Acids with Reactive Metals (SL IB Chemistry)

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

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