3.6 Intermolecular forces and bond properties
- Syllabus
- 9701–2028–2029
- Topic
- 3.6
- Level
- AS
Learning objectives
3.6.1Hydrogen bonding• Ammonia and water as simple examples- (a) describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including- (b) use the concept of hydrogen bonding to explain the anomalous properties of H₂O (ice and water):- its relatively high melting and boiling points- its relatively high surface tension- the density of solid ice compared with liquid water3.6.2Electronegativity and bond polarity• Use the concept of electronegativity to explain bond polarity and dipole moments of molecules3.6.3Van der Waals’ forces as the intermolecular• Those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces- (a) describe van der Waals’ forces as the intermolecular forces between molecular entities other than- (b) describe the types of van der Waals’ forces:- instantaneous dipole–induced dipole (id-id) forces, also called London dispersion forces- permanent dipole–permanent dipole (pd-pd) forces, including hydrogen bonding- (c) describe hydrogen bonding and understand that hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom3.6.4That, in general, ionic, covalent and metallic• State that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces
Hydrogen bonding links molecules containing N–H or O–H groups
Within the assessed scope, hydrogen bonding is a strong intermolecular attraction between a δ⁺ hydrogen covalently bonded to N or O in one molecule and a lone pair on N or O in a neighbouring molecule. It is a special case of permanent dipole–permanent dipole attraction.
Ammonia and water both hydrogen-bond because their N–H or O–H bonds are strongly polar and the N or O atoms carry lone pairs. In water, each molecule has two donor hydrogens and two acceptor lone pairs, so an extended intermolecular network can form.
| Anomalous property of H₂O | Hydrogen-bond explanation |
|---|---|
| relatively high melting and boiling points | extra energy is required to overcome attractions between water molecules |
| relatively high surface tension | surface molecules are pulled together strongly by neighbouring water molecules |
| ice is less dense than liquid water | ice has a rigid open hydrogen-bonded lattice; partial collapse on melting lets molecules pack more closely |
The O–H or N–H link within a molecule is a covalent bond; the hydrogen bond acts between molecules. Hydrogen in a formula is not sufficient by itself—the hydrogen must be directly bonded to N or O for the cases assessed here.
Bond dipoles arise from uneven sharing and add according to molecular shape
Electronegativity is an atom's tendency to attract the shared electron pair in a covalent bond. Equal or nearly equal attraction gives a non-polar bond; unequal attraction shifts electron density towards the more electronegative atom, leaving that end δ− and the other end δ+.
A bond dipole records the direction of this charge separation, with the dipole arrow pointing towards the partially negative end. A larger electronegativity difference gives a more polar bond, but bond polarity is a property of the individual bond before the whole molecular arrangement is considered.
To decide whether a molecule has an overall dipole, identify the polarity and direction of every polar bond, then use the molecular shape to see whether the dipoles cancel. In CH₃Cl they do not cancel, so the molecule is polar; in the symmetric CCl₄ arrangement they cancel, so the molecule is non-polar despite its polar C–Cl bonds.
Do not infer overall molecular polarity from one bond or from the presence of polar bonds alone. Keep δ+ and δ− on the correct atoms, and distinguish a polar bond from a polar molecule whose net dipole depends on three-dimensional arrangement.
Van der Waals’ forces include id–id and pd–pd attractions
Van der Waals’ forces is the syllabus's generic term for intermolecular forces between molecular entities other than forces due to bond formation. The assessed categories are instantaneous dipole–induced dipole (id–id) forces and permanent dipole–permanent dipole (pd–pd) forces; hydrogen bonding is a special pd–pd case.
| Type | How the dipoles arise | Where it acts |
|---|---|---|
| id–id (London dispersion) | electron motion creates an instantaneous dipole, which induces a dipole in a neighbour | between all atoms and molecules |
| pd–pd | polar bonds and molecular shape give neighbouring molecules permanent dipoles | between polar molecules |
| hydrogen bonding | an especially δ⁺ H bonded to N or O is attracted to a lone pair on N or O nearby | between suitable N–H/O–H molecules |
Id–id forces generally strengthen when a species has more electrons because its electron cloud is more polarisable; greater molecular contact can also strengthen the total attraction. Pd–pd forces additionally align opposite permanent partial charges.
An instantaneous dipole fluctuates and is not a permanent molecular dipole. Do not call every intermolecular attraction a hydrogen bond, and do not confuse any van der Waals force between molecules with the covalent bonds within them.
Bonding is generally stronger than intermolecular forces
In general, ionic, covalent and metallic bonding are stronger than intermolecular forces. Bonding holds ions, atoms or metal ions and delocalised electrons together; intermolecular forces attract separate molecular entities.
| Interaction | Particles it connects | Typical role |
|---|---|---|
| ionic, covalent or metallic bonding | particles within a lattice or atoms within a molecular entity/network | determines the substance's bonded structure |
| intermolecular forces | separate molecular entities | influences their separation in melting, boiling and other physical changes |
When a simple molecular substance melts or boils, intermolecular attractions are overcome while the covalent bonds inside each molecule normally remain intact. This is why a low boiling point does not mean the molecule's covalent bonds are weak.
The syllabus statement is a general comparison, not a universal numerical ranking of every individual interaction. First identify which particles are being separated and whether the relevant attraction is bonding or intermolecular.