INVESTIGATING IRON THIOCYANATE

Investigating Iron Thiocyanate
Revised: 4/28/15
INVESTIGATING IRON THIOCYANATE
Adapted from "Chemistry with Computers"
Vernier Software, Portland OR, 1997
&
from “Journal of Chemical Education”
Vol. 75, No. 12, December 1998
http://jchemed.chem.wisc.edu/Journal/Issues/1998/Dec/abs1628.html
REPORT INSTRUCTIONS
Create the following prelab sections on a new page in Experiment #2’s folder in your electronic
lab notebook (ELN):
1) Objectives. In the Objectives section, clearly state the main chemical event, the central theory
or theories used, and the technique(s) to be employed in the experiment.
2) Equipment and chemical tables. Follow the guidelines given with the last experiment.
3) Procedures and Observations. Follow the guidelines given with the last experiment.
4) Data Tables. A data table is given within the procedures. This data table should be created
and completely filled out in the lab notebook. In future labs, it is your responsibility to
determine what data should be included in a table or tables and create those tables yourself.
Calculations / Postlab Questions / Error Analysis. Follow the guidelines given with the last
experiment, creating a new page in the same experiment folder.
INTRODUCTION
When Fe3+ and SCN− are combined, an equilibrium is established between the two ions and their
product, the FeSCN2+ ion, and an equilibrium constant can be written (1).
Fe3+(aq) + SCN-(aq) ! FeNCS2+(aq)
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Kc = [FeNCS2+]eq (1)
[Fe3+]eq[SCN-]eq
Investigating Iron Thiocyanate
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You will calculate the equilibrium constant of the reaction, Kc, by finding the equilibrium
concentrations of the reactants and product. Determining initial concentrations is typically
straightforward – the mass of solute and volume of solution or the concentration of stock
solution and volume of dilute solution is known. Finding equilibrium concentrations requires a
more complex approach.
Equilibrium Product Concentration [FeNCS2+]eq.
Both reactants (Fe3+ and SCN-) create clear, colorless aqueous solutions when dissolved.
However, the dissolution of the product (FeNCS2+) results in a red-orange aqueous solution.
Like all dilute colored solutions, FeNCS2+(aq) absorbs visible radiation and obeys Beer’s Law (A
= εCl ). In this experiment A Visible to Near InfraRed Spectrometer (VIS-NIR) will be used to
measure absorbance of FeNCS2+. The absorbance will be correlated with concentration using
Beer’s Law. The observed red-orange color of an aqueous FeNCS2+ solution results from the
absorption of photons with the energy of the complementary color, blue/green. Therefore, λmax,
the wavelength with the highest absorbance, should be around 480 nm. As typical for many
experiments utilizing visible spectroscopy, the absorbance values in this experiment will be
recorded at λmax. (Note: λmax does not change during the experiment. The absorbance of all
trials must be measured at the same wavelength.)
By comparing the absorbance of each trial solution, Aeq, to the absorbance of the standard
solution, Astd, [FeNCS2+]eq can be determined. Two Beer’s Law equations can be written, one
for the trial (Aeq = ε[FeNCS2+]eql ), one for the standard (Astd = ε[FeNCS2+]stdl ). Set these two
equations equal to each other to derive equation (2).
[FeNCS2+]eq =
Aeq
Astd
2
X [FeNCS +]std
(2)
The concentration of FeNCS2+ ([FeNCS2+]std ) in the standard can be found by making an
assumption based on Le Chatelier’s principle. The standard solution contains a very large
concentration of Fe3+ ([Fe3+]i) and a small initial concentration of SCN– ([SCN ]i). (The [Fe3+]
−
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in the standard is 100 times more than [Fe3+] in the trials.) The high [Fe3+] forces the reaction
far to the right, using up nearly 100% of the SCN– ions. Therefore, for every one mole of SCN–
reacted, one mole of FeNCS2+ is produced and [FeNCS2+]std can be assumed equal to [SCN–]i,
the initial concentration of SCN in the standard solution.
−
Equilibrium Reactant Concentration [Fe3+]eq & [SCN–]eq.
Knowing [FeNCS2+]eq allows you to determine the concentrations of the other two ions at
equilibrium. For each mole of FeNCS2+ ions produced, one less mole of Fe3+ ions will be found
in the solution (see the 1:1 ratio of coefficients in the equation on the previous page). The [Fe3+]
can be determined by:
[Fe3+]eq = [Fe3+]i – [FeNCS2+]eq
(3)
Because one mole of SCN- is used up for each mole of FeNCS2+ ions produced, [SCN–]eq can be
determined by:
[SCN–]eq = [SCN–]i – [FeNCS2+]eq
(4)
Knowing the values of [Fe3+]eq, [SCN–]eq, and [FeNCS2+]eq, the value of Kc, the equilibrium
constant, can be calculated.
The thiocyanate ion acts as an isothiocyanate ligand to Fe3+, in other words, the iron binds to the
nitrogen atom of the ligand not the sulfur atom. This is the reason that the order of SCN is
changed to NCS for the product FeNCS2+ not FeSCN2+.
SAFETY PRECAUTIONS
Safety goggles and aprons must be worn at all times. Iron(III) nitrate (Fe(NO3)3) is a strong
oxidizer; skin and tissue irritant. Nitric acid (HNO3) is corrosive and can cause burns.
Potassium thiocyanate (KSCN) is moderately toxic by ingestion; emits toxic fumes of cyanide if
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strongly heated or in contact with concentrated acids. Always wash hands frequently while in
lab.
Before starting the experiment, the TA will asks one of you to do a quick demonstration on:
1) How to properly handle a cuvette (how to wipe it down, what to wipe it down with, how to
hold it)
2) How to use a volumetric flask
3) How to use a Mohr (graduated) pipet
Make sure you watch the videos on the course website and read the documents to prepare. These
demonstrations will be done every week. Everyone will have presented at least one topic by the
end of the quarter. The demonstrations should be short (>1 min) and will be graded.
PROCEDURE
Part A. Solution Preparation.
1. Work in pairs.
2. Place ~0.8 g Fe(NO3)3·9 H2O into a 10 mL volumetric flask and dissolve to the volumetric
line with 0.1 M HNO3(aq). (Make sure you prerinse the volumetric flask with the 0.1 M
HNO3(aq) solution before making the solution. Also, 9 waters of hydration (·9H2O) are
present in this chemical, and must be included in the formula weight calculation.) Record the
exact mass of iron nitrate used and show the calculation for the solution concentration in
your observations.
3. Before lab, create a plan for the dilution of the Fe(NO3)3(aq) solution created above with DI
water (this is different from tap water, ask your TA if you do not know where this is located).
This plan should be clearly outlined in your ELN. Create enough 0.0020 M Fe(NO3)3(aq) for
step 5 and a blank cuvette sample (You will lose points if you make too much solution).
Notes:
•
Do not try to make the solution’s concentration exactly 0.0020 M – just get close.
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Investigating Iron Thiocyanate
•
Revised: 4/28/15
A 1 mL Mohr (graduated) pipet is available, so volumes less than 1 mL can be delivered
with reasonable accuracy (not as accurate as a regular volumetric pipet, though). Pay
attention to the numbering on the Mohr pipet – the first marking is for zero, so think
about how you will deliver the desired volume of solution and be sure to express exactly
what you did in the procedures.
•
Make sure you record the procedure actually followed in lab if it deviates from the plan
you created.
4. The stockroom has prepared an aqueous KSCN solution for you by placing ~0.400 g of
KSCN into a 2.00 L volumetric flask and diluting to volume with deionized H2O. Record the
exact mass shown in the display case next to the stockroom window and show the calculation
for the solution concentration in your observations.
Part B. Absorbance Measurements.
5. Label 4 test tubes #1-4.
•
Trials (Test Tubes #1-3): Pipet 5.0 mL of 0.0020 M Fe(NO3)3(aq) (created in step 3) into
each test tube. Pour no more than 15 mL of the KSCN solution into another clean, dry
50-mL beaker. Pipet 2, 3, and 4 mL of this solution into Test Tubes #1-3, respectively.
Obtain about 15 mL of deionized water in a 100-mL beaker. Then pipet 3, 2, and 1 mL
of deionized water into Test Tubes #1-3, respectively, to bring the total volume of each
test tube to 10 mL. Mix each solution thoroughly with a stirring rod. Be sure to clean
and dry the stirring rod after each mixing.
•
Standard (Test Tube #4): Prepare a solution of FeNCS2+ by pipetting 9 mL of the more
concentrated Fe(NO3)3(aq) (created in step 2) into a test tube #4. Pipet 1 mL of the
KSCN solution into the same test tube. Mix the solution thoroughly with a stirring rod.
6. Copy the following table into your ELN and summarize the volumes and concentrations
used, as well as the absorbances measured in the table below. Take note of how the table is
constructed - headings contain a thorough description of the data being collected, including
the units where appropriate.
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Investigating Iron Thiocyanate
Test Tube
Number
Volume of
Fe(NO3)3 (mL)
Revised: 4/28/15
Concentration of
Fe(NO3)3 (M)
Volume of
Volume of H2O
_____ M KSCN
(mL)
(mL)
Absorbance
(λmax = ____nm)
1
2
3
4
7. Measure and record the temperature of one of the solutions.
8. READ VISIBLE SPECTROSCOPY BEFORE COMING TO LAB. Obtain a VIS-NIR
spectrometer from the stockroom. Use the USB cable to connect the VISNIR Spectrometer to the LabQuest 2.
9. Then calibrate the spectrometer by clicking
. The calibration dialog box
will display the message: “Waiting….seconds for lamp to warm up.” (The
minimum warm up time is one minute.) Note: For best results, allow the spectrometer to
warm up for at least three minutes. Insert the blank cuvette (filled with 0.0020 M
Fe(NO3)3) in the sample compartment. This step will remove any background absorbance
from the solvent (which is water). Click Finish Calibration and then OK.
10. Remove the blank cuvette from the spectrometer. Pour out the Fe(NO3)3 in the waste bottles
in the hood, prerinse the cuvette with solution #1 to remove any impurities (which could be
excess 0.0020 M Fe(NO3)3 solution), and then fill it 3/4 full with solution #1. Wipe the
outside of the cuvette with a kim wipe, place in the cuvette holder of the Spectrometer. Click
once the data collection is complete. Examine the graph and note the wavelength where
the maximum absorbance occurs (λmax). Remove the "rainbow" background spectrum by
double clicking the rainbow background. Click
and store lastest run.
11. Repeat this process with solution #2, #3, and the standard solution. Record the absorbance
values for each of the four solutions in a data table. To see all the spectra at once, click on
“run 4” and change it to “All runs.” Don’t forget to e-mail this data to your ELN. Remember
that any plot must be scaled appropriately – therefore, the spectrum of the standard should
not be combined with the spectra for the trials.
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12. Discard solutions in the waste bottles in the hood. Before leaving lab, copy another pair’s
data table for this experiment. Make sure that you also record their masses and volumes
from Part A, step 2 and 3 so you can double check the concentration of Fe(NO3)3 reported in
their data table.
Make sure to clear your email address and password of the LabQuest2 so others can’t access
your email account. Shutdown the LabQuest2 and not simply put it to sleep. To shutdown the
LabQuest2: press the home key, select System ! Shut Down ! OK.
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CALCULATIONS
Calculations should be performed on the data you collected and the data that the other pair
collected. (Therefore, each trial and the standard have two sets of data.) It should be clear to
the TA which data is yours.
1. For all trials (including those trials from the other pair of students), calculate the initial
concentrations of the reactants, [Fe3+]i and [SCN–]i. (Note: The initial concentrations must be
calculated from the concentrations found in the procedure. A dilution occurs when Fe3+(aq),
SCN-(aq), and H2O(l) are combined.) Show one sample calculation for each reactant.
2. Show the derivation of equation (2).
3. Show the derivation of equations (3) and (4) by creating an ICE chart (Initial Change
Equilibrium).
4. Calculate [FeNCS2+]std.
5. For all trials (yours and the other pair’s), calculate the equilibrium concentrations of the
reactants and products, [Fe3+]eq, [SCN–]eq, and [FeNCS2+]eq. Show one sample calculation for
each reactant and the product.
6. For all trials (yours and the other pair’s), calculate Kc. Provide the Kc expression and one
sample calculation.
ERROR ANALYSIS
1. Perform a Q-test on either the lowest or highest K value (choose the one most likely to be an
outlier). If the data point fails the test, throw it out. Then, determine an average value for Kc
and the estimated standard deviation.
2. What modifications could be made to the procedure to better account for random
(indeterminate) errors?
3. List three potential systematic (instrumental, methodological, or personal) errors that could
be made in this experiment. (Note: Be specific, systematic errors are in the details. For
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example, losing your solution because you knocked over the cuvette is not a systematic error
– it’s a gross one.)
4. a. Why did you blank with the 0.0020 M Fe(NO3)3 solution as opposed to just DI water?
What does it do to your spectrum?
b. What chemicals create the absorbance you observe from the standard solution (Test Tube
#4)?
c. How can you find the absorption of only the FeSCN in the standard solution?
5. Did any gross errors occur? Did you mess up? Did the equipment or instrumentation fail?
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