Reactions of acids

Important formulae

Acid + alkali / base → salt + water

Acid + metal → salt + hydrogen

Acid + metal oxide → salt + water

Acid + metal hydroxide → salt water

Acid + metal carbonate → salt + water + carbon dioxide

4.2.1 Reactions of acids with metals

(Reactivity: see table in 4.1.2 The reactivity series)

Acid + metal → salt + hydrogen

Acid	Hydrochloric acid	Sulfuric acid
Metal	Metal + hydrochloric acid → metal chloride + hydrogen M + 2HCl → MCl2 + H­2	Metal + sulfuric acid → metal sulfate + hydrogen M + H2SO4 → MSO4 + H­2

Knowledge of reactions limited to those of magnesium, zinc and iron with hydrochloric and sulfuric acids

Acid	Hydrochloric acid	Fe + 2HCl → FeCl2 + H­2
Magnesium	Mg + 2HCl → MgCl2 + H­2	Mg + H2SO4 → MgSO4 + H­2
Zinc	Zn + 2HCl → ZnCl2 + H­2	Zn + H2SO4 → ZnSO4 + H­2
Iron	Fe + 2HCl → FeCl2 + H­2	Fe + H2SO4 → FeSO4 + H­2

4.2.2 Neutralisation of acids and salt production

(See 4.2.1 Reactions of acids with metals – Important formulae)

(See 4.2.4  The pH scale and neutralization)

Redox reaction – reduction and oxidation happens at the same time

Zinc displaces copper from a solution of copper(II) sulfate. Using ionic equations, determine which species undergoes oxidation and which species undergoes reduction. (2)

• Equation: Zn + CuSO₄ → ZnSO₄ + Cu
• Ionic equation: Zn + Cu²⁺ + SO₄^2- → Zn²⁺ + SO₄^2- + Cu
• Zn → Zn²⁺ + 2e^–
  •       Zn oxidised as lost electron
  • Cu²⁺ + 2e^– → Cu = reduction
    •        H reduced as gained electron
    • 4.2.3  Soluble salts
    • Acid + metal oxide → salt + waterFor exampleCopper oxide + sulfuric acid → copper sulfate + waterCuO + H₂SO₄ → CuSO₄ + H­₂ONotes
      • Metal oxides are often insoluble in water – reaction mixture needs warming
      • Ionic equation is always H⁺(aq) + OH^– → H₂O(l)
      • This is a neutralisation reaction as H⁺ are used up
      • Acid + metal hydroxide → salt waterFor exampleSodium hydroxide + nitric acid → sodium nitrate + waterNaOH + HNO₃ → NaNO₃ + H­₂ONotes
        • Ionic equation is always H⁺(aq) + OH^– → H₂O(l)

        This is a neutralisation reaction as H⁺ are used upAcid + metal carbonate → salt + water + carbon dioxide

         For example

        Calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide

        CaCO₃ + 2HCl → CaCl₂ + H­₂O + CO₂Notes

        • When acids react with carbonates, there is fizzing as CO₂(g) is formed
        • Ionic equation is always H⁺(aq) + OH^– → H₂O(l)
        • This is a neutralisation reaction as H⁺ are used up
        • Describe how to make pure, dry sample of named salt from information provided.
        •  Describe a safe method for making pure crystals of copper sulfate from copper carbonate and dilute sulfuric acid. Use the information in the figure above to help you. In your method you should name all of the apparatus you will use. (6)
          1. Pour sulfuric acid in the beaker
          2. Add copper carbonate one spatula at a time until copper carbonate is in excess or until no more effervescence occurs
          3. Filter excess copper carbonate using filter paper and funnel
          4. Pour solution into evaporating basin
          5. Heat using Bunsen burner
          6. Leave to crystallise / for water to evaporate
          7. Decant solution
          8. Pat dry using filter paper

          9.    Wear safety goggles

        • Level 3 (5 – 6 marks): A coherent method is described with relevant detail, and in correct sequence which demonstrates a broad understanding of the relevant scientific techniques and procedures. The steps in the method are logically ordered. The method would lead to the production of valid results.4.2.4 The pH scale and neutralization

4.2.4 The pH scale and neutralization

Acid	Alkalis
Produce H+ in aq solution Proton donors	When dissolved in water, produce OH- Proton acceptors

pH scale (0-14)

• A measure of the acidity / alkalinity of a solution using universal indicator or pH probe
• pH 7 = neutral solution
• pH less than 7 = acids(aq)
• pH more than 7 = alkali(aq)
• Stronger the acid, lower the pH
• As pH decreases by 1 unit, H⁺ concentration of solution increases by factor of 10

Neutralisation reaction

Acid + alkali → salt + water

• Ionic equation is always H⁺(aq) + OH^– → H₂O(l)
• H⁺ are used up

4.2.5 Titrations (chemistry only)

Describe how to carry out tutrations using strong acids and strong alkalis (limited to sulfuric, hydrochloric and nitric acids only) to find the reacting volumes accurately (6)

1. Use pipette to measure 25cm³ of alkali into conical flask.
2. Fill burette with acid using funnel
3. Add 3 drops of indicator eg phenolphthalein to conical flask & swirl
4. Place conical flask on white tile directly below burette
5. Record initial reading of acid in burette
6. Open the burette tap & add acid dropwise from burette with swirling towards endpoint
7. Close burette when colour changes

• Phenolphthalein – pink → colourless
• Methyl orange – yellow → red
• Litmus – blue → red
•   8. Record final reading of acid in burette & calculate titre.
•   9. Repeat 1-8 till obtain concordant result (within 0.1cm³ of each other)The student used phenolphthalein as an indicator rather than universal indicator. Explain why a mixed indicator should not be used for an acid-base titration (2)
  • Mixed indicators change colour gradually over a pH range
  • During an acid-base titration, you want to see a sudden colour change at the endpoint, so you need to use a single indicatorWhy is pipette used to measure HCl? (1)
    • Pipette can measure out a known volume accurately

    Why is burette used to measure sodium carbonate? (1)

    Burette can measure out an unknown volume accurately

• 4.2.6 Strong and weak acids (HT only)Strong acid – completely ionised in aqueous solution
  • Eg hydrochloric, nitric and sulfuric acids

  Weak acid – partially ionised in aqueous solution

  • Eg ethanoic, citric and carbonic acidsDilute – has less acid molecules in a give volume of solutionConcentrated – has more acid molecules in a give volume of solution