Chemistry skills

Life and Health Sciences 0008 · CCEA AS Level · 9 min read · Updated: 31 Aug 2026

Introduction to Chemistry Skills in AS 1

Welcome to the Chemistry Skills study guide for Unit AS 1: Experimental Techniques! In CCEA Life and Health Sciences, this unit is assessed through an internally marked portfolio that accounts for 33.34% of your AS Award (and 13.34% of your overall A Level). Developing excellent laboratory technique and confident calculation skills is the key to creating a top-scoring portfolio.

Don't worry if quantitative chemistry or lab calculations have felt intimidating in the past. We will break down every technique step by step, showing you exactly what to do at the lab bench, how to process your data, and how to avoid the common errors that cost students easy marks.

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1. Volumetric Analysis: Standard Solutions & Titrations

Volumetric analysis involves using accurately measured volumes of solutions of known concentration to find the concentration of an unknown solution.

Step 1: Preparing a Standard Solution

A standard solution is a solution whose concentration is accurately known. To prepare one, a primary standard (a stable, pure solid) is weighed and dissolved to an exact volume.

Procedure:
1. Weighing by difference: Place your solid in a weighing boat or dish on a 2-decimal or 3-decimal analytical balance. Record the initial mass. Tip the solid into a clean beaker and re-weigh the empty dish. The actual mass transferred is calculated as:
$m_{\text{transferred}} = m_{\text{dish + substance}} - m_{\text{dish residue}}$
Why? This accounts for any tiny grains left behind in the dish!
2. Dissolving: Add a small volume of deionised water to the beaker and stir with a glass rod until the solid has fully dissolved.
3. Quantitative transfer: Pour the solution into a volumetric flask using a funnel. Rinse the beaker, glass rod, and funnel with deionised water into the flask so not a single molecule of solute is lost.
4. Making up to the mark: Add deionised water until the bottom of the meniscus sits exactly on the calibration line at eye level. If you overshoot the mark, you must start again!
5. Inversion: Stopper the flask and invert it repeatedly (at least 10 to 15 times) to ensure the concentration is uniform throughout.

Step 2: Carrying Out an Acid-Base Titration

A titration allows you to find the volume of one solution (the titrant) needed to react completely with a known volume of another (the analyte).

Setting Up the Apparatus Correctly:
• Burette: Rinse with deionised water, then rinse with the titrant (the solution going inside it). Fill the burette and open the tap briefly to ensure the jet space below the tap is completely filled with liquid and contains no air bubbles.
• Volumetric Pipette: Rinse with deionised water, then rinse with the analyte solution. Use a pipette filler to measure the exact aliquot into a clean conical flask.
• Conical Flask: Rinse only with deionised water. (Rinsing with the analyte would add extra moles and ruin your results!). Place the flask on a white tile so you can see the indicator colour change clearly.

Step 3: Titration Technique & Concordant Titres

• Carry out a rough titration first to find the approximate endpoint.
• Carry out subsequent accurate titrations, adding titrant dropwise near the endpoint while continuously swirling the conical flask.
• Burette readings must be recorded to 2 decimal places, with the second decimal place ending in either .00 or .05 $\text{cm}^3$.
• Continue titrating until you achieve concordant titres—meaning titres that are within $0.10\text{ cm}^3$ of each other.
• Golden Rule: Only calculate the mean (average) using the concordant titres. Never include the rough titration or non-concordant runs in your mean titre calculation!

Step 4: Titration Calculations

Use these fundamental formulae to work through your titration data:

$$\text{Moles } (n) = \text{Concentration } (c, \text{mol dm}^{-3}) \times \text{Volume } (V, \text{dm}^3)$$

Remember: To convert $\text{cm}^3$ to $\text{dm}^3$, divide by 1000:
$V\text{ (dm}^3) = \frac{V\text{ (cm}^3)}{1000}$

To convert molar concentration ($\text{mol dm}^{-3}$) to mass concentration ($\text{g dm}^{-3}$):

$$\text{Concentration } (\text{g dm}^{-3}) = \text{Concentration } (\text{mol dm}^{-3}) \times M_r$$

Quick Review: Rinse burettes and pipettes with the solutions they will contain. Rinse the conical flask with deionised water only. Average only titres within $0.10\text{ cm}^3$.

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2. Calorimetry & Enthalpy of Reaction

Calorimetry is the experimental measurement of heat energy changes during a chemical reaction.

The Calorimeter Setup

In the laboratory, an expanded polystyrene cup fitted with a lid is used as a simple calorimeter.
• Why polystyrene? It is an excellent thermal insulator with a low heat capacity, minimising heat loss to the surroundings.
• Why a lid? It reduces heat loss via evaporation and convection.

Calculating Heat Transferred ($q$)

When a reaction occurs in aqueous solution, we measure the temperature change ($\Delta T = T_{\text{final}} - T_{\text{initial}}$). The heat energy exchanged ($q$) is calculated using:

$$q = mc\Delta T$$

• $q$ = heat energy transferred in Joules ($\text{J}$)
• $m$ = mass of the solution in grams ($\text{g}$), assuming an aqueous density of $1\text{ g cm}^{-3}$ (e.g., $50\text{ cm}^3$ of solution = $50\text{ g}$)
• $c$ = specific heat capacity of water ($4.18\text{ J g}^{-1}\text{ K}^{-1}$)
• $\Delta T$ = temperature change in $\text{K}$ or $^\circ\text{C}$

Calculating Molar Enthalpy Change ($\Delta H$)

To convert the energy change $q$ into standard molar enthalpy change ($\Delta H$) in $\text{kJ mol}^{-1}$:

$$\Delta H = -\frac{q}{n \times 1000}$$

• $n$ = number of moles of the limiting reactant.
• We divide by $1000$ to convert Joules ($\text{J}$) to kilojoules ($\text{kJ}$).
• The Sign Convention: If the reaction is exothermic (temperature rises), $\Delta H$ must be negative ($-$). If the reaction is endothermic (temperature falls), $\Delta H$ must be positive ($+$).

Key Takeaway: Always check your sign on $\Delta H$! A temperature increase means heat was released, requiring a minus sign in your final answer.

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3. Separation and Analytical Techniques

Thin Layer Chromatography (TLC) & Paper Chromatography

Chromatography is used to separate and identify components within a mixture based on their relative affinities for a stationary phase (the plate or paper) and a mobile phase (the liquid solvent).

Essential Practical Rules:
• Origin line: Always draw the baseline in pencil, never pen (pen ink contains dyes that will dissolve and travel up the plate).
• Spotting: Apply small, concentrated spots of sample onto the pencil line and allow them to dry.
• Solvent level: The solvent depth in the developing beaker must sit below the pencil origin line. If the solvent covers the spots, the samples will simply dissolve into the solvent pool rather than travel up the plate.
• Chamber lid: Cover the beaker with a watch glass or lid to maintain a solvent-saturated atmosphere and prevent evaporation.

Retention Factor ($R_f$):
Once the solvent front reaches near the top, remove the plate and immediately mark the solvent front with a pencil. Measure distances from the origin line to calculate $R_f$:

$$R_f = \frac{\text{Distance moved by solute}}{\text{Distance moved by solvent front}}$$

Note: $R_f$ values have no units and are always between $0$ and $1$.

Colorimetry and Spectrophotometry

Colorimetry measures the amount of light absorbed by a coloured solution to determine its concentration (based on Beer-Lambert principles).

Step-by-Step Procedure:
1. Filter selection: Select a complementary filter colour (the colour absorbed most strongly by the sample) to maximise sensitivity.
2. Setting the zero / blank: Place a cuvette filled with distilled water or solvent into the colorimeter and calibrate the instrument to read zero absorbance.
3. Standard calibration curve: Measure the absorbance of a series of standard solutions of known concentrations. Plot a graph of Absorbance vs. Concentration and draw a line of best fit through the origin.
4. Determining unknown concentration: Measure the absorbance of your unknown sample, locate this value on the y-axis of your calibration curve, and interpolate (read across and down) to find its concentration on the x-axis.

Key Takeaway: In chromatography, keep the solvent line below your spots. In colorimetry, always calibrate to zero using a solvent blank before reading standard or unknown samples.

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4. Measurement Conventions & Experimental Uncertainties

In Unit AS 1, high marks depend on correctly presenting numerical data and evaluating experimental errors.

Equipment Resolution & Decimal Places

• Analytical Balance: Typically records to $\pm 0.01\text{ g}$ (or $\pm 0.001\text{ g}$).
• Burette: Resolution allows readings to $\pm 0.05\text{ cm}^3$. All readings must end in .00 or .05 $\text{cm}^3$ (e.g., $22.40\text{ cm}^3$ or $22.45\text{ cm}^3$).

Two-Reading Absolute Uncertainties

Whenever a value is calculated by finding the difference between two readings, the uncertainty of the instrument is doubled:
• Titration titre: $\Delta V = V_{\text{final}} - V_{\text{initial}}$. Since each reading has an uncertainty of $\pm 0.05\text{ cm}^3$, the total absolute uncertainty is $2 \times 0.05 = \pm 0.10\text{ cm}^3$.
• Mass by difference: Two balance readings mean the absolute uncertainty is $2 \times \text{balance uncertainty}$.

Calculating Percentage Uncertainty

To evaluate the impact of an instrument's limitation on your measurement, calculate the percentage uncertainty:

$$\% \text{ Uncertainty} = \frac{\text{Absolute uncertainty}}{\text{Quantity measured}} \times 100$$

Example: If you measure a titre of $25.00\text{ cm}^3$ using a burette ($\text{total uncertainty} = \pm 0.10\text{ cm}^3$):
$\% \text{ Uncertainty} = \frac{0.10}{25.00} \times 100 = 0.40\%$

Tip to minimise uncertainty: Increase the quantity measured (e.g., use a larger titre volume or a greater mass of solid) to reduce the percentage uncertainty!

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5. Summary of Common Mistakes to Avoid

Keep these examiner-reported pitfalls in mind when writing up your AS 1 portfolio:

• Averaging Discordant Titres: Never include the rough titre or any titre that differs by more than $0.10\text{ cm}^3$ from the others when calculating your mean.
• Meniscus Errors: Always read volumes from the bottom of the meniscus at eye level.
• Burette Air Bubbles: Forgetting to fill the space below the burette tap before starting a titration leads to a falsely large titre reading.
• Wrong Glassware Rinsing: Never rinse a volumetric pipette or burette with water right before filling (it dilutes the solution). Never rinse the conical flask with the chemical solution (it adds extra unmeasured moles).
• Submerged Chromatography Spots: Placing too much solvent in the developing tank so that the liquid level is above the pencil baseline causes samples to wash off into the solvent.
• Single Uncertainty on Differences: Forgetting to double the instrument uncertainty when calculating differences (titre volume or mass by difference).

Quick check

Can you answer these now?

Open each question to check the key ideas from this chapter.

What is the method of weighing by difference and why is it used?

Weighing by difference involves recording the initial mass of the container with solid, transferring the solid, and re-weighing the empty container to calculate $m_{\text{transferred}} = m_{\text{initial}} - m_{\text{residue}}$. It accounts for any solid residue left behind.

How should a volumetric flask be filled and inverted when making a standard solution?

Add deionised water until the bottom of the meniscus touches the calibration line at eye level, stopper the flask, and invert it 10 to 15 times to ensure uniform concentration.

How should the burette, volumetric pipette, and conical flask be rinsed before a titration?

Rinse the burette with deionised water then the titrant; rinse the pipette with deionised water then the analyte; rinse the conical flask with deionised water only.

What are concordant titres and which values should be included in the mean titre calculation?

Concordant titres are titration volumes that agree within $0.10\text{ cm}^3$ of each other. Only concordant titres are used to calculate the mean titre.

How do you convert molar concentration ($\text{mol dm}^{-3}$) into mass concentration ($\text{g dm}^{-3}$)?

Multiply the molar concentration by the relative formula mass ($M_r$): $\text{Concentration } (\text{g dm}^{-3}) = \text{Concentration } (\text{mol dm}^{-3}) \times M_r$.

How is heat energy transferred ($q$) calculated in an aqueous calorimetry experiment?

Using the equation $q = mc\Delta T$, where $m$ is the mass of solution, $c = 4.18\text{ J g}^{-1}\text{ K}^{-1}$, and $\Delta T$ is the temperature change.

How is standard molar enthalpy change ($\Delta H$) calculated from heat transferred ($q$)?

$\Delta H = -\frac{q}{n \times 1000}$, where $n$ is the moles of the limiting reactant, giving $\Delta H$ in $\text{kJ mol}^{-1}$ with a negative sign for exothermic reactions and positive for endothermic.

How is the retention factor ($R_f$) calculated in chromatography?

$R_f = \frac{\text{Distance moved by solute}}{\text{Distance moved by solvent front}}$. It is a dimensionless ratio between 0 and 1.

How is an unknown concentration determined using colorimetry?

Calibrate to zero absorbance using a solvent blank, measure standard solutions to plot a calibration curve of absorbance against concentration, and interpolate the unknown sample's concentration.

Why is the absolute uncertainty doubled for measurements taken by difference?

Two readings are taken (initial and final), so the total absolute uncertainty is the sum of both individual reading uncertainties (e.g., $2 \times 0.05 = \pm 0.10\text{ cm}^3$ for a burette).

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