Determining Equilibrium Constant for Thiocyanatoiron
Objective
Determine the equilibrium constant Kc for the chemical reaction:
Fe3+(aq) + SCN–(aq) ⇌ FeSCN2+(aq)
Procedure
| 0.0020M Fe(NO3)3 (ml) | 0.0020M KSCN (ml) | H2O (ml) |
| 5 | 2 | 3 |
| 5 | 3 | 2 |
| 5 | 4 | 1 |
| 5 | 5 | 0 |
| 18 (0.2M) | 2 | 0 |
(1) Calibrate the colorimeter.
(2) Prepare five test tubes with the designated solutions:
(3) Measure each solution using the colorimeter and computer software.
(4) Record an observation of the solution temperature.
(5) Calculate the average equilibrium constant for the reaction being studied, using the 5th test as the standard solution.
Result Sheet
| Trial: | 1 | 2 | 3 | 4 | 5 (standard) |
| Absorption | 0.032 | 0.027 | 0.043 | 0.058 | 0.107 |
| [Fe3+]i | 0.001 | 0.001 | 0.001 | 0.001 | 0.0018 |
| [SCN–]i | 0.0004 | 0.0006 | 0.0008 | 0.001 | 0.0002 |
| [FeSCN2+]eq | 5.98×10-5 | 5.05×10-5 | 8.04×10-5 | 1.08×10-4 | 0.0002 |
| [Fe3+]eq | 9.4×10-4 | 9.4×10-4 | 9.4×10-4 | 9.4×10-4 | |
| [SCN–]eq | 3.40×10-4 | 5.49×10-4 | 7.20×10-4 | 8.92×10-4 | |
| Kc (M-1) | 187.0 | 97.86 | 118.79 | 128.81 | |
| Average Kc | 133.1 | ||||
| Average Temperature: 23.5 C | |||||
Sample Calculations
1. [Fe3+]i and [SCN–]i @ Trial 1
[Fe3+]i= 0.0020M Fe(NO3)3 x (5ml/10ml) = 0.001M, [SCN–]i = 0.0020M KSCN x (2ml/10ml) = 0.0004M
2. [FeSCN2+]std
[FeSCN2+]std = 0.0020M x (2ml/20ml) = 0.0002M
3. [FeSCN2+]eq @ Trial 1
[FeSCN2+]eq= (Aeq/Astd) x [FeSCN2+]std = [(0.032/0.107) x 0.0002M=5.98×10-5M
4. [Fe3+]eq and [SCN–]eq @ Trial 1
[Fe3+]eq= 0.001 – (5.98×10-5) = 9.4 x 10-4M , [SCN–]eq = 0.0004 – (5.98×10-5) = 3.4 X 10-4M
5. KC @ Trial 1 (23.5 degrees C)
Kc= [FeSCN2+]eq / [Fe3+]eq [SCN–]eq = (5.98×10-5)/(9.4 x 10-4)( 3.4 X 10-4) = 187.0
6. Average Kc
Kc(avg) = ∑(Kc(trials 1->n))/ (n) = (187+97.86+118.79+128.81) / 4 = 133.1
Analysis
In this experiment, the chemical reaction Fe3+(aq) + SCN–(aq) ⇌ FeSCN2+(aq) was studied to determine the equilibrium constant, Kc. To determine this value, the absorptivity of several solutions were recorded using a colorimeter. Inclusion of a standard solution allowed for equilibrium calculations of the reactant and product concentrations. Experimentally, the average constant Kc was calculated at (133.1). Linear analysis of the Kc range has revealed flaws in the first trial. The source of this error is probably due to an error in calibrating the colorimeter or poor handling of the solution.
Editor’s note (October 29, 2026): The student’s recorded values are left as written. Three points need clarifying.
- [Fe3+]eq changes from trial to trial. In the ICE table each trial uses up x = [FeSCN2+]eq of both reactants, so [Fe3+]eq = 0.001 M − x and [SCN–]eq = [SCN–]i − x (LibreTexts). The table reuses trial 1’s 9.4×10-4 M for trials 2–4. Recomputed from the recorded absorbances (x = A/0.107 × 0.0002 M, unrounded): trial 2, [Fe3+]eq = 9.50×10-4 M and Kc = 96.7; trial 3, 9.20×10-4 M and 121.4; trial 4, 8.92×10-4 M and 136.4. Trial 1 (187.0) is unchanged, so the average is 135.4 M-1, not 133.1 (the mean of the recorded values).
- Standard tube. 18 mL of 0.2 M Fe(NO3)3 in 20 mL total gives [Fe3+]i = 0.18 M, not the 0.0018 M shown in the table. That large excess of Fe3+ drives essentially all the SCN– into FeSCN2+, which is why [FeSCN2+]std can be taken as the 0.0002 M of SCN– added (LibreTexts).
- Symbols. The equilibrium arrows, which displayed as garbled characters (“ß à”), are restored to ⇌, and “1O” (letter O) in the sample calculations now reads 10.
Related absorbance-based labs: Equilibrium and Determination of the stability constant for an iron (iii) sulfosalicylate complex.