Chemistry 5070/32 — October/November 2024

Walkthrough

This part is about taking and recording burette readings correctly. Each experiment gives a final and an initial burette reading. The volume of acid X used is found by subtracting the initial reading from the final reading. All readings and calculated volumes must be recorded to 1 decimal place, so a reading of exactly 25 cm3 must be written as 25.0, and an initial reading of exactly zero must be written as 0.0. The four experiments give four volumes; these should be close to one another and to the supervisor's value.

Key Takeaways

• Burette readings are recorded to 1 decimal place.
• Volume used = final reading - initial reading.
• Repeating the titration allows a reliable average to be calculated.

Common Mistakes

• Writing readings without a decimal place, such as 25 instead of 25.0.
• Writing 0 instead of 0.0 for an initial reading.
• Forgetting to calculate the volume used for each experiment.
• Recording readings in the wrong row of the table.

Things to Be Careful About

• Keep the same precision (1 decimal place) for every reading and every calculated volume.
• Use the correct units, cm3, for all volumes.
• Make sure the final reading is larger than the initial reading for each experiment.

Walkthrough

Look at the four volumes of X used. The two best results are the two that agree most closely with each other. Tick exactly these two in Table 1.1. The explanation should state that the two chosen volumes are concordant, meaning they are within 0.2 cm3 of each other. This shows the results are reliable and can be averaged.

Key Takeaways

• Concordant results are those that agree closely.
• Only the two best results should be used to calculate the average.

Common Mistakes

• Ticking three results instead of two.
• Ticking an anomalous result that is far from the others.
• Giving no explanation for the choice.

Things to Be Careful About

• The mark scheme requires exactly two ticks.
• The explanation must mention concordance or the 0.2 cm3 agreement.

Walkthrough

Use only the two volumes that were ticked in part (b). Add them together and divide by 2. This gives the average volume of X needed to neutralise 25.0 cm3 of Y. Do not include the other two volumes, because they are less reliable.

Key Takeaways

• The average is calculated from the concordant results only.
• The mean gives a more reliable value than any single titration.

Common Mistakes

• Averaging all four volumes instead of only the two ticked ones.
• Including an anomalous result in the average.
• Forgetting to divide by 2.

Things to Be Careful About

• Use exactly the two values ticked in part (b).
• Give the answer in cm3.

Walkthrough

The concentration of X is 0.500 mol / dm3. To find moles, the volume must be in dm3, so divide the average volume in cm3 by 1000. Then multiply by the concentration. This gives the number of moles of HCl in the average volume used.

Key Takeaways

• Moles = concentration × volume (in dm3).
• To convert cm3 to dm3, divide by 1000.

Common Mistakes

• Forgetting to divide the volume by 1000.
• Using the wrong concentration.
• Mixing up cm3 and dm3.

Things to Be Careful About

• The volume must be in dm3 before multiplying by the concentration.
• The answer is in moles.

Walkthrough

The balanced equation shows that 2 mol of HCl react with 1 mol of Na2CO3. So the moles of Na2CO3 in the 25.0 cm3 sample are half the moles of HCl found in part (d). The concentration is then moles divided by volume in dm3; 25.0 cm3 is 25.0/1000 dm3. Finally, round the answer to 3 significant figures.

Key Takeaways

• Use the mole ratio from the balanced equation.
• Concentration = moles ÷ volume in dm3.
• Round only at the end, to 3 significant figures.

Common Mistakes

• Forgetting to divide the moles of HCl by 2.
• Using 25.0 cm3 without converting to dm3.
• Rounding before the final step.

Things to Be Careful About

• The volume of Y is 25.0 cm3, not 150 cm3.
• The answer must be given to 3 significant figures with units mol / dm3.

Walkthrough

First calculate the relative molecular mass of sodium carbonate: 2 × 23 for sodium, plus 12 for carbon, plus 3 × 16 for oxygen, giving 106. Then find the number of moles of Na2CO3 in 150 cm3 of Y by multiplying the concentration from part (e) by 150/1000 (since 150 cm3 = 0.150 dm3). Finally multiply the moles by the M_r to get the mass in grams.

Key Takeaways

• M_r is found by adding the relative atomic masses of all atoms in the formula.
• Moles = concentration × volume in dm3.
• Mass = moles × M_r.

Common Mistakes

• Using 25.0 cm3 instead of 150 cm3.
• Using the wrong M_r for Na2CO3.
• Forgetting to convert 150 cm3 to dm3.

Things to Be Careful About

• 150 cm3 = 0.150 dm3.
• The final answer must be in grams.

Walkthrough

Swirling the conical flask while acid is added mixes the reactants thoroughly. This ensures the acid and carbonate react evenly and that the colour change at the end-point is reliable. Without swirling, local excess of acid could give a false end-point.

Key Takeaways

• Swirling ensures thorough mixing during a titration.
• It helps give an accurate end-point.

Common Mistakes

• Saying swirling cools the mixture.
• Saying swirling makes the reaction faster without mentioning mixing.

Things to Be Careful About

• The mark scheme accepts 'to mix the contents thoroughly/evenly'.

Walkthrough

Near the end-point the colour change happens with a single drop. Adding the acid drop by drop lets the candidate stop exactly when the colour just changes, so the end-point is not overshot. This gives an accurate volume of acid used.

Key Takeaways

• Dropwise addition near the end-point prevents overshooting.
• The end-point is the first permanent colour change.

Common Mistakes

• Saying dropwise addition makes the titration faster.
• Saying it cools the solution.

Things to Be Careful About

• The mark scheme accepts 'to ensure the end-point is not exceeded' or 'so the volume is accurate'.

Walkthrough

When conducting a flame test on an unknown sample containing sodium ions ($\text{Na}^+$), the characteristic persistent yellow (or yellow-orange) flame is observed.

Key Takeaways

• $\text{Na}^+$ gives a yellow flame in a flame test.
• Other common flame test colours to recall at O Level: $\text{Li}^+$ (red), $\text{K}^+$ (lilac), $\text{Ca}^{2+}$ (orange-red), $\text{Cu}^{2+}$ (blue-green), $\text{Ba}^{2+}$ (apple-green).

Common Mistakes

• Confusing the yellow flame of sodium with the lilac flame of potassium or orange-red of calcium.

Things to Be Careful About

• Ensure the flame colour recorded matches the standard syllabus descriptors.

Walkthrough

A yellow flame test result indicates the presence of sodium ions ($\text{Na}^+$).

Key Takeaways

• Cations can be deduced directly from their characteristic flame colours.

Common Mistakes

• Writing an incorrect charge for the ion (e.g. $\text{Na}^{2+}$).

Things to Be Careful About

• Either the full name 'sodium' or the correct chemical formula '$\text{Na}^+$' is accepted.

Walkthrough

Adding aqueous sodium hydroxide and aluminium foil to a solution followed by gentle warming is the standard qualitative test for nitrate ions ($\text{NO}_3^-$). In this reaction, nitrate ions are reduced by aluminium under alkaline conditions to produce ammonia gas ($\text{NH}_3$).
To earn all 3 marks:

1. State the test reagent/medium used: damp red litmus paper.
2. State the observed result: turns blue.
3. Identify the evolved gas: ammonia ($\text{NH}_3$).

Key Takeaways

• Nitrate test: warm with aqueous $\text{NaOH}$ and $\text{Al}$ foil $\rightarrow$ ammonia gas produced.
• Ammonia test: turns damp red litmus paper blue.

Common Mistakes

• Omitting that the litmus paper must be damp or moist.
• Naming the anion (nitrate) here instead of describing the gas test, observation, and the gas itself as prompted by the stem.

Things to Be Careful About

• The stem explicitly reminds candidates to test and identify any gases evolved and describe the test used.

Walkthrough

The production of ammonia upon heating with aqueous sodium hydroxide and aluminium foil confirms the presence of nitrate ions ($\text{NO}_3^-$).

Key Takeaways

• $\text{Al} + \text{NaOH}_{(\text{aq})} + \text{heat}$ reducing an unknown to $\text{NH}_3$ is the specific qualitative test for nitrate ($\text{NO}_3^-$).

Common Mistakes

• Writing nitrite instead of nitrate.

Things to Be Careful About

• Both the name 'nitrate' and formula '$\text{NO}_3^-$' are acceptable.

Walkthrough

Many gases produced in chemical reactions (such as chlorine, sulfur dioxide, or ammonia) are toxic, irritant, or harmful when inhaled in significant amounts. Therefore, chemical and indicator tests must always be used instead of smelling unknown gases.

Key Takeaways

• Smelling unknown gases poses severe health and safety risks due to potential toxicity and irritation.

Common Mistakes

• Vague answers like 'it smells bad' or 'it is dangerous' without specifying toxicity/poisonous nature or respiratory irritation.

Things to Be Careful About

• Keep the answer crisp and focused on safety/toxicity.

Walkthrough

For a solution containing chromium(III) ions ($\text{Cr}^{3+}$):

1. On dropwise addition of aqueous sodium hydroxide, a grey-green (or green) precipitate of chromium(III) hydroxide, $\text{Cr}(\text{OH})_3$, forms.
2. In excess aqueous sodium hydroxide, the amphoteric precipitate dissolves to form a green solution containing chromate(III) / tetrahydroxochromate(III) complex ions.

Key Takeaways

• $\text{Cr}^{3+}$ with $\text{NaOH}_{(\text{aq})}$: green precipitate $\rightarrow$ soluble in excess $\rightarrow$ green solution.
• Compare with $\text{Fe}^{2+}$ which forms a green precipitate insoluble in excess.

Common Mistakes

• Forgetting to state the colour of the solution formed in excess.
• Confusing $\text{Cr}^{3+}$ with $\text{Fe}^{2+}$.

Things to Be Careful About

• All three points (precipitate colour, solubility in excess, and colour of final solution) must be stated clearly.

Walkthrough

When aqueous ammonia is added to a solution containing $\text{Cr}^{3+}$:

1. A green (or grey-green) precipitate of $\text{Cr}(\text{OH})_3$ forms.
2. On adding excess aqueous ammonia, the precipitate remains insoluble (or is only very slightly soluble under specific conditions, but standard 5070 observation is insoluble in excess).

Key Takeaways

• $\text{Cr}^{3+}$ with $\text{NH}_{3(\text{aq})}$: green precipitate, insoluble in excess.

Common Mistakes

• Stating that it dissolves in excess ammonia.

Things to Be Careful About

• Distinguish clearly between behaviour with $\text{NaOH}$ (soluble in excess) and with $\text{NH}_3$ (insoluble in excess).

Walkthrough

1. Adding dilute nitric acid to a solution containing chloride ions results in no observable change (no effervescence or precipitation, confirming the absence of carbonate or sulfite ions).
2. Subsequent addition of aqueous silver nitrate results in a precipitation reaction between silver ions and chloride ions:

forming an insoluble white precipitate of silver chloride ($\text{AgCl}$).

Key Takeaways

• Acidification with dilute $\text{HNO}_3$ prevents false positives from carbonates.
• $\text{AgNO}_{3(\text{aq})}$ gives a white precipitate with $\text{Cl}^-$, cream with $\text{Br}^-$, and yellow with $\text{I}^-$.

Common Mistakes

• Leaving the first observation line blank instead of recording 'no change'.
• Writing 'white solution' instead of 'white precipitate'.

Things to Be Careful About

• Ensure 'precipitate' is explicitly written for the silver nitrate observation.

Walkthrough

• Cation identification: In parts (f) and (g), solution Q gave a green precipitate with aqueous $\text{NaOH}$ that dissolved in excess to a green solution, and a green precipitate with aqueous $\text{NH}_3$ that was insoluble in excess. These observations uniquely identify the cation as chromium(III) ($\text{Cr}^{3+}$).
• Anion identification: In part (h), solution Q acidified with dilute nitric acid gave a white precipitate with aqueous silver nitrate, confirming the anion is chloride ($\text{Cl}^-$).

Key Takeaways

• Green ppt soluble in excess $\text{NaOH}$ (green solution) and insoluble in excess $\text{NH}_3$ $\rightarrow$ $\text{Cr}^{3+}$.
• White ppt with acidified $\text{AgNO}_3$ $\rightarrow$ $\text{Cl}^-$.

Common Mistakes

• Omitting the oxidation state for the cation (writing just 'chromium' or 'iron(II)').
• Writing 'chlorine' instead of 'chloride'.

Things to Be Careful About

• State the oxidation state for transition metals with multiple oxidation states: chromium(III) / $\text{Cr}^{3+}$.