Notes on Bonding, Polarity, and Intermolecular Forces (Review)

Electronegativity, Bonding, and Polarity

• Electronegativity (EN) is a measure of how well an element competes for electrons in a bond. The atom that attracts electron density more strongly holds onto more of the bonding electrons, creating partial charges in the bond.

• Bond type depends on the difference in electronegativity between the two atoms, ΔEN.

  • If ΔEN < 0.4, the bond is essentially pure covalent (very little charge separation).
  • If 0.4 ≤ ΔEN ≤ 2.0, the bond is polar covalent (significant, but not full, charge separation).
  • If ΔEN > 2.0, the bond becomes highly polar and can be viewed as strongly ionic in character (not the focus here, but indicates large charge separation).

• Example: HF

  • Fluorine is much more electronegative than hydrogen (F has a very high EN). For HF, the difference is $riangle EN \,=\, EN(F) - EN(H) \approx 4.0 - 2.1 = 1.9$
  • This places HF in the polar covalent region (near the high end).
  • The more electronegative atom (F) hogs electron density, creating a partial negative charge on F (denoted $\boldsymbol{\delta^-}$) and a partial positive charge on H (denoted $\boldsymbol{\delta^+}$).

• Polar vs nonpolar concepts are relative: some bonds are very polar (e.g., HF), some are weakly polar, and some can be considered nonpolar depending on ΔEN and molecular geometry.

• Example comparisons:

  • CO₂: Two C=O bonds are polar, but the molecule is linear and the bond dipoles cancel, yielding a nonpolar molecule.
  • H₂O: Bent geometry prevents cancellation of bond dipoles, giving a net molecular dipole (polar molecule).

• Key electronegativity values used in examples:

  • Oxygen: EN ≈ $3.5$
  • Hydrogen: EN ≈ $2.1$
  • Carbon: EN ≈ $2.5$ (contextual; used to compare with O and H)

• Dipole moments and partial charges:

  • The more electronegative atom in a bond carries the partial negative charge; the less electronegative partner carries the partial positive charge.
  • The partial charge is often represented as $\delta^-$ on the more electronegative atom and $\delta^+$ on the less electronegative atom.

• Linear vs bent vs tetrahedral geometries:

  • Geometry determines whether bond dipoles cancel or reinforce, affecting net molecular polarity.
  • CO₂ is linear; water is bent; CH₄ is tetrahedral.

• Bond strength and energy concepts (contextual):

  • Intermolecular attractions (IMFs) are weaker than intramolecular bonds.
  • Example: the energy to separate a mole of HCl in the liquid phase into gas (to overcome neighbor attractions) is about $\Delta H_{vap} \approx 17\ \text{kJ mol}^{-1}$.
  • Contrast with covalent bond energy: breaking a covalent bond typically requires on the order of $\sim 400\ \text{kJ mol}^{-1}$ per bond.

• How to relate to states of matter:

  • Solids: strong attractions to neighbors; molecules are relatively immobile and only vibrate in place.
  • Liquids: still attracted to neighbors but can slide past each other; more mobility than a solid.
  • Gases: negligible or zero intermolecular attractions (ideal gas assumption), allowing particles to move freely apart.
  • Temperature increases kinetic energy (motion) and can overcome IMFs, leading to phase changes (solid → liquid → gas).

• Intermolecular vs intramolecular forces:

  • Intramolecular forces: covalent and ionic bonds within a molecule (strongest attractions).
  • Intermolecular forces (IMFs): attractions between neighboring molecules (weaker, noncovalent); include dipole–dipole, London dispersion forces, and hydrogen bonding (a subset of dipole–dipole).
  • The term Van der Waals forces is often used synonymously with IMFs, though it is a broad umbrella term.

• Visualizing IMFs:

  • Dipole–dipole attractions: occur between molecules with a net dipole (partial + and partial − ends).
  • Hydrogen bonds: a particularly strong form of dipole–dipole attraction; requires a hydrogen attached to N, O, or F in one molecule and a lone pair (or partially negative site) on a neighboring highly electronegative atom in another molecule.
  • Hydrogen bonding is not a true chemical bond; it is a strong IMF, illustrated by dotted lines between neighboring molecules.

• Hydrogen bonding specifics:

  • Conditions: when a hydrogen is covalently bonded to F, O, or N, the hydrogen can form a strong attraction to electronegative lone pairs on neighboring molecules.
  • Example molecules: water (H₂O), ethanol (CH₃CH₂OH), ammonia (NH₃), HF, etc.
  • The strength of hydrogen bonds is aided by a large EN difference (e.g., H–F, H–O, H–N bonds) and the small size of these atoms concentrates charge.
  • Multiple hydrogen bonds can form per molecule and between multiple molecules, contributing to higher boiling points and specific liquid structures.

• The concept of a homologous series in hydrogen-containing compounds:

  • Consider molecules of the form H–X where X is a group 15–17 element (e.g., S, Se, Te down the group). As X becomes larger down the group, the mass and number of electrons increase, enhancing London dispersion forces.
  • Boiling points tend to rise with heavier X in H–X series due to stronger dispersion forces and greater molecular size, despite similar polarity trends within a given row.
  • Example progression: H–S, H–Se, H–Te show increasing boiling points as you move down the group.
  • This contrast with the polar features introduced by oxygen, nitrogen, and fluorine, which can dominate water, ammonia, and related compounds through hydrogen bonding.

• Quick recaps you should be able to do after this material:

  • Given two elements with a known electronegativity difference, determine whether their bond is pure covalent, polar covalent, or predominantly ionic.
  • Predict whether a molecule will be polar or nonpolar based on bond dipoles and molecular geometry.
  • Identify the different types of intermolecular forces present in a substance and explain how they influence state of matter and phase changes.
  • Explain why hydrogen bonding strengthens certain liquids (e.g., water, ethanol) and how it differs from ordinary covalent bonds.
  • Distinguish between intramolecular bonds and intermolecular forces, and relate those to energy requirements for phase transitions (e.g., vaporization vs bond dissociation).

• Connections to foundational ideas:

  • This content builds on the idea from Gen Chem I that electronegativity differences drive bond polarity and that geometry determines the net dipole of a molecule.
  • It connects to VSEPR theory (tereactions about electron pairs around central atoms) to explain why CO₂ is nonpolar while H₂O is polar.
  • It ties into thermodynamics and kinetics: the magnitude of IMFs affects boiling points and phase stability, and temperature (kinetic energy) controls the ability of molecules to overcome IMFs.

• Notation and terminology to be comfortable with:

  • $\Delta EN = EN<em>A - EN</em>B$: electronegativity difference between two bonded atoms.
  • $\delta^- , \delta^+$: partial negative and partial positive charges in a polar bond.
  • Intermolecular forces (IMFs): attractions between neighboring molecules (not chemical bonds).
  • Dipole–dipole attractions: IMF between molecules with permanent dipoles.
  • Hydrogen bonding: a strong subset of dipole–dipole interactions involving H–N, H–O, or H–F bonding scenarios.
  • London dispersion forces: weak IMFs arising from instantaneous dipoles, present in all molecules, especially significant in nonpolar species.
  • Van der Waals forces: a broad term often used synonymously with IMFs.

• Practical implications (what this means in the real world):

  • Polarity and IMF strength explain why water is a liquid at room temperature and has unusually high boiling point for a small molecule, while CO₂ is a gas at room temperature despite having polar bonds (cancellation due to linear geometry).
  • Hydrogen bonding underpins the structure of DNA and many biological molecules, as well as the properties of alcohols and water as solvents.
  • When teaching, it’s important to distinguish a hydrogen bond (an IMF) from a chemical bond; breaking it is not the same as breaking a covalent bond, and phase changes often involve breaking IMFs rather than chemical bonds.

• Quick example recap linking to the transcript:

  • CH₄: ΔEN ≈ 0.4, tetrahedral geometry, overall nonpolar due to dipole cancellation.
  • CO₂: two C=O bonds with ΔEN ≈ 1.4 each, linear geometry, dipoles cancel → nonpolar molecule.
  • H₂O: ΔEN per O–H ≈ 1.4, bent geometry, net dipole → polar molecule with strong hydrogen bonding.
  • I₂: ΔEN ≈ 0, nonpolar diatomic, solid at room temperature, main features due to dispersion forces and symmetry.

• End of summary: be prepared to apply these concepts to new molecules by assessing EN differences, predicting bond type, evaluating geometry, and estimating the relative strength of IMFs between neighboring molecules.