Unit 3 · Properties of Substances and Mixtures

AP Chemistry Hydrogen Bonding

What a hydrogen bond is, how to identify donors and acceptors, how strong it really is, and how it explains boiling points, solubility, and floating ice, in four worked examples.

Hydrogen bonding is the intermolecular force that shows up most often in AP free-response explanations. The intermolecular forces overview covers the two-part test for when it occurs; this page goes further, into which molecule gives and which receives, how strong the attraction is, and the properties of water and other substances that only hydrogen bonding can explain.

What a Hydrogen Bond Is and Why It Forms

A hydrogen bond is an attraction between a hydrogen atom covalently bonded to N, O, or F in one molecule and a lone pair on N, O, or F in a neighboring molecule. Despite the name, it is an intermolecular attraction, not a chemical bond.

It is unusually strong for an intermolecular force for two reasons. N, O, and F are very electronegative, so they pull electron density away from the hydrogen and leave it with a large partial positive charge. And hydrogen has no inner electrons, so that small positive center can get very close to a lone pair on the neighbor. By Coulomb’s law, a large charge at a short distance gives a strong attraction.

Hydrogen Bond Donors and Acceptors

• Donor: the molecule that supplies the H, which must be bonded to N, O, or F (O–H in water and alcohols, N–H in ammonia and amines, H–F).
• Acceptor: the molecule that supplies the lone pair on N, O, or F.
• Both: water, alcohols, ammonia, and HF can donate and accept, so they hydrogen bond with themselves.
• Acceptor only: acetone, ethers, and esters have O with lone pairs but no O–H, so they cannot hydrogen bond with themselves but can accept hydrogen bonds from water.

How Strong a Hydrogen Bond Is

A hydrogen bond between water molecules is roughly 20 kJ/mol. The covalent O–H bond inside a water molecule is about 463 kJ/mol, more than twenty times stronger. That gap is why boiling water separates whole H₂O molecules without breaking them apart into atoms.

Among the attractions between neutral molecules, though, hydrogen bonding is usually the strongest, which is why it so often decides boiling point comparisons between molecules of similar size. For large molecules, London dispersion forces can add up to more, so always compare the whole picture.

What Hydrogen Bonding Does for Water

• High boiling point: water boils at 100°C, while H₂S, a heavier molecule that cannot hydrogen bond, boils near −60°C.
• High specific heat: energy goes into disrupting hydrogen bonds, so water warms slowly; see the specific heat of water.
• Ice floats: hydrogen bonds lock ice into an open lattice that is less dense than the liquid (Example 4).
• Strong solvent for polar substances: water hydrogen bonds to alcohols, sugars, and other molecules with O or N.

Worked Example 1: Finding the Donor and Acceptor Between Ethanol and Water

Ethanol (CH₃CH₂OH) dissolves in water. Describe two different hydrogen bonds that can form between an ethanol molecule and a water molecule.

Ethanol as donor: the H of ethanol’s O–H is attracted to a lone pair on water’s O.
Water as donor: an H of water is attracted to a lone pair on ethanol’s O.
Not possible: the H atoms bonded to carbon in ethanol cannot donate, because C is not electronegative enough.

On a diagram, draw each hydrogen bond as a dashed line from the H to a lone pair on the other molecule’s oxygen, and keep the solid covalent O–H line separate. A common error is drawing the dashed line to the oxygen the H is already bonded to.

Worked Example 2: Explaining Why Ethanol Boils Far Above Dimethyl Ether

Ethanol (CH₃CH₂OH) and dimethyl ether (CH₃OCH₃) share the formula C₂H₆O, yet ethanol boils at 78°C and dimethyl ether at −24°C. Explain.

Same size: the molecules have the same molar mass and the same number of electrons, so their London dispersion forces are similar.
Different attractions: ethanol has an O–H bond, so its molecules hydrogen bond to one another. Dimethyl ether’s H atoms are all bonded to C, so it has only dipole-dipole and dispersion forces.
Conclusion: more energy is needed to separate ethanol molecules, so ethanol has the much higher boiling point.

Fig. 1 shows the same pattern for acetic acid (118°C) and methyl formate (32°C), another isomer pair. Comparing isomers isolates hydrogen bonding, because size and formula cancel out. The vapor pressure consequences are in vapor pressure.

Worked Example 3: Explaining Why Acetone Dissolves in Water

Acetone, (CH₃)₂CO, cannot hydrogen bond with itself, yet it mixes with water in all proportions. Explain.

Acetone alone: no H is bonded to N, O, or F, so acetone molecules cannot donate hydrogen bonds to each other.
With water: acetone’s carbonyl oxygen has lone pairs, so it can accept hydrogen bonds from water’s O–H.
Conclusion: strong acetone–water attractions replace some water–water hydrogen bonds, so the two mix readily.

The donor-acceptor distinction is the key: a molecule needs only one of the two roles to hydrogen bond with water.

Worked Example 4: Explaining Why Ice Floats

Ice has a density of 0.917 g/cm³, and liquid water at 0°C has a density of about 1.00 g/cm³. Explain at the particle level why ice floats.

In ice: each water molecule hydrogen bonds to four neighbors in a rigid, open, hexagonal lattice that holds the molecules relatively far apart.
In liquid water: hydrogen bonds constantly break and re-form, so the molecules can pack more closely.
Conclusion: the same mass takes up more volume as ice, so ice is less dense and floats.

Most substances are denser as solids, so ice is an exception, and it is a direct result of hydrogen bonding. For the density calculation itself, see density.

Five Hydrogen Bonding Practice Questions With Answers

1. Which of these hydrogen bond with themselves: CH₄, NH₃, HF, CH₃Cl?
   Show answer

   NH₃ and HF. CH₄ and CH₃Cl have no H bonded to N, O, or F.
2. Can dimethyl ether form hydrogen bonds with water? Explain.
   Show answer

   Yes, as an acceptor. Its oxygen lone pairs accept hydrogen bonds from water’s O–H, even though dimethyl ether cannot donate.
3. Why does HF boil at a much higher temperature than HCl, even though HCl has more electrons?
   Show answer

   HF molecules hydrogen bond to one another (H bonded to F), while HCl has only dipole-dipole and dispersion forces.
4. When water boils, which bonds or attractions are broken?
   Show answer

   Hydrogen bonds and other intermolecular attractions between molecules; the covalent O–H bonds inside each molecule stay intact.
5. Rank by boiling point: CH₃OH, CH₃CH₂OH, CH₃OCH₃.
   Show answer

   CH₃CH₂OH > CH₃OH > CH₃OCH₃. Both alcohols hydrogen bond, and ethanol is larger; the ether cannot hydrogen bond with itself.

Common Hydrogen Bonding Mistakes

• Calling a hydrogen bond a covalent bond. It is an intermolecular attraction, about twenty times weaker than O–H.
• Counting any molecule with H. The H must be bonded to N, O, or F; C–H does not count.
• Drawing the dashed line to the wrong atom. It runs from the H to a lone pair on the neighboring N, O, or F.
• Saying boiling breaks O–H bonds. Boiling separates whole molecules.
• Ignoring acceptor-only molecules. Acetone and ethers can still hydrogen bond with water.
• Assuming hydrogen bonding always wins. In large molecules, dispersion forces can outweigh it.

Hydrogen bonding explanations are worth points in nearly every Unit 3 free-response question, and Unit 3 is the heaviest unit on the exam. To see how stronger Unit 3 answers would move your overall estimate, try the AP Chem Score Calculator.

Related Resources

• AP Chem Score Calculator
• Unit 3 Review: Properties of Substances and Mixtures
• AP Chemistry Intermolecular Forces
• AP Chemistry London Dispersion Forces
• AP Chemistry Vapor Pressure
• AP Chemistry Molecular Polarity
• AP Chemistry Specific Heat and Calorimetry
• AP Chemistry Study Guide

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