Skip to content
Dabour HubLEARN · PRACTICE · EXPLORESign inStudent

2.2Unit 2. Energetics, Group Chemistry, Halogenoalkanes and Alcohols

Intermolecular Forces — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

5 min read 12 parts Lesson 2 of 53 1 quiz

📚 IAL Edexcel Linear (Chemistry) Hub / Topical Revision
🩺
100% Curriculum Specification Match Lesson Doctor Verified
The only platform with 100% curriculum specification match and verified interactive question banks.
Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution
pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴
Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14
Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.
0%read
Lesson summary

2.2 Intermolecular Forces — types, hydrogen bonding, physical properties and solvents — IAL EDEXCEL LINEAR (CHEMISTRY) notes: Distinguish intramolecular from intermolecular forces, explain London, permanent dipole and hydrogen bonding, and use them to explain boiling points, the …

  • Intramolecular forces
  • London force
  • Permanent dipole-dipole force
  • Surface tension
  • Intermolecular forces — weaker forces between molecules
  • There are three types.Induced dipole–dipole (London, van der Waals) forces

Part 1 of 10Intramolecular and intermolecular forces

Intramolecular forces
  • Intramolecular forces forces within a molecule, usually covalent bonds.
    • (single, double, triple and dative)
Intermolecular forces
  • Intermolecular forces weaker forces between molecules.
There are three types.
  • Induced dipole–dipole (London, van der Waals) forces.
  • Permanent dipole–dipole forces.
  • Hydrogen bonding, a special permanent dipole–dipole force.
Intramolecular forces are much stronger than intermolecular forces; a hydrogen bond is about one tenth the strength of a covalent bond.
Boiling a simple molecular substance breaks intermolecular forces, not covalent bonds.
In giant covalent structures such as diamond, covalent bonds must be broken to melt the solid, so melting points are very high.

Part 2 of 10Induced dipole–dipole (London) forces

Induced dipole–dipole (London) forces
They exist between all atoms and molecules.
The electron cloud is constantly moving, so at any instant one side of a molecule can have more electron density.
  • This creates a temporary (instantaneous) dipole.
  • It induces a dipole in a neighbouring molecule, and the δ+ and δ− ends attract.
  • The dipoles are only temporary because the electrons keep moving.
For small molecules with the same number of electrons, permanent dipole forces are stronger than induced dipole forces, e.g. propanone boils higher than butane.

Part 3 of 10Permanent dipole–dipole forces

Permanent dipole–dipole forces
Polar molecules have a permanent dipole, with always a δ+ end and a δ− end.
The δ+ end of one molecule attracts the δ− end of a neighbour, e.g. HCl, CH3Cl, propanone.
These forces add to the London forces already present.

Part 4 of 10Hydrogen bonding

Hydrogen bond
  • Hydrogen bond the strongest intermolecular force, between a δ+ hydrogen bonded to O, N or F and a lone pair on O, N or F of another molecule.
The O-H, N-H or H-F bond is so polarised that the hydrogen is strongly δ+ and small.
Water can form two hydrogen bonds per molecule because oxygen has two lone pairs.
Compounds with hydrogen bonds include alcohols, ammonia, amines, carboxylic acids, hydrogen fluoride and proteins.
Draw a hydrogen bond as a dashed straight line from the lone pair to the δ+ hydrogen.
ForceFound inRelative strengthExample
London forcesAll moleculesWeakest (increase with electrons)Noble gases, alkanes
Permanent dipole-dipolePolar moleculesIntermediateHCl, propanone
Hydrogen bondingH bonded to N, O or FStrongest intermolecularH2O, NH3, HF

Part 5 of 10Anomalous properties of water

Anomalous properties of water
High melting and boiling points
  • Hydrogen bonds in ice and liquid water are strong, so a lot of energy is needed to separate the molecules.
  • The enthalpy of vaporisation of H2O is almost three times what its electrons alone would give (about 17 kJ mol−1), so more than London forces act.
  • Down H2S to H2Te the enthalpy of vaporisation rises because more electrons give stronger London forces.
High surface tension
  • Surface molecules are hydrogen bonded to molecules below and beside them and are pulled inwards.
  • Inner molecules are pulled in all directions, so the surface is compressed.
Ice is less dense than liquid water
  • Water molecules are held in an open lattice by hydrogen bonds with relatively long bond lengths.
  • The molecules are slightly further apart than in the liquid, so ice is about 9% less dense.
  • The H-O-H angle in ice approaches 109.5° because lone pairs are drawn towards neighbouring δ+ hydrogens, reducing repulsion.

Part 6 of 10Effect of molecular size and shape on boiling point

Effect of molecular size and shape on boiling point
Number of electrons
  • More electrons (greater Mr) allow larger and more frequent temporary dipoles.
  • London forces, and so enthalpy of vaporisation, melting and boiling points, increase.
  • Noble gas boiling points rise down the group for this reason.
Branching
  • Branching reduces the surface area of a molecule, so there is less contact with neighbours.
  • Less induced dipole means weaker London forces and a lower boiling point than the straight-chain isomer.

Part 7 of 10Hydrogen bonding in alcohols and hydrogen halides

Hydrogen bonding in alcohols and hydrogen halides
Alcohols
  • The O-H bond lets δ+ hydrogen bond to lone pairs on oxygen in other molecules.
  • Alkanes have only London forces, so alcohols boil higher, e.g. propane −42 °C and propanol 97 °C.
Hydrogen halides
  • Boiling points of HCl, HBr and HI rise as the molecules get larger and have more electrons.
  • HF is out of line because it also forms hydrogen bonds: H is strongly δ+ and fluorine has small, intense lone pairs.

Part 8 of 10Solvent choice

1
Like dissolves like: new solute–solvent attractions must be similar in strength to those broken.
2
Non-polar substances dissolve in non-polar solvents such as hydrocarbons through London forces.
3
Polar covalent substances dissolve in polar solvents through dipole–dipole attractions or hydrogen bonds, e.g. ethanol in water.
4
Solubility falls as the non-polar part grows.
  • Ethanol is soluble in water; hexanol is not.
5
Polar substances dissolve poorly in non-polar solvents because their dipole attractions cannot interact with the solvent.
6
Giant covalent substances do not dissolve because the strong covalent bonds cost too much energy to break.
7
Halogenoalkanes are only partly soluble in water.
  • The C-X bond is polar, but there are no H-F, H-O or H-N bonds, so they cannot hydrogen bond with water.

Part 9 of 10Dissolving ionic compounds

1
Many ionic compounds dissolve in polar solvents such as water.
2
Two changes occur.
  • The ionic lattice is broken down.
  • Polar molecules surround the ionsδ+ hydrogens surround anions and δ− oxygens surround cations.
3
Solubility depends on the strength of attractions within the lattice compared with those between ions and solvent.
4
In general, the greater the ionic charge, the less soluble the compound, e.g. 356.9 g of NaCl but 74.4 g of CaCl2 dissolve in 1 dm3 of water; there are many exceptions.

Part 10 of 10More from the specification

More from the specification
London dispersion forces
  • London dispersion forces temporary dipoles from electron movement induce dipoles in neighbouring atoms,
  • giving weak attraction that grows with the number of electrons
Permanent dipole interactions
  • Permanent dipole interactions a polar molecule has positive and negative ends that attract
Hydrogen bonding
  • Hydrogen bonding strong attraction involving H bonded to N, O or F
Solvent choice
  • Solvent choice a solute dissolves if the energy released by new interactions compensates the energy required to break the old ones,
    • so like dissolves like
Dispersion force
  • Another name for induced dipole–induced dipole force.

Key terms

TermMeaning
Intermolecular forceA weak attraction between separate molecules.
Intramolecular forceA force within a molecule or lattice, such as a covalent bond.
London forceAn attraction between an instantaneous dipole and the dipole it induces in a neighbour.
Permanent dipole-dipole forceAn attraction between the δ+ end of one polar molecule and the δ− end of another.
Hydrogen bondAn attraction between δ+ H bonded to N, O or F and a lone pair on N, O or F of another molecule.
Surface tensionThe ability of a liquid surface to resist external forces.
HydrationThe surrounding of ions by water molecules in solution.

Avoid these mistakes

1
Saying covalent bonds break when a simple molecular substance boils — only intermolecular forces are overcome.
2
Saying hydrogen bonding needs only a hydrogen atom — hydrogen must be bonded to O, N or F and there must be a lone pair on the other molecule.
3
Using hydrogen bonds to explain why HCl boils higher than HBr — it is the number of electrons, since HCl has no hydrogen bonds.
4
Ignoring surface area when comparing isomers — branching lowers boiling point.

Last updated:

\n