Advanced Physical Chemistry Core Practicals (Edexcel International A Level (IAL) Chemistry): Flashcards

1. H₂O₂ (aq) + 2I^- (aq) + 2H⁺ (aq) → I₂ (aq) + 2H₂O (l)

2. 2S₂O₃^2- (aq) + I₂ (aq) → 2I^- (aq) + S₄O₆^2- (aq)

Sodium thiosulfate consumes the iodine produced until it is used up, then the excess iodine turns the starch blue-black.

In the iodine clock reaction, the timer is stopped when the solution turns blue-black due to starch reacting with excess iodine.

The sodium thiosulfate consumes all iodine produced until it is completely used up. The moment the thiosulfate is exhausted, any further iodine produced reacts immediately with starch, giving an abrupt blue-black colour change.

The rate of reaction is calculated by taking the reciprocal of time (1/t). Each recorded time gives a rate value that can be plotted against concentration to show the order of reaction.

False.

The sulfuric acid is measured in a measuring cylinder as it is in excess. The other solutions, including potassium iodide, are measured in burettes.

The graph shows a directly proportional relationship: as concentration doubles, the rate doubles. This confirms that the reaction is first order with respect to potassium iodide.

In the iodine clock experiment, the iodine produced is in excess compared to the thiosulfate, which means it is only when the thiosulfate is used up that the colour change occurs.

It provides a sharp, easily timed endpoint (the sudden blue-black colour), avoiding the need for continuous sampling. The time to colour change directly reflects the rate of the reaction under the initial conditions.

The Arrhenius plot method involves plotting ln k against 1/T (in K^-1) to give a straight line. The gradient of that line equals -Ea/R, from which the activation energy Ea can be calculated.

The reaction between bromide ions and bromate(V) ions in the presence of phenol and an indicator. Bromine produced bleaches the indicator, and the time for bleaching is recorded at each temperature.

In Core Practical 10, the gradient of the Arrhenius plot equals -Ea/R, where R is the gas constant and Ea is the activation energy.

Methyl red (or methyl orange) acts as the indicator. The end of the reaction is signalled when the indicator decolourises, which occurs once the phenol is consumed and excess bromine bleaches the dye.

The rate is assumed proportional to 1/t, so k is proportional to 1/t at each temperature. The ln k values (or ln(1/t) values) are then plotted against 1/T to construct the Arrhenius plot.

True.

The x-axis of the Arrhenius plot is 1/T where T must be in Kelvin. Using Celsius would give incorrect gradient and hence an incorrect Ea.

To ensure both reactant solutions reach thermal equilibrium at the water bath temperature before mixing, so that the temperature at the start of the reaction is accurately known and controlled.

To achieve the lowest temperature in Core Practical 10, ice is added to the water bath.

Activation energy is the minimum energy that colliding particles must possess for a reaction to occur. It is determined from the gradient of an Arrhenius plot using Ea = -gradient × R.

The half-equivalence point is the point in a titration at which exactly half the acid has been neutralised. At this point, Ka = [H⁺], so pH = pKa.

At the half-equivalence point, the concentrations of the weak acid and its conjugate base are equal, so they cancel in the Ka expression, leaving Ka = [H⁺].

In Core Practical 11, the half-equivalence point is reached by first titrating 25 cm³ of ethanoic acid to the endpoint, then adding a further 25 cm³ of acid.

The [H⁺] is found using [H⁺] = 10^-pH, and since Ka = [H⁺] at the half-equivalence point, Ka equals that value in mol dm^-3.

Phenolphthalein is used. The endpoint is marked by the solution turning pink, indicating that all the ethanoic acid in the first 25 cm³ portion has been neutralised by sodium hydroxide.

True.

At the half-equivalence point, Ka = [H⁺], and taking the negative logarithm of both sides gives pKa = pH.

Before use in Core Practical 11, the pH meter or probe must be calibrated to ensure accurate readings.

Ka = [H⁺] = 10^-4.75 = 1.8 × 10^-5 mol dm^-3.

1. Measurements made by the pipette (single reading uncertainty)

2. Measurements made by the burette (double uncertainty: initial and final readings)

Judging the endpoint of the titration also contributes uncertainty.

A salt bridge is a strip of filter paper soaked in saturated potassium nitrate solution. It maintains electrical contact between the two half-cells while preventing direct mixing of the half-cell solutions.

Metal strips must be cleaned (e.g. with sandpaper) to remove oxide coatings that would act as a barrier to electron flow and give inaccurate or inconsistent EMF readings.

E^ꝋ(cell) = E^ꝋ(positive electrode) - E^ꝋ(negative electrode)

The half-cell with the more positive E^ꝋ value is the positive electrode.

In Core Practical 12, if the voltmeter gives a negative reading, you should swap the terminals around to obtain a positive value.

E^ꝋ(cell) = E^ꝋ(positive) - E^ꝋ(negative) = (+0.34) - (-0.76) = +1.10 V.

Copper is the positive electrode as it has the more positive E^ꝋ value.

False.

The salt bridge must be changed each time to prevent cross-contamination of ions between half-cells, which would affect the measured EMF.

Standard electrode potentials apply under standard conditions (1.0 mol dm^-3, 298 K), which are difficult to achieve precisely in a school laboratory. However, the relative EMF values between cells should still match theoretical predictions.

A high resistance voltmeter is used to measure EMF in Core Practical 12 to prevent current flow which would alter the cell voltage.

Standard electrode potentials are voltage values assigned to half-cell reactions measured under standard conditions (1.0 mol dm^-3, 298 K, 100 kPa). They are used to calculate the theoretical EMF of electrochemical cells.