Spectrophotometric Determination of The Equilibrium Constant

Categories: Chemistry


To find the equilibrium constant for the reaction between iron (III) and thiocyanate, absorbances of different equilibrium systems of FeSCN2+(aq) were measured by a spectrophotometer set at wavelength 470 nm. These absorbances were used to calculate the concentration of FeSCN2+(aq) at equilibrium, using Beer's Law, after which an ICE table was used to calculate the equilibrium concentration of reactants and products and to obtain the equilibrium constant. The equilibrium constant was found to be 124 at 22.7". Inaccuracy of results may occur from the inability to maintain the solutions at a constant temperature and leftover deionized water droplets from rinsing causing dilution errors.


The purpose of this experiment is to determine the equilibrium constant, Kc, for the reaction between iron (III) and thiocyanate ions (Fe3+(aq) + SCN-(aq)) by finding the absorbances of different concentrations of FeSCN2+ using a spectrophotometer.


The equilibrium constant is an expression that originates from the combination of the concentration of substances in a solution that react and reach their equilibrium point.

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The magnitude of the equilibrium constant acts as an estimate of the extent to which a reaction will proceed and the concentration of products and reactants when the reaction reaches equilibrium. This experiment utilizes Beer's Law and spectrophotometry to determine the equilibrium constant. Beer's Law states that the absorbance of a solution is proportional to its concentration as indicated by the equation:

Absorbance = e L c.

This suggests that the ratio of the absorbances is proportional to the ratio of concentrations of solutions at the same wavelength.

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Therefore, the equation for Beer's Law can be used to derive an alternative equation:

Concentration1 / Concentration2 = Absorbance1 / Absorbance2.

This equation is used to determine the concentration of FeSCN2+ at equilibrium. These absorbance values are obtained using a spectrophotometer, which indirectly measures the concentration of the orange-red colored iron (III) thiocyanate complex by measuring its absorbance in each solution. Concentrations of reactants at equilibrium can be calculated by subtracting the equilibrium concentration of the product from the initial concentration of the reactants using an Initial-Change-Equilibrium (ICE) table.

Materials and Methods

Materials: burette, distilled water, test tubes, graduated cylinder, 0.002 M Fe(NO3)3, 0.002 M KSCN, 0.200 M Fe (NO3)3, spectrophotometer, plastic cuvette, retort stand and clamp, thermometer, beakers


  1. Set up four burettes with 50 mL of solutions: 0.002 M Fe(NO3)3, 0.002 M KSCN, 0.200 M Fe(NO3)3, and distilled water.
  2. Label four clean, dry test tubes 1-4.
  3. Collect solutions and distilled water as specified in Table 1.
  4. Tap each test tube to mix the solutions thoroughly.
  5. Measure and record the temperature of the solution in Test tube 1.
  6. Prepare a standard solution by mixing 9.00 mL of 0.200 M Fe(NO3)3 and 1.00 mL of 0.002 M KSCN in Test tube 5. Stir thoroughly.
  7. Calibrate the spectrophotometer at a wavelength of 470 nm.
  8. Prepare a blank by filling a cuvette ѕ full with distilled water, and place it in the spectrophotometer. Set the reference to 0.
  9. Rinse the cuvette twice with the solution from Test tube 1, then fill it ѕ full with the solution. Measure and record the absorbance.
  10. Repeat the absorbance measurements for test tubes 2, 3, 4, and 5.



Upon mixing, the solutions in all test tubes changed from colorless to orange-red.


Absorbance Trial 1 Trial 2 Trial 3 Trial 4
Trial 5 (Standard) 0.911
Temperature 22.7°C


Kc expression: Kc = [FeSCN2+]/([Fe3+] [SCN-])

[Fe3+] (M) [SCN-] (M) [FeSCN2+] (M)
Initial 1.00 * 10^-3 4.00 * 10^-4 0.00
Change -4.15 * 10^-5 -4.15 * 10^-5 +4.15 * 10^-5
Equilibrium 9.59 * 10^-4 3.59 * 10^-4 4.15 * 10^-5


The equilibrium constant Kc is a measure of the extent to which a reaction will proceed and the concentration of products and reactants at equilibrium. In this experiment, the equilibrium constant was determined to be 124 at 22.7°C. The reaction involves the formation of a complex ion, FeSCN2+, which contributes to the orange-red color of the solution.

The concentrations of FeSCN2+ complex ions at equilibrium were proportional to the absorbance values, indicating that as concentrations increased, so did the absorbances. However, despite the concentration increase, the equilibrium constant remained within the range of 121 to 126. This suggests that changing the equilibrium concentration at a constant temperature does not significantly affect the equilibrium constant.

Possible sources of error in the experiment include the inability to maintain a constant temperature throughout all trials, which could influence the equilibrium constant since it is temperature-dependent. Additionally, leftover deionized water droplets from cuvette rinsing may have caused dilution errors, affecting absorbance values and the calculated Kc.


In conclusion, the equilibrium constant for the reaction between iron (III) and thiocyanate was determined to be 124 at 22.7°C. The experiment utilized spectrophotometry and Beer's Law to calculate equilibrium concentrations and obtain the equilibrium constant. Possible errors in the experiment could have affected the accuracy of the results, including temperature variations and dilution errors caused by leftover water droplets during cuvette rinsing.


Based on the findings and limitations of this experiment, the following recommendations are proposed:

  1. Ensure better temperature control during the experiment to minimize the impact of temperature variations on the equilibrium constant.
  2. Take extra care when rinsing cuvettes to avoid dilution errors caused by leftover water droplets.

Updated: Dec 29, 2023
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Spectrophotometric Determination of The Equilibrium Constant. (2019, Aug 20). Retrieved from https://studymoose.com/document/chem-lab-reporttt-1

Spectrophotometric Determination of The Equilibrium Constant essay
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