Intermolecular Forces

Intramolecular Forces:
Strong forces within individual molecules that hold atoms together. These are chemical bonds (e.g., covalent, ionic) and determine a substance's chemical properties.

Electronegativity:
An atom's ability to attract shared electrons in a covalent bond. Differences in electronegativity ($\Delta\text{EN}$) create
bond dipoles.

Bond Polarity:

Non-polar Covalent: $\Delta\text{EN} \approx 0$ (e.g., $\text{Cl-Cl}$).

Polar Covalent: 0 < \Delta\text{EN} < 1.7 (e.g., $\text{H-Cl}$).

Ionic: \Delta\text{EN} > 1.7 (e.g., $\text{Na}^+\text{Cl}^-$).

Molecular Polarity (Net Dipole):
An entire molecule can be polar if its individual bond dipoles do not cancel out due to molecular geometry.

Non-polar molecules:
Symmetrical shapes (e.g., linear)
where bond dipoles cancel, even if individual bonds are polar.

Polar molecules:
Asymmetrical shapes (e.g., bent)
where bond dipoles combine to create an overall net dipole moment.

Intermolecular Forces (IMFs):
Weaker attractive forces between separate molecules. These are responsible for a substance's physical properties, such as melting point, boiling point, and solubility. They are much weaker than intramolecular forces.

Dispersion Forces

Also called London forces; present in all molecular substances, regardless of whether they are polar or non-polar.

Arise from the random movement of electrons, creating short-lived temporary dipoles.

Temporary dipole in one molecule induces a dipole in a neighbouring molecule $\rightarrow$ weak electrostatic attraction.

Continually forms, collapses, and reforms elsewhere.

Strength scales with:

Number of electrons / molar mass (larger electron cloud = more easily polarised).

Molecular shape: linear or elongated shapes allow closer contact $\rightarrow$ stronger dispersion; compact/branched $\rightarrow$ weaker dispersion.text{C}$

omplexity of molecules: larger and more complex molecules often exhibit stronger van der Waals forces due to the increased possibilities for temporary dipoles.

Within groups of similar compounds (e.g. halogens), MP/BP rise smoothly with molar mass due to increasing dispersion:

Dipole–Dipole Forces

Require molecules with permanent dipoles (net polarity), which are polar molecules.

The positive end of one molecule attracts the negative end of another.

Produce modest BP/MP elevation vs purely dispersion substances of similar mass.

Hydrogen Bonding

Extreme special case of dipole–dipole; directional, up to $\approx12\,\%$ strength of a C–C covalent bond.

Occurs only with arrangements $\text{H–F}$, $\text{H–O}$, $\text{H–N}$.

Requires:

H covalently bonded to highly electronegative F, O or N $\Rightarrow$ strong $\delta^+$ on H.

Lone pair on another F/O/N (same or different molecule) providing strong $\delta^-$ site.

Greatly enhanced solubility between substances that can H-bond ("like dissolves like").

Physical-Property Trends vs Intermolecular Forces

Phase change requires overcoming intermolecular attractions $\Rightarrow$ stronger forces $\rightarrow$ higher MP/BP.

Order of typical strength (small molecules):

dispersion < dipole–dipole < hydrogen bonding < ionic/metallic/covalent network.

Equilibrium Vapour Pressure

In a closed vessel, rate of evaporation = rate of condensation $\rightarrow$ constant pressure (kPa).

Higher temperature $\Rightarrow$ higher molecular EK $\Rightarrow$ higher vapour pressure for all liquids.

At a fixed T, weaker intermolecular forces $\Rightarrow$ higher vapour pressure.This is because weaker forces allow more molecules to escape the liquid phase into the vapour phase, increasing the number of molecules in the vapour above the liquid.

Normal Boiling Point

Defined where $p_{\text{vapour}} = 101.3\,\text{kPa}$ (standard atmosphere). A substance with low intermolecular forces will have a lower normal boiling point due to its ability to vaporize more easily, as seen in the comparison of different liquids.

Unique Properties of Water

• Highly polar + each molecule can form 4 hydrogen bonds.

• Despite small molar mass, MP $0\,^\circ\text{C}$ & BP $100\,^\circ\text{C}$ (compare $\text{CH}4$, $\text{H}2\text{S}$ — both gases).

• Ice density lower than liquid water ($\approx8\,\%$ expansion on freezing; $\rho{\text{ice}} = 0.918\,\text{g mL}^{-1}$, max $\rho{\text{liquid}} = 1.000\,\text{g mL}^{-1}$ at $4\,^\circ\text{C}$):

  • Open tetrahedral H-bonded lattice in ice occupies more volume.

  • Environmental significance: ice floats, insulates aquatic life.

• Very high surface tension: inward net