1.1Unit 2 · Bonding & Electrolysis
Ionic bonding — Edexcel IGCSE Chemistry Modular - UNIT 2 Revision Notes & Topical Questions
10 min read 4 parts Lesson 1 of 4 1 quiz
Lesson summary
Explain how ions form and predict their charges.Group 1 → 1+, Group 2 → 2+, Group 3 → 3+; Group 5 → 3−, Group 6 → 2−, Group 7 → 1−.Ionic compounds have zero overall chargetotal positive charge = total negative charge.Formula of aluminium oxide2 × Al3+ (+6) and 3 × O2− (−6) = Al2O…
- Formation of ions
- Common ions
- Finding formulae
- Dot and cross diagrams
- Properties
- Ions and subatomic particles
Section 1 of 4Ionic bonding
Formation of ions
atoms lose or gain electrons to reach the stable electron arrangement of a noble gas.- Metals, with 1–3 outer electrons, lose them to form positive ions; non-metals, with 5–7 outer electrons, gain electrons to form negative ions.
- Loss of electrons is oxidation and gain of electrons is reduction, e.g. Mg → Mg2+ + 2e− and Cl2 + 2e− → 2Cl−.
Common ions
learn charges from group numbers, plus transition metal ions shown by Roman numerals (iron(II) Fe2+, iron(III) Fe3+, copper(II) Cu2+), silver Ag+, zinc Zn2+,- and compound ions: ammonium NH4+, hydroxide OH−, nitrate NO3−, carbonate CO32−, sulfate SO42−.
Finding formulae
choose numbers of each ion so charges cancel.- Calcium chloride: Ca2+ needs two Cl− → CaCl2.
- Ammonium sulfate: two NH4+ to one SO42− → (NH4)2SO4.
- Endings: -ide = element only, -ate = contains oxygen.
Dot and cross diagrams
show electron transfer: use dots for one atom's electrons and crosses for the other's.- Final ions are drawn in square brackets with the charge outside; often only outer shells are shown.
- For MgO, two electrons move from Mg to O.
Structure
ionic compounds form a giant lattice, where each ion is surrounded by ions of the opposite charge, and attractions act in all directions throughout the structure.Properties
high melting and boiling points because lots of energy is needed to overcome the many strong electrostatic attractions.- They conduct electricity only when molten or dissolved, because then the ions can move.
- Solids are hard but brittle.
Ions and subatomic particles
a cation has more protons than electrons; an anion has more electrons than protons. The size of the charge equals the number of electrons lost or gained, e.g. an aluminium atom losing 3 electrons forms Al3+.Predicting charges
atoms with 1–3 outer electrons usually lose them; atoms with 5–7 outer electrons usually gain enough to reach 8. Group 4 elements rarely form simple ions.More common ions
also learn lead(II) Pb2+ and hydrogen H+. Like ammonium NH4+, H+ is a positive ion that does not come from a metal.Swap-and-drop
to find a formula quickly, use the size of each ion's charge as the number of the other ion, then simplify. Al3+ and O2− give Al2O3; Mg2+ and O2− give Mg2O2, which simplifies to MgO.Dot and cross: other groups
in lithium nitride, three lithium atoms (2.1) each give one electron to one nitrogen atom (2.5). This forms three Li+ ions (2) and one N3− ion (2.8), so the formula is Li3N.Conduction: ions, not electrons
molten or dissolved ionic compounds conduct because their ions move and carry charge. Saying that electrons move through the compound is a common mistake that loses marks.
Section 2 of 4Process & sequence
- Write the electronic configuration of each atom (e.g. Na 2.8.1, Cl 2.8.7).
- The metal atom transfers its outer electron(s) to the non-metal atom.
- The metal becomes a positive ion with a full outer shell (Na+ 2.8).
- The non-metal becomes a negative ion with a full outer shell (Cl− 2.8.8).
- Draw each ion in square brackets with its charge at the top right.
- The oppositely charged ions attract, forming an ionic bond in a giant lattice.
Forming sodium chloride — write these steps in this order; each step is a marking point.
- Na 2.8.1 and Cl 2.8.7
- Na transfers one electron
- Na+ 2.8 formed
- Cl− 2.8.8 formed
- Opposite charges attract
- Giant ionic lattice
Section 3 of 4Compare
| Ion | Formula | Ion | Formula |
|---|---|---|---|
| Sodium | Na+ | Chloride | Cl− |
| Magnesium | Mg2+ | Oxide | O2− |
| Aluminium | Al3+ | Nitride | N3− |
| Ammonium | NH4+ | Hydroxide | OH− |
| Copper(II) | Cu2+ | Sulfate / Carbonate / Nitrate | SO42− / CO32− / NO3− |
| Iron(III) | Fe3+ | Sulfide | S2− |
| Property | Reason |
|---|---|
| High melting and boiling points | Much energy needed to overcome strong electrostatic forces in the lattice |
| Do not conduct when solid | Ions are fixed in position |
| Conduct when molten or dissolved | Ions are free to move and carry charge |
| Often soluble in water | Water molecules attract and separate the ions |
Section 4 of 4Worked examples
Work out the formula of iron(III) sulfate.
- Fe3+ and SO42−. Lowest common multiple of 3 and 2 is 6: 2 Fe3+ = +6, 3 SO42− = −6. Formula: Fe2(SO4)3.
Explain why magnesium oxide has a much higher melting point than sodium chloride.
- Both are giant ionic lattices, but MgO contains Mg2+ and O2− ions with double the charges of Na+ and Cl−.
- The electrostatic attractions are stronger, so more energy is needed to overcome them.
- Both are giant ionic lattices, but MgO contains Mg2+ and O2− ions with double the charges of Na+ and Cl−.
Marked practiceQuizzes for this lesson
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