Electrical Conductivity: Metals are excellent conductors because the delocalised electrons are free to move through the structure when a potential difference is applied.
Malleability and Ductility: Metals can be hammered into sheets or drawn into wires because the layers of positive ions can slide over each other without breaking the metallic bond, as the 'sea' of electrons adjusts to the new positions.
High Melting and Boiling Points: The strong electrostatic attraction between the cations and the delocalised electrons requires significant thermal energy to overcome.
Luster: The presence of free electrons on the surface allows metals to reflect light, giving them a shiny appearance.
Oxide Formation: Metals typically react with oxygen to form basic oxides (e.g., Magnesium Oxide), though some can form amphoteric oxides that react with both acids and bases.
Reaction with Acids: Many metals react with dilute acids to produce a salt and hydrogen gas, a process driven by the metal's tendency to lose electrons.
Group Trends: Reactivity increases down a group because the outer electrons are further from the nucleus and shielded by more inner shells, making the attraction to the nucleus weaker and electron loss easier.
| Feature | Metals | Non-Metals |
|---|---|---|
| Ion Formation | Form positive cations by losing electrons | Form negative anions by gaining electrons |
| Bonding Type | Metallic bonding (lattice of ions/electrons) | Covalent bonding (shared electron pairs) |
| Conductivity | High electrical and thermal conductivity | Poor conductors (insulators) |
| Oxide Nature | Basic or amphoteric oxides | Acidic or neutral oxides |
| State at RT | Mostly solids (except Mercury) | Solids, liquids, or gases |
Explain, Don't Just State: When asked why metals conduct electricity, always mention 'delocalised electrons' and their 'ability to move through the structure'.
Link Structure to Property: For malleability questions, use the phrase 'layers of ions can slide over each other' while emphasizing that the metallic bond remains intact.
Reactivity Justification: When comparing reactivity down a group, always reference the 'distance from the nucleus' and 'atomic radius' to explain why electrons are lost more easily.
Check the Ion Charge: Remember that the number of protons remains constant; an ion becomes positive only because it has fewer negative electrons than positive protons.
