Acids with Reactive Metals
Metals and acids
- The typical reaction of a metal and an acid can be summarised as
acid + metal → salt + hydrogen
- For example:
2HCl (aq) + Zn (s) → ZnCl2 (aq) + H2 (g)
hydrochloric acid + zinc → zinc chloride* + hydrogen
H2SO4 (aq) + Fe (s) → FeSO4 (aq) + H2 (g)
sulfuric acid + iron → iron(II) sulfate* + hydrogen
*zinc chloride and iron(II) sulfate are salts
- A salt is an ionic compound formed when the hydrogen of an acid is replaced by a metal or another positive ion
- Clearly, the extent of the reaction depends on the reactivity of the metal and the strength of the acid
- Very reactive metals would react dangerously with acids and these reactions are not usually carried out
- Metals low in reactivity do not react at all
- For instance, copper does not react with dilute acids
- Stronger acids will react more vigorously with metals than weak acids
- What signs of reaction would be expected to be different between the two?
- Faster reaction seen as:
- more effervescence
- the metal dissolves faster
- Faster reaction seen as:
Ionic Equations
- The reactions of acids and metals can be written as ionic equations showing only the species that has changed in the reaction
- Zinc reacts with hydrochloric acid to give a salt and hydrogen
- Zinc metal is being oxidised to a zin ion as shown in the table
| Full equation | 2HCl (aq) + Zn (s) → ZnCl2 (aq) + H2 (g) |
| Ionic equation |
2H+ (aq) + Zn (s) → Zn2+ (aq) + H2 (g) 2H+ (aq) + 2Cl– (aq) + Zn (s) → Zn2+ (aq) + 2Cl– (aq) + H2 (g) |
| Reducing agent |
Zn (s) Zn is being oxidised to Zn2+ (0 to +2) |
| Oxidising agent |
H+ (aq) in HCl (aq) H+ is being reduced to H2 (+1 to 0) |
Table to show the relative reducing power of metals
| Mg |
strongest reducing agent - most readily becomes oxidised |
| Al |  |
| Zn | |
| Fe | |
| Pb | |
| H | |
| Cu | |
| Ag |
weakest oxidising agent - least readily becomes oxidised |
