Chapter 1: The Basics (Organic Chemistry)

The Basics of Organic Chemistry

• Organic chemistry focuses on compounds that contain carbon; inorganic chemistry largely excludes carbon except for exceptions.

• Why carbon? Because carbon is central to living organisms, forms strong bonds with itself and with elements like H, N, O, and S, and enables vast numbers of compounds due to versatile bonding.

Basic Atomic Structure

• Nucleus: dense, positively charged; contains protons (Z) and neutrons (N).

• Electron cloud surrounds the nucleus and matches the number of protons in a neutral atom.

• Isotopes: same number of protons (Z) but different numbers of neutrons; atomic mass units defined as 1/12 the mass of a reference isotope (usually carbon-12).

• The periodic table lists atomic mass units (amu) for isotopes, e.g., C has a standard mass of approximately 12.011 amu.

• Question to consider: Which isotope is most prevalent and why? (discussion point: abundance depends on nuclear stability and natural abundance)

Basic Bonding Concepts

• Octet Rule: valence electrons occupy the outermost shell; seeking noble-gas configuration.

  • Ionic bonds: transfer of electrons from one atom to another.

  • Covalent bonds: sharing of electrons.

  • Most main-group elements tend to achieve an octet (8 electrons) in their valence shell; some use a duet (2 electrons) for hydrogen and sometimes boron/others.

• Why atoms form bonds: to reach stable electron configurations similar to noble gases.

Electronegativity and Ionic vs Covalent Bonding

• Electronegativity is the ability of an atom to attract electrons in a bond.

• Large difference in electronegativity between atoms ➔ ionic bond (formation of salts) and noble-gas configurations.

• Ionic bonding often leads to high melting solids and characteristic properties.

• Common trend: electronegativity increases across a period from left to right and decreases down a group.

• Example context (general): Hydrogen is less electronegative than halogens; fluorine is among the most electronegative elements.

Representing Covalent Bonds

• Covalent bonds arise when electronegativity difference is small; bonds are shown as lines in Lewis/dash formulas.

• Carbon prefers to form four bonds (C likes to have 4 bonds).

Representing Multiple Covalent Bonds

• Double bonds and triple bonds are shown with multiple lines or with dash/dot representations.

• Examples:

  • Ethene (C2H4): C=C with each carbon bonded to two hydrogens.

  • Ethyne (C2H2): C≡C with each carbon bonded to one hydrogen.

• Dash/dot representations: illustrate bond order and connectivity.

Covalent Bonding: Ammonium Ion

• Ammonium ion (NH4+) can be drawn with a central N bonded to four H atoms and a formal positive charge indicate on nitrogen:

  • Lewis/dash representation shows H–N(–H)4 with a formal charge of + on N in appropriate resonance conditions.

Lewis Structures: Building a Lewis Diagram

• Step-by-step approach:

  • Determine valence electrons for each atom (using group number/HONC rule for bonds made).

  • Sum valence electrons; adjust for any ionic charge (add electrons for anions, subtract for cations).

  • Connect bonded atoms with lone electron pairs; distribute remaining electrons to satisfy octet rule around each atom.

  • If octets are not satisfied, form additional bonds (double/triple) as needed to satisfy octet.

Formal Charge Calculations

• Formal charge formula: $F = V - rac{1}{2}S - U$

  • Where V = number of valence electrons of the atom, S = number of electrons in bonds (shared electrons, counted as half for the atom), U = number of unshared electrons (lone pairs) on the atom.

• Example question: what are the formal charges for each oxygen atom in carbonate ion (CO3^{2-})? (Think in terms of valence, shared vs unshared electrons, and total charge adjustment.)

• Governing principle: a neutral molecule distributes charges to minimize formal charges; resonance often distributes charge.

A Summary of Formal Charges and Bonding Patterns

• Become familiar with common bonding patterns and typical formal-charge distributions.

Constitutional Isomers

• Constitutional isomers of a given molecular formula (example: C3H6O) differ in connectivity/ bonding pattern.

• These isomers have different physical properties (boiling point, melting point, density) and chemical properties (reactivity).

Structural Formulas for 1-Propanol

• Multiple representational forms for the same molecule include:

  • Ball-and-stick model

  • Electron-dot (Lewis) structure

  • Dash (bond-line) formula

  • Condensed formula (fully/partially condensed)

  • Observing actual connectivity: CH3-CH2-CH2-OH (1-propanol)

• Each representation highlights different aspects (geometry vs. connectivity vs. energy states).

Equivalent Dash Formulas and Rotations

• Dash formulas substitute lines for bonding electron pairs; non-bonding electrons may still be shown.

• The same structure can be represented by different dash formulas depending on rotation about single bonds.

• Hand-held models help visualize three-dimensionality and alternate conformations.

Writing Condensed Formulas

• Partially condensed: hydrogens attached to a carbon are written directly after that carbon (e.g., CH3).

• Fully condensed: all atoms attached to a carbon appear directly after that carbon; hydrogens on carbons are often implicit.

• Parentheses indicate branching groups; subscripts denote multiple identical groups.

Bond-Line (Skeletal) Structures

• Most common for larger molecules due to speed of drawing.

• Rules:

  • Each line represents a bond.

  • Each bend or endpoint represents a carbon atom; assume carbon has four bonds unless shown otherwise.

  • Hydrogens attached to carbons are not shown (assumed).

  • H atoms on heteroatoms (O, N, S, etc.) are shown explicitly.

• Important to be able to convert between skeletal, condensed, and Lewis structures.

Drawing Bond-Line Structures: Practical Rules

• Carbons are shown as zig-zags unless in rings or triple bonds.

• Hydrogens on carbons are not shown.

• The hydrogen on oxygen is shown explicitly when relevant.

• Carbons are not labeled in the drawing.

Additional Bond-Line Structures and Practice

• Practice with model kits for:

  • Cycloalkanes: ring structures (bond-line representations)

  • Alkenes: double bonds (C=C) with appropriate geometry

  • Alkynes: triple bonds (C≡C) with linear geometry

Ideal Tetrahedral Geometry (Methane as Example)

• Tetrahedral arrangement around carbon arises from sp3 hybridization; bond angles close to 109.5°.

• Bond representation:

  • Two bonds in the plane (solid lines).

  • Two bonds away from the plane (solid wedge and hashed wedge).

• C with a double bond is trigonal planar (120°) and C with a triple bond is linear (180°).

• Energy context: the nuclear repulsion vs covalent bonding energy creates the characteristic geometry.

Three-Dimensional Representations in Bond-Line Form

• Bond-line diagrams can imply 3D geometry when combined with wedge/dash conventions.

• Examples show trigonal planar and linear arrangements.

Resonance Structures

• Resonance structures are alternate Lewis structures for the same compound that differ only in electron placement, not in atom positions.

• Electrons in lone pairs or in bonds can be moved using curved arrows to show resonance.

• Real molecule is a resonance hybrid, a weighted average of contributors.

• Key notes:

  • The double-headed arrow indicates relationship, not a chemical equilibrium.

  • All resonance structures must be valid Lewis structures with correct valence and obey octet rules where applicable.

Drawing and Interpreting Resonance

• Start with a source of electrons and move electrons to where they can go; assign formal charges to each structure.

• All resonance forms contribute to the actual molecular structure (hybrid), but individual resonance forms are not equivalent in energy.

Interconversion of Nitrate Resonance Forms (NO3−)

• In nitrate, the number of bonds to N and the net charge are conserved across resonance forms.

• The actual structure shows delocalized negative charge over the oxygens (resonance hybrid).

Resonance Hybrids

• The resonance hybrid depicts charge delocalization and electron distribution more accurately than any single Lewis structure.

Quantum Mechanics: Wave Functions and Atomic Orbitals

• Electrons behave as both particles and waves; wave functions (psi) describe energy states.

• The probability of finding an electron at a point in space is given by the square of the wave function's magnitude.

• Solutions can have positive, negative, or zero phase; phase indicates sign relative to a chosen reference.

• Atomic orbital wave functions give probability distributions in 3D space, yielding orbital shapes where electrons are most likely found (~90–95% of the time).

Covalent Bonding: A Goldilocks Distance

• Energy diagram concept:

  • Nuclear repulsion rises at very short distances.

  • Electron-nucleus attraction lowers energy when orbitals overlap.

  • An optimal internuclear distance r exists where bonding is strongest (approximately 0.74 Å for C–C single bonds).

  • Bond energy at this distance is substantial (e.g., about $E_ ext{bond} \,\approx\, 436\ \text{kJ mol}^{-1}$ for the C–C single bond shown in the diagram).

Constructive and Destructive Interference in Waves

• Constructive interference: wave crests/troughs align, increasing amplitude.

• Destructive interference: wave crests and troughs cancel, reducing amplitude.

• Notation: wavelength $\lambda$, amplitude $a$; interference patterns depend on phase relationships.

Bonding and Antibonding Molecular Orbitals

• Constructive (in-phase) overlap leads to bonding orbitals with electron density between atoms.

• Destructive (out-of-phase) overlap leads to antibonding orbitals with reduced or no density between atoms.

• Bonding can occur when orbitals overlap constructively; antibonding destabilizes the molecule if occupied.

Formation of Molecular Orbitals from Atomic Orbitals

• Atomic orbitals combine to form molecular orbitals when electrons occupy them; this can involve interactions that are considered in excited states when energy is supplied.

Electron Configuration, Aufbau, Pauli, and Hund’s Rules

• Aufbau principle: fill lowest energy orbitals first (1s, 2s, 2p, …).

• Pauli exclusion principle: each orbital holds at most two electrons with opposite spins.

• Hund’s rule: electrons fill degenerate orbitals singly before pairing, to maximize total spin.

Carbon Electronic Configuration and Hybridization Context

• Ground-state configuration and energy diagram for carbon; excited-state configurations illustrate potential promotions between levels.

• sp3 hybridization (tetrahedral) mixes 25% s and 75% p to form four equivalent orbitals.

• sp2 hybridization (trigonal planar) mixes 33% s and 67% p to form three equivalent orbitals with one unhybridized p orbital.

• sp hybridization (linear) mixes 50% s and 50% p to form two equivalent orbitals with two remaining unhybridized p orbitals.

Tetrahedral Geometry and Sigma Bonding in Methane

• Methane (CH4) displays tetrahedral geometry; the C–H bonds arise from overlap of carbon sp3 orbitals with hydrogen 1s orbitals.

• Bonding MO perspective shows head-on overlap leading to strong sigma bonds.

Sigma Bonding in Ethane

• Ethane features sigma (σ) bonds formed by end-to-end overlap of sp3 hybridized orbitals from adjacent carbons and from carbon-hydrogen interactions.

• Question to consider: what other sigma bonds are formed from sp3 orbitals?

sp2 Hybridization and Ethene (Alkene) Bonding

• Ethene (C2H4) is trigonal planar around each carbon due to sp2 hybridization.

• Bonding orbitals involve in-plane σ bonds and unhybridized p orbitals that participate in π bonding.

• Which orbitals compose the C=C bond? Answer: one σ bond from sp2/sp2 overlap and one π bond from lateral overlap of unhybridized p orbitals.

Pi Bonding in Ethene

• Pi (π) bonding arises from sideways overlap of p orbitals; it is weaker than the σ bond.

• A double bond consists of one σ and one π bond.

• Electron count in π bond: 2 electrons reside in the π bond; the σ bond also holds 2 electrons.

Rotation About the C=C Double Bond and Stereoisomerism

• Rotation about the carbon–carbon double bond would disrupt the π overlap, requiring energy to break the bond; this creates a barrier to rotation.

• Restricted rotation around C=C leads to stereoisomerism (cis/trans or E/Z configurations) when there are different substituents around the double bond.

sp Hybridization and Ethyne (Acetylene)

• Ethyne is linear with sp hybridization for each carbon.

• Bonding orbitals: a σ bond from sp–sp overlap and two π bonds from two sets of side-on p orbital overlaps, giving a triple bond: 1 σ + 2 π.

• Number of electrons: the triple bond contains 4 bonding electrons in π bonds and 2 in the σ bond.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

• VSEPR considers central atoms bonded to two or more atoms/groups; counts all valence electron pairs (bonding and nonbonding).

• Electron pairs repel to stay as far apart as possible; nonbonding pairs exert greater repulsion than bonding pairs.

VSEPR Geometries and Example Angles

• Tetrahedral (4 electron pairs, 0 nonbonding): 109.5°; examples: CH4, NH4+.

• Trigonal pyramidal (3 bonding, 1 nonbonding): ~107°; examples: NH3, CH3− attached species.

• Bent (2 bonding, 2 nonbonding): ~105°.

• Trigonal planar (3 electron pairs, 0 nonbonding): 120°; examples: BF3, CH3− groups.

• Linear (2 electron pairs, 0 nonbonding): 180°; examples: BeH2, CO2 (for the linear arrangement around the central atom).

Shapes of Molecules and Ions from VSEPR Theory

• Table-style summary (examples):

  • 2 bonding, 0 nonbonding: Linear, BeH2, CO2, sp

  • 3 bonding, 0 nonbonding: Trigonal Planar, BF3, CH3

  • 4 bonding, 0 nonbonding: Tetrahedral, CH4, NH4+

  • 3 bonding, 1 nonbonding: Trigonal Pyramidal, NH3, CH3

  • 2 bonding, 2 nonbonding: Angular (Bent), H2O

Quick Reference: Key Equations and Numbers

• Octet Rule reference: target 8 valence electrons around most main-group atoms; duet for H.

• Formal charge: $F = V - frac{1}{2}S - U$

• Bond types:

  • Single bond: one σ bond

  • Double bond: one σ + one π

  • Triple bond: one σ + two π

• Bond length reference (example): $r \approx 0.74\ \text{Å}$ for C–C single bond.

• Bond energy example: $E_ ext{bond} \approx 436\ \text{kJ mol}^{-1}$ for illustrative C–C bond context.

• Hybridization patterns:

  • sp3: 25% s, 75% p

  • sp2: 33% s, 67% p

  • sp: 50% s, 50% p

• Angles associated with common geometries: 109.5°, 107°, 105°, 120°, 180° (approximate values shown in the slides).

Practical Takeaways for Study and Exams

• Be able to convert between Lewis structures, condensed formulas, and bond-line skeletal structures.

• Recognize and predict VSEPR geometries based on the number of bonding and lone pairs.

• Understand the concept of resonance and be able to draw resonance forms and a resonance hybrid.

• Identify when π bonds restrict rotation and how this leads to cis/trans isomerism.

• Distinguish between sigma and pi bonding and how orbital overlap creates covalent bonds.

• Apply the Aufbau, Pauli, and Hund rules to predict ground-state electronic configurations and how hybridization arises from electron promotion.

• Memorize representative bond angles and basic geometries associated with sp3, sp2, and sp hybridizations for quick problem solving.

• Recognize common constitutional isomers for a given formula (e.g., C3H6O) and how connectivity changes properties.

Note

• Specific charts and visual diagrams from the slides (electronegativity scales, 2D vs 3D drawings, and model-kit practice pictures) are valuable for quick recall and should be reviewed alongside these notes to reinforce visual understanding and spatial intuition.

The Basics of Organic Chemistry

• Organic chemistry is carbon-focused due to carbon's central role in living organisms and its versatile bonding with itself and other elements (H, N, O, S).

Basic Atomic Structure & Bonding Concepts

• Atoms consist of a nucleus (protons, neutrons) and an electron cloud (electrons equal to protons in a neutral atom).

• Isotopes have the same protons but different neutrons.

• Bonds form to achieve stable electron configurations (octet/duet rule).

• Ionic bonds (electron transfer) form with large electronegativity differences, leading to salts.

• Covalent bonds (electron sharing) form with small electronegativity differences, represented by lines in Lewis structures. Carbon typically forms four bonds.

• Electronegativity: An atom's ability to attract electrons in a bond, generally increasing across a period and decreasing down a group.

Representing Molecules

• Multiple Covalent Bonds: Double (C=C) and triple (C

• Lewis Structures: Drawn by summing valence electrons, connecting atoms, distributing remaining electrons to satisfy octets, and forming multiple bonds if necessary.

• Formal Charge: $F = V - \frac{1}{2}S - U$ helps evaluate electron distribution; neutral molecules minimize formal charges.

• Constitutional Isomers: Compounds with the same molecular formula but different atomic connectivity, resulting in different physical and chemical properties.

• Structural Formulas: Can be represented as electron-dot, dash (bond-line), or condensed formulas. Bond-line structures are common for larger molecules, where carbons and their hydrogens are implied at bends/endpoints, but hydrogens on heteroatoms are shown.

• Equivalent Dash Formulas and Rotations: Single bonds allow rotation, leading to different conformations. Double bonds restrict rotation due to $\pi$ overlap.

Three-Dimensional Geometry & Hybridization

• VSEPR Theory: Electron pairs repel, determining molecular geometry. Nonbonding pairs exert greater repulsion.

• Hybridization explains molecular geometry:

  • sp3 (tetrahedral): 4 equivalent orbitals, 109.5° angles (e.g., CH4).

  • sp2 (trigonal planar): 3 equivalent orbitals + 1 unhybridized p, 120° angles (e.g., C2H4).

  • sp (linear): 2 equivalent orbitals + 2 unhybridized p, 180° angles (e.g., C2H2).

• Sigma ($\sigma$) Bonds: Formed by head-on orbital overlap (strong).

• Pi ($\pi$) Bonds: Formed by sideways overlap of unhybridized p orbitals (weaker).

• Bond Composition:

  • Single bond: one $\sigma$

  • Double bond: one $\sigma$ + one $\pi$

  • Triple bond: one $\sigma$ + two $\pi$

Resonance Structures

• Resonance structures are alternate Lewis structures differing only in electron placement (lone pairs or $\pi$ bonds), not atom positions.

• The resonance hybrid is the true representation, showing delocalized electron density, indicated by a double-headed arrow between contributing structures.

Quantum Mechanics & Molecular Orbitals

• Electrons behave as waves, described by wave functions ($\psi$). The probability of finding an electron is $\psi^2$.

• Molecular Orbitals: Formed by combining atomic orbitals.

  • Bonding orbitals: Constructive (in-phase) overlap, lower energy, electron density between nuclei.

  • Antibonding orbitals: Destructive (out-of-phase) overlap, higher energy, reduced density between nuclei.