Hydrogen Bonding Misconceptions

The name “hydrogen bond” is unhelpful in two ways. It contains the word hydrogen, which tempts students to look for it wherever a hydrogen atom appears, and it contains the word bond, which suggests something as strong as the bonds inside a molecule. Neither is quite right. A hydrogen bond is a specific kind of attraction, with specific requirements, and it sits between ordinary intermolecular forces and true chemical bonds in strength.

The eight misconceptions below are the ones that cost the most marks. Each comes with the reason it’s persuasive and the chemistry that replaces it. For the basics, start with hydrogen bonding; for the wider family of forces, see intermolecular force myths.

Misconception 1: “A hydrogen bond is the bond between hydrogen and oxygen inside a water molecule”

Why it’s believed. Water is H₂O, it has bonds to hydrogen, and it’s the standard example of hydrogen bonding. Putting those facts together, students label the O–H bonds in the molecule as “hydrogen bonds”.

The real picture. The O–H bonds inside a water molecule are covalent: a shared pair of electrons between oxygen and hydrogen. A hydrogen bond is the attraction between the hydrogen of one molecule and a lone pair on an oxygen of a different molecule (or a different part of a large molecule). In diagrams, the covalent bond is a solid line and the hydrogen bond is a dashed or dotted line: O–H···O.

The strengths are very different. An O–H covalent bond takes roughly 460 kJ/mol to break. A hydrogen bond between water molecules is closer to 20 kJ/mol. The full comparison is in hydrogen bond vs covalent bond.

Misconception 2: “Any molecule containing hydrogen can form hydrogen bonds”

Why it’s believed. The name suggests that hydrogen is the only ingredient.

The real picture. Two things are needed:

1. A donor: a hydrogen atom covalently bonded to nitrogen, oxygen or fluorine. These three are electronegative enough to pull electron density away from the hydrogen, leaving it with a significant partial positive charge. Because hydrogen has no inner electrons, that exposed nucleus can get very close to a neighbouring lone pair.
2. An acceptor: a lone pair on an N, O or F atom, usually in a neighbouring molecule.

Methane, CH₄, has four hydrogens and no hydrogen bonding at all. Carbon and hydrogen have similar electronegativities, so the C–H bond is only slightly polar and the hydrogen isn’t positive enough. That’s why methane boils at about −162 °C while water, a molecule of similar size, boils at 100 °C.

Chemists do describe weak C–H···O interactions in some crystals and proteins, but these are much weaker than classic hydrogen bonds, and school and university courses do not count C–H as a hydrogen-bond donor. For exams, stick to H on N, O or F.

Misconception 3: “Hydrogen molecules hydrogen bond to each other”

Why it’s believed. H₂ is made entirely of hydrogen, so surely it’s the best example?

The real picture. In H₂, both atoms have exactly the same electronegativity, so the bond is completely non-polar. There’s no δ+ hydrogen and no lone pair anywhere. The only attraction between H₂ molecules is a very weak London force, which is why hydrogen boils at about −253 °C, one of the lowest boiling points of any substance. You can compare the data on the hydrogen element page. Hydrogen bonding depends on what the hydrogen is attached to, not on how many hydrogens there are.

Misconception 4: “A molecule must be able to both donate and accept to take part”

Why it’s believed. The textbook examples (water, ammonia, alcohols, HF) all do both, so students assume that’s a requirement.

The real picture. Some molecules can only accept. Propanone (acetone), CH₃COCH₃, has an oxygen with lone pairs but no hydrogen on N, O or F. Its molecules can’t hydrogen bond to each other, which is part of why it’s so volatile (it boils at about 56 °C). But put propanone in water and water molecules can donate hydrogen bonds to propanone’s oxygen. That’s why the two liquids mix completely.

Ethers behave the same way. So when you’re asked whether hydrogen bonding occurs, look at the pair of molecules involved: a donor from one and an acceptor from the other is enough. See bonding and solubility for how this shapes what dissolves in what.

Misconception 5: “All hydrogen bonds have the same strength”

Why it’s believed. Textbooks often quote a single number, which makes the hydrogen bond sound like a fixed quantity.

The real picture. Typical hydrogen bonds fall across a range of roughly 5 to 40 kJ/mol, depending on the atoms involved, the charges nearby and the geometry. They’re strongest when the three atoms (donor, H and acceptor) lie close to a straight line, and they weaken as that line bends. Hydrogen bonds involving ions, such as those in some charged species, can be stronger still.

This range explains a famous puzzle. The F–H···F hydrogen bond is individually stronger than O–H···O, because fluorine is the most electronegative element. Yet water (bp 100 °C) boils far higher than hydrogen fluoride (bp about 20 °C). The answer is counting: each water molecule has two hydrogens and two lone pairs, so it can take part in up to four hydrogen bonds. Each HF molecule has three lone pairs but only one hydrogen, so on average it forms far fewer hydrogen bonds than water does. The total matters more than the strength of any single interaction. Why water has a high boiling point goes through the numbers.

Misconception 6: “Hydrogen bonds are weak, so they don’t really matter”

Why it’s believed. Next to covalent bonds, 20 kJ/mol does look small.

The real picture. Weak individually, hydrogen bonds are powerful in large numbers, and their weakness is often exactly what makes them useful.

• Water’s physical properties. Without hydrogen bonding, a molecule as small as H₂O would be expected to boil well below 0 °C, as hydrogen sulfide does (about −60 °C). Liquid water on Earth depends on hydrogen bonds.
• DNA. The two strands of the double helix are held together by hydrogen bonds between base pairs. They’re strong enough to keep the helix stable, but weak enough for enzymes to unzip it during copying. See hydrogen bonding in DNA.
• Proteins. Helices and sheets in proteins are held in shape by hydrogen bonds along the backbone.
• Everyday materials. Paper and cotton owe much of their strength to hydrogen bonds between cellulose chains.

Misconception 7: “Ice floats because it has air trapped in it”

Why it’s believed. Ice cubes from a freezer often look cloudy, and the cloudiness is air bubbles. It seems a reasonable explanation for why ice is less dense than water.

The real picture. Even perfectly clear, bubble-free ice floats. In ice, each water molecule is hydrogen bonded to four neighbours in a regular tetrahedral arrangement. That fixed geometry holds the molecules further apart than they are, on average, in the liquid, creating an open lattice with a lot of empty space. When ice melts, part of this network collapses and the molecules can pack more closely, so the liquid is denser.

Ice has a density of about 0.92 g/cm³, compared with about 1.00 g/cm³ for liquid water. For most substances the solid sinks in its own liquid; water is unusual, and hydrogen bonding is the reason. This is why lakes freeze from the top down, insulating the water below.

Misconception 8: “Liquid water has no hydrogen bonds; they only exist in ice”

Why it’s believed. Diagrams of ice show a neat hydrogen-bonded network, while diagrams of liquid water often show scattered molecules. Melting is then described as “breaking the hydrogen bonds”.

The real picture. Liquid water is still extensively hydrogen bonded. The difference is that the network in the liquid is constantly changing: individual hydrogen bonds break and re-form extremely quickly, so molecules can slide past each other. Melting ice breaks only a fraction of the hydrogen bonds, which is why the enthalpy of fusion of water (about 6 kJ/mol) is much smaller than its enthalpy of vaporisation (about 41 kJ/mol). Boiling is the step that separates molecules almost completely, and that’s where most of the hydrogen bonding is overcome.

Drawing hydrogen bonds for marks

Many exam answers lose marks on diagrams rather than ideas. A good hydrogen-bond diagram shows:

• the δ+ on the hydrogen and δ− on the N, O or F atoms;
• at least one lone pair on the acceptor atom;
• a dashed line from the hydrogen to that lone pair (not to another hydrogen, and not to a carbon);
• the three atoms donor–H···acceptor in a roughly straight line.

Key takeaways

• A hydrogen bond is an attraction between molecules (or parts of a large molecule), not the covalent O–H bond inside one.
• It needs H bonded to N, O or F as the donor, and a lone pair on N, O or F as the acceptor.
• C–H is not a classic donor; methane and H₂ show no hydrogen bonding.
• Some molecules, such as propanone, can only accept, but they can still hydrogen bond with water.
• Strength varies, roughly 5–40 kJ/mol; the number of hydrogen bonds per molecule explains why water out-boils HF.
• Ice floats because hydrogen bonding creates an open lattice, not because of trapped air.
• Liquid water is still heavily hydrogen bonded; the bonds constantly break and re-form.