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Chemistry 223 Spring 2004
Spectrophotometric Determination of an
Equilibrium Constant
Introduction:
In this experiment, you will study the reaction between aqueous iron (III) nitrate,
Fe(NO 3
3 , and potassium thiocyanate, KSCN. They react to produce the blood-red
complex [Fe(SCN)
2+
] according to the following net ionic equation:
Fe
3+
[FeSCN
2+
]
The equilibrium constant expression may be expressed as:
K =
[Fe ][SCN ]
[FeSCN ]
3
2
You will prepare a series of standard solutions that contain known concentrations of
[FeSCN
2+
]. Because the red solutions absorb blue light very well, the blue LED setting
on the colorimeter is used and you will determine their absorbances at 470
nanometers (blue). The concentrations and absorbance values will be used to
construct a Beer’s Law calibration curve for [FeSCN
2+
].
Part I: The Beer’s Law Calibration Curve
One question you might have is “How can I know the equilibrium concentration of
FeSCN
2+
before I make a Beer’s Law calibration curve?” The “trick” lies in utilizing
Le Chatelier’s principle as follows: A very large concentration of Fe
3+
will be added to
a small initial concentration of SCN
(hereafter referred to as [SCN
] i
). The [Fe
3+
] in
the standard solution is 100 times larger than the [SCN
] i
in the equilibrium
mixtures. According to Le Chatelier’s principle , this high concentration forces the
reaction far to the right, using up essentially all of the SCN
ions. According to the
balanced equation, for every one mole of SCN
reacted, one mole of FeSCN
2+
is
produced. Thus equilibrium moles FeSCN
2+
eq
are assumed to be equal to
initial moles SCN
-
i
for the Beer’s Law standard solutions only.
To reiterate: equilibrium moles FeSCN
2+
eq =
initial moles SCN
-
i
For Beer’s Law standard solutions only
The equilibrium concentration of FeSCN
2+
is simply this number of moles divided by
the total volume after mixing.
Part II: Determining K eq
In the second part of the experiment, various combinations of Fe(NO 3
3 and KSCN will
be combined. The amount of product formed at equilibrium, [FeSCN
2+
], will
be determined from the calibration graph prepared earlier. Knowing the
[FeSCN
2+
] eq
allows you to determine the concentrations of the other two ions at
equilibrium. For each mole of FeSCN
2+
ions produced, one less mole of Fe
ions will
be found in the solution (see the 1:1 ratio of coefficients in the balanced chemical
equation). The [Fe
3+
] eq
can be determined by:
[Fe
3+
] eq
= [Fe
3+
] i
2+
] eq
In words: The equilibrium concentration of Fe
3+
is the initial concentration of Fe
3+
(after dilution) minus the amount that reacted to form FeSCN
2+
(determined by
spectroscopy).
The [FeSCN
2+ ] eq
is determined by using spectroscopy and the Beer’s Law calibration
curve.
Because one mole of SCN
is used up for each mole of FeSCN
2+
ions produced,
[SCN
]
eq
can be determined by:
[SCN
] eq
= [SCN
] i
2+
] eq
Knowing the values of [Fe
3+ ] eq
, [SCN
, and [FeSCN
2+ ] eq
, you can then calculate
the value of K c
, the equilibrium constant. When these concentrations are substituted
into the equation for the equilibrium constant, numerical values for the equilibrium
constant are determined. An average value for the constant will then be determined.
Purpose :
The purpose of this experiment is to determine a value for the equilibrium constant
for the reaction between iron (III) nitrate and potassium thiocyanate.
MATERIALS
computer 0.0020 M KSCN
Vernier computer interface 0.0020 M Fe(NO 3
3 (in 1.0 M HNO 3
Logger Pro 0.200 M Fe(NO 3
3 (in 1.0 M HNO 3
Vernier Colorimeter pipets and bulbs
1 plastic cuvette buret
five 20 150 mm test tubes 25-mL volumetric flasks
thermometer tissues (preferably lint-free)
Safety and Disposal Guidelines :
Goggles (and apron) should always be worn in the lab.
The Fe(NO 3
3 solutions are dissolved in 1.0 M HNO 3 , a strong acid. 0.050 M
HNO
3 is used as well. In case of eye contact, flush with water for 15 minutes
in an eye-wash. In case of skin contact, flush with water for 15 minutes.
Dispose of all solutions in the waste bottle in the hood. Do not flush solutions
down the drain.
Procedure :
Special Note on Data Collection : At this point in the year, you will be expected to
prepare well-organized tables to contain your data. Pay attention to detail and what
you have learned in the past about column headings (quantity and unit) as well as
proper attention to significant figures. In addition, a good chemist will always write
down qualitative observations with every step. If you have been lacking this in the
past, now is the time to start. Think ahead: If you do not wish to re-write your data
section in the final report, be neat and organized now, it will save you time later.
Also, since you are expected to keep good lab notebooks by this point in the year, a
poor notebook may result in a reduced score for the lab report.
PART I: BEER’S LAW CALIBRATION CURVE
1. Clean, dry, and label 4 small beakers to contain the necessary reagent solutions. 2. Obtain the necessary reagents and dispense about 30 mL of each into a small
beaker. DO NOT pipet directly from a reagent stock bottle.
3. NOTE : You may prepare the following standard solutions to be used by 2 pairs of
students at the same table. Divide up the work of preparing the solutions among
the 4 people at a table. However, each pair of students wil be performing the
experiment below (you are just sharing solutions!!)
a. Open the Colorimeter lid.
b. Holding the cuvette by the upper edges, place it in the cuvette slot of the
Colorimeter. Close the lid.
c. The newer colorimeters only have a CAL button. Press the < or > button on the
Colorimeter to select a wavelength of 470 nm (Blue) for this experiment. Press
and hold the CAL button until the red LED begins to flash. Then release the
CAL button. When the LED stops flashing, the calibration is complete. Proceed
directly to Step 7.
If your Colorimeter does not have a CAL button, continue with this step to
calibrate your Colorimeter.
First Calibration Point
d. Choose Calibrate CH1: Colorimeter (%T) from the “Experiment” menu and
then click.
e. Turn the wavelength knob on the Colorimeter to the “0% T” position.
f. Type “0” in the edit box.
g. When the displayed voltage reading for Reading 1 stabilizes, click.
Second Calibration Point
h. Turn the knob of the Colorimeter to the Blue LED position (470 nm).
i. Type “100” in the edit box.
j. When the displayed voltage reading for Reading 2 stabilizes, click , then
click.
8. You are now ready to collect absorbance data for the Beer’s Law standard
solutions.
a. Click to begin data collection.
b. Empty the water from the cuvette. Rinse it twice with ~1 mL portions of the
Test Tube 1 solution.
c. Wipe the outside of the cuvette with a tissue and then place the cuvette in the
Colorimeter. After closing the lid, wait for the absorbance value displayed in
the meter to stabilize. Then click , type the FeSCN
2+
concentration in
edit box, and press the ENTER key.
d. Discard the cuvette contents as directed by your teacher. Rinse the cuvette
twice with the second solution and fill the cuvette 3/4 full. Follow the Step-c
procedure to find the absorbance of this solution. Type the FeSCN
2+
concentration in the edit box and press ENTER.
e. Repeat the Step-d procedure to find the absorbance of the remaining standard
solutions.
f. Record the concentration and absorbance values in your notebook.
g. Determine the best fit line by clicking the Linear Fit button,.
h. Record the equation of the line and print a copy of the graph.
i. Dispose of all solutions as directed by your instructor.
PART II: DETERMINATION OF EQUILIBRIUM CONSTANT.
9. Clean and dry 5 large test tubes and label them 1-5. Use a buret or pipet to
measure the target volumes of the reactants listed below. Note that this set of
Note: If you
have an
older “box”
type
colorimeter,
you will
need to do
steps d. – j.
combinations uses the more dilute Fe(NO 3
3 solution. As usual, record
actual volumes in your data table.
Trial
Volume
0.0020 M KSCN (mL)
Volume
0.0020 M Fe(NO 3
3
(mL)
Volume
0.050 M HNO
3 (mL)
10. Re-calibrate the Colorimeter, but this time use 0.0020 M Fe(NO 3
3
11. Measure the absorbance of each equilibrium solution following the same steps as
in Part I. NOTE: You are not making a graph this time. Simply record the
absorbance values in your notebook for further analysis.
Analysis
- Using you’re equation of the Beer’s
Law calibration curve, determine the concentration of [Fe(SCN)
2+
] eq for each of the
trials.
- From the concentration of
[Fe(SCN)
2+
] eq produced and the original concentrations of the reactants, construct
tables to determine the equilibrium concentrations of all species. NOTE: The
initial concentrations should be calculated based on the fact that they are diluted
(what they would be right after mixing). Since each solution is mixed with the
other solution and HNO 3 (aq), they are no longer at 0.0020 M.
- Use these values to calculate the
equilibrium constant (K c ) for each trial.
- Calculate your standard deviation
(STD) for the K c values, and your relative (%) STD. Comment on your level of
precision. What sources of error might there have been?
Lab Report
Prepare a lab report in the standard scientific manner. Your report must be typed
and grammatically correct. You may hand-write equations instead of typing them if
you prefer. Consider the following points when creating your report, they
are the basis for which your grade will be assigned:
Abstract: a 1 paragraph overall summary of the experiment. NOT an
explanation of the procedure! This paragraph would be similar to how you
would explain the overall purpose and theory of the lab to another person.
Introduction: describe and explain Beer’s Law, show equations, discuss
equations used. Describe a calibration curve, discuss what a calibration curve
is used for and why they are useful, describe and discuss chemical
