Intermolecular Forces (VCE SSCE Chemistry): Revision Notes

Intermolecular Forces

Understanding intermolecular forces

Intermolecular forces are the attractive forces that exist between molecules. These forces are much weaker than the covalent bonds that hold atoms together within molecules (intramolecular bonds). Understanding the difference between these two types of forces is essential for explaining the physical properties of molecular substances.

When you heat water, the molecules gain kinetic energy. Some molecules gain sufficient energy to overcome the forces holding them to neighbouring molecules and escape from the liquid surface. The resulting water vapour still consists of $\text{H}_2\text{O}$ molecules - the covalent bonds within each molecule remain intact. What has been disrupted are the forces between the water molecules. This demonstrates that covalent bonds within molecules are considerably stronger than intermolecular forces between molecules.

The strength of intermolecular forces depends on several factors, including molecular size, molecular shape, and molecular polarity. These factors determine not only how strong the forces are, but also which type of intermolecular force exists.

Types of intermolecular forces

There are three main types of intermolecular forces:

• Dispersion forces (also called London forces)
• Dipole-dipole attraction
• Hydrogen bonding

Each type has different characteristics and strengths, which affect the physical properties of substances such as melting points and boiling points.

Dispersion forces

Dispersion forces are present between all molecules, regardless of whether they are polar or non-polar. These forces arise from temporary fluctuations in electron distribution around molecules.

How dispersion forces work

Electrons are constantly moving within atoms and molecules. Even in non-polar molecules like fluorine ($\text{F}_2$), the electrons occasionally cluster more densely at one end of the molecule. This creates a temporary imbalance in charge distribution - one end becomes slightly negative ($\delta^-$) while the other becomes slightly positive ($\delta^+$). These temporary charge separations are called temporary dipoles or instantaneous dipoles.

When a temporary dipole forms in one molecule, it can induce similar dipoles in neighbouring molecules. These induced dipoles then attract each other, creating dispersion forces throughout the substance.

Because these dipoles are only temporary and constantly changing, dispersion forces are the weakest of the three types of intermolecular forces.

Factors affecting dispersion force strength

Molecular mass and electron count

The strength of dispersion forces increases as the molecular mass increases. Larger molecules contain more electrons, making it easier for temporary dipoles to form. Substances with stronger dispersion forces have higher melting and boiling points because more energy is needed to overcome these forces.

The table below shows this trend for the halogen group. All halogens form non-polar diatomic molecules with only dispersion forces between them. As molecular mass and electron count increase down the group, boiling points rise significantly.

Molecular shape

The shape of a molecule also affects dispersion force strength. Molecules with elongated, chain-like structures can form stronger dispersion forces than more compact, spherical molecules of similar molecular mass.

Consider butane ($\text{C}_4\text{H}_{10}$) and $2$-methylpropane ($\text{C}_4\text{H}_{10}$). Both have identical molecular formulas but different shapes. Butane has a linear structure while $2$-methylpropane is branched. The linear butane molecule has a larger surface area for contact with neighbouring molecules, allowing stronger dispersion forces to form. This explains why butane has a boiling point of $-0.5°\text{C}$ compared to $-11°\text{C}$ for $2$-methylpropane.

Dipole-dipole attraction

In addition to dispersion forces, polar molecules experience dipole-dipole attraction. This occurs because polar molecules have permanent charge separations - one end of the molecule is permanently slightly positive while the other end is permanently slightly negative.

These permanent dipoles cause polar molecules to align so that the positive end of one molecule attracts the negative end of a neighbouring molecule. Because these dipoles are permanent (rather than temporary), dipole-dipole attractions are stronger than dispersion forces.

Polarity and strength

The more polar a molecule is, the stronger the dipole-dipole attraction between molecules. Polarity increases when:

• There is a large difference in electronegativity between bonded atoms
• The molecule has an asymmetric shape

Substances with stronger dipole-dipole attractions have higher melting and boiling points. More energy (higher temperature) is required to overcome these stronger attractive forces when the substance changes state.

Hydrogen bonding

Hydrogen bonding is a particularly strong type of dipole-dipole attraction. It occurs in specific circumstances and is approximately ten times stronger than regular dipole-dipole attraction (though still only about one-tenth the strength of covalent or ionic bonds).

Requirements for hydrogen bonding

For hydrogen bonding to occur, molecules must meet two specific requirements:

These three elements (nitrogen, oxygen, and fluorine) are small and highly electronegative. When bonded to hydrogen, they strongly attract the bonding electrons, creating a large dipole. The hydrogen atom's single electron is pulled away, leaving the hydrogen nucleus (a proton) exposed. This exposed proton is strongly attracted to lone pairs of electrons on neighbouring molecules, forming the hydrogen bond.

The small size of the hydrogen atom allows neighbouring molecules to approach very closely, making the attractive force relatively strong.

Why other atoms don't form hydrogen bonds

You might wonder why other electronegative atoms like chlorine don't form hydrogen bonds. Although chlorine has high electronegativity, chlorine atoms are larger than nitrogen, oxygen, and fluorine. The electron density is more spread out and less concentrated, resulting in weaker attraction to hydrogen atoms on neighbouring molecules - not strong enough to be classified as hydrogen bonding.

Effect on physical properties

Hydrogen bonding significantly increases melting and boiling points.

Molecular size and intermolecular forces

An important principle to remember is that dispersion forces exist between all molecules - both polar and non-polar. In substances with large molecular masses, dispersion forces can become very strong and may even exceed the strength of dipole-dipole attraction or hydrogen bonding.

The table below compares four hydrogen halides. All except hydrogen fluoride follow an expected trend: as molecular mass increases, boiling point increases due to stronger dispersion forces. Hydrogen chloride is more polar than hydrogen iodide, but hydrogen iodide has the higher boiling point because its larger molecular mass creates stronger dispersion forces that outweigh the polarity effect.