6.1Unit 6. Practical Skills in Chemistry II

Advanced Physical Chemistry Core Practicals — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

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6.1 Advanced Physical Chemistry Core Practicals: Rates, Activation Energy, Ka and Cells — IAL EDEXCEL LINEAR (CHEMISTRY) notes: Describe the core practicals for the titrimetric and clock methods of rate measurement, activation energy, Ka of a weak acid and electrochemical cell EM…

Order in iodine
Other orders
Rate equation
What the symbols mean
Clock reaction
Hydrogen peroxide strength

Part 1 of 9Core Practical 9a: titrimetric method for the order in iodine

Reaction

CH₃COCH₃ + I₂ → CH₃COCH₂I + H⁺ + I⁻, catalysed by acid.

Design

Design a large excess of propanone and acid keeps their concentrations effectively constant.

Mix 25 cm³ of 1.0 mol dm⁻³ propanone with 25 cm³ of 1.0 mol dm⁻³ sulfuric acid.

Add 50 cm³ of 0.02 mol dm⁻³ iodine solution and start the timer.

At regular intervals pipette a 10 cm³ sample into a conical flask.

Quench with a spatula of sodium hydrogencarbonate, which removes the H⁺ catalyst.

Titrate with 0.01 mol dm⁻³ sodium thiosulfate, using starch near the end point: 2S₂O₃²⁻ + I₂ → 2I⁻ + S₄O₆²⁻.

Record the titre against time in a table.

Timing
- keep the timer running and record the time at which the hydrogencarbonate is added, because that is when the reaction stopped.

Part 2 of 9Analysis of Core Practical 9a

Graph

Graph plot titre (proportional to [I₂]) against time.

Shape

Shape a straight line with a negative gradient means a constant rate.

Order in iodine

Order in iodine the rate is independent of [I₂], so the reaction is zero order in iodine.

Other orders

Other orders the reaction is first order in propanone and in H⁺.

Rate equation

Rate equation rate = k[CH₃COCH₃][H⁺], with no [I₂] term.
What the symbols mean: rate (mol dm⁻³ s⁻¹) · k = rate constant · [ ] = concentration (mol dm⁻³) · each power is the order with respect to that reactant
- (units depend on the overall order)

Part 3 of 9Core Practical 9b: iodine clock reaction

Clock reaction

Clock reaction gives a sharp, sudden colour change after a period of time.

Reaction

H₂O₂ + 2I⁻ + 2H⁺ → I₂ + 2H₂O.

Timer

Timer sodium thiosulfate uses up the iodine until it is exhausted, then free iodine gives a blue-black colour with starch.

Measure the solutions in burettes into a small beaker, and the excess sulfuric acid in a measuring cylinder.

Start the reaction by adding 1 cm³ of 0.25 mol dm⁻³ hydrogen peroxide and start the timer.

Stop the timer when the blue-black colour appears.

Vary the volume of potassium iodide, keeping the total volume constant.

Core Practical 9b: iodine clock reaction

Rate

Rate taken as the reciprocal of the time, 1/t.

Graph

Graph rate against [KI] is a straight line through the origin, so the reaction is first order in iodide.

Hydrogen peroxide strength

"10 vol" means 1 cm³ gives 10 cm³ of oxygen, equal to about 3% or 0.979 mol dm⁻³.

Part 4 of 9Core Practical 10: activation energy

Reaction

BrO₃⁻ + 5Br⁻ + 6H⁺ → 3Br₂ + 3H₂O.

Timer

Timer the bromine reacts with phenol, C₆H₅OH + 3Br₂ → C₆H₂Br₃OH + 3HBr,
- and once the phenol is used up bromine bleaches an indicator such as methyl red or methyl orange.

Pipette 10.0 cm³ of phenol solution and 10.0 cm³ of bromide/bromate solution into one boiling tube and add four drops of methyl red.

Pipette 5.0 cm³ of sulfuric acid into a second boiling tube.

Stand both tubes in a water bath at 37 °C (±1 °C) until they reach its temperature.

Mix the contents by pouring from one tube into the other and back, and start the clock.

Time until the methyl red disappears.

Repeat at several other temperatures, using ice for the lowest.

Part 5 of 9Analysis of Core Practical 10

Rate

Rate taken as 1/t at each temperature.

Arrhenius plot

Arrhenius plot ln(1/t) against 1/T (T in kelvin) gives a straight line.

Gradient

Gradient equals −E_a/R.

Activation energy

E_a = −gradient × R, with R = 8.31 J mol⁻¹ K⁻¹.

Arrhenius equation

Arrhenius equation ln k = ln A − E_a/RT.
What the symbols mean: k = rate constant · A = pre-exponential (frequency) factor · E_a = activation energy (J mol⁻¹) · R = 8.31 J mol⁻¹ K⁻¹ · T = temperature (K)

Part 6 of 9Core Practical 11: Ka of a weak acid

Principle
- at the half-equivalence point [HA] = [A⁻], so K_a = [H⁺] and pK_a = pH.

Calibrate the pH meter.

Pipette 25 cm³ of 0.1 mol dm⁻³ ethanoic acid into a conical flask with a few drops of phenolphthalein.

Titrate with 0.1 mol dm⁻³ sodium hydroxide until the indicator just turns pink.

Add a further 25 cm³ of the acid to give a half-neutralised mixture.

Measure the pH of this mixture.

Calculation

[H⁺] = 10^−pH, and a pH of 4.75 gives K_a = 1.8 × 10⁻⁵ mol dm⁻³.

Uncertainty

Uncertainty arises from the pipette, burette and end-point judgement;
- the burette has double the reading uncertainty of a single pipette reading because it has initial and final readings.

Part 7 of 9Core Practical 12: EMF of electrochemical cells

Apparatus

Apparatus two small beakers, clean metal strips (Cu, Zn, Fe, Ag), 1.0 mol dm⁻³ metal ion solutions, a high-resistance voltmeter, crocodile clips and a salt bridge.

Salt bridge

Salt bridge a strip of filter paper soaked in saturated potassium nitrate, changed each time to prevent cross-contamination.

Clean the metal strips with sandpaper to remove oxide coatings.

Fill each beaker about two-thirds with the metal ion solution.

Dip filter paper in saturated KNO₃ and place it between the beakers with both ends immersed.

Connect the metals to the voltmeter and record the steady reading.

Swap the terminals if the reading is negative.

Part 8 of 9Theoretical EMF and analysis

Calculation

E°_cell = E° − E°(negative electrode), the more positive E° minus the more negative E°.
- (positive electrode)

Specimen EMFs

Zn/Cu 1.10 V, Zn/Fe 0.32 V, Fe/Cu 0.78 V, Zn/Ag 1.56 V and Cu/Ag 0.46 V.

Deviation

Deviation measured values differ from theory because standard conditions (298 K, 1.0 mol dm⁻³, 100 kPa) are hard to achieve.

Trend

Trend the relative EMFs still match theory, and the larger the EMF, the bigger the difference in reactivity between the metals.

Part 9 of 9More from the specification

1Electrochemical Cell EMF Determination

Following the experimental procedure for standard half-cells, measured cell potentials are compared with theoretical values calculated using E_cell = E_cathode - E_anode to evaluate typical results.

2Determining Order of Reaction

In kinetics experiments, initial rate measurements are performed with respect to different reactant concentrations,
and a graph plotted of rate versus concentration identifies zero, first,
or second order.

3Continuous Monitoring by Gas Collection

In an experiment where gas is evolved, measuring volume compositions over time provides data for concentration-time curves, finding the instantaneous rate from tangents.

4Evaluation of Experimental and Specimen Results

Comparing obtained lab results against theoretical benchmarks enables identification of experimental errors,
such as incomplete equilibrium or temperature fluctuations.

5Electrochemical Cell EMF Determination

Following the experimental procedure for standard half-cells, measured cell potentials are compared with theoretical values calculated using E_cell = E_cathode - E_anode to evaluate typical results.

6Determining Order of Reaction

In kinetics experiments, initial rate measurements are performed with respect to different reactant concentrations,
and a graph plotted of rate versus concentration identifies zero, first,
or second order.

7Continuous Monitoring by Gas Collection

In an experiment where gas is evolved, measuring volume compositions over time provides data for concentration-time curves, finding the instantaneous rate from tangents.

8Evaluation of Experimental and Specimen Results

Comparing obtained lab results against theoretical benchmarks enables identification of experimental errors,
such as incomplete equilibrium or temperature fluctuations.

Key terms

Term	Meaning
Quenching	Stopping a reaction in a sample so its composition can be analysed.
Clock reaction	A reaction showing a sudden colour change after a set amount of product forms, with 1/t proportional to rate.
Arrhenius plot	A graph of ln k (or ln 1/t) against 1/T whose gradient is −E_a/R.
Half-equivalence point	The point where half of a weak acid has been neutralised so that pH = pK_a.
Salt bridge	A strip soaked in saturated KNO₃ that completes the circuit by allowing ions to move.
Zero order	The rate does not depend on the concentration of that reactant.

Avoid these mistakes

Not recording the quench time in 9a — the reaction stops when hydrogencarbonate is added, so that time is used.

Using temperature in °C for the Arrhenius plot — T must be in kelvin.

Using 1/t as time directly — the rate is proportional to 1/t.

Using a low-resistance voltmeter for cell EMF — a high-resistance one stops current flowing and changing the cell.

Last updated: 15 October 2026

Part 1 of 9Core Practical 9a: titrimetric method for the order in iodine

Part 2 of 9Analysis of Core Practical 9a

Part 3 of 9Core Practical 9b: iodine clock reaction

Part 4 of 9Core Practical 10: activation energy

Part 5 of 9Analysis of Core Practical 10

Part 6 of 9Core Practical 11: Ka of a weak acid

Part 7 of 9Core Practical 12: EMF of electrochemical cells

Part 8 of 9Theoretical EMF and analysis

Part 9 of 9More from the specification

Key terms

Avoid these mistakes

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