Data-Based Questions (Paper 1B) (DP IB Chemistry: HL): Exam Questions

A student investigated how the distance between two electrodes affects the current in a voltaic cell. She used a magnesium electrode in a solution of magnesium nitrate and a copper electrode in a solution of copper(II) nitrate, connected by a salt bridge containing aqueous potassium nitrate (KNO₃).

The experiment was repeated with two different concentrations of KNO₃: 0.50 mol dm^-3 and 1.00 mol dm^-3. For each trial, the student increased the distance between the electrodes from 2.0 cm to 10.0 cm and recorded the current.

A diagram of the setup is shown below.

Identify the direction of electron flow in the external circuit and explain your answer.

i) Write the half-equation for the reaction occurring at the magnesium electrode.

[1]

ii) Explain whether the magnesium electrode acts as the anode or cathode.

[1]

Explain why the current is higher in the 1.00 mol dm^-3 KNO₃ solution than in the 0.50 mol dm^-3 solution.

The student's results are shown in the graph below.

i) Use the graph to estimate the current for the 1.00 mol dm^-3 solution when the electrodes are 6.0 cm apart.

[1]

ii) Calculate the percentage increase in current at 4.0 cm when using the 1.00 mol dm⁻³ solution instead of the 0.50 mol dm^-3 solution.

[2]

Explain the shape of the graph using ideas about ion movement and solution resistance.

The student considered replacing the Mg and Cu electrodes with platinum.

i) Explain whether this change would improve the reliability of the results.

[2]

ii) Suggest one controlled variable (other than distance, electrode material, and concentration) that should remain the same in each trial.

[1]

A student investigates the buffering effect of different concentrations of a weak acid and its conjugate base. The acid used is ethanoic acid (CH₃COOH), and the conjugate base is sodium ethanoate (CH₃COONa).

The student prepares three buffer solutions, each with a total concentration of 0.200 mol dm^-3, but with varying ratios of acid to conjugate base. The pK_a of ethanoic acid is 4.76.

The pH of a buffer can be calculated using the Henderson-Hasselbalch equation:

pH = pK_a + log$\frac{\left[\mathrm{conjugate}\mathrm{base}\right]}{\left[\mathrm{weak}\mathrm{acid}\right]}$

Calculate the pH of buffer A.

(pK_a of CH₃COOH = 4.76)

Explain why all three buffer solutions have the same total concentration but different pH values.

The student tests the ability of each buffer to resist changes in pH.
A small volume of 1.00 mol dm^-3 HCl is added to each buffer, and the new pH is recorded:

Use the data to identify which buffer had the greatest buffering capacity against added acid, and explain your reasoning.

Buffer B is tested by adding a small amount of sodium hydroxide. The pH increases from 4.76 to 4.94.

Explain why the pH of this buffer changes only slightly.

The student prepares a fourth buffer using 0.200 mol dm^-3 of CH₃COONa only, with no CH₃COOH.

Predict whether this solution can act as a buffer and justify your answer.

Every winter, road salt is spread on icy surfaces to lower the freezing point of water. A student investigates how the mass of calcium chloride (CaCl₂) added to water affects the freezing point of the solution.

In each trial, a measured mass of solid CaCl₂ is added to 100 g of distilled water. The solution is stirred and cooled, and the freezing point is recorded.

Write an ionic equation to show how CaCl₂ dissociates in water.

Explain how the dissociation of CaCl₂ helps to lower the freezing point of water.

The table below shows the freezing points recorded during the investigation.

(i) Describe the trend shown in the data.

[1]

(ii) Calculate the average decrease in freezing point per gram of CaCl₂ added.

[2]

Estimate the freezing point if 30.0 g of CaCl₂ were added.

Explain one reason why your estimate in (d) may not be completely accurate.

The student later discovers the CaCl₂ used was slightly damp and not pure.

Explain how this would affect the results of the experiment.

A second student suggests:

“The bigger the molar mass of the salt, the more it lowers the freezing point.”

Evaluate this suggestion.

A student investigates the effectiveness of four commercial antacids in reducing stomach acidity. The active ingredients of the antacids are shown below:

Each tablet is crushed and added to 25.0 cm³ of 1.00 mol dm^-3 hydrochloric acid (HCl). After 5 minutes, the final pH is recorded.

(i) Write a balanced chemical equation for the reaction between calcium carbonate and hydrochloric acid.

[1]

(ii) State the formula of one reactant ion in antacid tablets that causes the pH to increase..

[1]

Suggest two experimental variables, other than time, that should be controlled to ensure a fair comparison between the antacid tablets.

The results of the experiment are shown below.

(i) Assuming that the initial pH of the acid was 1.00, calculate the change in pH for antacid A.

[1]

(ii) Calculate the uncertainty in the pH change for antacid A, using ±0.02 for each pH value.

[1]

Explain one reason why antacid B may appear more effective than C, even though both contain calcium carbonate.

The student concludes that “Antacid B is the most effective.”

Use the data to evaluate this conclusion.

The student later discovers the antacid B tablet was slightly damp.

Explain how this might affect the result.

Suggest one environmental concern with using excess antacid tablets that contain carbonate or hydroxide compounds.

A student investigates how temperature affects the position of equilibrium in the following reversible reaction:

CO (g) + H₂O (g) ⇌ CO₂ (g) + H₂ (g)

ΔH^Ө = –41.2 kJ mol^-1

The student collects data at different temperatures and uses the equilibrium concentrations to calculate the equilibrium constant, K, for each trial.

The table below shows her results:

The student notes that the value of ln K decreases as temperature increases.

Using the data, describe how the position of equilibrium changes with increasing temperature. Support your answer with reference to the values of K.

Explain this effect using Le Châtelier’s principle and the sign of ΔH^Ө.

Use the data to calculate the standard Gibbs free energy change (ΔG^Ө) at 600 K.

The student then plots a graph of ln K against temperature (T) using the data.

Use the graph to estimate the temperature at which ΔG^Ө for the reaction is zero. Justify your answer.

Evaluate whether this reaction would be thermodynamically feasible at 750 K.

A student investigated the effect of temperature on the rate of hydrogen peroxide (H₂O₂) decomposition using manganese(IV) oxide (MnO₂) as a catalyst. The volume of oxygen gas produced in the first 20 seconds was measured at five different temperatures.

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g)

(i) Describe the trend shown in the data.

[1]

(ii) Use particle theory to explain the effect of increasing temperature on the rate of reaction.

[2]

(i) Calculate the average rate of reaction at 50 ^oC in cm³ s^-1.

[1]

(ii) Calculate the percentage increase in rate when the temperature is raised from 40 ^oC to 50 ^oC.

[2]

Explain why there is only a small increase in rate between 50 ^oC and 60 ^oC.

The student repeats the experiment without MnO₂.

(i) Sketch a second curve on the grid below to show how the reaction rate would differ without a catalyst. Label the curve.

[1]

(ii) Explain why the catalyst affects the rate.

[2]

Suggest one procedural improvement the student could make to increase accuracy in measuring the volume of gas produced.

The graph below shows how the solubility of carbon dioxide (CO₂) in water changes with pressure at a constant temperature of 25 ^oC.

(i) Describe the trend shown by the graph.

[1]

(ii) State the type of relationship shown between pressure and solubility.

[1]

Explain the trend using ideas about particle behaviour.

(i) Use the graph to determine the solubility of CO₂ at 2.5 atm.

[1]

(ii) Use the graph to determine the pressure needed for a solubility of 0.580 g per 100 g H₂O.

[1]

Explain why the graph passes through the origin.

(i) Explain why increasing the temperature would reduce the solubility of CO₂ in water.

[2]

(ii) Predict how the graph would look if temperature increased.

[1]

Suggest one real-world situation that relies on the solubility of gases in liquids.

Explain why the dissolution of CO₂ in water is considered an exothermic process.

A student investigates how different haloalkanes react with aqueous sodium hydroxide. Each halogenoalkane has the molecular formula C₄H₉Br, but a different structure:

Equal volumes of NaOH (aq) are added to each compound in separate test tubes, and the mixture is warmed. The student records the time taken for a white precipitate (AgBr) to appear.

Identify the compound that reacted the fastest and suggest the type of mechanism involved.

Explain why compound Z reacts faster than compound W, using ideas about carbocation stability and mechanism.

The student uses polar protic solvents in all trials.

Explain how this affects the likely substitution mechanism.

Compound W is tested again using an aprotic solvent. The rate increases significantly. Suggest a reason for this observation.

Predict the major organic product formed when compound Z reacts with NaOH (aq). State the type of reaction.

Suggest one experimental method, other than measuring rate, that could help distinguish between the S_N1 and S_N2 mechanisms.