Molecular Shape and Polarity — Comprehensive Notes

Intermolecular vs Intramolecular Forces

• Intermolecular forces: forces that act between molecules; generally weaker than intramolecular forces
• Intramolecular forces: forces that hold atoms together within a molecule (e.g., covalent bonds, polar covalent bonds)
• In the context of the notes: categorization used is
  • Intramolecular: polar covalent bonds within a molecule
  • Intermolecular: dipole–dipole, London dispersion (van der Waals), hydrogen bonding, etc.
• Dative covalent bond (coordinate covalent bond) is a type of covalent bond where both electrons in the bond are donated by one atom. It is a bonding interaction that can be considered within the covalent (intramolecular) framework, but it often manifests in intermolecular adducts as well (e.g., donor–acceptor interactions between molecules).

Dipole Moment and Bond Polarity

• Dipole moment measures net molecular polarity and is associated with an unequal distribution of electron density
• Dipole moment definition (simplified):
  • oldsymbol{\mu} = q \; d
  • where q is the charge and d is the distance between charges (in Debye units, typically).
• Polar bonds arise when electrons are shared unequally due to differences in electronegativity between the bonded atoms
• A bond will be polar if there is a wide electronegativity difference between the bonded atoms, leading to a bond dipole
• Electronegativity difference (Δχ) drives bond polarity:
  • \Delta\chi = \chiA - \chiB
• Example for hydrogen chloride (HCl):
  • Given electronegativities: H ≈ 2.2, Cl ≈ 3.16
  • \Delta\chi = 3.16 - 2.2 \approx 0.96 \approx 1.0
  • This sizable difference results in a permanent dipole with Cl carrying partial negative charge and H partial positive charge
• In molecules like water (O–H): oxygen is more electronegative, so O bears a partial negative charge and H partial positive charges, giving a dipole moment
• In summary:
  • Polar bonds → dipole moments
  • Greater electronegativity difference → larger bond dipole → higher molecular polarity (if geometry allows a net dipole)

Dative Covalent Bond (Coordinate Covalent Bond)

• Definition: A dative bond is formed when both electrons of the bond are donated by one atom (the donor) to an atom with an available empty orbital (the acceptor)
• Representation: coordinate bond is often shown with an arrow from donor to acceptor (donor → acceptor)
• Requirements to form a dative bond:
  • Donor atom must possess at least one lone pair to donate
  • Acceptor atom must have at least one vacant orbital in its outer shell to receive the electron pair
• Key terms:
  • Electron donor: atom that provides lone pair electrons
  • Electron acceptor: atom with vacant orbital(s) to receive the pair
• Examples:
  • Ammonia–boron trifluoride adduct: \mathrm{NH3\cdot BF3}
  • Ammonia reacting with hydrogen chloride to form ammonium chloride: \mathrm{NH3 + HCl \rightarrow NH4^+ Cl^-}
  • Ammonia–aluminum chloride interaction: \mathrm{NH3 + AlCl3 \rightarrow NH3!:!AlCl3} (donor NH3 donates a lone pair to AlCl3)

Example: Ammonia and hydrogen chloride forming ammonium chloride

• Reaction: \mathrm{NH3 + HCl \rightarrow NH4^+ + Cl^-}
• Mechanism: ammonia donates its lone pair to the proton (H+) from HCl, forming the coordinate bond NH3→H+; the resulting ammonium ion is stabilized by ionic association with Cl− to yield ammonium chloride (NH4Cl)
• Visual concept: H+ is the acceptor, NH3 is the donor (via its lone pair)

Example: Ammonia–boron trifluoride adduct

• Reaction: \mathrm{NH3 + BF3 \rightarrow NH3\cdot BF3}
• Donor: NH3 supplies a lone pair to BF3 (empty p-orbital on B)
• Result: a stable coordinate covalent bond between N (donor) and B (acceptor)

Example: Dative bonding in AlCl3 systems

• In the Al2Cl6 dimer, two AlCl3 monomers are linked through two dative bonds
• Donor atoms: two chloride ions (Cl−) donate electron pairs
• Acceptor atoms: two aluminum centers (Al) accept electron pairs
• Structural implication: polymeric linkage via coordinate bonds between Cl donors and Al acceptors
• Diagrammatic summary (textual):
  • Cl --- Al ← coordinate bond from Cl to Al
  • Repeated to form the dimer/polymeric network

BeCl2 Dimer and Electron-Deficient Species

• BeCl2 is electron deficient and tends to form dative bonds to stabilize
• In the solid state, BeCl2 exists as a polymer where each Be is bonded to two Cl atoms by covalent bonds and two Cl atoms by coordinate bonds
• Why this occurs:
  • Be has an incomplete octet in its covalent bonding framework, driving the formation of additional dative bonds with donated lone pairs from Cl to Be to achieve greater stability
• Checkpoint example: Draw BeCl2 dimer with BeCl2 units linked by two dative bonds
• Answer concept: BeCl2 forms an electron-deficient dimer/polymer to achieve more stable coordination through coordinate bonds from chloride donors to Be centers

Van der Waals Forces (Intermolecular Forces)

• General term for intermolecular forces including:
  • Dipole–dipole interactions: occur between polar molecules; medium strength
  • Dipole-induced dipole interactions: a polar molecule induces temporary dipoles in nearby nonpolar molecules
  • London dispersion (dispersion) forces: present in all particles (atoms, ions, molecules); generally the weakest
• Key characteristics:
  • Dipole–dipole strength increases with larger dipole moments; greater energy required to separate the molecules
  • Dispersion forces exist even in nonpolar systems; strength increases with molar mass and polarizability
• Examples:
  • Hydrogen chloride (HCl) shows dipole–dipole interactions due to its polar bond
  • He and H2 have very weak dispersion forces due to their small size; I2 and Xe have stronger dispersion forces due to larger electron clouds
• Qualitative rule: strength of van der Waals forces generally increases with increasing molar mass/molecular size

Dipole–Dipole Interaction (Detailed)

• Definition: Permanent dipole–permanent dipole attractions between polar molecules
• Example evidence: HCl has a dipole because Δχ ≈ 1.0, giving H partial positive and Cl partial negative ends
• Consequence: contributes to boiling/melting points of polar substances compared with nonpolar ones

London Dispersion Forces (Dispersion/Van der Waals)

• Mechanism: momentary fluctuations in electron density create instantaneous dipoles in atoms; these induce dipoles in neighboring particles, resulting in attraction
• Present in all particles (polar or nonpolar)
• Strength depends on molar mass and size of the particle; larger electron clouds yield stronger dispersion forces
• Visual sequence (textual):
  • A (Ar) atoms are nonpolar → instantaneous dipole in one induces dipole in neighbor → temporary attractions
• Practical consequence: higher boiling points for heavier nonpolar species (e.g., I2, Xe) due to stronger dispersion forces

Hydrogen Bonding

• Definition: strongest type of weak intermolecular force; a special dipole–dipole interaction
• Condition: occurs when a hydrogen atom is covalently bonded to a highly electronegative atom (N, O, or F) and interacts with a lone pair on O, N, or F of another molecule
• Strength trend based on electronegativity (of the atom bonded to hydrogen):
  • F > O > N, so H–F hydrogen bonds are the strongest among these, followed by H–O and then H–N
  • Expressed as: H–F > H–O > H–N in hydrogen-bond strength
• General depiction: A–H…B where A and B are N, O, or F; hydrogen bonds show up as donor (A–H) and acceptor (B) interactions
• Consequences of hydrogen bonding:
  • Significantly influences boiling points and melting points (e.g., water, NH3, HF show higher boiling points than expected for nonpolar analogs)
  • Affects solubility of small covalent molecules in water (e.g., ammonia, methanol, ethanoic acid)
  • Impacts density of water and ice (ice less dense than water, due to open hydrogen-bond network)
• Common schematic indicators in text: A–H…B with B often being O, N, or F

Intermolecular vs Intramolecular Connections and Real-World Relevance

• Polarity and dipole moments influence:
  • Physical properties such as boiling/melting points, solubility, and phase behavior
  • Solvent choices in chemical reactions and separations
  • Structure of hydrogen-bonded networks in water and biological systems
• Real-world relevance examples:
  • Water’s high boiling point relative to other small molecules due to extensive hydrogen bonding
  • Solubility of alcohols and carboxylic acids in water linked to their ability to form hydrogen bonds with water
  • The density difference between liquid water and ice arising from hydrogen-bonding structure

Numerical References and Formulas to Remember

• Electronegativity values used in examples:
  • Hydrogen: \chi_{H} \approx 2.2
  • Chlorine: \chi_{Cl} \approx 3.16
  • Be: \chi_{Be} = 1.5
  • Chlorine (in BeCl2 context): \chi_{Cl} = 3.0
• Electronegativity differences (examples):
  • For H–Cl: \Delta\chi = \chi{Cl} - \chi{H} \approx 3.16 - 2.2 \approx 0.96 \approx 1.0
• Dipole moment: \boldsymbol{\mu} = q \cdot d (or equivalently, polarization-dependent definitions in Debye)
• General relationship notes:
  • Larger electronegativity differences yield stronger bond dipoles and higher molecular polarity
  • Molecules with permanent dipoles interact via dipole–dipole forces; nonpolar molecules interact mainly via dispersion forces unless induced dipoles occur

Summary of Key Concepts (Quick Reference)

• Intramolecular vs Intermolecular: within a molecule vs between molecules
• Polar bonds arise from electronegativity differences; dipole moment measures net polarity
• Dative (coordinate) covalent bonds: both electrons supplied by donor atom; arrow notation; common in adduct formation (NH3·BF3, NH4+ formation)
• BeCl2 tends to polymerize via dative bonding due to electron deficiency
• Van der Waals forces include dipole–dipole, dipole–induced dipole, and London dispersion; strength scales with molar mass and molecular size
• Hydrogen bonding is a strong dipole–dipole interaction involving N, O, or F with hydrogen; strongest among the weak interactions; affects boiling points, solubility, and density of water/ice