Organic Chemistry Exam 1 Study Guide: Chapters 1–3

Electronic Structure of Atoms

Electron Configurations

The arrangement of electrons in an atom is described by its electron configuration. This can be written in spectroscopic notation (e.g., 1s2 2s2 2p6) or shown using orbital box notation, where each box represents an orbital and arrows represent electrons.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

Shapes of Atomic Orbitals: s (spherical), p (dumbbell), d (cloverleaf). Each orbital has nodes:

• Radial nodes: Regions where the probability of finding an electron is zero, along the radius.

• Planar nodes: Planes where the probability is zero (e.g., the nodal plane in a p orbital).

• Valence electrons: Electrons in the outermost shell, involved in bonding.

• Core electrons: Electrons in inner shells, not involved in bonding.

Lewis dot symbols represent valence electrons as dots around the element symbol.

Chemical Bonding

Ionic and Covalent Bonds

Ionic bonds form between atoms with large differences in electronegativity, resulting in electron transfer. Covalent bonds involve sharing of electrons between atoms.

• Electronegativity: The tendency of an atom to attract electrons in a bond.

• Polarity: Covalent bonds are polar if there is a difference in electronegativity between atoms.

Lewis Structures

Lewis structures show the arrangement of atoms, bonds, and lone pairs in a molecule.

• Single, double, and triple bonds; lone pairs.

• Formal charge: Calculated as: $\text{Formal charge} = \text{Valence electrons} - (\text{Nonbonding electrons} + \frac{1}{2}\text{Bonding electrons})$

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

• Pi (π) bonds: Formed by side-on overlap of p orbitals.

• Typical bonding patterns for C, N, O, F (Group 14–17 elements).

• Resonance structures: Multiple valid Lewis structures for a molecule; the true structure is a resonance hybrid.

• Curved arrow formalism: Shows electron movement between resonance forms.

Kekule structure: Lewis structure with lone pairs often omitted. Condensed structural formula: Atoms and bonds written in a compact form.

Valence Bond Theory and Hybridization

Bond Formation and Hybrid Orbitals

Valence bond theory explains bonding as overlap of atomic orbitals.

• Sigma bonds: Overlap of s or hybrid orbitals.

• Pi bonds: Overlap of unhybridized p orbitals.

• Hybrid orbitals: Formed by mixing atomic orbitals (e.g., sp3, sp2, sp).

• Bond angles depend on hybridization: sp3 (109.5°), sp2 (120°), sp (180°).

• Rotation is possible about sigma bonds, not pi bonds.

• Lone pairs affect hybridization and geometry.

• Methyl cation (sp2), methyl anion (sp3), methyl radical (sp2).

Molecular Geometry and VSEPR Theory

Geometry and Bond Angles

The VSEPR (Valence Shell Electron Pair Repulsion) model predicts molecular geometry based on electron groups repelling each other.

• Electron groups (bonds and lone pairs) arrange as far apart as possible.

• Geometry and bond angles determined from Lewis structure.

• Lone pairs reduce bond angles.

• Polarity of molecules: Determined by bond polarities and geometry (dipole moment).

• Bond length and strength: Multiple bonds are shorter and stronger than single bonds.

Acids and Bases (Chapter 2)

Bronsted-Lowry Acids and Bases

Bronsted-Lowry acid: Proton donor. Bronsted-Lowry base: Proton acceptor.

• Strong acids/bases: Completely dissociate in water.

• Weak acids/bases: Partially dissociate.

• Acid-base equilibrium: Reversible reaction; identify acid/base on each side.

• Conjugate acid/base pairs: Acid becomes conjugate base after donating a proton.

Mechanistic Curved Arrows

Curved arrows show electron movement in reactions.

• Arrow from lone pair to H: Bond formation.

• Arrow from bond to atom: Bond breaking.

• Mechanistic arrows show reaction steps; resonance arrows show electron delocalization.

Acid Strength, pKa, and Ka

• pKa: Negative logarithm of acid dissociation constant ($pK_a = -\log K_a$).

• Lower pKa = stronger acid.

• Relative strength of conjugate bases: Compare Ka or pKa values.

Structural Effects on Acid Strength

• Electronegativity: More electronegative atoms stabilize negative charge.

• Resonance stabilization: Delocalization of charge increases acid strength.

• Size: Larger atoms stabilize charge better.

• Inductive effects: Electronegative substituents withdraw electron density, stabilizing charge.

Predict acid-base reaction outcome by comparing acid strengths; equilibrium favors weaker acid/base.

Resonance Structures and Electron Delocalization

Resonance structures show delocalized electrons. Draw arrows to indicate electron movement.

• Resonance explains relative acid strengths.

Lewis Acids and Bases

Lewis acid: Electron pair acceptor. Lewis base: Electron pair donor.

• Electrophiles: Lewis acids; attract electrons.

• Nucleophiles: Lewis bases; donate electrons.

• Base vs. nucleophile: Base accepts proton; nucleophile attacks other atoms.

pH and Structure of Organic Compounds

• Compare pH to pKa to determine protonation state.

• Weak acid in solution: If pH < pKa, acid form dominates; if pH > pKa, conjugate base dominates.

• Some molecules (e.g., amino acids) have multiple sites affected by pH.

Buffer Solutions

• Made from weak acid and its conjugate base.

• Resist changes in pH by neutralizing added acid or base.

Hydrocarbons and Alkanes (Chapter 3)

Types of Hydrocarbons

• Alkanes: Saturated hydrocarbons (single bonds).

• Straight chain alkanes: Unbranched; first 10: methane, ethane, propane, butane, pentane, hexane, heptane, octane, nonane, decane.

• Branched alkanes: Contain alkyl substituents.

• Cycloalkanes: Ring structures; general formula $C_nH_{2n}$.

Constitutional Isomers

Isomers with same molecular formula but different atom connectivity.

Nomenclature of Alkanes

• Base names for 1–10 carbons.

• Alkyl substituents: methyl, ethyl, propyl, isopropyl, sec-butyl, tert-butyl, etc.

• Draw structure from name; name from structure.

• IUPAC rules: Longest chain, lowest numbers for substituents, alphabetical order.

Degree of Alkyl Substitution

• Primary (1°) carbon: Attached to one other carbon.

• Secondary (2°) carbon: Attached to two carbons.

• Tertiary (3°) carbon: Attached to three carbons.

• Quaternary (4°) carbon: Attached to four carbons.

• Hydrogens classified similarly.

Alkyl Halides, Ethers, and Alcohols

• Structure and common names.

• IUPAC nomenclature for alkyl halides: Name halogen as prefix (e.g., chloromethane).

Cycloalkanes

• General formula: $C_nH_{2n}$

• Kekule and condensed structural formulas.

• Skeletal structures: Lines represent bonds; vertices are carbon atoms.

• Nomenclature of cyclohexanes; cis/trans stereochemistry.

Noncovalent Interactions (Intermolecular Forces)

• Types: London dispersion, dipole-dipole, hydrogen bonding.

• Relative boiling points: Alcohols > ethers > alkanes.

• Melting points: Not always straightforward; depends on structure.

• Solubility: Polar compounds (alcohols, amines) soluble in water; nonpolar (alkanes) soluble in hydrocarbons.

• Isomers: Branched alkanes have lower boiling points than straight chain.

• Classification of amines and alcohols (primary, secondary, tertiary).

Conformations of Molecules

Rotation and Conformations

Rotation about C–C single bonds allows molecules to adopt different shapes (conformations).

• Conformations of straight chain alkanes: Staggered (gauche, anti), eclipsed.

• Relative energies: Staggered lower energy than eclipsed.

• Rotation about sigma bonds; pi bonds restrict rotation.

• Torsional strain: Resistance to rotation due to electron repulsion.

• Steric strain: Repulsion due to atoms being too close.

• Dihedral angle: Angle between two bonds on adjacent carbons.

• Newman projections: Visualize conformations.

Conformations of Cycloalkanes

• Cyclohexane: Chair, boat, twist conformations.

• Angle strain: Deviation from ideal bond angles.

• Side view of chair conformation: Axial (vertical) and equatorial (horizontal) positions.

• Cis/trans isomerism on rings; equatorial vs. axial substituents.

• Relative stability: Substituents prefer equatorial positions; 1,3-diaxial interactions cause steric strain.

Table: Types of Intermolecular Forces and Their Effects

Type	Example	Relative Strength	Effect on Boiling Point
London Dispersion	Alkanes	Weak	Low
Dipole-Dipole	Ethers, Alkyl Halides	Moderate	Medium
Hydrogen Bonding	Alcohols, Amines	Strong	High

Table: First 10 Straight Chain Alkanes

Name	Formula	Structure
Methane	CH4	CH4
Ethane	C2H6	CH3CH3
Propane	C3H8	CH3CH2CH3
Butane	C4H10	CH3CH2CH2CH3
Pentane	C5H12	CH3(CH2)3CH3
Hexane	C6H14	CH3(CH2)4CH3
Heptane	C7H16	CH3(CH2)5CH3
Octane	C8H18	CH3(CH2)6CH3
Nonane	C9H20	CH3(CH2)7CH3
Decane	C10H22	CH3(CH2)8CH3

Example: Predicting the outcome of an acid-base reaction: Compare pKa values of reactants; the reaction proceeds toward the weaker acid/base.

Example: Drawing Newman projections for butane: Anti conformation (180° dihedral angle) is most stable; gauche (60°) is less stable due to steric strain.

Example: Naming a branched alkane: 2-methylpropane (isobutane) has three carbons in the main chain and a methyl group on the second carbon.

Example: Identifying intermolecular forces: Alcohols can hydrogen bond, leading to higher boiling points than alkanes.

Additional info: Classification of amines and alcohols (primary, secondary, tertiary) is introduced but covered in detail later in the semester.