Required Practical 8 • Unit 4 / Unit 5

RP8: Measuring Reaction Rates

Investigate chemical reaction kinetics using continuous monitoring (gas syringe volume vs time with initial rate tangents at t=0) and initial rate clock reactions (iodine clock kinetics and disappearing cross Arrhenius analysis).

Revise the theory

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1. Continuous Monitoring vs Initial Rate Methods

In OxfordAQA A-Level Chemistry (Section 6.1, RP8), reaction kinetics are evaluated using two complementary experimental approaches:

Method A: Continuous Monitoring

Tracks how a physical property changes continuously over the entire course of a single reaction run:

  • Gas Syringe Collection: Measuring volume of gas evolved (e.g. H₂ from Mg + HCl, or CO₂ from CaCO₃ + HCl) at regular time intervals (every 15 s).
  • Colorimetry: Tracking absorbance of a colored ion or halogen over time.
  • Continuous Mass Loss: Placing a reaction flask on an analytical balance to measure mass drop as gas escapes.

Method B: Initial Rate (Clock Reactions)

Measures the time taken for an observable chemical event to occur at the very beginning of the reaction across multiple separate runs with varied concentrations or temperatures:

  • Iodine Clock Reaction: Persulfate (S₂O₈²⁻) oxidising iodide in the presence of fixed thiosulfate and starch indicator, timing the sudden switch from colourless to blue-black.
  • Disappearing Cross: Timing sulfur turbidity from Na₂S₂O₃ + HCl to evaluate activation energy (Arrhenius).
  • Initial Rate: Evaluated at t = 0 before reactant concentrations change appreciably.

2. Continuous Gas Volume Monitoring Protocol

To investigate the reaction between magnesium ribbon and hydrochloric acid:

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
  1. Clamp a clean, dry 100 cm³ glass gas syringe horizontally. Push the plunger fully in to read exactly 0 cm³. Check that the plunger moves freely with zero friction.
  2. Measure 25.0 cm³ of 1.0 mol dm⁻³ hydrochloric acid into a conical flask using a volumetric pipette.
  3. Clean a 5 cm strip of magnesium ribbon with emery paper to remove unreactive magnesium oxide (MgO) surface coating.
  4. Add the magnesium ribbon to the acid, immediately replace the rubber bung tightly, and start the stopwatch simultaneously.
  5. Swirl continuously to prevent hydrogen bubbles from adhering to the metal surface.
  6. Record the gas syringe volume reading every 15 seconds until the volume reaches a constant plateau (reaction complete).
Continuous gas collection with a gas syringe 020406080100 volume / cm³ H₂ gas Gas syringe Magnesium ribbon Delivery tube Airtight bung clamped horizontally in dilute HCl(aq) Mg + 2HCl → MgCl₂ + H₂ plunger pushed out clamp stand

3. Tangent Construction at t = 0 for Initial Rate

Reaction rate changes continuously as reactants are consumed. The initial rate at t = 0 provides the cleanest kinetic data because concentrations are known precisely with zero product inhibition:

  1. Plot recorded gas volume (V in cm³) on the y-axis against time (t in s) on the x-axis.
  2. Draw a smooth best-fit curve starting from the origin (0,0).
  3. Place a clear ruler tangent to the curve exactly at t = 0. The ruler must match the initial slope of the curve before curvature begins.
  4. Extend the tangent line across the graph grid to form a large gradient triangle spanning at least half the length of the line.
  5. Calculate gradient: Initial Rate = ΔV / Δt (in cm³ s⁻¹).
Initial rate from a tangent at t = 0 020406080100 050100150200 time / s volume of gas / cm³ Tangent at t = 0 touches the curve only at the start ΔV = 60 cm³ Δt = 30 s plateau: reaction complete initial rate = ΔV ÷ Δt = 60 ÷ 30 = 2.0 cm³ s⁻¹

4. The Iodine Clock Initial Rate Method

The classic clock reaction between peroxodisulfate(VI) ions and iodide ions:

S₂O₈²⁻(aq) + 2I⁻(aq) → 2SO₄²⁻(aq) + I₂(aq)    (slow reaction)

To measure the initial rate, a known, small, fixed volume of sodium thiosulfate (Na₂S₂O₃) and starch indicator are added at the start. As iodine forms, it is immediately scavenged and reduced back to iodide by thiosulfate:

2S₂O₃²⁻(aq) + I₂(aq) → S₄O₆²⁻(aq) + 2I⁻(aq)    (Instantaneous scavenger)

Clock Reaction Mechanism & Time Relationship

  • As long as thiosulfate remains in the flask, no free I₂ can accumulate; the solution remains completely colourless.
  • The exact instant all thiosulfate is completely consumed, newly produced I₂ reacts with starch to produce an abrupt, dramatic color change from colourless to blue-black.
  • Because the amount of thiosulfate is constant across all runs, the amount of iodine formed is constant: Δ[I₂] = constant.
  • Therefore: Initial Rate = Δ[I₂] / Δt ∝ 1 / t.

5. Sodium Thiosulfate Disappearing Cross & Temperature Dependence

To investigate the effect of temperature on reaction rate and determine the activation energy (Ea) via the Arrhenius equation:

Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + SO₂(g) + S(s) + H₂O(l)
  1. Mark a bold black cross (an 'X') on white paper. Place a clean 100 cm³ conical flask directly on top.
  2. Add 25.0 cm³ of 0.050 mol dm⁻³ sodium thiosulfate solution to the flask. In a boiling tube, measure 5.0 cm³ of 1.0 mol dm⁻³ hydrochloric acid.
  3. Place both containers into a thermostated water bath set to target temperatures (e.g. 20, 30, 40 and 50 °C).
  4. Once both reach thermal equilibrium, record the actual temperature of the thiosulfate solution.
  5. Add the acid to the flask, swirl once, place over the cross, and immediately start the stopwatch.
  6. Look vertically down through the neck of the flask. Stop the stopwatch the instant the cross disappears. Record time t.
  7. Immediately pour the contents into a sodium carbonate stop bath to neutralise acid and dissolve SO₂.
Disappearing cross experimental setup Look vertically down through the neck Na₂S₂O₃(aq) + HCl(aq) colloidal sulfur forms: mixture turns cloudy Paper marked with a bold black cross View down the neck start cross clear stop the clock cross hidden

Safety Hazard: Sulfur Dioxide (SO₂) Gas Containment

Sulfur dioxide is a toxic, choking respiratory irritant that can trigger severe asthma attacks. To ensure laboratory safety:

  • Work in a well-ventilated laboratory or fume cupboard.
  • Do not heat solutions above 55 °C to minimise excessive volatilisation of SO₂.
  • Sodium Carbonate Stop Bath: Conical flasks must be dumped immediately into a stop bath containing excess sodium carbonate or sodium hydrogencarbonate solution. The carbonate neutralises residual acid and neutralises dissolved SO₂: SO₂ + H₂O + Na₂CO₃ → Na₂SO₃ + CO₂ + H₂O.

6. Deducing Reaction Orders & Arrhenius Activation Energy

Determining Reaction Orders

In a clock experiment where Rate ∝ 1/t, comparing runs allows deduction of individual reactant orders:

  • Zero Order (order = 0): Doubling reactant concentration has no effect on rate (time t remains constant).
  • First Order (order = 1): Doubling reactant concentration doubles the rate (time t is halved: t → t/2).
  • Second Order (order = 2): Doubling reactant concentration quadruples the rate (time t is quartered: t → t/4).

Arrhenius Equation for Activation Energy

The relationship between rate constant k (approximated by 1/t) and absolute temperature T (in Kelvin):

ln(k) = -Ea / (R * T) + ln(A)   ⇒   ln(1/t) = -Ea / (R * T) + constant

Plotting ln(1/t) on the y-axis against 1/T (in K⁻¹) on the x-axis yields a straight line with gradient = -Ea / R. Given gas constant R = 8.314 J K⁻¹ mol⁻¹, the activation energy is calculated as:

Ea = -gradient * R   (J mol⁻¹)  ⇒  divide by 1000 for kJ mol⁻¹

7. Worked Kinetics Problem

Example: Persulfate-Iodide Clock Reaction Data

A series of experiments was conducted at 20 °C with fixed amounts of thiosulfate and starch:

Run [S₂O₈²⁻] / M [I⁻] / M Time t / s Rate ∝ 1/t / s⁻¹
1 0.040 0.020 88 0.0114
2 0.080 0.020 44 0.0227
3 0.040 0.040 44 0.0227

Step 1: Order with respect to [S₂O₈²⁻]: Comparing Run 1 and Run 2, [I⁻] is constant at 0.020 M. [S₂O₈²⁻] doubles (0.040 → 0.080 M) and rate doubles (0.0114 → 0.0227 s⁻¹). Therefore, order = 1.

Step 2: Order with respect to [I⁻]: Comparing Run 1 and Run 3, [S₂O₈²⁻] is constant at 0.040 M. [I⁻] doubles (0.020 → 0.040 M) and rate doubles (0.0114 → 0.0227 s⁻¹). Therefore, order = 1.

Step 3: Rate Equation: Rate = k [S₂O₈²⁻] [I⁻]. Overall order = 2.

8. Practice Questions

Question 1: In the continuous monitoring of Mg + HCl using a gas syringe, why must the initial rate be calculated from a tangent at t = 0 rather than simply dividing total volume by total time?

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Correct Answer: B

Reactant concentrations fall continuously during the reaction, meaning the rate slows down. An initial rate tangent taken at t = 0 measures the true rate when the reactant concentrations match their initial, precisely measured values.

Question 2: In an Arrhenius plot of ln(1/t) against 1/T for the disappearing cross reaction, a student obtains a line with a gradient of -4850 K. What is the experimental activation energy (R = 8.314 J K⁻¹ mol⁻¹)?

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Correct Answer: A

Gradient = -Ea / R. Therefore, Ea = -(-4850 K) * 8.314 J K⁻¹ mol⁻¹ = +40,323 J mol⁻¹ = +40.3 kJ mol⁻¹.

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Describe how a concentration-time graph for a zero-order reactant differs from one for a first-order reactant.

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  • Zero order: a straight line with a constant (negative) gradient, because the rate does not depend on concentration [1]
  • First order: a curve whose gradient gets less steep as the concentration falls [1]
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