Physical and chemical properties of metals explained by metallic bonding, reactions with oxygen, water, and dilute acids, thermal decomposition of metallic carbonates, nitrates, and hydroxides, the reactivity series, and displacement reactions.
The reactivity of a metal determines how it behaves chemically and what it can be used for. All of a metal's physical properties flow from one structural feature — metallic bonding — and all of its chemical behaviour can be predicted from its position in the reactivity series.
Most metals share a set of distinctive physical properties, all explained by metallic bonding (a lattice of positive metal ions surrounded by a sea of delocalised electrons):
| Property | Explanation |
|---|---|
| Good electrical conductors | Delocalised electrons carry charge freely |
| Good thermal conductors | Delocalised electrons transfer kinetic energy rapidly |
| Malleable | Layers of ions slide without breaking the metallic bond |
| Ductile | Ions can rearrange without bonds breaking |
| Lustrous (shiny) | Delocalised electrons reflect light |
| High melting and boiling points | Strong metallic bonding requires much energy to overcome |
| Usually dense | Closely packed cations |
Notable exceptions: mercury is a liquid metal at room temperature; sodium and potassium are soft, low-density, and low-melting because their metallic bonding is unusually weak.
Metals react with oxygen to form metal oxides. More reactive metals react vigorously; less reactive metals react slowly or not at all.
Gold does not react with oxygen.
Highly reactive metals react with cold water; moderately reactive metals react with steam; unreactive metals do not react with water at all.
Metals above hydrogen in the reactivity series displace hydrogen from dilute acids:
Copper, silver, and gold do not react with dilute acids.
Metals also form compounds that decompose on heating. Less reactive metals' compounds break down more easily, and the products follow a pattern that depends on position in the reactivity series.
Metal carbonates decompose on heating to give the metal oxide and CO₂. Less reactive metals' carbonates decompose more easily:
Sodium and potassium carbonates do not decompose at Bunsen burner temperatures.
| Metal | Products of heating nitrate |
|---|---|
| K, Na | Metal nitrite + oxygen: 2KNO₃ → 2KNO₂ + O₂ |
| Ca, Mg, Al, Zn, Fe, Pb, Cu | Metal oxide + NO₂ + oxygen: 2Cu(NO₃)₂ → 2CuO + 4NO₂ + O₂ |
| Ag, Au | Metal + NO₂ + oxygen: 2AgNO₃ → 2Ag + 2NO₂ + O₂ |
Most metal hydroxides decompose on heating to give the oxide and water:
NaOH and KOH are stable to heat and do not decompose at lab temperatures.
The patterns across all the reactions above fall into the same order. The reactivity series makes that order explicit.
The reactivity series ranks metals (and hydrogen) in order of how readily they lose electrons. Metals at the top are most reactive.
Using the series, you can predict:
The series also predicts reactions that haven't been observed yet. If a more reactive metal is placed in a solution of a less reactive metal's salt, it displaces it.
A more reactive metal displaces a less reactive metal from a solution of its salt:
Observations: blue solution fades; reddish-brown copper deposits on the zinc; zinc dissolves.
Net ionic equation: Zn + Cu²⁺ → Zn²⁺ + Cu (zinc is oxidised; copper ions are reduced)
Observations: blue solution turns pale green; copper deposited.
If the added metal is less reactive than the metal in solution, no reaction occurs. For example, copper placed in zinc sulfate solution produces no change.