Intermolecular Forces

students, by the end of this lesson you should be able to explain what intermolecular forces are, compare the main types, and predict how they affect the properties of substances. You will also see how these forces connect to the wider ideas in Structure 2 — Models of Bonding and Structure, including why some substances are gases at room temperature, why others are liquids, and why some solids have unusually high melting points. 🌡️

Intermolecular forces are the attractions between particles. They are not the same as the chemical bonds inside particles. That distinction is essential in IB Chemistry SL. A molecule may have strong covalent bonds within it, but the attraction between separate molecules may be much weaker. Those weaker attractions still matter a lot because they help decide boiling point, melting point, volatility, viscosity, and solubility.

What intermolecular forces are

Intermolecular forces act between molecules, atoms in noble gases, or ions and molecules in some solutions. They are always weaker than ionic, covalent, and metallic bonds. However, “weaker” does not mean “unimportant.” A substance with relatively strong intermolecular forces usually needs more energy to separate its particles, so it often has a higher boiling point and lower volatility.

The key idea is that every substance has particles that attract each other to some degree. The type and strength of those attractions depend on the structure of the particles, especially whether the particles are polar or non-polar, and whether hydrogen is bonded to certain highly electronegative atoms.

In IB Chemistry SL, the three main intermolecular forces you should know are:

London dispersion forces
Permanent dipole–dipole forces
Hydrogen bonding

You should also know that in some systems, especially solutions, ion–dipole attractions can be important, though they are often treated separately because they involve ions and polar molecules rather than neutral molecules only.

London dispersion forces: present in all substances

London dispersion forces are attractions caused by temporary changes in electron distribution. Electrons are always moving, so at any moment a molecule or atom may have a temporary dipole. That temporary dipole can induce a dipole in a nearby particle, creating a weak attraction.

These forces are present in all atoms and molecules, including non-polar ones such as $\mathrm{N_2}$, $\mathrm{O_2}$, and $\mathrm{CH_4}$. This is why even noble gases such as argon can be liquefied at very low temperatures. Without dispersion forces, non-polar substances would not stick together at all.

The strength of dispersion forces increases when particles have more electrons and larger electron clouds. Bigger particles are more polarizable, meaning their electron clouds are easier to distort. For example, the boiling point of the halogens increases down Group 17 because $\mathrm{I_2}$ has stronger dispersion forces than $\mathrm{F_2}$.

Another important factor is surface area. Straight-chain molecules usually have stronger dispersion forces than compact, branched molecules with the same formula because they can touch over a larger area. For example, a straight-chain hydrocarbon may have a higher boiling point than a branched isomer with the same molecular formula.

Example

Consider $\mathrm{CH_4}$ and $\mathrm{C_4H_{10}}$. Both are non-polar, so the main intermolecular forces are London dispersion forces. But $\mathrm{C_4H_{10}}$ has more electrons and a larger surface area, so its dispersion forces are stronger. As a result, $\mathrm{C_4H_{10}}$ has a much higher boiling point than $\mathrm{CH_4}$.

Permanent dipole–dipole forces: between polar molecules

A polar molecule has an uneven distribution of charge because its bonds have different electronegativities and the shape of the molecule does not cancel the bond dipoles. This creates a permanent dipole, with a partial positive end and a partial negative end.

Permanent dipole–dipole forces are attractions between the positive end of one polar molecule and the negative end of another. These forces occur in addition to dispersion forces. They are stronger than dispersion forces for molecules of similar size when the polarity is significant.

A common mistake is to assume that any molecule with polar bonds must be polar overall. That is not always true. The shape matters. For example, $\mathrm{CO_2}$ has polar $\mathrm{C=O}$ bonds, but the molecule is linear and symmetrical, so the bond dipoles cancel. Therefore, $\mathrm{CO_2}$ is non-polar overall and does not have permanent dipole–dipole forces.

By contrast, $\mathrm{HCl}$ is polar because the bond dipole does not cancel. This gives it dipole–dipole attractions in addition to dispersion forces.

Example

Compare $\mathrm{HCl}$ and $\mathrm{Cl_2}$. Both have similar relative molecular mass, so dispersion forces are present in both and are of comparable size. But $\mathrm{HCl}$ is polar, so it also has dipole–dipole forces. That means the particles in $\mathrm{HCl}$ attract each other more strongly, giving it a higher boiling point than $\mathrm{Cl_2}$.

Hydrogen bonding: a particularly strong type of dipole–dipole force

Hydrogen bonding is a special type of intermolecular force. It occurs when hydrogen is covalently bonded to a very electronegative atom: nitrogen, oxygen, or fluorine. The hydrogen becomes strongly $\delta^+$, and it is attracted to a lone pair on $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$ in a nearby molecule.

To identify hydrogen bonding, check for both conditions:

A hydrogen atom bonded directly to $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$
A nearby lone pair on $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$

Hydrogen bonding is stronger than ordinary dipole–dipole forces, but it is still much weaker than covalent bonding. Its effects are easy to see in water, alcohols, and carboxylic acids.

Water is the classic example. Each $\mathrm{H_2O}$ molecule can form multiple hydrogen bonds because it has two hydrogen atoms and two lone pairs. This network of attractions helps explain why water has an unusually high boiling point for such a small molecule.

Example

Compare $\mathrm{H_2O}$ and $\mathrm{H_2S}$. Both are group 16 hydrides, but only $\mathrm{H_2O}$ forms hydrogen bonds because oxygen is sufficiently electronegative. As a result, $\mathrm{H_2O}$ has a much higher boiling point than $\mathrm{H_2S}$. This is a clear example of how intermolecular forces affect physical properties.

Connecting intermolecular forces to properties

Intermolecular forces help explain several observable properties. This is one of the main reasons they matter in chemistry.

Boiling point

Boiling happens when particles gain enough energy to escape from the liquid into the gas phase. Stronger intermolecular forces mean more energy is needed, so the boiling point is higher. This is why liquids with hydrogen bonding often boil at higher temperatures than similar molecules without hydrogen bonding.

Melting point

Melting involves weakening the ordered arrangement of particles in a solid so they can move more freely. Stronger attractions usually lead to higher melting points, although crystal packing and molecular symmetry also matter.

Volatility

Volatility is how easily a liquid evaporates. A volatile liquid has weak intermolecular forces and can enter the gas phase easily. Gasoline is more volatile than water because its non-polar molecules mainly experience dispersion forces, while water has strong hydrogen bonding.

Viscosity

Viscosity is resistance to flow. Stronger intermolecular forces usually increase viscosity because molecules cling to each other more strongly. This is why syrup flows more slowly than water. 🍯

Solubility

A useful IB idea is “like dissolves like.” Polar substances often dissolve in polar solvents because the intermolecular forces between solute and solvent can be similar in type and strength. Non-polar substances usually dissolve better in non-polar solvents. For example, ethanol mixes well with water because it can hydrogen bond, while hexane does not mix well with water because it is non-polar.

Applying the ideas in IB Chemistry SL

When answering exam questions, students, you often need to identify the strongest intermolecular force present and then link it to a property.

A useful step-by-step method is:

Identify whether the substance is ionic, molecular, metallic, or atomic.
If it is molecular or atomic, decide whether it is polar or non-polar.
Check whether hydrogen bonding is possible.
Compare the strength of the intermolecular forces.
Link the forces to boiling point, melting point, volatility, or solubility.

Worked example

Which has the higher boiling point: $\mathrm{CH_3OH}$ or $\mathrm{CH_3OCH_3}$?

Both molecules have similar size, so dispersion forces are similar. Both are polar, so both have dipole–dipole forces. However, $\mathrm{CH_3OH}$ has an $\mathrm{O-H}$ bond, so it can form hydrogen bonds. $\mathrm{CH_3OCH_3}$ cannot hydrogen bond to itself because it has no hydrogen directly bonded to oxygen. Therefore, $\mathrm{CH_3OH}$ has the higher boiling point.

This type of explanation is exactly the kind of reasoning expected in IB Chemistry SL: identify structure, identify forces, and connect them to a property using evidence.

Intermolecular forces in the wider topic of Structure 2

This lesson fits into Structure 2 — Models of Bonding and Structure because chemistry uses models to explain how structure controls properties. Ionic bonding, covalent bonding, metallic bonding, molecular shape, and intermolecular forces all work together to explain the behavior of matter.

A substance’s properties are not determined by one idea alone. For example:

Ionic solids have strong electrostatic attractions between ions, so they usually have high melting points.
Molecular substances may have much lower melting points because the forces between molecules are weaker.
Metallic substances have metallic bonding, which explains conductivity and malleability.

Intermolecular forces are especially important because they help explain why substances with similar covalent structures can behave very differently. Two molecules may both contain covalent bonds, but if one has hydrogen bonding and the other does not, their boiling points can be very different.

Conclusion

Intermolecular forces are the attractions between particles that help determine the physical properties of substances. The main types are London dispersion forces, permanent dipole–dipole forces, and hydrogen bonding. These forces are weaker than chemical bonds, but they are still powerful enough to affect boiling point, melting point, viscosity, solubility, and volatility. By identifying molecular structure, polarity, and the presence of hydrogen bonding, you can predict and explain many patterns in IB Chemistry SL. Understanding intermolecular forces also strengthens your understanding of the larger Structure 2 topic because it shows how models of structure connect to real-world properties.

Study Notes

Intermolecular forces are attractions between particles, not bonds within a particle.
All substances have London dispersion forces.
Dispersion forces increase with more electrons, larger molecules, and greater surface area.
Permanent dipole–dipole forces occur between polar molecules.
Hydrogen bonding happens when $\mathrm{H}$ is bonded to $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$ and attracted to a lone pair on $\mathrm{N}$, $\mathrm{O}$, or $\mathrm{F}$ in another molecule.
Stronger intermolecular forces usually mean higher boiling point, higher melting point, lower volatility, and greater viscosity.
“Like dissolves like” is a useful rule for solubility.
Shape matters because it can cancel or strengthen molecular polarity.
$\mathrm{CO_2}$ is non-polar overall even though its bonds are polar.
$\mathrm{H_2O}$ has unusually strong intermolecular forces because of hydrogen bonding.
In exam questions, identify the particle type, determine polarity, check for hydrogen bonding, then connect the force to the property.
Intermolecular forces are a key part of Structure 2 because they link structure to observable behavior.