Chemical Bonding, Polarity & Bond Properties

Electronegativity Scale

  • Unit-less, relative measure of how strongly a nucleus pulls on the shared electron cloud.
    • Li = 1.0 (weak pull)
    • N = 3.0 (strong pull, but weaker than F)
    • F = 4.0 (strongest pull; ~4× that of Li)
  • Compare atoms by ratio or difference, never by an absolute unit.

Dipole Moments & Vector Notation

  • Unequal pulls create a dipole moment (a tiny electric force).
  • Represented with a vector arrow:
    • Tail (with \sigma^+) on the electro-positive atom (electron deficiency).
    • Arrowhead points toward the electro-negative atom (electron surplus, \sigma^-).
  • “Pseudo” charge: partial charges are not full \pm 1 coulombs but measurable disparities.

Charge-Cloud Symmetry & Bond Polarity

  • Examine symmetry of the drawn electron cloud:
    • Symmetrical cloud → non-polar covalent.
    • Asymmetrical cloud → polar covalent (vector drawn).
    • Complete charge separation → ionic (always polar).

Quick Polarity Heuristics

  • Look for symmetry first; if absent, look up electronegativity values.
  • Biology vs. Chemistry convention:
    • Chemists treat “polar vs. non-polar” as a binary ("pregnant or not").
    • Biologists track degree of polarity because it affects solubility/reactivity in water.

Electronegativity Difference (Bond Character)

  • Formula: \Delta EN = |EN1 - EN2|.
  • Classification guideline:
    • \Delta EN > 1.66 → predominantly ionic (>50 % ionic character).
    • \Delta EN < 1.66 → predominantly covalent (small differences rely on orbital overlap).
  • No bond is 100 % ionic or covalent—think of a continuum.
  • Example exceptions:
    • Mg\,(1.2) – I\,(2.5) → \Delta EN = 1.3 ⇒ covalent despite being metal–non-metal.

Coulomb’s Law Context

  • F = k \dfrac{q1 q2}{d^2} governs all bond attractions.
    • Ionic focus: magnitude of q1, q2.
    • Covalent focus: small q but very small d (overlap distance).

Bond Strength & Environmental Dependence

  • Physical definition (textbook/Pauli): energy or temperature required to break one mole of bonds via heat.
  • Typical benchmark: NaCl melts/bonds break ≈ 850\ ^\circ C.
  • Water environment in biology reverses the trend:
    • Ionic lattices dissociate quickly (sphere of hydration); low energy cost.
    • Covalent bonds (e.g.
      C–H) remain intact → biologically “strong.”
  • Chemistry exam viewpoint: polar/ionic bonds are stronger (HF stronger than CH).
  • Biology exam viewpoint: covalent non-polar bonds appear strongest inside aqueous cells.

Chemical Reactivity Trend

  • Elements at extreme EN values (far left metals + far right non-metals) react most vigorously → highly exothermic.

Ionic Compounds

  • Requirements
    • Large \Delta EN, generally metal + non-metal.
  • Properties
    • High Coulombic attraction → high melting point.
    • Form 3-D crystal lattices (≈ 20 lattice types; “simple cubic” most common).
    • Conduct electricity only when molten or dissolved.
  • Crystal Lattice Example (NaCl)
    • Each Na^+ contacts 6 Cl^- ions (N, S, E, W, front, back).
  • Mechanism (simplified)
    1. Atoms approach.
    2. Electron(s) “transferred” from metal (cation) to non-metal (anion).
    3. +/- Coulombic attraction holds lattice.

Covalent Compounds

  • Occur with small \Delta EN (often non-metals, but also similar-EN metal/non-metal pairs).
  • Properties
    • Low melting/boiling points in dry heat.
    • Insoluble or poorly soluble in water ⇒ stronger inside aqueous environments.
    • Do not form lattices; instead form discrete molecules with varied shapes.
  • Molecular Geometry Terms
    • Bond length: distance between two nuclei along the bond axis.
      Notation: r_{AB}.
    • Bond angle: angle (θ) between two bond axes.
  • Shapes vary because bond lengths and angles adjust with orbital overlap → thousands of possible geometries (critical in biology: structure = function).
  • Mechanical property: brittle (electrons localized; stress causes fissures).
  • Mechanism
    1. Atoms get close.
    2. Atomic orbitals overlap → molecular orbital; two electrons locked in the overlap.
    3. Electrons simultaneously attracted to both nuclei (shared pair).

Metallic Bonds

  • Usually involve only one type of metal atom (intro chemistry assumption).
  • “Sea of electrons” model
    • Orbitals overlap in many directions; valence electrons delocalize throughout the crystal.
  • Properties
    • Metallic luster (shiny).
    • Excellent conductor of heat & electricity (free electron movement).
    • Malleable & ductile (layers slide without breaking bonds because electrons re-delocalize).
  • All atoms are electro-positive; bonding resembles covalent (overlap) but with maximal delocalization.

Numerical & Example Reference List

  • Li = 1.0, N = 3.0, F = 4.0 … used in all comparisons.
  • \Delta EN_{HF} = 4.0 - 2.2 = 1.8 → strong polar bond.
  • \Delta EN_{CH} = 2.4 - 2.2 = 0.2 → slightly polar; treated non-polar in biology.
  • \Delta EN_{MgI} = 2.5 - 1.2 = 1.3 → covalent.
  • Threshold for predominant ionic character: 1.66.
  • Bond break temp for NaCl ≈ 850\ ^\circ C (dry heat).

Ethical / Practical Implications

  • Understanding environment-dependent bond strength crucial in medicine (drug design relies on covalent stability in water) vs. industrial metallurgy (heat stability).
  • Energy release from highly exothermic ionic reactions has safety and industrial relevance (e.g.
    sodium metal + chlorine gas).

Connections to Prior Topics

  • Linus Pauling’s electronegativity scale (same Pauling of the Pauli exclusion principle & quantum numbers).
  • Revisits Coulomb’s law introduced in electrostatics.
  • Molecular geometry concepts tie into VSEPR and later spectroscopy discussions.