Ions & Ionic Bonds
Section: Chemistry | Syllabus: Cambridge IGCSE Combined Science (0653)
Forming Ions
Atoms can gain or lose electrons to become charged particles called ions, achieving a stable, full outer shell like the nearest noble gas.
- Cation (positive ion): formed when an atom loses electrons - mainly formed by metals (Groups I-III)
- Na (2.8.1) loses 1 electron → Na⁺ (2.8)
- Mg (2.8.2) loses 2 electrons → Mg²⁺ (2.8)
- Anion (negative ion): formed when an atom gains electrons - mainly formed by non-metals (Groups V-VII)
- Cl (2.8.7) gains 1 electron → Cl⁻ (2.8.8)
- O (2.6) gains 2 electrons → O²⁻ (2.8)
For main groups, the charge can be predicted from the group number: Group I → 1+, Group II → 2+, Group III → 3+, Group V → 3-, Group VI → 2-, Group VII → 1-.
Ionic Bond Formation
An ionic bond is a strong electrostatic force of attraction between oppositely charged ions. It forms between a metal and a non-metal, when electrons are transferred (not shared) from the metal to the non-metal.
Worked Example: Sodium Chloride
- Question: use electronic configurations to explain the formation of the ionic bond in sodium chloride, NaCl. [3]
- Answer: sodium (2.8.1) transfers its 1 outer electron to chlorine (2.8.7) [1]; sodium becomes Na⁺ (2.8), with a stable configuration like neon [1]; chlorine becomes Cl⁻ (2.8.8), with a stable configuration like argon; the oppositely charged Na⁺ and Cl⁻ ions are then held together by strong electrostatic attraction [1]
In ionic bonding electrons are transferred, not shared - each ion is drawn separately in its own brackets with its charge shown outside
Formulas of Ionic Compounds
Ionic compounds are electrically neutral overall, so the total positive charge must equal the total negative charge. This fixes the ratio of ions in the formula.
| Cation | Anion | Balancing | Formula |
|---|---|---|---|
| Na⁺ | Cl⁻ | 1:1 | NaCl |
| Mg²⁺ | Cl⁻ | 1:2 (two Cl⁻ needed) | MgCl₂ |
| Al³⁺ | O²⁻ | 2:3 (2×3 = 3×2 = 6) | Al₂O₃ |
Common polyatomic ions to know: hydroxide OH⁻, nitrate NO₃⁻, carbonate CO₃²⁻, sulfate SO₄²⁻, and ammonium NH₄⁺ (the only common positive polyatomic ion). Brackets are used when more than one polyatomic ion is needed, e.g. Ca(OH)₂.
Naming: the metal ion keeps its element name; the non-metal ion's ending changes to -ide (chloride, oxide) - except polyatomic ions, which keep their own names (sulfate, carbonate, nitrate).
Common Mistakes
- ❌ Thinking ionic bonding involves sharing electrons, like covalent bonding
- ✓ ionic bonding involves the transfer of electrons from metal to non-metal, forming two separate charged ions - covalent bonding is what involves sharing
The Giant Ionic Lattice (Supplement)
Ionic compounds do not exist as separate pairs of ions or as individual "molecules". Instead, ions pack together into a giant ionic lattice: a regular, repeating three-dimensional arrangement of alternating positive and negative ions, held together by electrostatic attraction acting in every direction.
- The pattern of ions repeats regularly throughout the entire crystal
- Each ion is attracted to several neighbouring ions of opposite charge, not just one
- A single crystal contains a huge number of ions - the chemical formula (e.g. NaCl) shows only the simplest whole-number ratio of ions, not a separate unit
In solid NaCl, each Na⁺ ion is surrounded by 6 Cl⁻ ions, and each Cl⁻ ion is surrounded by 6 Na⁺ ions, arranged in a repeating cubic pattern.
Each ion in the lattice is surrounded by six oppositely charged neighbours, so the attraction acts throughout the whole structure, not just between one pair of ions
Properties of Ionic Compounds (Supplement)
- High melting and boiling points (e.g. NaCl: 801°C, MgO: 2852°C)
- The giant lattice is held together by strong electrostatic forces acting throughout the whole structure, so a large amount of energy is needed to separate the ions
- Ions with a greater charge (e.g. Mg²⁺ and O²⁻) attract each other more strongly than singly-charged ions (e.g. Na⁺ and Cl⁻), giving an even higher melting point
- Generally soluble in water
Electrical Conductivity
Whether an ionic compound conducts electricity depends entirely on whether its ions are free to move.
| State | Are ions free to move? | Conducts electricity? |
|---|---|---|
| Solid | No - fixed in the lattice | No |
| Molten | Yes - lattice has broken down | Yes |
| Aqueous (dissolved) | Yes - ions separated in solution | Yes |
The same ionic compound conducts when molten or dissolved but not when solid, because only then are the ions free to move and carry charge
Worked Example: Comparing Melting Points
- Question: explain why magnesium oxide (melting point 2852°C) has a much higher melting point than sodium chloride (melting point 801°C). [2]
- Answer: Mg²⁺ and O²⁻ ions have a greater charge than Na⁺ and Cl⁻ ions [1]; this produces stronger electrostatic forces of attraction between the ions in MgO, so more energy is needed to separate them, giving a higher melting point [1]
Common Mistakes
- ❌ Thinking solid ionic compounds don't conduct because they contain no charged particles
- ✓ solid ionic compounds do contain ions (charged particles) - they simply cannot move, since electrical conduction requires charged particles that are free to move
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