Notes on Relative Strengths of Interparticle Forces (Unit 10 Part 2)

Covalent Bonds

• The strongest interparticle forces among common types. They hold atoms together in networks or within molecules.
• Evidence of strength: water boils after covalent bonds within water break; hydrogen bonds break before O–H covalent bonds do.
• We will not rank covalent bonds across different covalent-network substances.

Relative Strengths – Quick Summary

• Covalent bonds → strongest
• Ionic bonds
• Metallic bonds
• Strong hydrogen bonds
• Dipole-dipole forces (DDF)
• London dispersion forces (LDF)
• Weakest overall
• Reminder: "charge" and "distance" are recurring themes in comparing strengths.

Ionic Bonds – Key Rules

• Strength grows with higher ion charges: strength ∝ product of charges |q{ ext{cation}}| imes|q{ ext{anion}}|.
  • Example: Na⁺–NO₃⁻: $(1)(1)=1$; Ca²⁺–S²⁻: $(2)(2)=4$; Fe²⁺–S²⁻: $(2)(2)=4$; Al³⁺–O²⁻: $(3)(2)=6$.
• When charges are the same in two compounds, the bond is stronger for the compound with smaller ions (shorter distance).
  • Compare radii of cations and anions; smaller radii → ions closer together → stronger ionic bond.
• Polyatomic ions: sketch likely structure to estimate sizes for comparison.

Example #1: CaS vs MgO

• Charges: both have +2 and -2 → product = 4 (same).
• Cation radii: Ca²⁺ larger than Mg²⁺; Anion radii: S²⁻ larger than O²⁻.
• Result: MgO ions are smaller → ionic bond stronger in MgO.
• Evidence: MgO mp = 2852^\circ\mathrm{C} vs CaS mp = 2525^\circ\mathrm{C}.

Example #2: KCl vs KClO₃

• Charges: +1 and -1 → product = 1 (same).
• Cations: same (K⁺).
• Anions: Cl⁻ vs ClO₃⁻; Cl⁻ is smaller.
• Predicted strength: KCl stronger than KClO₃.
• Evidence: mp(KCl) = 770^\circ\mathrm{C}; mp(KClO_3) = 356^\circ\mathrm{C}.

Metallic Bonds – Key Rules

• Forces arise from cations in a sea of delocalized electrons.
• What differs among metals: cation charge, number of delocalized electrons, and cation radius.
• General trend: higher cation charges and more delocalized electrons → stronger metallic bonding; smaller cation radii can contribute.
• Common charges (example set):
  • K: +1
  • Ca: +2
  • Cr: +3
• Predicted metallic bond strength (relative): K < Ca < Cr.

London Dispersion Forces (LDF) – Key Rules

• Strength increases with size of atoms/molecules (more electrons → larger electron cloud → stronger instantaneous dipoles).
• Shape/surface area matters: more surface area → stronger LDF.
• For a set of nonpolar molecules with the same formula, larger molecules have stronger LDF and higher mp/bp.
• Typical ranking for a group like H₂, F₂, Cl₂, Br₂, I₂: H₂ < F₂ < Cl₂ < Br₂ < I₂ (weakest to strongest LDF).
• Physical state at room temperature often follows size: H₂, F₂, Cl₂: gas; Br₂: liquid; I₂: solid (as size increases).

London Dispersion vs Polarity in Polar Molecules

• Dipole-dipole forces (DDF) are stronger for more polar molecules.
• Polar molecules also have LDF; the dominant interaction depends on size and polarity.
• Rules of thumb:
  • If two molecules are similar in size (LDF similar) but differ in polarity, DDF can dominate.
  • If molecule sizes differ (one much larger) and polarity differs, LDF often dominates.
• Example: NF₃ vs CHF₃
  • Molar masses: NF₃ (71) vs CHF₃ (70); LDF strength similar; CHF₃ is more polar (DDF stronger); Boiling points reflect this (CHF₃ bp higher).
• Example: PF₃ vs PCl₃
  • PF₃ is smaller and more polar (DDF stronger); PCl₃ is larger (LDF stronger); Boiling point data show LDF can dominate when size difference is significant.
• Example: HF vs HCl/HBr/HI
  • HF forms hydrogen bonds, which are much stronger than typical DDF or LDF, yielding a much higher boiling point (HF bp ≈ 19 °C, much higher than the others).

Practical Takeaways for Quick Recall

• Strength order (strongest to weakest): Covalent bonds > Ionic bonds > Metallic bonds > Hydrogen bonds > Dipole-dipole forces > London dispersion forces.
• For ionic compounds: higher charges and smaller ion radii generally yield stronger bonds; use charge product and ion sizes to compare.
• For metallic bonds: higher cation charges and more delocalized electrons strengthen the bond; cation radius also plays a role.
• For nonpolar molecules: larger size and greater surface area → stronger LDF and higher mp/bp.
• For polar molecules: compare DDF vs LDF by size and polarity; hydrogen bonding can override other forces (HF example).

Flow Chart and Practice Context (Study Aid)

• A flow chart exists to determine whether a substance is metallic, ionic, covalent-network, or molecular, and to identify the IPF (LDF, DDF, H-bonding, etc.).
• Practice problems help apply these rules to rank seven substances from strongest to weakest IPF.

What’s Next

• Next presentation covers Properties Related to Interparticle Force Strength.
• For extra practice: Unit 10 – Part 3: Properties Related to Interparticle Force Strength and Q&A Sheet #17, problems 16-33.