1. Measuring K for

2. The Reaction
But we only want
Make sure [Fe3+]o >> [SCN-]o

3. The Reaction
Fe3+ can be very insoluble because it can act as a Lewis Acid
Drive equilibrium left by adding HNO3(aq)

4. Measuring Equilibrium Concentrations
Deep red color
The complex absorbs in the blue/green/yellow, that is why it is red
We can monitor the concentration of the product using Beer’s Law

5. Beer’s Law Absorbance Path Length Molar Extinction Coefficient
“The stronger the brew, the less light that goes through”
Absorbance
Molar Extinction Coefficient
Path Length

6. Beer’s Law
If we know ε and l we can measure the concentration by measuring the absorbance
We can do this by calibrating the spectrometer, by measuring the absorbance at known concentrations of [FeSCN2+]
The slope of the calibration curve is ε.l

7. Calibrating the Spectrometer
But hang on how can we get known [FeSCN2+]eq to calibrate?
Add a LOT of Fe3+, this will drive the reaction to completion (Le Chatelier’s Principle)
Then [SCN-]o ≈ [FeSCN2+]eq

8. Calibrating the Spectrometer
Flask #
0.200 M Fe(NO3)3
(mL)
M KSCN
0.1 M HNO3 (mL) 1
5.00
0.00
20.00 2
15.00 3
4.00
16.00 4
3.00
17.00 5
2.00
18.00 6
1.00
19.00
Pipette the reactants into a 25mL volumetric flask, and fill up to the calibration mark on the flask with the 0.1M HNO3

9. Calibrating the Spectrometer
Flask #
0.200 M Fe(NO3)3
(mL)
0.0020M KSCN
0.1 M
HNO3
[Fe3+]o
[SCN-]o
[FeSCN2+]
Abs 1
5.00
0.00
20.00 2
15.00 3
4.00
16.00 4
3.00
17.00 5
2.00
18.00 6
1.00
19.00
Calculate the initial concentrations [Fe3+]o, [SCN-]o = [FeSCN2+]eq
Measure A for each sample