4.1 Reactivity of metals

4.1.1 Metal oxides

- Metals react with oxygen to produce metal oxides. The reactions are oxidation reactions because the metals gain oxygen.
- Oxidation is the gain of oxygen by a substance.
- Reduction is the loss of oxygen from a substance.

Iron Example

- When iron reacts with oxygen, it forms iron oxide (rust).
- The reaction is:
4Fe + 3O₂ → 2Fe₂O₃
- In this reaction, iron is oxidised (gains oxygen).
- In the opposite reaction:
2Fe₂O₃ → 4Fe + 3O₂
iron oxide is reduced (loses oxygen).

4.1.2 The reactivity series

- When metals react with other substances the metal atoms form positive ions.
- This is because in order to obtain a full outer shell of electrons (8 electrons), metal atoms need to lose electrons.
- The reactivity of a metal is related to its tendency to form positive ions.
The reactivity series is a list of metals (and some non-metals) arranged in order of their reactivity, from most reactive to least reactive. It helps predict how metals will react with water, acids, and other substances.

Potassium (most reactive)
Sodium
Lithium
Calcium
Magnesium
Carbon (non-metal, included for comparison)
Zinc
Iron
Nickel
Hydrogen (non-metal, included for comparison)
Lead
Copper
Silver
Gold

- Metals higher in the series are more reactive and can displace those below them from compounds.

Reactions of metals with water:

- Potassium, sodium, and lithium react vigorously with cold water, producing hydrogen gas and a metal hydroxide.
- Calcium reacts less vigorously with water, forming calcium hydroxide and hydrogen gas.
- Magnesium reacts very slowly with cold water, but more rapidly with hot water.
- Zinc, iron, and copper do not react with cold water.

Reactions of metals with dilute acids:

- Potassium, sodium, and lithium react explosively with dilute acids (these reactions are not usually done for safety reasons).
- Calcium reacts rapidly with dilute acids, producing hydrogen gas.
- Magnesium reacts quickly with dilute acids, producing hydrogen gas.
- Zinc reacts moderately with dilute acids.
- Iron reacts slowly with dilute acids.
- Copper does not react with dilute acids.

Deduction from experiments:

- By observing how quickly metals react with water or dilute acids (e.g., rate of hydrogen gas production), you can deduce their order of reactivity.
- If a metal displaces another from a solution (e.g., zinc displacing iron from iron(II) sulfate), it is more reactive.

4.1.3 Extraction of metals and reduction

- Unreactive metals such as gold are found in the Earth as the metal itself but most metals are found as compounds that require chemical reactions to extract the metal.
- Metals less reactive than carbon can be extracted from their oxides by reduction with carbon.
- Metals more reactive than carbon are extracted by electrolysis, this is covered later.
- For example, iron(III) oxide can be reduced by carbon to extract iron:
2Fe₂O₃ + 3C → 4Fe + 3CO₂
- In this reaction, iron(III) oxide is reduced (loses oxygen) and carbon is oxidised (gains oxygen).
- To interpret extraction processes, identify which substances gain or lose oxygen:

The substance that gains oxygen is oxidised.
The substance that loses oxygen is reduced.

- For example, in the above reaction:

Iron(III) oxide is reduced to iron (loss of oxygen).
Carbon is oxidised to carbon dioxide (gain of oxygen).

- You may be asked to evaluate extraction methods based on reactivity, cost, or environmental impact when given information.
- For example, extracting metals using carbon is often cheaper than electrolysis, but it can produce carbon dioxide, a greenhouse gas.
- The environmental damage of that carbon dioxide is often outweighed by the huge amount of energy required for electrolysis, though.

4.1.4 Oxidation and reduction in terms of electrons

- Oxidation is the loss of electrons and reduction is the gain of electrons.
- In displacement reactions, a more reactive metal displaces a less reactive metal from its compound. These reactions can be represented by ionic equations.
Example: Zinc displaces copper from copper(II) sulfate:
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
The ionic equation is:
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Half equations:

Zn(s) → Zn²⁺(aq) + 2e^- (oxidation, zinc loses electrons)
Cu²⁺(aq) + 2e^- → Cu(s) (reduction, copper ions gain electrons)

- To identify which species are oxidised and which are reduced:

The species that loses electrons is oxidised.
The species that gains electrons is reduced.

- In the above example, zinc is oxidised and copper ions are reduced.

Half Equations

- Half equations show either the oxidation or reduction part of a redox reaction.
- They include the electrons lost or gained in the process.
- For example, in the reaction where magnesium reacts with oxygen to form magnesium oxide:
2Mg + O₂ → 2MgO
The half equations are:
Oxidation: Mg → Mg²⁺ + 2e^-
Reduction: O₂ + 4e^- → 2O^2-
- In this case, magnesium is oxidised (loses electrons) and oxygen is reduced (gains electrons).