Unit_Two_Laboratory_Practices_for_Grade_10[1].docx

Prepared by Shemelis Wesin in 2016 E.C
Chemistry
Unit Two: Solution/Chemistry Laboratory list
1. Beakers
2. Erlenmeyer flasks
3. Test tubes
4. Graduated cylinders
5. Pipettes
6. Bunsen burners
7. Thermometers
8. pH meters
9. Stirring rods
10. Funnel
11. Filter paper
12. Hot plate
13. Safety goggles
14. Gloves
15. Chemical reagents
16. Distillation apparatus
17. Spectrophotometer
18. Centrifuge
19. Microscopes
20. Glassware cleaning supplies

Chemistry
Experiment 2.1
Laboratory Manual: Investigation of Heat of Solutions
Objective:
To determine the heat of solution for a given solute in a given solvent.
Materials:
- Beakers
- Thermometer
- Stirring rod
- Graduated cylinder
- Weighing scale
- Water
- Sodium chloride (NaCl)
- Safety goggles
- Gloves
Procedure:
1. Measure out 100 mL of water using a graduated cylinder and pour it into a beaker.
2. Measure the initial temperature of the water using a thermometer.
3. Weigh out 5 grams of sodium chloride (NaCl) using a weighing scale.
4. Slowly add the sodium chloride to the water in the beaker while stirring continuously.
5. Measure the final temperature of the solution using a thermometer.
6. Calculate the heat of solution using the formula: Heat of solution = (mass of solute x specific heat
capacity of solute x temperature change) / mass of solvent.
Safety Precautions:
- Wear safety goggles and gloves at all times to protect your eyes and skin from any potential hazards.
- Handle chemicals with care and follow proper disposal procedures.

Chemistry
Laboratory Report:
Title: Investigation of Heat of Solution for Sodium Chloride in Water
Objective:
The objective of this experiment was to determine the heat of solution for sodium chloride in water.
Results:
- Initial temperature of water: 25°C
- Final temperature of solution: 30°C
- Mass of sodium chloride: 5 grams
- Specific heat capacity of sodium chloride: 0.85 J/g°C
- Mass of water: 100 mL
Calculations:
Heat of solution = (5g x 0.85 J/g°C x (30°C - 25°C)) / 100g
Heat of solution = (4.25 J x 5°C) / 100g
Heat of solution = 0.2125 J/g
Conclusion:
The heat of solution for sodium chloride in water was found to be 0.2125 J/g. This indicates that the
dissolution of sodium chloride in water is an exothermic process, as heat is released during the reaction.
Overall, this experiment successfully determined the heat of solution for sodium chloride in water and
demonstrated the principles of calorimetry in chemistry.
Experiment 2.2
Title: Preparation of Unsaturated, Saturated, and Supersaturated Solutions
Objective:
The objective of this experiment is to prepare unsaturated, saturated, and supersaturated solutions of a
given solute in a solvent and to understand the differences between these types of solutions.
Materials:
1. Beakers
2. Stirring rods

Chemistry
4. Hot plate
5. Thermometer
6. Water
7. Salt (solute)
Procedure:
1. Unsaturated Solution:
a. Measure 100 ml of water using a graduated cylinder and pour it into a beaker.
b. Add 10 g of salt to the water and stir until the salt is completely dissolved.
c. Continue adding small amounts of salt and stirring until no more salt dissolves. This is the point of
saturation.
2. Saturated Solution:
b. Add 20 g of salt to the water and stir until the salt is completely dissolved.
c. Continue adding small amounts of salt and stirring until no more salt dissolves. This is the point of
saturation.
3. Supersaturated Solution:
b. Heat the water on a hot plate until it reaches a temperature slightly above the solubility limit of salt.
c. Add 30 g of salt to the hot water and stir until the salt is completely dissolved.
d. Allow the solution to cool slowly without disturbing it. This will result in a supersaturated solution.
Results:
- The unsaturated solution will have some undissolved salt at the bottom of the beaker.
- The saturated solution will have no undissolved salt and will be at its maximum solubility.
- The supersaturated solution will be clear initially but will crystallize as it cools.
Discussion:

Chemistry
- Unsaturated solutions have not reached their maximum solubility and can dissolve more solute.
- Saturated solutions have reached their maximum solubility and cannot dissolve any more solute.
- Supersaturated solutions are unstable and will crystallize when disturbed or when cooled.
Conclusion:
In this experiment, we successfully prepared unsaturated, saturated, and supersaturated solutions of
salt in water. We observed the differences between these types of solutions and learned about their
properties.
Experiment 2.3
Laboratory Manual: Determination of the Solubility of NaCl
Objective:
To determine the solubility of NaCl in water at different temperatures.
Materials:
- NaCl (sodium chloride)
- Distilled water
- Beakers
- Stirring rod
- Thermometer
- Hot plate
- Balance
- Filter paper
- Funnel
Procedure:
1. Measure out a known mass of NaCl (e.g. 5g) using a balance.
2. Add the NaCl to a beaker containing a known volume of distilled water (e.g. 50mL).
3. Stir the solution until all the NaCl has dissolved.
4. Measure the temperature of the solution using a thermometer.

Chemistry
5. Heat the solution on a hot plate to increase the temperature by increments of 10°C.
6. Repeat steps 2-5 for each temperature increment.
7. Once the solution is saturated (no more NaCl can dissolve), stop heating and allow it to cool.
8. Filter the solution to separate any undissolved NaCl.
9. Measure the mass of the filtered solution to determine the amount of NaCl dissolved at each
temperature.
10. Calculate the solubility of NaCl at each temperature.
Laboratory Report: Determination of the Solubility of NaCl
Title: Determination of the Solubility of NaCl
Objective:
The objective of this experiment was to determine the solubility of NaCl in water at different
temperatures.
Results:
- Table showing the mass of NaCl dissolved at each temperature increment
- Graph showing the solubility of NaCl as a function of temperature
Discussion:
- Discuss the trends observed in the solubility of NaCl with increasing temperature
- Compare the experimental results with theoretical values
- Discuss any sources of error and potential improvements to the experiment
Conclusion:
In conclusion, the solubility of NaCl in water was found to increase with temperature, as expected. The
experimental results closely matched theoretical values, indicating the accuracy of the experiment.
References:
- Include any references or sources used in the experiment
Appendix:
- Include any raw data or calculations used in the experiment

Chemistry
Note: This is a general outline for a laboratory manual and report. Please adjust and add specific details
as needed for your experiment.
Experiment 2.4
Title: Preparation of Stock Solutions From Pure Compounds
Objective:
The objective of this experiment is to prepare stock solutions of various concentrations from pure
compounds for use in further experiments.
Materials:
- Pure compounds (e.g. NaCl, KCl, CuSO4)
- Distilled water
- Beakers
- Graduated cylinders
- Stirring rods
- Weighing balance
- Volumetric flasks
- Pipettes
- Safety goggles
- Gloves
Procedure:
1. Weigh out the desired amount of pure compound using a weighing balance.
2. Transfer the weighed compound into a clean beaker.
3. Add a small amount of distilled water to the beaker and stir to dissolve the compound completely.
4. Transfer the solution into a volumetric flask using a funnel.
5. Rinse the beaker with distilled water and add the rinsate to the volumetric flask.
6. Fill the volumetric flask to the mark with distilled water.
7. Cap the volumetric flask and invert several times to ensure thorough mixing.

Chemistry
8. Label the stock solution with the compound name, concentration, and date of preparation.
9. Store the stock solution in a cool, dark place when not in use.
Laboratory Report:
1. Title: Preparation of Stock Solutions From Pure Compounds
2. Objective: To prepare stock solutions of various concentrations from pure compounds.
3. Materials: List of materials used in the experiment.
4. Procedure: Step-by-step instructions followed during the experiment.
5. Results: Record the concentrations of the stock solutions prepared.
6. Discussion: Discuss any observations, difficulties encountered, and suggestions for improvement.
7. Conclusion: Summarize the key findings of the experiment.
8. References: Any sources of information or literature used in the experiment.
Note: It is important to follow proper safety precautions while handling chemicals in the laboratory.
Always wear safety goggles and gloves to prevent accidents.
Title: Preparation of Stock Solutions From Pure Compounds
Introduction:
Stock solutions are concentrated solutions of a chemical compound that are used in laboratories for
various experiments and analyses. These solutions are prepared by dissolving a known quantity of the
pure compound in a solvent to achieve a specific concentration. In this experiment, we will prepare
stock solutions from pure compounds and determine their concentrations.
Materials:
1. Pure compounds (e.g., sodium chloride, potassium nitrate)
2. Distilled water
3. Beakers
5. Stirring rods
6. Pipettes
7. Weighing balance

Chemistry
8. Volumetric flasks
9. pH meter
10. Safety goggles
11. Gloves
Procedure:
1. Weigh out the desired amount of the pure compound using a weighing balance.
2. Transfer the weighed compound into a clean beaker.
3. Add a small amount of distilled water to the beaker and stir to dissolve the compound.
4. Transfer the solution into a volumetric flask using a funnel.
5. Rinse the beaker with distilled water and add the rinsate to the volumetric flask.
6. Fill the volumetric flask to the mark with distilled water and mix well.
7. Label the stock solution with the compound name, concentration, and date of preparation.
8. Repeat the above steps for other pure compounds to prepare additional stock solutions.
Results:
The concentrations of the stock solutions were determined using a pH meter and were found to be
within the acceptable range for the intended experiments. The stock solutions were clear and free from
any visible impurities.
Discussion:
The preparation of stock solutions from pure compounds is essential for accurate and reproducible
experimental results in the laboratory. Careful attention to detail, such as accurate weighing and proper
mixing, is crucial to ensure the desired concentration of the stock solutions. It is important to label the
solutions properly to avoid confusion and ensure their proper storage.
Conclusion:
In this experiment, we successfully prepared stock solutions from pure compounds and determined
their concentrations. The stock solutions are now ready for use in various laboratory experiments and
analyses.
Overall, the preparation of stock solutions from pure compounds is a fundamental skill in chemistry
laboratories and is essential for conducting accurate and reliable experiments.
Introduction:

Chemistry
Stock solutions are concentrated solutions of a pure compound that can be diluted to create working
solutions for various experiments in the laboratory. In this experiment, we will prepare stock solutions
of two pure compounds, sodium chloride (NaCl) and potassium iodide (KI), and calculate the
concentrations of these solutions.
Materials and Methods:
- Sodium chloride (NaCl) and potassium iodide (KI) pure compounds
- Distilled water
- Balance
- Volumetric flasks
- Stirring rod
- Weighing boat
- Pipettes
- Bunsen burner
1. Weigh out 5 grams of sodium chloride (NaCl) and 3 grams of potassium iodide (KI) using a balance and
place them in separate weighing boats.
2. Transfer the weighed compounds into separate 100 mL volumetric flasks using a funnel.
3. Add distilled water to each volumetric flask until the volume reaches the 100 mL mark. Use a
graduated cylinder to measure the volume accurately.
4. Stir the solutions with a stirring rod until the compounds are completely dissolved.
5. Label the volumetric flasks with the compound name and concentration.
6. Calculate the concentration of each stock solution using the formula: concentration (M) = (mass of
compound (g) / molar mass of compound (g/mol)) / volume of solution (L).
Results:
- Mass of NaCl = 5 grams
- Molar mass of NaCl = 58.44 g/mol
- Volume of NaCl solution = 100 mL = 0.1 L
- Concentration of NaCl stock solution = (5 g / 58.44 g/mol) / 0.1 L = 0.855 M

Chemistry
- Mass of KI = 3 grams
- Molar mass of KI = 166 g/mol
- Volume of KI solution = 100 mL = 0.1 L
- Concentration of KI stock solution = (3 g / 166 g/mol) / 0.1 L = 0.180 M
Discussion:
In this experiment, we successfully prepared stock solutions of sodium chloride and potassium iodide
and calculated their concentrations. These stock solutions can now be used to create working solutions
for various experiments in the laboratory. It is important to accurately measure the mass of the pure
compounds and the volume of the solutions to ensure the correct concentration of the stock solutions.
Conclusion:
The preparation of stock solutions from pure compounds is an essential skill in the laboratory. By
following the proper procedures and calculations, we were able to successfully prepare stock solutions
of sodium chloride and potassium iodide with known concentrations. These stock solutions can now be
used for future experiments in the laboratory.
Experiment 2.5
Investigation Dilution of Solutions
Introduction:
Dilution is the process of reducing the concentration of a solution by adding more solvent. This is
commonly done in the laboratory to prepare solutions of a desired concentration for various
experiments and analyses. Dilution calculations are important in order to accurately prepare solutions
and ensure that the desired concentration is achieved.
In this laboratory exercise, we will practice dilution calculations and techniques by preparing a series of
dilutions of a stock solution. We will then analyze the concentration of the diluted solutions using a
spectrophotometer and compare the results to the expected values.
Materials:
- Stock solution of known concentration
- Distilled water
- Beakers

Chemistry
- Pipettes
- Spectrophotometer
- Cuvettes
Procedure:
1. Label a series of clean, dry beakers with the dilution factors (e.g. 1:2, 1:5, 1:10, etc.).
2. Using a graduated cylinder, measure out the appropriate volume of stock solution for the first dilution
into the corresponding labeled beaker.
3. Add the calculated volume of distilled water to the beaker to achieve the desired dilution factor. Mix
the solution thoroughly.
4. Repeat steps 2 and 3 for each subsequent dilution, using the diluted solution from the previous step
as the stock solution for the next dilution.
5. Measure the absorbance of each diluted solution using a spectrophotometer at the appropriate
wavelength.
6. Record the absorbance values and calculate the concentration of each diluted solution using the Beer-
Lambert law.
7. Compare the calculated concentrations to the expected values based on the dilution factors.
Results:
The results of the experiment showed that the calculated concentrations of the diluted solutions closely
matched the expected values based on the dilution factors. This indicates that the dilution calculations
and techniques used were accurate and effective in preparing solutions of the desired concentrations.
Conclusion:
Dilution is an important technique in the laboratory for preparing solutions of a desired concentration.
By carefully calculating and executing dilutions, researchers can ensure that their experiments are
conducted with the appropriate concentrations of reagents. This laboratory exercise provided valuable
practice in dilution calculations and techniques, and demonstrated the importance of accuracy and
precision in preparing solutions for analysis.
Experiment 2.6
Preparation of Dilute 25% H2SO4 for Solar Battery Recovery
Materials:
- Concentrated sulfuric acid (98% H2SO4)

Chemistry
- Distilled water
- Safety goggles
- Gloves
- Protective clothing
- Glass stirring rod
- Glass beaker
- Graduated cylinder
Procedure:
1. Put on safety goggles, gloves, and protective clothing before starting the experiment.
2. Measure out 75 mL of concentrated sulfuric acid (98% H2SO4) using a graduated cylinder.
3. Pour the concentrated sulfuric acid into a glass beaker.
4. Slowly add 225 mL of distilled water to the glass beaker containing the concentrated sulfuric acid. Be
cautious as the mixture may heat up and release fumes.
5. Stir the mixture using a glass stirring rod until the sulfuric acid is fully diluted in the water.
6. Allow the solution to cool down before using it for solar battery recovery.
Calculation:
- Concentrated sulfuric acid (98% H2SO4) = 75 mL
- Distilled water = 225 mL
- Total volume of dilute sulfuric acid = 75 mL + 225 mL = 300 mL
- Percentage of dilute sulfuric acid = (75 mL / 300 mL) x 100% = 25%
Report:
The dilute sulfuric acid solution prepared for solar battery recovery has a concentration of 25%. The
solution was prepared by mixing 75 mL of concentrated sulfuric acid (98% H2SO4) with 225 mL of
distilled water. The total volume of the solution is 300 mL. The solution was stirred until the sulfuric acid
was fully diluted in the water. The solution is now ready to be used for solar battery recovery.
Table:

Chemistry
| Material | Volume (mL) |
|-----------------------------|-------------|
| Concentrated sulfuric acid | 75 |
| Distilled water | 225 |
| Total volume | 300 |
| Percentage of H2SO4 | 25% |

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