You are provided with solution $\mathbf{R}$. You will do a series of experiments. You should: - record your observations and conclusions for each of these experiments - test and name any gases evolved. To prepare for the experiment in (d) , place $1\text{ cm}$ depth of $\mathbf{R}$ in a test-tube. Place a wooden splint into the test-tube and leave it while doing the experiments in (a) , (b) and (c) . (a) To $1\text{ cm}$ depth of $\mathbf{R}$ in a test-tube, add a few drops…

(a) No precipitate; no sulfate or carbonate ions present · (b) White precipitate; chloride ions present · (c)(i) Precipitate dissolves / colourless solution formed · (c)(ii) To prevent other ions (e.g. carbonate) from precipitating with silver nitrate, making the test specific for chloride · (d)(i) Orange-red flame; calcium or lithium ions present · (d)(ii) Calcium and lithium give similar flame colours, so no definite conclusion · (e) White precipitate, insoluble in excess…

Chemistry 5070/31 — May/June 2024

Cambridge O-Level · Practical Test · worked solutions for every part, with the mark scheme

questions

marks

minutes

Topics Experimental Contexts · Analysis, Conclusions and Evaluation · Planning Experiments and Investigations · Observations and Measurements · Qualitative Analysis · Use of Techniques, Apparatus and Materials

Q117MMediumExperimental ContextsObservations and MeasurementsAnalysis, Conclusions and EvaluationPlanning Experiments and Investigations

(a)

(i)

You are going to investigate the reactions of three metals, $\mathbf{X}$ , $\mathbf{Y}$ and $\mathbf{Z}$ , with aqueous copper(II) sulfate.

Read all the instructions carefully before starting the experiments.

Instructions

You are going to do four experiments.

Experiment 1

Use a measuring cylinder to add $25\text{ cm}^3$ of aqueous copper(II) sulfate to a $100\text{ cm}^3$ beaker.
Use a thermometer to measure the initial temperature of the aqueous copper(II) sulfate in the beaker.
Record this temperature to the nearest $0.5\text{ }^\circ\text{C}$ in Table 1.1.
Add all the sample of metal $\mathbf{X}$ to the beaker. Carefully stir the mixture.
The temperature of the mixture will increase. Continue stirring until there is no further temperature increase.
Measure the highest temperature of the mixture. Record this temperature to the nearest $0.5\text{ }^\circ\text{C}$ in Table 1.1.
Leave the beaker to stand while you complete Experiments 2 and 3. You will need it for Experiment 4.

Experiment 2

Use a second beaker to repeat Experiment 1 using $\mathbf{Y}$ instead of $\mathbf{X}$ .

Experiment 3

Use a third beaker to repeat Experiment 1 using $\mathbf{Z}$ instead of $\mathbf{X}$ .

Determine the temperature increase in each reaction and write your answers in Table 1.1.

Table 1.1

experiment	initial temperature / $^\circ\text{C}$	highest temperature / $^\circ\text{C}$	temperature increase / $^\circ\text{C}$
1
2
3

Experiment 4

Slowly add $25\text{ cm}^3$ of dilute sulfuric acid to the contents of the beaker from Experiment 1.

5M Similar

DifficultyMedium-Easy

Worked solution

Answer

Complete Table 1.1 with:

initial and highest temperatures recorded to one decimal place, ending in .0 or .5
every highest temperature greater than the corresponding initial temperature
temperature increase = highest temperature $-$ initial temperature, correctly calculated
Experiment 2 has the greatest temperature increase
Experiment 3 has the lowest temperature increase

Final answer

Completed Table 1.1 with readings to 1 decimal place (.0 or .5), highest temperatures above initial, increases calculated, Experiment 2 greatest and Experiment 3 lowest.

Detailed explanation

Walkthrough

This part is about carrying out the temperature measurements and recording them properly. Read the initial and highest temperatures to the nearest 0.5 °C and write them in the table. Then subtract the initial temperature from the highest temperature for each experiment to get the temperature increase. The mark scheme requires all values to one decimal place ending in .0 or .5, and all highest temperatures must be greater than the initial temperatures. Since the reaction is exothermic, the temperature rises. Experiment 2 should give the greatest temperature increase and Experiment 3 the lowest, because the temperature increase is a measure of how much reaction occurred and therefore of the reactivity of the metal.

Key Takeaways

Temperature readings for a thermometer should be recorded to the precision required, here nearest 0.5 °C.
Temperature increase = highest temperature − initial temperature.
A larger temperature increase means a more exothermic reaction and usually a more reactive metal.

Common Mistakes

Recording temperatures as whole numbers instead of to one decimal place ending in .0 or .5.
Writing a highest temperature that is lower than the initial temperature, which is impossible for an exothermic reaction.
Arithmetic errors in subtracting the initial temperature.
Not making sure Experiment 2 has the greatest increase and Experiment 3 the lowest, which is needed for later parts.

Things to Be Careful About

Use the same thermometer and read it at eye level to avoid parallax error.
Stir continuously until the temperature stops rising, so you record the true highest temperature.
Use the units given in the table header ( $^\circ\text{C}$ ) and do not add units to every cell.

Techniques used

measure temperature to nearest 0.5 °Ccalculate temperature increaserecord results in a table with units

(ii)

Describe the initial appearance of the aqueous copper(II) sulfate used in Experiment 1.

______

1M Similar

(iii)

Describe the final appearance of the mixture in the beaker in Experiment 4.

______

2M Similar

(iv)

Explain how your observations show that $\mathbf{X}$ is in excess in Experiment 1.

______

1M Similar

(b)

Use your results to arrange $\mathbf{X}$ , $\mathbf{Y}$ , and $\mathbf{Z}$ in decreasing order of reactivity.

Explain how the results give this order of reactivity.

most reactive ______
______
least reactive ______

explanation ______

2M Similar

(c)

A student repeats the experiment using a fourth metal.

This metal is the second most reactive of the four metals.

Suggest a temperature increase for this experiment.

______

1M Similar

(d)

The temperature increases calculated are less than the true values for these experiments.

Suggest a reason for this.

Describe an improvement to the method which makes the results closer to the true values.

reason ______
improvement ______

2M Similar

(e)

State and explain the effect of using half the concentration of aqueous copper(II) sulfate on the temperature increase in Experiment 2.

effect ______
explanation ______

3M Similar

Q217MMediumExperimental ContextsQualitative AnalysisAnalysis, Conclusions and Evaluation

You are provided with solution $\mathbf{R}$ .

You will do a series of experiments.

You should:

record your observations and conclusions for each of these experiments
test and name any gases evolved.

To prepare for the experiment in (d), place $1\text{ cm}$ depth of $\mathbf{R}$ in a test-tube. Place a wooden splint into the test-tube and leave it while doing the experiments in (a), (b) and (c).

(a)

To $1\text{ cm}$ depth of $\mathbf{R}$ in a test-tube, add a few drops of dilute nitric acid.

Add $1\text{ cm}$ depth of aqueous barium nitrate.

observations ______
conclusions ______

2M Similar

(b)

To $1\text{ cm}$ depth of $\mathbf{R}$ in a test-tube, add a few drops of dilute nitric acid.

Add $1\text{ cm}$ depth of aqueous silver nitrate.

observations ______
conclusions ______

2M Similar

(c)

To $1\text{ cm}$ depth of aqueous sodium carbonate in a test-tube, add $1\text{ cm}$ depth of aqueous silver nitrate. A white precipitate should form.

Add dilute nitric acid a drop at a time until no further change is seen.

Effervescence of a colourless gas should be observed. The gas turns limewater milky.

(i)

Describe one other observation.

______

1M Similar

(ii)

Suggest why it is important to add dilute nitric acid in (b).

______

1M Similar

(d)

(i)

Place the end of the wooden splint which has been in $\mathbf{R}$ into the flame of a Bunsen burner with the air hole open. Record the first flame colour seen.

first flame colour seen ______
conclusions ______

2M Similar

(ii)

Explain why it is difficult to make a definite conclusion from the flame colour in (d)(i).

______

1M Similar

(e)

To $1\text{ cm}$ depth of $\mathbf{R}$ in a boiling tube, add aqueous sodium hydroxide drop by drop until a change is seen.

Then add excess aqueous sodium hydroxide.

Keep the mixture for use in (f).

observations ______
conclusions ______

3M Similar

(f)

Gently warm the mixture from (e).

observations ______
conclusions ______

3M Similar

(g)

Solution $\mathbf{R}$ is made from a mixture of two different ionic compounds.

Suggest the names of these two compounds.

______

2M Similar

Q36MMedium-HardPlanning Experiments and InvestigationsUse of Techniques, Apparatus and MaterialsAnalysis, Conclusions and EvaluationExperimental Contexts

You are not expected to do any experimental work for this question

Barium carbonate decomposes when heated. The word equation for the reaction is shown.

\text{barium carbonate} \rightarrow \text{barium oxide} + \text{carbon dioxide}

Plan an experiment to determine the percentage loss in mass when barium carbonate is heated.

Your plan must include the use of common laboratory apparatus and a sample of barium carbonate. No other chemicals should be used.

Your plan must include:

the apparatus needed
the method to use and the measurements to take
procedures to ensure that the percentage determined is as accurate as possible
how the measurements are used to determine the percentage loss in mass.

You may draw a diagram to help answer the question.

DifficultyMedium-Hard

Worked solution

Answer

Apparatus

crucible and lid, or a heat-proof dish/boiling tube
balance
Bunsen burner, tripod and pipeclay triangle
tongs and a heat-proof mat

Method and measurements

Place the empty crucible on the balance and record its mass, or set the balance to zero (tare) with the empty crucible on it.
Add a sample of barium carbonate and record the mass of the crucible plus sample.
Heat the crucible strongly with the Bunsen burner for several minutes.
Allow the crucible to cool, then reweigh the crucible and its contents.

Procedures to improve accuracy

Reheat and reweigh repeatedly until the mass is constant (no further change), showing that all the barium carbonate has decomposed.
Repeat the whole experiment and use an average, or at least repeat to check reliability and identify anomalous results.
Let the crucible cool before weighing to avoid affecting the balance reading.
Keep the lid partly on to prevent loss of solid, while still allowing carbon dioxide to escape.

Using the measurements

Mass of barium carbonate used = mass of (crucible + sample) − mass of empty crucible.
Mass of barium oxide left = mass after heating − mass of empty crucible.
Mass loss = mass of barium carbonate used − mass of barium oxide left.
Percentage loss = $\frac{\text{mass loss}}{\text{initial mass of sample}} \times 100$ .

Since the gas escapes, the loss in mass is the loss of carbon dioxide, so:

\text{percentage loss} = \frac{\text{initial mass of sample} - \text{final mass of sample}}{\text{initial mass of sample}} \times 100

Final answer

Plan: weigh a sample in a crucible, heat strongly, cool and reweigh, reheating to constant mass; percentage loss = (initial mass of sample − final mass of sample) / initial mass of sample × 100

Detailed explanation

Walkthrough

This is a planning question, so no actual readings are taken. The answer must describe an experiment that measures the loss in mass when barium carbonate is heated. The word equation tells us that barium carbonate decomposes into barium oxide (solid) and carbon dioxide (gas). Because the carbon dioxide escapes, the mass of the container plus solid decreases. The loss in mass is therefore the mass of carbon dioxide given off.

The mark scheme divides the marks into four sections: apparatus, method and measurements, accuracy, and calculation. One mark is available for each section, and two further marks can be gained from any suitable points, making six marks in total.

Apparatus: A named container that can be heated safely is needed, such as a crucible (with lid), heat-proof dish or boiling tube. A balance is needed to measure mass. A Bunsen burner, tripod, pipeclay triangle and tongs are the usual apparatus for heating a crucible.

Method and measurements: The key measurements are the mass of the container, the mass of the container plus sample before heating, and the mass of the container plus contents after heating. If the balance is tared (set to zero) with the empty container on it, the displayed masses are directly the masses of the sample before and after heating. This is simpler to calculate with.

Accuracy: The most important accuracy point is heating to constant mass. If the sample is heated only once, some barium carbonate may remain undecomposed. Reheating and reweighing until two consecutive readings are the same proves that no further mass is being lost, so the decomposition is complete. Repeating the whole experiment and averaging the results checks reliability and helps identify anomalous results. Cooling the crucible before weighing avoids damaging the balance and prevents convection currents above the hot crucible from affecting the reading. A lid can be used partially open to prevent solid from spitting out while still letting carbon dioxide escape.

Calculation: The mass of the sample used is initial mass of sample. After heating, the mass of the sample is the final mass (mass of barium oxide). Since the only mass lost is the carbon dioxide gas, mass loss = initial mass − final mass. The percentage loss is then:

\frac{\text{initial mass of sample} - \text{final mass of sample}}{\text{initial mass of sample}} \times 100

The question says no other chemicals should be used, and this plan uses only barium carbonate with common laboratory apparatus, so it satisfies that instruction.

Key Takeaways

Thermal decomposition of a carbonate produces a metal oxide and carbon dioxide gas.
When a gas escapes, the loss in mass is equal to the mass of gas given off.
Heating to constant mass is the standard way to ensure a thermal decomposition is complete.
Repeating experiments and averaging results improves reliability and helps spot anomalies.
A percentage change is always calculated as change divided by the original value, multiplied by 100.

Common Mistakes

Forgetting to reheat to constant mass, so the decomposition may be incomplete and the calculated percentage loss is too low.
Using the mass of the container + sample as if it were the mass of the sample alone, without subtracting the mass of the empty container.
Calculating percentage loss from final mass divided by initial mass, rather than loss in mass divided by initial mass.
Closing the lid completely, which traps carbon dioxide and prevents the mass from falling correctly.
Not allowing the crucible to cool before weighing, which can damage the balance or give inaccurate readings.
Not repeating the experiment, so reliability cannot be judged.

Things to Be Careful About

The calculation must use the mass of the sample before heating, not the mass of the container plus sample.
If the balance is tared with the empty container, the displayed masses are already the masses of the sample, which avoids an extra subtraction.
Ensure the final answer expresses the result as a percentage, so include the multiplication by 100.
The phrase “constant mass” is a key accuracy point; heating once is not sufficient.
No numerical answer is expected because this is a plan, not an actual experiment.

Techniques used

select apparatus for heating a solid and measuring its massweigh the sample before and after heating using a balanceheat repeatedly to constant mass to ensure complete decompositioncalculate percentage loss in mass from initial and final massesrepeat readings and take an average to improve reliability