Chemistry 5070/32 — October/November 2014

Walkthrough

This part is the practical titration itself. The candidate puts the sodium thiosulfate solution (Q) in the burette and pipettes a known volume of the iodine solution (P) into a flask. The iodine solution is red-brown. As thiosulfate is added from the burette, it reacts with the iodine and the colour fades. When the colour is pale yellow, a few drops of starch indicator are added — starch forms a dark blue complex with iodine, making the end point much sharper. The titration is then continued drop by drop until one drop of thiosulfate makes the blue colour disappear, leaving a colourless solution. This is the end point: all the iodine has reacted.

The volume of Q used in each titration is the final burette reading minus the initial reading. The titration is repeated until at least two results are concordant (within $0.2\text{ cm}^3$ of each other). The best (most concordant) results are ticked, and the average volume of Q is calculated from the ticked values only.

Key Takeaways

• Burette readings are taken to one decimal place (e.g. $25.2\text{ cm}^3$, not $25\text{ cm}^3$), and a zero initial reading is recorded as 0.0.
• The volume of Q used is the final reading minus the initial reading.
• Starch is added near the end point (when the colour is pale yellow) because it gives a sharper end point with iodine.
• Repeating titrations and using concordant results improves the reliability of the average.

Common Mistakes

• Reading the burette to the wrong precision — readings must be to 0.1 cm3 (e.g. 25.0, not 25).
• Forgetting to subtract the initial reading from the final reading.
• Adding the starch indicator too early, when the solution is still dark — the end point is then harder to see.
• Averaging all results instead of only the ticked (best) results.

Things to Be Careful About

• The initial reading is often 0.0 cm3 — record it as 0.0, not 0.
• The end point is when one drop of Q causes the blue colour to disappear.
• The average is taken only over the ticked values.
• The mark scheme awards accuracy marks for how close the titres are to the supervisor's value, and concordance marks for how close the ticked values are to each other.

Walkthrough

The titration tells us how much thiosulfate reacts with the iodine in P. From the equation $2\text{Na}_2\text{S}_2\text{O}_3 + \text{I}_2 \rightarrow \text{Na}_2\text{S}_4\text{O}_6 + 2\text{NaI}$, two moles of thiosulfate react with one mole of iodine. So the moles of thiosulfate used are found first: (volume in dm3) × concentration. Then halve this to get the moles of iodine. Finally divide by the volume of P (in dm3) to get the concentration of iodine.

Using the mark scheme's example (titre 25.2 cm3, pipette 25.0 cm3):

Key Takeaways

• The stoichiometric ratio from the balanced equation is essential: 2 mol thiosulfate to 1 mol iodine.
• Concentration = moles / volume, with the volume in dm3.
• Volumes in cm3 must be converted to dm3 by dividing by 1000.

Common Mistakes

• Forgetting the 2:1 stoichiometric ratio and not dividing the moles of thiosulfate by 2.
• Using the volume in cm3 instead of dm3 in the concentration formula.
• Using the titre volume instead of the pipette volume when calculating the concentration of iodine.

Things to Be Careful About

• The mark scheme gives the expression (titre × 0.100)/(2 × 25.0) — the 2 is the stoichiometric factor from the equation.
• The candidate uses their own titre value from part (a).
• The unit is mol/dm3.

Walkthrough

From the first equation, $\text{Ca(ClO)}_2 + 4\text{I}^- + 4\text{H}^+ \rightarrow \text{CaCl}_2 + 2\text{H}_2\text{O} + 2\text{I}_2$, one mole of calcium hypochlorite produces two moles of iodine. So the moles of calcium hypochlorite required to produce the iodine in 1 dm3 of P is half the concentration of iodine found in (b).

Using the example: $\frac{0.0504}{2} = 0.0252\text{ mol}$.

Key Takeaways

• The stoichiometric ratio from the equation: 1 mol Ca(ClO)2 to 2 mol I2.
• Moles of Ca(ClO)2 = moles of I2 / 2.

Common Mistakes

• Dividing by the wrong factor (e.g. using a 1:1 ratio instead of 1:2).
• Confusing the two equations in the question.

Things to Be Careful About

• This uses the answer from (b) — the mark scheme applies error carried forward (ecf), so a wrong (b) value still earns the mark here if the ratio is applied correctly.

Walkthrough

Convert the moles of calcium hypochlorite to a mass using $\text{mass} = \text{moles} \times M_r$. The relative formula mass of $\text{Ca(ClO)}_2$ is given as 143.

Using the example: $\text{mass} = 0.0252 \times 143 = 3.60\text{ g}$.

This mass was present in 10.0 g of bleaching powder, so the percentage by mass is $\frac{3.60}{10.0} \times 100 = 36.0\%$.

Key Takeaways

• mass = moles × Mr.
• Percentage by mass = (mass of component / total mass) × 100%.

Common Mistakes

• Using the wrong Mr.
• Forgetting to multiply by 100 for the percentage.
• Using the total mass of the powder instead of the mass of the component.

Things to Be Careful About

• The Mr of Ca(ClO)2 is given as 143 — use it directly.
• The percentage is out of 10.0 g of powder.
• Units: g for mass, % for percentage.