Bonding, Structure & Properties (Ionic, Covalent, Metallic)

Key Definitions

Ionic bond: Electrostatic attraction between oppositely charged ions, formed by electron transfer.
Covalent bond: A shared pair of electrons between two non-metal atoms.
Metallic bond: Electrostatic attraction between positive metal ions and delocalised electrons.
Polymer: A very large molecule made from many repeating smaller units (monomers).
Alloy: A mixture of a metal with one or more other elements (usually metals).
Intermolecular forces: Weak forces of attraction between molecules (not the bonds within them).

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

Electron Transfer: The metal atom loses outer electrons → positive ion (cation). The non-metal atom gains electrons → negative ion (anion).
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.

Formation of Magnesium Oxide (MgO)

Magnesium (2.8.2) loses 2 electrons → Mg²⁺ (2.8)

Oxygen (2.6) gains 2 electrons → O²⁻ (2.8)

Two electrons are transferred. The Mg²⁺ and O²⁻ ions have a strong electrostatic attraction.

Formation of Magnesium Chloride (MgCl₂)

Magnesium (2.8.2) loses 2 electrons - one to each of two chlorine atoms.

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

This gives the formula MgCl₂ - one Mg²⁺ ion for every two Cl⁻ ions.

Dot-and-Cross Diagrams for Ionic Bonds

In dot-and-cross diagrams, dots (•) represent electrons from one atom and crosses (×) represent electrons from the other. For ionic bonds, the transferred electrons are shown in the outer shell of the ion that gained them. Square brackets and charges are used around each ion.

When drawing dot-and-cross diagrams for ionic compounds, remember to show: the electron transfer, square brackets around each ion, and the charge on each ion (e.g. [Na]⁺ and [Cl]⁻).

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.

The transfer of an electron from a sodium atom to a chlorine atom forms oppositely charged ions, which arrange into a massive 3D giant ionic lattice.

A complete answer on ionic bonding must describe both the transfer of electrons AND the resulting electrostatic attraction between oppositely charged ions. For MgO, state that two electrons are transferred.

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.

Common Covalent Molecules

Single Bonds (one shared pair)

H₂ (hydrogen) - each H shares 1 electron
HCl (hydrogen chloride) - H and Cl each share 1 electron
H₂O (water) - oxygen shares 1 electron with each of 2 hydrogens
CH₄ (methane) - carbon shares 1 electron with each of 4 hydrogens

Double Bonds (two shared pairs)

O₂ (oxygen) - each oxygen shares 2 electrons (O=O)

Triple Bonds (three shared pairs)

N₂ (nitrogen) - each nitrogen shares 3 electrons (N≡N)

Dot-and-Cross Diagrams for Covalent Bonds

In covalent dot-and-cross diagrams, show the outer shell electrons only. Dots represent electrons from one atom and crosses from the other. The shared pair sits in the overlapping region between the two atoms.

When drawing covalent dot-and-cross diagrams, you only need to show the outer shell electrons. For double bonds, show two shared pairs, and for triple bonds, three shared pairs. Do NOT use square brackets (those are only for ionic diagrams).

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.

Simple molecules like water (H₂O) and methane (CH₄) are formed by atoms sharing electron pairs. Strong covalent bonds hold the atoms together within each molecule.

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 - the covalent bonds are NOT broken.

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).

Bonding & Structure Summary

The type of bonding and structure determines the physical properties of a substance. Use this table to compare all four types:

Property	Ionic	Simple Molecular	Giant Covalent	Metallic
Bonding	Ionic (transfer)	Covalent (sharing)	Covalent (sharing)	Metallic
Structure	Giant lattice of ions	Small molecules	Giant lattice of atoms	Giant lattice of ions + e⁻
Melting point	High	Low	Very high	High (varies)
Conductivity (solid)	No	No	No (except graphite)	Yes
Conductivity (liquid)	Yes	No	N/A	Yes
Examples	NaCl, MgO	H₂O, CH₄, CO₂	Diamond, graphite, SiO₂	Iron, copper, steel

Summary of bonding types, structures and key properties.

Predicting Properties: Why does NaCl conduct electricity when molten but not as a solid?

Solid NaCl: The Na⁺ and Cl⁻ ions are held in fixed positions in a giant ionic lattice. They cannot move, so they cannot carry charge. → Does not conduct.

Molten NaCl: When heated above 801°C, the lattice breaks down. The ions become free to move and carry charge through the liquid. → Conducts electricity.

Questions asking you to "explain the properties" of a substance always need three things: (1) name the type of bonding/structure, (2) describe the forces involved, (3) link to the specific property. The comparison table above is a great revision tool!

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.

Left: Pure metals form layers of identical positive ions that can slide easily. Right: Alloys contain mixed atoms of different sizes, distorting the regular layers and making them harder.

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.

Common Alloys

Alloy	Composition	Use
Steel	Iron + carbon	Construction, tools, vehicles
Stainless steel	Iron + chromium + nickel	Cutlery, surgical instruments
Bronze	Copper + tin	Statues, medals, ship propellers
Brass	Copper + zinc	Musical instruments, door handles
Gold alloys	Gold + copper/silver	Jewellery (harder than pure gold)

Common alloys, their compositions and uses.

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

Particle models of the three states of matter. Solids are rigidly packed, liquids are loosely packed and free to flow, and gases are widely spaced and highly energetic.

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.

Changes of State

When a substance changes state, the particles gain or lose energy. The chemical bonds are not broken - only the forces between particles change.

Melting: Solid → Liquid (particles gain energy, overcome some forces, start to move)
Boiling/Evaporating: Liquid → Gas (particles gain enough energy to overcome all intermolecular forces)
Freezing: Liquid → Solid (particles lose energy, forces pull them into fixed positions)
Condensing: Gas → Liquid (particles lose energy, forces bring them closer)
Sublimation: Solid → Gas directly (e.g. Dry ice, iodine)

During a change of state, the mass is conserved - no atoms are created or destroyed, they simply rearrange. The temperature stays constant during a change of state as energy is used to break or form intermolecular forces.

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)

Remember: during changes of state, the substance does not become a new substance - it is a physical change. The particles stay the same; only the arrangement and energy change.

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.

Left: Diamond's 3D tetrahedral structure. Right: Graphite's layered structure with weak intermolecular forces between layers.

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 continuous 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 Chemistry Only

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.

Why Surface Area to Volume Ratio Matters

As particles get smaller, their SA:V ratio increases dramatically. This means more of the material is exposed on the surface, making it more reactive and giving it different properties.

Comparing SA:V ratios

A 2 cm cube: SA = 6 × (2 × 2) = 24 cm².
Volume = 2 × 2 × 2 = 8 cm³.
SA:V = 24 ÷ 8 = 3:1

A 1 cm cube: SA = 6 × (1 × 1) = 6 cm².
Volume = 1 cm³.
SA:V = 6 ÷ 1 = 6:1

Halving the size doubled the SA:V ratio. At the nanoscale, this ratio becomes enormous.

Size Comparison

Nanoparticles: 1–100 nm (1 nm = 1 × 10⁻⁹ m)
Fine particles (dust): 100–2500 nm
Coarse particles (sand grains): 2500–10,000 nm

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.
More research is needed before nanoparticles can be considered completely safe.

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."

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