To neutralise something means to nullify or cancel its properties. In chemical terms, this means to remove either the acidic or basic properties of a substance rendering it neutral, i.e. not acidic or basic.
Syllabus reference R3.1.7Reactivity 3.1.7 - Acids react with bases in neutralization reactions.
- Formulate equations for the reactions between acids and metal oxides, metal hydroxides, hydrogencarbonates and carbonates.
Guidance
- Identify the parent acid and base of different salts.
- Bases should include ammonia, amines, soluble carbonates and hydrogencarbonates; acids should include organic acids.
Tools and links
- Tool 1, Structure 1.1 - How can the salts formed in neutralization reactions be separated?
- Reactivity 1.1 - Neutralization reactions are exothermic. How can this be explained in terms of bond enthalpies?
- Reactivity 3.2 - How could we classify the reaction that occurs when hydrogen gas is released from the reaction between an acid and a metal?
Neutralisation
Acids react with bases in neutralisation reactions. These are fast, stoichiometric reactions. Bases are the chemical opposites of acids. This means that they contain ions that can neutralise (react with and cancel out) the H+ ions of the acids.
H+ + OH- → H2O
The following types of compounds are classified as bases:
Neutralisation reactions
Below is a list of neutralisation reactions showing the behaviour of acids with different types of bases. Notice that water is always produced in neutralisation.
| sulfuric acid | + | magnesium oxide | → | magnesium sulfate | + | water | ||
| H2SO4 | MgO | MgSO4 | H2O | |||||
| hydrochloric acid | + | calcium oxide | → | calcium chloride | + | water | ||
| 2HCl | CaO | CaCl2 | H2O | |||||
| sulfuric acid | + | sodium hydroxide | → | sodium sulfate | + | water | ||
| H2SO4 | 2NaOH | Na2SO4 | 2H2O | |||||
| hydrochloric acid | + | calcium hydroxide | → | calcium chloride | + | water | ||
| 2HCl | Ca(OH)2 | CaCl2 | 2H2O | |||||
| sulfuric acid | + | zinc carbonate | → | zinc sulfate | + | water | + | carbon dioxide |
| H2SO4 | ZnCO3 | ZnSO4 | H2O | CO2 | ||||
| hydrochloric acid | + | potassium carbonate | → | potassium chloride | + | water | + | carbon dioxide |
| 2HCl | K2CO3 | 2KCl | H2O | CO2 | ||||
| hydrochloric acid | + | sodium hydrogen carbonate | → | sodium chloride | + | water | + | carbon dioxide |
| HCl | NaHCO3 | NaCl | H2O | CO2 |
Note: sulfuric acid always makes salts called sulfates, hydrochloric acid always makes salts called chlorides ( nitric acid makes nitrates, ethanoic acid makes ethanoates - etc.)
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 proceed 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.