Physical Chemistry Practicals (AQA A Level Chemistry): Flashcards

A high-resistance voltmeter draws negligible current from the cell, allowing a true EMF reading to be obtained rather than a terminal voltage reduced by internal resistance.

The EMF of the Zn/Zn²⁺ || Cu²⁺/Cu cell under standard conditions is +1.10 V. A higher EMF indicates a greater difference in standard electrode potential (E^ꝋ) between the two half-cells.

False.

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

Swap the voltmeter terminals around. A negative reading means the connections are reversed relative to the actual polarity of the cell.

A device consisting of two half-cells connected by a salt bridge and an external circuit, in which a spontaneous redox reaction drives the flow of electrons and produces a measurable EMF.

An elimination reaction is a reaction in which atoms or groups are removed from a molecule to form a double bond, typically producing an alkene and a small molecule such as HBr or H₂O.

In a pH titration curve experiment, a pH probe is used alongside a magnetic stirrer. The pH is recorded every 2 cm³ of NaOH until close to the equivalence point, when 1 cm³ portions are used.

True.

Phenolphthalein's pH range falls within the steep vertical section of the weak acid/strong base titration curve, so it changes colour sharply at the equivalence point.

Find the volume at the equivalence point, halve it, and read the corresponding pH from the curve. At the half-equivalence point, pH = pK_a of the weak acid.

The buffer region on a pH titration curve is identified as the flattest, most horizontal section of the graph, where pH changes very little on addition of acid or base.

False.

The indicator's pH range must fall within the steep vertical section of the curve. This ensures a sharp colour change exactly at the equivalence point.

pH values must be given to exactly 2 decimal places. Examiners penalise answers given to 1 d.p. or more than 2 d.p.

Entropy of vaporisation is the entropy change per mole when a liquid vaporises at its boiling point. It can be calculated from ΔS^ꝋ = ΔH^ꝋ / T, where T is the boiling point in kelvin.

At the boiling point of water, liquid and vapour are in equilibrium, so ΔG^ꝋ = 0. Rearranging the Gibbs equation gives ΔS^ꝋ = ΔH^ꝋ / T.

False.

After initial boiling, the kettle's automatic cut-off is depressed to re-boil the water for 100 s. The mass lost by evaporation is found by re-weighing the kettle before and after this timed boiling period.

Energy supplied = 3 × 100 = 300 kJ

Moles evaporated = 130 ÷ 18 = 7.22 mol

ΔH_vap = 300 ÷ 7.22 = 41.55 kJ mol^-1

Using an enthalpy of vaporisation of 41 550 J mol^-1 and a boiling point of 373 K, the entropy of vaporisation of water = 111.4 J mol^-1 K^-1.

False.

The actual volume used does not matter greatly, as long as it covers the heating element. The key measurement is the mass of water lost by evaporation during the timed boiling period.

The kettle itself plus around 1 dm³ of water (approximately 1 kg) must both be weighed together. A balance with insufficient capacity cannot measure the combined mass.

A clock reaction is a reaction that shows a sudden, dramatic colour change after a set period of time has elapsed, making it ideal for studying kinetics by timing the reaction.

In the iodine clock reaction, sodium thiosulfate is added to use up the iodine as it forms. Once the thiosulfate is consumed, any excess iodine reacts with starch to produce a blue-black colour.

False.

The rate is calculated as the reciprocal of the time (rate = 1/t). This approximates the initial rate and allows a rate–concentration graph to be plotted.

First order with respect to KI. If doubling [KI] doubles the rate (halves the time), the rate is directly proportional to [KI], confirming first order dependence.

In the iodine clock experiment, solutions are measured using burettes for accuracy. Distilled water is included in the mixture to keep the total volume constant while varying the concentration of KI.

False.

The iodine produced is in excess compared to the sodium thiosulfate. The colour change appears as soon as the thiosulfate is used up, so a precisely timed end point is observed.

A graph of rate (1/t) against concentration of KI is plotted. A straight line through the origin confirms that the reaction is first order with respect to KI.

The blank contains only sulfuric acid catalyst and water (no reactants). Its titre against NaOH accounts for the acid catalyst so it can be subtracted from the reaction mixture titre, giving the moles of ethanoic acid at equilibrium.

In the Kc esterification experiment, 6.0 g of ethanoic acid and 6.2 g of ethanol are used (each equivalent to 0.1 mol). The equilibrium is established over approximately one week(s), after which the mixture is titrated against 1.0 mol dm^-3 NaOH.

True.

When all species are in the same volume, the concentration ratios equal the mole ratios. However, this should always be checked since it depends on the equation and may not always cancel.

NaOH is a strong base, and the equivalence point of the titration against ethanoic acid is above pH 7. Phenolphthalein (pH range 8.3–10) changes colour in this alkaline range. Methyl orange (pH 3.1–4.4) changes colour in acidic solution and would give an inaccurate end point.

In the Kc esterification, reactants are weighed directly into the flask rather than measured by volume because this is more accurate and avoids loss of liquid as drops in measuring cylinders.

False.

Ethanoic acid has a very strong pungent smell and is corrosive. It must be handled in a fume cupboard, and safety spectacles and gloves should be worn.

Multiple flasks allow several independent titrations to be carried out. An average value for K_c can then be calculated, improving the reliability of the result.

A value that expresses the ratio of equilibrium concentrations of products to reactants for a reversible reaction at a given temperature, with each concentration raised to the power of its stoichiometric coefficient.

A structural formula is a representation that shows how atoms are bonded together in a molecule, giving enough detail to identify the arrangement of atoms without drawing every bond explicitly.

In simple cell experiments, metal foil electrodes are cleaned with sandpaper and rinsed, then immersed in solutions of their own ions. The cell is completed using a salt bridge soaked in saturated potassium chloride or nitrate.

False.

Measured results are generally lower than calculated values because the experiment uses non-standard conditions (not 1 mol dm^-3, not 298 K), whereas standard electrode potentials are defined under standard conditions.

Cell EMF decreases as the concentration of a solution decreases. By Le Chatelier's principle (or the Nernst equation), diluting a half-cell solution reduces the concentration of the ions involved in the electrode reaction, shifting the equilibrium to oppose the change. This lowers the electrode potential of that half-cell, reducing the overall cell EMF.

In the changing conditions experiment, one solution is diluted by taking 25 cm³ and diluting to 250 cm³ with distilled water. This is repeated for a total of five dilutions.

False.

Metals with a sufficiently large difference in electrode potential should be chosen, so the range of EMF results is meaningful and the effect of dilution on cell EMF can be clearly observed.

Temperature and electrode spacing (depth of immersion) can both be investigated using the same equipment to see their effect on cell EMF.

A skeletal formula is a simplified structural representation showing the carbon skeleton as a zigzag line, with carbon atoms implied at each vertex and end, and hydrogen atoms omitted unless attached to a heteroatom.

To find K_a experimentally, 25 cm³ of 0.1 mol dm^-3 ethanoic acid is titrated with NaOH until the indicator turns pink. A further 25 cm³ of the same acid is then added, and the pH at this half-equivalence point is measured.

True.

At the half-equivalence point, [acid] = [conjugate base], so the K_a expression simplifies to K_a = [H⁺]. The pH at this point therefore equals the pK_a.

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

Therefore K_a = 1.8 × 10⁻⁵ mol dm^-3

Also: pK_a = pH at the half-equivalence point = 4.75

The pH probe is calibrated before use using three buffer solutions: pH 4.00, pH 7.00 and pH 9.20. This is known as a three-point calibration.

True.

The pipette is read only once (one measurement), giving a single uncertainty. The burette requires an initial and a final reading, giving double the uncertainty.

The second 25 cm³ portion of acid has the same number of moles as the NaOH used to neutralise the first portion. Half the acid has been converted to its conjugate base, so [acid] = [salt], creating the half-equivalence condition.

A displayed formula is a structural representation showing every atom and every bond in a molecule, with each bond drawn as a line between the bonded atoms.

In the titration curves practical, 25 cm³ of acid is measured using a volumetric pipette. Alkali is added in 5 cm³ portions and the pH is recorded after each addition until 50 cm³ total has been added.

False.

The four curves differ in starting pH, the pH at the equivalence point, and the size of the vertical section. Strong acid/strong base gives the most pronounced S-shape; weak acid/weak base gives no sharp equivalence point.

The pH changes very rapidly near the equivalence point. Smaller additions give more data points in the steep region, making it easier to draw the shape of the curve accurately and locate the equivalence point precisely.

The four acid–base combinations investigated are: strong acid + strong base, strong acid + weak base, weak acid + strong base, and weak acid + weak base.

True.

When a strong acid and strong base are fully neutralised, only water and a neutral salt are produced, so the equivalence point is at pH 7.

The starting pH and the pH at the equivalence point reveal the type. A weak acid starts at a higher pH than a strong acid of the same concentration. A weak base gives an equivalence point below pH 7; a strong base gives one above pH 7.

It relates the pH of a buffer to the pK_a of the weak acid and the ratio of salt to acid concentrations: pH = pK_a + log₁₀([salt]/[acid]). It allows the correct salt–acid ratio to be calculated for a target pH.

A buffer is made by the direct method by dissolving 2.05 g of sodium ethanoate in 250 cm³ of 0.1 mol dm⁻³ ethanoic acid. When the salt-to-acid ratio is 1:1, the pH equals the pK_a of the acid.

False.

That describes the direct method. The indirect method uses NaOH solution in a burette to partially neutralise the weak acid, producing a mixture of acid and salt in situ.

pH = pK_a + log([salt]/[acid])

4.25 = 4.75 + log([salt]/[acid])

log([salt]/[acid]) = −0.50

[salt]/[acid] = 0.316

After preparing the buffer, its effectiveness is tested by adding 1 cm³ of 1.0 mol dm⁻³ HCl and then 1 cm³ of 1.0 mol dm⁻³ NaOH to separate 100 cm³ samples. The pH should change very little in both cases.

True.

A 250 cm³ volumetric flask is used. The solution is transferred to it with rinsings and made up to the graduation mark with distilled water to give an accurate final volume.

A calibrated pH probe gives a precise numerical reading (to 2 decimal places). Universal indicator gives only an approximate pH range from a colour comparison and lacks the accuracy needed for buffer characterisation.