Welcome to Practical Skills in AS Chemistry!
Practical work is at the very heart of chemistry. At AS Level (OCR Chemistry A - H032), practical skills are assessed in written examination papers (Breadth in chemistry and Depth in chemistry). In fact, at least 15% of the total exam marks test your direct understanding of laboratory experiments, apparatus, data handling, and evaluation.
Don't worry if laboratory questions have felt intimidating in the past. This guide breaks down the core concepts into simple, bite-sized principles that will help you secure maximum marks.
1. Planning and Experimental Design
Good experiments start with clear planning. When an exam question asks you to design or evaluate an experiment, always structure your thoughts around variables and method validity.
Understanding Experimental Variables
• Independent Variable: The factor you deliberately change (e.g., the concentration of an acid).
• Dependent Variable: The factor you measure as a result (e.g., the volume of gas produced in 30 seconds).
• Control Variables: All other conditions that must be kept strictly constant to ensure a fair test (e.g., temperature, total reaction volume, mass of catalyst, surface area of solid reactants).
Evaluating Apparatus and Method
When choosing equipment, match the tool to the precision required. For example, measuring \(25.0\text{ cm}^3\) of a solution requires a volumetric pipette, not a standard measuring cylinder, because pipettes have significantly lower percentage uncertainty.
Quick Summary: Change only the independent variable, measure the dependent variable, and keep all control variables identical.
2. Measurements, Units, and Data Presentation
Standard Units in Chemistry
Always record measurements with appropriate scientific units:
• Volume: \(\text{cm}^3\) or \(\text{dm}^3\) (Recall: \(1\text{ dm}^3 = 1000\text{ cm}^3\))
• Mass: \(\text{g}\)
• Amount of substance: \(\text{mol}\)
• Temperature: \(^{\circ}\text{C}\) or \(\text{K}\)
• Energy / Enthalpy: \(\text{kJ mol}^{-1}\)
• Concentration: \(\text{mol dm}^{-3}\)
• Pressure: \(\text{kPa}\)
• Time: \(\text{s}\)
Tables and Data Recording
When presenting data in tables:
1. Every column heading must contain both the quantity and the unit (e.g., Time / s or Volume / \(\text{cm}^3\)).
2. Data within a column must be recorded consistently to the same number of decimal places matching the precision of the measuring instrument.
3. Data Analysis and Graphing
Plotting Graphs Accurately
Graph questions carry easy marks if you follow these rules:
• Axes & Scales: Use regular, linear scales. Your plotted points must occupy more than half of the graph grid in both horizontal and vertical directions.
• Best-Fit Line: Draw a single, smooth straight line or curve through the points. Never join dots with short, jagged lines ("dot-to-dot").
• Calculating Gradients: Use the formula \(\text{Gradient} = \frac{\Delta y}{\Delta x}\). Always construct a large triangle where the hypotenuse covers at least 50% of the drawn line.
4. Evaluation: Precision, Accuracy, and Uncertainties
Accuracy vs Precision: What is the Difference?
• Accuracy: How close your measured value is to the true value.
• Precision: How close repeated, independent measurements are to one another under identical conditions.
Calculating Percentage Uncertainty
Every piece of laboratory apparatus has a built-in margin of error (uncertainty). You calculate its impact on your result using:
\(\text{Percentage Uncertainty} = \left( \frac{\text{Apparatus Uncertainty} \times \text{Number of Readings}}{\text{Measured Value}} \right) \times 100\)
Single vs Double Readings (A Crucial Exam Rule!)
• Single Reading (\(\text{Number of Readings} = 1\)): Used when the instrument has a fixed zero or is pre-calibrated. Example: Measuring a single volume in a \(250\text{ cm}^3\) volumetric flask.
• Double Reading (\(\text{Number of Readings} = 2\)): Used whenever you take an initial and a final reading to determine a change. Example: A burette (initial reading and final reading), a thermometer (initial temperature and final temperature), or weighing by difference on a balance (mass of boat + solid, and mass of empty boat).
Example: A student uses a balance with an uncertainty of \(\pm 0.005\text{ g}\) to measure a sample by difference. The recorded mass of solid is \(2.50\text{ g}\).
\(\text{Percentage Uncertainty} = \left( \frac{0.005 \times 2}{2.50} \right) \times 100 = 0.40\%\)
How to Reduce Percentage Uncertainty
• Use measuring equipment with greater precision / smaller uncertainty (e.g., a balance reading to 3 decimal places instead of 2).
• Increase the size of the measurement (e.g., use a larger mass of solid or a larger titre volume).
Key Takeaway: If a measurement involves calculating a difference (Final \(-\) Initial), always multiply the apparatus uncertainty by \(2\).
5. Core AS Practical Techniques
Technique 1: Volumetric Analysis (Titrations)
Titration is used to determine the concentration of an unknown solution.
Preparing a Standard Solution:
1. Weigh the solid accurately in a weighing boat, transfer it to a beaker, and reweigh the empty boat (weighing by difference).
2. Dissolve the solid in a beaker using a minimum volume of distilled water and stir with a glass rod.
3. Transfer the solution into a volumetric flask using a funnel.
4. Rinse the beaker, glass rod, and funnel with distilled water and add all washings to the flask.
5. Make up to the graduation mark with distilled water until the bottom of the meniscus touches the line. Invert the flask repeatedly to mix thoroughly.
Carrying Out the Titration:
• Rinse the pipette with the solution it will measure, and rinse the burette with the titrant.
• Read the burette at eye level at the bottom of the meniscus.
• Always record burette readings to 2 decimal places ending in \(.00\) or \(.05\text{ cm}^3\) (e.g., \(22.45\text{ cm}^3\) or \(24.00\text{ cm}^3\)).
Calculating the Mean Titre:
• Only use concordant titres — results within \(\pm 0.10\text{ cm}^3\) of each other.
• Never include the initial rough / trial titre or non-concordant values in your average calculation.
Technique 2: Organic Synthesis and Purification
1. Heating under Reflux
Reflux allows prolonged heating of organic mixtures to boiling point without losing volatile reactants or products through evaporation.
• The condenser is fitted vertically above the round-bottom or pear-shaped flask.
• Water enters at the bottom of the condenser jacket and exits at the top to ensure the jacket fills completely.
• Anti-bumping granules are added to the flask before heating to promote smooth boiling and prevent large bubbles.
• Vital Safety Point: The top of the condenser must remain open to prevent dangerous pressure build-up.
2. Simple Distillation
Used to separate liquids with significantly different boiling points.
• The thermometer bulb must sit level with the condenser side-arm entrance to record the exact temperature of condensing vapour.
• Water enters at the bottom of the condenser and leaves at the top.
3. Using a Separating Funnel & Drying Agents
• Used to separate two immiscible liquids (e.g., an aqueous layer and an organic layer).
• Invert the funnel and open the tap periodically to release built-up gas pressure.
• Collect the organic layer and add an anhydrous inorganic salt (such as anhydrous \(\text{MgSO}_4\), \(\text{CaCl}_2\), or \(\text{Na}_2\text{SO}_4\)) to remove traces of dissolved water.
Technique 3: Calorimetry (Enthalpy Determinations)
Calorimetry measures the heat exchanged during a chemical reaction using the formula:
\(q = mc\Delta T\)
Where \(q\) is heat energy exchanged (\(\text{J}\)), \(m\) is the mass of the solution (\(\text{g}\)), \(c\) is the specific heat capacity (\(4.18\text{ J g}^{-1}\text{ K}^{-1}\)), and \(\Delta T\) is the temperature change (\(\text{K}\) or \(^{\circ}\text{C}\)).
Correcting for Heat Loss (Cooling Curves):
In real experiments, heat is lost to the surroundings immediately as the reaction proceeds. To find the true theoretical temperature change:
1. Record temperature every minute before mixing (e.g., minutes 1 to 3).
2. Mix reactants at minute 4 (do not take a reading at minute 4).
3. Record temperature every minute from minute 5 onward as the mixture cools.
4. Plot temperature against time and extrapolate the cooling curve back to the exact minute of mixing (minute 4) to determine the theoretical maximum temperature change \(\Delta T\).
6. Qualitative Inorganic Analysis
Identifying unknown ions in solution requires systematic testing. Memorise the specific tests and observations below.
1. Carbonate Test (\(\text{CO}_3^{2-}\))
• Procedure: Add dilute acid (e.g., \(\text{HNO}_3\) or \(\text{HCl}\)).
• Observation: Effervescence (bubbling) of carbon dioxide gas, \(\text{CO}_2\), which turns limewater cloudy.
2. Sulfate Test (\(\text{SO}_4^{2-}\))
• Procedure: Add dilute acid (\(\text{HCl}\) or \(\text{HNO}_3\)), then aqueous barium chloride, \(\text{BaCl}_2\) (or barium nitrate, \(\text{Ba(NO}_3)_2\)).
• Observation: A dense white precipitate of barium sulfate, \(\text{BaSO}_4\), is formed.
3. Halide Ion Tests (\(\text{Cl}^-\), \(\text{Br}^-\), \(\text{I}^-\))
• Procedure: Add dilute nitric acid (\(\text{HNO}_3\)), then add aqueous silver nitrate (\(\text{AgNO}_3\)).
• Chloride (\(\text{Cl}^-\)): White precipitate (\(\text{AgCl}\)) \(\rightarrow\) Dissolves in dilute aqueous ammonia (\(\text{NH}_3\)).
• Bromide (\(\text{Br}^-\)): Cream precipitate (\(\text{AgBr}\)) \(\rightarrow\) Insoluble in dilute \(\text{NH}_3\), but dissolves in concentrated \(\text{NH}_3\).
• Iodide (\(\text{I}^-\)): Yellow precipitate (\(\text{AgI}\)) \(\rightarrow\) Insoluble in both dilute and concentrated \(\text{NH}_3\).
4. Ammonium Ion Test (\(\text{NH}_4^+\))
• Procedure: Warm the sample with aqueous sodium hydroxide (\(\text{NaOH}\)).
• Observation: Ammonia gas (\(\text{NH}_3\)) is evolved, which turns damp red litmus paper blue.
The Correct Sequence for Testing Unknown Anions
When testing a single unknown solution or mixture for anions, you must carry out the tests in this strict order:
1. Carbonate Test \(\rightarrow\) 2. Sulfate Test \(\rightarrow\) 3. Halide Test
Mnemonic: Carbonate, Sulfate, Halides \(\rightarrow\) "Chemistry Students Help!"
Why is the order essential?
• Why Carbonate first? Neither barium ions (\(\text{Ba}^{2+}\)) nor silver ions (\(\text{Ag}^+\)) can be added first because both \(\text{BaCO}_3\) and \(\text{Ag}_2\text{CO}_3\) are insoluble white precipitates that create false positive results.
• Why Sulfate second? Silver sulfate (\(\text{Ag}_2\text{SO}_4\)) is partially insoluble. If you added silver nitrate first, sulfate ions might form a precipitate and give a false positive for halides.
7. Common Exam Pitfalls & Examiner Tips
• Burette Readings Format: Never write \(23\text{ cm}^3\) or \(23.2\text{ cm}^3\). Always record to two decimal places: \(23.00\text{ cm}^3\) or \(23.20\text{ cm}^3\).
• Reflux Diagrams: Never draw a stopper closing the top of a reflux condenser. This creates a dangerous sealed vessel and automatically loses diagram marks.
• Condenser Water Connections: Always show water entering at the bottom and leaving from the top.
• Thermometer Position in Distillation: Ensure the bulb is drawn exactly opposite the entrance to the condenser side-arm.
• Vague Method Improvements: Avoid saying "be more careful" or "repeat it better". Always give specific improvements: "Use a 3-decimal-place balance to lower percentage uncertainty" or "Use a polystyrene cup with a lid to minimise heat loss to the surroundings."