Why This Chemistry Question Trips People Up
Here's a question that shows up constantly in organic chemistry courses, on exams, and in study forums: which of the following compounds can form intermolecular hydrogen bonds?* It looks simple at first glance. Most students think they know the answer. Then they look at the answer key, realize they got it wrong, and walk away quietly wondering what they missed And that's really what it comes down to. Turns out it matters..
The thing is, hydrogen bonding isn't actually that complicated once you strip away the textbook fog. It's just a specific type of attraction between molecules, and only certain molecules can do it. The trick is knowing which ones — and why the rules work the way they do.
The official docs gloss over this. That's a mistake.
Let me walk you through how to actually think about this question so it stops feeling like guesswork Still holds up..
What Intermolecular Hydrogen Bonds Actually Are
A hydrogen bond is what happens when a hydrogen atom that's already bonded to a highly electronegative atom gets pulled toward another electronegative atom on a neighboring molecule. It's not a "real" bond in the covalent sense — it's an intermolecular attraction, meaning it happens between* molecules, not within* them.
For a hydrogen bond to form between two molecules, you need two things:
- A hydrogen atom that's attached to nitrogen (N), oxygen (O), or fluorine (F) on one molecule
- A lone pair of electrons on N, O, or F on another molecule
That's it. No nitrogen, no oxygen, no fluorine? No hydrogen bond. This is the rule that eliminates most of the wrong answers right away.
One thing students often miss: it's not enough for a molecule to contain* an N, O, or F. Think about it: that atom has to be bonded directly to a hydrogen that's exposed enough to interact. Some molecules have oxygen or nitrogen buried in a structure where it's not available, and that changes the answer Simple as that..
The Big Three: N, O, F
When you're scanning a list of compounds looking for hydrogen bonders, your eyes should jump straight to anything with N–H, O–H, or F–H bonds. If you see one of those, it's almost certainly capable of intermolecular hydrogen bonding. If you don't see one, it almost certainly isn't.
Fluorine is a fun case. F–H bonds are found in hydrogen fluoride (HF), which is famously one of the strongest hydrogen-bonding systems around. But you rarely see fluorine in organic compounds outside of a few specialized cases, so in most exam questions, you're really looking for O–H or N–H Simple, but easy to overlook..
Why It Matters (Especially in Real-World Chemistry)
Why do instructors love this question so much? So because hydrogen bonding explains a huge amount of physical behavior. Because of that, water's high boiling point. Alcohols being more viscous than similar-sized alkanes. Why DNA's two strands stick together. Why sugar dissolves in water but oil doesn't Worth keeping that in mind..
This is where a lot of people lose the thread.
When you understand hydrogen bonding, a lot of chemistry starts to make intuitive sense. Compounds that can hydrogen bond tend to have higher boiling points, higher melting points, and greater water solubility than compounds of similar size that can't. It's not magic — it's just electrostatics doing their thing on a molecular level.
The flip side? On the flip side, get this wrong, and you'll mispredict boiling points, misjudge solubility, and miss easy points on exams. It's one of those foundational ideas that pays off everywhere.
How to Actually Answer the Question Step by Step
Let's go through a real example. Say you're given a list like this:
- Ethanol (CH₃CH₂OH)
- Dimethyl ether (CH₃OCH₃)
- Methane (CH₄)
- Ammonia (NH₃)
- Trimethylamine (N(CH₃)₃)
Your job is to figure out which of these can form intermolecular hydrogen bonds with themselves* (or with other similar molecules). Here's the move:
Step 1: Look for O–H, N–H, or F–H bonds
- Ethanol: yes, O–H. It's a hydrogen bonder.
- Dimethyl ether: has oxygen, but no O–H bond. The oxygen only has C–H neighbors. It can't donate a hydrogen bond. It can accept* one, but that's a different question (more on that below).
- Methane: no N, O, or F at all. Out.
- Ammonia: N–H bonds present. Hydrogen bonder.
- Trimethylamine: has nitrogen, but no N–H bond. Same situation as dimethyl ether — it can accept but not donate.
So ethanol and ammonia can form intermolecular hydrogen bonds. The others cannot, at least not with identical molecules No workaround needed..
Step 2: Don't Forget the "With What" Question
Here's where students lose points. Some questions ask whether a compound can hydrogen bond with water* (or with another specific molecule). That's a different and easier question.
Dimethyl ether can absolutely hydrogen bond with water — the oxygen in the ether acts as an acceptor, and water provides the O–H donor. Same with trimethylamine Easy to understand, harder to ignore..
But the classic phrasing — "can it form intermolecular hydrogen bonds" — usually means with itself*. And for that, the molecule needs to be both a donor (have an N–H, O–H, or F–H) and an acceptor (have a lone pair on N, O, or F). If it's only one or the other, no self-association through hydrogen bonding.
Step 3: Watch the Borderline Cases
Some molecules look like they should hydrogen bond but can't for structural reasons. Worth adding: tert-butanol, for instance, can definitely hydrogen bond — it's got an O–H, even if it's crowded. But bulky groups can sometimes make it harder for molecules to actually line up the right way, which affects how strong the network is, even if the bonding is technically possible Which is the point..
Acetic acid is an interesting one. It can hydrogen bond, and it does so so enthusiastically that it actually forms dimers* — two acetic acid molecules pairing up through two simultaneous hydrogen bonds. This is a special case worth knowing about.
Common Mistakes That Cost Points
Mistake 1: Counting any O or N as enough
Just because a molecule has oxygen doesn't mean it can hydrogen bond with itself. And diethyl ether, ketones, esters — all have oxygen, all lack an O–H bond. They can't form hydrogen bonds with other identical molecules Nothing fancy..
Mistake 2: Forgetting that "intermolecular" means between molecules
Hydrogen bonds inside a single molecule (intramolecular) are a different thing. They exist and matter in some contexts, but the question almost always asks about intermolecular.
Mistake 3: Ignoring the size and shape factor
Even when hydrogen bonding is possible, the strength depends on geometry. So linear molecules pack differently than branched ones. Two molecules that want to hydrogen bond have to physically get close enough for it to happen, and bulky groups can get in the way.
Mistake 4: Mixing up intramolecular and intermolecular effects
A molecule with an internal hydrogen bond might have a lower* boiling point than you'd predict, because its hydrogens are busy bonding internally and aren't available to bond with neighboring molecules. This is a real phenomenon and shows up in things like ortho*-nitrophenol Less friction, more output..
Practical Tips for Getting It Right Every Time
When you see this question on a test, here's what I'd actually do:
- Scan for O–H, N–H, F–H. If you see it, circle the compound. It's almost always a yes.
- If there's no O, N, or F in the molecule, it's an immediate no. Alkanes, alkenes, alkynes, simple halides without those atoms — none of them hydrogen bond with themselves.
- If there's an N or O but no H attached, it's a partial yes. The molecule can accept hydrogen bonds but not donate them, so it can't self-associate.
- Watch for sneaky questions. "With water" and "with itself" are not the same question. Read the wording carefully.
- Practice with real examples. The more structures you look at, the faster this becomes pattern recognition. After a while, you'll spot hydrogen bonders instantly.
FAQ
Can molecules with only C–H bonds form hydrogen bonds?
No. Plus, carbon isn't electronegative enough to create a strong enough dipole for a true hydrogen bond. C–H bonds can participate in very weak "hydrogen bond-like" interactions in special cases, but in a standard chemistry course, you treat C–H as a non-hydrogen-bonder.
Does water always hydrogen bond with everything?
Water can hydrogen bond with anything that has a lone pair on N, O,
F, or an O–H/N–H bond. That said, the strength and geometry of that bond will vary dramatically. A molecule like ammonia (NH₃) will form a strong hydrogen bond with water, while a simple ether like diethyl ether will form a much weaker one, primarily through the oxygen's lone pairs Surprisingly effective..
Is hydrogen bonding the same as a dipole-dipole interaction?
Hydrogen bonding is a special, strong type* of dipole-dipole interaction. All hydrogen bonds are dipole-dipole interactions, but not all dipole-dipole interactions are hydrogen bonds. The key is the involvement of a hydrogen atom bonded to a highly electronegative atom (N, O, F), which creates a sufficiently large partial positive charge on the hydrogen to attract a lone pair on another electronegative atom.
What about molecules with multiple functional groups?
We're talking about where it gets tricky. A molecule like glucose, with multiple –OH groups, is a prolific hydrogen bonder, both with itself and with water. Consider this: a molecule like an amino acid, with both an amine (–NH₂) and a carboxylic acid (–COOH) group, will have complex intramolecular and intermolecular hydrogen bonding networks. The rule of thumb still applies: identify all the O–H, N–H, and F–H bonds, and consider both their potential to donate and their lone pairs' potential to accept.
Conclusion
Mastering the concept of hydrogen bonding isn't about memorizing a list of molecules. That's why it's about developing a simple, systematic approach to molecular structure. This skill does more than just help you answer a specific question; it builds a foundational understanding of why water is a liquid at room temperature, why DNA has its double-helix structure, and why proteins fold into their specific shapes. By focusing on the presence of the key electronegative atoms—nitrogen, oxygen, and fluorine—and the crucial hydrogen atoms attached to them, you can reliably predict a substance's ability to engage in these vital interactions. It is, without exaggeration, one of the most important concepts bridging molecular structure and macroscopic properties in chemistry.
