Hydrogen Bonding: Why Water Is Special

What Is a Hydrogen Bond?

A hydrogen bond is an electrostatic attraction between a hydrogen atom bonded to a very electronegative atom (N, O, or F) and a lone pair on another N, O, or F atom in a nearby molecule.

It is the strongest type of intermolecular force — roughly 10 times stronger than London dispersion forces, but still about 10 times weaker than a typical covalent bond.

Learning Goals: By the end of this guide, you should be able to:

1. Explain how and why hydrogen bonds form.
2. Identify which molecules can form hydrogen bonds.
3. Use hydrogen bonding to explain water's anomalous properties.
4. Describe the role of hydrogen bonds in biology.
5. Answer comparison and explanation questions confidently.

How Hydrogen Bonds Form

The Three Requirements

For a hydrogen bond to form, you need:

1. A hydrogen atom covalently bonded to N, O, or F (the hydrogen bond donor).
2. A lone pair on a nearby N, O, or F atom (the hydrogen bond acceptor).
3. The molecules must be close enough for the electrostatic attraction to act.

Why Only N, O, and F?

These three elements are the most electronegative in the periodic table (after fluorine at 4.0, oxygen at 3.5, nitrogen at 3.0). When bonded to hydrogen:

• They pull the shared electrons strongly toward themselves → large $\delta^-$ on N/O/F.
• Hydrogen is left with a concentrated $\delta^+$ charge (it has no inner electron shells to shield it).
• The result: an unusually strong dipole-dipole interaction — strong enough to earn its own name.

Can vs. Cannot Form Hydrogen Bonds

Water's Anomalous Properties

Hydrogen bonding explains nearly every unusual property of water. Without it, water would boil at about –80°C and life as we know it could not exist.

1. Anomalously High Boiling Point

The trend $H_2S < H_2Se < H_2Te$ follows increasing molecular size (stronger London forces). But $H_2O$ breaks this trend because hydrogen bonding adds much more intermolecular attraction, requiring far more energy to boil.

2. Ice Is Less Dense Than Water

In liquid water, hydrogen bonds constantly break and reform, allowing molecules to move relatively close together. When water freezes:

• Each molecule forms 4 hydrogen bonds in a fixed, open hexagonal lattice.
• This lattice has more empty space than the liquid.
• Result: ice is less dense than water → ice floats.

This is biologically critical: floating ice insulates lakes and oceans, allowing life to survive beneath during winter.

3. High Specific Heat Capacity

Water absorbs a lot of heat energy before its temperature rises because energy is used to break hydrogen bonds rather than increase kinetic energy. This makes water an excellent thermal buffer for living organisms and climate regulation.

4. High Surface Tension

Water molecules at the surface are pulled inward by hydrogen bonds with molecules below. This creates a "skin" on the water surface, strong enough for insects to walk on.

Hydrogen Bonds in Biology

DNA Base Pairing

The two strands of DNA are held together by hydrogen bonds between complementary base pairs:

• Adenine (A) – Thymine (T): 2 hydrogen bonds
• Guanine (G) – Cytosine (C): 3 hydrogen bonds

This means G–C pairs are stronger than A–T pairs, which affects DNA melting temperature (high G–C content → higher melting temperature).

Protein Folding

Hydrogen bonds between the N–H and C=O groups along the polypeptide backbone create secondary structures:

• α-helices: stabilised by H-bonds between every 4th amino acid.
• β-sheets: stabilised by H-bonds between adjacent strands.

Worked Examples

Example 1: Identifying Hydrogen Bonding

Question: Which of the following can form hydrogen bonds? $CH_3OH$, $CH_3OCH_3$, $CH_3F$, $CH_4$

Solution:

• $CH_3OH$ (methanol): Has an O–H bond → can donate H-bonds. Has lone pairs on O → can accept H-bonds. ✅ Yes
• $CH_3OCH_3$ (dimethyl ether): Has lone pairs on O but no O–H, N–H, or F–H bond → can accept H-bonds from other molecules but cannot donate its own. Partially (acceptor only)
• $CH_3F$: Has lone pairs on F but no F–H bond → acceptor only, no self-hydrogen bonding. Partially
• $CH_4$: No electronegative atom bonded to H → ❌ No

Example 2: Explaining a Boiling Point

Question: Explain why ethanol ($C_2H_5OH$, bp 78°C) has a much higher boiling point than dimethyl ether ($CH_3OCH_3$, bp –24°C) despite having the same molecular formula $C_2H_6O$.

Solution: Ethanol has an O–H group that can form hydrogen bonds with other ethanol molecules. Dimethyl ether has no O–H bond and can only experience weaker dipole-dipole and London forces. The extra energy needed to break hydrogen bonds in ethanol gives it a much higher boiling point.

Example 3: Number of Hydrogen Bonds

Question: How many hydrogen bonds can one molecule of water form?

Solution: Each $H_2O$ molecule has:

• 2 O–H bonds → can donate 2 hydrogen bonds
• 2 lone pairs on oxygen → can accept 2 hydrogen bonds
• Maximum: 4 hydrogen bonds per molecule (2 donated + 2 accepted)

Common Mistakes

1. "Hydrogen bonds are covalent bonds" — No. Hydrogen bonds are intermolecular forces, not chemical bonds. They are electrostatic attractions, roughly 1/10th the strength of a covalent bond.

2. "Any molecule with hydrogen can form hydrogen bonds" — The hydrogen must be bonded to N, O, or F specifically. $CH_4$ and $H_2S$ cannot form hydrogen bonds.

3. "Hydrogen bonds are between atoms in the same molecule" — While intramolecular H-bonds exist in some large molecules, the term usually refers to intermolecular forces between different molecules.

4. Confusing H-bond donors and acceptors — The donor provides the H (from an N–H, O–H, or F–H bond). The acceptor provides the lone pair. A molecule like $CH_3OCH_3$ can accept but not donate.

5. Forgetting that hydrogen bonding includes London forces too — Molecules with H-bonds also have London forces and dipole-dipole interactions. H-bonding is the dominant force, but it's not the only one.

Exam Tips (A-Level / AP / IB)

• When asked to draw hydrogen bonds, use a dashed line between the $\delta^+$ H and the lone pair on N/O/F. Label the partial charges.
• Always mention the three conditions: H bonded to N/O/F, lone pair on nearby N/O/F, close proximity.
• For boiling point comparisons, state: "[molecule] has hydrogen bonding which is the strongest IMF, requiring more energy to overcome, hence a higher boiling point."
• Know the number of H-bonds a molecule can form (e.g., water: 4, ethanol: 3, HF: 2).

Frequently Asked Questions

What is the difference between a hydrogen bond and a covalent bond?

A covalent bond is a strong intramolecular force formed by sharing electrons between atoms (bond energy ~200-800 kJ/mol). A hydrogen bond is a much weaker intermolecular force (~5-40 kJ/mol) caused by electrostatic attraction between a $\delta^+$ H and a lone pair on N, O, or F.

Why does ice float on water?

When water freezes, hydrogen bonds lock molecules into an open hexagonal lattice with more space between molecules than in liquid water. This makes ice less dense than liquid water, so it floats.

Can hydrogen bonds form within the same molecule?

Yes, this is called intramolecular hydrogen bonding and occurs in some molecules like salicylic acid and certain proteins. However, most hydrogen bonding discussed in chemistry courses is intermolecular.

Why can't $H_2S$ form hydrogen bonds?

Sulfur has an electronegativity of 2.58, which is not high enough to create the strong $\delta^+$ on hydrogen needed for a hydrogen bond. Only N (3.04), O (3.44), and F (3.98) are electronegative enough.

Related Topics

• Intermolecular Forces — Compare hydrogen bonding with London forces and dipole-dipole interactions.
• Chemical Bonds — Understand the intramolecular forces (ionic, covalent, metallic) that are much stronger than hydrogen bonds.
• VSEPR Theory — Molecular shape determines polarity, which determines whether hydrogen bonding can occur.