In the Laboratory
Colorimetric Determination of the Iron(III)-Thiocyanate Reaction Equilibrium Constant with Calibration and Equilibrium Solutions Prepared in a Cuvette by Sequential Additions of One Reagent to the Other Frazier Nyasulu* and Rebecca Barlag Department of Chemistry and Biochemistry, Ohio University, Athens, Ohio 45701, United States *
[email protected] Many general chemistry lab manuals (1-4) include the colorimetric determination of the equilibrium constant for the aqueous reaction Fe3þ SCN - h FeðSCNÞ2þ
ð1Þ
The Fe(SCN)2þ complex has a λmax of ∼460 nm. The equilibrium constant, K, is K ¼
½FeðSCNÞ2þ ½Fe3þ ½SCN -
where A is the absorbance, Cs = [SCN-] þ [FeSCN2þ], Cf = [Fe3þ] þ [FeSCN2þ], b is the path length, and ε is the molar absorptivity. A plot of A/CsCf versus A(Cf þ Cs)/CfCs is linear with a slope of -K. We consider the colorimetric determination of Fe(SCN)2þ in equilibrium solutions to be more instructive in the general chemistry lab because it reinforces important concepts that are discussed in the lecture. Among these are (i)
Calculation of amount of reagent and product from molarity and volume added. (ii) Consideration of reaction stoichiometry to identify and use a limiting reagent to calculate concentration of a product. (iii) Use of the calibration plot equation and the absorbance to calculate the concentration of FeSCN2þ in the equilibrium solutions. (iv) Use (i) to (iii) to calculate amounts of Fe3þ added, SCNadded, FeSCN2þ formed, Fe3þ unreacted, and SCN- unreacted and the equilibrium concentrations of Fe3þ and SCN-.These concepts are “must learn” items in general chemistry and the opportunity to work with these in a lab can be very helpful. A procedural improvement to the calibration approach is described. The differences between a standard procedure and the procedure described herein are highlighted in Table 1.
ð2Þ
and can be determined in general chemistry labs in one of two methods. In the first method, the most common approach, the colorimetric determination of Fe(SCN)2þ in equilibrium solutions is based on a Fe(SCN)2þ absorbance versus concentration calibration plot (1-4). The absorbance-concentration calibration plot is constructed with [Fe3þ] . [SCN-], a condition that allows the calculation of the Fe(SCN)2þ concentrations. In the second method (4, 5), a calibration graph is not required because the equilibrium constant is determined from eq 3 A AðCf þ Cs ÞK = εbK Cf Cs Cf Cs
ð3Þ
Table 1. A Comparison between Standard and the New Procedure To Determine the Equilibrium Constant of FeSCN2þ Standard Procedure
New Procedure
Solutions are made using pipets or burets. The volumes are larger than volumes needed for the measurements.
Solutions are prepared using a 0.10-1.00 mL autopipettor and solution volumes are low. All the solution prepared is used in the measurement.
Individual calibration solutions are prepared and transferred into the cuvette for measurement.
Calibration solutions are prepared by sequential additions of 0.10 mL of a SCN- solution to 4.00 mL of an Fe3þ solution in a cuvette.
The volume of each of the calibration solutions is fixed.
The total volume of the calibration solution is varied.
Cuvette is loaded with new solution for each measurement.
Additions are made to the cuvette already placed in the colorimeter. Cuvette is not handled between measurements.
Calibration is performed typically with 5 standards.
Calibration is performed with 10 or more standards.
Much of the solution prepared to determine the equilibrium constant is not used in the measurement. Typically four solutions are prepared and analyzed.
To determine the equilibrium constant, sequential additions of 0.50 mL of an SCN- solution to 4.00 mL of a Fe3þ solution are made. Measurements with sequential addition of 0.50 mL of Fe3þ to 4.00 mL of SCN- are also performed. Typically 10 or more equilibrium solutions are analyzed.
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r 2011 American Chemical Society and Division of Chemical Education, Inc. pubs.acs.org/jchemeduc Vol. 88 No. 3 March 2011 10.1021/ed1009927 Published on Web 01/11/2011
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Journal of Chemical Education
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In the Laboratory Table 2. Absorbance Obtained with Sequential Additions of KSCN to Fe(NO3)3 Volume of KSCN/mLa
[FeSCN2þ]/Mb
0.00
0.00
0.10
2.65 10-5
0.20
-5
0.215
106.3
0.354
1.00
0.379
106.0
1.50
0.472
104.9
0.674
111.9
-4
0.844
4.00
2.50
0.599
111.4
1.014
0.50
4.00
0.215
111.5
4.00
0.389
109.1
1.158
1.00
-4
1.296
1.50
4.00
0.494
110.5
-4
1.428
2.00
4.00
0.557
110.6
1.555
2.50
4.00
0.593
110.1
-4
1.62 10 1.81 10 1.99 10
-4
2.17 10
1.00
0.50
4.00
0.563
1.42 10
0.90
4.00
2.00
-4
0.80
0.162
4.00
1.21 10
0.70
Kd
0.521
9.86 10
0.60
Absorbancec
-5
7.57 10
0.50
Volume of KSCN/mLb
4.00
-5
0.40
Volume of Fe(NO3)3/mLa
0.000
5.17 10
0.30
Absorbancec
a
b
The concentration of the KSCN aqueous solution was 0.00108 M. The concentration of the Fe(NO3)3 aqueous solution was 0.200 M Fe(NO3)3 in 1.0 M HNO3 and the volume was 4.00 mL. c Absorbance was measured at 468 nm.
Experiment Materials and Equipment Datalogger, colorimeter, cuvette (∼7.5 mL), small stir bar, 0.10-1.00 mL variable autopipettor, 0.200 M Fe(NO3)3 in 1.0 M HNO3, 0.0020 M Fe(NO3)3 in 1.0 M HNO3, 0.0010 M KSCN in 1.0 M HNO3, 0.0020 M KSCN in 1.0 M HNO3. All solutions are freshly prepared by the instructor. Calibration Plot The colorimeter is calibrated (0% T and 100% T) with distilled water. A small stir bar is placed in a dry cuvette and 4.00 mL of 0.200 M Fe(NO3)3 is added. A 0.10 mL aliquot of 0.0010 M KSCN is added, the solution is stirred, and the absorbance recorded. Additional 0.100 mL aliquots are added until a total of 1.00 mL of 0.0010 M KSCN has been added. Equilibrium Constant The equilibrium constant is determined with sequential additions of 0.50 mL of 0.0020 M KSCN to 4.00 mL of 0.0020 M Fe(NO3)3 in a cuvette. The procedure is repeated with sequential additions of 0.50 mL of 0.0020 M Fe(NO3)3 to 4.00 mL of 0.0020 M KSCN. Hazards FeCl3 is an oxidizer, is harmful if swallowed or inhaled, and causes irritation to skin, eyes, and respiratory tract. HNO3 is corrosive and causes burns to all body tissue. KSCN is harmful if swallowed or inhaled, and causes irritation to skin, eyes, and respiratory tract. Results and Discussion The colorimeter measures the absorbance at four fixed wavelengths: 468, 565, 610, and 660 nm. On the basis of these, λmax is 468 nm. Typical student calibration data obtained by adding 0.10 mL increments of 0.00108 M KSCN(aq) to 4.00 mL of 0.200 M Fe(NO3)3(aq) in 1.0 M HNO3 are shown in Table 2. The plot of absorbance versus concentration of FeSCN2þ gives a slope of (7.26 ( 0.05) 103 L/mol, a y intercept of -0.02 ( 0.07, and R2 = 0.9994. 314
Journal of Chemical Education
Table 3. Typical Student Results for Solution Composition, Absorbance, and Calculated Equilibrium Constants
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Vol. 88 No. 3 March 2011
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a
The concentration of the Fe(NO3)3 aqueous solution was 0.00200 M Fe(NO3)3 in 1.0 M HNO3. b The concentration of the KSCN aqueous solution was 0.00200 M. c Absorbance was measured at 468 nm. d The average value for K is 109.2 ( 2.6.
The equilibrium constants calculated for the various mixtures are shown in Table 3. Students perform calculations in an Excel spreadsheet; a template is provided in the supporting information. The average value reported by students is 113 ( 3 (N = 120) at ∼20 °C. The literature values for the equilibrium constant are in the range 138-146 (6, 7). Our values are lower because activity coefficients are ignored and the less rigorous calibration approach is adopted. Conclusion A well-known and highly practiced lab to determine an equilibrium constant is improved. The improvement is based on the ability of the autopipettor to deliver small and precise volumes of solutions directly into a cuvette. In today's world of calculators and computers, it is not essential to keep the volumes of the calibration solutions and the equilibrium solutions fixed because these are easy to calculate. For those instructors who prefer the graphical approach based on eq 3, the same procedure described here can be used. Literature Cited 1. Stanton, S.; Zhu, L.; Atwood, C. H. Experiments in General Chemistry Featuring MeasureNet; Thompson Brooks/Cole: Belmont, CA, 2006; p 311. 2. Beran, J. A. Laboratory Manual for Principles of General Chemistry, 7th ed.; John Wiley: New York, 2004; p 293. 3. Postma, J. M.; Roberts, J. L.; Hollenberg, J. L. Chemistry in the Laboratory, 5th ed.; Freeman: New York, 2004; p 24-1. 4. Bramwell, F. B.; Dillard, C. R.; Wieder, G. M. Basic Laboratory Principles in General Chemistry with Quantitative Techniques; Kendal/Hunt: Dubuque, IA, 1990; p 233. 5. Ramette, R. W. J. Chem. Educ. 1963, 40, 71–72. 6. Lawrence, G. S. Trans. Faraday Soc. 1956, 52, 236. 7. Cobb, C. L.; Love, G. A. J. Chem. Educ. 1998, 75, 90.
Supporting Information Available Notes for the instructor; handouts for the students; Excel spreadsheet for the calculations. This material is available via the Internet at http://pubs.acs.org.
pubs.acs.org/jchemeduc
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r 2011 American Chemical Society and Division of Chemical Education, Inc.
