Intermolecular forces
Section: 3 Chemical Bonding | Syllabus: Cambridge AS Level Physics 9702
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 Tip Always 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 …
Interactive revision notes, videos and practice questions load below.