Intermolecular forces, including hydrogen bonding

Covalent bonds explain why atoms remain joined inside a molecule, but they do not explain why water is a liquid at room temperature, why iodine is a solid or why some liquids evaporate much more readily than others. To understand those observations, we need to look at the attractions acting between separate particles.

These attractions are called intermolecular forces. Each one is relatively weak compared with a covalent bond, but a sample contains an enormous number of particles. Together, the attractions can therefore have a large effect on how much energy is needed to separate the particles.

The key skill in this chapter is therefore not simply memorising the names of three forces. You need to look at molecular structure, decide which attractions are present, compare their strength and then connect that microscopic picture to observable physical properties.

Intermolecular forces and intramolecular bonds

The prefix inter- means between. Intermolecular forces act between separate molecules or particles.

The prefix intra- means within. Intramolecular bonds hold atoms together within a molecule or structure.

Consider liquid water. Each H₂O molecule contains strong covalent O–H bonds. When water boils, the individual H₂O molecules remain intact. What changes is the distance between the molecules.

The attractions between neighbouring H₂O molecules are overcome sufficiently for the molecules to move apart and form steam.

Why intermolecular forces are electrostatic

Every intermolecular force in this chapter ultimately involves attraction between regions of opposite charge.

The charge separation may arise in different ways:

• a temporary uneven electron distribution creates an instantaneous dipole;
• unequal sharing of electrons can produce a permanent molecular dipole; or
• an especially polar N–H, O–H or F–H bond can allow hydrogen bonding.

Where do van der Waals forces fit?

The phrase van der Waals forces can cause confusion because it is a collective term rather than the name of an additional force.

In this H2 Chemistry treatment, the two types of van der Waals force that we need to distinguish are:

• instantaneous dipole–induced dipole forces, abbreviated as id-id forces; and
• permanent dipole–permanent dipole forces, abbreviated as pd-pd forces.

Hydrogen bonding is normally discussed separately because it is an unusually strong intermolecular attraction with specific structural requirements.

Instantaneous dipole–induced dipole forces

Electrons are continually moving. Even in a perfectly non-polar atom or molecule, the electron distribution is not frozen into one unchanging arrangement.

At a particular instant, slightly more electron density may happen to be found on one side of a particle than the other.

The electron-rich side becomes temporarily δ−, while the electron-poor side becomes temporarily δ+. The particle therefore possesses a temporary separation of charge called an instantaneous dipole.

How one temporary dipole affects its neighbour

1. The electrons in one particle momentarily become distributed unevenly.
2. An instantaneous δ+ end and δ− end appear.
3. This temporary dipole distorts the electron distribution in a neighbouring particle.
4. An oppositely orientated dipole is induced in the neighbour.
5. The adjacent δ+ and δ− regions attract electrostatically.

The exact positions of the temporary dipoles change continuously, but new dipoles are constantly being generated. The overall result is a persistent attractive interaction between the particles.

Alternative names for id-id forces

Instantaneous dipole–induced dipole forces may also be described as:

• London dispersion forces;
• dispersion forces; or
• id-id forces.

These names describe the same basic interaction.

What determines the strength of id-id forces?

Not all id-id attractions are equally strong. The most important idea is polarisability.

Polarisability describes how easily the electron cloud of a particle can be distorted.

Electron-cloud size and polarisability

A larger electron cloud generally contains more electrons and extends further from the nuclei. The outer regions of the cloud can be distorted more easily.

Larger instantaneous dipoles and induced dipoles can therefore form, creating stronger electrostatic attractions between neighbouring particles.

The number of electrons is therefore often a useful comparison tool within a related series.

However, an exam explanation should go beyond simply saying that one molecule has a larger Mᵣ. Relative molecular mass may correlate with the number of electrons, but mass itself is not the origin of the attraction.

Example: the halogens

F₂, Cl₂, Br₂ and I₂ are all non-polar. Their molecules therefore do not experience pd-pd attractions or hydrogen bonding.

Down the group, the electron clouds become progressively larger and more polarisable. The id-id forces become stronger, so progressively more energy is needed to separate the molecules.

Exam-quality reasoning:
I₂ has more electrons and a larger, more polarisable electron cloud than Cl₂. Larger instantaneous and induced dipoles can therefore form, producing stronger id-id forces between I₂ molecules. More energy is required to overcome these attractions, so I₂ has the higher boiling point.

Molecular shape also matters

Two molecules can contain the same number of electrons yet still have different boiling points.

This occurs because the overall strength of dispersion attractions also depends on how effectively neighbouring electron clouds can come into contact.

An extended molecule usually has a larger effective area of contact with nearby molecules. A compact, highly branched molecule generally makes less surface contact.

Example: isomers of pentane

All three molecules have formula C₅H₁₂ and therefore contain the same total number of electrons.

Pentane is less compact and can make greater surface contact with neighbouring molecules. Its overall id-id attractions are therefore stronger than those of the more highly branched 2,2-dimethylpropane.

Permanent dipole–permanent dipole forces

A polar molecule has a permanent uneven distribution of electron density. One region of the molecule is permanently δ+, while another is permanently δ−.

When neighbouring polar molecules adopt suitable orientations, the δ+ region of one molecule is attracted to the δ− region of another.

This attraction is called a permanent dipole–permanent dipole force, commonly shortened to pd-pd force.

The molecule must actually be polar

A common mistake is to identify a polar bond and immediately conclude that the whole molecule is polar.

Molecular polarity depends on both:

• the polarity of the individual bonds; and
• the three-dimensional arrangement of those bond dipoles.

Bond dipoles are vectors. If they are arranged symmetrically, they can cancel.

CO₂ and BF₃ still experience id-id forces because every particle does. They simply lack the additional pd-pd attraction.

A useful controlled comparison

Pd-pd attractions are easiest to identify when comparing particles with similar numbers of electrons.

For example, HCl molecules and Ar atoms each contain 18 electrons. Their id-id contributions are therefore of broadly comparable scale, but HCl possesses a permanent dipole while Ar does not.

HCl therefore has the additional pd-pd attraction between its molecules.

Hydrogen bonding

Hydrogen bonding is a particularly strong type of intermolecular attraction.

It occurs when hydrogen is covalently bonded to a very electronegative and small atom, causing the hydrogen atom to become strongly δ+.

At this level, the important atoms are:

• nitrogen, N;
• oxygen, O; and
• fluorine, F.

The two requirements for hydrogen bonding

For a hydrogen bond to form between two particles, there must be:

1. a hydrogen atom covalently bonded directly to N, O or F; and
2. a lone pair of electrons on an N, O or F atom in a neighbouring particle.

The group containing the N–H, O–H or F–H bond acts as the hydrogen-bond donor.

The lone-pair-bearing atom acts as the hydrogen-bond acceptor.

Here:

• X–H is a covalent bond;
• Y provides a lone pair; and
• H···Y represents the hydrogen bond.

Why N, O and F are special

N, O and F are highly electronegative. When H is bonded to one of these atoms, the bonding electrons are drawn strongly away from H.

The hydrogen atom therefore develops a substantial δ+ charge.

N, O and F are also relatively small atoms. Their lone-pair electron density is concentrated in a small region, allowing a strong electrostatic attraction to the δ+ hydrogen of a neighbouring molecule.

Which familiar molecules form hydrogen bonds?

Water can form an extensive hydrogen-bonded network

Each H₂O molecule contains:

• two O–H bonds that can donate hydrogen bonds; and
• two lone pairs on oxygen that can accept hydrogen bonds.

In a suitable arrangement, one H₂O molecule can therefore participate in up to four hydrogen bonds.

Hydrogen bonding in ammonia

NH₃ contains three N–H bonds and one lone pair on nitrogen, so ammonia molecules can also hydrogen-bond.

However, the hydrogen-bonded network in ammonia is less extensive than that in water. Oxygen is also more electronegative than nitrogen, making the O–H bond more polar than the N–H bond.

Why dimethyl ether cannot hydrogen-bond with itself

Dimethyl ether, CH₃OCH₃, contains an oxygen atom with lone pairs.

It can therefore accept hydrogen bonds from a suitable molecule such as water.

However, CH₃OCH₃ contains no O–H bond. A pure sample therefore contains no hydrogen-bond donor, so dimethyl ether molecules do not hydrogen-bond with one another.

Comparing the strengths of intermolecular forces

For molecules of broadly similar size and shape, the following order is a useful first approximation:

This does not mean that every molecule with pd-pd forces has a higher boiling point than every molecule with only id-id forces.

The total intermolecular attraction also depends on factors such as:

• the size and polarisability of the electron cloud;
• molecular shape and surface contact;
• the magnitude of the permanent dipole; and
• the number and strength of hydrogen bonds that can form.

Why the simple hierarchy can mislead

I₂ is non-polar and therefore has only id-id attractions between its molecules. HCl is polar and has both id-id and pd-pd attractions.

Yet I₂ has a much higher boiling point than HCl.

The reason is that I₂ has a much larger and more polarisable electron cloud. Its id-id attractions are therefore sufficiently strong to outweigh the additional pd-pd contribution found in the much smaller HCl molecules.

A reliable method for comparison questions

1. Identify the particles. Decide whether you are dealing with atoms, non-polar molecules or polar molecules.
2. Include id-id forces. They are always present.
3. Determine whether pd-pd forces are present. The molecule must have an overall permanent dipole.
4. Check for hydrogen bonding. Look for a suitable donor and acceptor.
5. Compare polarisability. More electrons and a larger electron cloud generally mean stronger id-id attractions.
6. Compare shape if necessary. Greater effective surface contact can strengthen the overall dispersion attraction.
7. Connect the difference to the property. State whether more or less energy is required to separate the particles.

How intermolecular forces affect physical properties

Once the intermolecular attractions have been identified and compared, we can use them to explain macroscopic properties.

Boiling point

When a liquid boils, its molecules must move from the liquid phase into the gas phase, where they are much further apart.

Stronger intermolecular attractions require more energy to overcome.

Melting point

Stronger intermolecular attractions can also increase melting point because more energy is required to disrupt the ordered arrangement of molecules in the solid.

However, melting-point trends are often less straightforward than boiling-point trends because the way molecules pack in a solid lattice also matters.

Molecular shape and symmetry can therefore make melting-point comparisons less regular.

Volatility

A volatile liquid evaporates readily.

If the intermolecular attractions are weak, molecules at the surface can escape into the gas phase relatively easily.

Vapour pressure

At the same temperature, a liquid with weaker intermolecular attractions generally has more molecules able to escape into the vapour phase.

Solubility

Dissolving a substance involves more than simply “pulling molecules apart”.

Existing attractions between solute particles and between solvent particles must be disrupted, while new attractions form between solute and solvent.

Dissolution is more favourable when the newly formed solute–solvent attractions are sufficiently favourable compared with the attractions being disrupted.

• Polar molecules can interact favourably with polar solvents through pd-pd attractions.
• Molecules with suitable donor or acceptor groups may form hydrogen bonds with water.
• Non-polar substances interact with non-polar solvents mainly through id-id forces.

The phrase “like dissolves like” is a useful shortcut, but in an explanation you should identify the intermolecular attractions responsible.

Example: alcohols and water

The –OH group in an alcohol can form hydrogen bonds with water molecules.

This favours solubility.

However, the hydrocarbon part of an alcohol is non-polar. As the hydrocarbon chain becomes longer, this non-polar region becomes increasingly important and the alcohol becomes progressively less soluble in water.

Worked intermolecular-force comparisons

Example 1: H₂O and H₂S

H₂S has a greater Mᵣ than H₂O, yet water has the much higher boiling point.

Both molecules experience id-id forces. Both are also polar, so they have pd-pd attractions.

The key difference is that H₂O molecules form hydrogen bonds because H is bonded directly to O and the O atom has lone pairs.

H₂S does not form conventional hydrogen bonds because H is bonded to S rather than N, O or F.

Answer:
H₂O molecules form strong intermolecular hydrogen bonds. More energy is required to overcome the attractions between H₂O molecules than those between H₂S molecules, so H₂O has the higher boiling point.

Example 2: ethanol and dimethyl ether

Ethanol, CH₃CH₂OH, and dimethyl ether, CH₃OCH₃, have the same molecular formula, C₂H₆O.

They therefore have the same number of electrons and the same Mᵣ.

Both substances experience id-id and pd-pd attractions.

However, ethanol contains an O–H group, so ethanol molecules form hydrogen bonds with one another.

Dimethyl ether has oxygen lone pairs but no O–H group, so its own molecules cannot donate hydrogen bonds to one another.

Answer:
Ethanol has additional strong hydrogen bonding between its molecules. More energy is therefore needed to separate ethanol molecules, giving ethanol the higher boiling point.

Example 3: Cl₂ and I₂

Both Cl₂ and I₂ are non-polar molecules. The intermolecular attractions between their molecules are therefore id-id forces.

I₂ contains more electrons and has a larger, more polarisable electron cloud than Cl₂.

Answer:
Larger instantaneous and induced dipoles form in I₂, so the id-id attractions between I₂ molecules are stronger. More energy is required to separate the I₂ molecules, giving I₂ the higher boiling point.

Example 4: pentane and 2,2-dimethylpropane

Both molecules have formula C₅H₁₂ and therefore the same number of electrons.

The important difference is molecular shape.

Pentane is more extended, allowing greater effective surface contact between neighbouring molecules. 2,2-Dimethylpropane is more compact.

Answer:
Pentane experiences stronger overall id-id attractions because its more extended shape permits greater surface contact between neighbouring electron clouds. It therefore has the higher boiling point.

Example 5: CO₂ and SO₂

Both molecules contain polar bonds.

CO₂ is linear and symmetrical, so the two C=O bond dipoles cancel. CO₂ is therefore non-polar and has no pd-pd attraction.

SO₂ is bent, so its bond dipoles do not cancel. SO₂ is polar and has pd-pd attractions in addition to id-id forces.

SO₂ also contains more electrons and has a more polarisable electron cloud.

Its overall intermolecular attractions are therefore substantially stronger.

Why ice is less dense than liquid water

Water shows an unusual property: solid ice is less dense than liquid water.

This behaviour is a direct consequence of hydrogen bonding.

The open hydrogen-bonded structure of ice

In ice, each H₂O molecule occupies a fixed position in an ordered hydrogen-bonded lattice.

Each water molecule is arranged approximately tetrahedrally relative to four neighbouring water molecules.

It can donate two hydrogen bonds through its two O–H bonds and accept two through the two lone pairs on oxygen.

This hydrogen-bonding arrangement creates an open structure containing significant empty space.

What happens when ice melts?

When ice melts, some of the ordered hydrogen bonds are disrupted and the rigid open structure partly collapses.

Water molecules can then move into some of the spaces that were present in the ice structure.

The molecules therefore pack more closely on average in liquid water than in ice.

Common mistakes with intermolecular forces

Intermolecular-forces checklist