Ionic, covalent, and metallic bonds — plus crystal structures, diamond, graphite, and the valency method for writing formulae.
Section 1 — Why Do Atoms Bond?
Atoms form chemical bonds to achieve a full outer electron shell — the same stable configuration as the noble gases (Group 0). This is called the noble gas configuration. Atoms achieve stability by losing, gaining, or sharing valence electrons.
Metal + Non-metal
Electrons are TRANSFERRED. Metal loses electrons → cation (+). Non-metal gains electrons → anion (−). Electrostatic attraction between ions.
Examples: NaCl, MgO, CaCl₂, Al₂O₃
Non-metal + Non-metal
Electrons are SHARED in pairs. Each shared pair = 1 covalent bond. Single, double, or triple bonds possible.
Examples: H₂O, CO₂, CH₄, N₂, HCl
Metal + Metal
Valence electrons become DELOCALISED — free to roam. 'Sea' of mobile electrons surrounds positive cations.
Examples: Cu, Fe, Na, Mg, Al, Au
Section 2 — Chemical Formulae & Valency
The valency of an element tells you how many bonds it forms. Use the valency method to write formulae: the number of atoms of each element is swapped from the other element's valency.
| Compound | Element 1 (valency) | Element 2 (valency) | Working | Formula |
|---|---|---|---|---|
| Magnesium nitride | Mg (2) | N (3) | Mg₃ × 2 = 6, N₂ × 3 = 6 | Mg₃N₂ |
| Aluminium oxide | Al (3) | O (2) | Al₂ × 3 = 6, O₃ × 2 = 6 | Al₂O₃ |
| Iron(III) bromide | Fe (3) | Br (1) | Fe × 3 = 3, Br₃ × 1 = 3 | FeBr₃ |
| Carbon dioxide | C (4) | O (2) | C₂O₄ → simplify → CO₂ | CO₂ |
| Calcium chloride | Ca (2) | Cl (1) | Ca × 2 = 2, Cl₂ × 1 = 2 | CaCl₂ |
Section 3 — Ionic Bonding
Ionic bonding involves the transfer of valence electrons from a metal to a non-metal. The metal loses electrons and becomes a positive cation; the non-metal gains electrons and becomes a negative anion. The strong electrostatic force between oppositely-charged ions IS the ionic bond.
Worked Examples of Ionic Bond Formation
| Compound | Metal Ion | Non-metal Ion | Formula |
|---|---|---|---|
| Sodium chloride | Na loses 1e⁻ → Na⁺ | Cl gains 1e⁻ → Cl⁻ | NaCl |
| Magnesium fluoride | Mg loses 2e⁻ → Mg²⁺ | F gains 1e⁻ → F⁻ (need 2F) | MgF₂ |
| Aluminium oxide | Al loses 3e⁻ → Al³⁺ (need 2 Al) | O gains 2e⁻ → O²⁻ (need 3 O) | Al₂O₃ |
| Potassium nitride | K loses 1e⁻ → K⁺ (need 3 K) | N gains 3e⁻ → N³⁻ | K₃N |
Section 4 — Covalent Bonding
Covalent bonding occurs between two or more non-metal atoms. Instead of transferring electrons, atoms share pairs of valence electrons. Each shared pair forms one covalent bond. The resulting particles are called molecules.
| Bond Type | Shared Pairs | Bond Symbol | Examples |
|---|---|---|---|
| Single covalent | 1 pair (2 electrons) | — | H₂, HCl, H₂O, CH₄, NH₃, Cl₂ |
| Double covalent | 2 pairs (4 electrons) | = | O₂, CO₂, C₂H₄ (ethene) |
| Triple covalent | 3 pairs (6 electrons) | ≡ | N₂, C₂H₂ (ethyne) |
Key Covalent Molecules — Quick Reference
| Molecule | Bond Type | Valence Electrons Used | Structural Formula |
|---|---|---|---|
| Cl₂ | Single | Cl(7) needs 1 more — shares 1 pair with Cl | Cl—Cl |
| H₂O | 2× Single | O(6) needs 2 — two H atoms share 1 pair each | H—O—H |
| CH₄ | 4× Single | C(4) needs 4 — four H atoms share 1 pair each | H₃C—H (4 bonds) |
| O₂ | Double | O(6) needs 2 — two O atoms share 2 pairs | O=O |
| CO₂ | 2× Double | C(4) needs 4 — double bond with each O | O=C=O |
| N₂ | Triple | N(5) needs 3 — two N atoms share 3 pairs | N≡N |
Section 5 — Metallic Bonding
In metals, valence electrons become delocalised — free to move throughout a 'sea' of electrons surrounding positive metal cations. The strong electrostatic attraction between the positive cations and the electron sea IS the metallic bond.
| Property | Explanation Using Metallic Bonding |
|---|---|
| High melting/boiling points | Strong electrostatic forces between cations and delocalised electrons require large amounts of energy to break. E.g. iron melts at 1538°C. |
| Good electrical conductors | Delocalised electrons are free to move. When voltage is applied, they flow, carrying electric current (e.g. copper wires). |
| Good heat conductors | Mobile electrons carry kinetic energy (heat) rapidly through the metal lattice. |
| Malleable (can be hammered flat) | All atoms are the same type and size. Layers can slide past each other without breaking metallic bonds. |
| Ductile (drawn into wires) | Layers of atoms slide along the direction of applied force without bonds breaking. |
| Shiny appearance | The delocalised electron sea reflects light from the surface. |
Section 6 — Four Crystal Types
| Crystal Type | Particles | Forces | Melting Point | Conductivity | Examples |
|---|---|---|---|---|---|
| Ionic | Cations & anions | Strong ionic bonds | HIGH (NaCl = 801°C) | Only when molten or dissolved in water | NaCl, KBr, MgO |
| Simple Molecular | Small molecules | WEAK intermolecular forces (strong covalent bonds within) | LOW (Ice = 0°C) | Never | Ice, I₂, CO₂, glucose |
| Giant Molecular | Non-metal atoms | Strong covalent bonds throughout 3D lattice | VERY HIGH (Diamond = 3550°C) | Usually none (except graphite) | Diamond, graphite, SiO₂ |
| Metallic | Cations + delocalised electrons | Strong metallic bonds | HIGH (varies) | Excellent in all states | Cu, Fe, Al, Na |
Section 7 — Diamond & Graphite: Allotropes of Carbon
| Property | Diamond | Graphite |
|---|---|---|
| Structure | Each C atom bonded covalently to 4 others in a tetrahedral giant molecular lattice | Each C atom bonded to 3 others in hexagonal rings forming flat layers. Weak van der Waals forces between layers |
| Hardness | EXTREMELY HARD — hardest known natural substance | SOFT & flaky — layers slide easily |
| Electrical conductivity | Does NOT conduct — all 4 valence electrons used in bonds, none free to move | CONDUCTS — 4th valence electron on each C is delocalised and free to move between layers |
| Lubricating power | None — atoms rigidly bonded in all directions | EXCELLENT lubricant — layers slide over each other |
| Melting point | ~3550°C — enormous energy to break all covalent bonds | ~3600°C — strong covalent bonds within layers |
| Real-world uses | Jewellery, diamond drill bits, glass cutters, surgical scalpels | Pencil 'lead', electrodes in electrolysis, crucibles, solid lubricant |
Section 8 — Resources & Simulations
Section 9 — CSEC Practice Questions
(A) NaCl (B) MgO (C) NH₃ (D) KF
1 NaCl = Na (metal) + Cl (non-metal) → ionic. MgO = Mg (metal) + O (non-metal) → ionic. KF = K (metal) + F (non-metal) → ionic.
2 NH₃ = N (non-metal) + H (non-metal) → non-metals only → covalent bonding.
1 Solid NaCl: Na⁺ and Cl⁻ ions are held in fixed positions in the ionic crystal lattice by strong electrostatic forces of attraction. They cannot move, so they cannot carry an electric charge — no conduction.
2 NaCl in water: Water molecules attract and separate the ions. The Na⁺ and Cl⁻ ions become free to move throughout the solution. Mobile charged particles (ions) can carry charge from one electrode to the other — conduction occurs.
1 Diamond: Each carbon atom is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement. A giant covalent 3D lattice forms throughout the crystal. All 4 valence electrons of each carbon are used in covalent bonds — there are no free electrons.
2 Graphite: Each carbon atom is covalently bonded to only 3 other carbon atoms, forming hexagonal rings in flat layers. The 4th valence electron from each carbon atom is delocalised — free to move between layers throughout the structure.
3 Conductivity: Diamond has no free (delocalised) electrons, so charge cannot flow — it does not conduct. Graphite has delocalised electrons that can move freely between layers when a voltage is applied, carrying electric charge — it conducts.
a Calcium nitride: Ca²⁺ and N³⁻. To balance: 2 Ca (losing 2e⁻ each = 4e⁻ total), 3 N (gaining... wait — 3 Ca × 2 = 6 and 2 N × 3 = 6). Formula: Ca₃N₂
b Iron(III) oxide: Fe³⁺ and O²⁻. 2 Fe × 3 = 6, 3 O × 2 = 6. Formula: Fe₂O₃
c Aluminium sulfide: Al³⁺ and S²⁻. 2 Al × 3 = 6, 3 S × 2 = 6. Formula: Al₂S₃
1 High melting point → strong forces of attraction between particles.
2 Does NOT conduct as solid but DOES conduct when dissolved → ions present, but only free to move in solution (not when fixed in lattice). This is the hallmark of ionic bonding.
3 Hard but brittle → ionic crystal lattice; layers shift under pressure, like charges come adjacent and repel, shattering the lattice.