Spectrophotometry Use in AnalysisPart 1: Calibration Curve Data:
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Part 2: Hydrate Data:
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1. Calculate the concentration of the stock solution in mol/L.
2. Using the dilution formula, C1 V1 = C2 V2 , calculate the concentration of copper(II) ion in each of the
four different stage solutions:
a. Stage 1
b. Stage 2
c. Stage 3
d. Stage 4
3. Graph absorbance a
s a unction of concentration of copper(II) ion on the Chemistry Lab Resources
you can get the slope and y intercept.
a. Does th is plot appear to be a straight line?
b. Does the data appear to lie along a straight line that extends back close to (0,0)?
c. According to your results, is Beer’s law valid for this case?
4. From the best-fit line of your data:
Slope (2 decimals)
Y-intercept (2 decimals) =
5. Graph the data on paper and include the best-fit line. Include your graph with your report.
6. Using the slope and y-intercept of your best-fit line, the absorbance of your hydrate, along with
y = mx + b, determine the concentration of copper(II) ion in your dilute unknown hydrate solution.
This was the solution that was in the cuvette. Hint solve for “x” and keep three significant figures.
7. Calculate the concentration of the original copper(!!) ion solution in the volumetric flask. Hint solve
for C1 when using the dilution formula, C1 V1 = C:N2. C2 is the copper(II) ion concentration calculated
in the previous step.
8. Now you need to calculate the concentration of copper(II) nitrate in the volumetric flask and you need
to calculate it using the concentration of copper(II) ion. Use the concentration from the previous step
and the mole ratio between copper(II) ion and copper(II) nitrate. Hint First write out the chemical
formula for copper(II) nitrate and then do a stoichiometry conversion.
9. Calculate the moles of copper(II) nitrate present in the volumetric flask.
10. Calculate the mass of copper(II) nitrate in the 100.00 ml volumetric flask.
11. Now that you have the mass of copper(II) nitrate, subtract it from the mass of hydrate you originally
weighed to get the mass of water in your hydrate.
12. Calculate the amount of water in your hydrate in moles.
13. Divide the moles of water (from #12) by the moles of copper(II) nitrate (from #9). This is your
experimental value for the number of waters of hydration.
14. Go to the Chemistry Lab Resources website, CHEM 1090 Labs/Experiment 8, enter your data,
and record the results for the number of waters of hydration attached to the copper(II) nitrate
formula unit (#13). Compare it to your calculated value in #13 above. If they are significantly
different then you need to fix any mistakes:
15. Take your value calculated in #13 and round it to the nearest whole number.
16. Write the chemical formula of your hydrate using the rounded whole number from #15.
CC”‘- C,J13J’i, . . . .. 1-h 0~
17. Using an SCC iPad, open the Google Drive app, and go to the “Reference” folder. Or, log in on an
computer and access the folder named either “ES-CHEM”, “LN-CHEM”, or “WB-CHEM” in the
“Z” drive. Or, you can find the file linked in the instructions page for this experiment module in your
Canvas shell. Open the pdf file, “Physical Constants of Inorganic Compounds.” Find the metal
name section of your compound. Under the metal look for the anion by name. Examine these
hydrates and write down the chemical formula closest to your own.
18. Using your unrounded value calculated in #13 as the EV and the value from the chemical formula
you just looked up in #17 as the SV, calculate the percent error.