Intermolecular Forces

SPEC 3.1.3.7 - BONDING

London forces, permanent dipole-dipole interactions, and hydrogen bonding

What are Intermolecular Forces?

Definition

Intermolecular forces are forces of attraction that exist between molecules. They are NOT the same as covalent bonds, which exist within molecules.

Key Distinction

Intramolecular forces = bonds WITHIN molecules (covalent bonds) - STRONG

Intermolecular forces = attractions BETWEEN molecules - WEAK

There are three types of intermolecular forces:

TypeAlso Known AsRelative Strength

London dispersion forcesInduced dipole-dipole, Van der WaalsWeakest

Permanent dipole-dipoleDipole-dipole interactionsIntermediate

Hydrogen bondingH-bondingStrongest

Why This Matters

Intermolecular forces determine physical properties: melting points, boiling points, solubility, and viscosity.

Strength Comparison

Intermolecular forces are much weaker than covalent bonds. For example, a hydrogen bond is only about one tenth the strength of a covalent bond.

London Dispersion Forces

Definition

London dispersion forces (also called induced dipole-dipole forces or Van der Waals forces) are weak attractions caused by temporary, instantaneous dipoles that induce dipoles in neighbouring molecules.

How they arise:

Key Point

London forces exist in ALL molecules - polar AND non-polar. They are the ONLY intermolecular force in non-polar molecules like O₂, CO₂, and alkanes.

Important

Because the electron clouds are constantly moving, the dipoles are only temporary. They are constantly forming and disappearing as the electrons move. However, the overall effect is a net attraction between the atoms or molecules.

Factors affecting strength:

Number of electrons - more electrons = larger electron cloud = more polarisable (easier to distort) = stronger London forces

Surface area - larger molecules = more contact = stronger forces

Key Term

Polarisability is the ease with which an electron cloud can be distorted to form a temporary dipole. Larger atoms/molecules with more electrons are more polarisable.

Worked Example01

Explain why the boiling point of Br₂ (59°C) is higher than that of Cl₂ (−34°C), even though both are non-polar molecules.

1

Both Cl₂ and Br₂ are non-polar molecules, so the only intermolecular forces present are London dispersion forces.

2

Br₂ has more electrons than Cl₂ (70 vs 34 electrons).

3

More electrons means the electron cloud is more polarisable, creating stronger instantaneous dipoles.

4

Stronger London forces require more energy to overcome, so Br₂ has a higher boiling point.

Answer: Br₂ has more electrons than Cl₂, leading to stronger London dispersion forces between molecules. More energy is needed to overcome these stronger forces, resulting in a higher boiling point.

Molecular Lattices and Noble Gases

Induced dipole-dipole forces can be strong enough to hold molecules together in a lattice structure.

Example: Solid Iodine (I₂)

Within each I₂ molecule, atoms are held together by strong covalent bonds

Between I₂ molecules, there are only weak induced dipole-dipole forces

These weak forces hold the molecules in a regular arrangement called a molecular lattice

Noble Gases

Induced dipole-dipole forces also explain why noble gases can exist as liquids and solids at very low temperatures:

Noble gas atoms have complete outer shells and do NOT form covalent, ionic, or metallic bonds

However, weak induced dipole-dipole forces allow them to condense into liquids and solids at very low temperatures

Larger noble gases (e.g., Xe) have more electrons, so they have stronger London forces and higher boiling points than smaller ones (e.g., He)

Noble GasElectronsBoiling Point (K)

Helium (He)24

Neon (Ne)1027

Argon (Ar)1887

Xenon (Xe)54165

Permanent Dipole-Dipole Forces

Definition

Permanent dipole-dipole forces are attractions between the permanent dipoles of polar molecules. The δ+ end of one molecule attracts the δ− end of another.

Requirements: The molecule must be polar (have a permanent dipole).

How Polar Bonds Arise

When two atoms in a covalent bond have different electronegativities, the bond is polar:

The more electronegative atom gets a partial negative charge (δ−)

The less electronegative atom gets a partial positive charge (δ+)

The greater the difference in electronegativity, the more polar the bond becomes

Important

Polar molecules have BOTH London forces AND permanent dipole-dipole forces. The permanent dipole-dipole forces are in addition to the London forces.

Examples of polar molecules: HCl, HBr, CHCl₃, propanone, ethanal

Exam Tip

For similar-sized molecules, permanent dipole-dipole forces are stronger than London forces alone. This is why polar molecules often have higher boiling points than non-polar molecules of similar size.

Worked Example02

Pentane (C₅H₁₂) has a boiling point of 36°C, while 2,2-dimethylpropane (also C₅H₁₂) has a boiling point of 10°C. Both have the same molecular formula. Explain why their boiling points are different.

1

Both molecules have the same molecular formula (C₅H₁₂), so they have the same number of electrons.

2

Both are non-polar, so the only intermolecular forces present are London dispersion forces.

3

Pentane is a straight-chain molecule with a larger surface area. More surface contact between molecules = stronger London forces.

4

2,2-dimethylpropane is a compact, spherical molecule with a smaller surface area. Less surface contact = weaker London forces.

Answer: Although both molecules have the same number of electrons, pentane has a larger surface area than the more compact 2,2-dimethylpropane. Greater surface area allows more contact between molecules, resulting in stronger London forces and a higher boiling point.

Polar Molecules Have Both Forces

Important

Polar molecules experience permanent dipole-dipole forces IN ADDITION TO induced dipole-dipole forces. This means polar molecules have stronger overall intermolecular forces than non-polar molecules of similar size.

For example, methanal (CH₂O) has a higher boiling point than ethane (C₂H₆):

MoleculeFormulaM_rPolarityBoiling Point (K)

MethanalCH₂O30Polar254

EthaneC₂H₆30Non-polar184

The difference in boiling points can be explained by the types of intermolecular forces present:

Methanal is polar due to the carbonyl group (C=O), so it experiences both permanent dipole-dipole forces AND induced dipole-dipole forces . The permanent dipole-dipole forces are stronger than the induced dipole-dipole forces alone.

Ethane is non-polar, so it only experiences induced dipole-dipole forces .

Key Conclusion

More energy is required to overcome the stronger intermolecular forces in methanal, leading to a higher boiling point, even though the molecules are of similar size (same M_r = 30).

Hydrogen Bonding

Definition

Hydrogen bonding is a special, strong type of intermolecular force. It occurs when a hydrogen atom bonded to F, O, or N is attracted to a lone pair on F, O, or N of another molecule.

Conditions for Hydrogen Bonding

BOTH conditions must be met:

A hydrogen atom directly bonded to F, O, or N

A lone pair available on F, O, or N of another molecule

Why is it so strong?

F, O, and N are highly electronegative - they pull electrons strongly away from H

H has no inner electron shells , giving it a very high charge density - when bonded to F/O/N, the δ+ proton is very exposed

The exposed δ+ H is strongly attracted to the lone pair on a nearby F/O/N

Common Mistake

NOT all molecules containing hydrogen can form hydrogen bonds! The hydrogen must be bonded directly to F, O, or N. For example, CH₄ contains hydrogen but CANNOT form hydrogen bonds.

Why Water Has Strong H-Bonding

Each water molecule can form up to 4 hydrogen bonds: 2 through its H atoms (as H-bond donors) and 2 through the lone pairs on oxygen (as H-bond acceptors). This is why H₂O has stronger hydrogen bonding than HF or NH₃.

Drawing Hydrogen Bonds

When drawing hydrogen bonds between molecules, you must show:

The hydrogen bond as a dotted/dashed line

The bond at approximately 180° (linear) to the O-H or N-H bond

The lone pair on the acceptor atom

Exam Tip

Make sure to use a dashed, straight line when drawing intermolecular forces! Hydrogen bonds should start at the lone pair and go right up to the delta positive atom - it must be really clear where your H-bond starts and ends.

Worked Example03

Identify which of the following molecules can form hydrogen bonds with each other: CH₄, NH₃, HF, H₂S

1

For hydrogen bonding, we need H bonded directly to F, O, or N.

2

CH₄: H is bonded to C, not F/O/N. Cannot form H-bonds.

3

NH₃: H is bonded directly to N. N has a lone pair. Can form H-bonds.

4

HF: H is bonded directly to F. F has lone pairs. Can form H-bonds.

5

H₂S: H is bonded to S, not F/O/N. Cannot form H-bonds.

Answer: NH₃ and HF can form hydrogen bonds. CH₄ and H₂S cannot.

Comparing IMF Strengths

Strength Order

London forces < Permanent dipole-dipole < Hydrogen bonding

(For molecules of similar size)

IMF TypePresent InTypical Strength

London dispersionALL molecules0.05 - 40 kJ mol⁻¹

Permanent dipole-dipolePolar molecules only5 - 25 kJ mol⁻¹

Hydrogen bondingMolecules with H-F, H-O, or H-N10 - 40 kJ mol⁻¹

Important Exception

Size can dominate! A large non-polar molecule can have stronger TOTAL intermolecular forces than a small polar molecule, because it has much stronger London forces.

Example: I₂ (non-polar, large) has a higher boiling point (184°C) than HCl (polar, small, −85°C).

Comparison Tip

For small molecules with the same number of electrons, permanent dipole-dipole forces are stronger than induced dipole-dipole forces. This is why polar molecules like propanone have higher boiling points than non-polar molecules like butane (both have similar numbers of electrons).

Effect on Boiling Points

Stronger intermolecular forces → higher boiling point and higher enthalpy of vaporisation

More energy is needed to overcome the attractions between molecules.

Anomalous Enthalpy of Vaporisation

The enthalpy changes increase from H₂S to H₂Te due to the increased number of electrons in the Group 6 elements. This causes an increase in instantaneous dipole-induced dipole forces as the molecules become larger.

Based on this trend, H₂O should have a much lower enthalpy change of vaporisation (around 17 kJ mol⁻¹). However, the actual value is almost 3 times larger (~41 kJ mol⁻¹) because H₂O can form hydrogen bonds, while H₂S, H₂Se, and H₂Te cannot.

Worked Example04

Explain why the boiling points of the hydrogen halides increase in the order: HCl (−85°C) < HBr (−67°C) < HI (−35°C), but HF (20°C) does not follow this trend.

1

HCl, HBr, and HI are all polar molecules with permanent dipole-dipole forces AND London forces.

2

From HCl → HBr → HI, the number of electrons increases, so London forces get stronger.

3

Although polarity decreases (weaker permanent dipole-dipole), the increase in London forces is the dominant factor, causing boiling points to increase.

4

HF is different: The H is bonded directly to F (highly electronegative), so HF can form hydrogen bonds.

5

Hydrogen bonds are much stronger than the dipole-dipole forces in HCl/HBr/HI, giving HF an anomalously high boiling point.

Answer: HCl → HBr → HI: increasing London forces (more electrons) dominate over decreasing polarity. HF has a much higher bp because it can form hydrogen bonds, which are stronger than the dipole-dipole forces in other hydrogen halides.

Why Does Ice Float?

Most solids are denser than their liquids. Water is unusual: ice is LESS dense than liquid water.

Explanation

In ice, hydrogen bonds hold water molecules in an open hexagonal lattice structure. The relatively long bond lengths of the hydrogen bonds mean that the water molecules are slightly further apart than in the liquid form.

This creates empty space within the lattice, making ice about 9% less dense than liquid water.

Why This Matters

Because ice floats, lakes freeze from the top down. The ice layer insulates the water below, allowing aquatic life to survive in winter. This is essential for life on Earth!

Practice

Question 1: State the type(s) of intermolecular forces present in each of the following: (a) Ar (b) HBr (c) CH₃OH

(a) Ar: London dispersion forces only (non-polar, monatomic)

(b) HBr: London forces AND permanent dipole-dipole forces (polar molecule, but H not bonded to F/O/N so no H-bonding)

(c) CH₃OH: London forces, permanent dipole-dipole forces, AND hydrogen bonding (H is bonded directly to O, and O has lone pairs)

Practice

Question 2: Explain why ethanol (C₂H₅OH, Mr = 46) has a higher boiling point (78°C) than methoxymethane (CH₃OCH₃, Mr = 46, bp = −24°C), despite having the same molecular formula C₂H₆O.

Ethanol (C₂H₅OH): Contains an O-H group, so hydrogen bonding can occur between molecules. It also has London forces and permanent dipole-dipole forces.

Methoxymethane (CH₃OCH₃): Has no O-H, N-H, or F-H bonds. Although it contains oxygen and is polar, it can only form permanent dipole-dipole forces and London forces - NO hydrogen bonding.

Conclusion: The hydrogen bonding in ethanol requires more energy to overcome than the weaker dipole-dipole forces in methoxymethane, giving ethanol a higher boiling point.

Practice

Question 3: Draw a diagram showing hydrogen bonding between two molecules of ammonia (NH₃). Label the partial charges and lone pair involved.

Key features: H-bond shown as dashed line, approximately linear (180°), from δ+ H to lone pair on δ− N

Properties of Simple Molecular Substances

The type and strength of intermolecular forces present in a simple molecular substance influence its physical properties:

1. Low Melting and Boiling Points

Weak intermolecular forces require little energy to overcome, so covalent compounds often have low melting and boiling points and may be liquid or gaseous at room temperature. Stronger intermolecular forces lead to higher melting and boiling points.

2. Solubility in Water

Polar molecules, especially those capable of hydrogen bonding (e.g., ethanoic acid), can interact favourably with water molecules and are soluble.

Non-polar molecules that only have induced dipole-dipole forces (e.g., hexane) are hydrophobic and insoluble in water.

3. Electrical Conductivity

Covalent compounds do NOT conduct electricity, regardless of polarity. Even though polar molecules have permanent dipoles, they are electrically neutral overall and do not carry charge.

Common Exam Question

Explain why simple molecular substances have low melting points but giant covalent structures have high melting points.

Answer: In simple molecular substances, only weak intermolecular forces must be overcome (NOT the strong covalent bonds within molecules). In giant covalent structures, many strong covalent bonds must be broken.

Summary

Three Types of Intermolecular Forces

TypePresent InCaused By

London dispersionALL moleculesTemporary/induced dipoles

Permanent dipole-dipolePolar moleculesAttraction between δ+ and δ−

Hydrogen bondingH-F, H-O, H-N bondsH attracted to lone pair on F/O/N

Key Exam Points

Strength order (similar-sized molecules): London < dipole-dipole < H-bonding

Larger molecules can have stronger London forces than smaller polar molecules

H₂O and HF have anomalously high boiling points due to hydrogen bonding

Ice floats because hydrogen bonds create an open hexagonal structure

Draw H-bonds as dashed lines at ~180° to the covalent bond