Chemistry Exam Solutions – Paper 42

Chemistry Exam Solutions – Paper 42

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Pro Tip: In titration questions, always record burette readings to 2 decimal places and look for concordant results (within 0.10 cm³).

Key Study Areas

Titration Techniques

Master the use of burettes, pipettes, and indicators

Energy Changes

Understand exothermic and endothermic reactions

Qualitative Analysis

Learn tests for ions and gases

Experimental Planning

Practice designing experiments with proper controls

Question 1

A student finds the amount of iron(II) ions in a solution by titration.

The student uses a volumetric pipette to add 25.0 cm³ of aqueous iron(II) sulfate to a conical flask, adds approximately 20 cm³ of dilute sulfuric acid, and slowly adds aqueous potassium manganate(VII) until the solution just turns pink.

Solution

(a)(i) The apparatus shown in Fig. 1.1 is a measuring cylinder.

(a)(ii) The student does not need to use a volumetric pipette because the exact volume of acid is not critical to the titration calculation; it only provides acidic conditions.

(a)(iii) To use the volumetric pipette safely: Use a pipette filler, ensure the bottom of the meniscus is at the calibration mark, and allow the pipette to drain freely.

(b) The apparatus used to add potassium manganate(VII) is a burette.

(c) No indicator is needed because potassium manganate(VII) is self-indicating – it changes from purple to colorless.

(d) The student knows they have done enough titrations when they obtain concordant results (titres within 0.10 cm³ of each other).

Strategy & Key Concepts

Redox Titration: This is an example of a redox titration where potassium manganate(VII) oxidizes Fe²⁺ to Fe³⁺. The manganate(VII) ion is reduced from MnO₄⁻ (purple) to Mn²⁺ (colorless).

Self-Indicating Titrations: Potassium manganate(VII) titrations are unique because the titrant itself acts as an indicator. The first permanent pink color indicates the endpoint.

Acidic Conditions: Sulfuric acid is added to provide H⁺ ions needed for the redox reaction and to prevent precipitation of manganese(IV) oxide.

Redox Reaction

Click to reveal the equation

MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

Tips & Tricks

Quick Tip: In manganate(VII) titrations, the first permanent pink color (lasting about 30 seconds) indicates the endpoint.

Common Mistake: Adding too much acid can sometimes interfere with the reaction. Use approximately 20 cm³ as stated.

Memory Aid: “Manganate needs acid to oxidize” – remember the acidic conditions are crucial.

Related Problems & Solutions

Problem 1: Why is hydrochloric acid not used to acidify the solution in manganate(VII) titrations?

Solution: HCl would be oxidized by manganate(VII) to chlorine gas, interfering with the titration.

Problem 2: If the solution turns brown at the endpoint, what might have happened?

Solution: This indicates the formation of manganese(IV) oxide, usually due to insufficient acid.

Problem 3: How would you calculate the concentration of iron(II) ions from titration data?

Solution: Use the mole ratio from the balanced equation (1:5 for MnO₄⁻:Fe²⁺) and the formula: concentration = (moles × 1000) / volume.

Question 2

A student investigates the temperature change when a solid completely dissolves in water.

The student measures temperature changes over time as a solid dissolves in water.

Solution

(a)(i) Completed table:

Time (s) Temperature (°C)
0 19.5
60 13.0
120 13.5
180 14.0
240 14.5
300 15.0

(a)(ii) Maximum temperature change = 6.5°C (19.5 – 13.0)

(a)(iii) The temperature decreases rapidly at first, reaches a minimum, then gradually increases toward room temperature.

(a)(iv) Temperature after 60 minutes: Approximately 19.5°C (room temperature)

Explanation: The solution will eventually reach thermal equilibrium with the surroundings.

(b) Description: Energy is absorbed from the surroundings (endothermic process)

Explanation: The temperature decrease shows the reaction is taking in heat energy.

(c) Stirring is important to ensure even temperature distribution and complete dissolution.

(d) To obtain more precise temperature measurements: Use a digital thermometer with higher precision and ensure it’s properly calibrated.

(e) Reason: Heat exchange with the surroundings

Improvement: Use a polystyrene cup with a lid to insulate the reaction mixture.

Strategy & Key Concepts

Energy Changes in Dissolution: When a solid dissolves, two processes occur:

  • Breaking solute-solute bonds (requires energy – endothermic)
  • Forming solute-solvent bonds (releases energy – exothermic)

The overall energy change determines whether the process is exothermic or endothermic.

Temperature Measurement: Always:

  • Stir before reading the thermometer
  • Allow time for the thermometer to equilibrate
  • Read at eye level to avoid parallax error

Heat Loss Prevention: Polystyrene is a better insulator than glass, reducing heat exchange with surroundings.

Tips & Tricks

Quick Tip: If temperature decreases → endothermic process. If temperature increases → exothermic process.

Common Mistake: Not stirring consistently can lead to uneven temperature readings.

Memory Aid: “Endo-in” – endothermic processes take energy in (temperature drops).

Related Problems & Solutions

Problem 1: How would the graph differ for an exothermic dissolution?

Solution: Temperature would increase rapidly, reach a maximum, then gradually decrease toward room temperature.

Problem 2: Why does the temperature eventually return to room temperature?

Solution: The system reaches thermal equilibrium with the surroundings through heat exchange.

Problem 3: How would using a larger beaker affect the results?

Solution: A larger beaker has greater surface area, leading to faster heat loss and a larger apparent temperature change.

Question 3

A student does a series of experiments to investigate solution R through various chemical tests.

Solution

(a)(i) The observation is a yellow flame.

(a)(ii) The air hole must be open to produce a hot, non-luminous flame that doesn’t interfere with the flame color.

(b)(i) Conclusions:

  • R contains carbonate ions (CO₃²⁻)
  • The gas produced is carbon dioxide

(b)(ii) The observation is a white precipitate forms.

(b)(iii) R does not contain chloride, bromide, or iodide ions.

(b)(iv)

Test: Hold damp red litmus paper near the mouth of the test tube

Observation: The litmus paper does not turn blue

(c) The two compounds are sodium carbonate and sodium sulfate.

(d) The student could prepare known samples of chloride and bromide ions and compare the precipitate colors side by side with solution P.

(e) To pass gas through limewater: Use a delivery tube from the reaction vessel into a test tube containing limewater, ensuring the end of the tube is below the surface of the limewater.

Strategy & Key Concepts

Systematic Qualitative Analysis: Always follow a logical sequence:

  1. Flame tests for metal ions
  2. Tests for carbonate first (as it interferes with other tests)
  3. Tests for sulfate, then halides
  4. Tests for ammonium last

Flame Test Colors:

  • Yellow – Sodium
  • Lilac – Potassium
  • Red – Lithium
  • Orange-red – Calcium

Precipitation Tests:

  • White precipitate with Ba²⁺ – Sulfate
  • White precipitate with Ag⁺ – Chloride
  • Cream precipitate with Ag⁺ – Bromide

Tips & Tricks

Quick Tip: Always acidify before testing for sulfate or halide ions to remove carbonate interference.

Common Mistake: Confusing cream (bromide) and yellow (iodide) precipitates. Compare with known samples.

Memory Aid: “Cats Bite Yellow Insects” – Chloride (white), Bromide (cream), Yellow (iodide).

Related Problems & Solutions

Problem 1: How would you distinguish between sodium chloride and sodium sulfate?

Solution: Add barium nitrate solution – sulfate gives a white precipitate, chloride does not.

Problem 2: Why is nitric acid used for acidification instead of hydrochloric acid?

Solution: Hydrochloric acid would introduce chloride ions that could give false positives in halide tests.

Problem 3: A solution gives a white precipitate with both NaOH and ammonia. What cation might be present?

Solution: Zinc or aluminum ions (both form white precipitates with NaOH, soluble in excess).

Question 4

Plan an experiment to determine the volume of carbon dioxide formed when a known mass of copper(II) carbonate completely reacts with dilute sulfuric acid.

Copper(II) carbonate + sulfuric acid → copper(II) sulfate + carbon dioxide + water

Solution

Apparatus Needed:

  • Conical flask
  • Delivery tube with bung
  • Gas syringe (100 cm³)
  • Measuring cylinder (for acid)
  • Balance (accurate to 0.01 g)
  • Stand and clamp

Method:

  1. Set up apparatus: conical flask connected via delivery tube to gas syringe.
  2. Measure a known mass of copper(II) carbonate (e.g., 1.00 g) using the balance.
  3. Add the copper(II) carbonate to the conical flask.
  4. Measure a fixed volume of dilute sulfuric acid (e.g., 50 cm³) using a measuring cylinder.
  5. Add the acid to the flask and quickly seal with the bung and delivery tube.
  6. Record the final volume of gas in the syringe when no more gas is produced.
  7. Repeat with different masses to verify results.

Procedures for Accuracy:

  • Ensure all connections are airtight to prevent gas leakage
  • Use excess acid to ensure all carbonate reacts
  • Allow the reaction to complete before taking final reading
  • Measure gas volume at room temperature and pressure
  • Repeat the experiment and calculate mean volume

Strategy & Key Concepts

Gas Collection Methods:

  • Gas syringe – most accurate for measuring gas volumes
  • Downward displacement of water – for gases insoluble in water
  • Upward displacement of air – for gases less dense than air

Stoichiometry Application: The volume of CO₂ produced can be used to verify the reaction equation and calculate the molar volume of gases.

Experimental Design Principles:

  • Control variables: temperature, pressure, acid concentration
  • Minimize errors: use precise measurements, ensure complete reaction
  • Safety: work in well-ventilated area, handle acids carefully

Tips & Tricks

Quick Tip: To check if the reaction is complete, gently swirl the flask – if no more fizzing, the reaction is done.

Common Mistake: Not ensuring airtight connections, leading to gas loss and inaccurate volume measurements.

Memory Aid: “Gas collection needs perfect connection” – emphasize airtight apparatus.

Related Problems & Solutions

Problem 1: How would you calculate the theoretical volume of CO₂ from a given mass of carbonate?

Solution: Use the formula: Volume = (mass/Mr) × 24,000 cm³ (at RTP)

Problem 2: Why is a gas syringe preferred over collecting gas over water for CO₂?

Solution: CO₂ is slightly soluble in water, which would lead to inaccurate volume measurements.

Problem 3: How would using powdered carbonate instead of lumps affect the reaction?

Solution: Powder has larger surface area, leading to faster reaction but the same final gas volume.