Explain any changes that you made, either intentionally or unintentionally, to the experimental procedure that was provided to you.

• Perform chemical reactions that reach equilibrium
• Measure concentrations of reactants and products
• Calculate equilibrium constants

Introduction
When chemical substances react, the reaction typically does not go to completion. Rather, the system reaches a state in which the reactants and the products have concentrations which do not change with time. Such a system is said to be in chemical equilibrium. A ratio of concentrations of products to the concentrations of reactants will be a constant, represented by KC for the reaction.

In this experiment we will study the equilibrium properties of the reaction between iron(III) ion and thiocyanate ion:

Fe3+ (aq) + SCN-– (aq) ⇄ FeSCN2+ (aq) Reaction 1

When solutions containing Fe3+ ion and thiocyanate ion are mixed, they react to some extent, forming FeSCN2+, which has a deep red color. As a result of the reaction, the equilibrium amounts of Fe3+ and SCN- will be less than they would have been if no reaction had occurred; for every mole of FeSCN2+ formed, one mole of Fe3+ and one mole of SCN- will react. The equilibrium constant expression KC for Reaction 1 is:

KC = [FeSCN2+] Equation 1
[Fe3+] [SCN-]

The value of Kc is constant at a given temperature. This means that mixtures containing Fe3+ and SCN– will come to equilibrium with the same value of Kc, no matter what initial amounts of Fe3+ and SCN- were used. Our purpose in this experiment will be to find Kc for this reaction for several mixtures that have been made up in different ways, and to show that Kc indeed has the same value in each of the mixtures.

The mixtures will be prepared by mixing solutions containing known concentrations of iron (III) nitrate, Fe(NO3)3, and potassium thiocyanate, KSCN. The color of the FeSCN2+ ion formed will allow us to determine its equilibrium concentration. Knowing the initial composition of a mixture and the equilibrium concentration of FeSCN2+, we can calculate the equilibrium concentrations of the rest of the pertinent species and then determine KC.

Method for Determining [FeSCN2+]
In this experiment, you will use a spectrophotometer to determine [FeSCN2+] in the equilibrium mixtures. The spectrophotometer measures absorbance, the amount of light absorbed by the solution at a given wavelength. Beer’s law (equation 2) expresses the relationship between absorbance, A, and concentration of a colored species, c.

A =  b c Equation 2

Here,  is a constant that depends on the wavelength of light and on the substance that is absorbing the light, b is the distance that the light travels through the sample of the absorbing substance. The value of b is typically 1 cm. Once the value of  is known, then you can easily convert an absorbance value to a concentration.

The FeSCN2+ complex absorbs blue light. That is why it has a reddish orange color. It absorbs the most light at a wavelength of 447 nm. Therefore, at this wavelength absorbance measurements will have the highest sensitivity to [FeSCN2+]. In this experiment, you will measure the absorbance of all solutions at 447 nm.

In the first part of the experiment, you will determine the relationship between the absorbance and [FeSCN2+] at 447 nm. This will allow you to determine the value of  for FeSCN2+. You will do this by measuring the absorbance of three standard solutions, in which [FeSCN2+] is known. There is a problem here: how can known concentrations of FeSCN2+ be obtained? FeSCN2+ participates in the equilibrium with Fe3+ and SCN- ions. Known amounts of the reactants will not necessarily yield a known amount of the product since the reaction reaches equilibrium before all the reactants are consumed.

This difficulty can be avoided. According to Le Châtelier’s principle, a net reaction from left to right (that is, in the forward direction) occurs when more of a reactant is added. As more and more of the same reactant is added, more and more of the product will be formed. It is possible to add so much of this reactant that essentially all of the other reactant will be converted to product. You will use limiting quantities of SCN- and overwhelming amounts of Fe3+ to produce known amounts of FeSCN2+ in your standard solutions. The amount of FeSCN2+ that is formed will then be essentially equal to the starting amount of the limiting reactant.

Once you have measured the absorbances of the standard solutions, you will plot the absorbances against the concentrations of FeSCN2+ on a graph. The points on the calibration curve will fall on a straight line, which has a slope of b, the constant of proportionality between absorbance and concentration. You can use this calibration curve to find [FeSCN2+] in other solutions.

In the second part of the experiment, you will measure the absorbance of a different set of solutions, in which substantial amounts of both reactants and the product are present. You will use your calibration curve to convert the measured absorbance to the equilibrium concentration of FeSCN2+ in each solution. From the initial concentrations of the reactants and the equilibrium concentration of the product, you will calculate the equilibrium constant for the reaction.

In preparing each of the mixtures in this experiment you will maintain the concentration of H+ ions at 0.5 M. The hydrogen ion does not participate directly in the reaction, but its presence is necessary to avoid the formation of brown-colored species such as FeOH2+, which would interfere with the analysis of [FeSCN2+].

Safety

Always wear your goggles and lab apron or lab coat.

Nitric acid (HNO3) is corrosive. If any solutions containing nitric acid splash onto your hands or clothing, wash them off immediately with copious amounts of running water.

All waste from this experiment should be poured into the appropriate containers.

Experimental Procedure

Part 1: Finding the relationship between absorbance and [FeSCN2+]

Table 1 shows the amounts of the 0.20 M Fe(NO3)3 (in 1 M HNO3) solution, 2.00 × 10-3 M KSCN solution, and DI water to combine. Add these chemicals to your three test tubes. Note that your test tubes are labeled #1, #2 and #3. Be sure to add the solutions to the correct tube. Use the pipets Pipet the approximate amount of each solution into each tube. On your data sheet, record the exact amount of each solution that you add. You will need to use these actual amounts in your calculations.

Mix each solution thoroughly by inverting the test tube several times with a stirring rod. Be sure to wipe the stirring rod with a paper towel in before stirring the next solution. Record your observations.

Zero the spectrophotometer with water as a blank. Measure and record the absorbance of each solution at 447 nm. Rinse out the cuvette before adding the next solution.

Table 1 Composition of solutions for preparing the calibration curve.
Test Tube Vol of 0.20 M Fe(NO3)3 in 1 M HNO3 (mL) Vol of 2.00  10-3 M KSCN (mL) Vol of DI water (mL)
1 5.0 0.50 4.5
2 5.0 1.0 4.0
3 5.0 1.5 3.5

Graph your data using the spreadsheet on Canvas for this experiment. Record the value of  on your data sheet. Include the correct units for this value.

After you are satisfied with your results, dispose of your solutions in the appropriate waste container. Wash and dry your three test tubes since you will use them in the next section.

Part 2: Determining the equilibrium constant
Table .2 shows the amounts of 2.00  10–3 M Fe(NO3)3 (in 1 M HNO3) solution, 2.00  10-3 M KSCN solution, and DI water that you should add to each tube. Add these chemicals to your three test tubes. Note that your test tubes are labeled by number. Be sure to add the solutions to the correct tube. Use the pipets Pipet the approximate amount of each solution into each tube. On your data sheet, record the exact amount of each solution that you add. You will need to use these actual amounts in your calculations.

Zero the spectrophotometer with water as a blank. Measure and record the absorbance of each solution at 447 nm. Rinse out the cuvette before adding the next solution.

Table 2 Composition of solutions for determining the equilibrium constant.
Test Tube Vol of 2.00  10-3 M Fe(NO3)3 in 1 M HNO3 (mL) Vol of 2.00  10-3 M KSCN (mL) Vol of DI water
(mL)
1 5.0 1.0 4.0
2 5.0 2.0 3.0
3 5.0 3.0 2.0
4 5.0 4.0 1.0
5 5.0 5.0 0

After you are satisfied with your results, dispose of your solutions in the appropriate waste container.

Calculations

Part 1. Graphing the relationship between absorbance and [FeSCN2+]
Assuming that all of the SCN- ions have been converted to FeSCN2+ ions, calculate [FeSCN2+] in each of the solutions in Part 1. Remember to account for the dilution that occurs when all three solutions are mixed together.

Part 2: Determining the equilibrium constant
Calculate the value of the equilibrium constant for Reaction 1 from your data for each of the five solutions in Part 2. Remember to account for the dilution that occurs when all three solutions are mixed together.

Data Sheet

Name Section

Lab Partner Name Date

Part 1: Finding the relationship between absorbance and [FeSCN2+]

Complete the table below with the volumes of solutions you prepared.

Test Tube Vol 0.20 M Fe(NO3)3
in 1 M HNO3 (mL) Vol 2.00  10-3 M KSCN (mL) Vol DI water (mL) [FeSCN2+] (M) ABS at 447 nm
1

Value of  =

Part 2: Determining the equilibrium constant

Show your work for calculating one equilibrium constant in the table below.

Fe3+ (aq) + SCN- (aq) ⇄ FeSCN2+ (aq)
Initial concentration (M)

– Change in concentration (M)

= Equilibrium concentration (M)

Record all your calculated equilibrium constants and calculate the average value.

Test Tube 1 2 3 4 5 Average

KC
Report
Each student will write a short report for this two-week experiment. Write everything in your own words. Do not copy or rephrase what somebody else writes. Type the report using double spacing and 12 point font. It should include the sections described below. Lengths for each section are just suggestions.

Procedure (< 1/2 page)
Explain any changes that you made, either intentionally or unintentionally, to the experimental procedure that was provided to you. This includes using different equipment (e.g. 100 mL graduated cylinder instead of 50 mL) or amounts of chemicals. If you followed the procedure without any changes, simply state that you followed the procedure that was given to you.

Data and Results (length varies)
1. Include your data sheet, your Excel graph and data, and calculations in this section.

2. What differences did you see in the solutions that you prepared in Part 1?

3. You have been told that the reaction you are observing is:

Fe3+ (aq) + SCN-– (aq) ⇄ FeSCN2+ (aq)

Another possibility is that the reaction is:

Fe3+ (aq) + 2 SCN-– (aq) ⇄ Fe(SCN)2+ (aq)

Write the equilibrium equation for this new reaction.

Using your data from the mixtures in test tubes 1, 3, and 5 of Part 2, recalculate three new equilibrium constants for this new possible reaction. Hint: Notice that the change in concentration of SCN- in your equilibrium table will change because of the new stoichiometry. Also, the exponent on [SCN-] in the equilibrium constant expression will change for the same reason.

What are the three new equilibrium constants?

Why do the results of your calculations show that the reaction you are observing is more likely to be the first one (one mole of SCN- with one mole of Fe3+)?

4. In Part 1 of this experiment, you made the assumption that the FeSCN2+ concentrations in your standard solutions were equal to the initial concentrations of SCN- ions. Now that you know the value of the equilibrium constant, calculate the actual equilibrium concentrations of Fe3+, SCN-, and FeSCN2+ in mixture in test tube 2. How good was the assumption? Did this satisfy the 5% rule (the loss of reactant is less than 5% of its initial concentration)?

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