QUESTIONS
1. A student is investigating the rate of reaction between magnesium ribbon and dilute hydrochloric acid by measuring the volume of hydrogen gas produced over time using a gas syringe.
The apparatus consists of a conical flask containing 50cm³ of 1.0 mol/dm³ HCl, into which a 3cm strip of magnesium ribbon is dropped, with the flask immediately connected to a gas syringe. The volume of gas collected is recorded at 30-second intervals for 5 minutes.
(a) Sketch and label a suitable diagram of the apparatus you would set up for this experiment. (8 marks)
(b) State THREE precautions you would take to obtain accurate and reliable results in this experiment. (6 marks)
(c) The table below shows a typical set of results a student might obtain. Using this data:
| Time (s) | 0 | 30 | 60 | 90 | 120 | 150 | 180 | 210 | 240 |
|---|---|---|---|---|---|---|---|---|---|
| Volume of H₂ (cm³) | 0 | 18 | 31 | 40 | 46 | 49 | 50 | 50 | 50 |
i) Plot a graph of volume of gas against time. (6 marks)
ii) From your graph, determine the initial rate of reaction (the gradient at t=0). (5 marks)
iii) Determine the time taken for the reaction to reach half of its total gas volume, and explain what this value represents. (4 marks)
iv) Explain, in terms of collision theory, why the rate of gas production decreases as the reaction proceeds. (6 marks)
(d) If the experiment were repeated with the same mass of magnesium but powdered instead of ribbon form, sketch (on the same axes, or describe) how the new curve would differ, and explain your reasoning. (5 marks)
2. Two unlabelled bottles, P and Q, contain aqueous solutions of separate iron salts — one is iron(II) tetraoxosulphate(VI) and the other is iron(III) trioxonitrate(V). Without being told which is which, you are to identify P and Q using the observations in the table below, which have already been recorded from a set of preliminary tests. Study the table carefully and answer the questions that follow.
| Test performed | Observation with P | Observation with Q |
|---|---|---|
| Colour of solution | Pale green | Yellow-brown |
| Add NaOH dropwise, then excess | Dirty green precipitate, insoluble in excess, turns brown on standing in air | Red-brown precipitate, insoluble in excess |
| Add KI solution, then a few drops of starch | No colour change | Solution turns blue-black |
| Add acidified KMnO₄ solution dropwise | Purple colour is decolourised | Purple colour persists (no decolourisation) |
a) Using the evidence in the table, identify solution P and solution Q, giving the ion responsible for each key observation. (8 marks)
b) Write balanced ionic equations for:
i) the reaction of P with NaOH (16 marks total for b, allocate across sub-parts)
ii) the reaction of Q with NaOH
iii) the reaction of Q with KI
c) Explain, in terms of electron transfer, why P decolourises acidified KMnO₄ but Q does not. (6 marks)
d) The precipitate formed from P is observed to "turn brown on standing in air" (as noted in the table). Write a balanced equation for this change and name the type of reaction occurring. (5 marks)
ANSWERS
Question 1
(a) Apparatus diagram (described): A conical flask fitted with a bung/delivery tube leading to a gas syringe. The flask sits on a stand, contains 50cm³ of dilute HCl; the delivery tube runs horizontally/upward from the bung to the gas syringe, which is mounted on a clamp stand and reads the volume of gas collected. A stopwatch is used alongside to record time.
(b) Three precautions:
- Start the stopwatch at the exact instant the magnesium is added and the flask is sealed, to ensure accurate timing from the true start of the reaction.
- Ensure the bung fits tightly and no gas escapes before reaching the syringe, as any leakage would give an artificially low volume reading.
- Keep the temperature of the acid constant (e.g., using a water bath) throughout repeated trials, since temperature significantly affects reaction rate and could introduce inconsistency between runs.
(c) i) (Graph: volume of H₂ (y-axis, cm³) plotted against time (x-axis, s); curve rises steeply at first, then levels off/plateaus at 50cm³ from around t=180s onward, showing a typical rate-decay curve.)
ii) Initial rate = gradient of the tangent at t=0. Using the early data points (0,0) and (30,18) as an approximation:
Initial rate ≈ (18−0)/(30−0) = 0.6 cm³/s
iii) Half of total volume = 50/2 = 25cm³. From the table, this occurs between t=60s (31cm³, already past) — interpolating between t=30 (18cm³) and t=60 (31cm³):
25cm³ falls at approximately t = 30 + [(25−18)/(31−18)]×30 = 30+16.15 ≈ 46 seconds
This value represents the half-life of the reaction with respect to gas evolution — the time taken for the reaction to reach the halfway point of its total extent, which is a useful indicator of overall reaction rate under first-order-like decaying conditions.
iv) As the reaction proceeds, the concentration of HCl decreases (magnesium is in excess or being consumed alongside acid, and acid particles are being used up). By collision theory, a lower concentration of acid means fewer H⁺ ions are present per unit volume, reducing the frequency of effective collisions between H⁺ ions and the magnesium surface, and hence the rate of gas production decreases progressively until the reaction stops (either reactant exhausted).
(d) With powdered magnesium (same mass, but much larger total surface area), the curve would rise more steeply at the start (faster initial rate) and reach the same final plateau volume of 50cm³ sooner, since the same total moles of Mg (same mass) react with the same moles of HCl overall, producing the same total volume of gas — but the increased surface area allows more frequent collisions between acid particles and the metal surface, increasing the rate without changing the total amount of product formed.
Question 2
a) Solution P is iron(II) tetraoxosulphate(VI) (FeSO₄) — identified by its pale green colour (characteristic of Fe²⁺), its precipitate with NaOH being dirty green and darkening on standing (Fe(OH)₂ oxidising to Fe(OH)₃ in air), its lack of reaction with KI (Fe²⁺ cannot oxidise I⁻), and its ability to decolourise acidified KMnO₄ (Fe²⁺ is a reducing agent, readily oxidised to Fe³⁺).
Solution Q is iron(III) trioxonitrate(V) (Fe(NO₃)₃) — identified by its yellow-brown colour (characteristic of Fe³⁺), its red-brown precipitate with NaOH (Fe(OH)₃), its oxidation of iodide to iodine (giving the blue-black starch-iodine complex, confirming Fe³⁺ as an oxidising agent), and its failure to decolourise KMnO₄ (Fe³⁺ cannot reduce MnO₄⁻ further, as it is already in its higher, oxidised state).
b) i) Reaction of P (Fe²⁺) with NaOH:
Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s) (dirty green precipitate)
ii) Reaction of Q (Fe³⁺) with NaOH:
Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) (red-brown precipitate)
iii) Reaction of Q (Fe³⁺) with KI:
2Fe³⁺(aq) + 2I⁻(aq) → 2Fe²⁺(aq) + I₂(aq) (I₂ formed gives blue-black colour with starch)
c) P (Fe²⁺) is readily oxidised to Fe³⁺, losing one electron per ion (Fe²⁺ → Fe³⁺ + e⁻); this electron is transferred to and reduces MnO₄⁻ (Mn⁷⁺→Mn²⁺), causing the purple colour of KMnO₄ to fade (decolourise). Q (Fe³⁺) is already in its higher, more stable oxidation state and has no further electron to donate to MnO₄⁻ — it cannot act as a reducing agent toward KMnO₄, so no electron transfer occurs and the purple colour of KMnO₄ persists unchanged.
d) The pale green/dirty green Fe(OH)₂ precipitate is oxidised by atmospheric oxygen (and moisture) to reddish-brown Fe(OH)₃ on standing in air:
4Fe(OH)₂(s) + O₂(g) + 2H₂O(l) → 4Fe(OH)₃(s)
This is an oxidation reaction (specifically, aerial/atmospheric oxidation, in which Fe²⁺ is oxidised to Fe³⁺ by dissolved/atmospheric oxygen).