A Guide to A level Chemistry Practicals

by Muhammad Haris | 17 Oct 2025

So there you are, taking your first ever lab practical lesson at school; the room reeks of something pungent, and you’ve got 12 different instruments laid out before you, most of which you can’t name. The question paper tells you to add some “FA1” to a beaker, but you’re scared you’ll spill it all…

There’s no denying that practicals can seem intimidating for a lot of newcoming AS students, but lucky for you, we’ve got you all sorted! Presenting, Mojza’s carefully crafted guide to A Level Chemistry practicals!

Paper Pattern

Your practical component for Chemistry is your paper 3; it’s worth 40 marks and allotted a time of 2 hours, often consisting of 3 questions.

Titrations

This is perhaps the most crucial skill to master for your chemistry practicals, as its question shows up on the exam every single year.

You will be given an acidic and an alkaline solution, one of a known concentration and the other of an unknown one. The purpose of a titration is to react the acid with the base in exact volumes in the presence of an indicator till neutralization, and then determine the concentration of the unknown solution with the help of stoichiometric calculations.

There are 3 types of titrations you can be asked to perform:

Acid-Base

Reacting an Acidic solution with an Alkaline one, e.g. HCl with NaOH.

Redox

Involves the use of acidified aqueous Potassium Manganate (VII) solution (KMnO4) in place of an acid.

Iodometric

Rather different from the other two; its purpose is to determine the concentration of an aqueous iodide salt, such as Potassium Iodide, by reacting it with a thiosulfate solution. This type specifically uses a starch indicator.

Here are some important guidelines for titrations:

1. Reading The Meniscus Correctly

When taking measurements from a burette or pipette, make sure to read their scales at eye level to avoid parallax error.

2. Rinsing Instruments

Make sure to rinse your burette, pipette and conical flask with the correct liquid; first with distilled water and then with the solution you are about to use, in order to avoid residual water from diluting your solution and thus affecting your molar calculations.

3. Adding solution dropwise near the end point

As you approach a colour change, continuously swirl the conical flask to which you’re adding the burette solution, adding it drop by drop to avoid excess. Stop as soon as you notice a complete color change in the indicator.

4. Table of Results

Cambridge is very particular with the format of the table of results required in a titration question; out of 7 marks, 2-3 marks are for your table alone. It should look something like this:

As you can see, your table must include rows for your initial and final burette readings, your titre volume (the volume of burette solution used), as well as one for your concordant readings.

Your concordant readings are the readings which are within 0.1 cm3 of each other. You will be using these to calculate an average volume to later use in your stoichiometric workings.

Note that all readings for volume are written correct to 2 decimal places, as that is the maximum accuracy of a burette’s scale. Make sure to include this step, as not doing so would cost you marks.

Additionally, and this cannot be stressed enough: ALWAYS include units for your rows/columns, like the cm3 used above for all volume readings. These are the easiest marks to gain on your exam.

For a more in-depth understanding of titrations in context of past paper questions, here are some great tutorials:

Enthalpy Changes

Enthalpy change questions have you perform one or more chemical reactions, to then calculate their Enthalpy change yourself. To do this, you should know your formula for Enthalpy change, that is

Q = (mCΔT) / (n × 1000)

Where Q is energy gained/lost (in Joules),

m is the mass (in grams) or volume (in cm3) of solution,

C is the specific heat capacity of water (4.18 J/g°C),

and Δϑ is the change in temperature (in Kelvin—note that the change is the same in any unit),

n is the number of moles of the limiting reactant; that is completely reacted first,

and we divide by 1000 to convert from Joules to Kilojoules.

Here are some crucial tips for such questions, as well as commonly made mistakes:

1. Using a Thermometer

To ensure your reading for temperature is accurate, try to

Keep the thermometer’s bulb fully submerged in the solution
Keep stirring the mixture with the thermometer until you reach a steady temperature
Read its scale at eye level to avoid parallax error

Additionally, when noting your readings for temperature, make sure to write them to the correct number of decimal places, depending on your thermometer’s accuracy. This would most likely be one decimal place, so record them accordingly.

2. Determining if a reaction is Exo/Endothermic

To know this, you simply have to note how the reaction mixture’s temperature changes as the reaction proceeds; if it drops from the original temperature, the reaction is Endothermic. If it rises above the original temperature, then the reaction is exothermic.

An Endothermic enthalpy change has a positive sign, while an Exothermic one always has a negative sign.

3. Using the Formula

There are a number of guidelines to keep in mind when substituting values into this formula, which students often get wrong.

Firstly, remember that ‘m’ represents the mass of your solution, i.e. your reaction mixture, and NOT the mass of any single reactant used in the procedure which is what a lot of students erroneously plug in. The mass of your solution is simply equal to its volume, as 1g of water (your solvent) is taken to be equal to 1 cm3 of water.

Other than that, also keep in mind that the product from this formula isn’t your final answer; you still have to add a sign ( + or – ) to signify if the Enthalpy change is Exo/Endothermic. + if Endothermic, – if Exothermic.

4. Sources of Heat Loss/Error

It is common for the last part of this question to ask you to list ways in which heat energy could’ve escaped the container, and/or ways to prevent this. A common source of error in such an experiment is as follows:

Heat energy escapes from the cup’s top as well as its sides; use a lid to cover the cup, and use a cup stacked into another in order to reduce heat being lost to the surroundings.

Gravimetric Analysis

While the name might sound intimidating, gravimetric analysis simply means to determine the amount of a required substance after isolating it through a method like heating.

These questions would often ask you to heat a solid, hydrated salt (e.g. Hydrated Copper (II) Sulfate ) in a crucible over a bunsen flame, recording its mass before and after, and possibly repeating this process before finally calculating moles of the final anhydrous salt.

Alternatively, such questions may also ask you to calculate the value—the mass or the moles—if the water of Crystallisation in the originally presented hydrated salt. Therefore, once again, make sure to keep your stoichiometric skills sharp.Here’s another tutorial to see what these questions look like yourself.

Rate of Reaction

Questions testing this topic are rather straightforward; you’d be demanded to measure how the concentration of a reagent impacts its rate of reaction. You’ll have to prepare multiple concentrations of a reagent by dilution, and perform the reaction with each one, recording the time taken for the reaction mixture to cloud-up until it is fully opaque.

If the experiment is conducted correctly, you’ll observe a trend in the time taken for the solution to become opaque, as the reagent’s concentration changes; the question often goes on to ask the candidate to include a column for the rate of reaction in their table of results, as well as plot them on a graph to draw a line of best fit, representing the trend.

You can refer to this past paper walkthrough to build a better understanding.

Salt Analysis

Much like titrations, this is another chapter highly likely to make an appearance on your exam.

You will be provided with multiple aqueous and solid salts, and by conducting all the tests as directed in the question, you will eventually have to determine the identity of the given compounds.

This question specifically tests your ability to note observations; whether a solution “turns clear” or “goes cloudy/milky”; difference in observations when adding in a reagent dropwise Vs adding it in excess; recognizing pungent smells of gases like Ammonia and Sulfur Dioxide, as well as being familiar with observations for tests of common gases (e.g. how Hydrogen gas makes a burning matchstick go out with a squeaky pop sound).

Be sure to go through your salt analysis notes, going over the observations for the tests for different anions and cations. While these will be provided during your exam, it’s always better to be familiar with them beforehand.

Organic Analysis

This question type shows up in place of the Salt Analysis mentioned above or makes up a part of it; it’s a direct practical extension of the same Organic Chemistry you will go over for your theory components. Therefore, be sure to have a strong grip over it as you can be tested over a multitude of concepts, mostly encircling Aldehydes, Ketones, Alcohols and Carboxylic Acids and their many typical chemical reactions such as oxidation and reduction.

Fortunately though, this question rarely makes an appearance, since the last time it did come up on the exam was around 10 years ago.

General Tips

Time management is crucial in a practical exam. It’s easy to get caught up in repeating a step or rechecking your results, but remember that every minute counts. Plan your work in stages: set a target for when you should finish setting up, when to start taking results, and when to move on to calculations or graph plotting. This will help you avoid a last-minute rush and ensure you have enough time to check your answers before handing in your paper.

Always place your exam paper or any written material on a clean, dry part of your bench, well away from chemicals, water, or flame sources. Spilled liquids or stains can damage your paper, making it unreadable for examiners, and in worst cases, you could lose marks. Keeping your workspace organized will also reduce confusion, minimize accidents, and help you focus better.

Glassware, sharp instruments, and electrical equipment are delicate and can be hazardous if mishandled. Always set up your apparatus steadily, avoid unnecessary movements that could knock things over, and check that everything is clamped securely before starting.

One of the most common mistakes in practical exams is forgetting to include units in your tables. Every column should have both the quantity and its unit clearly labeled, for example: Time / s or Volume of gas / cm³. Omitting these steps cost you arguably the easiest marks on the paper.

All the nerd-talk aside, in the end you also have to accept that practicals are quite prone to mishaps and accidents; you may spill some chemicals all over your question paper, or break some apparatus. Thankfully, paper 3 only holds a weightage of 11.5% on your final A level grade, so calm yourself down if you mess something up on the practical, and be at ease about your results if you did well on the theoretical components, since those hold the most weightage.

Good luck with your A levels, people! You’ve got this!

Author: Muhammad Haris

April 30, 2026

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