Introduction to Intermolecular Forces
Welcome to one of the most exciting and essential topics in AS Chemistry! Have you ever wondered why water is a liquid at room temperature while carbon dioxide is a gas, even though carbon dioxide is heavier? Or why water expands when it freezes, allowing icebergs to float? The answer lies in intermolecular forces.
Don't worry if bonding felt tricky at GCSE. In this chapter, we will break down the forces that act between molecules into simple, easy-to-understand steps. Master this, and you will unlock high-scoring marks in exam questions on boiling points, physical properties, and molecular structures!
What you will learn:
• The difference between intramolecular bonds and intermolecular forces
• Van der Waals' forces (instantaneous dipole–induced dipole attractions)
• Permanent dipole–dipole attractions
• Hydrogen bonding (the strongest intermolecular force)
• How these forces explain physical properties like boiling points and the unique behavior of ice
1. The Big Picture: Intramolecular vs. Intermolecular
Before diving into the types of forces, let's make sure we never confuse these two words:
• Intramolecular bonds are the strong bonds inside a molecule holding its atoms together (such as covalent bonds in \( \text{H}_2\text{O} \) or \( \text{Cl}_2 \)). Breaking these requires a chemical reaction.
• Intermolecular forces are the much weaker attractions between separate molecules. When you boil water, you do not break the covalent \( \text{O}-\text{H} \) bonds inside the molecules; you only overcome the weak intermolecular forces between the water molecules!
Analogy Time: Think of a classroom of students. The bones holding each student's body together are like covalent bonds (very strong!). Two students high-fiving or gently holding hands across the aisle is like an intermolecular force (much weaker and easier to separate).
Quick Review: Relative Strengths
Covalent bonds \( \gg \) Hydrogen bonds \( > \) Permanent dipole–dipole \( > \) Van der Waals' forces
Key Takeaway: Physical changes (like melting and boiling) only break intermolecular forces, not covalent bonds.
2. Van der Waals' Forces (Instantaneous Dipole–Induced Dipole)
Van der Waals' forces (also called London dispersion forces) are the weakest type of intermolecular force, but they exist between all molecules and single atoms (like noble gases), whether they are polar or non-polar.
How Do They Form? (Step-by-Step)
Step 1: Electrons are in constant, random motion around atomic nuclei.
Step 2: At any single 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 instantaneous dipole (one end becomes slightly negative, \( \delta^- \), and the other slightly positive, \( \delta^+ \)).
Step 4: This temporary dipole repels or attracts electrons in a neighboring molecule, creating an induced dipole.
Step 5: The opposite charges attract each other briefly. This weak attraction is a van der Waals' force.
What Affects the Strength of Van der Waals' Forces?
Two main factors determine how strong these forces are:
1. Number of Electrons (Molecular Size):
The more electrons a molecule has, the larger its electron cloud, and the more easily the electron cloud can be distorted (we say it is more polarisable). Larger temporary dipoles are formed, leading to stronger van der Waals' forces and higher boiling points.
Example: Boiling points of the halogens increase down Group 7:
\( \text{F}_2 \) (\( 18 \) electrons) is a gas at room temp.
\( \text{Cl}_2 \) (\( 34 \) electrons) is a gas.
\( \text{Br}_2 \) (\( 70 \) electrons) is a liquid.
\( \text{I}_2 \) (\( 106 \) electrons) is a solid.
2. Surface Area and Molecular Shape:
Long, straight-chain molecules can pack closely together and have more points of surface contact with neighboring molecules. Branched molecules are more spherical, meaning they cannot pack as closely, resulting in fewer points of contact.
Example: Both pentane and 2,2-dimethylpropane have the formula \( \text{C}_5\text{H}_{12} \). However, straight-chain pentane has a boiling point of \( 36^\circ\text{C} \), while spherical 2,2-dimethylpropane boils at only \( 9.5^\circ\text{C} \) because of reduced surface contact!
Common Mistake to Avoid: Never say "iodine has higher boiling point because covalent bonds are stronger." The covalent \( \text{I}-\text{I} \) bond is NOT broken when iodine sublimes or melts; only the van der Waals' forces between \( \text{I}_2 \) molecules are overcome!
Key Takeaway: More electrons and longer, unbranched chains \( = \) stronger van der Waals' forces \( = \) higher boiling points.
3. Permanent Dipole–Dipole Attractions
Permanent dipole–dipole attractions occur only between polar molecules (molecules with a permanent dipole).
What Causes a Permanent Dipole?
When two atoms in a covalent bond have significantly different electronegativities, the shared electrons are pulled closer to the more electronegative atom. This creates a permanent polar bond with \( \delta^+ \) and \( \delta^- \) ends.
If the shape of the molecule is asymmetrical, the dipoles do not cancel out, leaving the entire molecule polar.
Example: In hydrogen chloride, \( \text{HCl} \), chlorine is more electronegative than hydrogen. The molecule has a permanent dipole: \( \text{H}^{\delta+}-\text{Cl}^{\delta-} \). The \( \delta^+ \) hydrogen of one \( \text{HCl} \) molecule attracts the \( \delta^- \) chlorine of a neighboring molecule.
Comparing Van der Waals' and Dipole–Dipole
For molecules of similar size and electron count (e.g., non-polar \( \text{F}_2 \), \( 18 \) electrons, vs polar \( \text{HCl} \), \( 18 \) electrons), the polar molecule will usually have a higher boiling point because permanent dipole–dipole forces provide an extra attractive force on top of the van der Waals' forces.
Key Takeaway: Permanent dipole–dipole forces occur between polar molecules and add extra attractive strength alongside van der Waals' forces.
4. Hydrogen Bonding
Hydrogen bonding is not a true chemical bond; it is the strongest type of intermolecular force. It is an extreme form of permanent dipole–dipole attraction.
When Does Hydrogen Bonding Occur?
For a hydrogen bond to form, two strict conditions must be met:
1. A hydrogen atom must be covalently bonded directly to a small, highly electronegative atom: Nitrogen (\( \text{N} \)), Oxygen (\( \text{O} \)), or Fluorine (\( \text{F} \)).
2. There must be an electronegative atom (\( \text{N} \), \( \text{O} \), or \( \text{F} \)) on a neighboring molecule with at least one lone pair of electrons.
Memory Trick: Remember "NOF" – hydrogen bonding is NO Fun without Nitrogen, Oxygen, or Fluorine!
Why is Hydrogen Bonding So Strong?
• \( \text{N} \), \( \text{O} \), and \( \text{F} \) are so electronegative that they pull electron density strongly away from \( \text{H} \), leaving the hydrogen nucleus almost entirely unshielded (a tiny, concentrated \( \delta^+ \) charge).
• Because hydrogen is tiny, it can get very close to the lone pair of electrons on a neighboring \( \text{N} \), \( \text{O} \), or \( \text{F} \) atom, creating an exceptionally strong electrostatic attraction.
Drawing a Hydrogen Bond for the Exam
CCEA examiners love asking you to draw hydrogen bonds (for example, between two water molecules). To get full marks, make sure your diagram includes all of the following:
1. Partial charges (\( \delta^+ \) and \( \delta^- \)) clearly labelled on all relevant atoms.
2. Lone pairs of electrons drawn on the \( \text{O} \), \( \text{N} \), or \( \text{F} \) atoms.
3. The hydrogen bond represented as a dashed or dotted line extending directly from a lone pair to the \( \delta^+ \) hydrogen.
4. A straight bond angle of approximately \( 180^\circ \) along the \( \text{O}-\text{H}\cdots\text{O} \) line.
5. The Anomalous Properties of Water
Hydrogen bonding gives water several unusual ("anomalous") properties that are essential for life on Earth:
1. Unusually High Boiling and Melting Points
Group 6 hydrides follow the sequence \( \text{H}_2\text{S} \), \( \text{H}_2\text{Se} \), \( \text{H}_2\text{Te} \). Their boiling points increase down the group because they have more electrons and therefore stronger van der Waals' forces.
However, \( \text{H}_2\text{O} \) completely breaks this trend with a much higher boiling point of \( 100^\circ\text{C} \) (compared to \( -60^\circ\text{C} \) for \( \text{H}_2\text{S} \)). This is because water forms extensive hydrogen bonds (each \( \text{H}_2\text{O} \) molecule can form up to 4 hydrogen bonds), requiring a large amount of thermal energy to separate the molecules.
2. Ice is Less Dense than Liquid Water
Unlike most substances, solid water (ice) is less dense than liquid water, which is why ice floats!
Why? In ice, each water molecule is held in a rigid, open 3D tetrahedral lattice by four hydrogen bonds. This open cage-like structure holds the molecules further apart than in liquid water. When ice melts, the crystal lattice collapses, and water molecules pack closer together, making liquid water denser than ice.
Did you know? Because ice floats, it insulates lakes and rivers during freezing winters, allowing aquatic life to survive underneath!
Key Takeaway: Hydrogen bonding gives water an exceptionally high boiling point and makes ice less dense than liquid water due to an open tetrahedral structure.
6. Summary & Exam Quick-Check Guide
Intermolecular Forces Summary Table
1. Van der Waals' Forces:
• Found in: All atoms and molecules.
• Relative Strength: Weakest.
• Key Feature: Caused by temporary dipoles; strength increases with number of electrons and surface contact area.
2. Permanent Dipole–Dipole:
• Found in: Polar molecules only (e.g., \( \text{HCl} \), \( \text{CH}_3\text{Cl} \)).
• Relative Strength: Medium.
• Key Feature: Electrostatic attraction between permanent \( \delta^+ \) and \( \delta^- \) ends of molecules.
3. Hydrogen Bonding:
• Found in: Molecules with \( \text{H}-\text{N} \), \( \text{H}-\text{O} \), or \( \text{H}-\text{F} \) bonds (e.g., \( \text{H}_2\text{O} \), \( \text{NH}_3 \), \( \text{HF} \), alcohols).
• Relative Strength: Strongest intermolecular force.
• Key Feature: Attraction between lone pair on \( \text{N} \), \( \text{O} \), or \( \text{F} \) and a \( \delta^+ \) hydrogen atom.
How to Answer Boiling Point Comparison Questions (Step-by-Step)
When asked to explain why substance A has a higher boiling point than substance B:
1. Identify the type of force present in each substance (e.g., "Substance A has hydrogen bonds, whereas Substance B only has van der Waals' forces").
2. Compare their relative strengths (e.g., "Hydrogen bonds are stronger than van der Waals' forces").
3. State energy requirement (e.g., "Therefore, more thermal energy is needed to overcome the forces between molecules in Substance A").