CHEM 231: Molecular Structure and Acid–Base Fundamentals

1.1 Orbitals take up three-dimensional space

• Electron cloud surrounds the nucleus; electrons occupy orbitals that have specific shapes and orientations in three dimensions.
• Hybridization concepts are central to understanding geometry and bonding in this course.
• Hybrids of orbital types (e.g., s and p) are the primary focus for predicting bonding patterns and shapes.

1.2 Ionic bonds are not the focus of this class

• Ionic bonds arise from electrostatic attraction between ions with opposite charges.
• Components: cations (+) and anions (-).
• Characteristics:
  • High melting points (example: NaCl ≈ 800°C).
  • Tend to be soluble in water; insoluble in most organic solvents.
  • Crystalline in the solid state.

1.3 Bonding: Valence Bond model vs Molecular Orbital model

• Valence Bond (VB) picture for dihydrogen (H₂):
  • Uses the 1s orbitals of two hydrogen atoms.
  • Each 1s orbital contains one electron.
  • In-phase overlap of two 1s orbitals forms a bonding orbital that encompasses both hydrogens and contains two electrons.
• Molecular Orbital (MO) picture for H₂:
  • Construct MOs from the combination of atomic orbitals (AOs).
  • Bonding MO (in-phase) results in energy stabilization.
  • Antibonding MO (out-of-phase) costs energy; node between A and B.
  • Energy diagrams indicate increasing energy for antibonding configurations.
• Key concepts:
  • Bond order, stabilization vs destabilization via constructive vs. destructive interference of AOs.
  • Nodes and energy of MOs determine bond strength and reactivity.

1.4 Molecular geometries are dictated by orbitals

• VSEPR-inspired geometry arises from repulsions between electron domains (bonding and lone pairs).
• Example geometries with approximate angles:
  • Methane (CH₄): tetrahedral, ~109.5°.
  • Water (H₂O): bent, ~104.5° (due to two lone pairs).
  • Ammonia (NH₃): trigonal pyramidal, ~107° (lone pair repulsion).
  • Other illustrated examples show ~107°, ~105°, etc., depending on lone-pair/bonding-pair repulsions.
• Concept of repulsions:
  • Bonded pair–bonded pair: least repulsive
  • Bonded pair–unshared pair: intermediate
  • Unshared pair–unshared pair: most repulsive
• Context: these geometries reflect the arrangement of electron domains around the central atom.

1.5 Polar covalent bonds, electronegativity, and bond dipoles

• Electronegativity (Pauling scale) governs how electrons are shared in a bond.
• Bond dipole moments arise from differences in electronegativity between atoms in a bond.
• Examples of dipole moments (in Debye, D):
  • H–F ≈ 1.7 D
  • C–F ≈ 1.4 D
  • H–C ≈ 0.3 D (C more electronegative end is negative in polar C–H)
• Increasing number of bonds generally increases the overall molecular dipole moment (more polarity contributions).
• Direction of a bond dipole is toward the more electronegative atom.
• Implication: bond polarity contributes to molecular polarity when dipoles do not cancel in a symmetric molecule.

1.6 Polar molecules and net dipole moment

• Molecular dipole moment μ is the vector sum of individual bond dipoles.
• If molecule is highly symmetric (e.g., CO₂), dipoles cancel and μ = 0 (nonpolar molecule).
• Example: H–C–Cl has a net dipole toward Cl; Cl–C dipole magnitude cited as μ ≈ 1.87 D (illustrative example).
• Nonzero net dipole → polar molecule; zero net dipole → nonpolar molecule.

1.7 Resonance; electron delocalization: a first approach

• Ozone (O₃) as an example of resonance:
  • Two major resonance contributors depict electron delocalization across the O–O–O framework.
  • Resonance hybrid represents a blended distribution of electrons.
• Resonance notation:
  • Dashed-line structures indicate resonance contributors.
  • The actual structure is a resonance hybrid of all valid contributors.

1.8 Curved arrows (Electron movement)

• Curved arrows depict movement of electrons, not a reaction step.
• Rules (summary):
  • Move electrons from electron-rich sources (lone pairs or bonds) to electron-poor targets.
  • Do not move electrons from a positively charged center to a more charged center inappropriately.
  • Correct arrow-pushing is developed with practice.
• Example themes:
  • Representing resonance contributors (major vs minor) via arrow-pushing.
  • Proton transfer or bond-making/breaking steps in mechanisms.

1.9 Curved arrows: additional practice and mechanism steps

• Illustrative bimolecular substitutions (SN2) and acid–base reactions show ARROW- pushing during bond formation/breakage.
• Example sequences:
  • HO⁻ attacking a carbon center in S_N2-like steps to form new bonds while displacing leaving groups.
  • Proton transfers where a conjugate base abstracts a proton from an acid.

1.10 Formal charge; who owns the electrons?

• Formal charge concept helps assign charges to atoms in a Lewis structure.
• Nitromethane example: shows charges on N and O to balance electron ownership.
• Formal charge calculation:
  • Electron ownership per atom = valence electrons − electrons owned by the atom (bond electrons are split between bonded atoms).
  • Formula: $ext{Formal charge} = V - E_{owned}$ where V = valence electrons.
• Example calculations illustrate how to assign charges to achieve the most reasonable Lewis structure.

1.11 Practical Lewis structures: formal charge and octet considerations

• Steps to draw reasonable Lewis structures:
  • Count total valence electrons from all atoms (adjust for ionic species if applicable).
  • Draw a skeleton structure (connect atoms with bonds) to satisfy octets as far as possible.
  • Distribute remaining electrons as lone pairs to satisfy octets.
  • Calculate formal charges and adjust if needed to minimize charge separation.
• Example: C₂H₆O
  • Total valence electrons: from C (group 4A) and O (group 6A) and H (group 1A).
  • Two major skeletons: H–C–C–O–H vs H–C–O–C–H; decide based on octets and formal charges.
  • If octets are satisfied without charges, that structure is preferred; otherwise assign charges appropriately.
• Key principle: the atom prefers an arrangement that minimizes formal charge separation and obeys octet rules for second-row elements where applicable.

1.12 Lewis formulae: double and triple bonds; line-angle drawings

• Double and triple bonds are used to depict multiple bonds between the same two atoms.
• Line-angle drawings: simplified representation of the carbon skeleton.
  • Each line endpoint or bend represents a carbon atom.
  • Hydrogens bonded to carbons are inferred and not drawn explicitly.
  • Explicit heteroatoms (O, N, halogens, etc.) and their hydrogens may be drawn.
• Examples show transformations between Lewis structures and line-angle notations.

1.13 Structure and acidity strength; periodic trends in acidity

• Trend 1: In a column (group), larger conjugate base anion → stronger acid (lower conjugate base stability).
  • Examples (pKa values, approximate):
  • HF: pKa ≈ 3.1
  • HCl: pKa ≈ -3.9
  • HBr: pKa ≈ -5.8
  • HI: pKa ≈ -10.4
  • Strongest acid among the listed in the group is HI (lower pKa).
  • Note: The hydrogen is directly bonded to the atom being compared.
• Trend 2: Across a row (period), electronegativity is a strong predictor of acid strength.
  • Examples across period: CH₄ (pKa ≈ 50) < NH₃ (pKa ≈ 38) < H₂O (pKa ≈ 15.7) < HF (pKa ≈ 3.1) – more electronegative central atom → stronger acid.
• Additional context examples (pKa values, approximate):
  • Hydronium (H₃O⁺) = 0
  • Water (H₂O) = 15.7
  • Ammonia (NH₃) = 38
  • Ammonium ion (NH₄⁺) = 9.3
  • Acetic acid (CH₃COOH) = 4.7
  • Benzene conjugate acid ≈ 43 (very weak acid)
• Takeaway: acid strength is governed by both electronegativity and conjugate-base stability; the trend is reinforced by inductive effects (neighboring substituents) and delocalization.

1.13 Environmental trend 1: Inductive effects

• Inductive withdrawing/donating groups near the conjugate base atom affect acidity.
• Example: Ethanol vs. 2,2,2-trifluoroethanol (CF₃CH₂OH)
  • Ethanol: pKa ≈ 16
  • 2,2,2-Trifluoroethanol: pKa ≈ 11.3
• Principle: Electron-withdrawing groups stabilize conjugate bases and lower pKa (increase acidity) via inductive effects.
• Diagrammatic note: inductive effects propagate through sigma bonds and influence stability of charged species.

1.13 Environmental trend 2: Delocalization in the conjugate base (CB)

• Delocalization of negative charge in the CB stabilizes the CB.
• If CB is able to delocalize electrons (resonance, conjugation), the conjugate base is weaker and the corresponding acid is stronger.
• Example schematic ideas (not all shown here): delocalized CB charges across multiple atoms lead to greater stability.

1.14 Acid–Base Equilibria

• Concept: Reactions proceed to form the weaker acid or base (weaker acid + stronger base on one side).
• Rule: Compare pKa values of the acids involved to predict direction of equilibrium.
• Example 1 (bromide and water):
  • Reaction: HBr + H₂O ⇄ Br⁻ + H₃O⁺
  • pKa(HBr) ≈ -4.0 to -5.0; pKa(H₃O⁺) ≈ 0
  • Equilibrium favors the side with the weaker acid (HBr has the stronger acidity, so the equilibrium lies to the left, toward Br⁻ and H₂O).
• Example 2 (acetic acid vs. water):
  • Reaction: CH₃COOH + H₂O ⇄ CH₃COO⁻ + H₃O⁺
  • pKa(CH₃COOH) ≈ 4.7; pKa(H₂O) ≈ 15.7
  • Equilibrium favors the weaker acid side (acetic acid is weaker than hydronium, so the reaction lies toward products if comparing conjugate partners in a generalized sense; in a comparative acid-base sense, acetate is a stronger base than water).
• Example 3 (ammonia vs. acetic acid):
  • Ammonia (NH₃) has pKa ≈ 38 as an acid (very weak); ammonium (NH₄⁺) has pKa ≈ 9.3 as an acid.
  • Acid-base strength comparison can be used to reason about equilibrium direction in proton transfers.

1.16 Lewis acid–base equilibria

• Definitions:
  • Lewis acid: electron pair acceptor.
  • Lewis base: electron pair donor.
• Example: Boron trifluoride (BF₃) acts as a Lewis acid.
  • BF₃ can accept electron density from a Lewis base such as diethyl ether (Et₂O) to form a Lewis acid–base adduct, commonly referred to as the "boron trifluoride etherate":
  • BF₃ + Et₂O → BF₃·Et₂O
• Example reaction:
  • A-B (Lewis acid A and Lewis base B) forms a Lewis acid–base complex A–B.
  • Illustration: BF₃ accepting an electron pair from an ether gives a stable adduct.

1.17 Practice: curved arrows and resonance reasoning (summary ideas)

• Practice questions involve drawing curved arrows to show plausible electron flow leading to products.
• Typical tasks:
  • Predict major resonance contributor for a given structure by applying resonance rules.
  • Decide if an attempted resonance form is valid by checking octet rule, charge distribution, and electron counts.
  • Determine if a proposed curved-arrow sequence leads to stable products without violating fundamental rules (no movement of electrons to create impossible charges, etc.).

1.4 Double and triple bonds: depiction and examples

• Double bonds (alkenes) and triple bonds (alkynes) are depicted by multiple lines between two atoms.
• Examples show how to share electrons to form C=C and C≡C bonds.
• Oxidized species such as carbon dioxide can be shown with multiple bond representations (e.g., O=C=O).

Summary: Connections to concepts and real-world relevance

• Orbitals and hybridization explain three-dimensional shapes of molecules, reactivity, and bonding patterns observed in organic and inorganic chemistry.
• Electronegativity and dipole moments underpin bond polarity, molecular polarity, and properties like solubility and boiling points.
• Resonance and curved arrows explain electron delocalization, stabilization of structures, and reaction mechanisms.
• Formal charges and Lewis structures provide a practical framework for assessing stability and predicting reactivity in molecules.
• Acidity trends (periodic and inductive effects) guide predictions about proton transfer, reaction equilibria, and the strength of acids and bases in different environments.
• Lewis acids and bases extend beyond Bronsted definitions to include electron-pair interactions, enabling explanations of complexation and catalysis.

Key equations and constants to remember

• Formal charge: $ext{Formal charge} = V - E_ ext{owned}$
  • where V is the number of valence electrons for the atom and E_owned is the number of electrons assigned to that atom in the Lewis structure.
• pKa relation: $pK<em>a = -\, ext{log}</em>{10} K_a$
• Dipole moment direction: toward the more electronegative atom in a bond; vector sum determines molecular polarity.
• Resonance rules (summary):
  • Connectivity remains the same
  • Total number of electrons and net charge unchanged
  • Number of unpaired electrons unchanged
  • Octet rule should be satisfied for second-row elements
• Lewis acid–base adducts: BF₃ + Et₂O → BF₃·Et₂O (illustrative example of Lewis acid-base interaction)