DAT General Chemistry: Ranking Intermolecular Forces and Boiling Points

DAT High-Yield Concept

Intramolecular vs. Intermolecular: Intramolecular forces (covalent, ionic) bind atoms within a molecule. Intermolecular forces (IMFs) act between separate molecular species.
IMF Relative Strength Ranking: Ion–Dipole > Hydrogen Bonding > Dipole–Dipole > London Dispersion Forces.
Physical Property Impact: Stronger IMFs require greater thermal energy to break, leading to higher boiling points ( $T_b$ ) and melting points ( $T_m$ ).
Hydride Trend Exceptions: $\ce{H2O}$ , $\ce{HF}$ , and $\ce{NH3}$ exhibit abnormally high boiling points relative to heavier group hydrides due to strong hydrogen bonding.

Intramolecular versus Intermolecular Forces

Intramolecular forces are true chemical bonds (covalent, polar covalent, ionic) that hold atoms together inside a molecule or formula unit. Breaking intramolecular forces involves a chemical reaction and requires substantial energy (typically $\pu{150 kJ/mol}$ to $\pu{1000 kJ/mol}$ ).

Intermolecular forces (IMFs) are non-covalent attractive forces occurring between adjacent molecules. IMFs are weaker (typically $\pu{2 kJ/mol}$ to $\pu{50 kJ/mol}$ ) and dictate physical state transitions (boiling, melting, vapor pressure) without breaking chemical bonds or altering molecular identity.

Ranking Intermolecular Force Strength

To predict physical properties, rank intermolecular interactions from strongest to weakest:

Ion–Dipole Interactions: Attraction between a formal ion and a polar molecule's partial charge (e.g., $\ce{Na+}$ solvated by $\ce{H2O}$ ).
Hydrogen Bonding: A strong, directional dipole–dipole attraction occurring when hydrogen is directly covalently bonded to $\ce{F}$ , $\ce{O}$ , or $\ce{N}$ and interacts with a lone pair on a neighboring $\ce{F}$ , $\ce{O}$ , or $\ce{N}$ .
Dipole–Dipole Interactions: Electrostatic attraction between permanent dipoles of polar molecules.
London Dispersion Forces (LDFs): Transient induced dipole attractions present in all molecules. Their strength scales with polarizability, electron count, and accessible surface area.

London forces in alkanes (Source: Danielchemik · CC BY-SA 3.0 · Wikimedia Commons)

Worked Example: Identifying Predominant IMFs

Problem: Identify the predominant IMF in pure samples of $\ce{CH4}$ , $\ce{CH3Cl}$ , and $\ce{CH3OH}$ .

Step-by-step Solution:

$\ce{CH4}$ : Tetrahedral and nonpolar. Primary interaction: London dispersion forces.
$\ce{CH3Cl}$ : Polar due to the $\ce{C-Cl}$ dipole, but lacks $\ce{H-F}$ , $\ce{H-O}$ , or $\ce{H-N}$ bonds. Primary interaction: Dipole–dipole interactions.
$\ce{CH3OH}$ : Contains a hydroxyl ( $\ce{O-H}$ ) group capable of hydrogen bonding. Primary interaction: Hydrogen bonding.

Exam Trap: Assuming any molecule with hydrogen can form hydrogen bonds. Hydrogen bonding strictly requires $\ce{H}$ covalently bound directly to $\ce{F}$ , $\ce{O}$ , or $\ce{N}$ .

Effect of Intermolecular Force Strength on Boiling and Melting Points

Boiling occurs when kinetic energy overcomes liquid-phase cohesive forces. Stronger IMFs require higher thermal energy to separate molecules into the gas phase, increasing boiling point ( $T_b$ ). Melting point ( $T_m$ ) also increases with stronger IMFs, though solid lattice symmetry and packing efficiency play additional roles.

When comparing boiling points:

Identify the highest-ranking IMF present.
If primary IMFs are identical, compare molecular weights and surface areas; greater polarizability increases dispersion strength.
For structural isomers, straight-chain compounds exhibit greater surface contact and higher boiling points than highly branched spherical isomers.

Graph showing the boiling points of alkanes, alkenes, ethers, halogenoalkanes, aldehydes, ketones, alcohols and carboxylic acids as a function of molar mass. In all cases, boiling points of straight-chain compounds are used. All the curves show approximately the same gradient. Grouping of the curves shows the difference in strength between induced dipole-induced dipole forces (alkanes, alkenes, ethers and halogenoalkanes), permanent dipole-permanent dipole forces which raise the boiling points of aldehydes and ketones, and finally the strongest intermolecular hydrogen bonding in alcohols and carboxylic acids. Carboxylic acids have higher boiling points than alcohols due to the further presence of the polar C=O bond. (Source: Jpta · CC BY-SA 3.0 · Wikimedia Commons)

Binary Hydride Boiling Point Anomalies

In Group 14 to 17 hydrides, boiling points generally increase down a period due to increasing molar mass and electron cloud polarizability. However, Period 2 hydrides $\ce{H2O}$ , $\ce{HF}$ , and $\ce{NH3}$ deviate dramatically with anomalously high boiling points because of intermolecular hydrogen bonding.

Boiling points of the binary hydrides: dispersion trend, hydrogen-bond spikes — Within the group 14 hydrides, which cannot hydrogen bond, the boiling point climbs steadily from methane to stannane as the molecules get larger and more polarizable, so dispersion forces strengthen. The other three series follow the same rise from period 3 onward but start far too high at period 2: water, hydrogen fluoride and ammonia boil hundreds of degrees above where their size alone would put them, because their molecules must also be pulled apart from hydrogen bonds. Water shows the largest jump of the three, since it can form the most hydrogen bonds per molecule.

Worked Example: Ranking Compound Boiling Points

Problem: Rank $\ce{CH4}$ , $\ce{CH3OCH3}$ , $\ce{CH3CH2OH}$ , and $\ce{KCl}$ in order of increasing boiling point.

Step-by-step Solution:

$\ce{CH4}$ : Nonpolar alkane; weak LDFs only (lowest $T_b$ ).
$\ce{CH3OCH3}$ : Polar ether; permanent dipole–dipole interactions.
$\ce{CH3CH2OH}$ : Alcohol; strong intermolecular hydrogen bonding.
$\ce{KCl}$ : Ionic lattice; strong electrostatic ion–ion bonds (highest $T_b$ ).

Correct Ranking: $\ce{CH4}$ < $\ce{CH3OCH3}$ < $\ce{CH3CH2OH}$ < $\ce{KCl}$

Summary Comparison of Intermolecular Forces

Interaction Type	Relative Strength	Structural Requirement	Example Species
Ionic / Ion–Dipole	Strongest (> $\pu{100 kJ/mol}$ )	Full ionic charges in polar medium	$\ce{NaCl(aq)}$ , $\ce{Mg^{2+}}$ in $\ce{H2O}$
Hydrogen Bonding	Strong ( $\pu{10-40 kJ/mol}$ )	$\ce{H}$ bonded to $\ce{F}$ , $\ce{O}$ , or $\ce{N}$	$\ce{H2O}$ , $\ce{HF}$ , $\ce{NH3}$ , $\ce{CH3OH}$
Dipole–Dipole	Moderate ( $\pu{5-25 kJ/mol}$ )	Permanent dipole moment	$\ce{HCl}$ , $\ce{CH3CHO}$ , $\ce{CH3Cl}$
London Dispersion	Weakest ( $\pu{0.05-40 kJ/mol}$ )	All atoms and molecules	$\ce{F2}$ , $\ce{Cl2}$ , $\ce{CH4}$ , $\ce{Ar}$

DAT Mnemonic for IMF Strength: I Hate Dull Lectures!
Ion-Dipole > Hydrogen Bonding > Dipole-Dipole > London Dispersion

Critical Thinking Questions

Why does water ( $\ce{H2O}$ , $T_b = \pu{100 degC}$ ) have a higher boiling point than hydrogen fluoride ( $\ce{HF}$ , $T_b = \pu{19.5 degC}$ ), even though $\ce{F}$ is more electronegative than $\ce{O}$ ?
Comparing diethyl ether ( $\ce{CH3CH2OCH2CH3}$ ) and 1-butanol ( $\ce{CH3CH2CH2CH2OH}$ )—both with molecular formula $\ce{C4H10O}$ and mass $\pu{74.12 g/mol}$ —why is the boiling point of 1-butanol ( $\pu{117 degC}$ ) so much higher than diethyl ether ( $\pu{34.6 degC}$ )?
How does compact spherical symmetry affect boiling point versus melting point when comparing pentane and neopentane (2,2-dimethylpropane)?

References

Sources returned by Google Search grounding during lesson verification.

Intermolecular Forces - Chemistry LibreTexts — chem.libretexts.org
Trends in Boiling Points for p-block Binary Hydrides - Chemistry LibreTexts — chem.libretexts.org

Ranking Intermolecular Forces and Boiling Points

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