Describe how to test for the purity of water using melting point and boiling point

Chemistry of the Environment – Water: Testing Purity

1. Objective

To describe how to test the purity of water using qualitative colour‑change tests and quantitative melting‑point (freezing‑point) and boiling‑point determinations, and to relate these techniques to the wider Cambridge IGCSE 0620 syllabus.

2. Syllabus Context (IGCSE 0620)

• 10.1 Water – chemical tests & purity
• 10.1 Water – sources & contaminants
• 10.1 Water – treatment processes
• 10.2 Fertilisers – colligative‑property applications
• 10.3 Air quality & climate – environmental monitoring
• 12 Experimental techniques & chemical analysis
• Core units 1‑9 (states of matter, atoms, stoichiometry, electrochemistry, energetics, reactions, acids & bases, periodic table, metals) – required background knowledge for understanding the tests.

3. Core Chemistry Refresher (required for the water‑purity topic)

3.1 States of Matter

• Solids have fixed shape and volume; particles vibrate about fixed positions.
• Liquids have fixed volume but take the shape of their container; particles slide past one another.
• Gases have neither fixed shape nor volume; particles move freely and diffuse rapidly.
• Phase changes (melting, freezing, boiling, condensation) involve energy transfer and are reversible.

3.2 Atoms, Ions & Bonding

• Atoms consist of a nucleus (protons + neutrons) surrounded by electrons.
• Isotopes differ in neutron number; ions form when atoms lose or gain electrons.
• Bond types:
  • Ionic – transfer of electrons (e.g., NaCl).
  • Simple covalent – sharing of electrons between non‑metals (e.g., H₂O).
  • Giant covalent – network of covalent bonds (e.g., SiO₂, diamond).

3.3 Stoichiometry (Mole Concept)

• Relative formula mass (Mᵣ) → molar mass (M, g mol⁻¹).
• 1 mol contains 6.02 × 10²³ particles (Avogadro’s constant).
• Basic calculations:
  1. Convert mass ↔ moles using \(n = \frac{m}{M}\).
  2. Use balanced equations to relate moles of reactants and products.

3.4 Electrochemistry

• Electrolysis – non‑spontaneous redox driven by electricity.
  • General set‑up: two electrodes, electrolyte, power source.
  • Typical products:
    • Molten NaCl → Na (cathode) + Cl₂ (anode).
    • Molten PbBr₂ → Pb (cathode) + Br₂ (anode).
    • Dilute H₂SO₄ → H₂ (cathode) + O₂ (anode).
• Fuel cells – spontaneous redox producing electricity; e.g. H₂ + ½ O₂ → H₂O.

3.5 Chemical Energetics

• Enthalpy change (ΔH): exothermic (ΔH < 0) releases heat; endothermic (ΔH > 0) absorbs heat.
• Activation energy (Eₐ) – minimum energy required for a reaction to proceed.
• Bond‑energy calculations: ΔH ≈ Σ(broken bonds) – Σ(formed bonds).

3.6 Chemical Reactions

• Rate of reaction – increased by higher concentration, temperature, surface area, or a catalyst.
• Reversible reactions – written with a double arrow ↔; equilibrium is reached when forward and reverse rates are equal.
• Redox reactions – identified by changes in oxidation number; electrons are transferred.

3.7 Acids, Bases & Salts

• Acids release H⁺ in water; bases release OH⁻ or accept H⁺.
• pH scale: 0 – 14; strong acids/bases dissociate completely, weak acids/bases only partially.
• Neutralisation: acid + base → salt + water.
• Typical titration for water hardness (Ca²⁺/Mg²⁺) uses EDTA or a strong acid‑base indicator.

3.8 Periodic Table Trends (Core 8)

Group	General Trend	Typical Example
1 (alkali metals)	Increasing atomic radius down the group; very reactive, form +1 ions.	Na, K
2 (alkaline earths)	Form +2 ions; reactivity less than group 1.	Ca, Mg
13‑16 (p‑block)	Metal‑non‑metal character changes across the period; covalent bonding dominates.	Si, P, S, Cl
17 (halogens)	High electronegativity; form –1 ions.	Cl, Br
18 (noble gases)	Very low reactivity; complete valence shells.	Ar, Ne

3.9 Metals (Core 9)

• Typical properties: metallic luster, good conductors of heat/electricity, malleable, ductile, high melting points (except alkali metals).
• Reactivity series (most → least reactive): K > Na > Ca > Mg > Al > Zn > Fe > Sn > Cu > Ag > Au.
  • More reactive metals displace less reactive metals from their salts.
  • Corrosion (rusting) of iron: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ → Fe₂O₃·nH₂O.
• Extraction example: iron from Fe₂O₃ in a blast furnace using coke (C) as a reducing agent.

4. Water – Purity Testing

4.1 Why Test Purity?

• Pure water freezes at 0 °C and boils at 100 °C (1 atm).
• Any dissolved non‑volatile solute lowers the freezing point (freezing‑point depression) and raises the boiling point (boiling‑point elevation) – both are colligative properties.
• Magnitude of the change is proportional to the total molality of solutes: \[ \Delta T_f = iK_f m \qquad \Delta T_b = iK_b m \] where K_f = 1.86 °C·kg mol⁻¹, K_b = 0.512 °C·kg mol⁻¹ for water and i is the van t Hoff factor.

4.2 Typical Contaminants (Cambridge list)

Group	Examples	Effect on Colligative Properties?
Inorganic salts	NaCl, CaSO₄, Mg(NO₃)₂	↓ T_f, ↑ T_b
Organic matter	Sugars, ethanol, pesticides	↓ T_f, ↑ T_b
Dissolved gases	O₂, CO₂	Negligible (very low concentration)
Heavy‑metal ions	Fe³⁺, Pb²⁺ (as soluble salts)	↓ T_f, ↑ T_b
Micro‑organisms	Bacteria, algae	No direct effect on T_f/T_b (requires microbiological tests)
Micro‑plastics	Polyethylene fragments	Usually negligible for colligative properties

4.3 Water‑Treatment Processes (syllabus requirement)

• Sedimentation – removes suspended solids; does not change T_f or T_b.
• Filtration – mechanical removal of fine particles; similarly no effect on colligative properties.
• Activated‑carbon adsorption – removes many organic solutes; may reduce ΔT_f/ΔT_b if organics are the main impurity.
• Chlorination – adds Cl₂/HOCl; the resulting chloride ions act as a very dilute non‑volatile solute, causing a <0.1 °C> rise in boiling point – essentially negligible for most school‑level measurements.

4.4 Qualitative Colour‑Change Tests (rapid indication of water & some ions)

Dry reagent	Colour change on addition of water	Interpretation
Anhydrous cobalt(II) chloride, CoCl₂·6H₂O	Blue → pink (monohydrate) → purple (excess water)	Detects presence of water (hydrolysis).
Anhydrous copper(II) sulfate, CuSO₄	White → blue (pentahydrate)	Detects water; persistent blue after drying also indicates sulphate ions.

4.5 Quantitative Purity Tests – Melting‑Point (Freezing‑Point) & Boiling‑Point

Apparatus

1. Digital temperature probe (range –20 °C to 110 °C, accuracy ±0.1 °C) or calibrated thermometer.
2. Ice‑salt bath (or ice‑water mixture for a 0 °C reference).
3. Thin‑walled glass tube (≈10 mL) or sealed ampoule.
4. Boiling set‑up: clean beaker (≈100 mL), Bunsen burner or hot plate, stand & clamp.
5. Polystyrene cup or insulated container for the freezing‑point test.
6. Distilled water (control) and the water sample to be examined.
7. Stopwatch, lab notebook, ruler (to note thermometer bulb height).

Experimental Design (Core 12)

• Independent variable: type of water (distilled vs. sample).
• Dependent variables: observed freezing point and boiling point.
• Controlled variables: atmospheric pressure (~1 atm), volume of water, heating/cooling rate, thermometer calibration, composition of ice‑salt bath.
• Include a control (distilled water) in each series of measurements.
• Perform at least three repetitions per sample; report the mean ± standard deviation.
• Record any anomalies (super‑cooling, premature bubbling, thermometer lag).

Procedure

4.5.1 Freezing‑Point (Melting‑Point of Ice) Determination

1. Prepare an ice‑salt bath in a polystyrene cup; stir until a stable temperature (e.g., –15 °C) is reached. Record the bath temperature.
2. Fill a thin‑walled glass tube with ~2 mL of the water sample, seal, and tap gently to remove trapped air.
3. Immerse the tube in the ice‑salt bath, ensuring the thermometer bulb lies within the liquid but does not touch the tube wall.
4. Allow the system to equilibrate. Observe the temperature at which the first ice crystals appear on cooling (or the first liquid droplets appear on warming). This temperature is the **freezing point**.
5. Repeat the same steps with distilled water as a reference.
6. Carry out three trials for each water type and calculate the average.

4.5.2 Boiling‑Point Determination

1. Measure 50 mL of the water sample into a clean beaker.
2. Place the beaker on a pre‑heated hot plate (or Bunsen burner). Insert the thermometer so the bulb is ~1 cm below the surface.
3. Heat gently, then increase to a steady boil. Record the temperature when a **continuous stream of bubbles** forms and the temperature stops rising – this is the boiling point.
4. Maintain the boil for at least 30 s and note the stable temperature.
5. Repeat with distilled water under identical conditions.
6. Perform three repetitions per sample and compute the mean.

Data‑Recording Table

Test	Pure (Distilled) Water	Sample Water (mean ± SD)	Interpretation
Freezing point (°C)	0.0
Boiling point (°C) at 1 atm	100.0

Calculations (optional – quantitative analysis)

Use the average temperature deviation to estimate the total concentration of dissolved non‑volatile solutes.

Freezing‑point depression

\[ \Delta T_f = T_{\text{pure}} - T_{\text{sample}} = i\,K_f\,m \]

Boiling‑point elevation

\[ \Delta T_b = T_{\text{sample}} - T_{\text{pure}} = i\,K_b\,m \]

Solve for molality \(m\) (mol kg⁻¹). Set \(i = 1\) for a non‑electrolyte; for a 1:1 electrolyte use \(i = 2\); adjust for other electrolytes accordingly.

Interpretation of Results

• Freezing point < 0 °C → dissolved solutes present.
• Boiling point > 100 °C → same solutes raise the boiling temperature.
• The magnitude of the deviation gives an estimate of total solute concentration, but not the identity of the impurity.
• If both temperatures agree with the pure‑water values (within experimental error), the water can be considered chemically pure (microbial contamination would need separate tests).
• Colour‑change tests that show no colour shift support the conclusion of purity.

Limitations of the Temperature Methods

1. Only detects **non‑volatile** solutes; gases and very low‑concentration contaminants may be missed.
2. Cannot distinguish between different solutes – a mixture of salts and organics gives the same overall ΔT.
3. Super‑cooling or premature nucleation can give falsely low freezing‑point readings.
4. Accurate boiling‑point measurement requires stable atmospheric pressure; altitude corrections are needed for field work.

5. Complementary Tests

• Acid–base titration – determines hardness (Ca²⁺, Mg²⁺) by complexometric titration with EDTA or by neutralising carbonate hardness with a strong acid.
• Conductivity measurement – a quick indicator of total ionic content; higher conductivity correlates with larger ΔT values.
• Spectroscopic or chromatographic methods (advanced) – identify specific organic pollutants, but are beyond the IGCSE scope.

6. Safety Precautions

• Wear safety goggles, lab coat and heat‑resistant gloves.
• Handle hot equipment (boiling water, Bunsen burner) with tongs or heat‑proof gloves.
• When preparing the ice‑salt bath, avoid skin contact with the extremely cold mixture.
• Work in a well‑ventilated area if using a Bunsen burner.
• Dispose of coloured reagent waste according to school safety guidelines.

7. Suggested Diagram

8. Summary

Testing water purity in the IGCSE curriculum combines rapid qualitative colour‑change tests (anhydrous CoCl₂ and CuSO₄) with quantitative melting‑point (freezing‑point) and boiling‑point determinations. Pure water freezes at 0 °C and boils at 100 °C (1 atm). Any deviation indicates dissolved non‑volatile solutes, and the size of the deviation can be used to estimate total solute concentration via the colligative‑property equations. The method fits within a broader environmental context that includes typical contaminants, water‑treatment processes, related fertilizer analysis, and links to air‑quality monitoring. A solid understanding of the core chemistry topics (states of matter, atoms, stoichiometry, electrochemistry, energetics, reactions, acids & bases, periodic trends, and metals) underpins the interpretation of the results.

9. Assessment Questions

1. Explain why a non‑volatile solute lowers the freezing point of water, including a brief discussion of vapour‑pressure lowering.
2. A water sample freezes at -2.5 °C. Assuming the solute is a non‑electrolyte, calculate its molality using K_f = 1.86 °C·kg mol⁻¹.
3. During a boiling‑point test the temperature stabilises at 101.2 °C. Estimate the molality of the solute if it is a 1:1 electrolyte (i = 2) using K_b = 0.512 °C·kg mol⁻¹.
4. List two limitations of using melting‑point and boiling‑point determinations for assessing water purity.
5. Describe how an acid–base titration could be used to complement the temperature tests when checking for water hardness.
6. Briefly discuss why chlorination does not significantly affect the boiling point of treated water.