Year 13 required and core practicals

7.1 · Free notes

Practical: measuring reaction rates by initial-rate and continuous-monitoring methods (including iodine–propanone and clock reactions)

Key points

• Rate = change in concentration ÷ time, with units mol dm⁻³ s⁻¹. The rate equation (rate = k[A]^m[B]^n) can only be found by experiment, never from the balanced equation.
• Continuous monitoring: follow one property against time in a single run. Options are gas volume (gas syringe), colorimetry (absorbance), conductivity, mass loss, or quenching and titrating aliquots at intervals.
• On a concentration–time graph the gradient of a tangent gives the rate at that concentration. A constant half-life shows first order; a straight line of constant gradient shows zero order.
• Initial-rate method: run several experiments, changing one concentration at a time and keeping the others and the temperature constant. Use the initial rate so that reactant concentrations are still known and products or the reverse reaction have no effect.
• On a rate–concentration graph, a horizontal line means zero order, a straight line through the origin means first order, and a curve of increasing gradient means second order. Orders can also be found from ratios, e.g. doubling [A] gives ×4 rate for second order.
• Iodine–propanone: CH₃COCH₃ + I₂ → CH₃COCH₂I + HI, catalysed by H⁺. The rate equation is rate = k[CH₃COCH₃][H⁺], so the reaction is zero order in iodine, and [I₂] falls linearly.
• To follow iodine–propanone, withdraw aliquots at timed intervals and add them to sodium hydrogencarbonate, which neutralises the acid catalyst and quenches the reaction. Then titrate the I₂ with sodium thiosulfate using starch near the end-point (blue-black → colourless).
• Clock reaction (persulfate–iodide): S₂O₈²⁻ + 2I⁻ → 2SO₄²⁻ + I₂. A fixed small amount of thiosulfate (2S₂O₃²⁻ + I₂ → 2I⁻ + S₄O₆²⁻) with starch gives a sudden blue-black colour when it runs out. Because a fixed amount of I₂ forms in time t, the initial rate ∝ 1/t.
• Temperature: k = Ae^(−Ea/RT), so ln k = ln A − Ea/RT. A plot of ln k against 1/T has gradient −Ea/R, so Ea = −gradient × 8.31 J K⁻¹ mol⁻¹, which is then converted to kJ mol⁻¹.
• Good practice: use a thermostatted water bath, start the timer when the last reagent is added, and use measuring devices of suitable precision. Calculate % uncertainty = (absolute uncertainty ÷ measured value) × 100, and double the burette uncertainty for a titre (two readings).
• The units of k depend on overall order: zero = mol dm⁻³ s⁻¹, first = s⁻¹, second = dm³ mol⁻¹ s⁻¹. For first order, k = ln2 ÷ t½.

In more detail

• Iodine–propanone rate equation: rate = k[CH₃COCH₃][H⁺], zero order in I₂, so iodine is not involved in the rate-determining step.
• Quenching: adding the sample to sodium hydrogencarbonate neutralises the H⁺ catalyst, stopping the reaction so the iodine can be titrated with thiosulfate.
• Worked Ea: gradient of ln k vs 1/T = −6000 K → Ea = 6000 × 8.31 = 49 900 J mol⁻¹ ≈ 49.9 kJ mol⁻¹.
• Clock reactions give 1/t: the time for a fixed small amount of product to form is inversely proportional to the initial rate, as long as only a small proportion of reactant is used.