Chemistry 5070/31 — October/November 2025

Walkthrough

In this practical step, the candidate measures the temperature of the reaction mixture as successive portions of acid Y are added to a fixed volume of alkali X.

To score full marks on this table:

1. Completeness: All 9 rows must be filled with recorded temperatures and calculated temperature changes.
2. Precision: Thermometer readings must be consistently recorded to the nearest $0.5\,^\circ\text{C}$ (i.e. ending in $.0$ or $.5$).
3. Trend: Because neutralisation is exothermic, adding acid causes the temperature to rise initially until all $\text{OH}^-$ is neutralised (the equivalence point). Further addition of cold acid then cools the solution down, so the readings must rise for at least the first three additions before decreasing.
4. Calculations: The temperature change in each row is found by subtracting the initial temperature (at $0.0\text{ cm}^3$) from the measured temperature.
5. Accuracy: The maximum temperature change achieved should be close (within $1.0\,^\circ\text{C}$) to the Supervisor's recorded value.

Key Takeaways

• Always record measurements to the precision specified by the apparatus and instructions (here, $+/- 0.5\,^\circ\text{C}$).
• Thermometric titrations show a characteristic rise to a maximum followed by a decrease.

Common Mistakes

• Omitting the $.0$ on whole-number readings (e.g. writing $24$ instead of $24.0$).
• Incorrect subtraction when calculating the temperature change from the initial value.

Things to Be Careful About

• Ensure the initial temperature is accurately read before any acid is added, as every temperature change depends on this baseline value.

Walkthrough

To construct the thermometric titration graph:

1. Axis label and unit: Label the vertical axis as $\text{temperature change / }^\circ\text{C}$.
2. Scale: Select a sensible linear scale (e.g. 1 major division = $2\,^\circ\text{C}$ or $5\,^\circ\text{C}$) so the plotted data covers more than half of the vertical grid.
3. Plotting: Plot each coordinate $(\text{volume of } \mathbf{Y}, \text{temperature change})$ with a small neat cross or dot within $\pm\frac{1}{2}$ small square.
4. Lines of best fit: Use a ruler to draw a straight line of best fit through the rising points, and a second straight line of best fit through the falling points.
5. Extrapolation: Extend both straight lines until they cross each other. The intersection corresponds to the exact endpoint of the titration.

Key Takeaways

• In thermometric titrations, two intersecting straight lines are used to determine the equivalence point rather than a smooth curved line.
• Graph axes must always include both the quantity and its unit.

Common Mistakes

• Drawing a single continuous curve over the peak instead of two intersecting straight lines.
• Drawing a scale that uses less than half of the available grid height.
• Forgetting to write the unit on the $y$-axis label.

Things to Be Careful About

• Do not force the lines to pass through an anomalous point; aim for a balanced distribution of points on either side of each straight line.

Walkthrough

The point of intersection of the two best-fit lines represents the theoretical maximum temperature change, which occurs at the exact neutralisation point (equivalence point). Project vertically downwards from this intersection to the $x$-axis to read the exact volume of acid Y required.

Key Takeaways

• The intersection of the rising and falling best-fit lines gives the neutralisation volume without requiring an exact measurement at that specific volume.

Common Mistakes

• Reading the $y$-axis value (temperature change) instead of the $x$-axis value (volume of Y).
• Misreading the grid subdivisions on the horizontal axis.

Things to Be Careful About

• Ensure the reading reflects the intersection of the drawn lines, not simply the data point with the highest recorded temperature.

Walkthrough

To find the number of moles of hydroxide ions in the sodium hydroxide solution:

1. Sodium hydroxide ($\text{NaOH}$) dissociates completely to provide one $\text{OH}^-$ ion per formula unit, so the concentration of $\text{OH}^-$ is $1.60\text{ mol / dm}^3$.
2. Convert the volume from $\text{cm}^3$ to $\text{dm}^3$:

1. Multiply concentration by volume in $\text{dm}^3$:

Key Takeaways

• $\text{moles} = \text{concentration (mol / dm}^3\text{)} \times \text{volume (dm}^3\text{)}$.
• Always convert $\text{cm}^3$ to $\text{dm}^3$ by dividing by $1000$.

Common Mistakes

• Forgetting to divide the volume by $1000$, resulting in $40\text{ mol}$.

Things to Be Careful About

• Give the answer to an appropriate number of significant figures (e.g. $0.04$ or $0.0400$).

Walkthrough

1. According to the neutralisation equation $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$, one mole of $\text{H}^+$ reacts exactly with one mole of $\text{OH}^-$.
2. Therefore, the moles of $\text{H}^+$ in the neutralisation volume of Y equals the moles of $\text{OH}^-$ calculated in part (d) ($0.0400\text{ mol}$).
3. Calculate the concentration of $\text{H}^+$ by dividing the number of moles by the volume of Y from part (c) in $\text{dm}^3$:

Key Takeaways

• In a neutralisation reaction between $\text{H}^+$ and $\text{OH}^-$, $\text{moles of } \text{H}^+ = \text{moles of } \text{OH}^-$.
• Concentration is given by $\text{moles} / \text{volume (dm}^3\text{)}$.

Common Mistakes

• Inverting the formula (e.g. dividing volume by moles).
• Forgetting to convert volume from $\text{cm}^3$ to $\text{dm}^3$.

Things to Be Careful About

• Carry forward values from parts (c) and (d) accurately.

Walkthrough

A burette has smaller scale divisions ($0.1\text{ cm}^3$) compared to a standard measuring cylinder ($1\text{ cm}^3$ or $0.5\text{ cm}^3$), allowing volumes of liquid to be delivered and measured with greater precision and smaller percentage uncertainty.

Key Takeaways

• Burettes and volumetric pipettes offer higher precision and lower experimental error than measuring cylinders.

Common Mistakes

• Stating that a burette is "easier to use" or "faster" rather than referring to accuracy/precision.

Things to Be Careful About

• Use clear scientific terminology: "more accurate", "more precise", or "has higher resolution".

Walkthrough

When acid is added to alkali, the neutralisation reaction and subsequent heat transfer throughout the solution and into the thermometer bulb take time. Waiting 30 seconds while stirring ensures that the solution is thoroughly mixed and that the thermometer has reached thermal equilibrium with the liquid, registering the true maximum temperature.

Key Takeaways

• In thermometric experiments, stirring ensures uniform temperature distribution and waiting ensures the maximum temperature is attained before reading.

Common Mistakes

• Claiming it is "to let the solution cool down" (the aim is to measure the maximum temperature before significant heat loss occurs).

Things to Be Careful About

• Clearly state that waiting allows the temperature to reach its maximum or stop changing.

Walkthrough

Universal indicator changes colour depending on pH. Q gives a red colour, which means Q is strongly acidic. There is no need to name the acid yet; just record the colour you see.

Key Takeaways

• Universal indicator is a mixture of dyes that gives a pH colour.
• Red = strongly acidic, purple/blue = alkaline, green = neutral.
• In practical papers, observations must be recorded exactly as seen.

Common Mistakes

• Writing "acid" instead of the colour. The mark is for the observation, not the conclusion.
• Saying "red litmus" – the test uses universal indicator, not litmus.

Things to Be Careful About

• Record the colour of the solution, not the colour of the indicator bottle.
• The answer must be a colour, and the mark scheme accepts "red (solution)".

Walkthrough

Magnesium is a reactive metal. When added to an acid (Q is hydrochloric acid), it reacts to form a salt and hydrogen gas:

The bubbles are effervescence, the magnesium ribbon gets used up, and the reaction is exothermic so the tube warms. To identify the gas, hold a lighted splint at the mouth of the tube: hydrogen burns with a squeaky 'pop'.

Key Takeaways

• Metal + acid → salt + hydrogen.
• Hydrogen is tested with a lighted splint: 'pop' sound.
• The reaction is exothermic.

Common Mistakes

• Writing "bubbles" without "effervescence" – both are accepted, but record what you see.
• Testing with a glowing splint (that is for oxygen); hydrogen needs a lighted splint.
• Forgetting to name the gas as hydrogen.

Things to Be Careful About

• The mark scheme gives any two from the first three observations, then one mark each for the test, the result and the gas name.
• "Colourless solution" is accepted because the magnesium chloride solution is colourless.

Walkthrough

After the magnesium has reacted with the acid, the solution contains magnesium ions. Adding aqueous sodium hydroxide to a solution containing Mg²⁺ gives a white precipitate of magnesium hydroxide:

Magnesium hydroxide is insoluble in excess sodium hydroxide, so the precipitate stays.

Key Takeaways

• Aqueous sodium hydroxide is used to test for cations.
• Mg²⁺ gives a white precipitate that is insoluble in excess NaOH.
• The observation must include both the initial precipitate and its behaviour in excess.

Common Mistakes

• Writing "white precipitate" only – the mark scheme awards one mark for the precipitate and one for "insoluble in excess".
• Confusing with Al³⁺, which gives a white precipitate soluble in excess NaOH.

Things to Be Careful About

• Add the sodium hydroxide drop by drop first; record the change before adding excess.
• "Insoluble in excess" is the key second mark.

Walkthrough

Silver nitrate is used to test for halide ions. The solution from part (b) contains chloride ions. Adding silver nitrate gives a white precipitate of silver chloride:

Dilute nitric acid is added first to remove any carbonate or sulfite ions that could also give precipitates.

Key Takeaways

• $\text{Ag}^+ + \text{Cl}^- \rightarrow \text{AgCl}$, a white precipitate.
• Dilute nitric acid is added before silver nitrate to prevent false positives from carbonate ions.
• Chloride gives a white precipitate; bromide gives cream; iodide gives yellow.

Common Mistakes

• Writing "cream precipitate" – that is for bromide.
• Forgetting the nitric acid step, though the observation mark only needs the precipitate.
• Saying "silver chloride" instead of "white precipitate" – both are fine, but the mark scheme wants the observation.

Things to Be Careful About

• The mark is for the observation, not the name of the precipitate.
• "White precipitate" is the accepted answer.

Walkthrough

Q gives a red colour with universal indicator, so it is an acid. When magnesium is added, hydrogen gas is produced, which confirms the presence of hydrogen ions. The silver nitrate test gives a white precipitate, which identifies chloride ions. Therefore Q is hydrochloric acid, HCl, containing H⁺ and Cl⁻.

Key Takeaways

• An acid contains H⁺ as its cation.
• The white precipitate with silver nitrate identifies Cl⁻.
• The evidence from each test must be combined to name the ions.

Common Mistakes

• Writing "magnesium" as the cation – the magnesium is added to the acid and is not part of Q.
• Writing "hydrochloric acid" instead of the ions – the question asks for the cation and anion.
• Confusing the anion test: chloride gives white, bromide cream, iodide yellow.

Things to Be Careful About

• The cation is H⁺, not Mg²⁺, because Q is the original solution before magnesium is added.
• The mark scheme accepts "hydrogen / H⁺" and "chloride / Cl⁻".

Walkthrough

Aqueous ammonia is a test for cations. With copper(II) ions it first gives a light blue precipitate of copper(II) hydroxide. On adding excess ammonia, the precipitate dissolves to form a dark blue solution containing the complex ion [Cu(NH3)4]2+. This is a characteristic test for Cu²⁺.

Key Takeaways

• Cu²⁺ + ammonia → light blue precipitate, soluble in excess to a dark blue solution.
• This distinguishes Cu²⁺ from Fe²⁺ (green) and Fe³⁺ (brown).
• The observation must include both stages.

Common Mistakes

• Writing "blue precipitate" only – need "light blue" and "dark blue solution in excess".
• Saying "insoluble in excess" – for copper it is soluble.
• Confusing with sodium hydroxide test: with NaOH, Cu²⁺ gives a blue precipitate insoluble in excess.

Things to Be Careful About

• The mark scheme has three points: light blue ppt, soluble in excess, dark blue solution.
• Record the colour change clearly.

Walkthrough

The flame test gave a blue-green colour, which is typical of copper(II). The ammonia test confirmed it: a light blue precipitate that dissolves in excess to a dark blue solution. Both tests together identify the cation as Cu²⁺.

Key Takeaways

• Flame colour and ammonia test are both used to identify Cu²⁺.
• Blue-green flame + light blue ppt soluble in excess ammonia = copper(II).
• The conclusion must be a cation, not a compound.

Common Mistakes

• Writing "copper" instead of "copper(II)" – the charge matters.
• Writing "copper sulfate" – the question asks for the cation.
• Confusing with other blue-flame metals.

Things to Be Careful About

• The mark scheme accepts "Copper(II) / Cu²⁺".
• Use the Roman numeral to show the +2 charge.