Hydrogen Bonding

Introduction: Why does water act so strangely? 💧

students, have you ever noticed that water can form droplets on a leaf, climb up a paper towel, or take a long time to boil compared with many other liquids? These behaviors are not caused by ionic bonding or covalent bonding alone. They are strongly influenced by a special intermolecular force called hydrogen bonding.

In this lesson, you will learn:

what hydrogen bonding is and how it forms,
how to identify molecules that can hydrogen bond,
why hydrogen bonding affects properties such as boiling point, viscosity, and surface tension,
how hydrogen bonding fits into the IB Chemistry HL topic of Structure 2 — Models of Bonding and Structure.

Hydrogen bonding is a key idea in understanding how molecular structure affects real-world properties. It helps explain the behavior of water, DNA, proteins, alcohols, and many other substances 🌍.

What is hydrogen bonding?

Hydrogen bonding is a strong intermolecular force that occurs when a hydrogen atom is covalently bonded to a very electronegative atom, usually nitrogen, oxygen, or fluorine, and is attracted to a lone pair of electrons on a nearby molecule.

The important idea is that hydrogen bonding is not a bond inside the molecule like a covalent bond. Instead, it is an attraction between molecules. In IB Chemistry terms, it is stronger than ordinary dipole-dipole forces and London dispersion forces, but weaker than covalent or ionic bonding.

A hydrogen bond forms when two conditions are met:

A hydrogen atom is bonded to $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$.
A nearby $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$ atom has a lone pair that can attract the $\delta^+$ hydrogen.

For example, in water, the oxygen atom pulls electron density toward itself, so the hydrogen atoms become partially positive, written as $\delta^+$. Another water molecule’s oxygen, which has lone pairs, is partially negative, written as $\delta^-$. The attraction between $\delta^+$ hydrogen and a lone pair on oxygen is a hydrogen bond.

This can be shown as:

$$\mathrm{O-H \cdots O}$$

The dotted line shows the hydrogen bond, not a covalent bond.

Why are $\mathrm{N}$, $\mathrm{O}$, and $\mathrm{F}$ special?

Hydrogen bonding is strongest when hydrogen is attached to nitrogen, oxygen, or fluorine because these atoms are very electronegative. Electronegativity is the ability of an atom to attract the shared pair of electrons in a covalent bond.

Because $\mathrm{N}$, $\mathrm{O}$, and $\mathrm{F}$ are so electronegative:

the $\mathrm{H}$ atom becomes highly $\delta^+$,
the bond becomes very polar,
and the hydrogen can strongly attract a lone pair on another molecule.

Fluorine is the most electronegative element, but hydrogen bonding is not simply about electronegativity alone. The atom must also have small size and lone pairs available. That is why hydrogen bonding is especially important in molecules such as $\mathrm{H_2O}$, $\mathrm{HF}$, and $\mathrm{NH_3}$.

Examples of molecules that can hydrogen bond include:

water, $\mathrm{H_2O}$
hydrogen fluoride, $\mathrm{HF}$
ammonia, $\mathrm{NH_3}$
ethanol, $\mathrm{CH_3CH_2OH}$
carboxylic acids such as ethanoic acid, $\mathrm{CH_3COOH}$

Examples of molecules that cannot hydrogen bond with themselves include:

methane, $\mathrm{CH_4}$
hydrogen chloride, $\mathrm{HCl}$
ethane, $\mathrm{C_2H_6}$

Even though $\mathrm{HCl}$ is polar, chlorine is not one of the atoms that produces strong hydrogen bonding in IB Chemistry.

How hydrogen bonding affects properties

Hydrogen bonding has major effects on the physical properties of substances. These effects are a direct example of the Structure 2 idea that structure determines properties.

1. Higher boiling points

A liquid boils when its particles gain enough energy to overcome intermolecular forces and escape into the gas phase. If a substance has hydrogen bonding, more energy is needed to separate its molecules.

That is why water has a much higher boiling point than expected for such a small molecule. Compare:

$\mathrm{H_2O}$ boils at $100^\circ\mathrm{C}$
$\mathrm{H_2S}$ boils at about $-60^\circ\mathrm{C}$

Both are group 16 hydrides, but water forms hydrogen bonds while hydrogen sulfide does not. This is a classic IB example of how hydrogen bonding changes physical properties.

2. Higher viscosity and surface tension

Viscosity is a liquid’s resistance to flow. Surface tension is the tendency of a liquid surface to resist being broken.

In water, hydrogen bonds make molecules stick together more strongly than in many other liquids. This gives water relatively high surface tension, which helps insects walk on water and allows droplets to form. It also increases viscosity compared with similar small molecules that do not hydrogen bond.

3. Unusual density of ice

Ice is less dense than liquid water. This is unusual because most substances become denser when they solidify.

When water freezes, hydrogen bonds arrange molecules into an open hexagonal structure. This structure spaces molecules farther apart than in liquid water, so solid ice floats. This property is crucial for life in lakes and oceans because ice forms a protective layer at the surface ❄️.

4. Solubility in water

Hydrogen bonding helps explain why many small molecules dissolve well in water if they can form hydrogen bonds with water molecules. These solutes are often called hydrophilic.

For example, ethanol dissolves in water because its $\mathrm{-OH}$ group can hydrogen bond with water. In contrast, hydrocarbons such as hexane do not hydrogen bond and are poorly soluble in water.

Recognizing hydrogen bonding in IB questions

A common IB Chemistry skill is deciding whether hydrogen bonding is present in a substance or between two substances. Use this step-by-step method, students:

Look for a hydrogen atom directly bonded to $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$.
Check whether a nearby molecule has a lone pair on $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$.
Decide whether the molecules are arranged so the attraction can occur.

For example:

Water molecules hydrogen bond with each other.
Ethanol molecules hydrogen bond with each other and with water.
Ammonia molecules can hydrogen bond, but the interactions are weaker than in water because nitrogen is less electronegative than oxygen and each ammonia molecule has fewer lone pairs and a different shape.

A useful comparison is between methanol, $\mathrm{CH_3OH}$, and dimethyl ether, $\mathrm{CH_3OCH_3}$. Both have the same molecular formula $\mathrm{C_2H_6O}$, but methanol has an $\mathrm{O-H}$ bond and can hydrogen bond with itself, while dimethyl ether cannot hydrogen bond with itself because it has no $\mathrm{O-H}$ hydrogen. This helps explain why methanol has a higher boiling point.

Hydrogen bonding in larger structures and materials

Hydrogen bonding is not only important in small molecules. It also matters in large biological and chemical systems.

DNA

In DNA, the two strands are held together by hydrogen bonds between complementary bases. These bonds are weak individually, but many together provide enough stability while still allowing the strands to separate during replication.

Proteins

Hydrogen bonding helps stabilize the shapes of proteins. It contributes to the formation of structures such as alpha helices and beta sheets. The specific folding of a protein is essential for its function.

Materials and everyday products

Hydrogen bonding is important in polymers, adhesives, and cellulose. For example, cellulose in plant cell walls contains many $\mathrm{-OH}$ groups that form hydrogen bonds, giving plant fibers strength. This is one reason wood and cotton have their characteristic structure and mechanical properties.

Comparing hydrogen bonding with other intermolecular forces

To understand hydrogen bonding well, students, it helps to compare it with other forces in Structure 2.

London dispersion forces exist between all molecules and are caused by temporary dipoles.
Dipole-dipole forces occur between polar molecules.
Hydrogen bonding is a stronger special case of dipole-dipole attraction when hydrogen is bonded to $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$.

In many substances, more than one type of intermolecular force exists at the same time. The overall physical behavior depends on the combined effect of these forces.

For example, water has:

London dispersion forces,
dipole-dipole interactions,
hydrogen bonding.

However, hydrogen bonding is the dominant reason for many of water’s unusual properties.

Conclusion

Hydrogen bonding is a powerful intermolecular force that occurs when hydrogen bonded to $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$ is attracted to a lone pair on a nearby molecule. It explains many important properties such as high boiling point, surface tension, viscosity, solubility, and the low density of ice.

For IB Chemistry HL, hydrogen bonding is essential because it shows how molecular structure affects macroscopic behavior. It also connects to biological molecules, materials, and real-world chemistry. If you can identify when hydrogen bonding occurs and explain its effects, you have mastered a major part of Structure 2 — Models of Bonding and Structure ✅.

Study Notes

Hydrogen bonding is a strong intermolecular force, not a covalent bond.
It occurs when $\mathrm{H}$ is covalently bonded to $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$.
A hydrogen bond forms between the $\delta^+$ hydrogen of one molecule and a lone pair on $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$ in another molecule.
Hydrogen bonding raises boiling point, viscosity, and surface tension.
Water’s high boiling point and the floating of ice are classic examples.
Molecules such as $\mathrm{H_2O}$, $\mathrm{NH_3}$, $\mathrm{HF}$, and alcohols can hydrogen bond.
Molecules like $\mathrm{CH_4}$ and $\mathrm{C_2H_6}$ cannot hydrogen bond with themselves.
Hydrogen bonding helps explain solubility, biological structure, and material strength.
In IB questions, always check for both the correct hydrogen atom and an available lone pair.
Hydrogen bonding is a central example of the link between structure and properties in chemistry.