Intermolecular forces
Section: 3 Chemical Bonding | Syllabus: Cambridge AS Level Chemistry (9701)
Bond Polarity and Dipole Moments
A bond is polar when the two bonded atoms have different electronegativities - the more electronegative atom pulls the shared electrons closer, developing a partial negative charge (δ−) while the other atom becomes δ+.
- The separation of charge creates a bond dipole, shown as δ+ → δ− or with an arrow pointing towards the more electronegative atom.
- Whether a molecule is polar depends on both the bond dipoles and the molecular shape.
- In a symmetrical molecule, bond dipoles cancel → non-polar molecule (e.g. CO₂, CCl₄, BF₃).
- In an asymmetric molecule, bond dipoles do not cancel → polar molecule with a net dipole moment (e.g. H₂O, NH₃, HCl).
Figure 3.15: Polar vs Non-polar Molecules
(Left: CO₂ - linear, two C=O dipoles pointing in opposite directions, cancel → non-polar. Centre: H₂O - V-shaped, two O–H dipoles both point toward O, net dipole upward → polar molecule. Right: CCl₄ - tetrahedral, four C–Cl dipoles cancel symmetrically → non-polar. Label δ+ and δ− on each bond and show net dipole arrow where applicable.)
Exam TipAlways draw the shape of the molecule first (use VSEPR), then assess whether bond dipoles cancel. A molecule can have polar bonds but be overall non-polar if the shape is symmetrical (e.g. BF₃, CO₂, CCl₄).
Van der Waals' Forces - Overview
Van der Waals' Forces A generic term for all intermolecular forces between molecular entities, other than those due to bond formation. Includes instantaneous dipole–induced dipole (id–id) forces, permanent dipole–permanent dipole (pd–pd) forces, and hydrogen bonding.| Type | Also called | Acts between | Relative strength |
|---|---|---|---|
| id–id forces | London dispersion forces | All molecules (polar and non-polar) | Weakest (but strongest for large molecules) |
| pd–pd forces | Dipole–dipole interactions | Polar molecules | Moderate |
| Hydrogen bonding | Special case of pd–pd | Molecules with N–H or O–H | Strongest intermolecular force |
Instantaneous Dipole–Induced Dipole (id–id) Forces
Also called London dispersion forces. These exist between all molecules - polar or non-polar.
- Electrons in a molecule are in constant motion. At any instant, their distribution may be uneven, creating a temporary (instantaneous) dipole.
- This instantaneous dipole induces a dipole in a neighbouring molecule - the electrons in the neighbour are attracted or repelled, creating a corresponding dipole.
- The instantaneous and induced dipoles attract each other - this is the id–id force.
- id–id forces increase with: more electrons (larger molecule/atom), larger surface area of the molecule.
Figure 3.16: Instantaneous Dipole–Induced Dipole Forces
(Three sequential diagrams: (1) two non-polar molecules, electron clouds uniform. (2) Electron cloud in left molecule fluctuates → instantaneous δ+/δ− dipole. (3) Dipole induces opposite dipole in right molecule; attraction shown between δ− of left and δ+ of right.)
Why do boiling points increase down Group 17?F₂ (bp −188°C) < Cl₂ (bp −34°C) < Br₂ (bp +59°C) < I₂ (bp +184°C)
Down the group, the number of electrons increases → larger, more polarisable electron clouds → stronger id–id forces → more energy needed to overcome them → higher boiling points.
Permanent Dipole–Permanent Dipole (pd–pd) Forces
- Exist between polar molecules that have a permanent dipole.
- The δ+ end of one molecule is attracted to the δ− end of a neighbouring molecule.
- Stronger than id–id forces between molecules of similar size, because the dipole is permanent rather than temporary.
- Examples: HCl, propanone (CH₃COCH₃), SO₂.
Figure 3.17: Permanent Dipole–Permanent Dipole Forces in HCl
(Row of HCl molecules aligned with δ+ H end of one molecule pointing toward δ− Cl end of the next. Dotted lines between molecules indicating pd–pd attraction. Label δ+ on H and δ− on Cl for each molecule.)
Hydrogen Bonding
Hydrogen Bond A special, stronger type of permanent dipole–permanent dipole force. It occurs when hydrogen is bonded to a highly electronegative atom (N, O, or F), making the H atom strongly δ+. This δ+ H is attracted to a lone pair on an electronegative atom (N, O, or F) of a neighbouring molecule.- Hydrogen bonding only occurs when H is directly bonded to N, O, or F - the most electronegative elements.
- The H atom must also be attracted to a lone pair on N, O, or F of a neighbouring molecule.
- Represented as a dotted line: X–H···Y (X and Y are N, O, or F; ··· is the hydrogen bond).
- Hydrogen bonds are much stronger than ordinary pd–pd forces but weaker than covalent bonds.
Figure 3.18: Hydrogen Bonding in Water
(Three or four water molecules. Each O has two lone pairs shown. Dotted lines from δ+ H of one molecule to lone pair on O of adjacent molecule. Each water molecule can form up to 4 hydrogen bonds (2 as donor, 2 as acceptor). Label the O–H covalent bond and the H···O hydrogen bond.)
Figure 3.19: Hydrogen Bonding in Ammonia
(Two NH₃ molecules. Dotted line from δ+ H of one molecule to lone pair on N of adjacent molecule. Note: NH₃ can form fewer H-bonds than H₂O (1 lone pair vs 2 on O), so is weaker overall.)
Anomalous Properties of Water Explained by Hydrogen Bonding
1. Relatively High Melting and Boiling Points
- Water has unusually high mp (0°C) and bp (100°C) for such a small molecule.
- Comparing Group 16 hydrides: H₂S bp = −60°C, H₂Se bp = −41°C - both far lower than H₂O.
- Extra energy is needed to break the hydrogen bonds between water molecules in addition to id–id forces.
Figure 3.20: Boiling Points of Group 16 Hydrides
(Line graph: boiling point (°C) on y-axis vs group 16 hydrides H₂Te, H₂Se, H₂S, H₂O on x-axis (right to left). H₂S, H₂Se, H₂Te show a smooth downward trend. H₂O is anomalously high - far above the expected extrapolated trend line. Annotate: "anomalously high due to hydrogen bonding".)
2. Relatively High Surface Tension
- Water molecules at the surface experience hydrogen bonding only inward and sideways - not upward.
- This creates a net inward force, pulling surface molecules together and creating a high surface tension.
- This is why small objects (e.g. paper clips) can rest on the water surface without sinking.
3. Ice is Less Dense than Liquid Water
- In liquid water, hydrogen bonds are constantly forming and breaking - the structure is partially disordered. Molecules are relatively close together.
- In ice, water molecules form a rigid, open hexagonal lattice - each molecule forms 4 hydrogen bonds in a fixed arrangement.
- The open lattice structure of ice has more space between molecules than liquid water → ice is less dense than liquid water (density of ice ≈ 0.917 g cm⁻³ vs liquid water ≈ 1.000 g cm⁻³).
- This is why ice floats on water - vital for aquatic life in cold climates.
Figure 3.21: Structure of Ice vs Liquid Water
(Left: ice - ordered hexagonal lattice with large empty spaces between molecules, each O forming 4 H-bonds shown as dotted lines. Right: liquid water - disordered arrangement, molecules closer together on average, H-bonds constantly breaking/forming. Label the open spaces in ice vs the compact arrangement in liquid water.)
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