Ch 6 · Intermolecular forces, electronegativity and bond properties · AS & A Level Chemistry 9701

Topic 6 of 7 in Chemical bonding — AS & A Level Chemistry 9701 — 58 lessons.

Introduction

Intermolecular forces are much weaker than the bonds holding a molecule together, yet they control physical properties such as melting point, boiling point and solubility. Cambridge questions require you to name the correct type of force from a structure, and to use hydrogen bonding specifically to explain the unusual behaviour of water and ice.

Key ideas — Bond polarity and dipole moments

When two atoms of different electronegativity form a covalent bond, the shared electron pair lies closer to the more electronegative atom, giving it a small negative charge (delta-) and leaving the other atom slightly positive (delta+). This separation of charge is a bond dipole. If a molecule's bond dipoles do not cancel by symmetry, the molecule has a net dipole moment and is described as polar; H_2O is polar because its two O-H dipoles do not cancel in its bent shape, whereas CO_2 is non-polar because its two identical, opposing C=O dipoles cancel exactly in its linear shape.

🧑‍🏫 Teacher's explanation

Imagine a game of tug-of-war where one person is much stronger than the other. Even though they are both pulling on the rope, the rope moves slightly toward the stronger person. In a covalent bond, the 'rope' is the pair of shared electrons, and the 'people' are the atoms. If one atom is much more electronegative, it pulls the electrons closer to itself.

📖 Definitions

🔍 Worked example — Easy

Explain why the ceHCl molecule is polar using the concept of electronegativity.

🔍 Worked example — Standard

Compare the polarity of ceCCl4 and ceCH3Cl.

🔍 Worked example — Exam-style

A molecule has two bond dipoles, each with a magnitude of 0.5 Debye, oriented at an angle of 180^° to each other. Calculate the net dipole moment.

⚠️ Common errors

🧠 Quick check 1

🧠 Quick check 2

🎯 Exam question

Key ideas — Van der Waals' forces

Van der Waals' forces is a generic term for the weak intermolecular forces of attraction that exist between molecules for reasons other than covalent, ionic or metallic bond formation. Two types are distinguished by their origin:

🧑‍🏫 Teacher's explanation

Imagine a crowded dance floor. Most people are moving around randomly, but sometimes, two people happen to walk past each other in a way that they momentarily lean towards one another. This is like an instantaneous dipole—a random, fleeting moment of attraction. Now, imagine two people who are both wearing magnets on their clothes. They will naturally feel a pull towards each other every time they get close. This is like a permanent dipole.

📖 Definitions

🔍 Worked example — Easy

Explain why methane, CH_4, is a gas at room temperature while water, H_2O, is a liquid, despite both having similar molar masses.

🔍 Worked example — Standard

Predict the relative boiling points of Cl_2 and I_2.

🔍 Worked example — Exam-style

Calculate the energy required to vaporize 1.00 mol of a substance if the enthalpy of vaporization Δ H_vap is 35.0 kJ/mol.

⚠️ Common errors

🧠 Quick check 1

🧠 Quick check 2

🎯 Exam question

Key ideas — Hydrogen bonding

Hydrogen bonding is a particularly strong type of permanent dipole–permanent dipole force. It occurs when a hydrogen atom is covalently bonded directly to a small, highly electronegative atom — fluorine, oxygen or nitrogen — giving the hydrogen atom a large delta+ charge; this hydrogen atom is then strongly attracted to a lone pair on an F, O or N atom of a neighbouring molecule. In ammonia, hydrogen bonds form between the delta+ hydrogen of one N-H bond and the lone pair on the nitrogen of another molecule. In water, each molecule can form up to four hydrogen bonds, since oxygen has two lone pairs and two O-H hydrogens.

🧑‍🏫 Teacher's explanation

Imagine a group of people trying to hold hands. Most people just lightly touch fingertips (weak van der Waals forces). However, if one person is incredibly magnetic and reaches out to grab the other person's hand firmly, that is like hydrogen bonding. It is a much stronger version of a standard dipole-dipole interaction.

📖 Definitions

🔍 Worked example — Easy

Explain why NH_3 has a higher boiling point than PH_3.

🔍 Worked example — Standard

Calculate the number of hydrogen bonds a single water molecule can form in liquid water.

🔍 Worked example — Exam-style

Explain why ice is less dense than liquid water.

⚠️ Common errors

🧠 Quick check 1

🧠 Quick check 2

🎯 Exam question

Key ideas — Anomalous properties of water and ice

Hydrogen bonding gives water and ice properties that are anomalous compared with similarly sized molecules with no hydrogen bonding. Water has an unusually high melting and boiling point, since extra energy is needed to break the hydrogen bonds between molecules in addition to overcoming van der Waals' forces. Water has an unusually high surface tension, since surface molecules are pulled inward by hydrogen bonds with no molecules above them to balance the force. Most importantly, ice is less dense than liquid water: in ice, each molecule forms four hydrogen bonds to neighbours in a rigid, open hexagonal lattice with relatively large gaps; on melting, some hydrogen bonds break and the lattice partially collapses, allowing molecules to pack more closely, so liquid water at 0 °C is denser than ice.

🧑‍🏫 Teacher's explanation

Imagine a crowded dance floor. In most liquids, the dancers (molecules) are like people moving randomly, bumping into each other but staying close together. However, water is like a group of dancers performing a highly synchronized, rigid square dance. Because they are all holding hands very tightly (hydrogen bonds), they can't get close to each other; they are stuck in a fixed, spacious formation. This is why ice has more empty space between molecules than liquid water.

📖 Definitions

🔍 Worked example — Easy

Explain why the boiling point of H_2O is significantly higher than that of H_2S, despite H_2S having a higher molar mass.

🔍 Worked example — Standard

Calculate the density of ice if its volume is 1.06 dm^3 and its mass is 0.999 g (assume 1 dm^3 = 1000 cm^3).

🔍 Worked example — Exam-style

In a crystal lattice of ice, each water molecule is hydrogen-bonded to 4 other molecules. If a sample of ice contains 3.01 × 10^22 water molecules, calculate the total number of hydrogen bonds in the lattice (assume each bond is shared by 2 molecules).

⚠️ Common errors

🧠 Quick check 1

🧠 Quick check 2

🎯 Exam question

Key ideas — Relative strength of bonding and intermolecular forces

Ionic, covalent and metallic bonding are, in general, considerably stronger than intermolecular forces, including hydrogen bonding. This is why simple molecular substances such as ice or iodine melt at low temperatures — only the weak intermolecular forces between molecules are broken on melting, while the strong covalent bonds within each molecule remain intact.

🧑‍🏫 Teacher's explanation

Imagine you are at a crowded party. The people in the room are like molecules. The 'strong bonds' (ionic, covalent, metallic) are like the individual people themselves—their bones, muscles, and internal structure. The 'intermolecular forces' are like the social connections between people, such as a handshake or a conversation. It takes a lot of energy to physically break a person apart, but it takes very little energy to make them stop talking to each other and walk away.

Ch 6 · Intermolecular forces, electronegativity and bond properties · AS & A Level Chemistry 9701

All chapters in Chemical bonding