Intermolecular Forces in Covalent Substances
Quick Notes
- Intermolecular forces (IMFs) are forces of attraction between molecules.
- Intermolecular forces determine the physical properties (e.g. boiling point, solubility) of covalent compounds.
- Types of forces:
- London (dispersion): present in all molecules; stronger in larger/heavier ones.
- Dipole-induced dipole: between a polar molecule and a non-polar one.
- Dipole–dipole: between polar molecules.
- Hydrogen bonding: strong dipole–dipole force when H is bonded to N, O, or F and interacts with a lone pair on another electronegative atom.
- Van der Waals forces is an umbrella term covering London, dipole-induced dipole, and dipole–dipole forces
- The polarity and size of the molecule influence which forces are present and how strong they are.
Full Notes
What Are Intermolecular Forces?
Intermolecular forces (IMFs) are forces of attraction between molecules.
They are weaker than covalent bonds but crucial for determining boiling/melting points, volatility, solubility, and viscosity of molecular substances.
Types of Intermolecular Forces
There are different types of intermolecular force, with each type forming in different ways.
London (Dispersion) Forces
London Dispersions Forces are caused by temporary dipoles due to random movement of electrons
- They are present in all molecules, even non-polar ones.
- Strength increases with:
- More electrons (bigger atoms and molecules)
- Larger molecular surface area
Examples:Noble gases: He < Ne < Ar (increasing boiling point)
Dipole–Induced Dipole Forces
Dipole-induced dipole forces occur when a polar molecule induces a dipole in a non-polar molecule.
- Temporary and generally weak.
- Important in mixtures of polar + non-polar substances.
Example: O2 (non-polar) dissolved in H2O (polar)
Dipole–Dipole Forces
Dipole-dipole attractions occur between permanent dipoles in polar molecules.
They are stronger than London forces (but weaker than hydrogen bonds).
Example: HCl molecules attract each other via permanent dipole–dipole forces.
Hydrogen Bonding
Hydrogen bonding occurs between the H from an N-H, O-H or F-H bond and the lone pair of electrons on another N, O or F atom.
Examples: H2O, HF, NH3
Consequences of hydrogen bonding:
- High boiling points for H2O, HF
- Ice is less dense than water (open hydrogen-bonded structure)
Hydrogen bonding is always stronger than London or permanent dipole forces for comparably sized molecules.
Using “Van der Waals Forces”
Van der Waals forces is a general term used to describe:
- London dispersion forces
- Dipole–dipole forces
- Dipole-induced dipole forces
Hydrogen bonding is not included as an example of a van der Waals force.
Deduce Intermolecular Forces
To identify IMFs in a molecule:
- Check for polarity: Use electronegativity + molecular shape
- Look for H-bonding: Is H bonded to N, O, or F?
- Size and electron count: Larger molecules have stronger dispersion forces
Summary
| Force type | Present in | Relative strength | Key factors | Typical examples | Notes |
|---|---|---|---|---|---|
| London (dispersion) | All molecules and atoms | Weak to moderate | More electrons and larger surface area increase strength | Ar, I2, long-chain alkanes | Only IMF in non-polar substances |
| Dipole-induced dipole | Polar with non-polar | Weak | Polarizability of non-polar molecule and dipole strength | O2 in H2O | Important in solutions and mixtures |
| Dipole–dipole | Polar molecules | Moderate | Magnitude and orientation of dipoles | HCl, CH3Cl, propanone | Higher b.p. than similar non-polar molecules |
| Hydrogen bonding | H bonded to N, O, or F | Strongest IMF | Availability of lone pairs and linear H-bond geometry | H2O, HF, alcohols, amides | Leads to anomalies such as ice density, high b.p. of water |