Chapter: Intermolecular Forces

Welcome to one of the most important and exciting topics in AS Chemistry! Have you ever wondered why water is a liquid at room temperature while oxygen is a gas, even though oxygen molecules are heavier? Or why ice floats in your drink instead of sinking? The answer lies in intermolecular forces.

Don't worry if this topic sounds a bit abstract at first. We will break it down step-by-step using simple everyday analogies, clear explanations, and helpful memory tricks.


1. Intramolecular Bonds vs. Intermolecular Forces

Before diving in, let's make sure we do not confuse two very similar-sounding terms:

Intramolecular bonds are the strong chemical bonds inside a molecule that hold the atoms together (such as covalent bonds, ionic bonds, and metallic bonds).

Intermolecular forces (often abbreviated as IMFs) are the weaker attractive forces between separate molecules.

Everyday Analogy: Think of a classroom full of students holding their own hands tightly together. The grip of your left hand holding your right hand is an intramolecular bond (very strong, keeps you as one unit). If you lightly tap shoulders with the person sitting next to you, that weak attraction is an intermolecular force.

Key Takeaway: When you boil water, you are NOT breaking the strong covalent \( \text{O-H} \) bonds inside the water molecule. You are only overcoming the weaker intermolecular forces between different water molecules!


2. The Three Types of Intermolecular Forces

In CCEA AS Chemistry, you need to know three main types of intermolecular forces, ranked here from generally weakest to strongest:

1. van der Waals' forces (also called London dispersion forces or instantaneous dipole–induced dipole forces)
2. Permanent dipole–dipole forces
3. Hydrogen bonds (the strongest type of intermolecular force)


3. van der Waals' Forces (Instantaneous Dipole–Induced Dipole)

How do they form?

van der Waals' forces exist between all molecules and noble gas atoms, whether they are polar or non-polar. Here is how they happen step-by-step:

Step 1: Electrons are in constant, random motion around atomic nuclei.
Step 2: At any given instant, the electron cloud may become unevenly distributed, with more electrons on one side than the other.
Step 3: This creates a temporary, short-lived charge imbalance called a temporary dipole (or instantaneous dipole), with a slight negative charge \( \delta^- \) on one side and a slight positive charge \( \delta^+ \) on the other.
Step 4: This temporary dipole repels or attracts electrons in a neighbouring molecule, creating an induced dipole.
Step 5: The opposite charges on adjacent molecules attract each other weakly. These weak attractions are van der Waals' forces.

What factors affect the strength of van der Waals' forces?

1. Number of Electrons (Molecular Size):
The more electrons a molecule or atom has, the larger and more polarisable its electron cloud. This leads to larger temporary dipoles and stronger van der Waals' forces, requiring more thermal energy to overcome.

Example: Noble gases boiling points increase down Group 0:
\( \text{He} \) (\( 2 \) electrons) has a boiling point of \( -269^\circ\text{C} \)
\( \text{Xe} \) (\( 54 \) electrons) has a boiling point of \( -108^\circ\text{C} \)

2. Molecular Shape and Surface Area:
Straight-chain molecules have a larger surface area of contact with neighbouring molecules. They can pack closely together, leading to stronger van der Waals' forces.
Branched-chain isomers are more spherical and have smaller surface areas of contact. They cannot pack as closely together, resulting in weaker van der Waals' forces and lower boiling points.

Example: Pentane (\( \text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_2\text{CH}_3 \), straight chain) boils at \( 36^\circ\text{C} \), while its branched isomer 2,2-dimethylpropane (\( \text{C}(\text{CH}_3)_4 \)) boils at only \( 9.5^\circ\text{C} \).

Key Takeaway: More electrons = stronger temporary dipoles = stronger van der Waals' forces = higher boiling point.


4. Permanent Dipole–Dipole Forces

What are they?

Permanent dipole–dipole forces occur exclusively between polar molecules (molecules with a permanent dipole). A permanent dipole arises when two bonded atoms have a significant difference in electronegativity and the molecule's shape is asymmetrical.

In a sample of polar molecules, the slightly positive end (\( \delta^+ \)) of one polar molecule attracts the slightly negative end (\( \delta^- \)) of another polar molecule.

Example: In hydrogen chloride (\( \text{HCl} \)), chlorine is more electronegative than hydrogen. The bond is polar: \( \text{H}^{\delta+} - \text{Cl}^{\delta-} \). The \( \delta^+ \) hydrogen of one \( \text{HCl} \) molecule attracts the \( \delta^- \) chlorine of a neighbouring \( \text{HCl} \) molecule.

Comparing van der Waals' and Permanent Dipole–Dipole Forces

If two molecules have similar numbers of electrons (and therefore similar van der Waals' forces), the one with permanent dipole–dipole forces will have a higher boiling point because extra energy is needed to break these permanent electrostatic attractions.

Example Comparison:
Propane (\( \text{C}_3\text{H}_8 \)): \( 26 \) electrons, non-polar \( \rightarrow \) only van der Waals' forces \( \rightarrow \) boiling point = \( -42^\circ\text{C} \)
Methoxymethane (\( \text{CH}_3\text{OCH}_3 \)): \( 26 \) electrons, polar \( \rightarrow \) van der Waals' + permanent dipole–dipole \( \rightarrow \) boiling point = \( -24^\circ\text{C} \)

Key Takeaway: Permanent dipole–dipole attractions add extra holding strength on top of van der Waals' forces in polar molecules.


5. Hydrogen Bonding

What is a Hydrogen Bond?

A hydrogen bond is an especially strong type of permanent dipole–dipole attraction. It is the strongest of all intermolecular forces (though still significantly weaker than a covalent bond, typically around \( 10\% \) as strong).

The Two Golden Conditions for Hydrogen Bonding:

For hydrogen bonding to occur, you must have:

1. A hydrogen atom directly bonded to a highly electronegative atom with a small atomic radius: Nitrogen (\( \text{N} \)), Oxygen (\( \text{O} \)), or Fluorine (\( \text{F} \)).
2. A lone pair of electrons on that electronegative atom (\( \text{N} \), \( \text{O} \), or \( \text{F} \)) in a neighbouring molecule.

Memory Trick: Remember the phrase "Hydrogen bonding is NOF fair!" (N, O, F).

Why does this happen?

\( \text{N} \), \( \text{O} \), and \( \text{F} \) are so electronegative that they pull electron density strongly away from the small hydrogen atom. This leaves the hydrogen nucleus (a single proton) almost completely unshielded and carrying a high concentration of positive charge (\( \delta^+ \)). This exposed \( \delta^+ \) hydrogen is strongly attracted to the lone pair of electrons on an \( \text{N} \), \( \text{O} \), or \( \text{F} \) atom of an adjacent molecule.

How to Draw a Hydrogen Bond (CCEA Exam Skill):

When drawing hydrogen bonds between molecules (e.g. water \( \text{H}_2\text{O} \)), always include:

1. Partial charges (\( \delta^+ \) and \( \delta^- \)) on all relevant atoms.
2. The lone pair of electrons on the \( \text{N} \), \( \text{O} \), or \( \text{F} \) atom.
3. A dashed or dotted line representing the hydrogen bond going directly from the lone pair to the \( \delta^+ \) hydrogen atom.
4. A linear arrangement (\( 180^\circ \) bond angle) around the hydrogen bond: the \( \text{O-H}\cdots\text{O} \) grouping should form a straight line.

Key Takeaway: Hydrogen bonds only occur when \( \text{H} \) is bonded directly to \( \text{N} \), \( \text{O} \), or \( \text{F} \), and they interact with a lone pair on another \( \text{N} \), \( \text{O} \), or \( \text{F} \).


Boiling Points of Group 14, 15, 16, and 17 Hydrides

When we plot the boiling points of the hydrides of Group 14, 15, 16, and 17, we observe two important patterns:

General Trend: Going down each group (e.g. \( \text{H}_2\text{S} \rightarrow \text{H}_2\text{Se} \rightarrow \text{H}_2\text{Te} \)), boiling points increase because the molecules have more electrons, resulting in stronger van der Waals' forces.

The Anomalies: Three compounds have anomalously high boiling points that break the downward trend:
Ammonia (\( \text{NH}_3 \)) in Group 15
Water (\( \text{H}_2\text{O} \)) in Group 16
Hydrogen fluoride (\( \text{HF} \)) in Group 17

Why? These three molecules contain hydrogen bonds in addition to van der Waals' forces, requiring significantly more energy to separate the molecules.

Did you know? Water has a higher boiling point (\( 100^\circ\text{C} \)) than \( \text{HF} \) (\( 19.5^\circ\text{C} \)), even though fluorine is more electronegative than oxygen. This is because each water molecule can form an average of two hydrogen bonds per molecule (it has \( 2 \) hydrogens and \( 2 \) lone pairs), creating an extensive 3D network. \( \text{HF} \) can only form an average of one hydrogen bond per molecule because it is limited by having only one hydrogen atom.


7. The Unusual Properties of Water

Hydrogen bonding gives water two remarkable properties that are essential for life on Earth:

1. Relatively High Melting and Boiling Points

For such a small molecule (\( M_r = 18 \)), water has an exceptionally high boiling point (\( 100^\circ\text{C} \)) because of the extensive network of hydrogen bonds between molecules.

2. Ice is Less Dense than Liquid Water

Unlike almost all other substances, solid water (ice) is less dense than liquid water, which is why icebergs and ice cubes float.

Explanation: When water freezes into ice, the molecules arrange themselves into an open, rigid 3D tetrahedral lattice held together by fixed hydrogen bonds. This open structure holds the water molecules further apart than they are in the liquid state. When ice melts, this rigid lattice collapses, allowing molecules to pack closer together, making liquid water denser.

Biological Significance: Ice forms an insulating layer on top of lakes and ponds in winter, allowing aquatic life to survive underneath!


8. Solubility in Water

A substance dissolves in water if the attractions formed between the solute and water molecules are strong enough to replace the intermolecular forces originally present within the solute and within water.

Soluble substances: Small molecules that can form hydrogen bonds with water (such as methanol \( \text{CH}_3\text{OH} \), ethanol \( \text{CH}_3\text{CH}_2\text{OH} \), and ammonia \( \text{NH}_3 \)) are freely miscible and dissolve readily in water.

Insoluble substances: Non-polar substances like alkanes (e.g. hexane \( \text{C}_6\text{H}_{14} \)) cannot form hydrogen bonds with water. They only interact via weak van der Waals' forces, which cannot overcome the strong hydrogen bonding between water molecules. Hence, they are insoluble.

Key Takeaway: "Like dissolves like" — polar and hydrogen-bonding substances dissolve in polar solvents like water; non-polar substances dissolve in non-polar solvents.


9. Quick Summary & Common Mistakes to Avoid

Summary of Relative Strengths

Covalent Bonds \( \gg \) Hydrogen Bonds \( > \) Permanent Dipole–Dipole \( > \) van der Waals' Forces

Common Mistakes to Avoid in Exams:

Mistake 1: Saying "covalent bonds break when water boils".
Correction: Only intermolecular hydrogen bonds break during boiling; covalent \( \text{O-H} \) bonds remain intact.

Mistake 2: Thinking van der Waals' forces are only present in non-polar molecules.
Correction: van der Waals' forces exist between all molecules.

Mistake 3: Forgetting lone pairs when drawing hydrogen bonds.
Correction: Always draw the lone pair of electrons on the \( \text{N} \), \( \text{O} \), or \( \text{F} \) atom, and draw the dashed bond directly to that lone pair.

Mistake 4: Believing all molecules containing hydrogen exhibit hydrogen bonding.
Correction: Hydrogen MUST be bonded directly to Nitrogen, Oxygen, or Fluorine (e.g. \( \text{CH}_4 \) and \( \text{HCl} \) do NOT form hydrogen bonds!).