Organic Chemistry Chapter 1: Review of General Chemistry: Electrons, Bonds, and Molecular Properties

1. What is Organic Chemistry?

• Definition: The study of carbon-containing molecules and their reactions.
• What happens during a reaction:
  • Molecules collide.
  • Bonds are broken and bonds are formed.
• Why reactions occur: influenced by electron behavior; fundamental focus on electrons; learning often spans extensive coursework (e.g., "we will need at least 2 semesters of your time to answer this question").

2. The Atom, The Periodic Table, and Atomic Concepts

• Atomic number Z: number of protons in a nucleus; defines the element.
• Mass number A: sum of protons and neutrons, A = Z + N.
• In a neutral atom, number of protons = number of electrons; ext{#electrons} = Z.
• The periodic table organizes elements by groups (columns) and periods (rows); blocks (s, p, d, f) reflect valence electron configurations.
• The table includes synthetic elements and materials with different states at room temperature.
• Valence electrons in main-group elements correspond to the group number.

3. The Atom: Quantum View and Electron Configuration

• Aufbau Principle:
  • Electrons are assigned to shells (n value) of increasing energy.
  • Within a shell, electrons fill subshells of successively higher energy.
  • The arrangement minimizes total energy.
  • Summary: fill from lowest energy up, obeying Pauli and Hund rules.
  • ext{Aufbau: fill order } 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p <
    ext{…}
• Pauli Exclusion Principle:
  • No two electrons in an atom can share the same four quantum numbers.
  • Two electrons may occupy the same orbital only if their spins are opposite.
  • Maximum of two electrons per orbital.
• Orbital types (basic): s, p, d, f shells determine shapes and energies.
• Atomic orbitals encode electron density; a region where the probability density of finding an electron is high (90–95% of the electron density is contained within the orbital cloud).
• Orbital shapes (illustrative):
  • s orbitals: spherical (1s, 2s, …)
  • p orbitals: dumbbell-shaped (px, py, pz)
  • d and f orbitals: more complex shapes (described in later sections).

4. The Atom: Orbital Shapes and Electron Density (s, p, d, f)

• s orbitals: sphere; one per shell (e.g., 1s, 2s).
• p orbitals: three degenerate orbitals (px, py, pz) with lobes along x, y, z axes.
• d orbitals: clover/diagonal shapes; five distinct d orbitals per shell starting at n ≥ 3.
• f orbitals: more complex shapes; start at higher n.
• Node concept: wavefunctions may have nodes where the probability of finding an electron is zero; signs of wavefunctions (+ or −) are not charges but relate to phase and overlap in bonding.

5. Electron Configurations and Periodic Trends

• Ground-state electron configurations can be written in two ways:
  • Full notation: e.g., Na: 1s^2 \, 2s^2 \, 2p^6 \, 3s^1
  • Noble-gas shorthand: e.g., Na: [Ne] \, 3s^1
• Example patterns (selected):
  • He: 1s^2
  • Na: [Ne] \, 3s^1
  • Mg: [Ne] \, 3s^2
  • Al: [Ne] \, 3s^2 \, 3p^1
  • Si: [Ne] \, 3s^2 \, 3p^2
• Valence electrons for main-group elements correspond to the outer s and p electrons; this equals the group number.
• Practical exercise: pick a number between 3 and 17, identify the corresponding atom, and determine its valence electrons (illustrated as an exercise in the transcript).

6. Periodic Table and Electronegativity

• Electronegativity (EN): the tendency of an atom in a molecule to attract electrons toward itself.
• Pauling scale values (typical references):
  • H ≈ 2.1, C ≈ 2.5, N ≈ 3.0, O ≈ 3.5, F ≈ 4.0
  • Cl ≈ 3.0, S ≈ 2.5, P ≈ 2.1, Si ≈ 1.9, Na ≈ 0.9, K ≈ 0.8
• EN differences drive bond polarity and bond type.
• Trend: EN generally increases across a period from left to right and decreases down a group.

7. Bonding: Ionic, Covalent, Polar Covalent

• Types of bonds depend on EN difference (ΔEN):
  • Covalent bond: ΔEN < 0.5 (electrons shared fairly).
  • Polar covalent bond: 0.5 ≤ ΔEN < 1.7 (unequal sharing).
  • Ionic bond: ΔEN > 1.7 (electrons transferred, forming ions).
• Example relationships from the transcript:
  • Ionic bonding involves electrostatic attraction between oppositely charged ions (cation–anion), e.g., Na+ and Cl−.
  • Covalent bonding involves sharing electron pairs; H2, CH4, etc.
• Bond dipoles: dipole moment μ depends on EN difference and bond distance; examples:
  • ext{H}_3 ext{C}– ext{F} ext{ has } oldsymbol{ ext{μ}} ext{ around } 1.4 ext{ Debye (D)}
  • ext{H}_3 ext{C}– ext{H} ext{ has } oldsymbol{ ext{μ}} ext{ around } 0.3 ext{ D}
• Some bonds fall on a continuum (cusp between polar covalent and ionic) and are not fixed by a single cutoff.

8. Intermolecular Forces (IMF)

• Intermolecular attractions influence solubility, boiling/melting points, density, and state.
• Main types:
  • Hydrogen bonding: strong dipole-dipole interaction involving H bonded to N, O, or F and a lone pair on another electronegative atom.
  • Dipole-dipole interactions: attractions between polar molecules.
  • London dispersion forces (LDF): present in all molecules, arise from instantaneous dipoles.
• Prototypical example: acetone (polar) vs isobutylene (nonpolar); acetone has stronger dipole-dipole interactions and higher boiling point.
• Protic vs aprotic solvents:
  • Protic solvents can donate H-bonds; aprotic cannot participate in H-bonding.

9. Simple Lewis Structures and Formal Charge

• Simple Lewis-structure steps:
  1) Draw atoms with valence electrons as dots.
  2) Connect atoms to share electron pairs to complete octets.
  3) Add lone pairs as needed to satisfy valency.
• Example: NH3 shows a lone pair on nitrogen.
• Formal charge calculation (important for resonance and stability):
  • ext{Formal charge} = ( ext{valence electrons on free atom}) - [ ext{nonbonding electrons} + frac{1}{2} ext{bonding electrons}]
• Example checks from the transcript:
  • An oxygen with seven electrons owned (of eight around it) has a formal charge of −1 because it needs six valence electrons to be neutral and currently owns seven.
  • Carbon in a structure with eight electrons around it but only owning four has a formal charge of 0 if it matches valence; otherwise, charges arise depending on sharing.

10. Molecular Geometry: VSEPR Theory

• VSEPR: electron pairs (bonding and lone pairs) repel; molecular shape derived from steric number (SN).
• Steric number (SN) determines hybridization and geometry:
  • SN = 4 → sp3 → tetrahedral electron-pair geometry.
  • SN = 3 → sp2 → trigonal planar electron-pair geometry.
  • SN = 2 → sp → linear electron-pair geometry.
• Common geometries (from table in transcript):
  • CH4: SN 4 → Tetrahedral, geometry: tetrahedral
  • NH3: SN 4 with 1 lone pair → Trigonal pyramidal
  • H2O: SN 4 with 2 lone pairs → Bent
  • BF3: SN 3 → Trigonal planar
  • BeH2: SN 2 → Linear
• Key angles:
  • Tetrahedral: ~109.5°
  • Trigonal planar: ~120°
  • Linear: ~180°

11. Hybridized Atomic Orbitals and Carbon's Bonding Problem

• Carbon’s ground-state config: 1s^2 2s^2 2p^2; can’t explain four equivalent CH bonds.
• Hybridization resolves this: carbon undergoes hybridization to form four equivalent orbitals with the same energy.
• sp3 hybridization:
  • 25% s-character, 75% p-character.
  • Four sp3 orbitals arranged tetrahedrally to form four equivalent σ bonds (e.g., CH4).
  • Example: CH4 forms from overlap of C sp3 orbitals with H 1s orbitals.
• sp2 hybridization (e.g., ethene, CH2=CH2):
  • 33% s-character, 67% p-character.
  • Three sp2 orbitals form σ-bonds, with one unhybridized p orbital left for π-bonding.
• sp hybridization (e.g., acetylene, C≡CCH):
  • 50% s-character, 50% p-character.
  • Two sp orbitals form σ-bonds; remaining p orbitals form π-bonds.
• Pi bonds and sigma bonds:
  • σ (sigma) bonds result from end-to-end overlap of orbitals.
  • π (pi) bonds result from side-by-side overlap of p orbitals.
• Practice reference: SkillBuilder 1.7 focuses on identifying hybridization states.

12. Molecular Orbital (MO) Theory vs Valence Bond Theory

• Valence Bond (VB) theory:
  • A bond forms when atomic orbitals overlap; constructive interference yields a bond (σ).
  • Localized bonding orbitals describe bonds between specific atoms.
• Molecular Orbital (MO) theory:
  • Atomic orbitals combine to form molecular orbitals that extend over the entire molecule.
  • MOs account for constructive (bonding) and destructive (antibonding) interference.
  • Number of MOs equals the number of AOs used to form them (e.g., H2 has 2 AOs → 2 MOs).
• Key MO concepts:
  • Bonding MO is lower in energy than the corresponding antibonding MO.
  • Antibonding MO is higher in energy and less favorable for bond formation.
  • HOMO: Highest Occupied Molecular Orbital; LUMO: Lowest Unoccupied Molecular Orbital.
  • Reactions often involve electrons in HOMO/LUMO.
• Example: For CH3Br, there are multiple MOs; focus often on those participating in reactions.

13. Conjugation and Aromaticity (Overview)

• Conjugation involves alternating single and multiple bonds allowing π-electron delocalization across atoms.
• Aromatic compounds follow special stability rules (e.g., Huckel’s rule for planar rings with (4n + 2) π electrons); further details appear in later chapters.

14. Intermolecular Forces: Summary and Implications

• Important IMF implications for properties:
  • Solubility, boiling/melting points, density, and phase behavior.
• Major forces:
  • Hydrogen bonding (special case of strong dipole-dipole interaction).
  • Dipole-dipole interactions (between polar molecules).
  • London dispersion forces (present in all molecules; arise from instantaneous dipoles).

15. Practice and Conceptual Nuggets

• Octet Rule (recap): Most atoms prefer to have 8 electrons in their outermost shell to achieve stability.
• Ionic vs Covalent vs Polar covalent distinctions depend on ΔEN and electron distribution.
• Formal charge checks help determine the most stable Lewis structure and resonance forms.
• Molecular geometry (VSEPR) provides quick predictions of shapes based on electron-domain repulsion.
• Hybridization concepts explain observed bond lengths and angles in organic compounds.

16. Quick Reference: Key Formulas and Facts

• Mass number: A = Z + N
• Charge balance in neutral atoms: ext{#electrons} = Z
• Formal charge: ext{Formal charge} = V - (NBE + frac{1}{2}B) where
  • V = valence electrons on free atom
  • NBE = nonbonding electrons
  • B = bonding electrons
• Octet Rule: Atoms tend to achieve 8 valence electrons.
• Steric number (SN) and hybridization:
  • SN = number of bonded pairs + number of lone pairs
  • SN = 4 → sp3; SN = 3 → sp2; SN = 2 → sp
• Bond types by EN difference:
  • Covalent: ΔEN < 0.5
  • Polar covalent: 0.5 ≤ ΔEN < 1.7
  • Ionic: ΔEN > 1.7
• Dipole moment intuition: oldsymbol{ ext{μ}} ext{ depends on } ext{ΔEN} ext{ and bond distance } r
• Representative EN values (Pauling scale): H ≈ 2.1, C ≈ 2.5, N ≈ 3.0, O ≈ 3.5, F ≈ 4.0, Cl ≈ 3.0, Na ≈ 0.9, K ≈ 0.8
• MO vs VB highlights:
  • VB emphasizes localized bonds via orbital overlap (σ/π).
  • MO emphasizes delocalized orbitals across the molecule; HOMO/LUMO are key for reactivity.

17. Notable Structures and Examples (from the transcript)

• Common small species: CH3, CH4, -CH3, NH4+, NH3, -NH2, H3O+, H2O, H3C–CH3, H2C=CH2, HC≡CH (ethyne).
• Prototypical strong H-bonding networks occur in water and ammonia derivatives due to H attached to O/N/F.
• BeH2: linear geometry with sp hybridization; CO2 Lewis structure demonstration.
• BF3: trigonal planar geometry with sp2 hybridization on B and an empty p orbital.

Note: The above notes summarize the content from the provided transcript, including foundational concepts in atomic structure, bonding, orbitals, and introductory physical organic chemistry. Where equations and numeric relationships appeared in the source, they are reproduced in LaTeX format for clarity and exam use. If you want, I can convert these notes into a printable PDF or tailor a study plan around the specific chapters and problem types you expect to encounter on the exam.