Suggest advantages and disadvantages of experimental methods and apparatus

Cambridge IGCSE Chemistry (0620) – Experimental Methods & Apparatus

1. Scope Overview – What Is Covered in These Notes?

Link to Assessment Objectives

• AO1 – Knowledge & Understanding: definitions of variables, particle‑theory concepts, mole‑concept formulas, redox notation.
• AO2 – Application: selecting appropriate methods, performing calculations (e.g. $n = m/M$, $PV=nRT$), using correct indicators.
• AO3 – Analysis & Evaluation: designing experiments, recognising systematic vs random errors, suggesting realistic improvements.

2. Designing a Complete Experiment (AO3)

1. Write a clear hypothesis / aim (e.g. “The volume of hydrogen produced from the reaction of magnesium with hydrochloric acid increases with temperature”).
2. Identify variables:
   • Independent variable – what you change (temperature).
   • Dependent variable – what you measure (gas volume).
   • At least three controlled variables – e.g. mass of Mg, concentration of HCl, volume of water bath.
3. Choose a method that best answers the question (e.g. gas‑collection with a eudiometer) and list all required apparatus.
4. Explain how each piece of apparatus reduces systematic error and promotes repeatability (e.g. calibrated burette, digital thermometer).
5. Plan the record‑keeping:
   • Data tables for raw observations.
   • Calculation sheets (e.g. $PV=nRT$).
   • Error‑analysis template (see Section 6).
6. Carry out **≥ 3 trials**, record all data, and calculate the average result.
7. Analyse the data, discuss sources of error (systematic vs random) and suggest realistic improvements.

AO3 Checklist (tick when complete)

• ✓ Hypothesis / aim written.
• ✓ Independent, dependent and controlled variables listed.
• ✓ Method written in logical order, including safety precautions.
• ✓ All required apparatus named (including devices for time, temperature and pressure).
• ✓ Calibration / checking of equipment described.
• ✓ Data‑tables prepared (raw data, calculations, error table).
• ✓ Evaluation of systematic and random errors included.

3. Measuring Time, Temperature & Pressure – Advantages & Disadvantages

4. Common Experimental Methods – Advantages & Disadvantages

5. Typical Apparatus – Advantages & Disadvantages

6. Key Theory Boxes (Core Content Required Elsewhere in the Syllabus)

6.1 Kinetic Particle Theory & Diffusion (Topic 1)

• Particles are in constant motion; kinetic energy increases with temperature.
• State‑change diagrams illustrate how heating expands solids → liquids → gases.
• Diffusion: net movement of particles from high to low concentration; rate ↑ with temperature, ↓ with particle mass.
• Equation for average kinetic energy: $E_{\text{kin}} = \frac{3}{2}RT$ (useful for linking temperature to particle motion).

6.2 Mole Concept, Molar Mass & Gas Laws (Topics 3 & 4)

• Number of moles: $n = \dfrac{m}{M}$ where $m$ = mass (g) and $M$ = molar mass (g mol⁻¹).
• Avogadro’s constant: $N_{\text A}=6.02\times10^{23}$ particles mol⁻¹.
• Ideal‑gas equation: $PV = nRT$ (R = 8.314 J mol⁻¹ K⁻¹ or 0.0821 L atm mol⁻¹ K⁻¹).
• Standard temperature and pressure (STP): 0 °C, 1 atm → 22.4 L mol⁻¹.
• Example calculation: 0.50 g NaCl (M = 58.44 g mol⁻¹) contains $n = 0.50/58.44 = 8.55\times10^{-3}$ mol.

6.3 Stoichiometry, Limiting Reactant & Percentage Yield (Topic 3)

• Use balanced equations to relate moles of reactants and products.
• Limiting reactant: the reactant that produces the smallest amount of product.
• Percentage yield: $\displaystyle \%\,\text{yield}= \frac{\text{actual yield}}{\text{theoretical yield}}\times100$.
• Example: 2 g H₂ + 16 g O₂ → 2 g H₂O (theoretical). If 1.5 g H₂O is obtained, % yield = $1.5/2\times100 = 75\%$.

6.4 Basic Redox Notation (Topic 4)

• Oxidation number rules (e.g., O = –2, H = +1, alkali metals = +1).
• Redox reaction: separate into half‑equations.

          Example: $\ce{Zn + Cu^{2+} -> Zn^{2+} + Cu}$
          Oxidation:  $\ce{Zn -> Zn^{2+} + 2e^-}$
          Reduction: $\ce{Cu^{2+} + 2e^- -> Cu}$
          

• Electron balance ensures total electrons lost = total electrons gained.

7. Acid–Base Titrations (Syllabus 12.2)

• Typical apparatus: burette, stand & clamp, volumetric pipette, conical flask, white tile, indicator (phenolphthalein, methyl orange, bromothymol blue), pH‑meter (optional), distilled‑water rinse bottle.

Advantages of the Apparatus

• Burette – delivers titrant with high precision (±0.05 mL).
• Volumetric pipette – ensures a known, exact volume of analyte.
• White tile – enhances visibility of subtle colour changes.
• Indicator – gives a clear visual end‑point for many acid‑base reactions.
• pH‑meter (if used) – provides an instrumental end‑point, reducing subjectivity.

Disadvantages / Common Sources of Error

• Wrong indicator (pH range does not match equivalence point).
• Parallax error when reading the burette meniscus.
• Air bubbles trapped in the burette tip.
• Insufficient mixing of titrant and analyte (swirl gently after each addition).
• Temperature fluctuations alter solution volume and pH.

End‑Point Identification

• Visual*: colour change of the chosen indicator (e.g., phenolphthalein turns faint pink at pH≈8.2).
• Instrumental*: pH‑meter reading stabilises at the expected equivalence pH (e.g., pH ≈ 7 for a strong acid–strong base).
• Record the volume of titrant at the **first permanent** colour change (or pH value) – this is the experimental end‑point.