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

Class: Grade 10 Chemistry (SSC-II)
Board: Federal Board | FBISE | National Curriculum Pakistan
Topic: Practical Based Assessment (PBA)
Purpose: To help students understand experimental concepts, formulas, procedures, and viva questions for chemistry practical exams.
Difficulty Level: Conceptual + Exam Preparation Part: 1 Major Experiments

Fractional Distillation of Water and Alcohol

Objective: To separate a miscible mixture of water and alcohol (ethanol) using fractional distillation.

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

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2. Procedure

1. Setup: Assemble the apparatus as shown in the diagram, 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 highly volatile and flammable alcohol vapors from escaping.
4. Heating: Gradually heat the flask using a Bunsen burner.
5. Condensation: Turn on the water tap so cold water circulates through the condenser (entering at the lower inlet and exiting at the upper outlet).
6. Separation:
   • When the temperature reaches 78°C, alcohol vapors will pass through the condenser and be collected in the receiving flask.
   • Once the alcohol is fully distilled, the temperature will rise toward 100°C. At this point, replace the receiving flask to collect the pure water.

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3. Observations & Calculations

• Temperature Stability: During the collection of alcohol, the thermometer will remain steady at 78°C. It only begins to rise again once all alcohol has evaporated.
• Yield Calculation: In a mixture containing 30% alcohol, you should ideally expect a yield of approximately 30% of the initial volume, provided the distillation is efficient and there is no vapor loss.
• State Change: Inside the condenser, a physical change occurs where gas (vapor) turns back into a liquid due to cooling.

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4. Preventions & Safety Measures

• Bumping Prevention: Always add boiling chips to provide a "seed" for smooth boiling, preventing bumping (violent boiling) that could damage the flask.
• Fire Safety: Since alcohol is flammable, the heat source should be turned off or removed before the distillation flask runs completely dry.
• Volume Control: The distillation flask should only be partially filled to prevent the liquid from splashing into the condenser and contaminating the distillate.
• Cooling Efficiency: Cold water must always enter the lower inlet to ensure the entire condenser jacket remains full, maximizing the cooling effect.

Paper Chromatography

Objective: This study guide outlines the process of separating a green mixture of inks into its individual components using paper chromatography.

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Introduction

Paper Chromatography is a laboratory technique used to separate a mixture of chemical substances, particularly low molecular mass compounds, 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 (liquid) that moves through the paper.
• Capillary Action: The process that allows the solvent to travel up the paper strip.
• Principle: Separation occurs because different substances have different affinities for the stationary and mobile phases, causing them to move at different rates.

Apparatus and Chemicals

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

Procedure

1. Prepare the Strip: Cut a strip of chromatography paper and draw a horizontal baseline near the bottom using a pencil.
2. Apply Ink: Place small dots of the ink mixture onto the baseline. Ensure the dots are dry before proceeding.
3. Prepare the Chamber: Pour a small amount of solvent into the developing chamber. The level must be below the pencil baseline.
4. Develop the Chromatogram: Hang the paper strip vertically in the chamber so it does not touch the walls.
5. Seal the Container: Cover the chamber with a lid to create a saturated atmosphere, ensuring uniform solvent movement.
6. Separation: Allow the solvent to travel about 3/4 of the way up the paper. Remove the strip and immediately mark the "solvent front" with a pencil.

Calculation and Observation

The Retention Factor ($R_f$) is used to identify components. It is calculated using the following formula:

$$R_f = \frac{\text{Distance travelled by the solute (ink component)}}{\text{Distance travelled by the solvent front}}$$

Observation Table (Sample Data):

Component (Color)	Distance of Component (cm)	Distance of Solvent Front (cm)	$R_f$ Value
Blue	3	17	0.18
Red	6	17	0.35
Green	9	17	0.53

Precautions

• Always handle the paper by the edges to avoid contamination from skin oils.
• Use a pencil for the baseline, as ink from a pen would dissolve and interfere with results.
• The ink spots must be above the solvent level to prevent them from dissolving directly into the reservoir.
• The paper must be vertical and not touching the walls of the chamber.
• The ink dots should be small and controlled to prevent "tailing" or overlapping of 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, which prevents evaporation from the paper and ensures a smoother, more 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, while those with a higher affinity for the paper move slower.
• What can be concluded if two different ink samples produce identical chromatograms? They likely contain the same dyes and may come from the same source.
• Why is a water-ethanol mixture (7:3) preferred? This specific blend provides a balanced polarity that effectively dissolves and separates a wide range of ink dyes.

Separation of $Pb^{2+}$ and $Cd^{2+}$ Ions by Paper Chromatography

Objective: The goal of this experiment is to separate a mixture of Lead ($Pb^{2+}$) and Cadmium ($Cd^{2+}$) ions using paper chromatography and to determine their individual $R_f$ values.

Theory

• Definition: Paper chromatography is a technique used to separate and analyze the components of a mixture based on their distribution between a stationary phase (filter paper) and a mobile phase (solvent).
• Principle: Different substances move at different speeds because of their varying solubility in the solvent and their affinity for the chromatography paper.

Apparatus and Chemicals

• Materials: Chromatography paper, pencil, ruler, glass/plastic containers, clips, and a developing chamber with a lid.
• Solvents: Water, isopropyl alcohol, or acetone.
• Specific Solvent Mixture: n-butanol + 3 M $HNO_3$.
• Sample: Mixture of Lead and Cadmium chlorides or nitrates.
• Locating Agent: Concentrated Hydrogen sulphide ($H_2S$) gas or dithizone solution.

Procedure

1. Prepare the Strip: Cut chromatography paper to size and draw a pencil baseline near the bottom.
2. Spot the Sample: Apply a small volume of the mixture to the baseline using a capillary tube. Ensure the spot is dry before proceeding.
3. Prepare Chamber: Pour the solvent mixture into the chamber. Ensure the solvent level is below the baseline.
4. Development: Hang the paper vertically in the chamber. Close the lid to saturate the environment with solvent vapors.
5. Separation: Allow the solvent to travel up the paper. It will carry the $Pb^{2+}$ and $Cd^{2+}$ ions at different rates.
6. Mark the Front: Remove the paper before the solvent reaches the top and immediately mark the "solvent front" with a pencil.
7. Visualization: Spray the dried paper with a locating agent (like $H_2S$) to make the colorless spots visible.

Calculation and Observation

The Retention Factor ($R_f$) is calculated as follows:

$$R_f = \frac{\text{Distance travelled by the solute}}{\text{Distance travelled by the solvent}}$$

Ion	Color of Spot	Distance from Baseline	$R_f$ Value
Lead ($Pb^{2+}$)	Black	3.5 cm	0.2
Cadmium ($Cd^{2+}$)	Yellow	5.3 cm	0.3

*Note: Example distance from solvent front = 17.6 cm.

Precautions

• The applied sample spot must be very small and completely dry to prevent tailing or smearing.
• The baseline must never be submerged in the solvent; otherwise, the ions will dissolve into the bulk solvent.
• Do not let the paper touch the glass walls of the chamber to ensure a uniform solvent flow.
• The solvent front must be marked immediately before the solvent evaporates.

Concept-Based Questions

Q: Why is a locating agent necessary for $Pb^{2+}$ and $Cd^{2+}$?These ions are colorless in solution. A reagent like $H_2S$ reacts with them to form colored sulphides ($PbS$ and $CdS$) that are visible to the eye.

Q: Why use a pencil instead of a pen for the baseline?Pencil graphite is inert and insoluble. Ink contains dyes that would dissolve and travel with the solvent, ruining the chromatogram.

Q: Why does $Cd^{2+}$ travel further than $Pb^{2+}$?$Cd^{2+}$ has a higher affinity for the mobile phase (solvent), whereas $Pb^{2+}$ interacts more strongly with the stationary phase (paper).

Q: Why must the chamber be saturated with solvent vapor?Saturation prevents the solvent from evaporating off the paper, ensuring a constant and uniform flow for accurate $R_f$ calculations.

Titration Experiment

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

• Standard Solution: Hydrochloric acid ($HCl$) is used as the known standard to find the concentration of the $NaOH$ solution.

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Theory

• Chemical Equation: $NaOH + HCl \rightarrow NaCl + H_2O$
• Molar Ratio: The stoichiometric ratio of $HCl$ to $NaOH$ is $1:1$.
• Indicator: Phenolphthalein is used to signal the end point.
• End Point: The point where the solution changes from light pink to colorless.

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Apparatus and Chemicals

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

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Procedure

1. Preparation: Rinse the pipette with $NaOH$ and the burette with the given $HCl$ solution.
2. Sampling: Pipette $10\text{ cm}^3$ of $NaOH$ into a conical flask and add 1–2 drops of phenolphthalein (the solution will turn pink).
3. Setting the Burette: Fill the burette with $HCl$ using a funnel, then remove the funnel and ensure no air bubbles are trapped in the nozzle.
4. Rough Titration: Add $HCl$ from the burette while swirling the flask until the pink color just disappears.
5. Precise Titration: Repeat the process, adding the acid drop-wise near the end point until three concordant readings (results within $0.1\text{ cm}^3$) are obtained.

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Precautions

• Rinsing: Always rinse the burette and pipette with the specific solutions they will contain to prevent dilution by water.
• Funnel Removal: Remove the glass funnel from the burette before taking readings to prevent accidental drops from altering the volume.
• Visual Aid: Use a white tile or white paper under the conical flask to observe the color change more clearly.
• Reading the Burette: Always keep the eye at the same level as the lower meniscus to avoid parallax error.
• Nozzle Check: Ensure the burette nozzle is completely filled with solution and free of air bubbles before starting.

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Calculation and Observation

• Standard Formula: $$\frac{M_1 V_1}{n_1} = \frac{M_2 V_2}{n_2}$$
• Variables:
  • $M_1$: Molarity of acid solution ($HCl$)
  • $V_1$: Mean volume of acid used (from burette readings)
  • $n_1$: Moles of acid in the balanced equation (1)
  • $M_2$: Molarity of base solution ($NaOH$)
  • $V_2$: Volume of base solution used ($10\text{ cm}^3$)
  • $n_2$: Moles of base in the balanced equation (1)

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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 is 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.

Detection and Confirmation of Gases

Objective: This study guide covers the identification of three specific gases: Ammonia ($\text{NH}_3$), Carbon Dioxide ($\text{CO}_2$), and Chlorine ($\text{Cl}_2$). Identification is achieved by observing distinct chemical reactions or physical properties when the gases are exposed to specific reagents.

Theory

• Ammonia ($\text{NH}_3$): A basic gas. It reacts with water on damp litmus paper to form ammonium hydroxide ($\text{NH}_4\text{OH}$), which is alkaline and changes red litmus to blue.
• Carbon Dioxide ($\text{CO}_2$): Reacts with limewater (calcium hydroxide) to form calcium carbonate ($\text{CaCO}_3$), an insoluble white precipitate that makes the solution appear milky.
• Chlorine ($\text{Cl}_2$): An acidic gas and a strong oxidizing agent. It reacts with water to form hydrochloric acid ($\text{HCl}$) and hypochlorous acid ($\text{HClO}$), which first turns blue litmus red and then bleaches it white.

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Apparatus and Chemicals

• Equipment: Test tubes, gas jars, delivery tubes, and Bunsen burner.
• Indicators: Red litmus paper, blue litmus paper, and starch-iodide paper.
• Reagents: Limewater ($\text{Ca(OH)}_2$ solution), water (for damping litmus).

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Procedure

1. For Ammonia: Hold a piece of damp red litmus paper near the mouth of the container releasing the gas.
2. For Carbon Dioxide: 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 of the chlorine gas.

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Observations and Results

Gas	Observation	Chemical Equation
Ammonia ($\text{NH}_3$)	Damp red litmus paper turns Blue.	$\text{NH}_{3(g)} + \text{H}_2\text{O}_{(l)} \rightarrow \text{NH}_4^+ + \text{OH}^-$
Carbon Dioxide ($\text{CO}_2$)	Limewater turns Milky/Cloudy.	$\text{Ca(OH)}_{2(aq)} + \text{CO}_{2(g)} \rightarrow \text{CaCO}_{3(s)} + \text{H}_2\text{O}_{(l)}$
Chlorine ($\text{Cl}_2$)	Blue litmus turns Red then White (Bleached). Starch-iodide paper turns Blue.	$\text{Cl}_{2(g)} + \text{H}_2\text{O}_{(l)} \rightarrow \text{HCl}_{(aq)} + \text{HClO}_{(aq)}$

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Precautions

• Understand safety measures for handling toxic or irritating gases.
• Ensure proper ventilation in the laboratory to avoid exposure to fumes.
• Be aware of the hazards associated with specific chemicals before handling.
• Use damp litmus paper to allow the gas to dissolve and react chemically.

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Concept-Based Questions

• Why must litmus paper be damp? Water is required for the gas to dissolve and react (e.g., forming $\text{NH}_4\text{OH}$ or $\text{HClO}$) to produce the ions that cause a color change.
• How do you distinguish $\text{NH}_3$ from other gases? Only $\text{NH}_3$ is basic enough to turn red litmus paper blue.
• What causes the milkiness in the limewater test? The formation of insoluble calcium carbonate ($\text{CaCO}_3$) solid particles.
• Why does chlorine bleach litmus paper? Chlorine is a strong oxidizing agent. It reacts with water to form hypochlorous acid ($\text{HClO}$), which destroys the dye in the litmus paper.
• What are the daily life applications of gas detection?
  • Monitoring air quality and pollution.
  • Industrial safety to detect toxic leaks.
  • Detecting leaks in natural gas lines (Home Safety).
  • Medical diagnosis through respiratory function assessment.

Class: Grade 10 Chemistry (SSC-II)
Board: Federal Board | FBISE | National Curriculum Pakistan
Topic: Practical Based Assessment (PBA)
Purpose: To help students understand experimental concepts, formulas, procedures, and viva questions for Chemistry practical exams.
Difficulty Level: Conceptual + Exam Preparation Part: 2 Minor Experiments

Separation by Sublimation

Objective: This practical demonstrates the separation of naphthalene from a mixture of sand and salt using the method of sublimation.

Theory

• Definition: Sublimation is the process where a solid changes directly into vapors upon heating, and returns directly to a solid state upon cooling, without ever becoming a liquid.
• Chemical Behavior: The process is effective for substances like naphthalene, ammonium chloride, iodine, and camphor, which possess sublimable properties.

Apparatus and Chemicals

• Apparatus: China dish (or watch glass), Funnel, Cotton, Sand bath, Tripod stand, Burner (or spirit lamp), Wire gauze, Filter paper.
• Chemicals: Sand and Naphthalene.

Procedure

1. Place 5 to 8 g of the impure sample in a china dish.
2. Set the china 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 to trap vapors.
5. Heat the mixture very gently. Naphthalene will sublime and deposit as crystals on the filter paper.
6. Stop heating once all naphthalene has sublimed. Allow the apparatus to cool completely before removing the funnel to collect the pure crystals.

Precautions

• Gentle Heating: Always heat gently to avoid melting the naphthalene instead of subliming it.
• Safety & Yield: Ensure the funnel is cooled before removal to prevent loss of product or accidental damage.

Observations and Calculations

• Completion Signal: The process is complete when no more vapors appear inside the funnel and only sand remains in the china dish.
• Result: Pure naphthalene is recovered from the filter paper, leaving the non-volatile sand behind.

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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 would transfer 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.

Identification of Metal Ions by Flame Test

Objective: The primary objective of this experiment is 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.

Introduction

The flame test is a qualitative analysis technique used to identify the presence of specific metal ions in a sample based on the characteristic color they impart to a Bunsen burner flame.

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Theory

• Principle: When a metal ion is heated in a flame, its electrons absorb energy and jump to higher energy levels (excitation). As these electrons return to their original lower energy levels, they release energy in the form of light.
• Characteristic Colors: The wavelength of light emitted is unique to each element, producing a specific color that serves as a "chemical fingerprint."
• Metal Chlorides: Concentrated HCl is used because metal chlorides are more volatile than other salts, making them easier to evaporate and introduce into the flame.
• Electron Binding: Some metals do not impart color if their valence electrons are too tightly bound to be excited by the flame's heat.

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Apparatus and Chemicals

• Apparatus: Platinum or nichrome wire, watch glass, Bunsen burner, matches.
• Chemicals: Concentrated Hydrochloric Acid (HCl), salts of Sodium (Na⁺), Calcium (Ca²⁺), Copper (Cu²⁺), Potassium (K⁺), Barium (Ba²⁺), and Strontium (Sr²⁺).

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Procedure

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

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Observations

The following colors are observed for specific metal cations:

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

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Precautions

• Always use concentrated HCl to make the salt paste.
• Avoid Contamination: Ensure the wire is perfectly clean between tests to avoid mixed or unclear colors.
• Safety: Do not touch the paste or acid with bare skin (HCl is corrosive).
• Eye Protection: Use protective eyewear and ensure proper ventilation to avoid inhaling fumes.

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Concept-Based Questions

• Why use a cobalt glass? It is used to detect calcium ions in the presence of sodium, as the glass filters out the overpowering yellow flame of sodium.
• What does flame duration indicate? The duration of the color can provide information regarding the concentration of the metal ion in the sample.
• What are real-world applications? This principle is seen in the colorful light emissions of fireworks and in sodium vapor street lamps.
• Why might a test fail? Contamination of the wire or using an impure sample can lead to incorrect results.

Preparation of Pure Crystals of $CuSO_4 \cdot 5H_2O$

Objective: The goal of this practical is to prepare pure crystalline copper (II) sulfate pentahydrate from an impure sample using the technique of recrystallization.

• Principle: Recrystallization is based on the fact that impurities remain dissolved or do not crystallize when a saturated hot solution is cooled.
• Chemical Equation:

  $$CuSO_{4(s)} + 5H_2O_{(l)} \xrightarrow{heat} CuSO_4 \cdot 5H_2O_{(s)}$$

• Key Definitions:
  • Efflorescent: A substance (like $CuSO_4 \cdot 5H_2O$) that loses water of crystallization in dry air, turning into a white anhydrous powder.
  • Anhydrous: Copper (II) sulfate ($CuSO_4$) that does not contain any water molecules.

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Apparatus and Chemicals

• Bunsen burner / Spirit lamp
• Glass rod
• Funnel & Filter paper
• China dish & Sand bath

• Tripod stand
• Beakers
• Distilled water
• Impure copper sulfate

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Procedure

1. Dissolving: Take 40ml of water in a beaker and add small amounts of impure copper sulfate while stirring with a glass rod 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 into the hot solution; if a thin crust forms on the rod upon blowing gently, it is ready.
4. Cooling: Allow the solution to cool slowly at room temperature.
5. Collection: After some time, blue triclinic crystals of copper sulfate will appear.

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Precautions

• Slow Cooling: Always cool the saturated solution slowly to ensure the formation of large, pure crystals.
• Temperature: Saturated solutions should be prepared at room temperature before heating.
• Avoid Overheating: Do not evaporate to dryness, as this produces a powder rather than crystals and may decompose the compound.

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Calculations and Observations

• Determining Water Molecules ($x$):
  1. Weigh the hydrate ($CuSO_4 \cdot xH_2O$).
  2. Heat gently to drive off water.
  3. Weigh the remaining anhydrous $CuSO_4$.
  4. Subtract weights to find the mass of water lost.
  5. Use molar masses to calculate the ratio of moles of water to moles of $CuSO_4$.
• Visual Observation: Pure crystals are bright blue and triclinic in shape. Overheating results in a pale blue or white appearance.

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Concept-Based Questions

• Why filter the solution while hot? To prevent the copper sulfate from crystallizing prematurely in the filter paper along with impurities.
• Why is slow cooling necessary? Slow cooling allows for an orderly arrangement of molecules, resulting in large, well-formed, and pure crystals. Fast cooling creates many small, impure crystals.
• Why wash crystals with acetone instead of water? Acetone does not dissolve the crystals, removes impurities quickly, and evaporates fast. Water would partially dissolve the crystals, leading to a loss of yield.
• What is the role of a desiccator? It is a sealed container used to prevent crystals from losing their water of crystallization to the air, maintaining their hydrate form and structure.
• Why store crystals in a corked bottle? To prevent moisture loss (efflorescence) that would turn the blue crystals into a white anhydrous powder.

Determining the Melting Point of Naphthalene

Objective: To determine the melting point of Naphthalene.

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Theory (Principle)

• 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.

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Apparatus and 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).

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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 into the heating block/bath.
6. Heat gradually; record the temperature when the first drop of liquid appears.
7. Remove heat, let it cool, and record the temperature when the substance solidifies.
8. Calculate the mean of these two temperatures.
9. Repeat the experiment for accuracy.

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Calculations and Observations

The melting point is determined by taking the average of two readings:

• Observation 1 ($T_1$): Temperature when melting starts.
• Observation 2 ($T_2$): Temperature when solidification starts.
• Mean Value Calculation: $\frac{T_1 + T_2}{2}$

Reference Melting Points:

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

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Precautions

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

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Concept-Based Questions

• Why use a water bath instead of direct flame? Naphthalene is flammable and could catch fire; a bath ensures even, safe heating.
• Why does the temperature remain constant during melting? Heat energy is being used to overcome forces between particles to change the state from solid to liquid.
• 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? It indicates the substance is pure.
• Why heat gently? To allow sufficient time for the phase change to occur and be recorded accurately.

Determining the Boiling Point of Ethyl Alcohol

Objective:This study guide covers the experimental procedure for finding the boiling point of Ethanol ($C_2H_5OH$).

Introduction & Theory

• Definition: The boiling point is the temperature at which a liquid's vapor pressure equals the external (atmospheric) pressure.
• Key Principle: $\text{Boiling point} \propto \text{External pressure}$.
  • Higher external pressure increases the boiling point.
  • Lower external pressure decreases the boiling point.
• Impurity Factor: The presence of impurities generally increases the boiling point of a substance.

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Apparatus and Chemicals

• Apparatus: Glass beaker, Tripod stand, Fusion tube, Dropper, Thermometer, Wire gauze, Capillary tube, Thread/Rubber band, Bunsen burner, Glass stirrer, Iron stand with clamp, and Matchbox.
• Chemicals: Ethyl Alcohol ($C_2H_5OH$) and Water (for the heating bath).

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Procedure

1. Seal one end of a capillary tube by heating it in a flame.
2. Attach the fusion tube to a thermometer using thread or a rubber band, ensuring the bulb of the thermometer and the base of the fusion tube are at the same level.
3. Clamp the thermometer to an iron stand and suspend it in a beaker containing $160 \text{ cm}^3$ of water.
4. Fill the fusion tube one-third full with Ethanol using a dropper.
5. Place the capillary tube into the fusion tube with the closed end upward.
6. Heat the water beaker gently while stirring continuously to ensure uniform temperature.
7. Observe the capillary tube; the liquid will begin to fill it when vapor pressure matches atmospheric pressure.
8. Record the temperature the moment a continuous stream of bubbles emerges from the capillary tube.
9. Repeat the experiment twice to calculate a mean value for accuracy.

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Calculations and Observation

Observations are recorded to find the average boiling point:

No. of obs.	Temperature ($T$) at which bubbles start (°C)
1.	$T_1 = 78.6\text{°C}$
2.	$T_2 = 78.2\text{°C}$

Mean Calculation: $\frac{78.6 + 78.2}{2} = 78.4\text{°C}$

Conclusion: The boiling point of the given Ethanol is $78.4\text{°C}$.

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Precautions & Technical Details

• Stirring: Essential to maintain uniform temperature in the water bath and prevent localized overheating.
• Gentle Heating: Ensures a gradual temperature rise for a more accurate and stable reading.
• Alignment: The thermometer bulb must be level with the fusion tube base to measure the actual sample temperature rather than the surrounding bath.
• High Boiling Liquids: If a liquid's boiling point exceeds $100\text{°C}$, the water bath must be replaced with concentrated sulfuric acid or liquid paraffin/mineral oil.

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Concept-Based Questions

• What is the function of the capillary tube? It traps air/vapor so that internal pressure can be compared against atmospheric pressure.
• What is the precise sign that boiling point is reached? A continuous, steady stream of bubbles exiting the capillary tube.
• Why is the capillary placed with the closed end up? To trap a pocket of air/vapor; without this seal, the pressure comparison mechanism wouldn't work.
• Why repeat the experiment? To minimize random errors and ensure the result is reliable.
• Alternative Observation: During cooling, the boiling point can be noted when the liquid is suddenly drawn back into the capillary tube as external pressure overcomes vapor pressure.

Metal Displacement Reactions

Objective:This study guide covers the essential concepts and procedures for Finding the Critical Angle of a Prism

Introduction

This practical demonstrates a metal displacement reaction in an aqueous medium, focusing on how a more reactive metal displaces a less reactive metal from its salt solution.

Theory

The ability of one metal to displace another is determined by the Reactivity Series.

• Core Principle: A more reactive metal (e.g., Iron) will displace a less reactive metal (e.g., Copper) from its compound.
• The Reactivity Series: Highly reactive metals include Potassium (K), Sodium (Na), and Calcium (Ca). Less reactive metals include Gold (Au), Silver (Ag), and Copper (Cu).
• Chemical Equation:
  $$Fe_{(s)} + CuSO_{4(aq)} \longrightarrow FeSO_{4(aq)} + Cu_{(s)}$$ (Blue solution changes to a Green solution)

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Apparatus and Chemicals

• Apparatus: Beaker ($250 \text{ cm}^3$), stirrer (glass rod), and iron nails.
• Chemicals: Copper sulphate solution and distilled water.

Procedure

1. Add $100 \text{ cm}^3$ of distilled water to a beaker and add $3\text{--}4$ copper sulphate crystals.
2. Stir well with a glass rod to dissolve the $CuSO_4$ completely.
3. Add iron nails to the beaker and allow them to stand for some time.

Observations and Results

• Color Change: The blue copper sulphate solution gradually changes to light green due to the formation of $FeSO_4$.
• Physical Change: A pink precipitate of elemental copper deposits on the surface of the iron nails.

Precautions

• The iron nails used in the experiment must be free from rust.

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Concept-Based Questions

Q: What factors determine which metal will displace another?

• The reactivity series determines this; a more reactive metal is able to displace a less reactive metal from its compound.

Q: What are the general observations during a metal displacement reaction?

1. Precipitation: The less reactive metal falls out of the solution as a solid.
2. Color Change: The solution's color shifts based on the new metal salt formed.
3. Gas Evolution: In some cases, gases like hydrogen may be evolved.

Q: What happens if a less reactive metal (Copper) is placed in a solution of a more reactive metal (Zinc Sulfate)?

• No reaction occurs because copper is less reactive than zinc. Displacement only happens when the solid metal is more reactive than the metal ion in the solution.

Q: How does concentration affect the reaction?

• Higher concentrations can increase the reaction rate because more metal ions are available for displacement, while lower concentrations may slow the process.

Q: What are the industrial applications of these reactions?

• Extraction of metals from ores, corrosion processes, metal plating (like silver or gold), and purifying metals by replacing impurities.

Testing for Water with Anhydrous Copper (II) Sulfate

Objective: This practical demonstrates how to use anhydrous copper (II) sulfate as a chemical test to detect the presence of moisture or water in various substances.

Theory

The test relies on a reversible hydration reaction. Anhydrous copper (II) sulfate is a white powder. When it comes into contact with water, it absorbs the water molecules to become hydrated, resulting in a distinct color change from white to blue.

The Chemical Reaction:

$$\text{CuSO}_4 (\text{s}) + 5\text{H}_2\text{O} (\text{l}) \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O} (\text{s})$$

• Forward reaction: Hydration (white to blue).
• Reverse reaction: Dehydration via heating (blue to white).

Apparatus and Chemicals

• Chemicals: Anhydrous copper (II) sulfate ($\text{CuSO}_4$), Test substance (liquid or solid).
• Equipment: Test tubes or small beakers, Spatula, Heat source (Bunsen burner/hot plate), Tongs, Glass rod, Petri dish.

Procedure

1. Drying the Reagent: Heat a small amount of copper (II) sulfate gently in a dry test tube until it turns completely white. Allow it to cool before use.
2. Preparation: Place a small amount of the white anhydrous powder into a clean, dry test tube.
3. Testing Liquids: Add a few drops of the sample liquid to the powder.
4. Testing Solids: Mix a small amount of the solid sample with the powder.
5. Heating (if required): For solid samples, heat the mixture gently to release any trapped moisture.
6. Observation: Stir with a glass rod and watch for a color change.

Calculations and Observation

• Positive Result: If water is present, the powder turns from white to blue as it forms hydrated copper (II) sulfate pentahydrate ($\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$).
• Negative Result: If no water is present, the powder remains white.

Precautions

• Cooling: Always allow the heated anhydrous powder to cool before adding a sample to prevent premature evaporation.
• Dry Equipment: Ensure test tubes and glass rods are completely dry to avoid a false positive result.
• Gentle Heating: Use gentle heat when drying the sulfate to avoid decomposing the chemical.
• Safety: Always use tongs when handling hot test tubes to avoid burns.

Concept-Based Questions

• Why must the sulfate be anhydrous before the test? Only the anhydrous form is white; if it is already hydrated (blue), it cannot show a visible change when water is added.
• What factors affect the accuracy of this test? Low water concentration, impurities in the sample, or using sulfate that is already partially hydrated can make the result difficult to see.
• Is this a qualitative or quantitative test? It is a qualitative test because it indicates the presence of water through color change but does not measure the amount of water.
• Why is a control test recommended? Testing with pure water serves as a benchmark to confirm the anhydrous copper (II) sulfate is reacting correctly.
• What is the significance of the reaction's reversibility? It allows the copper (II) sulfate to be regenerated and reused by heating the blue hydrated form back into the white anhydrous form.

Testing Water Purity via Melting and Boiling Points

Objective: The purity of a water sample can be determined by measuring its melting point and boiling point. Pure substances have specific, fixed temperatures at which these phase changes occur at standard pressure (1 atm).

Theory

• Pure Water: Melts at 0°C and boils at 100°C.
• Impure Water: Impurities alter these fixed points:
  • Freezing Point Depression: Dissolved impurities (like salt or sugar) lower the melting/freezing point below 0°C.
  • Boiling Point Elevation: Non-volatile impurities (like salts or minerals) raise the boiling point above 100°C.
  • Volatile Impurities: Impurities like alcohol can lower the boiling point below 100°C.

Apparatus and Chemicals

• Thermometer (accuracy to 0.1°C preferred)
• Beakers (for holding water and ice)
• Heat source (Bunsen burner or hot plate)
• Tripod and gauze
• Stirring rod and glass rod
• Water sample and Crushed ice
• Heat-resistant gloves

Procedure

A. Melting Point Test

1. Prepare ice from the water sample and place it in a beaker with crushed ice.
2. Insert the thermometer into the center of the ice-water mixture without touching the beaker walls.
3. Stir gently with a glass rod to ensure uniform temperature.
4. Observe the temperature when the ice begins to melt. For pure water, this must be exactly 0°C and remain constant throughout the process.

B. Boiling Point Test

1. Fill a beaker with the water sample and set it on a tripod over a Bunsen burner.
2. Place the thermometer in the water (avoiding contact with the container).
3. Heat the water gently, stirring occasionally for uniform heating.
4. Monitor the temperature until the water boils. Pure water should boil at exactly 100°C at sea level.

Observations

Test	Pure Water	Impure Water
Melting Point	Exactly 0°C	Below 0°C (Depressed)
Boiling Point	Exactly 100°C	Above 100°C (Non-volatile) or Below 100°C (Volatile)

Precautions

• Thermometer Placement: Do not let the thermometer touch the bottom or sides of the beaker; it should measure the liquid/ice temperature, not the container's heat.
• Stirring: Stir gently to maintain thermal equilibrium and a uniform temperature throughout the sample.
• Ice Size: Use crushed ice rather than large cubes to provide a larger surface area for faster, more effective heat transfer.

Concept-Based Questions

• Why does the temperature stay constant at 100°C during boiling? It indicates only one substance is vaporizing.
• What does a melting point of -2°C indicate? It confirms the water is impure (contains solute) due to Freezing Point Depression.
• How do salt or sugar affect boiling? They reduce the water's ability to evaporate, requiring a higher temperature (Boiling Point Elevation).
• Why use a high-accuracy thermometer (0.1°C)? Melting points are sharp; small deviations are critical for identifying even minor impurities.