Ionic Bonding

Ever wondered why salt dissolves in water but doesn't conduct electricity until it's melted? Ionic bonding happens when metals give away electrons to non-metals, creating oppositely charged particles that attract each other.

Take sodium chloride - sodium loses an electron to become Na⁺, whilst chlorine gains one to become Cl⁻. These oppositely charged ions stick together through strong electrostatic forces, forming a giant lattice structure that's hard with high melting points.

Here's the key bit for your exams: ionic compounds only conduct electricity when molten or dissolved because the ions need to be free to move and carry the charge.

Quick Tip: Remember metals lose electrons (become positive), non-metals gain electrons (become negative) - opposites attract!

Covalent Bonding - Giant Structures

Covalent bonding occurs when non-metals share electrons, but the results can be dramatically different depending on the structure formed.

Diamond showcases the ultimate covalent structure - each carbon forms four strong bonds in a rigid 3D network. This makes it incredibly hard with high melting points, but it can't conduct electricity because all electrons are locked in bonds.

Graphite tells a different story. Each carbon bonds to only three others, leaving the fourth electron delocalised. These free electrons make graphite conduct electricity, whilst weak forces between layers make it slippery - perfect for pencil lead and lubricants.

Graphene, essentially one layer of graphite, conducts electricity even better and has exciting applications in electronics and composites.

Exam Focus: Diamond = hard, doesn't conduct. Graphite = soft layers, conducts electricity due to delocalised electrons.

Small Molecules, Polymers & Fullerenes

Small covalent molecules like water have low melting points because weak forces exist between molecules, not within them. Breaking these intermolecular forces requires little energy, explaining why water boils at just 100°C.

Polymers are long chains of repeating units with strong covalent bonds within the chain and significant forces between chains. When writing molecular formulas, put the repeating unit in brackets with 'n' outside, like (C₂H₄)ₙ.

Fullerenes are fascinating carbon structures with hollow centres and hexagonal rings. Buckminsterfullerene (C₆₀) was the first discovered, whilst carbon nanotubes are cylindrical versions with incredible strength-to-weight ratios.

These materials revolutionise technology - from drug delivery systems to computer chips, their unique structures create extraordinary properties.

Memory Trick: Think of polymers as molecular chains - strong links (covalent bonds) but the chains can still move past each other.

Metallic Bonding & Nanoparticles

Metallic bonding creates a "sea of electrons" - positive metal ions surrounded by delocalised electrons that can move freely. This explains why metals conduct electricity and heat so well, plus why they're malleable rather than brittle.

Metal alloys mix different metals $or metals with non-metals$ to create stronger materials. Different-sized atoms disrupt the regular arrangement, preventing layers from sliding over each other easily - that's why steel is harder than pure iron.

Nanoparticles $1-100nm diameter$ have massive surface area to volume ratios, making them incredibly useful. They're revolutionising everything from medicine to electronics because their tiny size gives them unique properties.

Applications include drug delivery, computer chips, antibacterial deodorants, and surgical masks - their high surface area makes them excellent catalysts too.

Key Concept: Size matters! As particles get smaller, their surface area to volume ratio increases dramatically, creating new properties and applications.