Mastering AS Chemistry Practical Skills: Required Practicals 1 to 6
Welcome to your complete guide to the AQA AS Chemistry practicals! Did you know that at least 15% of your total AS examination marks come directly from questions testing practical skills, apparatus, and data handling? Even though you might take these exams on paper, examiners will test whether you truly understand what happens at the lab bench.
Don't worry if practical chemistry has felt overwhelming before. In this guide, we break down every essential skill, the six required practical activities (RP1 to RP6), uncertainty calculations, and examiner warning points into step-by-step, bite-sized explanations.
---1. Essential Apparatus and Techniques (AT a to AT k)
Before diving into specific experiments, let's review the fundamental tools and techniques you need to recognize and explain:
• AT a: Measuring Quantities Accurately
Always select the right instrument: use a balance for mass, a stopwatch for time, a gas syringe or measuring cylinder for gas volume, a pipette or burette for liquid volume, and a thermometer for temperature changes.
• AT b: Safe Heating of Flammable Liquids
Many organic compounds (like alcohols) are highly flammable. Never heat flammable organic liquids directly with a naked Bunsen flame! Always use a water bath, an electric heating mantle, or a sand bath.
• AT c: Measuring pH
Measure \( \text{pH} \) using calibrated \( \text{pH} \) probes connected to data loggers, digital \( \text{pH} \) meters, or standard universal indicator \( \text{pH} \) colour charts.
• AT d & AT e: Core Laboratory Operations
Master volumetric glassware (pipettes, burettes, and volumetric flasks), heating under reflux, simple distillation, filtration (both standard gravity filtration with fluted filter paper and vacuum filtration under reduced pressure), and qualitative test-tube reactions.
• AT f & AT g: Indicators and Purification
Choose appropriate acid–base indicators (such as phenolphthalein or methyl orange) and carry out purification techniques including distillation, liquid–liquid separation, and recrystallisation.
• AT k: Chromatography
Use Thin Layer Chromatography (TLC) or paper chromatography to separate, identify, and monitor the purity of chemical mixtures.
2. RP1: Making a Standard Solution and Acid–Base Titration
Part A: Making a Standard Solution
A standard solution is simply a solution of known concentration. Preparing it accurately is crucial for reliable titration results.
Step-by-Step Procedure:
1. Weighing by difference: Weigh a sample bottle containing your solid on a precision 2 decimal place balance. Tip the solid into a clean beaker, re-weigh the empty sample bottle, and subtract the second mass from the first (\( \text{mass solid} = \text{mass dish + solid} - \text{mass emptied dish} \)). This accounts for any solid left sticking to the container.
2. Dissolve: Add deionised water (around \( 100\text{ cm}^3 \)) to the beaker and stir thoroughly with a glass rod until the solid has completely dissolved.
3. Quantitative transfer: Pour the solution into a \( 250\text{ cm}^3 \) volumetric flask using a funnel.
4. Rinsing: Rinse the beaker, glass rod, and funnel several times with deionised water, adding all washings into the volumetric flask.
5. Make up to the mark: Fill the flask with deionised water until the bottom of the meniscus sits exactly on the calibration line at eye level. Add the final drops with a dropping pipette.
6. Invert to mix: Insert the stopper and invert the flask repeatedly (at least 10–15 times) to ensure a uniform concentration throughout.
Part B: Carrying Out the Titration
Step-by-Step Technique:
1. Rinse glassware: Rinse the burette with the titrant (the solution going into it) and rinse the volumetric pipette with the analyte (the solution being sampled). Never rinse them with water immediately before filling, as residual water will dilute your solutions!
2. Fill and check: Fill the burette and ensure the jet space below the tap is filled with liquid and free from air bubbles.
3. Pipette analyte: Use a pipette filler to transfer exactly \( 25.0\text{ cm}^3 \) of your standard solution into a clean conical flask. Place the conical flask on a white tile to make the end-point colour change sharp and easy to spot.
4. Add indicator: Add just a few drops (\( 2 \text{ to } 3 \) drops) of indicator.
5. Titrate: Carry out a rough/trial titration first to find the approximate end-point. Then perform accurate titrations, adding titrant dropwise near the end-point while swirling continuously.
6. Readings: Read the burette at eye level from the bottom of the meniscus to two decimal places, ending in \( .00 \) or \( .05\text{ cm}^3 \).
7. Concordancy: Repeat until you obtain concordant titres—titres within \( 0.10\text{ cm}^3 \) of each other. Calculate your mean titre using only these concordant values.
Key Takeaway for RP1: Never include the rough trial or non-concordant titres in your mean. Always remove the funnel from the top of the burette before taking readings to avoid stray drops falling in!
---3. RP2: Measurement of an Enthalpy Change
In calorimetry, we measure the heat energy transferred during a chemical reaction using a simple calorimeter.
Why Use an Expanded Polystyrene Cup?
Polystyrene is a good thermal insulator and has a very low heat capacity. Adding a plastic lid further minimises heat loss to the surrounding air.
The Cooling Curve (Extrapolation) Method
For reactions in solution (such as a neutralisation or displacement reaction), heat is lost to the surroundings while the reaction proceeds. To correct for this:
1. Measure and record the initial temperature of your solution every minute for \( 3 \) minutes (at minutes 0, 1, 2, and 3).
2. At minute 4, add the second reagent, stir thoroughly, but do not record a temperature.
3. From minute 5 onwards, record the temperature every minute for up to 10–15 minutes.
4. Plot a graph of Temperature vs. Time.
5. Draw two lines of best fit: one through the initial temperatures and one through the cooling temperatures. Extrapolate both lines back to minute 4 to determine the theoretical instantaneous temperature change (\( \Delta T \)).
Calorimetry Calculations
• Calculate heat energy transferred: \( q = mc\Delta T \)
Where \( m \) is the mass of solution (assuming density \( = 1.00\text{ g cm}^{-3} \), so \( 1.0\text{ cm}^3 = 1.0\text{ g} \)), \( c \) is the specific heat capacity of water (\( 4.18\text{ J g}^{-1}\text{ K}^{-1} \)), and \( \Delta T \) is the extrapolated temperature change.
• Calculate molar enthalpy change: \( \Delta H = -\frac{q}{n \times 1000}\text{ kJ mol}^{-1} \)
Where \( n \) is the number of moles of the limiting reactant. Always remember the negative sign for exothermic reactions (temperature rise) and a positive sign for endothermic reactions (temperature drop)!
4. RP3: Investigating Rate of Reaction with Temperature
This investigation uses the classic "disappearing cross" reaction between sodium thiosulfate and hydrochloric acid:
\( \text{Na}_2\text{S}_2\text{O}_3\text{(aq)} + 2\text{HCl(aq)} \rightarrow 2\text{NaCl(aq)} + \text{SO}_2\text{(g)} + \text{S(s)} + \text{H}_2\text{O(l)} \)
Method:
1. Place a conical flask over a white piece of paper marked with a clear black cross.
2. Warm the solutions to the desired temperature using a water bath.
3. Mix the reactants and start the timer immediately.
4. Look down through the top of the flask and stop the timer the exact moment the yellow solid sulfur precipitate (\( \text{S(s)} \)) obscures the black cross.
5. Repeat at several different temperatures while keeping all volumes and concentrations constant.
Data Processing:
• The amount of sulfur formed to obscure the cross is fixed, so the initial rate is proportional to \( \frac{1}{t} \) (where \( t \) is time in seconds).
• Plotting \( \frac{1}{t} \) against temperature (\( T \)) shows an exponential curve.
• Plotting \( \ln(\text{rate}) \) or \( \ln\left(\frac{1}{t}\right) \) against \( \frac{1}{T} \) (where \( T \) is in Kelvin) produces an Arrhenius plot with a straight line.
5. RP4: Identifying Inorganic Ions by Simple Test-Tube Reactions
Master these qualitative tests to identify cations and anions in unknown samples:
A. Testing for Cations (Positive Ions)
• Magnesium (\( \text{Mg}^{2+} \)): Add dilute \( \text{NaOH}\text{(aq)} \) dropwise until in excess \( \rightarrow \) forms a thick white precipitate of \( \text{Mg(OH)}_2 \) (sparingly soluble). With dilute \( \text{H}_2\text{SO}_4\text{(aq)} \) \( \rightarrow \) remains colourless (no precipitate formed).
• Barium (\( \text{Ba}^{2+} \)): Add dilute \( \text{NaOH}\text{(aq)} \) \( \rightarrow \) remains colourless (no precipitate formed). With dilute \( \text{H}_2\text{SO}_4\text{(aq)} \) \( \rightarrow \) forms a thick white precipitate of \( \text{BaSO}_4 \) (insoluble).
• Ammonium (\( \text{NH}_4^+ \)): Add dilute \( \text{NaOH}\text{(aq)} \) and warm gently in a water bath \( \rightarrow \) releases pungent ammonia gas (\( \text{NH}_3\text{(g)} \)), which turns damp red litmus paper blue.
B. Testing for Anions (Negative Ions)
• Halides (\( \text{Cl}^- \), \( \text{Br}^- \), \( \text{I}^- \)): First acidify with dilute nitric acid (\( \text{HNO}_3 \)) to remove any carbonate ions, then add silver nitrate solution (\( \text{AgNO}_3\text{(aq)} \)):
- Chloride (\( \text{Cl}^- \)) \( \rightarrow \) White precipitate (\( \text{AgCl} \)); dissolves in dilute \( \text{NH}_3\text{(aq)} \).
- Bromide (\( \text{Br}^- \)) \( \rightarrow \) Cream precipitate (\( \text{AgBr} \)); insoluble in dilute \( \text{NH}_3 \), but dissolves in concentrated \( \text{NH}_3\text{(aq)} \).
- Iodide (\( \text{I}^- \)) \( \rightarrow \) Yellow precipitate (\( \text{AgI} \)); insoluble in both dilute and concentrated \( \text{NH}_3\text{(aq)} \).
Memory Trick: White, Cream, Yellow = Chloride, Bromide, Iodide (Alphabetical order by halide matches alphabetical order by precipitate color: C \(\rightarrow\) B \(\rightarrow\) I corresponds to W \(\rightarrow\) C \(\rightarrow\) Y).
• Carbonate (\( \text{CO}_3^{2-} \)): Add dilute hydrochloric acid (\( \text{HCl} \)) or nitric acid (\( \text{HNO}_3 \)) \( \rightarrow \) effervescence (bubbling). Bubbling the gas through limewater turns it cloudy due to \( \text{CO}_2 \).
• Sulfate (\( \text{SO}_4^{2-} \)): Acidify with dilute \( \text{HCl} \) (to destroy carbonates), then add barium chloride solution (\( \text{BaCl}_2\text{(aq)} \)) \( \rightarrow \) forms a dense white precipitate of \( \text{BaSO}_4 \).
• Hydroxide (\( \text{OH}^- \)): Alkaline solution turns red litmus paper blue or universal indicator solution purple.
6. RP5: Distillation of an Organic Product (Oxidation of Ethanol)
In this practical, primary alcohol ethanol is partially oxidised to ethanal (an aldehyde) using acidified potassium dichromate(\( \text{VI} \)):
\( \text{CH}_3\text{CH}_2\text{OH} + [\text{O}] \rightarrow \text{CH}_3\text{CHO} + \text{H}_2\text{O} \)
Key Experimental Setup Rules
• Apparatus: Pear-shaped or round-bottom flask, still head, thermometer, Liebig condenser, receiver adapter, and collecting vessel.
• Anti-bumping granules: Added to the flask before heating. They provide nucleation sites to promote smooth boiling and prevent violent boiling/spattering.
• Thermometer placement: The thermometer bulb must sit directly opposite the side-arm opening to measure the true boiling point of the vapor entering the condenser.
• Condenser water flow: Cold water must enter at the bottom and exit at the top. This ensures the condenser jacket remains completely filled with water, providing maximum cooling efficiency.
• Open system: The apparatus must never be sealed tightly; a closed system builds dangerous pressure upon heating.
• Ice bath collection: Place the receiving flask in an ice bath to condense and preserve the volatile ethanal vapor (boiling point \( 21^\circ\text{C} \)) and prevent it from evaporating.
7. RP6: Tests for Organic Functional Groups
• Alkene: Shake with bromine water (\( \text{Br}_2\text{(aq)} \)) \( \rightarrow \) colour changes from orange/brown to colourless (decolourisation).
• Aldehyde vs. Ketone:
- Tollens' Reagent (ammoniacal silver nitrate, warmed): Aldehydes form a reflective silver mirror on the inside of the tube; ketones show no reaction.
- Fehling's / Benedict's Solution (warmed): Aldehydes change the deep blue copper(\( \text{II} \)) solution into a brick-red precipitate of \( \text{Cu}_2\text{O} \); ketones show no reaction (remain blue).
• Carboxylic Acid: Add solid sodium hydrogencarbonate (\( \text{NaHCO}_3 \)) or aqueous sodium carbonate (\( \text{Na}_2\text{CO}_3 \)) \( \rightarrow \) effervescence produced (\( \text{CO}_2 \) gas turns limewater cloudy).
• Alcohols (Primary and Secondary vs. Tertiary): Warm with acidified potassium dichromate(\( \text{VI} \)) (\( \text{K}_2\text{Cr}_2\text{O}_7 / \text{H}_2\text{SO}_4 \)) \( \rightarrow \) orange solution (\( \text{Cr}_2\text{O}_7^{2-} \)) turns green (\( \text{Cr}^{3+} \)) for primary and secondary alcohols. Tertiary alcohols cannot be oxidised and remain orange.
8. Measurement Uncertainty and Data Handling
Calculating Percentage Uncertainty
Every laboratory measurement has an associated absolute uncertainty:
• Analogue scale: \(\pm \frac{1}{2}\) of the smallest scale division (e.g., a measuring cylinder with \( 1\text{ cm}^3 \) divisions has an uncertainty of \( \pm 0.5\text{ cm}^3 \)).
• Digital scale: \(\pm\) the last decimal place (e.g., a 2 d.p. balance has an uncertainty of \( \pm 0.01\text{ g} \)).
Use the formula:
\( \text{Percentage Uncertainty} = \left(\frac{\text{Absolute Uncertainty}}{\text{Measured Value}}\right) \times 100 \)
Measurements by Difference
When a reading requires an initial and a final value (such as a burette titre or weighing by difference), the uncertainty applies to both readings:
\( \text{Total Absolute Uncertainty} = 2 \times (\text{Uncertainty of a single reading}) \)
Example: If a burette has a single-reading uncertainty of \( \pm 0.05\text{ cm}^3 \), a titre of \( 20.00\text{ cm}^3 \) has a combined uncertainty of \( \pm 0.10\text{ cm}^3 \):
\( \text{Percentage Uncertainty} = \left(\frac{0.10}{20.00}\right) \times 100 = 0.50\% \)
How to Reduce Percentage Uncertainty
To reduce percentage uncertainty without buying more expensive equipment, increase the magnitude of the measured quantity (the denominator):
• In titrations: Use a larger mass of solid or lower titrant concentration to obtain a larger titre volume.
• In calorimetry: Use larger temperature changes.
Graphing and Significant Figures Rules
• Significant Figures: Always give your final calculated answers to the same number of significant figures as the raw measurement with the lowest number of significant figures.
• Graph Plotting: The independent variable belongs on the \( x \)-axis and the dependent variable on the \( y \)-axis. Plotted points must occupy at least 50% of the graph grid. Mark points with neat small crosses (\( + \) or \( \times \)), and identify and ignore anomalous points when drawing a smooth line or curve of best fit.
9. Common Exam Pitfalls to Avoid
1. Titrations: Averaging non-concordant titres. Always ignore the rough trial and only average concordant values within \( 0.10\text{ cm}^3 \) of each other.
2. Rinsing: Rinsing the burette or pipette with deionised water instead of the reagent being measured. This dilutes the solution and introduces systematic error.
3. Calorimetry: Forgetting the negative sign (\( - \)) for exothermic enthalpy changes (\( \Delta H \)).
4. Distillation Setup: Drawing water flowing into the top of a Liebig condenser instead of the bottom, or placing the thermometer bulb submerged in the boiling liquid.
5. Qualitative Tests: Acidifying halide tests with hydrochloric acid (\( \text{HCl} \)) or sulfuric acid (\( \text{H}_2\text{SO}_4 \)). This introduces chloride or sulfate ions, producing false positive precipitates! Always use nitric acid (\( \text{HNO}_3 \)).