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?Because of that, * 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.
The thing is, hydrogen bonding isn't actually that complicated once you strip away the textbook fog. And 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.
Let me walk you through how to actually think about this question so it stops feeling like guesswork.
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 Most people skip this — try not to..
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 Worth keeping that in mind. Nothing fancy..
One thing students often miss: it's not enough for a molecule to contain* an N, O, or F. Because of that, 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.
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. Think about it: 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 Most people skip this — try not to..
Not the most exciting part, but easily the most useful.
Fluorine is a fun case. In real terms, 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.
Why It Matters (Especially in Real-World Chemistry)
Why do instructors love this question so much? Why DNA's two strands stick together. On top of that, water's high boiling point. In real terms, alcohols being more viscous than similar-sized alkanes. And because hydrogen bonding explains a huge amount of physical behavior. Why sugar dissolves in water but oil doesn't Took long enough..
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? 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 Simple, but easy to overlook..
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 And it works..
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.
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. 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.
Acetic acid is an interesting one. Which means 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 Simple, but easy to overlook. Practical, not theoretical..
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. This leads to diethyl ether, ketones, esters — all have oxygen, all lack an O–H bond. They can't form hydrogen bonds with other identical molecules Small thing, real impact. And it works..
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 That's the part that actually makes a difference..
Mistake 3: Ignoring the size and shape factor
Even when hydrogen bonding is possible, the strength depends on geometry. 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 Not complicated — just consistent..
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.
This is the bit that actually matters in practice Easy to understand, harder to ignore..
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. 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 The details matter here..
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. Still, 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.
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 Turns out it matters..
What about molecules with multiple functional groups?
This is where it gets tricky. Here's the thing — a molecule like glucose, with multiple –OH groups, is a prolific hydrogen bonder, both with itself and with water. 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 Not complicated — just consistent. Nothing fancy..
Conclusion
Mastering the concept of hydrogen bonding isn't about memorizing a list of molecules. 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. Worth adding: 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. Think about it: it's about developing a simple, systematic approach to molecular structure. It is, without exaggeration, one of the most important concepts bridging molecular structure and macroscopic properties in chemistry Not complicated — just consistent. Less friction, more output..
Most guides skip this. Don't.
