This section continues a treatment of the redox reactions of metals.
Syllabus reference R3.2.4Reactivity 3.2.4 - Acids react with reactive metals to release hydrogen.
- Deduce equations for reactions of reactive metals with dilute HCl and H2SO4.
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Reaction with metals
The reaction of metals with acids is often called neutralisation, as the acid gets used up. However, it is nothing of the sort; it is a redox reaction (reduction oxidation). The metal loses its outer electrons and the hydrogen ions from the acid gain electrons to become hydrogen gas. The overall result is a transfer of electrons from the metal to the hydrogen.
M(s) + 2H+(aq) → M2+(aq) + H2(g)
This reaction can only take place if the metal is higher in the reactivity series than hydrogen. In other words the reaction does not work with metals less reactive than lead, such as copper and silver.
Example: The reaction of magnesium with dilute hydrochloric acid:
Mg + 2HCl → MgCl2 + H2
Metals high in the series react very violently with acids and this reaction must not be performed.
Care must be taken with nitric acid, as it does not behave like a typical acid in its reactions with metals. Nitric acid is a strong oxidising agent and preferentially gets reduced to oxides of nitrogen. It is able to react in this way with most metals, it does not depend on any reactivity series.
Example: The reaction of copper with dilute nitric acid:
3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O
It is easier for the metal to reduce the nitrogen atom than to reduce the hydrogen atom in nitric acid. Some metals, such as iron, become 'passive' when treated with nitric acid. This means that there is an initial reaction which soon ceases as the surface of the metal develops an impervious layer.
Finally, magnesium can cause nitric acid to behave as a normal acid (cf: behaviour of nitric acid with most metals). If the acid is very dilute, hydrogen is evolved.