Rates of Reaction
Aim: investigate how the concentration of a reactant affects the rate of reaction, following the reaction by two different methods. Like RP3, this practical is a named opportunity for developing a hypothesis: “increasing the concentration will increase the rate, because there are more particles in the same volume, so collisions are more frequent.”
Full theory: C6: Rates & Equilibrium
Sodium thiosulfate reacts with dilute hydrochloric acid to make a fine precipitate of sulfur, which slowly clouds the mixture:
Na2S2O3(aq) + 2HCl(aq) → 2NaCl(aq) + S(s) + SO2(g) + H2O(l)
- Measure thiosulfate solution into a conical flask standing on a black cross drawn on paper.
- Add the acid, swirl, and start the stop clock.
- Look down through the solution; stop the clock the moment the cross is no longer visible.
- Repeat with different concentrations of thiosulfate, made by diluting the same stock solution with water so the total volume is always the same (e.g. 10 cm3 thiosulfate + 40 cm3 water, then 20 + 30, 30 + 20…). Record every time in seconds, repeat the whole investigation and take means, leaving out anomalies.
A shorter time means a faster rate: rate is proportional to 1/time.
React magnesium (or marble chips) with dilute hydrochloric acid and collect the gas in a gas syringe (or an upside-down measuring cylinder full of water). Record the volume every 10 seconds, repeat with a different acid concentration, and plot volume against time. The steeper the initial gradient, the faster the rate.
Two of the ways of following the rate in required practical 5. Method 1: time how long the black cross takes to vanish as the solution clouds. Method 2: read the gas volume off a syringe at regular times and plot the curve.
If a reaction gives off a gas, you can follow it on a balance. Stand the flask of reactants on the balance, plug the neck with cotton wool, and record the mass every few seconds. As gas escapes into the air the mass of the flask falls. A faster fall means a faster reaction.
CaCO3 + 2HCl → CaCl2 + H2O + CO2
- The usual reaction is a carbonate with acid: the carbon dioxide leaves the flask, so the reading drops.
- The cotton wool plug earns marks of its own: it lets the gas escape (so the mass can change) while stopping acid spray being thrown out as the mixture fizzes. Spray leaving would make the mass fall too far.
- Mass only appears to be lost: no atoms are destroyed; the gas has simply left the flask. Seal the flask and the mass wouldn’t change at all.
- Mass loss suits a heavy gas like CO2 but is the wrong method for hydrogen: H2 is too light to move the balance reading reliably, so collect it in a gas syringe instead. Matching the method to the gas is itself a common exam point.
Following a reaction by loss of mass: a carbonate reacts with acid in a flask plugged with cotton wool. The CO2 escapes into the air so the balance reading falls, while the cotton wool stops acid spray being lost with it.
The concentration of the reactant solution (the thiosulfate, or the acid).
The time for the cross to disappear, or the volume of gas at set times.
The total volume (and so depth) of solution, the temperature, the amount of the other reactant, and the same flask, cross and observer.
The mean rate over a period is volumetime, but the rate at one moment is the gradient of the tangent to the curve at that time. Draw the tangent touching the curve at the point, make the gradient triangle large (small triangles magnify drawing errors), show your working on the graph, and quote the answer with its unit (e.g. cm3/s). Questions sometimes ask for the answer in standard form, e.g. 0.0025 = 2.5 × 10−3.
Tangent explorer: the rate at any moment
Magnesium + dilute hydrochloric acid, with the hydrogen collected in a gas syringe. Slide along the curve and watch the tangent, the triangle and the gradient calculation update. Notice when the reaction is fastest, and what the gradient becomes once it finishes.
- Same depth of solution every time. You look down through the liquid, so a deeper mixture hides the cross sooner. Topping up with water when diluting keeps the total volume, and so the depth, the same.
- Measure volumes with a measuring cylinder, not a beaker. Its finer scale (better resolution) makes the “same total volume” genuinely the same.
- State the observation, not an inference. Timing stops when “the cross is no longer visible”. “The reaction finished” is an inference, and a false one: the reaction is still going; there’s just enough sulfur to block your view.
- The end-point judgement is the weakness. Improve it with the same observer every time, or remove the judgement entirely with a light sensor and data logger.
- Safety. The reaction makes sulfur dioxide, a toxic gas: keep the room well ventilated.
Exam-style questions
A student reacts marble chips with dilute hydrochloric acid in a flask on a balance. The balance reading decreases during the reaction. Why? Tick (✓) one box.
A student investigates how concentration affects the rate of the sodium thiosulfate and acid reaction, using the disappearing cross. Which variables must be kept the same? Tick (✓) one box.
A reaction produces 48 cm3 of gas in the first 120 seconds. (a) Calculate the mean rate of reaction over this time. Give the unit. (b) The volume–time graph for this reaction is steepest at the start. Explain why, in terms of particles.
Show the mark scheme
- (a) Mean rate = 48120 = 0.4 1 mark
- Unit: cm3/s 1 mark
- (b) At the start the concentration of the reactants is greatest… 1 mark
- …so collisions are most frequent, giving the fastest rate (steepest gradient). 1 mark
Say more frequent collisions, not just “more collisions”. As the reactants are used up the collisions become less frequent, the curve flattens, and it goes horizontal when one reactant has run out.
An experiment with 20 g of large marble chips and excess acid is repeated with 20 g of smaller marble chips, with everything else unchanged. Describe how the volume–time curve for the smaller chips would compare, and explain why.
Show the mark scheme
- The curve is steeper at the start (faster rate). 1 mark
- It levels off at the same final volume of gas, sooner. 1 mark
- Smaller chips have a larger surface area (to volume ratio), so collisions at the surface are more frequent. 1 mark
The same mass of carbonate makes the same amount of CO2. Surface area changes how fast you get there, not how much you get. Sketch questions mark exactly these two features: steeper start, same plateau.
Chromatography
Aim: use paper chromatography to separate and identify a mixture of coloured substances (e.g. food colourings), and calculate Rf values.
Full theory: C8: Chromatography & Rf Values
The one rule the practical is built on: the start line and spots must sit above the solvent, or the samples wash straight off into the beaker.
- Draw a pencil start line about 2 cm from the bottom of the chromatography paper (pencil is insoluble, so it won’t run).
- Use a fresh capillary tube to put a small spot (2–3 mm) of each known colouring on the line, plus a spot of the unknown mixture. Label each in pencil.
- Add solvent to the beaker to a depth of less than 1 cm, so the level is below the start line. Hang the paper (e.g. from a glass rod resting on the beaker’s rim) so its bottom edge dips in. The sides of the paper must not touch the beaker wall.
- Leave it undisturbed until the solvent has risen about three-quarters of the way up.
- Remove the paper, immediately mark the solvent front in pencil, and let it dry.
- Measure, in mm, the distance from the start line to the centre of each spot, and from the start line to the solvent front. Calculate Rf for each.
Rf = distance moved by substancedistance moved by solvent
Analysis: compare each known colouring’s Rf (and spot position) with the unknown to see which dyes it contains. A substance that gives a single spot is pure; one that separates into several spots is a mixture. Rf values let you compare results between labs, but only if the same solvent was used, since Rf depends on the solvent.
The substance spotted on the start line (each different colouring or ink).
The distance each spot moves, used to work out its Rf value.
The same solvent, paper and start line for every sample; always measure to the centre of the spot.
- Pure / impure: categorical, not relative. A single spot means the substance is pure; more than one spot means impure. Writing “purer” or “less pure” misses the mark. Tie the word to the number of spots.
- Why a dye travels further. The marks want both halves of the balance: it is more soluble in the solvent and has a weaker attraction to the paper. Solubility alone is usually not enough.
- What Rf really says. For a given substance, solvent and paper, the spot and the solvent travel in a constant proportion: that ratio is the Rf, which is why it identifies the substance regardless of how long the paper runs.
- Start line in ink. The ink dissolves and travels up with the samples, ruining the chromatogram. Always pencil.
- Solvent above the start line. The spots dissolve straight into the solvent and are lost — the one rule the practical is built on.
- Measuring to the top of a spot. Always measure to the centre.
- An Rf greater than 1, or with units. Rf is a ratio of two distances: no units, always less than 1. If yours is bigger than 1, you’ve divided the wrong way round.
Try the common mistakes
Run the chromatography three ways and watch what goes wrong with each mistake.
Correct setup: a pencil start line, solvent below the spots. Press play.
Exam-style questions
Why is the start line drawn in pencil? Tick (✓) one box.
A spot moves 3.6 cm and the solvent front moves 8.0 cm. What is the Rf value? Tick (✓) one box.
A chromatogram of an unknown mixture U is run alongside four known colourings A–D. U separates into three spots: two line up exactly with the spots from A and C, and one matches none of the known colourings. What two conclusions can be drawn about U?
Show the mark scheme
- U contains (the same dyes as) A and C. 1 mark
- U also contains at least one other substance that is not A, B, C or D. 1 mark
Matching spots means same height / same Rf in the same solvent. Note the careful wording (“at least one”): two different substances can sit at the same height, so one unmatched spot doesn’t prove only one extra dye.
A student wants to find out whether a green ink contains the same blue dye as a sample of pure blue dye. Describe a method the student could use. You should name the technique and say how the result would show the answer.
Show the mark scheme
- Use paper chromatography: draw a pencil line near the bottom of the paper and put a small spot of the green ink and a spot of the pure blue dye on the line. 1 mark
- Stand the paper in a small depth of solvent, with the solvent level below the start line. 1 mark
- Leave until the solvent has nearly reached the top; the green ink separates into its dyes. 1 mark
- If one of the green ink’s spots is at the same height (same Rf) as the pure blue dye’s spot, the ink contains that dye. 1 mark
The comparison only works because both samples ran on the same paper, in the same solvent, from the same start line. That’s what makes the Rf values comparable.
Identifying Ions Triple Only
Aim: use chemical tests to identify the ions in a single unknown ionic compound: a cation test plus an anion test together name the compound.
Full theory: C8: Identifying Ions (sections 5–7)
Flame tests: metal cations
- Clean a nichrome (or platinum) wire loop by dipping it in dilute acid and holding it in a blue Bunsen flame until it adds no colour. This removes any ions left from before, so you test only your sample.
- Dip the clean loop in the sample and hold it in the edge of the blue flame.
- Observe the colour. Don’t let the wire glow red-hot. That orange glow can be mistaken for a flame colour.
The big limitation: in a mixture, a strong colour (sodium’s yellow especially) masks the others, which is exactly why flame emission spectroscopy is better for mixtures: it picks out every ion at once, quickly, sensitively and with a reading that can be matched to a database. If you name it, write flame emission spectroscopy in full. Abbreviations like “F.E.S.” are not credited.
Sodium hydroxide precipitates: metal cations
Add a few drops of sodium hydroxide solution first (slowly), then add it in excess:
| Ion | Precipitate | In excess NaOH |
|---|---|---|
| Copper(II), Cu2+ | Blue | No change |
| Iron(II), Fe2+ | Green | No change |
| Iron(III), Fe3+ | Brown | No change |
| Aluminium, Al3+ | White | Dissolves (colourless solution) |
| Calcium, Ca2+ | White | Does not dissolve |
| Magnesium, Mg2+ | White | Does not dissolve |
Three coloured precipitates (copper blue, iron(II) green, iron(III) brown) and three that are all white (aluminium, calcium, magnesium). The whites need a second test to tell apart.
Of the three whites, only aluminium re-dissolves in excess. To separate the remaining two, use a flame test: calcium is orange-red, magnesium gives no colour.
The three anion tests
| Anion | Add this… | Positive result |
|---|---|---|
| Carbonate (CO32−) | Dilute acid | Fizzes; the gas (CO2) turns limewater milky |
| Halide (Cl−, Br−, I−) | Dilute nitric acid, then silver nitrate | Silver halide precipitate: white (Cl), cream (Br) or yellow (I) |
| Sulfate (SO42−) | Dilute hydrochloric acid, then barium chloride | White precipitate (barium sulfate) |
The silver halides step through white → cream → yellow as you go down the group (Cl → Br → I). Barium sulfate is a separate white precipitate, made with barium chloride.
Carbonate ions would also form a precipitate with silver nitrate or barium chloride, giving a false positive. Adding acid first reacts away any carbonate (it fizzes off as CO2) so it can’t interfere. The catch is choosing an acid that doesn’t add the very ion you’re testing for: nitric acid for the halide test (hydrochloric would add chloride ions), and hydrochloric acid for the sulfate test (sulfuric would add sulfate ions).
- “Clear” when you mean “colourless”. “Clear” only means see-through: copper sulfate solution is clear and blue. Examiners penalise “clear” here.
- Adding NaOH too fast. Flood aluminium with excess straight away and the white precipitate forms and re-dissolves before you spot it. A few drops first.
- Chloride ions are not chlorine gas. The test for chloride ions is nitric acid + silver nitrate (white precipitate); the test for chlorine gas is damp blue litmus bleaching white (section 9). Examiners report students giving the gas test when the ion test is asked for. Read which one the question wants.
- Stopping at the cation. A full identification needs both halves: e.g. green flame + white precipitate with barium chloride = copper sulfate.
Six-mark plan questions here are marked on whether the method would actually deliver an answer for the sample in the question. Students who list correct tests in the abstract (without saying they would carry them out on the sample, or in a workable order) can score nothing despite knowing every test. Structure the plan: take (a solution of) the sample → flame test for the cation → one anion test at a time, naming the reagent and the expected result for a positive.
Put the whole strategy together: run the bench tests and name the unknown compound:
Identify the unknown salt
You’re handed an unknown ionic compound. Run bench tests to find its cation (metal ion) and anion (non-metal ion), then name it, in as few tests as you can. For each result, try to recall what it tells you before you tap Reveal. This is exactly the Required practical 7 skill.
Pick a test to start.
| Test | Observation | What does it tell you? |
|---|---|---|
| No tests run yet — click a test above. | ||
Exam-style questions
Sodium hydroxide solution is added to a solution of a metal compound. A white precipitate forms, which dissolves when excess sodium hydroxide is added. Which metal ion is present? Tick (✓) one box.
What is the test for sulfate ions? Tick (✓) one box.
An unknown ionic compound gives a yellow flame in a flame test. A solution of the compound is acidified with dilute nitric acid, and silver nitrate solution is added: a cream precipitate forms. Name the compound. Explain how the results identify each ion, and why the solution is acidified before adding the silver nitrate.
Show the mark scheme
- Yellow flame = sodium ions (Na+). 1 mark
- Cream precipitate with silver nitrate = bromide ions (Br−): silver bromide. 1 mark
- The compound is sodium bromide. 1 mark
- The nitric acid removes carbonate ions, which would also form a precipitate with silver nitrate and give a false positive. 1 mark
The silver halides step white → cream → yellow for Cl → Br → I. And note it must be nitric acid here. Hydrochloric acid would add chloride ions to the solution being tested.
Flame emission spectroscopy is an instrumental method for identifying metal ions. Give two advantages of instrumental methods compared with flame tests.
Show the mark scheme
- Any two of: more sensitive (detects tiny amounts) / more accurate / faster / works on mixtures, where one flame colour would mask another / gives the concentration as well as the identity. 1 mark each
The mixtures point is the strongest pairing with this practical: a flame test can’t separate overlapping colours (sodium’s yellow especially), but each element’s line spectrum shows up independently.
Water Analysis & Purification
Aim: analyse water samples for pH and dissolved solids, then purify a sample by distillation.
Full theory: C10: Potable Water
Part A: analysis.
- Measure the pH of each sample with universal indicator (or a pH probe).
- Find the dissolved solids: weigh an evaporating basin, add a known volume (e.g. 25 cm3) of the sample, evaporate the water, then re-weigh. The mass gained is the dissolved solid. Don’t overheat, or you decompose the solids.
Part B: purification by distillation.
- Heat the sample so the water boils off, then condense the vapour and collect it. The distillate is pure water; the dissolved solids stay behind in the flask.
- Test the distillate: pure water boils at exactly 100 °C and leaves no residue on evaporation.
Simple distillation: boil the salt water, condense the steam in a cooled tube, and collect pure water. The salt is left behind in the flask.
The water sample tested (fresh, sea or distilled) or, in Part B, before vs after distillation.
The pH, the mass of dissolved solids per known volume, and whether the distillate is pure (boils at 100 °C, no residue).
The same volume of each sample, and the same apparatus and method each time (same balance, same basin, no overheating).
AQA’s student sheet tests the salt water chemically, before and after distilling, a neat link to the ion tests of RP7: a flame test (yellow = sodium ions) and dilute nitric acid + silver nitrate (white precipitate = chloride ions). The salt water gives both positives; the distillate (collected through a delivery tube into a test tube standing in ice water to condense the steam) gives neither. No yellow flame, no precipitate: the ions have been left behind, so the water is purified. A likely exam angle: “How could the student show the distilled water no longer contains chloride ions?”
- “Potable means pure.” No. Potable water still contains dissolved substances; it’s just safe to drink. Pure water contains only H2O.
- Confusing the two treatment steps. Filtration removes insoluble solids; sterilisation (chlorine, ozone or UV) kills microbes. Filtering does not kill bacteria.
- Forgetting why desalination is a last resort. Distillation and reverse osmosis both need large amounts of energy, so they’re expensive, used only where fresh water is scarce.
Exam-style questions
Which result shows that a sample of water is pure? Tick (✓) one box.
In the treatment of fresh water for drinking, what is the purpose of sterilisation? Tick (✓) one box.
Sea water can be made into drinking water by distillation. Give one reason why distillation works for this, and one reason why it is only used where fresh water is scarce.
Show the mark scheme
- Only the water boils off and is condensed. The dissolved salt is left behind in the flask. 1 mark
- Boiling the water needs large amounts of energy, so it is expensive. 1 mark
The same applies to reverse osmosis: both desalination methods are energy-intensive, which is why they’re a last resort.
A student distils salt water in the lab. Describe how the student should collect the distilled water, and how they could show that the distilled water no longer contains chloride ions.
Show the mark scheme
- Heat the salt water (in a flask) so the water boils off as steam. 1 mark
- Condense the steam (e.g. lead it through a delivery tube into a test tube standing in ice water) and collect the distillate. 1 mark
- Add dilute nitric acid then silver nitrate solution to the distillate. 1 mark
- No white precipitate forms, showing the chloride ions have been left behind in the flask. 1 mark
The salt water itself gives a white precipitate with the same test: testing before and after is what proves the distillation worked. This is the RP7 halide test doing its job inside RP8.
The Four Gas Tests
Not a required practical, but these four tests are needed inside RP3 and RP5, and they earn marks across both papers. Learn each as a method-and-result pair.
Full theory: C8: Tests for Common Gases
| Gas | Test | Positive result |
|---|---|---|
| Hydrogen, H2 | Hold a lighted splint at the open end of the tube | Burns rapidly with a squeaky “pop” |
| Oxygen, O2 | Place a glowing splint inside the tube | The splint relights |
| Carbon dioxide, CO2 | Bubble the gas through limewater | Limewater turns milky / cloudy |
| Chlorine, Cl2 | Hold damp blue litmus paper in the gas | Litmus is bleached white |
- “A splint” isn’t enough for oxygen. It must be a glowing splint (one that has just gone out, so only the ember still glows) and the oxygen relights the flame. “Splint” or “lit splint” earns nothing here: a lit splint is the hydrogen test.
- “CO2 puts out a lit splint” is not the test. Other gases do this too. It’s a property, not an identification.
- Memory hooks: liGHted → Hydrogen; glOwing → Oxygen.
- State results precisely: “relights” not “glows brighter”; “milky/cloudy” not “white”; “bleached white” not “clear”.
Cracking & the Test for Alkenes
From C7: cracking itself is usually examined as conditions and equations, but the bromine water test is a hands-on test exams expect you to describe precisely.
Full theory: C7: Cracking & the Test for Alkenes
- Catalytic cracking: heat the fraction to vaporise it, then pass the vapours over a hot catalyst.
- Steam cracking: mix the vaporised fraction with steam at a very high temperature.
- Asked for two conditions? High temperature + catalyst, or high temperature + steam, are the safe pairings. And why must air be kept out of the reactor? The hot hydrocarbon vapour would combust.
- The test for alkenes: shake the sample with bromine water. An alkene turns it from orange to colourless; an alkane leaves it orange. (Colourless, not “clear”.)
Bromine water: orange → colourless with an alkene; stays orange with an alkane.