Solutions and titration

Unit 11 Solutions and titration

Key unit competence: Prepare standard solutions and use them in volumetric analysis (acid–base and redox titrations) to determine unknown concentrations.

Key ideas

• A standard solution is a solution whose concentration is known accurately. A primary standard is a substance pure and stable enough to be weighed directly to make one: anhydrous Na₂CO₃, potassium hydrogen phthalate, oxalic acid (H₂C₂O₄·2H₂O), K₂Cr₂O₇.
• A primary standard must be very pure, stable in air (not hygroscopic, deliquescent or efflorescent), soluble in water, have a fairly high molar mass (small weighing errors) and react completely and quickly. NaOH is NOT a primary standard: it absorbs water and CO₂ from air.
• To make a standard solution: mass = c × V × M. Weigh, dissolve in a beaker, transfer with rinsings into a volumetric flask, fill to the mark (bottom of the meniscus at eye level), stopper and invert. Dilution: c₁V₁ = c₂V₂.
• Acid–base titration: n = c × V; use the mole ratio of the balanced equation. Indicators: strong acid–strong base: methyl orange or phenolphthalein; weak acid–strong base: phenolphthalein; strong acid–weak base: methyl orange. Back titration is used for insoluble or slow-reacting solids (e.g. CaCO₃): add a known excess of acid, then titrate what is left.
• Redox titrations: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O (KMnO₄ is its own indicator: colourless to permanent pale pink); 2MnO₄⁻ + 5C₂O₄²⁻ + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O (warm to about 60 °C); 2MnO₄⁻ + 5H₂O₂ + 6H⁺ → 2Mn²⁺ + 5O₂ + 8H₂O; Cr₂O₇²⁻ + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O (orange to green, redox indicator needed).
• Iodine titrations: iodimetry uses a standard iodine solution directly; iodometry frees iodine from KI with an oxidant, then titrates it with thiosulfate: I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻. Starch is added near the end point: blue-black to colourless. Titration also finds water of crystallisation and relative atomic masses.

Common misconceptions

• "Any pure-looking solid can be weighed to make a standard solution." No: Only primary standards can; NaOH absorbs water and CO₂, so its solutions must be standardised by titration.
• "In a titration the moles of acid always equal the moles of base." No: Use the mole ratio from the equation: one H₂SO₄ reacts with two NaOH.
• "Potassium manganate(VII) titrations need an indicator." No: MnO₄⁻ is purple and Mn²⁺ almost colourless, so the first permanent pink shows the end point.

Worked examples

Practice questions

1. State four properties of a primary standard.

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   High purity; stable in air (not hygroscopic or efflorescent); soluble in water; high molar mass; reacts completely and quickly (any four).

   Common mistake: Listing 'cheap' or 'coloured' as required properties.

2. Calculate the mass of NaCl needed to make 500 cm³ of a 0.200 mol dm⁻³ solution (M = 58.5).

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   n = 0.100 mol; m = 0.100 × 58.5 = 5.85 g.

   Common mistake: Forgetting to change cm³ into dm³ (dividing by 1000).

3. What volume of 5.00 mol dm⁻³ HCl is diluted to make 250 cm³ of 0.200 mol dm⁻³ acid?

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   V₁ = 0.200 × 250 ÷ 5.00 = 10.0 cm³.

   Common mistake: Using c₁V₁ = c₂V₂ the wrong way round.

4. Why is a titration with KMnO₄ done in H₂SO₄ and not in HCl?

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   KMnO₄ would oxidise Cl⁻ to Cl₂, using extra titrant; sulfate is not oxidised.

   Common mistake: Saying HCl is too weak an acid.

5. Write the ionic equation for the reaction of dichromate(VI) with iron(II) in acid.

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   Cr₂O₇²⁻ + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O

   Common mistake: Writing a 1 : 1 ratio between Cr₂O₇²⁻ and Fe²⁺.

6. Why is starch added only near the end point of an iodine–thiosulfate titration?

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   Too much iodine forms a stable starch–iodine complex that releases iodine slowly, so the end point is unclear.

   Common mistake: Adding starch at the start.

7. Which indicator suits the titration of Na₂CO₃ with HCl to the second end point?

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   Methyl orange (the solution ends acidic because of dissolved CO₂).

   Common mistake: Choosing phenolphthalein, which changes at the first end point.

8. Define the term titre.

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   The volume of the solution added from the burette to reach the end point.

   Common mistake: Confusing it with the volume in the pipette.

9. Why is the conical flask rinsed with distilled water and not with the solution it will contain?

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   Adding water does not change the number of moles pipetted into it; rinsing with the solution would add extra moles.

   Common mistake: Rinsing the conical flask with the pipette solution.

10. Give one application of titration outside the school laboratory.

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    Quality control of vinegar, milk acidity, vitamin C in fruit juice, iron in tablets, water hardness (any one).

    Common mistake: Giving an unrelated use such as filtration.

Matching and drag-and-drop

1. Matching

   Match each item in column A with the correct answer in column B. Write the letter next to each number.

   Column A	Column B
   1. Before titrating oxalate with KMnO4, the oxalate solution is:	a. warmed to about 60 °C
   2. In a KMnO4 titration the end point is seen when the solution becomes:	b. an insoluble solid such as calcium carbonate
   3. A back titration is most useful when the sample is:	c. a standard iodine solution is used directly
   4. Iodimetry differs from iodometry because in iodimetry:	d. permanently pale pink

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   1 → a (warmed to about 60 °C) · 2 → d (permanently pale pink) · 3 → b (an insoluble solid such as calcium carbonate) · 4 → c (a standard iodine solution is used directly)

   • The reaction is slow at room temperature.
   • The first excess of MnO₄⁻ gives a pink colour.
   • The solid reacts with a known excess of acid; the excess is then titrated.
   • Iodometry titrates iodine set free by an oxidant.

2. Drag and drop

   Choose the correct word or phrase from the box for each gap. One word in the box is not needed.

   0.100 mol dm⁻³  ·  starch  ·  methyl orange  ·  the solution it will contain
   1. 25.0 cm³ of 2.00 mol dm⁻³ HCl is diluted to 500 cm³. The new concentration is .
   2. Before use, a burette is rinsed with .
   3. In iodometric titrations the indicator is .
   Show answer

   1: 0.100 mol dm⁻³ · 2: the solution it will contain · 3: starch

   • c₂ = 2.00 × 25.0 ÷ 500 = 0.100 mol dm⁻³.
   • Rinsing with water would dilute the solution.
   • Starch gives a blue-black colour with iodine.

3. Matching

   Match each item in column A with the correct answer in column B. Write the letter next to each number.

   Column A	Column B
   1. The apparatus used to make up a standard solution to an exact volume is a:	a. volumetric flask
   2. A standard solution is a solution:	b. whose concentration is known accurately
   3. The volume of titrant used is read at:	c. I2 + 2S2O32− → 2I− + S4O62−
   4. The equation for the thiosulfate–iodine reaction is:	d. the bottom of the meniscus at eye level

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   1 → a (volumetric flask) · 2 → b (whose concentration is known accurately) · 3 → d (the bottom of the meniscus at eye level) · 4 → c (I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻)

   • It has one accurate graduation mark.
   • Its exact concentration is known, so it can be used to find others.
   • This avoids parallax errors.
   • Two thiosulfate ions reduce one iodine molecule.

4. Drag and drop

   Choose the correct word or phrase from the box for each gap. One word in the box is not needed.

   2.65 g  ·  agree within 0.10 cm³  ·  are all exactly 25.0 cm³  ·  dilute sulfuric acid
   1. Concordant titres are titres that .
   2. Potassium manganate(VII) titrations are carried out in the presence of .
   3. The mass of Na₂CO₃ (M = 106 g mol⁻¹) needed for 250 cm³ of 0.100 mol dm⁻³ solution is .
   Show answer

   1: agree within 0.10 cm³ · 2: dilute sulfuric acid · 3: 2.65 g

   • Only concordant titres are averaged.
   • HCl would be oxidised to Cl₂; H₂SO₄ is not.
   • n = 0.100 × 0.250 = 0.0250 mol; m = 0.0250 × 106 = 2.65 g.

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