Rank The Following Compounds In Decreasing Order Of Boiling Point

You're staring at a chemistry problem set. Four compounds. One question: rank them in decreasing order of boiling point.

Your brain immediately reaches for molecular weight. Heavier molecules boil higher, right?

Not always. And that's exactly where most students lose points.

What Actually Determines Boiling Point

Boiling point isn't about how heavy a molecule is. Now, it's about how hard it is to pull molecules apart from each other. That's intermolecular forces — IMFs if you're in a hurry.

The hierarchy is straightforward once you stop overthinking it:

Hydrogen bonding > Dipole-dipole > London dispersion forces

But here's where it gets messy. London forces exist in everything*. Because of that, they're the baseline. And they scale with polarizability — which roughly tracks with electron count and molecular surface area. So a massive nonpolar molecule can absolutely out-boil a tiny polar one.

I've seen students argue that acetone (58 g/mol) should boil higher than hexane (86 g/mol) because acetone has a dipole. Also, 69°C vs 56°C. Hexane wins. The dispersion forces in hexane's larger electron cloud simply overpower acetone's dipole-dipole interactions.

The Hydrogen Bond Trump Card

If a molecule has H bonded to N, O, or F, it hydrogen bonds. Same electron count. But neon (20 g/mol) boils at -246°C. Water (18 g/mol) boils at 100°C. That's usually the dominant factor. Hydrogen bonding changes everything.

But — and this matters — not all H-bonders are equal. Worth adding: primary amines hydrogen bond. So do alcohols. Think about it: carboxylic acids dimerize, effectively doubling their molecular weight in the liquid phase. That's why acetic acid (60 g/mol) boils at 118°C while propanol (60 g/mol) hits 97°C.

Molecular Shape Matters More Than Textbooks Admit

Take pentane isomers. So n-Pentane (linear) boils at 36°C. Consider this: neopentane (spherical) boils at 9. 5°C. Same formula. Same molecular weight. Still, same IMFs in theory*. But neopentane's compact shape reduces surface contact between molecules. And weaker London forces. Lower boiling point.

This shows up constantly in exam questions. They'll give you constitutional isomers and expect you to recognize that branching lowers boiling point.

How to Approach Any Ranking Problem

Don't guess. Follow a decision tree.

Step 1: Identify Functional Groups

Scan for H-bond donors/acceptors. Alcohols, amines, carboxylic acids, amides — these sit at the top of any ranking unless something weird is happening.

Step 2: Check for Ionic Character

Salts don't "boil" in the normal sense — they decompose. But if you're comparing something like sodium acetate to organic molecules, the ionic compound wins by a landslide. Lattice energy dwarfs everything else.

Step 3: Compare Hydrogen Bonders Among Themselves

Carboxylic acids > alcohols ≈ amines > amides (sometimes — amides are weird because resonance reduces N-H polarity but they also have strong dipoles). Within each class: more H-bond sites = higher boiling point. Which means ethylene glycol (two OH groups) boils at 197°C. Ethanol (one OH) boils at 78°C.

Step 4: For Non-H-Bonders, Compare Dipole Moments

Polar molecules beat nonpolar ones of similar size. But "similar size" is doing heavy lifting here. And acetonitrile (41 g/mol, strong dipole) boils at 82°C. Butane (58 g/mol, nonpolar) boils at -0.In practice, 5°C. The dipole wins despite lower molecular weight.

Step 5: For Nonpolar Compounds, It's All About Size and Shape

More electrons = stronger London forces. More surface area = stronger London forces. Plus, linear > branched. That's the whole game.

Worked Examples That Actually Represent What You'll See

Example Set 1: CH₃CH₂OH, CH₃OCH₃, CH₃CH₂CH₃

Ethanol, dimethyl ether, propane.

Ethanol has an OH group. Hydrogen bonding. Winner.

Dimethyl ether has a dipole (C-O-C) but no H-bond donor. Middle.

Propane is nonpolar, 44 g/mol. Loser.

Ranking: Ethanol (78°C) > Dimethyl ether (-24°C) > Propane (-42°C).

Example Set 2: CH₃CH₂CH₂CH₂OH, CH₃CH₂CH₂CHO, CH₃CH₂COCH₃, CH₃CH₂CH₂CH₂CH₃

1-Butanol, butanal, 2-butanone, pentane.

1-Butanol: alcohol, H-bonds. Top.

Butanal and 2-butanone: both polar carbonyls, similar molecular weights (72 vs 72). Butanal has a slightly larger dipole moment (aldehyde vs ketone) but the difference is small. In practice they're close — butanal 75°C, 2-butanone 80°C. The ketone wins slightly due to better charge separation in the carbonyl.

Pentane: nonpolar, 72 g/mol. 36°C. Bottom.

Ranking: 2-Butanone ≈ Butanal > 1-Butanol? Wait.

Stop. Re-read.

1-Butanol is 74 g/mol and hydrogen bonds. But it boils at 118°C. On top of that, i just said it's top. But then I wrote the ranking wrong in my head.

1-Butanol (118°C) > 2-Butanone (80°C) > Butanal (75°C) > Pentane (36°C).

The alcohol wins despite similar molecular weight because hydrogen bonding dominates. This is the trap — students see similar MW and forget the IMF hierarchy.

Want to learn more? We recommend identify the values from the graph. amplitude period and energy needed to start a chemical reaction for further reading.

Example Set 3: HF, HCl, HBr, HI

Classic trick question.

HF hydrogen bonds. The others don't (Cl, Br, I aren't electronegative enough). So HF should be highest.

But HI is heavy* (128 g/mol) with massive electron clouds. Its London forces are enormous.

Actual boiling points: HF (19.5°C), HCl (-85°C), HBr (-67°C), HI (-35°C).

HF wins. Here's the thing — london forces scaling with size eventually overtook dipole-dipole. But notice the trend in the hydrogen halides: boiling point increases* down the group despite decreasing polarity. This is the only group where this happens — and it's a favorite exam topic.

Common Mistakes That Cost Points

Mistake 1: Confusing Intramolecular and Intermolecular

Covalent bond strength has nothing* to do with boiling point. Because of that, boiling doesn't break bonds. It overcomes attractions between* molecules.

stronger than O-H bonds. Wrong. Boiling point is about the forces between molecules*, not the bonds within them.

Mistake 2: The "Molecular Weight is Everything" Fallacy

Yes, for nonpolar compounds, MW is king. But the moment you have a polar functional group, the hierarchy kicks in. Worth adding: a 100 g/mol alcohol will always beat a 200 g/mol alkane. Always.

Mistake 3: Forgetting That Hydrogen Bonding Requires a Donor AND an Acceptor

A ketone has a lone pair (acceptor) but no H attached to N, O, or F (donor). In real terms, it cannot hydrogen bond with itself*. It can only accept H-bonds from a solvent like water or alcohol. This is why ketones have higher boiling points than ethers of similar weight, but lower than alcohols.

The Bottom Line

Predicting boiling points is a three-step thought process:

Identify the strongest IMF present. Hydrogen bonding > Dipole-dipole > London dispersion forces.
Compare within the same IMF category. For H-bonders, more OH groups = higher BP. For dipole-dipole, larger dipole moment = higher BP. For London forces, larger surface area and molecular weight = higher BP.
When comparing across categories, the stronger IMF always wins, regardless of molecular weight.

This isn't just for passing a test. These principles govern how chemists separate compounds in a lab, how drugs are designed to interact with biological targets, and how industries purify chemicals on a massive scale. The invisible forces between molecules are the silent directors of the molecular world.

Now, go forth and rank with confidence.

When Things Get Complicated: Mixed Functional Groups

Real-world molecules rarely present a single, clear intermolecular force. Both have the same molecular formula (C₂H₆O), yet ethanol boils at 78°C while dimethyl ether boils at -23°C. Practically speaking, consider a molecule like ethanol (CH₃CH₂OH) versus dimethyl ether (CH₃OCH₃). On the flip side, the difference? Ethanol can hydrogen bond with itself due to its -OH group, while dimethyl ether lacks this capability.

This principle extends to more complex scenarios. A carboxylic acid like acetic acid (MW 60 g/mol) will boil higher than pentane (MW 72 g/mol) because the carboxylic acid can form extensive hydrogen bonding networks, despite being lighter. Still, if you compare acetic acid to pentanol (both capable of hydrogen bonding), then molecular weight becomes the deciding factor.

Temperature Matters: Phase Changes and Critical Points

Boiling point predictions assume standard pressure, but real-world applications often involve varying conditions. Here's the thing — the relationship between pressure and boiling point follows a predictable pattern: increased pressure raises boiling points, while reduced pressure lowers them. This is why pressure cookers work—they force liquid water to temperatures above 100°C, cooking food faster.

Understanding these relationships becomes crucial when dealing with substances near their critical points. Beyond the critical temperature, no amount of pressure will create a liquid phase, fundamentally changing how we predict and manipulate molecular behavior.

Biological Implications

These intermolecular forces aren't confined to test tubes—they govern life itself. Hemoglobin's ability to bind and release oxygen depends on subtle changes in hydrogen bonding and dipole interactions. In practice, protein folding, enzyme specificity, and drug-receptor interactions all hinge on the same principles used to predict boiling points. A drug designer who understands that a small polar group can dramatically increase a molecule's water solubility—or that adding a methyl group might tip the balance toward membrane permeability—is working with the same fundamental concepts.

Final Thoughts: Building Intuition

Mastering boiling point prediction requires developing chemical intuition through practice. Start by identifying functional groups and their associated intermolecular forces, then systematically compare molecular weights and structural features. When you encounter exceptions, dig deeper—understanding why something behaves unexpectedly often reveals the most important insights.

Remember that chemistry builds upon itself. The same forces that determine whether a molecule is a liquid or gas at room temperature also control how that molecule behaves in solution, how it interacts with biological systems, and how we can manipulate it in the laboratory. Mastering these concepts isn't just about acing an exam—it's about understanding the invisible forces that shape our molecular world.

The next time you're asked to rank compounds by boiling point, approach it systematically: identify the strongest intermolecular forces, compare like with like within each category, and trust that hydrogen bonding will typically trump everything else. With practice, these patterns become second nature, freeing your mind to tackle more complex chemical challenges with confidence.