Reactions of acids

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수업 노트

Reactions of Acids with Metals

  • Acids react with some metals to produce a salt and hydrogen gas.
  • General equation: metal + acid → salt + hydrogen.
  • Only metals above hydrogen in the reactivity series react with dilute acids; copper does not react.
  • The more reactive the metal, the more vigorous the reaction; potassium and sodium react explosively.
  • The salt produced depends on the acid: hydrochloric acid gives chlorides, sulfuric acid gives sulfates.
  • Example: Mg (s) + 2HCl (aq) → MgCl₂ (aq) + H₂ (g).
  • Test for hydrogen: a lit splint gives a squeaky pop.

Metal-Acid Reactions as Redox (Higher Tier)

  • Metal-acid reactions are redox reactions: reduction and oxidation occur together.
  • Oxidation is loss of electrons; reduction is gain of electrons.
  • In the reaction Zn + 2HCl → ZnCl₂ + H₂, the ionic equation is Zn + 2H⁺ → Zn²⁺ + H₂.
  • Oxidation half-equation: Zn → Zn²⁺ + 2e⁻ (zinc is oxidised).
  • Reduction half-equation: 2H⁺ + 2e⁻ → H₂ (hydrogen ions are reduced).
  • Metal atoms tend to lose electrons and are usually the species oxidised.

Neutralisation and Salt Production

  • A base neutralises an acid; examples include metal oxides, hydroxides and carbonates.
  • An alkali is a soluble base; all alkalis are bases, but not all bases are alkalis.
  • General neutralisation: acid + base → salt + water.
  • With a metal carbonate: acid + carbonate → salt + water + carbon dioxide.
  • The salt name has two parts: the metal (from the base) and the acid ending (chloride, sulfate, nitrate).
  • Hydrochloric acid → chlorides; sulfuric acid → sulfates; nitric acid → nitrates.
  • Example: 2HCl + CuO → CuCl₂ + H₂O.
  • Test for carbon dioxide: bubble through limewater; it turns milky/cloudy.

Soluble Salts and Their Preparation

  • Soluble salts can be made by reacting an acid with an insoluble base (metal, oxide, hydroxide or carbonate).
  • The solid is added in excess to ensure all the acid reacts; otherwise unreacted acid would become dangerously concentrated during evaporation.
  • The excess solid is removed by filtration, leaving a solution of salt and water.
  • The solution is heated to evaporate water until small crystals begin to appear (about half the water removed).
  • Slow evaporation over days gives larger crystals.
  • Example: CuO (s) + H₂SO₄ (aq) → CuSO₄ (aq) + H₂O (l).
  • Acids can also react with metals if the metal is above hydrogen but not too reactive.

Required Practical: Preparation of a Soluble Salt

  • Aim: prepare a pure, dry sample of a soluble salt from an insoluble oxide or carbonate.
  • Add dilute acid to a beaker and warm gently.
  • Add the insoluble oxide slowly until it is in excess (no more dissolves).
  • Filter to remove excess solid, collecting the filtrate in an evaporating basin.
  • Heat gently (water bath or electric heater) to evaporate water and make the solution saturated.
  • Test saturation by dipping a cold glass rod; crystals form on it.
  • Leave to crystallise, then decant excess solution and dry the crystals.
  • Expected result: bright blue, regularly shaped hydrated copper(II) sulfate crystals.

The pH Scale and Neutralisation

  • Acids produce hydrogen ions (H⁺) in aqueous solution; alkalis produce hydroxide ions (OH⁻).
  • The pH scale runs from 0 to 14; pH 7 is neutral.
  • Acids have pH below 7; alkalis have pH above 7.
  • pH 1–3 = strong acid; pH 4–6 = weak acid; pH 8–11 = weak alkali; pH 12–14 = strong alkali.
  • pH can be measured with universal indicator (approximate) or a pH probe (more accurate).
  • Neutralisation: H⁺ (aq) + OH⁻ (aq) → H₂O (l).
  • Not all acid reactions are neutralisations: metal + acid produces a salt but no water.

The pH scale

The pH scale

Strong and Weak Acids

  • Strong acids are completely ionised in water, giving a high concentration of H⁺ and low pH (1–3).
  • Examples of strong acids: hydrochloric, nitric and sulfuric acids.
  • Weak acids are only partially ionised in water, giving a lower concentration of H⁺ and pH 4–6.
  • Examples of weak acids: ethanoic, citric and carbonic acids.
  • Weak acids set up an equilibrium, e.g. CH₃CH₂COOH ⇌ H⁺ + CH₃CH₂COO⁻; the equilibrium lies to the left.
  • Concentration (dilute/concentrated) is different from strength (strong/weak).
  • A dilute strong acid can have a lower pH than a concentrated weak acid.
  • pH is logarithmic: a change of 1 unit means a tenfold change in H⁺ concentration.

Titrations

  • Titrations measure the volume of one solution that reacts exactly with a known volume of another.
  • Acid-base titrations can determine the concentration of an acid or alkali.
  • A pipette measures a fixed volume (e.g. 25.0 cm³) of one solution into a conical flask.
  • A burette delivers the other solution; the volume used is the titre.
  • A suitable indicator (e.g. phenolphthalein or methyl orange) shows the end-point by a sharp colour change.
  • Universal indicator is not suitable for titrations because its colour change is not sharp.
  • Concordant results are within 0.1 cm³ of each other; only these are used to calculate the mean titre.

Titration apparatus set-up

Titration apparatus set-up

Required Practical: Strong Acid & Strong Alkali Titration

  • Aim: determine the reacting volumes of a strong acid and a strong alkali by titration.
  • Use a pipette to place exactly 25.0 cm³ of alkali into a conical flask.
  • Add a few drops of a suitable indicator (e.g. phenolphthalein).
  • Perform a rough titration, then repeat accurately, adding dropwise near the end-point.
  • Record the volume to the nearest 0.05 cm³, reading the bottom of the meniscus at eye level.
  • Repeat until two concordant results are obtained.
  • Calculate the mean titre, ignoring anomalous results.

Titration Calculations (Higher Tier)

  • Use the relationship: moles = concentration × volume (in dm³).
  • Convert cm³ to dm³ by dividing by 1000.
  • Use the balanced equation to find the mole ratio between acid and alkali.
  • Calculate moles of the known solution, then use the ratio to find moles of the unknown.
  • Concentration = moles ÷ volume (in dm³).
  • Example: 25.00 cm³ of 0.15 mol/dm³ Ba(OH)₂ neutralises 12.80 cm³ HNO₃; moles Ba(OH)₂ = 0.15 × 0.025 = 3.75 × 10⁻³ mol; moles HNO₃ = 7.5 × 10⁻³ mol; concentration = 7.5 × 10⁻³ ÷ 0.0128 = 0.59 mol/dm³.

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연습 문제

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  1. 1.Which ion is present in aqueous solutions of all acids?

    Easy
    • AHydroxide ion, OH⁻
    • BHydrogen ion, H⁺
    • CChloride ion, Cl⁻
    • DHydroxide ion, OH⁺

  2. 2.Which ions make a solution alkaline?

    Easy
    • AHydrogen ions, H⁺
    • BHydroxide ions, OH⁻
    • COxide ions, O²⁻
    • DHydroxide ions, OH⁺

  3. 3.All alkalis are bases, but not all bases are alkalis.

    Easy

    True or false?

  4. 4.Which acid should be used to produce magnesium sulfate from magnesium?

    Easy
    • AHydrochloric acid
    • BNitric acid
    • CSulfuric acid
    • DEthanoic acid

  5. 5.Which salt is formed when zinc carbonate reacts with nitric acid?

    Easy
    • AZinc
    • BZinc chloride
    • CZinc nitrate
    • DZinc oxide

  6. 6.Which of the following statements about neutralisation reactions are correct? (select all that apply)

    Medium
    • ANeutralisation always produces water.
    • BThe reaction between an acid and a metal is a neutralisation reaction.
    • CAcid + base → salt + water is the general equation for neutralisation.
    • DWhen the base is a metal carbonate, carbon dioxide is also produced.
    • ENeutralisation always produces hydrogen gas.

  7. 7.Match each acid to the type of salt it produces.

    Medium
    • Hydrochloric acid
    • Sulfuric acid
    • Nitric acid
    • Chlorides
    • Sulfates
    • Nitrates

  8. 8.Which statement correctly describes a strong acid?

    Medium
    • AIt is completely ionised in aqueous solution.
    • BIt is partially ionised in aqueous solution.
    • CIt always has a high concentration.
    • DIt always contains a high number of acid molecules per dm³.

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