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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

Section 1 of 7Analysis 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)

Section 2 of 7Analysis 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 −Ea/R.
Activation energy
  • Ea = −gradient × R, with R = 8.31 J mol−1 K−1.
Arrhenius equation
  • Arrhenius equation ln k = ln A − Ea/RT.
    What the symbols mean: k = rate constant · A = pre-exponential (frequency) factor · Ea = activation energy (J mol−1) · R = 8.31 J mol−1 K−1 · T = temperature (K)

Section 3 of 7Theoretical 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.

Section 4 of 7More 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 Ecell = Ecathode - Eanode 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 Ecell = Ecathode - Eanode 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.

Section 5 of 7Key terms

TermMeaning
QuenchingStopping a reaction in a sample so its composition can be analysed.
Clock reactionA reaction showing a sudden colour change after a set amount of product forms, with 1/t proportional to rate.
Arrhenius plotA graph of ln k (or ln 1/t) against 1/T whose gradient is −Ea/R.
Half-equivalence pointThe point where half of a weak acid has been neutralised so that pH = pKa.
Salt bridgeA strip soaked in saturated KNO3 that completes the circuit by allowing ions to move.
Zero orderThe rate does not depend on the concentration of that reactant.

Section 6 of 7Syllabus checklist

Reaction
  • CH3COCH3 + I2 → CH3COCH2I + H+ + I−, catalysed by acid.
Design
  • Design a large excess of propanone and acid keeps their concentrations effectively constant.
1
Mix 25 cm3 of 1.0 mol dm−3 propanone with 25 cm3 of 1.0 mol dm−3 sulfuric acid.
2
Add 50 cm3 of 0.02 mol dm−3 iodine solution and start the timer.
3
At regular intervals pipette a 10 cm3 sample into a conical flask.
4
Quench with a spatula of sodium hydrogencarbonate, which removes the H+ catalyst.
5
Titrate with 0.01 mol dm−3 sodium thiosulfate, using starch near the end point: 2S2O32− + I2 → 2I− + S4O62−.
6
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.
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.
1
Measure the solutions in burettes into a small beaker, and the excess sulfuric acid in a measuring cylinder.
2
Start the reaction by adding 1 cm3 of 0.25 mol dm−3 hydrogen peroxide and start the timer.
3
Stop the timer when the blue-black colour appears.
4
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.
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.
1
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.
2
Pipette 5.0 cm3 of sulfuric acid into a second boiling tube.
3
Stand both tubes in a water bath at 37 °C (±1 °C) until they reach its temperature.
4
Mix the contents by pouring from one tube into the other and back, and start the clock.
5
Time until the methyl red disappears.
6
Repeat at several other temperatures, using ice for the lowest.
  • Principle
    • at the half-equivalence point [HA] = [A−], so Ka = [H+] and pKa = pH.
1
Calibrate the pH meter.
2
Pipette 25 cm3 of 0.1 mol dm−3 ethanoic acid into a conical flask with a few drops of phenolphthalein.
3
Titrate with 0.1 mol dm−3 sodium hydroxide until the indicator just turns pink.
4
Add a further 25 cm3 of the acid to give a half-neutralised mixture.
5
Measure the pH of this mixture.
Calculation
  • [H+] = 10−pH, and a pH of 4.75 gives Ka = 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.
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.
1
Clean the metal strips with sandpaper to remove oxide coatings.
2
Fill each beaker about two-thirds with the metal ion solution.
3
Dip filter paper in saturated KNO3 and place it between the beakers with both ends immersed.
4
Connect the metals to the voltmeter and record the steady reading.
5
Swap the terminals if the reading is negative.

Section 7 of 7Exam technique

1
Not recording the quench time in 9a — the reaction stops when hydrogencarbonate is added, so that time is used.
2
Using temperature in °C for the Arrhenius plot — T must be in kelvin.
3
Using 1/t as time directly — the rate is proportional to 1/t.
4
Using a low-resistance voltmeter for cell EMF — a high-resistance one stops current flowing and changing the cell.

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