Pearson Edexcel International GCSE in Chemistry · 4CH1 · Free · no account needed
Chemistry required practicals
All 12 required practicals for Pearson Edexcel International GCSE Chemistry (4CH1) on one page — the method step by step, the independent, dependent and control variables with the reason each one is held constant, and the sources of error classified as systematic, random or judgement. Each is taught inside its lesson, linked below.
The specification lists them as practical investigations; Pearson's UK GCSE calls the equivalent set core practicals, and students call them required practicals — this is that list for the International GCSE. They are examined in writing: "state a control variable and explain why it must be kept constant", "suggest an improvement", "explain why the result is anomalous". Every mark in those questions is a reason, not a step, which is why each practical here carries the reason beside each variable and the fix beside each error. Learn the method here; then open the lesson, understand the science the practical is testing, prove it in five marked questions, teach it back to an AI pupil and defend it in a short debate — the four things a method sheet cannot ask of you. Completely free.
The 12 practicals
- paper chromatography of inks or food colourings — 1.13
- measuring the solubility of a solid at a given temperature — 1.7
- the formula of magnesium oxide — 1.36
- electrolysis of aqueous solutions — 1.60
- the percentage of oxygen in air — 2.14
- reactions of dilute acids with metals — 2.21
- hydrated copper(II) sulfate crystals — 2.42
- a pure dry sample of lead(II) sulfate — 2.43
- measuring temperature changes — 3.8
- marble chips and hydrochloric acid — 3.15
- catalytic decomposition of hydrogen peroxide — 3.16
- preparing a sample of an ester — 4.43
Topic 1 · Principles of chemistry
1. Required practical (1.13): paper chromatography of inks or food colourings
Specification statement 1.13 · Topic 1 · taught in Pure Substances, Mixtures and How to Separate Them
Method
- Draw the baseline in PENCIL about 1 cm from the bottom — ink would dissolve and travel.
- Spot the samples on the baseline, drying between drops to keep spots small.
- Stand the paper so the solvent level is BELOW the baseline, or the spots dissolve into the solvent.
- Let the solvent rise near the top, then mark the solvent front in pencil immediately.
- Measure to the centre of each spot and calculate Rf.
Variables. Independent: the substance spotted. Dependent: distance travelled. Control: same solvent, same paper, same baseline, same run time.
Sources of error and improvement. Spots that are too large smear and have no measurable centre — use a fine capillary and build up in small drops. If the solvent runs off the top, the front cannot be measured and no Rf can be calculated.
Understand the science behind it → · Prove it · Teach it · Defend it
2. Required practical 1.7 — measuring the solubility of a solid at a given temperature
Specification statement 1.7 · Topic 1 · taught in Solubility and Solubility Curves
Method
- Add solute to water in a boiling tube and stir at the chosen temperature until no more dissolves and some solid remains undissolved, so the solution is certainly saturated.
- Hold the tube at that temperature in a water bath and allow it to stand, so the solution reaches equilibrium with the excess solid.
- Weigh a clean dry evaporating basin.
- Decant or filter a sample of the saturated solution into the basin, taking none of the undissolved solid, and weigh the basin and solution.
- Evaporate the water carefully, then dry the residue, cool and reweigh. Reheat and reweigh until the mass is constant.
- Mass of solute = final mass − empty basin. Mass of water = solution mass − solute mass. Solubility = (solute ÷ water) × 100.
Variables
Independent (changed) — temperature of the water
Dependent (measured) — mass of solute dissolved per 100 g of water
| Control variable | Why it must be held constant |
|---|---|
| The solute used | every solute has its own solubility curve |
| Purity of the water | dissolved impurities change how much solute can dissolve |
| Time left to equilibrate | too short and the solution is not yet saturated |
| Mass of the basin | it must be dry before the first weighing |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | Undissolved solid is carried over with the solution. | Filter through a pre-warmed filter so nothing cools and crystallises. |
| Systematic | The solution was not fully saturated when sampled. | Leave longer with excess solid visible before taking the sample. |
| Judgement | Deciding when all the water has evaporated away. | Heat, cool and reweigh until two masses agree. |
| Random | Small errors in each of the three balance readings. | Repeat the whole determination and take a mean. |
Understand the science behind it → · Prove it · Teach it · Defend it
3. Required practical 1.36 — the formula of magnesium oxide
Specification statement 1.36 · Topic 1 · taught in Empirical and Molecular Formulae
Method
- Weigh a clean, dry crucible with its lid.
- Add a coil of cleaned magnesium ribbon and weigh again; the difference is the mass of magnesium.
- Heat the crucible strongly, lifting the lid briefly at intervals with tongs to admit air.
- Continue until the magnesium has fully reacted and no further change is seen.
- Allow the crucible to cool, then weigh it; the increase over the empty crucible is the mass of magnesium oxide.
- Reheat, cool and reweigh until two successive masses agree — heating to constant mass.
- Mass of oxygen = mass of oxide − mass of magnesium; divide each mass by its relative atomic mass and simplify the ratio.
Variables
Independent (changed) — The mass of magnesium ribbon used
Dependent (measured) — The mass of magnesium oxide formed
| Control variable | Why it must be held constant |
|---|---|
| The same crucible and lid throughout | a different crucible has a different mass, so the differences would be meaningless |
| Heating to constant mass | stopping early leaves unreacted magnesium and gives too little oxygen |
| Magnesium cleaned before use | an existing oxide layer adds oxygen that was never measured as part of the reaction |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | The lid is lifted too often and white smoke of magnesium oxide escapes. | Lift the lid only briefly and as seldom as possible. |
| Systematic | Heating stops before the reaction is complete, so some magnesium is unreacted. | Reheat, cool and reweigh until the mass is constant. |
| Random | Weighing the crucible while still warm gives a low reading from convection currents. | Cool the crucible fully before each weighing. |
Understand the science behind it → · Prove it · Teach it · Defend it
4. Required practical 1.60 — electrolysis of aqueous solutions
Specification statement 1.60 · Topic 1 · taught in Electrolysis
Method
- Half-fill a small beaker with the solution to be electrolysed, such as copper(II) sulfate or sodium chloride solution.
- Clamp two clean graphite electrodes in the solution so that they do not touch each other.
- Connect the electrodes to a d.c. power supply, noting which is joined to the negative terminal — that one is the cathode.
- If a gas is expected, invert a small water-filled test tube over each electrode to collect it.
- Switch on the supply and observe both electrodes, recording any coating, colour change or bubbling.
- Test any gas collected: a lit splint pops for hydrogen, a glowing splint relights in oxygen, and damp litmus paper is bleached white by chlorine.
- Switch off, and repeat with the other solutions so that the products can be compared.
Variables
Independent (changed) — The solution being electrolysed
Dependent (measured) — The products formed at the cathode and at the anode
| Control variable | Why it must be held constant |
|---|---|
| Inert graphite electrodes throughout | a reactive electrode takes part in the reaction and changes the products |
| The same voltage for each solution | a different voltage changes the rate, so the results would not be comparable |
| The same concentration of each solution | a very dilute halide can give oxygen at the anode instead of the halogen |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | The electrodes are not cleaned, so a previous product remains and is mistaken for a new one. | Clean both electrodes with emery paper before each run. |
| Judgement | A gas is identified by appearance alone when the bubbles look alike. | Use the splint and damp litmus tests every time. |
| Random | Gas escapes before the collecting tube is in position over the electrode. | Place the tubes before switching the supply on. |
Understand the science behind it → · Prove it · Teach it · Defend it
Topic 2 · Inorganic chemistry
5. Required practical 2.14 — the percentage of oxygen in air
Specification statement 2.14 · Topic 2 · taught in The Gases in the Air
Method
- Connect two gas syringes to either end of a silica tube containing copper turnings.
- Set one syringe to a measured volume of air, such as 100 cm³, and the other to zero, and record the total volume.
- Heat the copper strongly with a Bunsen burner.
- Pass the air slowly backwards and forwards over the hot copper using the two syringes, so that all of it makes contact repeatedly.
- Observe the copper turning black as copper(II) oxide forms.
- Continue until the total volume stops decreasing, showing that all the oxygen has been used.
- Allow the apparatus to cool completely, then read the final total volume.
- Calculate the percentage as the decrease in volume divided by the original volume, multiplied by 100.
Variables
Independent (changed) — The volume of air passed over the copper
Dependent (measured) — The decrease in the total volume of gas
| Control variable | Why it must be held constant |
|---|---|
| Excess copper present | too little copper leaves oxygen unreacted, so the answer comes out too low |
| Apparatus cooled before the final reading | a hot gas is expanded, so the decrease would appear smaller than it is |
| Sealed apparatus with no leaks | air drawn in during heating adds oxygen that was never in the original sample |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | The gas is passed over the copper too few times, so some oxygen never reacts. | Continue passing until the volume stops decreasing. |
| Systematic | The final volume is read while the apparatus is still hot, so the gas is expanded. | Cool the apparatus fully before the final reading. |
| Random | The syringe plunger sticks, so the volume readings vary between repeats. | Lubricate the plunger and repeat, taking a mean. |
Understand the science behind it → · Prove it · Teach it · Defend it
6. Required practical 2.21 — reactions of dilute acids with metals
Specification statement 2.21 · Topic 2 · taught in Redox
Method
- Measure equal volumes of the same dilute acid, such as 25 cm³ of hydrochloric acid, into several test tubes.
- Take equal masses of each metal — magnesium, zinc, iron and copper — in the same physical form, such as similar-sized pieces of ribbon or turnings.
- Add one metal to each tube and start a stopwatch.
- Observe and record the rate of fizzing in each tube, and note any temperature change by feeling the tube or using a thermometer.
- Test the gas produced with a lit splint; a squeaky pop confirms hydrogen.
- Rank the metals by how vigorously they react, most vigorous first.
- Note that any metal producing no fizzing at all lies below hydrogen in the reactivity series.
Variables
Independent (changed) — The metal added to the acid
Dependent (measured) — The rate of reaction, judged by how vigorously hydrogen is given off
| Control variable | Why it must be held constant |
|---|---|
| The same volume and concentration of acid | a more concentrated acid reacts faster, which would be mistaken for a more reactive metal |
| Surface area of the metal | a powder reacts faster than a lump of the same metal, so the comparison would be unfair |
| The same starting temperature | a warmer mixture reacts faster regardless of which metal is present |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Judgement | Rate of fizzing is judged by eye, so similar metals cannot be ranked reliably. | Time how long a fixed volume of gas takes to collect. |
| Systematic | An oxide layer on the metal delays the start of the reaction. | Clean each metal with emery paper before use. |
| Random | Pieces of metal differ in surface area between tubes, changing the rate. | Use pieces of the same size and shape. |
Understand the science behind it → · Prove it · Teach it · Defend it
7. Required practical 2.42 — hydrated copper(II) sulfate crystals
Specification statement 2.42 · Topic 2 · taught in Preparing a Soluble Salt
Method
- Measure about 25 cm³ of dilute sulfuric acid into a beaker and warm it gently with a Bunsen burner, without boiling.
- Add a spatula of black copper(II) oxide and stir; it dissolves and the solution turns blue.
- Continue adding copper(II) oxide a little at a time until some black solid remains undissolved even after stirring.
- This excess shows that all of the acid has reacted.
- Filter the mixture to remove the unreacted copper(II) oxide, collecting the blue filtrate.
- Pour the filtrate into an evaporating basin and heat gently until crystals begin to appear at the edges.
- Remove from the heat and leave the solution to cool and crystallise slowly.
- Filter off the blue crystals and dry them between two pieces of filter paper.
Variables
Independent (changed) — The mass of copper(II) oxide added to the acid
Dependent (measured) — Whether all the acid has reacted, judged by solid remaining undissolved
| Control variable | Why it must be held constant |
|---|---|
| Volume of sulfuric acid | a different volume of acid needs a different mass of oxide to use it all up |
| Concentration of the acid | a more concentrated acid would leave unreacted acid in the filtrate |
| Temperature of the acid | a hotter acid reacts faster, which changes how long the addition takes |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Judgement | Too little oxide is added, so unreacted acid remains and contaminates the crystals. | Add oxide until solid visibly remains undissolved. |
| Systematic | The solution is evaporated to dryness, driving off the water of crystallisation. | Stop heating when crystals first appear at the edges. |
| Random | Crystals are lost when transferring the filtrate between vessels. | Rinse each vessel into the next with a little water. |
Understand the science behind it → · Prove it · Teach it · Defend it
8. Required practical 2.43 — a pure dry sample of lead(II) sulfate
Specification statement 2.43 · Topic 2 · taught in Preparing an Insoluble Salt
Method
- Dissolve lead(II) nitrate in distilled water in one beaker, and sodium sulfate or magnesium sulfate in distilled water in another.
- Mix the two solutions together in a larger beaker and stir.
- A white precipitate of lead(II) sulfate forms immediately.
- Filter the mixture, so that the precipitate is collected as the residue on the filter paper.
- Wash the precipitate with distilled water while it is still in the filter funnel, to remove the soluble sodium nitrate solution.
- Repeat the washing so that all the soluble impurity is removed.
- Dry the precipitate in a warm oven or between two pieces of filter paper.
Variables
Independent (changed) — The two soluble solutions mixed together
Dependent (measured) — The mass of dry lead(II) sulfate obtained
| Control variable | Why it must be held constant |
|---|---|
| Volume of each solution | one reactant in excess is wasted and remains dissolved with the spectator ions |
| Use of distilled water | tap water contains dissolved ions that could precipitate and contaminate the product |
| Number of washings | too few leaves soluble sodium nitrate dried onto the crystals |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | The precipitate is not washed, so soluble spectator ions dry onto it. | Wash with distilled water in the funnel before drying. |
| Random | Fine precipitate passes through the filter paper and is lost. | Use finer filter paper and refilter the filtrate. |
| Judgement | Drying stops too soon, so the sample is still damp when weighed. | Dry to constant mass before weighing. |
Understand the science behind it → · Prove it · Teach it · Defend it
Topic 3 · Physical chemistry
9. Required practical 3.8 — measuring temperature changes
Specification statement 3.8 · Topic 3 · taught in Exothermic and Endothermic Reactions
Method
- Stand a polystyrene cup inside a glass beaker, which provides an extra layer of insulation.
- Measure a fixed volume of the first solution, such as 25 cm³, into the cup using a measuring cylinder.
- Record its starting temperature with a thermometer reading to at least 0.5 °C.
- Add the second reagent, put a lid on the cup and stir gently with the thermometer.
- Record the highest temperature reached for an exothermic reaction, or the lowest for an endothermic one.
- Calculate the temperature change by subtracting the starting temperature from the final one.
- Repeat for each of the four reaction types — dissolving, neutralisation, displacement and combustion — and compare.
Variables
Independent (changed) — The reaction being investigated
Dependent (measured) — The change in temperature of the solution
| Control variable | Why it must be held constant |
|---|---|
| Volume of solution used | a larger volume warms less for the same energy released, so the changes would not be comparable |
| Concentration of the solutions | a more concentrated solution releases more energy, which would be mistaken for a difference between reactions |
| Starting temperature | heat is lost faster from a mixture that starts further above room temperature |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | Heat escapes to the room, so every temperature rise measured is too small. | Use an insulated cup with a lid, inside a beaker. |
| Judgement | The final temperature is read before the mixture has been stirred, so it is uneven. | Stir gently and read the steady maximum or minimum. |
| Random | The thermometer is read at slightly different times between repeats. | Record the temperature at fixed intervals and take the peak. |
Understand the science behind it → · Prove it · Teach it · Defend it
10. Required practical 3.15 — marble chips and hydrochloric acid
Specification statement 3.15 · Topic 3 · taught in Rates of Reaction
Method
- Measure a fixed volume of dilute hydrochloric acid, such as 50 cm³, into a conical flask.
- Weigh out a fixed mass of marble chips of a known size.
- Add the chips to the acid, immediately stopper the flask with a delivery tube leading to a gas syringe, and start a stopwatch.
- Record the volume of carbon dioxide collected every 15 or 30 seconds until no more gas is produced.
- Repeat using the same mass of marble but in smaller pieces or as a powder, keeping the acid volume and concentration the same.
- Repeat again using acid of a different concentration, keeping the chip size the same.
- Plot volume of gas against time for each run on the same axes and compare the steepness of the curves.
Variables
Independent (changed) — The size of the marble chips, or the concentration of the acid
Dependent (measured) — The volume of carbon dioxide collected in a given time
| Control variable | Why it must be held constant |
|---|---|
| Mass of marble used | a greater mass produces more gas in total, which would be mistaken for a faster reaction |
| Volume of acid | a larger volume of acid contains more of it, so more gas is produced before it runs out |
| Temperature | a warmer mixture reacts faster whatever the chip size, so the comparison would be unfair |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | Gas escapes before the stopper is in place, so early readings are too low. | Have the stopper ready and fit it immediately. |
| Random | Chips described as the same size vary considerably in surface area. | Sieve the chips to a narrow size range before use. |
| Judgement | Readings are taken at slightly irregular intervals, distorting the curve. | Read the syringe on a fixed timer signal. |
Understand the science behind it → · Prove it · Teach it · Defend it
11. Required practical 3.16 — catalytic decomposition of hydrogen peroxide
Specification statement 3.16 · Topic 3 · taught in Catalysts
Method
- Measure a fixed volume of hydrogen peroxide solution, such as 50 cm³, into a conical flask.
- Run a control first with no solid added, collecting the oxygen produced in a fixed time so that the uncatalysed rate is known.
- Weigh out a fixed mass of the first solid, such as 1.0 g of manganese(IV) oxide, using a consistent particle size.
- Add the solid, immediately stopper the flask with a delivery tube to a gas syringe, and start the stopwatch.
- Record the volume of oxygen collected after a fixed time, or at regular intervals.
- Repeat with each solid in turn, keeping the volume and concentration of hydrogen peroxide, the mass and particle size of solid, and the temperature the same.
- At the end, filter off the solid, wash and dry it, weigh it and compare with the starting mass.
- Add the recovered solid to fresh hydrogen peroxide to confirm that it still works, showing it is chemically unchanged.
Variables
Independent (changed) — The solid added to the hydrogen peroxide
Dependent (measured) — The volume of oxygen collected in a fixed time
| Control variable | Why it must be held constant |
|---|---|
| Mass of solid added | more catalyst gives more surface for the reaction, so the comparison would be unfair |
| Particle size of the solid | a finer powder has more surface area and would appear the better catalyst for that reason alone |
| Concentration of hydrogen peroxide | a more concentrated solution decomposes faster whatever solid is present |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | No control without a solid is run, so any increase in rate cannot be demonstrated. | Always run the reaction with no solid first. |
| Random | Oxygen escapes while the stopper is being fitted after adding the solid. | Have the stopper ready and fit it immediately. |
| Judgement | Solids of different particle size are compared as though equivalent. | Grind and sieve every solid to the same size. |
Understand the science behind it → · Prove it · Teach it · Defend it
Topic 4 · Organic chemistry
12. Required practical 4.43 — preparing a sample of an ester
Specification statement 4.43 · Topic 4 · taught in Esters And How To Make One
Method
- Put about 2 cm³ of ethanol and about 2 cm³ of ethanoic acid into a boiling tube.
- Add three or four drops of concentrated sulfuric acid, using a teat pipette and wearing eye protection, since it is corrosive.
- Stand the tube in a beaker of hot water and leave it for about five minutes. Do not heat it over a flame, because ethanol and the ester are volatile and flammable.
- Pour the warm mixture into a beaker containing sodium carbonate solution, to neutralise the remaining acids.
- A colourless oily layer of ester floats on the surface.
- Smell it by wafting the vapour towards your nose with a hand; a sweet, fruity smell like pear drops confirms that ethyl ethanoate has formed.
Variables
Independent (changed) — The alcohol used, when the same preparation is repeated to make different esters
Dependent (measured) — The ester produced, identified by its characteristic smell
| Control variable | Why it must be held constant |
|---|---|
| Volume of carboxylic acid | using more acid gives more ester, so a difference in smell could not be attributed to the alcohol |
| Number of drops of catalyst | the catalyst sets the rate, so an unequal amount changes how much ester forms in the time given |
| Water-bath temperature | a warmer bath speeds the reaction, so more ester would form for reasons unconnected with the alcohol |
| Heating time | the reaction is reversible and approaches equilibrium slowly, so a longer time gives a larger yield |
Sources of error
| Type | What goes wrong | What to do |
|---|---|---|
| Systematic | Heating over a naked flame boils off the volatile ethanol and ester before they can react. | Heat the tube in a beaker of hot water instead. |
| Systematic | Unreacted ethanoic acid left in the mixture has a sharp smell that masks the ester. | Pour the product into sodium carbonate solution before smelling it. |
| Judgement | Identifying an ester by its smell is subjective, and people describe the same smell differently. | Have several people smell the sample independently and compare. |
Understand the science behind it → · Prove it · Teach it · Defend it
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