Preparation of Known Molar Solutions of Copper Sulfate Pentahydrate – ($$\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$$)

Preparation of Known Molar Solutions of Copper Sulfate Pentahydrate – ($\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$)

Prelab Requirements

• Read through the handout thoroughly.
• Complete the prelab quiz available on D2L.
• Finish the moles and molarity worksheet (provided as a separate document).
• Perform all calculations related to "molecular mass," "mole," and "molarity" as detailed in the document.
• Completely fill out Table 1.
• It is crucial to arrive in class prepared with the correct amount of $\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$ already calculated, to start dissolving it immediately.

Safety Notes

• Pipetting: Never pipet by mouth. Always use a pipet pump.
• Glassware: Handle glassware with care to prevent breakage and serious injury.
• Reporting: Report any broken glass or injuries to your instructor immediately.
• Eye Protection: Wear safety glasses at all times when working with solutions and during clean-up.
• Chemical Contact: Avoid getting Copper Sulfate ($\text{CuSO}_4$) on skin or in eyes. Do not inhale or ingest it.
• Waste Disposal: Dispose of all waste solutions in the designated appropriate containers, not down the sink.

Post-lab Clean Up Procedures

• Test Tubes: Empty contents into the specified waste beaker. Clean using a test tube brush and rinse thoroughly. Place clean test tubes upside down in a rack to drain.
• Group Waste: Empty your group's waste beaker into the large, labeled beaker for the entire class's waste.
• Glassware: Rinse all glassware, stir bars, scoopulas, and cuvettes thoroughly. Dry them gently.
• Glass Pipets: If a tub is provided, place glass pipets there to prevent contents from drying inside. Consult your lab instructor for specific instructions on glass pipet disposal/storage.
• Disposables: Discard disposable pipets, used weigh paper, and used Kimwipes.
• Equipment: Wipe down Labquests, colorimeters, electronic balances, and stir plates. Ensure they are turned off and returned neatly to their designated cart or counter space.
• Kit Items: Place all "kit" equipment back into your group’s tray. Ensure the tray contains all components, including any unused weigh paper, clean glass pipets, and clean disposable pipets (which may need to be replenished by your instructor, along with $\text{CuSO}_4$).
• General: Return all other used items (clean and neatly) to the cart or counter.

Purpose of the Experiment

• To understand the concept of Molecular Mass (also referred to as Molecular Weight) of a compound.
• To grasp the concept of Moles of a compound.
• To learn how to calculate Molarity.
• To use the Vernier/colorimeter to determine the light absorption of eight different molar concentrations of Copper Sulfate Pentahydrate ($\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$) at a wavelength of $635 \text{ nm}$.
• To create a graph, plot a standard curve using the molar standard solutions, and subsequently determine the concentration of three unknown solutions.
• To be able to define the following key terms:
  • Molecular formula
  • Atomic mass
  • Atomic number
  • Molecular mass
  • Mole
  • Molarity
  • Solvent
  • Solute
  • Solution
  • Standard curve
  • Absorbance
  • Colorimeter

Concepts: Moles and Molarity

Molecular Formula

• Definition: Represents the total number and types of each atom present in a molecule.
• Example: For Copper Sulfate Pentahydrate, the molecular formula is $\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$. This indicates:
  • One Copper (Cu) atom.
  • One Sulfur (S) atom.
  • Nine Oxygen (O) in total (four from $\text{SO}<em>4$ and five from $5\text{H}</em>2\text{O}$).
  • Ten Hydrogen (H) atoms (from $5\text{H}_2\text{O}$).

Molecular Mass

• Definition: The sum of the atomic weights of all atoms within a molecule.
• Calculation Process:
  1. Determine the atomic mass for each individual atom in the molecule by looking it up on the periodic table. The atomic mass is typically found at the bottom of each element's entry (or sometimes at the top), and it is always larger than the atomic number.
  2. Count the number of times each element appears in one molecule.
  3. Multiply the atomic mass of each element by its count in the molecule.
  4. Sum these values to get the total molecular mass.
  5. Important Note: For hydrate molecules like $\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$, the mass of the water molecules must be included in the total molecular mass calculation.
• Units: The units for Molecular Mass are $\text{g/mole}$.
• Calculation Example (Student Task, blanks to be filled):
  • Copper atomic mass = \text{_ g/mole}
  • Sulfur atomic mass = \text{_ g/mole}
  • Oxygen atomic mass = $\text{<strong><em><em></em></strong> g/mole} \times 9 \text{ oxygen atoms} = \text{<strong><em></em></em></strong> g/mole}$
  • Hydrogen atomic mass = $\text{<strong><em><em></em></strong> g/mole} \times 10 \text{ hydrogen atoms} = \text{<strong><em></em></em></strong> g/mole}$
  • Total molecular mass of \text{CuSO}4 \cdot 5\text{H}2\text{O} = \text{_ g/mole}

Mole

• Definition: A fundamental unit in stoichiometry used to quantify the amount of a substance.
• Relationship to Atomic Mass: One mole of a substance is equivalent to its atomic mass (or molecular mass) expressed in grams.
• Units: Grams/mole ($\text{g/mol}$).
• Calculation: To determine the number of moles, one needs to know the amount of the compound in grams and its molecular mass.
• Examples:
  • $\text{H}<em>2\text{O}$ (water) has an atomic mass of $18$, so $18 \text{ g}$ of $\text{H}</em>2\text{O}$ constitutes one mole of water.
  • $\text{CH}<em>4$ (methane) has an atomic mass of $16$, so $16 \text{ g}$ of $\text{CH}</em>4$ constitutes one mole of methane.
  • $\text{C}<em>6\text{H}</em>{12}\text{O}<em>6$ (glucose) has an atomic mass of $180$, so $180 \text{ g}$ of $\text{C}</em>6\text{H}<em>{12}\text{O}</em>6$ constitutes one mole of glucose.
• Calculation Examples (Student Task, blanks to be filled):
  • $0.5 \text{ mole of H}<em>2\text{O} = \text{<strong><em></em></em></strong> g of H}_2\text{O}$
  • $0.75 \text{ mole of CH}<em>4 = \text{<strong><em></em></em></strong> g of CH}_4$
  • $0.3 \text{ mole of C}<em>6\text{H}</em>{12}\text{O}<em>6 = \text{<strong><em></em></em></strong> g of C}<em>6\text{H}</em>{12}\text{O}_6$

Molarity

• Definition: Represents the concentration of a solution, defined as the number of moles of solute dissolved per liter of solution.
• Formula: $Molarity (M) = \frac{\text{moles}}{\text{liter}} \text{ (mol/L)}$.
• Experiment Components:
  • Solute: Copper Sulfate Pentahydrate ($\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$)
  • Solvent: Water
• Calculations (Student Task, blanks to be filled):
  • 1 M solution in 1 L: Using the molecular mass from the "Molecular Mass" section, calculate how many grams of $\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$ are required to create a $1 \text{ M}$ solution in $1 \text{ liter}$. Result: \text{_}
  • 1 M solution in 20 mL: Since experiments often require smaller volumes, calculate the grams needed to prepare a $1 \text{ M}$ solution of $\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$ in $20 \text{ mL}$. This can be done by first calculating the moles needed ($M = \text{moles/L}$) and then converting moles to grams. Alternatively, use the rearranged equation:
    $Molarity \left( \frac{\text{mol}}{\text{L}} \right) = \frac{\text{mass (g)}}{\text{molecular mass } \left( \frac{\text{g}}{\text{mol}} \right)} \times \frac{1}{\text{volume (L)}}$
    Solving for mass (g): $\text{mass (g)} = Molarity \left( \frac{\text{mol}}{\text{L}} \right) \times \text{molecular mass } \left( \frac{\text{g}}{\text{mol}} \right) \times \text{volume (L)}$
    Result for $20 \text{ mL}$ solution: \text{_}
• Verification: Always double-check calculations to ensure the correct amount of copper sulfate is weighed out.

Concepts: Absorption of Light and Standard Curves

Spectrophotometric Assays

• These assays are fundamentally based on the direct absorption of light by a substance.
• Light absorption can be quantitatively measured using a spectrophotometer or a colorimeter.

Calibration and Blanking

• Importance: Prior to use, a spectrophotometer or colorimeter must be properly calibrated, "zeroed," or "blanked." This step is critical because it accounts for and negates any absorbance that might originate from the solvent (in this experiment, water).
• Consequence of Omission: Failing to blank or zero the instrument will result in meaningless data.

Beer-Lambert Law

• This law describes the relationship between absorbance, concentration, and the molar extinction coefficient of a molecule.
• Formula: $A = \log<em>{10} \frac{I</em>o}{I} = \epsilon cl$
  • $A$: Absorbance of the solution (unitless).
  • $I_o$: Intensity of the incident (incoming) light.
  • $I$: Intensity of the transmitted light.
  • $\epsilon$ (Epsilon): Molar extinction coefficient (a constant specific to a molecule at a given wavelength, a measure of how strongly the substance absorbs light).
  • $c$: Concentration of the absorbing solute (in Molarity, mol/L).
  • $l$: Pathlength (the distance the light travels through the solution, typically the width of the cuvette, in cm).
• Utility & Limitations:
  • While not required to be used for calculations in this lab, understanding it is useful.
  • If the molar extinction coefficient ($\epsilon$) for a specific molecule at a certain wavelength is known, Beer's law allows for the calculation of the substance's concentration in a solution after measuring its absorbance with a spectrophotometer.
  • Conditions for Validity: Beer's law is only applicable if:
    1. The relationship between absorbance and concentration is linear.
    2. You are analyzing a pure substance with a single, known value of $\epsilon$.

Standard Curves

• Purpose: Often, the molar extinction coefficient ($\epsilon$) is unknown, or the linearity of the absorbance-concentration relationship is uncertain. In such cases, spectrophotometry/colorimetry can still be used for quantitative measurements by constructing a standard curve.
• Definition: A standard curve is a graph that illustrates the relationship between the known amount (or concentration) of a specific compound in a series of solutions and the absorbance readings of those solutions.
• Creation & Application:
  1. A series of solutions with known concentrations (standards) are prepared.
  2. Their absorbances are measured using the same instrument and wavelength.
  3. A graph (standard curve) is plotted, typically with concentration on the x-axis and absorbance on the y-axis.
  4. Once created, this curve can be used to determine the concentration of the compound of interest in an unknown solution.
  5. This is done by measuring the unknown solution's absorbance (using the identical instrument and wavelength as the standards) and then either plotting it on the graph or using the equation of the linear region of the standard curve to derive its concentration.

Procedure

Materials

• Shared Equipment:
  • Electronic balance
  • LabQuest 2
  • 3 Unknown solutions (choose from A, B, C, D, E) - $5 \text{ mL}$ each
  • Disposable gloves
  • Copper Sulfate ($\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$)
  • Stir plate with heat
  • Safety glasses
• Student Group Kits:
  • Colorimeter
  • 1 Cuvette
  • Kimwipes
  • Scoopula
  • Weigh paper
  • $100 \text{ mL}$ beaker filled with DI water
  • $100 \text{ mL}$ beaker for liquid waste
  • $50 \text{ mL}$ beaker to dissolve $\text{CuSO}_4$
  • Stir bar
  • $25 \text{ mL}$ graduated cylinder
  • $5 \text{ mL}$ pipets (2) (serological or measuring, review usage)
  • Green pipet pump
  • Disposable pipets (11)
  • Test tube rack
  • Test tubes (11)

Prepare Standard Solution

1. Put on disposable gloves and safety glasses.
2. After calculating the required gram amount of $\text{CuSO}<em>4 \cdot 5\text{H}</em>2\text{O}$, weigh this quantity using a weigh boat and transfer it to an empty $50 \text{ mL}$ beaker. Turn off the balance after use.
3. Add approximately $18 \text{ mL}$ of DI water to the $\text{CuSO}_4$ in the beaker. DO NOT ADD $20 \text{ mL}$ INITIALLY. The final total volume of the solution, including the solute, must be exactly $20 \text{ mL}$. The solution will be brought to its final volume after the solute is completely dissolved.
4. Place a magnetic stir bar into the beaker. Position the beaker on a stir plate, turn on the heat to approximately $50^\circ \text{C}$, and stir until the $\text{CuSO}_4$ is completely dissolved. This process will likely take $5 \text{ to } 10 \text{ minutes}$. Monitor the beaker to ensure it does not get too hot; it should feel comfortably warm to the touch.
   • While waiting: Begin preparing the test tubes for the standard curve dilutions (steps 1 and 2 under "Standard Curve Make Dilutions of Standard"), and add the appropriate amount of water to each tube.
5. Once the $\text{CuSO}<em>4$ has fully dissolved, transfer the solution to a $25 \text{ mL}$ graduated cylinder. Carefully add DI $\text{H}</em>2\text{O}$ to reach a final volume of exactly $20 \text{ mL}$. Pour this solution back into the $50 \text{ mL}$ beaker used for dissolving (to prevent accidental spillage from a graduated cylinder).
6. Turn off the stir plate and heat when finished.

Standard Curve: Make Dilutions of Standard

1. Use the $V<em>1M</em>1 = V<em>2M</em>2$ equation to calculate the required volumes of $\text{CuSO}<em>4$ (Solution 1 and $\text{M}</em>1$ being the $1 \text{ M}$ standard solution you prepared) and deionized water ($\text{dH}_2 ext{O}$) for each of the eight standard solutions. The equation components are:
   • $V_1$: Volume of solution 1
   • $M_1$: Concentration of solution 1
   • $V_2$: Volume of solution 2 (total final volume for the dilution)
   • $M_2$: Concentration of solution 2 (desired final concentration for the dilution)
   • Note: The volumes for tubes 1 (0 M) and 8 (1 M) have been provided in Table 1.
2. Label eight clean, dry test tubes from 1 to 8.
3. Using a graduated pipet, prepare each of the eight standard solutions according to the volumes determined from the $V<em>1M</em>1 = V<em>2M</em>2$ calculation and Table 1. Ensure precision in pipetting for accurate results. Each solution should have a total volume of $5 \text{ mL}$. Use separate pipets for the $\text{CuSO}_4$ solution and the water.
4. Thoroughly mix each solution using a clean disposable (plastic) pipet.
5. Save each disposable pipet by placing it in its corresponding test tube.

TABLE 1: Standard Curve Volumes (Student Task, blanks to be filled)

Measuring Absorbance using Vernier LabQuest/Colorimeter

Vernier and Colorimeter Set-Up

1. Plug the Colorimeter into the Vernier LabQuest, and turn on the LabQuest (power switch on the top right side). The LabQuest should automatically detect the Colorimeter connection.
2. Allow approximately $90 \text{ seconds}$ for the Colorimeter to warm up, if necessary.
3. Set the Colorimeter to $635 \text{ nm}$ by pressing the arrow buttons on the device.
4. Fill a cuvette approximately $3/4$ full with deionized water (this acts as your blank). Wipe away any fingerprints from the cuvette's exterior using a Kimwipe tissue.
5. Insert the cuvette into the Colorimeter, ensuring the unlined sides of the cuvette align with the yellow arrow. Close the lid of the Colorimeter.
6. Press and hold the "CAL" button until the red light stops blinking. This calibrates (blanks) the instrument. Do not recalibrate unless instructed by your instructor.
7. Remove the cuvette and empty its contents into the $100 \text{ mL}$ waste beaker.
8. On the right side of the LabQuest monitor, push "Mode" and set it to "Time Based." Select "s" (seconds) from the pull-down menu and set the duration for $10 \text{ sec}$. Push "OK."

Collect Absorbance Data for the Eight Standard Solutions

1. Using the solution from Test Tube 1 ($0 \text{ M}$), fill a new cuvette approximately $3/4$ full. Wipe the outside with a tissue and place it into the Colorimeter. Close the lid.
2. Record the displayed absorbance value in Table 2 for Test Tube 1.
3. Remove the cuvette and empty its contents into the waste beaker.
4. Repeat steps 1-3 for Test Tubes 2 through 8, using a new cuvette for each measurement.

Collect Absorbance Data for Unknown Solutions

1. Obtain about $5 \text{ mL}$ of each of the three unknown solutions you will test (choose 3 from the 5 available: A, B, C, D, or E). Place each unknown in a separate test tube.
2. Determine the absorbance of each unknown using the Colorimeter. Follow the same procedure as for the standard solutions, ensuring you use a new cuvette for each unknown.
3. Record the displayed absorbance value for each unknown in Table 3. Make sure to indicate which specific unknowns (e.g., A, B, C) you tested.

TABLE 2: Standard Curve Results (Student Task, blanks to be filled)

Results

Standard Curve

• Transfer the calculated concentration for each standard solution from Table 1 to Table 2.
• Record the corresponding Absorbance readings measured at $635 \text{ nm}$ in Table 2.

Graphing Questions

1. Question 1: For the data collected in Table 2, what type of graph should be used to chart the data (scatter, line, bar, histogram, pie)?
   • Answer: \text{__}
2. Question 2: What should be plotted on the x-axis (include units, if relevant)?
   • Answer: $\text{<strong><em><em></em></strong><strong><em></em></em></strong><strong><em></strong></em>}$
3. Question 3: What should be plotted on the y-axis (include units, if relevant)?
   • Answer: $\text{<strong><em><em></em></strong><strong><em></em></em></strong><strong><em></strong></em>}$

• Task: Use the provided graph paper on the last page of the report to plot the data from Table 2 and create a Standard Curve. Follow the graphing guidelines discussed in the introductory lab videos. This graph must be submitted as part of your report.

Unknowns

• Determine the concentration of each unknown solution by plotting its measured absorbance onto the created Standard Curve Graph.
• Record these determined concentrations in Table 3.

TABLE 3: Unknown Solutions Results (Student Task, blanks to be filled)

Conclusions

1. Can spectrophotometry be effectively used to determine the concentration of "colorless" solutes, such as common salt or sugar? Provide a full explanation for your answer.
2. Explain the importance of using standards or developing a comprehensive standard curve in spectrophotometry and how it aids in quantitative analysis.
3. Elaborate on the critical reason behind "reading" (blanking) a cuvette containing only the solvent, with no solute present, during the Colorimeter/spectrophotometer setup.