I/GCSE Chemistry Chapter Analysis - Chapter 2: Inorganic Chemistry - Reactivity Series
For I/GCSE Chemistry, you should know:
Reactivity Series:
Metals can be arranged in a reactivity series based on their reactions with water and dilute hydrochloric or sulfuric acid. This series is also known as the "activity series" or "reactivity series" of metals.
The reactivity series of metals is as follows, from the most reactive to the least reactive:
Potassium, Sodium, Lithium, Calcium, Magnesium, Aluminium, Zinc, Iron, Copper, Silver, Gold
This order is determined by the ease with which the metals can lose their valence electrons and form positive ions. The more easily a metal can lose its electrons, the more reactive it is.
Displacement Reactions:
Metals can also be arranged in a reactivity series based on their displacement reactions with metal oxides and aqueous solutions of metal salts. In these reactions, a more reactive metal can displace a less reactive metal from its compounds.
For example, if a piece of iron is placed in a copper(II) sulfate solution, the iron will displace the copper from the solution, forming iron(II) sulfate and copper metal:
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
The order of reactivity determined by these displacement reactions is the same as the order in the reactivity series.
Rusting of Iron:
Iron is a commonly used metal, but it is susceptible to a corrosion process called rusting. Rusting occurs when iron is exposed to air and moisture, and it involves a series of chemical reactions:
4Fe(s) + 3O2(g) + 2H2O(l) → 2Fe2O3 · xH2O(s)
The formation of rust, or iron(III) oxide hydrate, can be prevented or slowed down by using various methods:
- Barrier methods: Coating the iron surface with paint, varnish, or other protective coatings to prevent direct contact with air and moisture.
- Galvanizing: Coating the iron surface with a layer of zinc, which acts as a sacrificial anode and protects the iron from corrosion.
- Sacrificial protection: Connecting the iron to a more reactive metal, such as magnesium or zinc, which corrodes preferentially, protecting the iron.
Oxidation, Reduction, and Redox Reactions:
In the context of the reactivity series, we can understand the concepts of oxidation, reduction, and redox reactions:
- Oxidation is the loss of electrons by a substance, which results in an increase in its oxidation state.
- Reduction is the gain of electrons by a substance, which results in a decrease in its oxidation state.
- Redox (reduction-oxidation) reactions involve both oxidation and reduction, where one substance is oxidized, and another is reduced.
For example, in the reaction between iron and copper(II) sulfate, iron is oxidized, and copper(II) is reduced:
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
The iron acts as a reducing agent, donating electrons and being oxidized, while the copper(II) ion acts as an oxidizing agent, accepting electrons and being reduced.
Work hard for your I/GCSE Chemistry examination!
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