Chemistry Grade 10 Practical Based Assessment (PBA) Notes, Worksheets, Concept-Based Questions | Federal Board (FBISE)

⚗ SSC-II · PBA 2025–2026

Chemistry Grade 10 PBA Notes
Major & Minor Experiments · FBISE

Complete Practical Based Assessment notes for Chemistry Grade 10 — with step-by-step procedures, observation tables, precautions, and concept-based viva questions for every major and minor experiment.

9Experiments

2Parts

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PART 1

Major Experiments

5 experiments

1 Fractional Distillation of Water & Alcohol Separating a miscible mixture by boiling point difference ▶

Introduction & Theory

• Core Principle: When the mixture is heated, the component with the lower boiling point evaporates first. These vapors are cooled in a condenser to turn back into a liquid (distillate).
• Target Mixture: Alcohol (Ethanol) and Water
  • Boiling Point of Alcohol: 78°C
  • Boiling Point of Water: 100°C

Procedure

1. Setup: Assemble the apparatus, ensuring the thermometer bulb is positioned precisely in front of the side arm of the distillation flask to measure vapor temperature.
2. Preparation: Pour 100 cm³ of the mixture into the round-bottom flask and add a few boiling chips.
3. Sealing: Seal all joints with Plaster of Paris or grease to prevent volatile and flammable alcohol vapors from escaping.
4. Heating: Gradually heat the flask using a Bunsen burner.
5. Condensation: Circulate cold water through the condenser (entering at the lower inlet, exiting at the upper outlet).
6. Separation: At 78°C, alcohol vapors pass through the condenser and are collected. Once alcohol is fully distilled, temperature rises toward 100°C — replace the receiving flask to collect pure water.

Observations & Calculations

• Temperature Stability: During alcohol collection, the thermometer remains steady at 78°C. It rises again once all alcohol has evaporated.
• Yield Calculation: In a mixture containing 30% alcohol, expect ~30% of the initial volume as distillate (with efficient distillation).
• State Change: In the condenser, a physical change occurs — vapor turns back to liquid due to cooling.

Precautions

• Bumping Prevention: Always add boiling chips to prevent violent boiling.
• Fire Safety: Since alcohol is flammable, remove the heat source before the flask runs dry.
• Volume Control: The distillation flask should only be partially filled.
• Cooling Efficiency: Cold water must enter the lower inlet for maximum cooling.

2 Paper Chromatography Separating ink mixture into individual dye components ▶

Introduction

Paper Chromatography is a laboratory technique used to separate a mixture of chemical substances into its individual components based on their distribution between a stationary phase and a mobile phase.

Theory

• Stationary Phase: The chromatography paper.
• Mobile Phase: The solvent that moves through the paper.
• Capillary Action: Allows the solvent to travel up the paper strip.
• Principle: Different substances have different affinities for the stationary and mobile phases, causing them to move at different rates.

Apparatus & Chemicals

• Equipment: Filter/chromatography paper, ink/markers, pencil, ruler, containers, clips, developing chamber with lid.
• Chemicals: Solvent mixture (Water + Ethanol in 7:3 ratio) and the ink mixture.

Procedure

1. Prepare the Strip: Cut chromatography paper; draw a pencil baseline near the bottom.
2. Apply Ink: Place small dots on the baseline; ensure they dry.
3. Prepare Chamber: Pour solvent into chamber — level must be below the baseline.
4. Develop: Hang the paper vertically in the chamber (not touching walls).
5. Seal: Cover with a lid for a saturated atmosphere.
6. Separation: Allow solvent to travel ~¾ of the way up. Remove and mark the "solvent front" with a pencil.

Observations

The Retention Factor (R_f) identifies components:

R_f = Distance travelled by solute ÷ Distance travelled by solvent front

Component (Color)	Distance (cm)	Solvent Front (cm)	Rf
Blue	3	17	0.18
Red	6	17	0.35
Green	9	17	0.53

Precautions

• Handle paper by the edges — avoid contamination from skin oils.
• Use a pencil for the baseline (pen ink would dissolve and interfere).
• Ink spots must be above the solvent level.
• Paper must be vertical and not touching chamber walls.
• Ink dots should be small to prevent "tailing" or overlapping bands.

Concept-Based Questions

Why must the paper be removed before the solvent reaches the top?

If the solvent reaches the edge, the solvent front cannot be measured accurately, making R_f values invalid.

What is the purpose of the lid?

It maintains a solvent-vapor-rich environment, preventing evaporation and ensuring smooth, uniform flow.

Why does the ink separate into different colors?

Different components have different polarities and affinities. Those more soluble in the solvent move faster; those with higher affinity for the paper move slower.

What can be concluded if two ink samples produce identical chromatograms?

They likely contain the same dyes and may come from the same source.

Why is water-ethanol (7:3) preferred?

This balanced polarity effectively dissolves and separates a wide range of ink dyes.

3 Separation of Pb²⁺ & Cd²⁺ Ions by Paper Chromatography Isolating lead and cadmium ions using R_f analysis ▶

Objective

To separate a mixture of Lead (Pb²⁺) and Cadmium (Cd²⁺) ions using paper chromatography and determine their individual R_f values.

Theory

• Definition: Paper chromatography separates and analyzes mixture components based on distribution between a stationary phase (filter paper) and a mobile phase (solvent).
• Principle: Different substances move at different speeds due to varying solubility in the solvent and affinity for the paper.

Apparatus & Chemicals

• Materials: Chromatography paper, pencil, ruler, glass/plastic containers, clips, developing chamber with lid.
• Solvents: Water, isopropyl alcohol, or acetone. Specific: n-butanol + 3 M HNO₃.
• Sample: Mixture of Lead and Cadmium chlorides or nitrates.
• Locating Agent: Concentrated H₂S gas or dithizone solution.

Procedure

1. Prepare Strip: Cut chromatography paper; draw pencil baseline near bottom.
2. Spot Sample: Apply small volume of mixture with a capillary tube; dry.
3. Prepare Chamber: Pour solvent mixture — level below the baseline.
4. Development: Hang paper vertically; close lid to saturate.
5. Separation: Allow solvent to travel — it carries Pb²⁺ and Cd²⁺ at different rates.
6. Mark Front: Remove paper before solvent reaches top; mark the solvent front.
7. Visualization: Spray dried paper with H₂S to make colorless spots visible.

Observations

R_f = Distance travelled by solute ÷ Distance travelled by solvent

Ion	Color of Spot	Distance from Baseline	Rf
Lead (Pb²⁺)	Black	3.5 cm	0.2
Cadmium (Cd²⁺)	Yellow	5.3 cm	0.3

* Example distance of solvent front = 17.6 cm.

Precautions

• Applied sample spot must be very small and completely dry.
• Baseline must never be submerged in the solvent.
• Do not let paper touch the chamber walls — ensures uniform flow.
• Mark the solvent front immediately before it evaporates.

Concept-Based Questions

Why is a locating agent necessary for Pb²⁺ and Cd²⁺?

These ions are colorless in solution. H₂S reacts with them to form colored sulphides (PbS and CdS) that are visible.

Why use a pencil instead of a pen?

Pencil graphite is inert and insoluble. Pen ink contains dyes that would dissolve and travel with the solvent, ruining the chromatogram.

Why does Cd²⁺ travel further than Pb²⁺?

Cd²⁺ has a higher affinity for the mobile phase (solvent); Pb²⁺ interacts more strongly with the stationary phase (paper).

Why must the chamber be saturated with solvent vapor?

Saturation prevents solvent from evaporating off the paper, ensuring a constant, uniform flow for accurate R_f calculations.

4 Titration Experiment Determining molarity of NaOH using standard HCl ▶

Objective

To determine the exact molarity of a Sodium Hydroxide (NaOH) solution through volumetric analysis (titration).

Theory

• Standard Solution: Hydrochloric acid (HCl) is used as the known standard.
• Chemical Equation: NaOH + HCl → NaCl + H₂O
• Molar Ratio: HCl : NaOH = 1:1
• Indicator: Phenolphthalein
• End Point: Solution changes from light pink to colorless.

Apparatus & Chemicals

• Equipment: Burette, pipette, funnel, conical flask, beaker, dropper, burette stand.
• Chemicals: Standard HCl solution, NaOH solution, phenolphthalein indicator.

Procedure

1. Preparation: Rinse pipette with NaOH and burette with HCl solution.
2. Sampling: Pipette 10 cm³ of NaOH into a conical flask; add 1–2 drops of phenolphthalein (solution turns pink).
3. Setting Burette: Fill with HCl; remove funnel; ensure no air bubbles.
4. Rough Titration: Add HCl while swirling until pink just disappears.
5. Precise Titration: Repeat, adding acid drop-wise near the end point, until three concordant readings (within 0.1 cm³) are obtained.

Calculation

M₁V₁ / n₁ = M₂V₂ / n₂

• M₁: Molarity of HCl
• V₁: Mean volume of acid used
• n₁: Moles of acid (1)
• M₂: Molarity of NaOH (unknown)
• V₂: Volume of NaOH used (10 cm³)
• n₂: Moles of base (1)

Precautions

• Rinsing: Rinse burette and pipette with their specific solutions to avoid dilution.
• Funnel Removal: Remove glass funnel from burette before taking readings.
• Visual Aid: Use a white tile under the flask to see color change clearly.
• Reading the Burette: Keep eye at the level of the lower meniscus to avoid parallax error.
• Nozzle Check: Ensure nozzle is fully filled and free of air bubbles.

Concept-Based Questions

Why remove the funnel?

To prevent extra acid from dropping into the burette, which would cause inaccurate volume readings.

Why is phenolphthalein used?

It's suitable for strong acid–strong base titrations because it provides a sharp color change at the end point.

What is the purpose of standardization?

To find the exact concentration (molarity) of a solution by reacting it against a solution of known concentration.

Which meniscus should be read?

For clear solutions like HCl, the lower meniscus must be read at eye level.

5 Detection and Confirmation of Gases Identifying NH₃, CO₂ & Cl₂ by chemical reactions ▶

Objective

To identify three specific gases — Ammonia (NH₃), Carbon Dioxide (CO₂), and Chlorine (Cl₂) — by observing distinct chemical reactions or physical properties when exposed to specific reagents.

Theory

• Ammonia (NH₃): A basic gas. Reacts with water on damp litmus paper to form ammonium hydroxide (NH₄OH), which changes red litmus to blue.
• Carbon Dioxide (CO₂): Reacts with limewater to form calcium carbonate (CaCO₃), an insoluble white precipitate that makes the solution milky.
• Chlorine (Cl₂): An acidic gas and strong oxidizing agent. Reacts with water to form HCl and HClO, which first turns blue litmus red, then bleaches it white.

Apparatus & Chemicals

• Equipment: Test tubes, gas jars, delivery tubes, Bunsen burner.
• Indicators: Red litmus paper, blue litmus paper, starch-iodide paper.
• Reagents: Limewater (Ca(OH)₂ solution), water (for damping litmus).

Procedure

1. For Ammonia: Hold damp red litmus paper near the mouth of the container releasing the gas.
2. For CO₂: Pass the gas through a delivery tube into a test tube containing limewater.
3. For Chlorine: Hold damp blue litmus paper or starch-iodide paper near the source.

Observations

Gas	Observation	Chemical Equation
Ammonia (NH₃)	Damp red litmus → Blue	NH₃ + H₂O → NH₄⁺ + OH⁻
Carbon Dioxide (CO₂)	Limewater → Milky/Cloudy	Ca(OH)₂ + CO₂ → CaCO₃ + H₂O
Chlorine (Cl₂)	Blue litmus → Red → White (bleached). Starch-iodide → Blue	Cl₂ + H₂O → HCl + HClO

Precautions

• Follow safety measures for handling toxic/irritating gases.
• Ensure proper ventilation in the laboratory.
• Be aware of chemical hazards before handling.
• Use damp litmus paper to allow the gas to dissolve and react chemically.

Concept-Based Questions

Why must litmus paper be damp?

Water is required for the gas to dissolve and react (e.g., forming NH₄OH or HClO) to produce the ions that cause a color change.

How do you distinguish NH₃ from other gases?

Only NH₃ is basic enough to turn red litmus paper blue.

What causes milkiness in the limewater test?

The formation of insoluble calcium carbonate (CaCO₃) solid particles.

Why does chlorine bleach litmus paper?

Chlorine is a strong oxidizing agent. It reacts with water to form hypochlorous acid (HClO), which destroys the litmus dye.

Daily-life applications of gas detection?

Air quality monitoring, industrial safety, detecting natural gas leaks, and medical diagnosis through respiratory function assessment.

PART 2

Minor Experiments

4 experiments

1 Separation by Sublimation Separating naphthalene from a sand & salt mixture ▶

Objective

To separate naphthalene from a mixture of sand and salt using the method of sublimation.

Theory

• Definition: Sublimation is when a solid changes directly to vapor upon heating, and back to solid upon cooling, without becoming liquid.
• Chemical Behavior: Effective for substances like naphthalene, ammonium chloride, iodine, and camphor — all with sublimable properties.

Apparatus & Chemicals

• Apparatus: China dish, funnel, cotton, sand bath, tripod stand, burner, wire gauze, filter paper.
• Chemicals: Sand and Naphthalene.

Procedure

1. Place 5–8 g of impure sample in a china dish.
2. Set the dish on a sand bath over a tripod stand.
3. Moisten a cone-shaped filter paper with water and fix it inside the funnel (for a cooler condensation surface).
4. Invert the funnel over the china dish and plug the stem with cotton.
5. Heat very gently. Naphthalene sublimates and deposits as crystals on the filter paper.
6. Stop heating once all naphthalene has sublimed. Allow to cool completely before removing the funnel.

Observations

• Completion Signal: Process is complete when no more vapors appear and only sand remains.
• Result: Pure naphthalene is recovered from the filter paper; non-volatile sand remains behind.

Precautions

• Gentle Heating: Heat gently to avoid melting naphthalene instead of subliming it.
• Safety & Yield: Ensure funnel is cooled before removal to prevent product loss.

Concept-Based Questions

Why is naphthalene separated this way?

Because naphthalene is volatile (sublimes) while sand is non-volatile.

What is the purpose of the sand bath?

It distributes heat uniformly and prevents overheating or melting.

Why moisten the filter paper?

It provides a cooler surface to help vapors condense into solid crystals more effectively.

What happens if the funnel touches the sample?

Heat transfers directly to the funnel, reducing the temperature difference and inhibiting condensation.

Why use a cotton plug?

To prevent vapors from escaping, ensuring maximum condensation and yield.

2 Identification of Metal Ions by Flame Test Qualitative cation analysis using flame colors ▶

Objective

To identify specific metal cations in a given sample through qualitative analysis by observing the unique color characteristics they impart to a Bunsen burner flame.

Theory

• Principle: When a metal ion is heated, its electrons absorb energy and jump to higher energy levels (excitation). As they return, they release light of a specific wavelength.
• Characteristic Colors: The wavelength of light is unique to each element — a "chemical fingerprint."
• Metal Chlorides: Concentrated HCl is used because metal chlorides are more volatile, making them easier to evaporate into the flame.

Apparatus & Chemicals

• Apparatus: Platinum or nichrome wire, watch glass, Bunsen burner, matches.
• Chemicals: Concentrated HCl, salts of Na⁺, Ca²⁺, Cu²⁺, K⁺, Ba²⁺, Sr²⁺.

Procedure

1. Clean the Wire: Dip in concentrated HCl; heat in the oxidizing (blue) flame until it imparts no color.
2. Prepare Sample: Place a small amount of salt on a watch glass; add a few drops of concentrated HCl to make a thick paste.
3. Load Sample: Dip the clean loop into the paste so a small amount adheres.
4. Flame Observation: Hold the loop in the hottest part of the flame (upper edge of the blue cone).
5. Record and Repeat: Note the flame color; re-clean before testing a different salt.

Observations

Metal Ion	Flame Color
Sodium (Na⁺)	Persistent golden yellow
Calcium (Ca²⁺)	Brick red
Copper (Cu²⁺)	Bluish green
Potassium (K⁺)	Violet / Lilac
Barium (Ba²⁺)	Apple green / Grassy green
Strontium (Sr²⁺)	Crimson red

Precautions

• Always use concentrated HCl to make the salt paste.
• Avoid Contamination: Ensure the wire is perfectly clean between tests.
• Safety: Do not touch the paste or acid with bare skin — HCl is corrosive.
• Eye Protection: Use protective eyewear; ensure proper ventilation.

Concept-Based Questions

Why use a cobalt glass?

It is used to detect calcium in the presence of sodium — the glass filters out the overpowering yellow of sodium.

What does flame duration indicate?

The duration of color can provide information about the metal ion's concentration in the sample.

Real-world applications?

The colorful light emission of fireworks and sodium vapor street lamps.

Why might a test fail?

Contamination of the wire or an impure sample can lead to incorrect results.

3 Preparation of Pure Crystals of CuSO₄·5H₂O Recrystallization of copper (II) sulfate pentahydrate ▶

Objective

To prepare pure crystalline copper (II) sulfate pentahydrate from an impure sample using recrystallization.

Theory

• Principle: Recrystallization is based on the fact that impurities remain dissolved or do not crystallize when a saturated hot solution is cooled.
• Chemical Equation: CuSO₄(s) + 5H₂O(l) → CuSO₄·5H₂O(s) [heat]
• Key Definitions:
  • Efflorescent: A substance (like CuSO₄·5H₂O) that loses water of crystallization in dry air, becoming a white anhydrous powder.
  • Anhydrous: Copper (II) sulfate (CuSO₄) with no water molecules.

Apparatus & Chemicals

• Bunsen burner / Spirit lamp
• Glass rod
• Funnel & filter paper
• China dish & sand bath
• Tripod stand
• Beakers
• Distilled water
• Impure copper sulfate

Procedure

1. Dissolving: Take 40 ml of water; add small amounts of impure copper sulfate while stirring until some remains undissolved (saturated solution).
2. Filtration: Filter the hot solution to remove insoluble impurities.
3. Concentration: Heat the filtrate in a china dish over a sand bath until the crystallization point is reached. Test: dip a glass rod; if a thin crust forms, it's ready.
4. Cooling: Allow the solution to cool slowly at room temperature.
5. Collection: Blue triclinic crystals of copper sulfate appear.

Observations

• Determining Water Molecules (x): Weigh hydrate → heat gently → weigh anhydrous CuSO₄ → subtract to find mass of water → use molar masses to find moles ratio.
• Visual: Pure crystals are bright blue and triclinic. Overheating produces pale blue or white.

Precautions

• Slow Cooling: Always cool the saturated solution slowly for large, pure crystals.
• Temperature: Prepare saturated solutions at room temperature before heating.
• Avoid Overheating: Do not evaporate to dryness — produces powder, not crystals, and may decompose the compound.

Concept-Based Questions

Why filter the solution while hot?

To prevent copper sulfate from crystallizing prematurely in the filter paper along with impurities.

Why is slow cooling necessary?

Slow cooling allows for orderly arrangement of molecules, resulting in large, well-formed, pure crystals. Fast cooling creates many small, impure crystals.

Why wash crystals with acetone instead of water?

Acetone doesn't dissolve crystals, removes impurities quickly, and evaporates fast. Water would partially dissolve crystals, reducing yield.

What is the role of a desiccator?

It's a sealed container that prevents crystals from losing their water of crystallization, maintaining their hydrate form.

Why store crystals in a corked bottle?

To prevent moisture loss (efflorescence) that would turn blue crystals into a white anhydrous powder.

4 Determining the Melting Point of Naphthalene Purity test via sharp-phase transition temperature ▶

Objective

To determine the melting point of Naphthalene.

Theory

• The melting point is a characteristic property of a substance.
• Pure substances melt at a specific, sharp temperature.
• Impurities lower the melting point and broaden the temperature range.

Apparatus & Chemicals

• Chemicals: Naphthalene sample, Liquid Paraffin (or water for the bath).
• Glassware: Thermometer, capillary tubes, watch glass, boiling tube/beaker.
• Heating Tools: Bunsen burner, tripod stand, wire gauze, heating block.
• Misc: Matches, thread (to attach capillary).

Procedure

1. Place a small amount of Naphthalene on a watch glass.
2. Seal one end of a capillary tube by heating it.
3. Pack a small quantity of the sample into the capillary tube.
4. Attach the capillary tube to the thermometer.
5. Place the assembly in the heating block/bath.
6. Heat gradually; record the temperature when the first drop of liquid appears.
7. Remove heat, cool, and record the temperature when the substance solidifies.
8. Calculate the mean of these two temperatures.
9. Repeat for accuracy.

Observations

Mean = (T₁ + T₂) / 2

• T₁: Temperature when melting starts.
• T₂: Temperature when solidification starts.

Reference Melting Points

Compound	Melting Point
Naphthalene	80 °C
Acetamide	82 °C
Diphenylamine	54 °C
Benzophenone	48 °C
Citric Acid	69 °C

Precautions

• Do not let the thermometer bulb touch the walls or base of the beaker.
• Ensure the thermometer bulb and sample are fully immersed.
• Stir the bath constantly for uniform heating.
• Apply heat gently.
• The open end of the capillary must remain above the liquid level.

Concept-Based Questions

Why use a water bath instead of direct flame?

Naphthalene is flammable — a bath ensures even, safe heating.

Why does temperature remain constant during melting?

Heat energy is being used to overcome forces between particles to change state, not to raise temperature.

Why seal one end of the capillary?

To prevent the sample from falling out or mixing with the heating liquid.

What does a lower melting point (e.g., 75 °C instead of 80 °C) indicate?

It indicates the sample is impure (freezing point depression).

What does a sharp melting point at exactly 80 °C indicate?

The substance is pure.