Practical

2018 O Level Pure Chemistry Practical Paper

Try out the first ever practical exam in the new syllabus

Foreword: Comments on the 2018 O Level Chemistry Practical

This was the first time practical skills were tested in this format: one major examination at the end of two years whereby any topic and lab skill could be asked.

PartContent TestedLab Skills Tested
1. Metal DisplacementSolution concentration
Energy from chemicals
Metals
Measure temperature
Record data in a table
QA observations
Plot graph
2. Redox TitrationSolution concentration
Redox
Titration
3. Reacting CarbonatePercentage purity
Reactions of acids
Thermal decomposition
None, as you are not required to carry out your plan

At first blush, it seems like the 2018 O Level Chemistry practical paper was entirely quantitative. It had three parts involving some form of calculations: temperature changes, titration, and percentage purity.

There was no identification of ions or gases. However, the observations skills of qualitative analysis are sneaked into the paper. A few sub-questions in the paper require you to understand and explain visible observations.

Read on and try out the paper yourself (comfortably seated at home, instead of the hot lab)!


Part 1: Experiment on Metal Displacement

The redox reaction between zinc and copper(II) sulfate is exothermic. In this experiment, you will measure the increase in temperature when zinc is added to different concentrations of copper(II) sulfate.

Reagents and Apparatus

  • Solution A, 0.5 mol/dm3 of copper(II) sulfate
  • Solid B, zinc powder
  • Distilled water
  • Boiling tube, measuring cylinder, and burette

Procedure

  1. Use a measuring cylinder to transfer 10 cm3 of solution A into a boiling tube. Measure the temperature of A and record the value in the table below.
  2. Add about 0.5 to 0.6 g of B into the boiling tube. Use the thermometer to stir and measure the temperature throughout the reaction. Record the highest temperature reached in the table below.
  3. Empty the boiling tube and rinse it with water.
  4. Repeat steps 1 to 3, but with different volumes of A and distilled water as given in the table below, such that the total volume remains at 10 cm3.
  5. Complete the table by measuring and recording the initial and maximum temperatures for each set-up.
Volume of A/ cm3Volume of water/ cm3Initial temperature/ °CFinal temperature/ °CTemperature increase/ °C
10027.545.518.0
8227.542.014.5
6427.538.010.5
4627.535.07.5
2827.532.04.5

Note: The measurements written in blue are hypothetical student answers. However, they are theoretically valid readings that you might get, if you were to conduct the experiment on your own.

QUESTION 1
State two observations when solid B was added to solution A. Explain.

As you cannot conduct the experiment at home, watch the YouTube video above to answer this question.

The blue colour of the solution fades away. Brown copper coats the surface of black zinc powder.

This is because zinc is more reactive than copper, causing zinc to displace copper in copper(II) sulfate. Therefore, brown copper and colourless zinc sulfate solution are formed.

QUESTION 2
Plot a graph of temperature increase against volume of A.

Marker’s note: Your graph should contain (1) a best-fit line drawn with a ruler, (2) labelled axes with units and (3) title.

QUESTION 3
State and explain the trend of the graph plotted in Q2.

The temperature change increases proportionally with the volume of A.

As A is the limiting reactant, using more A causes more copper in copper(II) sulfate to be displaced by zinc. Thus, more heat is released during the displacement reaction to increase the surrounding temperature to a greater extent.

QUESTION 4
Determine the volume of A which contained 0.0035 mol of copper(II) sulfate. Hence, using the graph plotted in Q2, find the increase in temperature when zinc powder was added to the volume of A that contains 0.0035 mol of copper(II) sulfate.

Volume of A = no of moles ÷ concentration = 0.0035 ÷ 0.5 = 0.007 dm3 = 7 cm3

Using the graph, when the volume of A is 7 cm3, the temperature increases by 12.8 °C.

Marker’s Note: The graph is accurate to half the smallest division.

QUESTION 5
Predict the increase in temperature if excess zinc powder was added to 10 cm3 of 1 mol/dm3 of A.

Temperature increase = 2 × 18.0 = 36.0 °C

For the same volume of A, doubling the concentration doubles the number of moles of copper(II) sulfate. As copper(II) sulfate is the limiting reactant, this will double the amount of products formed. Therefore, the increase in temperature will be twice that of the reaction whereby 10 cm3 of 0.5 mol/dm3 of A was used.

QUESTION 6
Suggest a source of error. Suggest a change to improve the accuracy of the results.
  1. Measuring cylinder does not give a precise measurement of volume. Use a burette instead, which has a higher precision of 0.05 cm3.
  2. There is water left behind in the boiling tube after rinsing. Use different dry boiling tubes instead of reusing the same one.
  3. There is heat loss to surrounding, which is more severe when the final temperature reached is a lot higher than room temperature. Insulate the boiling tube with a Styrofoam cup.
QUESTION 7
Predict the increase a temperature if excess iron was added to 10 cm3 of 1 mol/dm3 of A.

Iron is less reactive than zinc. Therefore, it will react less vigorously and release less energy to produce a smaller increase in temperature.

Part 2: Experiment on Redox Titration

Iron supplement tablets contain iron(II) ions. In this experiment, you will find the amount of iron(II) present in the tablets by titrating it against potassium manganate(VII).

Reagents and Apparatus

  • Solution C, which is made by crushing and dissolving 9 tablets into 250 cm3 of water
  • Solution D, 0.01 mol/dm3 of potassium manganate(VII)
  • Dilute sulfuric acid
  • 25.0 cm3 pipette, conical flask, measuring cylinder, burette, small funnel, white tile

Procedure

  1. Use a pipette to transfer 25.0 cm3 of solution C into a conical flask.
  2. Use a measuring cylinder to add 10 cm3 of dilute sulfuric acid into the conical flask to acidify the solution.
  3. Slowly add solution D from a burette to the conical flask, swirling to mix. Stop when the colourless solution in the conical flask has just turned light pink.
  4. Repeat steps 1 to 3 until you have 2 consistent titre values.
QUESTION 8
Record your titration results in an appropriate format. Calculate the average volume of D used to titrate C.

Assume that you did three runs of titration. For your initial rough run, you started at 0.00 cm3 and the end-point was at 19.60 cm3. You continued with the first accurate run, starting at 19.60 cm3 and stopping at the end-point of 39.30 cm3. You then topped up the burette for the second accurate run, restarting at 0.00 cm3 and stopping at the end-point of 19.70 cm3.

Burette volumeRoughFirst AccurateSecond Accurate
Initial reading/ cm3019.400
Final reading/ cm319.4039.1019.60
Titre/ cm319.4019.70 19.60

Average titre value = (19.60 + 19.70) ÷ 2 = 19.65 cm3

Marker’s note: Give your readings in 2 decimal places, ending with either 0 or 5. Do also tick the two consistent readings that you will use to calculate the average.

QUESTION 9
Solution D contains 0.01 mol/dm3 of potassium manganate(VII), KMnO4.

Calculate the number of moles of potassium manganate(VII) in the average volume of solution D found in Q8.

No of moles = concentration × volume = 0.01 × (19.65/1000) = 0.0001965 mol

Marker’s note: To apply the formula, give the volume of solution D in dm3.

QUESTION 10
1 mol of potassium manganate(VII) in solution D react with 5 mol of iron(II) ions in solution C.

Using your answer in Q9, calculate the number of moles of iron(II) ions in 25.0 cm3 of solution C.

No of moles of iron(II) = (5/1) × 0.0001965 = 0.0009825 mol

QUESTION 11
Using your answer in Q10, calculate the number of moles of iron(II) ions in the 250 cm3 of solution C that the 9 tablets dissolved in.

No of moles of iron(II) in 250 cm3 of solution C = (250/25) × 0.0009825 = 0.009825 mol

QUESTION 12
Using your answer in Q11, calculate the mass of iron(II) ions in 1 tablet.

Molar mass of iron(II) = 56 g/mol
Mass of iron(II) ions in 9 tablets = 0.009825 × 56 = 5.502 g
Mass of iron(II) ions in 1 tablet = 5.502 ÷ 9 = 0.611 g

Part 3: Planning Question

The shell of cockles (of the ocean, not your heart) are made up of about 60% carbonate ions, CO32-.

QUESTION 13
Outline an experiment to determine the exact percentage by mass of carbonate ions in 1 g of cockle shells. State any assumption that you would make.

Method 1: Acid-Carbonate Reaction

2H+(aq) + CO32-(aq) ⟶ H2O(l) + CO2(g)

  1. Use mortar and pestle to ground a few shells into powder.
  2. Add about 1 g of powdered shell into a pre-weighed boiling tube.
  3. Measure the combined mass of the powdered shell and the boiling tube with a balance.
  4. Use a dropper to 5 cm dilute nitric acid into the boiling tube.
  5. Immediately connect the boiling tube to a gas syringe via a delivery tube. Measure the increase in volume in the gas syringe throughout the reaction.
  6. Repeat steps 4 and 5 until there is no further increase in volume, which shows that all carbonates have been reacted.
  7. Record the total increase in volume, which is the volume of carbon dioxide produced.
  8. By dividing the volume produced with the molar volume, find the number of moles of carbon dioxide. Use the mole ratio of 1:1 to then find the number of moles of carbonate ions. Multiply that with the molar mass of carbonate ions to find the mass of carbonate ions.
  9. Use the formula of percentage mass = mass of carbonate ions ÷ mass of powdered shell × 100% to find the percentage mass of carbonate ions in 1 g of cockle shells.

For this method, the assumption is that only the carbonate ions in the shell would react with nitric acid to produce carbon dioxide gas.

Method 2: Decomposition of Calcium Carbonate

MCO3(s) ⟶ MO(s) + CO2(g)

  1. Use mortar and pestle to ground a few shells into powder.
  2. Add about 1 g of powdered shell into a pre-weighed boiling tube.
  3. Measure the combined mass of the powdered shell and the boiling tube with a balance.
  4. Connect the boiling tube to a gas syringe via a delivery tube.
  5. Heat the boiling tube strongly with a Bunsen flame until there is no further increase in volume.
  6. Record the total increase in volume, which is the volume of carbon dioxide produced.
  7. By dividing the volume produced with the molar volume, find the number of moles of carbon dioxide. Use the mole ratio of 1:1 to then find the number of moles of carbonate ions. Multiply that with the molar mass of carbonate ions to find the mass of carbonate ions.
  8. Use the formula of percentage mass = mass of carbonate ions ÷ mass of powdered shell × 100% to find the percentage mass of carbonate ions in 1 g of cockle shells.

For this method, the assumption is that the metal carbonate in the shell can undergo thermal decomposition at the temperature of the Bunsen flame.

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