State the meaning of the term ‘ionic bond’.
Electrostatic attraction between oppositely charged ions formed by electron transfer.
Describe the bonding in platinum.
Metallic bonding: positive metal ions in a lattice held by delocalised electrons.
Describe the bonding in chromium and explain its high melting temperature.
Strong metallic bonding with many delocalised electrons
requires large energy to overcome.
Using data, explain how changes in the cation affect ionic bond strength.
Smaller and more highly charged cations produce stronger electrostatic attraction.
State all conditions under which MgBr₂ conducts electricity.
When molten or dissolved in water so ions are free to move.
Explain conductivity of solid vs aqueous KBr.
Solid: ions fixed in lattice so no conduction.
Aqueous: ions free to move and carry charge.
Complete electron configuration of Mg.
1s² 2s² 2p⁶ 3s².
Bonding in magnesium.
Metallic bonding: positive ions attracted to sea of delocalised electrons.
Correct equation for Na + F₂.
2Na(s) + F₂(g) → 2NaF(s).
Explain why Na⁺ and F⁻ are isoelectronic.
Both have 10 electrons after electron transfer, same electron configuration as Ne.
Trend in ionic radius from N³⁻ to Al³⁺.
Radius decreases as nuclear charge increases while electrons stay the same.
Formula of ion isoelectronic with Mg²⁺.
Na⁺ or F⁻ or O²⁻ (e.g., Na⁺).
Explain difference in melting temps of MgO and KBr.
MgO has higher charge ions and smaller ionic radii → much stronger ionic attraction.
Two physical properties of metals explained by bonding.
Good electrical/thermal conductivity (delocalised electrons).
Malleable/ductile (layers slide).
Equation for thermal decomposition of CaCO₃.
CaCO₃(s) → CaO(s) + CO₂(g).
Types of bonds in CaCO₃.
Ionic bonds and covalent bonds within carbonate ion.
Reason high decomposition temperature required.
Strong ionic lattice and strong covalent bonding require high energy to break.
