Reactions of MetalsAQA GCSE Chemistry: Revision notes
Section 1
How do metals react with oxygen?
Metals react with oxygen to form metal oxides. The general equation is:
2M + O₂ → 2MO (or 4M + O₂ → 2M₂O for some metals)
Key examples:
- Magnesium: 2Mg + O₂ → 2MgO (burns with bright white flame)
- Copper: 2Cu + O₂ → 2CuO (black product when heated)
- Iron: 3Fe + 2O₂ → Fe₃O₄ (or 2Fe + O₂ → 2FeO)
The reactivity of metals with oxygen generally increases down Group 1 (alkali metals are most reactive). Some metals like sodium burn vigorously in air at room temperature, whilst less reactive metals like copper must be heated strongly to react with oxygen.
The product is always a solid metal oxide at room temperature.
Examiners expect you to write balanced equations with correct state symbols. Always write (s) for solid metal oxides. Show the product clearly – some metals form different oxides depending on conditions.
Section 2
What happens when metals react with water?
Metals react with water to form a metal hydroxide and hydrogen gas. The general equation is:
2M + 2H₂O → 2MOH + H₂↑
Key examples:
- Sodium: 2Na + 2H₂O → 2NaOH + H₂↑ (reacts vigorously, floats and melts into a ball)
- Potassium: 2K + 2H₂O → 2KOH + H₂↑ (even more vigorous than sodium; hydrogen ignites with a lilac flame)
- Calcium: Ca + 2H₂O → Ca(OH)₂ + H₂↑ (vigorous reaction)
- Magnesium: Mg + 2H₂O → Mg(OH)₂ + H₂↑ (very slow with cold water; reacts faster with steam)
Reactivity trend with water: Reactivity increases down Group 1. Group 2 metals are less reactive with water than Group 1 metals. Most transition metals do not react with water at all.
Testing for hydrogen: Use a lighted splint at the mouth of the test tube – hydrogen makes a 'pop' sound as it burns.
When sodium reacts with water, use the equation 2Na + 2H₂O → 2NaOH + H₂↑. The hydrogen gas produced is shown with the ↑ arrow symbol to indicate a gas. Balance the equation by checking atoms: 2 Na on each side, 2 O on each side, 4 H on each side.
Students often write the wrong product formulas. Remember: sodium forms NaOH (not Na₂OH), and calcium forms Ca(OH)₂ (not CaOH). Check the charge of the metal ion before writing the hydroxide formula.
Section 3
How do metals react with dilute acids?
Metals react with dilute acids to form a salt and hydrogen gas. The general equation is:
M + 2HA → MA₂ + H₂↑ (for a metal with +2 charge)
Or more generally: Metal + Dilute Acid → Salt + Hydrogen↑
Key examples:
- Magnesium with hydrochloric acid: Mg + 2HCl → MgCl₂ + H₂↑
- Zinc with sulphuric acid: Zn + H₂SO₄ → ZnSO₄ + H₂↑
- Iron with dilute hydrochloric acid: Fe + 2HCl → FeCl₂ + H₂↑
- Calcium with dilute hydrochloric acid: Ca + 2HCl → CaCl₂ + H₂↑
Reactivity trend with dilute acids: Reactivity increases down the periodic table. More reactive metals react vigorously (fast and exothermic), whilst less reactive metals react slowly or not at all. Metals below hydrogen in the reactivity series do not react with dilute acids (e.g. copper, silver, gold).
Rate of reaction: Depends on the metal's position in the reactivity series – more reactive metals produce hydrogen faster. Temperature also affects rate; warming the acid speeds up the reaction.
Always balance equations carefully, paying attention to the charge on the metal ion. If the metal has a +2 charge, you need 2 acid molecules (or 2 H⁺ ions). Show hydrogen with the ↑ symbol to indicate a gas produced.
For magnesium reacting with dilute hydrochloric acid: Mg has a +2 charge, so it needs 2 HCl molecules. The products are MgCl₂ (salt) and H₂↑ (hydrogen gas). Balanced equation: Mg + 2HCl → MgCl₂ + H₂↑
Section 4
How can we explain displacement reactions?
A displacement reaction occurs when a more reactive metal removes a less reactive metal from a compound. The more reactive metal has a greater tendency to lose electrons and form positive ions.
General principle: A metal will displace another metal from its salt solution if the displacing metal is higher in the reactivity series.
Key examples:
- Iron displaces copper: Fe + CuSO₄ → FeSO₄ + Cu (iron is more reactive than copper, so it removes copper from the solution)
- Zinc displaces iron: Zn + FeSO₄ → ZnSO₄ + Fe (zinc is more reactive than iron)
- Copper does NOT displace zinc: Cu + ZnSO₄ → No reaction (copper is less reactive than zinc)
Why this happens: More reactive metals lose electrons more easily. In a displacement reaction, the more reactive metal (M1) donates electrons to the less reactive metal ions (M2⁺), reducing them to atoms:
- M1 → M1⁺ + e⁻ (oxidation of the more reactive metal)
- M2⁺ + e⁻ → M2 (reduction of the less reactive metal ion)
The reactivity series is therefore crucial for predicting whether a displacement will occur.
Think of the reactivity series as a 'strength ranking'. A stronger metal can 'push out' a weaker metal from its partner compound, just like a stronger player replacing a weaker player in a game.
To predict displacement, check the reactivity series: if the metal you're adding is higher (more reactive) than the metal in the salt, displacement will occur. Always write the salt of the displacing metal and show the displaced metal as the product.
Section 5
How do we write ionic equations for displacement reactions? (Higher Tier)
Ionic equations show only the particles that actually change during a reaction, removing spectator ions (ions that do not participate).
Steps to write an ionic equation:
- Write the full balanced equation first
- Write the ionic equation by breaking down all soluble ionic compounds into their ions
- Identify and remove spectator ions (ions present on both sides unchanged)
- Write the net ionic equation with only the reacting species
Example: Iron reacting with copper sulphate
- Full equation: Fe + CuSO₄ → FeSO₄ + Cu
- Ionic equation (showing all ions): Fe(s) + Cu²⁺(aq) + SO₄²⁻(aq) → Fe²⁺(aq) + SO₄²⁻(aq) + Cu(s)
- Net ionic equation: Fe(s) + Cu²⁺(aq) → Fe²⁺(aq) + Cu(s)
The SO₄²⁻ ion is a spectator ion because it appears unchanged on both sides.
Another example: Zinc displacing iron
- Full equation: Zn + FeSO₄ → ZnSO₄ + Fe
- Net ionic equation: Zn(s) + Fe²⁺(aq) → Zn²⁺(aq) + Fe(s)
Again, SO₄²⁻ is spectator and removed.
Key rule: Solids, liquids, and gases are written in molecular form; soluble ionic compounds are written as separate ions (aq).
Examiners look for correct identification of spectator ions and accurate cancellation. Write state symbols (s), (aq), (l), (g) carefully. Remember: all Group 1 hydroxides and nitrates are soluble; most other hydroxides are insoluble.
For magnesium reacting with hydrochloric acid: Full equation: Mg + 2HCl → MgCl₂ + H₂. Ionic form: Mg(s) + 2H⁺(aq) + 2Cl⁻(aq) → Mg²⁺(aq) + 2Cl⁻(aq) + H₂(g). Net ionic: Mg(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂(g) (Cl⁻ is spectator, removed).
Must Know
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Metal + oxygen → metal oxide (e.g. 2Mg + O₂ → 2MgO). Reactivity increases down the periodic table.
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Metal + water → metal hydroxide + hydrogen gas (e.g. 2Na + 2H₂O → 2NaOH + H₂↑). Group 1 metals are most reactive; hydrogen gas produces a 'pop' with a lighted splint.
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Metal + dilute acid → salt + hydrogen gas (e.g. Mg + 2HCl → MgCl₂ + H₂↑). Only metals above hydrogen in the reactivity series react with dilute acids.
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Displacement reactions: A more reactive metal displaces a less reactive metal from its compound (e.g. Fe + CuSO₄ → FeSO₄ + Cu). Use the reactivity series to predict whether displacement occurs.
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Ionic equations (HT): Remove spectator ions to show only the particles that change. Example: Fe(s) + Cu²⁺(aq) → Fe²⁺(aq) + Cu(s). Always use correct state symbols: (s), (aq), (l), (g).
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Reactivity series (top to bottom): Potassium, sodium, lithium, calcium, magnesium, zinc, iron, copper, silver, gold. More reactive metals lose electrons more easily and react faster.
That's the notes covered.
Carry on to the next subtopic.