Topic 2 of 10

Bonding, Structure & Properties of Matter

Discover how atoms link together through ionic, covalent, and metallic bonds — and how these structures determine the properties of every substance around you.

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Chemical Bonds

A chemical bond is a force of attraction that holds atoms together. Atoms form bonds to achieve a more stable electron arrangement — usually a full outer shell of electrons (the 'octet rule').

The Three Types of Chemical Bonds

  • Ionic Bonding: Forms between a metal and a non-metal. The metal atom loses electrons to become a positive ion; the non-metal gains electrons to become a negative ion. The bond is the strong electrostatic attraction between these oppositely charged ions.
  • Covalent Bonding: Forms between two non-metal atoms. Atoms share one or more pairs of electrons. This can form simple molecules (H₂O, CH₄) or giant covalent structures (diamond).
  • Metallic Bonding: Found in metals and alloys. Metal atoms lose outer electrons, creating a lattice of positive ions surrounded by a "sea" of delocalised electrons.
When asked to define a chemical bond, state that it's a force holding atoms together, allowing them to achieve a stable, full outer shell of electrons.

Ionic Bonding

Ionic bonding happens between a metal and a non-metal. It involves the transfer of electrons from the metal atom to the non-metal atom.

How Ionic Bonds Form

  1. Electron Transfer: The metal atom loses outer electrons → positive ion (cation). The non-metal atom gains elections → negative ion (anion).
  2. Electrostatic Attraction: The oppositely charged ions are strongly attracted to each other. This powerful force is the ionic bond.

Formation of Sodium Chloride (NaCl)

Sodium (2.8.1) loses 1 electron → Na⁺ (2.8)

Chlorine (2.8.7) gains 1 electron → Cl⁻ (2.8.8)

The Na⁺ and Cl⁻ ions are held together by strong electrostatic attraction.

The Giant Ionic Lattice

Ionic compounds form a giant ionic lattice — a regular, repeating 3D arrangement of cations and anions. Each ion is strongly attracted to all surrounding ions of the opposite charge.

A complete answer on ionic bonding must describe both the transfer of electrons AND the resulting electrostatic attraction between oppositely charged ions.

Properties of Ionic Compounds

High Melting & Boiling Points

Ionic compounds have very high melting and boiling points (e.g., NaCl melts at 801°C). A large amount of energy is needed to overcome the strong electrostatic forces between the ions.

Electrical Conductivity

Whether an ionic compound conducts depends on its state:

  • Solid: Does not conduct — ions are held in fixed positions and cannot move.
  • Molten/Dissolved: Conducts electricity — ions are free to move and carry charge.

Solubility

Many ionic compounds dissolve in water. Water molecules surround the individual ions, breaking down the lattice and allowing the ions to move freely.

When answering about ionic compound properties, always link to the structure. Use "strong electrostatic forces" for melting points, and "ions are free to move" for conductivity.

Covalent Bonding

Covalent bonding occurs when two non-metal atoms share pairs of electrons to achieve a full outer shell. Each atom contributes one or more electrons to form a shared pair.

Simple Molecular Substances

Made of individual, discrete molecules. Within each molecule, atoms are held by very strong covalent bonds. However, forces between molecules (intermolecular forces) are very weak.

When discussing simple molecules, always distinguish between the strong covalent bonds within molecules and the weak intermolecular forces between them. It's the weak forces that are broken during melting or boiling.

Properties of Simple Molecules

  • Low melting/boiling points: Only a small amount of energy is needed to overcome the weak intermolecular forces. Many are gases or liquids at room temperature.
  • Poor electrical conductors: The molecules are neutral with no free-moving electrons or ions to carry charge.

Examples: Hydrogen (H₂), Chlorine (Cl₂), Water (H₂O), Methane (CH₄).

Metallic Bonding

In metals, atoms lose outer electrons to form a regular lattice of positive ions surrounded by a "sea" of delocalised electrons. The metallic bond is the strong electrostatic attraction between the positive ions and the delocalised electrons.

How Metallic Bonding Explains Metal Properties

  • Good electrical conductors: Delocalised electrons can move throughout the structure and carry charge.
  • Good thermal conductors: Delocalised electrons carry kinetic energy from hotter to cooler regions.
  • Malleable & ductile: Layers of ions can slide over each other without breaking the metallic bond.
  • High melting points: Strong metallic bonds require a lot of energy to break.

Alloys

Alloys are mixtures of a metal with other elements. Different-sized atoms distort the layers, so they cannot slide as easily. This makes alloys harder and often stronger than pure metals.

Most everyday metals are actually alloys. Pure metals are often too soft because their identical atoms form neat layers that slide easily.
When explaining why alloys are harder than pure metals, you must mention that the different-sized atoms distort the regular layers, preventing them from sliding.

States of Matter

Solids

Particles packed tightly in a fixed, regular pattern (lattice). Strong forces hold them in place. Particles vibrate on the spot. Fixed shape and volume; cannot be compressed.

Liquids

Particles close together but arranged randomly. Weaker forces than solids. Particles can move and slide past each other. Fixed volume but take the shape of their container.

Gases

Particles very far apart, arranged randomly. Very weak forces. Particles move quickly and randomly. No fixed shape or volume; easily compressed.

State Symbols

In equations, state symbols show the physical state: (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).

AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Polymers

Polymers are very large molecules (macromolecules) built from many smaller, repeating units called monomers. The process of joining them is called polymerisation.

  • Within chains: Very strong covalent bonds.
  • Between chains: Intermolecular forces — stronger than in simple molecules (so polymers are solid at room temperature), but weaker than covalent/ionic bonds in giant structures.

Examples: Poly(ethene) – plastic bags; Poly(propene) – ropes and crates.

Do not confuse the strong covalent bonds within polymer chains with the weaker intermolecular forces between them. It is the intermolecular forces that are overcome when a polymer melts.

Giant Covalent Structures

Huge numbers of non-metal atoms joined by a continuous network of strong covalent bonds. No separate molecules and no weak intermolecular forces.

Diamond

Each carbon atom forms four strong covalent bonds in a rigid tetrahedral arrangement. Extremely hard, very high melting point, does not conduct electricity (no delocalised electrons).

Graphite

Each carbon forms three covalent bonds, creating flat layers of hexagonal rings. Layers held by weak forces — can slide (soft and slippery). One delocalised electron per carbon allows electrical conductivity along the layers.

Silicon Dioxide (SiO₂)

Similar structure to diamond. Each silicon bonded to four oxygens, each oxygen bonded to two silicons. Very hard, high melting point, does not conduct electricity.

When explaining properties of giant covalent structures, always link to the specific structure. E.g., "Diamond is hard because each carbon is held in a rigid lattice by four strong covalent bonds."

Graphene & Fullerenes

Graphene

A single, one-atom-thick layer of graphite. Each carbon forms three bonds, leaving one delocalised electron per atom. Extremely strong, lightweight, and an excellent electrical conductor. Used in electronics, touchscreens, and composites.

Buckminsterfullerene (C₆₀)

60 carbon atoms arranged in a hollow sphere (like a football). Forms a simple molecular structure with weak intermolecular forces between C₆₀ molecules. Uses: lubricants, drug delivery, catalysts.

Carbon Nanotubes

A sheet of graphene rolled into a seamless cylinder. Very high tensile strength, excellent conductors of electricity and heat, huge surface area to volume ratio. Used in composites (tennis rackets, bike frames), nanoelectronics, and as catalysts.

Nanoparticles

Nanoparticles have a diameter from 1 nm to 100 nm. Their extremely high surface area to volume ratio gives them unique properties different from the bulk material.

Uses of Nanoparticles

  • Sun creams: TiO₂ and ZnO nanoparticles block UV but appear transparent on skin.
  • Antibacterial coatings: Silver nanoparticles release ions toxic to bacteria.
  • Catalysts: Huge surface area makes industrial reactions more efficient.
  • Drug delivery: Can carry drugs to specific target cells in the body.
  • Self-cleaning surfaces: Break down dirt when exposed to sunlight.

Risks & Concerns

  • May enter the body through skin, inhalation, or ingestion — long-term health effects not fully understood.
  • Could accumulate in the environment and harm ecosystems.
When explaining nanoparticle uses, always link your answer to their high surface area to volume ratio. For example: "Silver nanoparticles are effective antibacterial agents because their high SA:V ratio allows them to release a steady stream of silver ions."