Chapter 1: Remembering General Chemistry – Electronic Structure and Bonding

What is Organic Chemistry?

Definition and Historical Context

Organic chemistry is the study of compounds containing carbon. Historically, organic compounds were thought to originate only from living organisms ("vital force" theory), while inorganic compounds came from minerals. The synthesis of urea from ammonium cyanate disproved this theory, showing that organic compounds can be synthesized from inorganic sources.

• Organic compounds: Contain carbon; originally thought to require a "vital force".

• Inorganic compounds: Typically derived from minerals; do not necessarily contain carbon.

• Example: Ammonium cyanate (inorganic) can be converted to urea (organic) by heating.

What Makes Carbon So Special?

Position in the Periodic Table and Electron Sharing

Carbon is unique due to its ability to share electrons, forming stable covalent bonds. In the second row of the periodic table, atoms to the left of carbon tend to give up electrons, those to the right accept electrons, while carbon shares electrons.

• Electron sharing: Enables carbon to form diverse and stable compounds.

• Versatility: Carbon can bond with many elements, leading to a vast array of organic molecules.

The Structure of an Atom

Subatomic Particles and Atomic Number

An atom consists of a nucleus (protons and neutrons) and an electron cloud. The atomic number is the number of protons in the nucleus, which defines the element.

• Protons: Positively charged

• Neutrons: No charge

• Electrons: Negatively charged

• Atomic number of carbon: 6

• Neutral carbon atom: 6 protons and 6 electrons

Isotopes

Atomic Number vs. Mass Number

All carbon atoms have the same atomic number but can have different mass numbers due to varying numbers of neutrons.

• Mass Number: $\text{Mass Number} = \text{Number of protons} + \text{Number of neutrons}$

• Isotopes: $^{12}\text{C}$, $^{13}\text{C}$, $^{14}\text{C}$ (all have 6 protons, but different neutrons)

The Distribution of Electrons in an Atom

Electron Shells and Atomic Orbitals

Electrons occupy shells around the nucleus, each with specific types and numbers of atomic orbitals. The energy of an orbital increases with distance from the nucleus.

Shell	Atomic Orbitals	Number of Orbitals	Max Electrons
First	s	1	2
Second	s, p	1, 3	8
Third	s, p, d	1, 3, 5	18
Fourth	s, p, d, f	1, 3, 5, 7	32

• First shell: Closest to nucleus, lowest energy

• Energy order within a shell: $s < p < d < f$

Electronic Configurations and Principles

Aufbau, Pauli Exclusion, and Hund's Rule

Electron configurations follow specific principles:

• Aufbau principle: Electrons fill the lowest energy orbitals first ($1s < 2s < 2p < 3s < 3p < 3d$)

• Pauli exclusion principle: Maximum two electrons per orbital

• Hund's rule: Electrons occupy empty degenerate orbitals before pairing

Formation of Ions and Covalent Bonds

Electron Loss and Gain

Atoms achieve stability by filling or emptying their outer electron shells:

• Atoms in Group 1 (e.g., Li, Na): Lose electrons to form cations

• Atoms in Group 17 (e.g., F, Cl): Gain electrons to form anions

• Hydrogen: Can lose or gain an electron to form $H^+$ (proton) or $H^-$ (hydride ion)

Achieving Filled Shells by Sharing Electrons

Atoms can also achieve filled shells by sharing electrons, forming covalent bonds.

• Covalent bond: A bond formed by sharing electrons

• Example: $H_2$, $F_2$, $HCl$

How Many Bonds Does an Atom Form?

Bonding Patterns of Common Elements

• Oxygen: Forms 2 covalent bonds (e.g., in water)

• Nitrogen: Forms 3 covalent bonds (e.g., in ammonia)

• Carbon: Forms 4 covalent bonds (e.g., in methane)

• Hydrogen: Forms 1 covalent bond

Nonpolar and Polar Covalent Bonds

Electronegativity and Bond Polarity

The polarity of a covalent bond depends on the difference in electronegativity between the bonded atoms.

• Nonpolar covalent bond: Electronegativity difference < 0.5 (e.g., $H-H$, $C-H$)

• Polar covalent bond: Electronegativity difference 0.5–1.9 (e.g., $H-Cl$, $H-O$)

• Ionic bond: Electronegativity difference > 1.9 (e.g., $Na^+Cl^-$)

Dipole Moment

Definition and Examples

Dipole moment quantifies the polarity of a bond:

• Formula: $\text{Dipole moment} = \text{size of charge} \times \text{distance between charges}$

• Greater electronegativity difference: Larger dipole moment, more polar bond

Bond	Dipole Moment (D)	Bond	Dipole Moment (D)
H–C	0.4	C–C	0
H–N	1.3	C–N	0.2
H–O	1.5	C–O	0.7
H–F	1.7	C–F	1.6
H–Cl	1.1	C–Cl	1.5
H–Br	0.8	C–Br	1.4
H–I	0.4	C–I	1.2

Electrostatic Potential Maps

Visualizing Charge Distribution

Electrostatic potential maps use color to show regions of positive and negative charge in molecules. Red indicates negative potential (attracts positive charge), blue indicates positive potential (attracts negative charge).

Lewis Structures and Formal Charge

Drawing and Interpreting Lewis Structures

Lewis structures represent molecules by showing all valence electrons as dots or lines. Formal charge helps identify the most stable structure.

• Formal Charge Formula: $\text{Formal Charge} = \text{Number of valence electrons} - (\text{Number of lone-pair electrons} + \text{Number of bonds})$

• Example: Water, hydronium ion, hydroxide ion, hydrogen peroxide

Bonding Patterns and Charges

Common Elements and Their Bonding

• Carbon: Forms 4 bonds; if not, it is charged or a radical (carbocation, carbanion, methyl radical)

• Nitrogen: Forms 3 bonds and has one lone pair; otherwise, it is charged (ammonium ion, amide anion)

• Oxygen: Forms 2 bonds and has two lone pairs; otherwise, it is charged

• Hydrogen and Halogens: Form 1 bond and have three lone pairs (halogens); otherwise, they are charged or radicals

Lewis Structures: Bonds and Lone Pairs

Bonding and Lone Pair Relationships

The sum of the number of bonds and lone pairs for second-row elements (C, N, O, F) always equals four.

• Carbon: 4 bonds, 0 lone pairs

• Nitrogen: 3 bonds, 1 lone pair

• Oxygen: 2 bonds, 2 lone pairs

• Fluorine: 1 bond, 3 lone pairs

Drawing Lewis Structures

Steps and Examples

• Count total valence electrons (add for negative charge)

• Avoid O–O bonds unless necessary

• Check for formal charges

• Example: $NO_3^-$ (nitrate ion), $CO_3^{2-}$ (carbonate ion)

Kekulé, Condensed, and Skeletal Structures

Different Ways to Represent Organic Molecules

• Kekulé structures: Show all atoms and bonds explicitly

• Condensed structures: Group atoms together, omitting some bonds

• Skeletal structures: Show carbon-carbon bonds as lines; carbons and hydrogens bonded to carbons are not shown

Kekulé Structure	Condensed Structure
CH3CH2CH2CH3	CH3(CH2)2CH3
CH3CH(CH3)CH2CH3	CH3CH(CH3)CH2CH3
CH3CH2CH2OH	CH3CH2CH2OH

Atomic Orbitals

s and p Orbitals

Atomic orbitals are regions of space where electrons are most likely to be found. s orbitals are spherical; p orbitals have two lobes with opposite phases.

• 1s, 2s orbitals: Spherical, differ in size and electron density

• p orbitals: Three orientations (x, y, z), each with two lobes

Wave Nature of Electrons

Standing Waves and Orbital Phases

Electrons behave as standing waves, with nodes and phases. Constructive interference leads to bonding; destructive interference leads to antibonding.

Molecular Orbital Theory

Formation of Molecular Orbitals

Atomic orbitals combine to form molecular orbitals. The number of molecular orbitals equals the number of atomic orbitals combined.

• Sigma (σ) bond: Formed by head-on overlap of orbitals

• Pi (π) bond: Formed by side-to-side overlap of p orbitals

Hybridization and Molecular Geometry

sp3, sp2, and sp Hybridization

Hybridization explains molecular shapes and bond angles:

• sp3 hybridization: Tetrahedral geometry, bond angle 109.5° (e.g., methane)

• sp2 hybridization: Trigonal planar geometry, bond angle 120° (e.g., ethene)

• sp hybridization: Linear geometry, bond angle 180° (e.g., ethyne)

Bond Strengths and Angles

Factors Affecting Bond Properties

• More s character: Shorter and stronger bonds, larger bond angles

• Bond strength: Double bonds are stronger than single bonds, but weaker than triple bonds

• Bond length: Shorter bonds are stronger

Summary: Hybridization, Bond Lengths, Bond Strengths, and Bond Angles

• sp3: Tetrahedral, 109.5°, longest and weakest C–H bonds

• sp2: Trigonal planar, 120°, intermediate bond length and strength

• sp: Linear, 180°, shortest and strongest C–H bonds

*Additional info: Some explanations and examples have been expanded for clarity and completeness, including the relationships between hybridization and molecular geometry, and the principles of electron configuration.*