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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

S2.2 - The covalent model

2.2.1 Covalent Bonds and Lewis Formulas 2.2.2 Bond Types 2.2.3 Co-coordination (Dative) Bonds 2.2.4 VSEPR Shapes of Molecules 2.2.5 Electronegativity and Bond Polarity 2.2.6 Polarity and Dipole Moments 2.2.7 Covalent Network Structures and Allotropes 2.2.8 Intermolecular Forces 2.2.9 Physical Properties of Covalent Substances 2.2.10 Chromatography and Intermolecular Forces 2.2.11 Resonance Structures (AHL) 2.2.12 Benzene and Resonance (AHL) 2.2.13 Expanded Octet and VSEPR (AHL) 2.2.14 Formal Charge (AHL) 2.2.15 Sigma and Pi Bonds (AHL) 2.2.16 Hybridization (AHL)

Physical Properties of Covalent Substances

Specification Reference S2.2.9

Quick Notes

  • Strength of intermolecular forces (IMFs):
    • London < Dipole–Dipole < Hydrogen Bonding
  • Volatility:
    • Weaker intermolecular forces make substances more volatile (more easily evaporate)
    • Stronger intermolecular forces make substances less volatile (higher boiling point)
  • Electrical conductivity:
    • Covalent compounds do not conduct electricity (no mobile ions or electrons)
    • Exceptions: substances like graphite (delocalised electrons)
  • Solubility:
    • Polar covalent compounds dissolve in polar solvents (e.g. water)
    • Non-polar covalent compounds dissolve in non-polar solvents (e.g. hexane)
    • Like dissolves like: solubility depends on intermolecular forces

Full Notes

Comparing Intermolecular Forces

For substances with similar molar mass, the strength of intermolecular forces follows this order:

London (dispersion) < Dipole–Dipole < Hydrogen Bonding

The bonding and intermolecular forces present in covalent substances explains their physical properties (such as volatility, solubility and electrical conductivity).

Volatility

Volatility is the tendency of a substance to vaporise.

A substance is more volatile if its intermolecular forces are weaker.

Stronger forces require more energy to overcome, leading to higher boiling points and making it harder to vapourise.

Examples:

IB Chemistry comparison diagram showing hydrogen bonding in H2O versus dipole–dipole forces in H2S and their impact on boiling point and volatility.

Solubility

Solubility depends on the intermolecular forces between the solute and solvent:

IB Chemistry diagram illustrating like dissolves like with polar water dissolving polar solutes and non-polar solvents dissolving non-polar solutes.

“Like dissolves like” is a helpful rule of thumb

Examples:

Dissolving Small Polar Molecules

Water also dissolves small molecules that can form hydrogen bonds, such as alcohols (e.g. ethanol), ammonia (NH3 and Simple carboxylic acids.

IB Chemistry comparison of hydrogen bonding interactions between water and small polar molecules such as alcohols, ammonia, and simple carboxylic acids.

The hydrogen bonds formed between these solutes and water make them soluble.

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Matt’s exam tip

Whether a substance dissolves in water depends on the relative strength of the interactions between solute and water molecules compared to the interactions within the solute itself and within water.

For example, ethanol is soluble in water because the intermolecular forces between ethanol and water molecules are similar in strength to the forces between only ethanol molecules and only water molecules. These comparable attractions make it energetically favourable for the substances to mix.

However, hexanol is not soluble in water. The hydrogen bonding between water molecules is stronger than the interactions between water and hexanol molecules. As a result, it is more energetically favourable for water molecules to stick together than to interact with hexanol – so the two do not mix well.

Electrical Conductivity

Covalent substances typically do not conduct electricity, because there are no free electrons or ions that can move and carry charge.

Exceptions:

Summary