+$).
• Oxygen of a carbonyl group is nucleophilic ($\delta-$).
• Halide ions ($\mathrm{Cl^-}$, $\mathrm{Br^-}$, etc.) are nucleophilic.
• Carbocations (positively charged carbon centers) are electrophilic.

This donor–acceptor view of bonds and reaction centers is the foundation for understanding organic reaction mechanisms.

Formal charges and lone pairs

To reason about electron distribution, you often:

• Count valence electrons for each atom
• Assign electrons in bonds and lone pairs
• Determine formal charge (not necessarily the real charge, but a bookkeeping tool)

Rules (for common atoms in organic chemistry):

• Carbon normally forms 4 bonds and has no lone pairs → formal charge 0
• Nitrogen typically: 3 bonds + 1 lone pair → formal charge 0
• Oxygen typically: 2 bonds + 2 lone pairs → formal charge 0
• Halogens typically: 1 bond + 3 lone pairs → formal charge 0

If an atom deviates from its usual bonding pattern, it often carries a formal charge and behaves as a strongly nucleophilic or electrophilic center.

Resonance (Delocalization) – Conceptual Introduction

Many organic molecules have $\pi$ electrons or lone pairs that are not confined to a single bond but spread over several atoms. This phenomenon is called delocalization, and we often describe it with resonance structures.

Key points:

• A resonance structure is one of several valid Lewis structures that differ only in electron placement (not atom positions).
• The real molecule is a hybrid of all reasonable resonance structures.
• Delocalization usually:
• Stabilizes the molecule (lower energy)
• Influences bond lengths and reactivity

Common examples in organic chemistry:

• Conjugated double bonds (e.g., in dienes)
• Aromatic systems (e.g., benzene)
• Carboxylate ions (delocalization of negative charge over two oxygens)

Aromaticity and the detailed consequences of resonance are treated more fully in chapters on aromatic hydrocarbons and electronic effects.

Organic Reaction Participants: Reagents and Substrates (Overview)

Organic reactions usually involve:

• A substrate: the molecule being transformed (e.g., an alkane, alkene, alcohol)
• A reagent: the chemical species used to bring about the change

Reagents can be:

• Nucleophilic (electron-pair donors)
• Electrophilic (electron-pair acceptors)
• Radical (species with unpaired electrons)
• Acidic or basic (in the Brønsted or Lewis sense)

Later chapters will classify and exemplify these (see “Reagents, Substrates, and Reactions” and “Reaction Types in Organic Chemistry”). Here it is enough to recognize the general roles.

Fundamental Ideas of Organic Reaction Types

Although detailed reaction types are treated in their own section, organic reactions broadly fall into a few recurring patterns, all based on the movement of electrons:

• Breaking and forming covalent bonds
• Reorganization of electron density
• Stabilization of charges or radicals

Common pattern ideas (only as orientation):

• Substitution: one group replaces another on a carbon atom.
• Addition: atoms or groups add across a multiple bond.
• Elimination: atoms or groups are removed, creating a double or triple bond.
• Rearrangement: the carbon skeleton or positions of substituents change internally.

These patterns will be repeatedly revisited in the context of specific functional groups.

Relationship Between Structure and Properties (Introductory View)

Organic chemists constantly relate structure to properties such as:

• Boiling and melting points
• Solubility (e.g., in water vs. in organic solvents)
• Acidity and basicity
• Reactivity toward specific reagents

Some general trends you will see repeatedly:

• Increasing chain length (more carbons) often increases boiling point (more surface area, stronger intermolecular interactions).
• Branching in the carbon chain often lowers boiling point (less surface area).
• Polar functional groups increase:
• Intermolecular attractions (e.g., hydrogen bonding)
• Solubility in polar solvents like water (up to a point)
• Electron-withdrawing groups (e.g. $-\mathrm{NO_2}$, $-\mathrm{CN}$, $-\mathrm{COOH}$) and electron-donating groups (e.g. $-\mathrm{OH}$, $-\mathrm{OCH_3}$, $-\mathrm{NH_2}$) systematically modify electron density and thereby reactivity (treated in detail under “Electronic Effects in Organic Compounds”).

Understanding such connections is the central goal of organic chemistry: from structure to properties and reactivity.

How to Read and Think in Organic Chemistry

To work effectively with organic chemistry, it helps to develop a certain way of thinking:

1. Identify the carbon skeleton
• How many carbons?
• Are they in a chain, ring, or multiple rings?
• Is the chain branched or unbranched?
2. Locate functional groups
• What types of functional groups are present?
• How many of each?
• Are they conjugated (connected by alternating single and double bonds)?
3. Assess electronic features
• Where are polar bonds?
• Where might partial charges or formal charges reside?
• Which atoms have lone pairs?
4. Predict likely reactive sites
• Where are nucleophilic centers (electron-rich)?
• Where are electrophilic centers (electron-poor)?
• Are there conjugated or aromatic systems that might affect stability and reactivity?
5. Connect to known patterns
• Which general class of reactions is likely (substitution, addition, elimination, etc.)?
• How might a given reagent interact with the functional groups present?

These habits will be refined as you study specific types of organic compounds and reactions in later chapters.