Practical Skills 1: Paper 3

Practical Weightage: Paper 3 Practical P3

“Paper 3 is a test of what your hands and eyes can do, and of whether you can show it on paper. Marks go to the careful reading, the clear table, the working shown line by line and the honest comment on error. None of these needs new chemistry, only habits you practise until they are automatic.”

— SCORECHEM ACADEMIC TEAM

0. What Examiners Want

This guide covers the practical skills tested in Paper 3 (Advanced Practical Skills) of Cambridge International AS & A Level Chemistry 9701 (2025–27). The practical section of the syllabus lists the expectations for each skill rather than numbered learning outcomes, so the table below gives them word for word, grouped by skill. The right-hand column shows where each is taught. Paper 5 (planning, analysis and evaluation) is covered in the next unit.

Paper 3 skill expectations. Candidates should be able to:	Section
Manipulation, measurement and observation (at least 12 marks)
Successful collection of data and observations. Candidates should be able to: set up apparatus; follow instructions given in the form of written instructions or diagrams; use apparatus to collect an appropriate quantity of data; make observations, including subtle differences in colour, solubility or quantity of materials; make measurements using pipettes, burettes, measuring cylinders, thermometers and other common laboratory apparatus; candidates should record burette readings to the nearest 0.05 cm³ and, when using a thermometer calibrated at 1 °C intervals, temperature readings should be recorded to the nearest 0.5 °C	1, 2, 3
Quality of measurements or observations. Candidates should be able to: make accurate and consistent measurements and observations, including achieving concordant titres (two titres within 0.10 cm³ of each other) and precise colour descriptions	2, 5
Decisions relating to measurements or observations. Candidates should be able to: decide how many tests or observations to perform; make measurements that span a range and have a distribution appropriate to the experiment; identify where repeated readings or observations are appropriate; replicate readings or observations as necessary, including where an anomaly is suspected; identify where confirmatory tests are appropriate and the nature of such tests; select reagents to distinguish between given ions	2, 3, 5
Presentation of data and observations (at least 6 marks)
Recording data and observations. Candidates should be able to: present numerical data, values or observations in a single table of results with headings and units that conform to accepted scientific conventions, i.e. volume / cm³, volume (cm³) or volume in cm³; record raw readings of a quantity to the same degree of precision; record observations to the same level of detail, using simple language such as ‘blue’ or ‘yellow’ and, where fine discrimination is required, terms such as ‘pale’ or ‘dark’	1, 2, 4, 5
Display of calculation and reasoning. Candidates should be able to: show working in calculations and key steps in reasoning; use the correct number of significant figures for calculated quantities (the same as, or one more than, the smallest number of significant figures in the provided or experimentally determined data)	2, 3
Data layout. Candidates should be able to: present data in a single table of results; draw an appropriate table in advance of taking readings; record all data in the table; plot appropriate variables on appropriate, clearly labelled x- and y-axes; choose scales such as 1, 2 or 5 units to a 20 mm square so that the points occupy at least half of the grid in both directions; plot all points using a cross × or circled dot to an appropriate accuracy; draw straight lines or smooth curves of best fit with an even distribution of points on either side, identifying anomalous points	4
Analysis, conclusions and evaluation (at least 10 marks)
Interpretation of data or observations. Candidates should be able to: describe the patterns and trends shown by data in tables and graphs; describe and summarise the key points of a set of observations; calculate quantities from data, or calculate the mean from repeated values; find an unknown value by using coordinates, a point of intersection or intercepts on a graph; determine the gradient of a straight-line graph, using two points that are more than half of the length of the axes apart; extrapolate the line of a graph	2, 3, 4
Drawing conclusions. Candidates should be able to: draw conclusions from an experiment, considering whether experimental data support a given hypothesis, and making further predictions; draw conclusions from interpretations of observations, data and calculated values; make scientific explanations of data, observations and conclusions that they have described	3, 5, 6
Identifying sources of error and suggesting improvements. Candidates should be able to: evaluate the effectiveness of control variables; comment on errors intrinsic in measuring devices and in the limitations of the method; show an understanding of the distinction between systematic errors and random errors; identify the most significant sources of error in an experiment; state the uncertainty in a quantitative measurement and express it as an actual or percentage error; suggest realistic modifications that will improve the accuracy of the experiment or the observations; suggest ways in which to extend the investigation to answer a new question	6
Practical procedures
Quantitative analysis: titration experiments (acid–alkali with the indicators listed in the syllabus; potassium manganate(VII) with hydrogen peroxide, iron(II) ions or ethanedioic acid or its salts; sodium thiosulfate and iodine); rates experiments; gravimetric experiments; thermometric experiments; gas volume experiments	2, 3
Qualitative analysis: carrying out and interpreting tests for the cations, anions and gases in the Qualitative analysis notes, and the organic tests for aldehydes, CH₃CO– or CH₃CH(OH)– groups and oxidisable compounds	5

How marks are earned in Paper 3

Measurement: readings to the precision of the instrument, results close to the Supervisor’s values, and concordant titres.
Observation: every test recorded in plain words, including ‘no change’, with fine colour differences described.
Presentation: a table drawn before the readings, with headings and units, and a graph with labelled axes, a sensible scale and a ruled line.
Calculation: each step shown, the right number of significant figures, and units on every answer.
Conclusion: an identification or a decision that follows from the data, supported by one piece of evidence.
Evaluation: the main source of error named precisely, its size as a percentage where asked, and a realistic improvement.

⚠️ Examiner Trap: Where the marks went. In 2024 (9701/31) the quantitative questions were a thermometric experiment leading to an enthalpy change by Hess’s law, and an iodine and thiosulfate titration to find the water of crystallisation in sodium sulfite. Marks in questions like these are most often lost through temperatures not recorded to 0.5 °C, titres that are not concordant, answers with the wrong number of significant figures and ‘human error’ given as the only source of error.

1. The Paper and Its Measurements

Paper 3 is a two-hour laboratory examination worth 40 marks. You use real apparatus, record your own results and answer on the question paper. The Qualitative analysis notes, the Periodic Table and the key constants are printed in the paper, so you do not need to memorise them, but you need to know them well enough to use them quickly.

Paper 3 marks: at least 12 for manipulation, 6 for presentation, 10 for analysis and evaluation, 12 shared, with the qualitative and quantitative questions

Paper 3 is 40 marks. Marks are assigned to skills, so a good table and a sound evaluation are worth as much as a good titre.

Planning your time. Read the whole method before you start. Draw each results table before taking readings, so that nothing needs to be copied out later. Start any slow step (such as a solution standing) first. Two or three questions share two hours, so allow about 35 to 40 minutes for each quantitative question and about 40 minutes for the qualitative one.

Reading apparatus. Cambridge expects you to read to the precision of the instrument. The rule is half the smallest division when you must estimate between marks, and the printed precision for a digital instrument such as a balance.

Thermometer read to 0.5 C, measuring cylinder read to 0.5 cm3 and balance read as displayed, with a rules box

Estimate between marks to half a division. A thermometer with 1 °C marks is read to 0.5 °C; a cylinder with 1 cm³ marks is read to 0.5 cm³.

A burette is marked every 0.1 cm³. You read the bottom of the meniscus with your eye level with it, and record to the nearest 0.05 cm³, which means the second decimal place is always 0 or 5.

Three burette readings of 23.45, 24.10 and 22.85 cm3 with the meniscus at eye level

Read the bottom of the meniscus at eye level. A reading on a line is written with a final 0 (24.10), between lines with 5 (23.45).

Instrument	Smallest division	Record to	Maximum uncertainty of one reading
burette	0.1 cm³	0.05 cm³ (e.g. 23.45)	±0.05 cm³
thermometer, 1 °C marks	1 °C	0.5 °C (e.g. 27.5)	±0.5 °C
measuring cylinder, 1 cm³ marks	1 cm³	0.5 cm³	±0.5 cm³
balance reading to 0.01 g	0.01 g	0.01 g (e.g. 5.20)	±0.005 g
stopwatch	as displayed	as displayed (e.g. 47 s)	depends on reaction timing

⚠️ Examiner Trap: Record, then compare. A reading of 24.1 is not wrong in value but is wrong in precision. Write 24.10. Every burette reading in the table must show two decimal places, every thermometer reading one decimal place ending in 0 or 5, and every mass the same number of decimal places as the balance shows. A titre is a calculated value, so it is also written to two decimal places.

What is given to you. The method may refer to solutions as FA 1, FA 2 and so on. Treat every unknown with caution, add only the volume specified and record all observations. If you cannot set up the apparatus, the Supervisor may help, but marks for collecting data may then be reduced.

2. Titrations

A titration finds the concentration of one solution by reacting a measured volume of it with a solution of known concentration. Cambridge expects you to be able to carry out acid–alkali titrations, potassium manganate(VII) titrations and sodium thiosulfate with iodine titrations, and to follow the same technique every time.

Titration apparatus: burette on a stand above a conical flask on a white tile, with a pipette, and the rough then accurate sequence

Fill the burette with the titrant, rinse it first, remove the funnel and fill the tip. Swirl the flask throughout; go dropwise near the end-point.

Technique and the marks it earns.

Rinse the burette with the titrant and the pipette with the solution it will deliver. Rinse the conical flask with distilled water only.
Fill the burette above the zero mark, run some out to fill the tip, and record the starting reading.
Use a pipette to measure the solution in the flask (25.0 cm³) and add a few drops of indicator. Place the flask on a white tile.
Carry out a rough titration first to find the approximate end-point. Then repeat accurately, adding the titrant dropwise within about 2 cm³ of the end-point.
Repeat until you have concordant titres: two titres within 0.10 cm³ of each other.

Recording the results. Use a single table with headings and units. Do not include the rough titration in the mean.

A titration results table with rough, 1, 2 and 3, final and initial burette readings and titres, with a mean titre of 23.15

Titre = final reading − initial reading. Choose the two titres that agree to 0.10 cm³ and show which you used for the mean.

When a result is far from the others, repeat it. If three titres are 23.70, 23.15 and 23.15, use the two that agree. If none agree within 0.10 cm³, the mean of the closest pair is acceptable only if you state it, but you will lose the mark for concordance.

Indicators. The pH range of the indicator must fall within the steep part of the titration curve, so that one drop changes the colour sharply.

Indicator pH ranges for methyl orange, bromophenol blue, thymol blue and thymolphthalein compared with the steep regions of four types of titration

Strong acid with strong base: any of these indicators. Weak acid with strong base: thymol blue or thymolphthalein. Strong acid with weak base: methyl orange or bromophenol blue.

Titration	Equation or ratio	End-point colour change
acid with alkali	e.g. HCl + NaOH 1 : 1; H₂SO₄ + 2NaOH 1 : 2	depends on the indicator (methyl orange: red to yellow in acid-to-alkali titrations)
MnO₄⁻ with Fe²⁺	MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺ (1 : 5)	colourless to the first permanent pale pink
MnO₄⁻ with H₂O₂	2MnO₄⁻ + 6H⁺ + 5H₂O₂ → 2Mn²⁺ + 8H₂O + 5O₂ (2 : 5)	colourless to pale pink
MnO₄⁻ with ethanedioate	2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 8H₂O + 10CO₂ (2 : 5), warmed to about 60 °C	colourless to pale pink
I₂ with S₂O₃²⁻	I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻ (1 : 2)	brown or yellow to pale yellow; add starch near the end; blue-black to colourless

⚠️ Examiner Trap: Manganate(VII) is its own indicator. Potassium manganate(VII) is read from the top of the meniscus because the solution is too dark to see the bottom. The end-point is the first permanent pale pink in the flask. Do not add an indicator, and do not use hydrochloric acid to acidify, because Cl⁻ is oxidised by MnO₄⁻. Use dilute sulfuric acid.

Calculations. Always follow the same four steps.

Four steps: moles of the known solution, mole ratio, moles of the unknown, concentration or mass

Work in moles, use the equation ratio, then convert. Show each line.

n = c × V ÷ 1000 (V in cm³) c = n × 1000 ÷ V M_r = mass ÷ moles c₁V₁ = c₂V₂ for a dilution

Worked example. In a titration, 25.0 cm³ of an acidified solution of iron(II) sulfate required a mean titre of 23.15 cm³ of 0.0200 mol dm⁻³ potassium manganate(VII). Calculate the concentration of the iron(II) solution. [3]

Step 1. Moles of MnO₄⁻ = 0.0200 × 23.15 ÷ 1000 = 4.630 × 10⁻⁴ mol [1]. Step 2. The ratio is MnO₄⁻ : Fe²⁺ = 1 : 5, so moles of Fe²⁺ = 5 × 4.630 × 10⁻⁴ = 2.315 × 10⁻³ mol [1]. Step 3. Concentration = 2.315 × 10⁻³ × 1000 ÷ 25.0 = 0.0926 mol dm⁻³ [1]. The titre has four significant figures and the concentration of manganate(VII) has three, so three or four significant figures are acceptable.

Worked example. A student titrates 25.0 cm³ of aqueous iodine with sodium thiosulfate, 0.100 mol dm⁻³. The mean titre is 20.00 cm³. Calculate the concentration of the iodine. [3]

Moles S₂O₃²⁻ = 0.100 × 20.00 ÷ 1000 = 2.00 × 10⁻³ mol [1]. Ratio I₂ : S₂O₃²⁻ = 1 : 2, so moles I₂ = 1.00 × 10⁻³ mol [1]. Concentration = 1.00 × 10⁻³ × 1000 ÷ 25.0 = 0.0400 mol dm⁻³ [1].

⚠️ Examiner Trap: The ratio line. The most common lost mark is the mole ratio. Write the balanced equation, take the coefficients from it and state them (1 : 5, 1 : 2). Using 1 : 1 by habit gives a wrong answer even when the arithmetic is perfect. The method marks may still be awarded if the working is shown.

3. Thermometric, Rate, Gravimetric and Gas Methods

The syllabus names five types of quantitative experiment. Titration is the first. The other four are each built on one measured quantity and one calculation.

3.1 Thermometric experiments

A thermometer and stirrer in a lidded polystyrene cup supported in a beaker, with the q = mc delta T and delta H = minus q over n equations

Insulate the reaction, cover the bulb, stir, and record the highest (or lowest) temperature. Then q = mcΔT and ΔH = −q ÷ n.

Measure the initial temperature of the solutions (or of one, if they start at room temperature), mix, stir and record the maximum or minimum temperature. Calculate the temperature change, then the energy transferred.

Worked example. 25.0 cm³ of 2.00 mol dm⁻³ hydrochloric acid and 25.0 cm³ of 2.00 mol dm⁻³ sodium hydroxide, both at 21.0 °C, are mixed in a polystyrene cup. The maximum temperature is 34.5 °C. Calculate ΔH for the neutralisation, and the maximum percentage error in the temperature change. [4]

ΔT = 34.5 − 21.0 = 13.5 °C. Energy transferred q = m × c × ΔT = 50 × 4.18 × 13.5 = 2822 J [1]. Moles of HCl = 2.00 × 25.0 ÷ 1000 = 0.0500 mol [1]. ΔH = −2822 ÷ 0.0500 = −56 440 J mol⁻¹ = −56.4 kJ mol⁻¹ [1]. The sign is negative because the temperature rose. Maximum percentage error in ΔT = (2 × 0.5) ÷ 13.5 × 100 = 7.41% [1].

⚠️ Examiner Trap: Two readings, two uncertainties. A temperature change is a difference between two readings, so the uncertainty is twice 0.5 °C = ±1.0 °C. Using ±0.5 °C for ΔT halves the percentage error and loses the mark. Use the total mass of solution (50 g here) in q = mcΔT, but the moles of the limiting reagent in ΔH = −q ÷ n.

If the reaction is slow or heat is lost, record the temperature at regular intervals and plot a graph. The maximum temperature is then found by extrapolating the cooling line back to the time of mixing. Section 4 explains this, and Paper 5 uses it again.

3.2 Rate experiments

Disappearing cross: a flask of reaction mixture on a paper with a cross, viewed from above, timed until the cross is obscured

The time for the cross to disappear is inversely proportional to the rate when the amount of sulfur needed is the same each time.

Mix sodium thiosulfate with hydrochloric acid in a flask over a cross, start the clock when the acid is added, and stop it when the cross can no longer be seen. The rate is proportional to 1 ÷ t. Change one variable at a time (concentration of thiosulfate, or temperature) and keep every other variable the same, including the total volume and the observer. Because the end-point is a judgement, this is a source of random error.

3.3 Gravimetric experiments

A crucible with a lid on a pipe-clay triangle over a Bunsen burner with the four-step procedure to constant mass

Heat, cool, weigh and repeat until two successive masses agree within 0.01 g. The solid has then lost all of its volatile component.

A typical experiment is the determination of the water of crystallisation of a hydrated salt. Weigh the empty crucible, add the salt, weigh, heat gently then strongly, cool and weigh. Repeat the heating until the mass is constant. Do not allow the solid to spit out of the crucible.

Worked example. A crucible has a mass of 21.36 g. With hydrated copper(II) sulfate, CuSO₄•xH₂O, the mass is 23.86 g. After heating to constant mass the crucible and solid have a mass of 22.96 g. Calculate x. [4]

Mass of hydrate = 23.86 − 21.36 = 2.50 g. Mass of anhydrous CuSO₄ = 22.96 − 21.36 = 1.60 g, so mass of water = 2.50 − 1.60 = 0.90 g [1]. Moles of CuSO₄ = 1.60 ÷ 159.6 = 0.01003 mol [1]. Moles of H₂O = 0.90 ÷ 18.0 = 0.0500 mol [1]. x = 0.0500 ÷ 0.01003 = 4.99, so x = 5 [1].

3.4 Gas volume experiments

Two ways to collect a gas: a gas syringe, and an inverted measuring cylinder over water

Use a gas syringe where the gas is soluble or the volume is small; the apparatus must be gas-tight and the bung replaced quickly.

React a known mass of a carbonate with excess acid, collect the carbon dioxide and read its volume. Add the solid last and replace the bung immediately, otherwise some gas escapes before the clock or the syringe is ready. Convert volume to moles using 24.0 dm³ mol⁻¹ at room conditions, which the Data Booklet gives.

Worked example. A 0.250 g sample of impure calcium carbonate reacts with excess hydrochloric acid and gives 52.0 cm³ of carbon dioxide at room conditions. Calculate the percentage purity. [4]

Moles of CO₂ = 52.0 ÷ 24 000 = 2.167 × 10⁻³ mol [1]. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂, so moles of CaCO₃ = 2.167 × 10⁻³ mol [1]. Mass = 2.167 × 10⁻³ × 100.1 = 0.2169 g [1]. Purity = 0.2169 ÷ 0.250 × 100 = 86.8% [1].

⚠️ Examiner Trap: A syringe has a limit. A gas syringe holds 100 cm³. If the calculation shows that the reaction would give more gas than the syringe can hold, the mass of solid must be reduced. Dissolved CO₂ also makes a gas volume collected over water smaller than expected, a systematic error.

4. Tables and Graphs

Tables. Draw the table before you start. Give each column a heading in the form quantity / unit, for example time / s or temperature / °C. Put only numbers in the body. Give all readings of one quantity to the same number of decimal places. The independent variable is in the first column.

Graphs. Marks are awarded for the axes, the scale, the points, the line and (when asked) the use of it.

A well-drawn straight-line graph with labelled axes, plotted crosses, one circled anomalous point, a ruled line of best fit and a large gradient triangle

Label the axes, use more than half the grid, plot with neat crosses, draw one ruled line and use a large triangle for the gradient.

Feature	What is expected
axes	independent variable on the x-axis; label as quantity / unit
scale	1, 2 or 5 (or multiples of 10) units to a 20 mm square; no awkward scales such as 3 or 7; the points use more than half the grid in both directions
points	a small cross (×) or a circled dot, plotted to within half a small square
line	a sharp pencil; a straight line with a ruler, or a smooth curve; about the same number of points above and below it
anomaly	circle it, state that it is anomalous, and ignore it when drawing the line
gradient	a right-angled triangle that covers more than half of the line, with the coordinates read from the line and not from the table
intercept or extrapolation	extend the line to the axis, dashed, and read the value

Using the graph. The gradient is the change in y divided by the change in x, with units. For a rate graph of 1 ÷ t against concentration, a straight line through the origin shows that the rate is directly proportional to concentration (first order with respect to that reagent). An intercept or the intersection of two lines gives a value that is not measured directly, such as the volume at the end-point of a thermometric titration. An extrapolation is used to find a temperature at a time when it could not be measured.

Worked example. A student plots 1 ÷ t (in 10⁻³ s⁻¹) against the concentration of sodium thiosulfate. A straight line passes through the origin and through the point (0.100, 50.0). Calculate its gradient, and state what it shows. [3]

Gradient = (50.0 − 0) × 10⁻³ ÷ (0.100 − 0) = 0.500 dm³ mol⁻¹ s⁻¹ [2: value, units]. The line is straight and passes through the origin, so the rate is directly proportional to the concentration of thiosulfate [1].

⚠️ Examiner Trap: Triangles and points. Take the gradient from two points on the line, far apart (more than half the line), not from two data points and not from a small triangle. A line of best fit does not have to start at the first point or pass through the origin unless the data show that it should. Join no points ‘dot to dot’.

5. Qualitative Analysis

One question gives you unknown substances, labelled FA 1 and so on, and a set of tests. You record observations, then use the Qualitative analysis notes printed in the paper to identify cations, anions and gases. Notes for the paper: do not carry out tests the method does not ask for, add only the volume specified (about 1 cm depth in a test-tube), and use a boiling tube to warm a solution and a hard-glass tube to heat a solid.

Observation language. Write what you see. Use simple terms: colourless solution, white ppt., effervescence, pale blue, dark green. Use ‘ppt.’ for precipitate. State when a ppt. dissolves in excess or does not dissolve, and whether the change is on standing. If nothing happens, write no change or no ppt.

Table of reactions of aluminium, ammonium, barium, calcium, chromium(III), copper(II), iron(II), iron(III), magnesium, manganese(II) and zinc ions with sodium hydroxide and ammonia

Add NaOH(aq) and NH₃(aq) a few drops first, then in excess. Whether the ppt. dissolves in excess separates ions that look alike.

Making the cation decision. First look at the colour of the solution and the ppt. Coloured ppt.: Cu²⁺ (pale blue), Fe²⁺ (green, turning brown), Fe³⁺ (red-brown), Cr³⁺ (grey-green), Mn²⁺ (off-white, turning brown). White ppt.: use the excess test. A white ppt. that dissolves in excess NaOH is Al³⁺ or Zn²⁺; in excess NH₃, Al³⁺ does not dissolve but Zn²⁺ does. A white ppt. that does not dissolve in excess is Mg²⁺ (ppt. with NH₃ as well) or Ca²⁺ (no ppt. with NH₃). A faint white ppt. with NaOH and none with NH₃ is Ba²⁺. If there is no ppt. with NaOH, warm; ammonia gas turning damp red litmus blue shows NH₄⁺.

Table of reactions of carbonate, chloride, bromide, iodide, nitrate, nitrite, sulfate, sulfite and thiosulfate and the tests for gases

The tests use dilute acid, silver nitrate, barium ions, hot alkali with aluminium and acidified manganate(VII).

Making the anion decision. Add dilute acid first: effervescence of CO₂ (limewater turns milky) shows carbonate; a pale yellow ppt. forming slowly shows thiosulfate. To separate the halides use silver nitrate, then ammonia: white and soluble is chloride, cream and partly soluble is bromide, pale yellow and insoluble is iodide. Barium ions give a white ppt. with sulfate (insoluble in acid) and sulfite (dissolves in acid). For nitrate and nitrite warm with NaOH and aluminium foil: ammonia gas turns damp red litmus blue. Acidified manganate(VII) is decolourised by nitrite (not nitrate) and by sulfite.

Organic test (Paper 3)	Positive result	What it shows
Fehling’s reagent, warm	orange / red ppt.	aldehyde (a ketone gives no change)
Tollens’ reagent, warm	silver mirror or black ppt.	aldehyde
alkaline aqueous iodine, warm	yellow ppt.	CH₃CO– or CH₃CH(OH)– group
acidified KMnO₄, warm	purple to colourless	a compound that can be oxidised (alcohol, aldehyde, alkene)

Worked example. Solution FA 1 gives a white ppt. with NaOH(aq) that dissolves in excess, and a white ppt. with NH₃(aq) that does not dissolve in excess. A second sample gives a white ppt. with Ba²⁺(aq) that remains when dilute nitric acid is added. Identify the ions in FA 1. [4]

White ppt. with NaOH that dissolves in excess is Al³⁺ or Zn²⁺ [1]. The white ppt. with NH₃ that does not dissolve in excess shows Al³⁺ (Zn²⁺ would dissolve) [1]. A white ppt. with Ba²⁺ that is insoluble in acid shows SO₄²⁻ [1]. FA 1 is aluminium sulfate, Al₂(SO₄)₃ [1].

⚠️ Examiner Trap: Observation first, name last. Writing ‘Al(OH)₃ forms’ or ‘sulfate present’ in the observation column does not earn the observation mark. State only what you see. Also, an answer such as ‘add acid’ for a test is not enough: name the reagent. If a substance contains an ion not in the notes, record the observations and give a general conclusion only.

Confirmatory tests and decisions. Cambridge awards marks for choosing a reagent that distinguishes two possibilities. To separate Al³⁺ from Zn²⁺ add excess NH₃. To separate Mg²⁺ from Ca²⁺ use NH₃. To separate Cl⁻ from Br⁻ from I⁻ use Ag⁺ then NH₃. To separate SO₄²⁻ from SO₃²⁻ add acid to the ppt. or use acidified KMnO₄.

6. Analysis, Errors and Evaluation

Systematic and random errors. A systematic error shifts every reading in the same direction, for example a balance with a zero error, so using the same balance for every weighing cancels it in a difference. A random error varies unpredictably, for example a change in room temperature or judging the end of the disappearing cross, and repeated readings with a mean reduce it.

Two graphs: random error shows scatter about the true line, systematic error shows all points shifted

Random errors scatter the points; systematic errors shift them all. Repeats reduce random errors but not systematic ones.

Uncertainty and percentage error. The maximum uncertainty of one reading is half the smallest division. Percentage error = uncertainty ÷ measured value × 100. For a difference (a titre, a temperature change), add the uncertainties of the two readings.

Percentage error of a temperature change of 5, 10, 14 and 25 degrees measured with a thermometer with 1 degree divisions

The same ±1.0 °C uncertainty gives a smaller percentage error when the change is bigger.

Worked example. A burette is read before and after a titration. The titre is 23.15 cm³. Calculate the maximum percentage error in the titre. [2]

Uncertainty in the titre = 0.05 + 0.05 = ±0.10 cm³ [1]. Percentage error = 0.10 ÷ 23.15 × 100 = 0.43% [1].

Worked example. In the gravimetric experiment in Section 3 the balance reads to 0.01 g. Calculate the maximum percentage error in the 2.50 g mass of hydrate, which is found as a difference of two weighings. [2]

Uncertainty = 0.005 + 0.005 = ±0.01 g [1]. Percentage error = 0.01 ÷ 2.50 × 100 = 0.40% [1].

Suggesting improvements. A good answer links the error to a change and to its effect. Choose the largest source of error first.

Source of error	Improvement	Why it helps
heat lost in a thermometric experiment	use a lid and a polystyrene cup; plot a cooling curve and extrapolate	reduces heat loss; finds the maximum temperature at the time of mixing
a small temperature rise measured on a thermometer with 1 °C marks	use more concentrated solutions or a larger mass of solid, or a thermometer reading to 0.1 °C	a larger ΔT gives a smaller percentage error
reaction too slow to see the end-point	use a colorimeter or a data logger	removes the judgement of the observer
gas lost before the syringe is attached	add the solid last and replace the bung immediately; use a tube to hold the solid	all the gas is collected
crucible not constant in mass	heat to constant mass and cool in a desiccator	all the water is lost and none is regained
few repeats	repeat the experiment and use a mean; check for anomalies	reduces the effect of random errors
burette reading error of ±0.05 cm³	use a larger titre (a more dilute titrant or a bigger sample)	a larger volume gives a smaller percentage error

Drawing conclusions. State the finding in terms of the data (‘the rate doubles when the concentration doubles’) and then give the chemical reason from the theory. When a hypothesis is given, say whether the data support it and quote one piece of evidence.

⚠️ Examiner Trap: Not ‘human error’. The syllabus says that ‘human error’ is not an acceptable source of error. Name the instrument or the effect: ‘heat loss to the surroundings’, ‘the thermometer reads to 0.5 °C’, ‘the end-point of the cross is hard to judge’. For improvements, use more accurate apparatus is not enough; say which, and why it is more accurate.

7. Quick Sheet and Checklist

Topic	What to remember
precision	burette 0.05 cm³; thermometer 0.5 °C; mass as the balance shows; one precision for each quantity
titration	rinse, rough first, dropwise near the end-point, concordant within 0.10 cm³, mean of concordant titres
ratios	MnO₄⁻ : Fe²⁺ 1 : 5; MnO₄⁻ : H₂O₂ 2 : 5; MnO₄⁻ : C₂O₄²⁻ 2 : 5; I₂ : S₂O₃²⁻ 1 : 2
thermometric	q = mcΔT; ΔH = −q ÷ n; ΔT uncertainty ±1.0 °C
gravimetric	heat to constant mass (within 0.01 g); x = moles water ÷ moles salt
gas volume	24.0 dm³ mol⁻¹ at room conditions; syringe 100 cm³
graph	axes labelled, scale over half the grid, crosses, ruled line, big gradient triangle
qualitative	a few drops, then excess; record ‘no change’; observation first, conclusion last
errors	systematic and random; % error = uncertainty ÷ value × 100; no ‘human error’

[ ] I can read a burette, thermometer, cylinder and balance to the correct precision and record every reading in a table that I drew first.
[ ] I can carry out a titration with a rough and accurate runs, identify concordant titres and calculate a mean.
[ ] I can calculate concentration, mass or M_r from a titre, with the mole ratio stated and the right number of significant figures.
[ ] I can find ΔH, the formula of a hydrated salt, a purity or a rate from thermometric, gravimetric, gas-volume and rate data.
[ ] I can draw a graph with labelled axes, a good scale, plotted crosses and a ruled line, and find a gradient, an intercept or an intersection.
[ ] I can identify the cations, anions and gases in the Qualitative analysis notes and choose a reagent that distinguishes two ions.
[ ] I can state the uncertainty of a reading, calculate a percentage error, distinguish random from systematic errors and suggest a specific improvement.