Intramolecular Bonding - The Strong Stuff Inside Molecules

Ever wondered what actually holds atoms together inside molecules? Covalent bonds are the electrostatic attraction between positively charged nuclei and the negatively charged electrons they share. Think of it like two people sharing a blanket - sometimes fairly, sometimes not so much.

Non-polar covalent bonds form when two non-metals with identical electronegativity values share electrons equally. Examples include H-H, P-H, and C-S bonds where both atoms have the same "pulling power" for electrons.

Polar covalent bonds occur when atoms with different electronegativities bond together. The more electronegative atom (like fluorine) pulls the shared electrons closer, creating a slightly negative charge (δ-), whilst the other atom becomes slightly positive (δ+). Examples include H-F and Cl-H bonds.

Quick Tip: Remember that electronegativity differences determine bond type - small differences create polar covalent bonds, whilst identical values create non-polar bonds.

Ionic Bonding and Properties Comparison

When the electronegativity difference becomes massive, ionic bonding takes over. The more electronegative element completely steals an electron, creating charged ions. Think Na-Cl becoming Na+ and Cl- - it's like one person taking the entire blanket!

Ionic bonds are simply the electrostatic attraction between these positive and negative ions. Examples include NaCl and MgO, where complete electron transfer occurs.

Different bond types create dramatically different properties. Ionic compounds are typically solid at room temperature with high melting points (>400°C) and conduct electricity when dissolved or melted. Covalent molecular substances are often gases or low-melting solids that don't conduct electricity.

Intermolecular bonding explains why covalent substances can change state - only weak forces between molecules break, not the strong covalent bonds within them.

Exam Focus: Learn the property differences between ionic, covalent molecular, and covalent network substances - they're exam favourites!

Van der Waals Forces - The Weak but Important Attractions

Between molecules, three types of Van der Waals (VDW) forces operate, ranging from weak to relatively strong. These determine how molecules interact when they're not actually bonded together.

London Dispersion Forces (LDF) are the weakest VDW forces, present in all molecules. Their strength increases with the number of electrons - more electrons mean stronger temporary dipoles and higher boiling points.

Permanent dipole-permanent dipole $PD-PD$ interactions occur between polar molecules. These are much stronger than LDF. For example, trichloromethane (CHCl₃) has a higher boiling point than carbon dioxide because its dipoles don't cancel out due to asymmetry.

Remember: hydrocarbons are always non-polar, and molecules with formula XY₄ (like CCl₄) are non-polar due to symmetry.

Memory Trick: Think "Like attracts like" - polar molecules stick to other polar molecules through their permanent dipoles.

Hydrogen Bonding - The Strongest Intermolecular Force

Hydrogen bonding is the strongest VDW force and occurs only when hydrogen bonds to highly electronegative atoms: fluorine, oxygen, or nitrogen. You'll spot it in molecules containing F-H, O-H, or N-H bonds.

Water and ammonia are classic examples. The hydrogen atoms become so positive (δ+) that they're strongly attracted to lone pairs on neighbouring molecules' oxygen or nitrogen atoms.

These intermolecular forces directly influence three key properties of covalent molecular substances: melting and boiling points, viscosity, and solubility. The stronger the intermolecular forces, the more energy needed to separate molecules.

Think of hydrogen bonding as molecular velcro - much stronger than ordinary attractions but still weaker than actual covalent bonds within molecules.

Key Point: Only F-H, O-H, and N-H can form hydrogen bonds - memorise this trinity for exam success!

Predicting Physical Properties

To predict melting and boiling points, consider both the number of electrons and the VDW forces present. Propane, methoxymethane, and ethanol all have 26 electrons, but their boiling points differ dramatically: -44°C, -22°C, and 78°C respectively.

The difference? Propane (hydrocarbon) has only LDF, methoxymethane has permanent dipole interactions, whilst ethanol has hydrogen bonding - the strongest intermolecular force.

Viscosity measures how thick liquids are. Stronger intermolecular forces create higher viscosity because molecules resist flowing past each other. Substances with hydrogen bonding are more viscous, and larger molecules with more hydrogen bonding sites become even thicker.

Compare methanol (CH₃OH) with ethanol (C₂H₅OH) - ethanol is more viscous due to its larger molecular size allowing more intermolecular interactions.

Real-world Connection: This explains why honey (lots of hydrogen bonding) pours slowly whilst petrol (only weak LDF) flows easily.

Solubility and Mixing Behaviour

The golden rule for solubility is "like dissolves like" - substances dissolve best in solvents with similar intermolecular forces. Sugar dissolves in water (both polar), whilst wax dissolves in pentane $both non-polar$.

Ionic compounds are often soluble in highly polar solvents like water because polar molecules are strongly attracted to both positive and negative ions. However, they're insoluble in non-polar solvents.

Miscible substances mix completely with no visible boundary (like ethanol and water), whilst immiscible substances form separate layers (like oil and water). You can test solvent polarity using a charged rod - polar liquids bend toward the charge, non-polar liquids don't respond.

This explains why you can't make salad dressing without shaking - oil and vinegar have completely different intermolecular forces!

Lab Insight: The charged rod test is a simple way to identify polar vs non-polar liquids in practical work.

Water's Special Properties

Water's unique properties all stem from its extensive hydrogen bonding network. Its high boiling point (100°C) seems unusual for such a small molecule, but hydrogen bonding requires extra energy to break.

Ice floats because it's less dense than liquid water - hydrogen bonding creates an expanded, ordered structure when water freezes. This is actually quite rare; most substances are denser as solids.

Water also has very high surface tension, allowing insects to walk on its surface. Again, this is down to those strong hydrogen bonds holding the surface molecules together like a flexible skin.

Environmental Impact: Ice floating on water insulates aquatic life below - if ice sank, lakes would freeze solid from bottom to top!